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Title: Chlorination of Water
Author: Race, Joseph
Language: English
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Copyright Status: Not copyrighted in the United States. If you live elsewhere check the laws of your country before downloading this ebook. See comments about copyright issues at end of book.

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  |                        Transcriber's notes:                         |
  |                                                                     |
  | * Text in italics in the original document is presented here        |
  |   between underscores, as in _text_.                                |
  | * Bold-face text in the original document is presented here between |
  |   equal signs, as in =text=.                                        |
  | * Old-English font is presented here by enclosing the text between  |
  |   ~, as in ~text~.                                                  |
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  |   respectively.                                                     |
  | * The oe-ligature is presented as [oe] as in diarrh[oe]a.           |
  | * Small-caps text in the original document is presented as all-caps |
  |   here.                                                             |
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  |   <=>.                                                              |
  | * Footnotes (indicated by [A], [B], etc.) have been moved to        |
  |   directly below the paragraph or table they refer to, references   |
  |   (indicated by [1], [2], etc.) are collected at the end of each    |
  |   chapter.                                                          |
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     =Examination of Milk for Public Health Purposes.=

     A practical handbook for those engaged in the chemical and
     bacteriological examination of milk for public health purposes. vi
     + 224 pages, 5-1/4 × 8, 4 diagrams. Cloth, $1.75 net.

     =Chlorination of Water.=

     In this book the various aspects and methods of chlorination are
     discussed with a view to stimulating research work in this field of
     science. viii + 158 pages, 5-1/4 × 8, 12 figures and 16 diagrams.
     Cloth, $1.50 net.

Chlorination of Water



_City Bacteriologist and Chemist, Ottawa; Capt. Canadian Army
Hydrological Corps; Associate Member of Committee on Water Supplies,
American Public Health Association; Member of Committee on Water
Standards and Standard Methods of Analysis, American Water Works
Association; Chairman of Committee on Standard Methods of Analysis,
Canadian Public Health Association_






Copyright, 1918









~Sir Alexander Houston, K.B.E., D.Sc., M.B., C.M.~


No apology is necessary for the publication of a book on the
chlorination of water. This method of treatment, practically unknown
fifteen years ago, has advanced in popularity during the last decade in
a most remarkable manner, and in 1918 over forty millions of people are
being supplied with chlorinated water.

It may justifiably be said that no other sanitary measure has
accomplished so much at so small a cost; and that civilization owes a
deep debt of gratitude to the pioneers in municipal water chlorination:
Dr. A. C. Houston in England, and Mr. G. A. Johnson and Dr. Leal in

In this volume I have endeavoured to collect and correlate the
information hitherto scattered in various journals and treatises and to
present it in a comprehensible manner. The various aspects and methods
of chlorination are discussed and suggestions have been made which, I
hope, will stimulate research work in this fertile field of science.

I wish to acknowledge my indebtedness to the engineering staff of the
Ottawa Water Works Department and to Lieut. W. M. Bryce for the
preparation of diagrams.



April, 1918.


  CHAPTER                                                           PAGE

  I. HISTORICAL                                                        1

  Sodium Chloride. Chlorine. Bleach. Eau de Javelle. Antiseptics.
  Hermite fluid. Webster's process. Electrozone. Chlorination of
  sewage in Germany, U. S. A., and England. Chlorination of
  water. Lincoln installation. Oxychloride. German experiments.
  European practice. Inception of chlorination in America.

  II. MODUS OPERANDI                                                  14

  Composition of bleach. Bleaching action. Nascent oxygen
  hypothesis. Hydrolysis of bleach. Effect of acids and salts on
  hydrolysis and germicidal action. Effect of ammonia. Direct
  toxic action. Hypochlorous acid. Sodium hypochlorite. Chlorine
  water. Nature of action.

  III. DOSAGE                                                         30

  Organic matter. Initial count. Viability of organisms. Mineral
  matter. Colour. Temperature. Admixture. Contact period.
  Turbidity. Light. Determination of dosage.

  IV. BACTERIA SURVIVING CHLORINATION                                 50

  Disinfectants. Antiseptics. Viability of bacteria. New York
  results. Reversed ratio of counts. Coliform organisms.
  Aftergrowths in water and sand.

  V. COMPLAINTS                                                       62

  Auto-suggestion. Tastes and odours. Sludge problem. Colic.
  Effect on fish and birds. Effect on plants and flowers.
  Corrosion of iron and lead pipes.

  VI. BLEACH TREATMENT                                                72

  Storage of bleach. Mixing tanks. Storage tanks. Dosing
  apparatus. Control. Analysis of liquor. Detection and
  estimation of free chlorine. Chlorometer. Cost of construction
  and operation. Antichlors. DeChlor filters.

  VII. LIQUID CHLORINE                                                89

  Historical. Leavitt-Jackson machine. Electro Bleaching Gas
  Co.'s types. Wallace and Tiernan's manual control types.
  Effect of temperature on gas pressure. Impurities in gas.
  Advantages. Comparison of liquid chlorine and bleach. Cost of
  treatment. Popularity. Chlorine water.


  Hermite fluid. Eau de Javelle. Chloros. Non-diaphragm cells:
  Dayton, Hermite, Mather and Platt, Haas and Oettel. Diaphragm
  cells: Hargreaves-Bird, Nelson, Allen-Moore. Montreal
  installation. Costs.

  IX. CHLORAMINE                                                     115

  Preparation. Absorption by water. Experimental results. Works
  results. Ratio of chlorine to ammonia. Economics. Advantages.
  Operation. Other chloramines. Halazone.

  X. RESULTS OBTAINED                                                132

  Object of chlorination. Effect on filter rates and algæ.
  Hygienic results. Typhoid rates. Typhoid reduction at
  Philadelphia, Chicago, and Ottawa. Abortive epidemics. Use and
  abuse of chlorination.

  APPENDIX                                                           147

  Estimation of free chlorine in water.

  NAME INDEX                                                         151

  SUBJECT INDEX                                                      155




Chlorine, although one of the most widely distributed elements known to
chemists, is never found in the free condition in nature; it exists in
enormous quantities in combination with sodium, potassium, calcium,
magnesium, etc. As sodium chloride, common salt, it occurs in
practically inexhaustible quantities in sea water together with smaller
quantities of other chlorides. In mineral form, enormous deposits of
sodium chloride are found in Galicia, Transylvania, Spain, in England
(particularly in Cheshire), and in sections of North America. The most
important deposits of potassium chloride are those at Stassfurt,
Germany, where it occurs either in the crystalline condition as sylvine
or combined with magnesium chloride as carnallite.

Chlorine was discovered by the Swedish chemist Scheele in 1774, but he,
like Lavoisier and his pupil Berthollet, who declared it an oxygenated
muriatic acid, was unaware of the elemental nature of the new substance.
Sir Humphrey Davy investigated this body in 1810 and definitely proved
it to be an element; Davy designated the element chlorine from the Greek
[Greek: chlôros] = green.

The first attempt to utilise chlorine, or its compounds, for bleaching
purposes, appears to have been due to James Watt, who noticed the
decolourising properties of chlorine during a visit to Berthollet. This
attempt ended in failure because of the destructive effect on the
fibres, but, in later trials, this was prevented by first absorbing the
gas in a solution of fixed alkali. These experiments proved the
possibility of bleaching by means of chlorine compounds but the high
cost of soda made the process unprofitable, and it was not until Henry
succeeded in preparing a combination with lime that could be reduced to
a dry powder that this mode of chemical bleaching became a commercial

The manufacture of chloride of lime (hypochlorite of lime, bleaching
powder, bleach) was taken up by Charles Tennant in 1799 at St. Rollox
near Glasgow, and in 1800 about 50 tons were sold at a price of $680
(£139) per ton.

Chlorine is produced as a by-product in the manufacture of soda by the
Leblanc process, but until 1865, when the British Alkali Act stopped the
discharge of hydrochloric acid vapours into the atmosphere, the
development of the bleaching powder industry was not rapid. The
hydrochloric acid that was formerly discharged into the air as a waste
product afterwards became a valuable asset that enabled the Leblanc
process to successfully compete with the newer ammonia-soda process. In
1890 another competitor to the Leblanc process was introduced when
caustic and chlorine were produced in Germany by electrolytic methods.
After the successful development of this method in Germany, it was taken
up in the United States of America and in 1912 more than 30,000
electrical horse-power were daily used in this industry. In 1914 the
almost complete cessation of exports of bleach from Europe raised the
price, which attained phenomenal heights in 1916 (cf. page 125), and
stimulated the production of bleach both in the U. S. A. and Canada.


      Year.    |    Bleach Manufactured,    |    Selling Price
               |        Short Tons.         |     Per 100 Lbs.
      1904     |          19,000            |
      1909     |          58,000            |
      1914     |         155,000            |        $1.63
      1915     |         180,000[A]         |         2.63
      1916     |         230,000[A]         |         6.56
      1917     |         260,000[A]         |         2.44
  [A] Estimated.

As a disinfectant, chlorine was first used about the year 1800 by de
Morveau, in France, and by Cruikshank, in England, who prepared the gas
by heating a mixture of hydrochloric acid and potassium bichromate or
pyrolusite; this is essentially the same as the original mixture used by

During the early part of the last century the efficacy of chlorine of
lime as a disinfectant, and particularly as a deodourant, was well
recognised and as early as 1854 an English Royal Commission used this
substance for deodourising the sewage of London. A committee of the
American Public Health Association reported in 1885 that chloride of
lime was the best disinfectant available when cost and efficiency were

Eau de Javelle, first made by Percy at the Javelle works near Paris in
1792, is another chlorine compound that has enjoyed a considerable
reputation as a disinfectant and deodouriser for over a century; it is
essentially a mixture of sodium chloride and sodium hypochlorite.

The discovery of electrolytic hypochlorites dates back to 1859, when
Watt found that chlorides of the fixed alkalies and alkaline earths
yielded hypochlorites on being submitted to the action of an electrical

Until the middle of the last century disinfection was regarded as a
process that arrested or prevented putrefactive changes but the nature
of these changes was imperfectly comprehended and micro-organisms were
not associated with them.

In 1839 Theodor Schwann,[1] who might be regarded as the founder of the
school of antiseptics, reported that "Fermentation is arrested by any
influence capable of killing fungi, especially by heat, potassium
arseniate, etc...."; but his results were not accepted by the adherents
of the theory of spontaneous generation and it was not until the
publication of the work of Schroder and Dusch[2] that Schwann's views
were even partially accepted. The final refutation to the spontaneous
generation theory was given by the monumental researches of Pasteur who,
in 1862, proved the possibility of preparing sterile culture media and
demonstrated the manner in which they could be protected from
contamination. Bacteria and other micro-organisms were shown to be
responsible for the phenomena that had been attributed previously to the
"oxygen of the air," and from this period the development of
bacteriology as a science proceeded rapidly.

The next important step, from the public health standpoint, was the
discovery by Koch, in 1876, that a specific bacterium (_B. anthracis_)
was the cause of a specific disease in cattle (anthrax or splenic
fever). In 1882 Koch made a further advance by developing a solid
culture medium which permitted disinfectants and antiseptics to be
studied quantitatively with a greater degree of accuracy than had been
possible previously.

Since 1845, when Semmelweiss succeeded in stamping out puerperal fever
in Vienna, where it had been so long established as to be endemic,
chlorine has been very generally employed in sanitary work and the
conditions necessary for obtaining successful results have been
partially elucidated. Baxter was the first to state that the
disinfecting action depended more upon the nature of the pabulum than
upon the specific organism present and this was confirmed later by Kuhn,
Bucholtz, and Haberkorn. The latter found that urine consumed large
quantities of chlorine before any disinfection occurred.

One of the earliest preparations used in sanitary work was an
electrolysed sea water, usually known as Hermite Fluid. This was
introduced by M. Hermite in 1889 and was employed for domestic purposes
and for flushing sewers and latrines. It was used at Brest for the
dissolution of fæcal matter and a prolonged trial was given to it at
Worthing in 1894. The report of Dupré and Klein, who conducted the
bacteriological examinations, was against the process, but Ruffer and
Roscoe reported more favourably and further trials were carried out at
Havre, l'Orient, and Nice. The _Lancet_ (May 26, 1894) reported at
length upon the Worthing experiments: it was found that during the
electrolysis of the sea water, the magnesium chloride was also partially
converted into hypochlorite, which then dissociated into magnesium
hydrate and hypochlorous acid; the former deposited in the electrolyser
and left the solution acid and unstable; urine was found to act upon it
at once with a consequent loss in strength of over 50 per cent.

Another electrolytic method was that of Webster,[3] who installed an
experimental plant at Crossness, near London, in 1889. A low-tension
direct current was passed between iron electrodes placed in the sewage
and although the process was largely one of chemical precipitation,
Webster noted the disinfecting value of the hypochlorite formed from the
chlorides normally present in the sewage. He also directed the attention
of sanitarians to the possibility of using sea water as a cheap source
of chlorides and a plant based on this principle was erected in Bradford
in 1890 and reported upon by McLintock.[4]

Strong salt solutions were substituted for sea water by Woolf and the
product was commercially known as "Electrozone." A plant of this
description was installed at Brewster, N. Y., in 1893[5] for
chlorinating the sewage from a small group of houses. The sewage was
discharged into a small creek which polluted Croton Lake. Successful
results led to a similar treatment near Tonetta Creek.[6] This was
apparently the first occasion on which the specific object was the
destruction of bacteria.

Electrozone was used at Maidenhead, on the Thames, in 1897 and the
installation was reported upon by Robinson, Kanthack, and Rideal in
1898. Kanthack found that a dosage 3-3.6 p.p.m. reduced the organisms in
a sewage effluent to 10-50 per c.cm. whilst Rideal found that about 18
p.p.m. of chlorine produced a condition of sterility in 1 c.cm.

Chloride of lime had previously been used in the London sewage as a
deodourant by Dibden in 1884 but the treatment was not successful and
was abandoned in favour of other oxidisers.

During the last decade of the twentieth century the use of bleach for
the disinfection of both sewage and water received the attention of many
well-known German sanitarians and many important results were obtained.

In the earlier experiments made at Hamburg, Proskauer and Elsner[7]
obtained satisfactory results with 3-4 p.p.m. of chlorine on a clarified
sewage with 10 minutes contact. Dunbar and Zirn (_ibid._) used crude
sewage and found that 17 p.p.m. of available chlorine were required to
remove _B. typhosus_ and cholera vibria with a contact period of two
hours. A striking feature of all the German work on chlorination is the
very high degree of purification aimed at: quantities as large as one
litre were tested for specific organisms and in many of the experiments
with sewage _B. coli_ was found to be absent from a considerable
percentage of the samples.

The importance of previously removing suspended matter, which could not
be penetrated by the germicide, was emphasised by Schwartz[8] although
it had been previously noted by Schumacher.

At the Royal Testing Station in Berlin, numerous experiments on sewage
chlorination were made by Kranejuhl and Kurpjuivut.[9] The results were
judged by the _B. coli_ content, which was taken as an index of
pathogenicity because this typical intestinal bacillus was found to be
more frequent and less viable than the majority of the pathogenic

Other important work on this subject was carried out, in connection with
the pollution of the Hooghly River, by a Bengal Government Commission in
1904; and by the State Board of Health of Ohio in co-operation with the
Bureau of Plant Industry of the United States Department of Agriculture
in 1907. The chlorination experiments of the latter were reported by
Kellerman, Pratt, and Kimberly.[10]

The most valuable contribution to the disinfection of sewage was that of
Phelps,[11] who critically examined the work of previous experimenters
and directed attention to the unnecessary stringent standards adopted in
European practice. His work at Boston in 1906, at Red Bank, N. J., and
at Baltimore in 1907, demonstrated in an indubitable manner the economic
possibilities of sewage chlorination. The dosages necessary for crude
sewage and filter effluents were indicated and also the necessary
contact periods. This work marks the commencement of a new era in
sanitary science.

The first occasion on which chlorine compounds were first used for the
disinfection of water cannot be definitely ascertained. It has been
stated to the author that bleach was used for treating wells as early as
1850 but this treatment was apparently made without definite knowledge
of the destruction of micro-organisms.

In 1897, Sims Woodhead employed bleach solutions for the sterilisation
of the distribution mains at Maidstone, Kent, subsequent to an epidemic
of typhoid fever.

The credit for the first systematic use of chlorine in water
disinfection is due to A. C. Houston with whom McGowan was associated in
the work carried out at Lincoln in 1904-1905.[12] The reservoirs,
filters, and distribution system, owing to flood conditions, became
infected with typhoid bacilli which caused a severe epidemic amongst the
consumers. The storage and purifications works were thoroughly treated
with a solution of "chloros" (sodium hypochlorite containing
approximately 10 per cent of available chlorine) which was regulated to
give an approximate dosage of 1 part per million. The bacteriological
results were entirely satisfactory but many complaints were received
that the treatment had imparted a mawkish taste to the water. This was
attributed to the action of the alkaline chloros on the organic
impurities in the water. It was also stated that the water injured
plants, fish, and birds and extracted abnormal amounts of tannin from
tea but no substantiating evidence was produced in support of these
complaints. Houston made a continuous physiological test of the effect
of the chlorinated water on small fish by suspending a cage of gold fish
in the filter effluent chamber and also proved that the treatment had no
appreciable effect on the plumbo-solvency of the supply.

Nesfield, of the Indian Army Medical Service,[13] reported in 1903 the
results of numerous experiments on the destruction of pathogenic
organisms by chlorine compounds and suggested their use in military work
to prevent a recurrence of the appalling loss of life from water-borne
diseases (especially enteric fever) such as took place during the Boer
War. Nesfield proposed to use about 100 p.p.m. of available chlorine and
to remove the excess after a contact period of 10 minutes. This work is
especially interesting from the historical standpoint because it
contains the first suggestion of the possibilities of compressed
chlorine gas in steel cylinders.

A few years later, electrolytic hypochlorite (oxychloride) was used at
Guildford by Rideal and various chlorine compounds were tried on the
water of the Seine and Vanne, in France, and at Middlekerke and Ostend,
in Belgium. Experimental work on water chlorination was also reported by
Thresh and by Moor and Hewlett.[14]

During the nineties many experiments on water chlorination were made by
Traube, Sickenberger, Kauffman, Berge, Bassenge, and others. Traube[15]
was able to completely sterilise water rich in bacteria in 2 hours by
the addition of bleach equal to 1.06 p.p.m. of available chlorine. At
the end of the contact period about 90 per cent of the added chlorine
was unabsorbed and was destroyed by the addition of sodium bisulphite.
Bassenge[16] followed up the work of Traube and that of Sickenberger and
Kauffman, who had shown that it was possible to destroy cholera vibrio
in Nile water by means of sodium hypochlorite. Bassenge used higher
concentrations than Traube and found it possible to destroy _B.
typhosus_ and _B. coli_ in ten minutes with 60-90 p.p.m. of available
chlorine. The excess was destroyed by adding calcium bisulphite.
Lode[17] experimented with waters seeded with _B. coli_, _B. typhosus_,
and _B. tetani_ and found, contrary to Traube, that 1-2 p.p.m. of
chlorine did not sterilise in two hours. _B. coli_ was usually destroyed
by 4 p.p.m. of chlorine in ten minutes and even better results were
obtained with _B. typhosus_ and cholera vibrio: the former was destroyed
in one hour by 1 p.p.m. and in ten minutes by 2 p.p.m.; the latter
organism required 1-2 p.p.m. with a twenty-minute contact period. Lode
noted that organic matter lowered the bactericidal activity of chlorine
and recommended the use of 30 p.p.m. of chlorine to ensure rapid and
complete sterilisation. Berge[18] used chlorine peroxide, generated by
the action of hydrochloric acid on potassium chlorate, for the
sterilisation of water and this process was afterwards used at Ostend at
a plant having a capacity of about 1,300,000 gallons per day. The dosage
was equal to 0.53 p.p.m. of available chlorine and coke filters were
used to destroy the excess although they were not found to be
indispensable as the free chlorine disappeared spontaneously. This
process appears to have been tried on the Brussels supply and also for
the treatment of a hospital supply at Petrograd.

The object of German sanitarians seems to have been to obtain
practically instantaneous sterilisation of water for the use of
travellers and troops in the field. Until the commencement of the
European War they did not have a high opinion of chlorination and
generally regarded it as inefficient. Schumberg[19] expressed the
opinion that no chemical method of disinfection could be absolutely
relied upon, under all circumstances, to prove fatal to bacteria.
Plucker[20] stated that several investigators, particularly Schuder, had
shown that chlorine, even in the proportion of 40 p.p.m. did not
invariably destroy cholera vibrio and _B. typhosus_; and that with
smaller doses the destruction was still less complete. He also stated
that the bacteriological experiments of American workers were open to
criticism and that they employed antiquated methods.

By 1916 the German sanitarians appeared to have realised that their
bacteriological standards were too stringent (Langer[21]) and that the
process had proved its value in an indisputable manner.

European practice, in the comparatively few instances in which it has
been used, has been to employ large doses of chlorine and to remove the
excess by chemicals or by filtration through special media. In 1916,
however, London commenced to chlorinate a portion of its supply and the
following year practically the whole supply was chlorinated. A dosage of
approximately 0.5 p.p.m. is used and the bleach solution is added to
the pre-filtered water. Worcester is also proposing to chlorinate the
supply to maintain the purity of the water without extending the slow
sand filtration plant.

In North America, hypochlorite of soda and chlorine were used on the
Jewell Filter at the Louisville Experimental Station in about 1896 by
George W. Fuller and a year later they were used at Adrian by Jewell.
The first commercial successful attempt was made by G. A. Johnson. In
1908 the Union Stock Yards Company of Chicago were proceeded against by
the City of Chicago regarding the condition of the effluent of the
Bubbly Creek filter plant. Copper sulphate had been previously used in
conjunction with the filters but stock shippers complained that the
water had a deleterious effect upon the animals consuming it. Johnson
eliminated the copper treatment and substituted bleach which was added
seven and a half hours previous to filtration, with a dosage of 1.5
p.p.m. The results were very satisfactory.

About the same time, Johnson and Leal commenced the treatment of the
Boonton supply of Jersey City, N. J., consumed about 40 million gallons
per day. The water was first treated with 36 pounds of bleach per
million gallons (1.4 p.p.m. of available chlorine) but this quantity was
gradually reduced until only 5 pounds per million gallons (0.2 p.p.m. of
chlorine) were being used in April, 1909. The ability of the process to
adequately purify water became the cause of a lawsuit and the decision
of the Court was:

"From the proofs taken before me, of the constant observation of the
effect of this device, I am of the opinion and find that it is an
effective process which destroys in the water the germs, the presence of
which is deemed to indicate danger, including the pathogenic germs, so
that the water after this treatment attains a purity much beyond that
attained in water supplies of other municipalities. The reduction and
practical elimination of such germs from the water was shown to be
substantially continuous.

"Upon the proofs before me, I find that the solution described leaves no
deleterious substances in the water. It does produce a slight increase
in the hardness but the increase is so slight as in my judgment to be

"I do therefore find and report that this device is capable of rendering
the water delivered in Jersey City pure and wholesome, for the purposes
for which it is intended and is effective in removing from the water
those dangerous germs which were deemed by the decree to possibly exist
therein at certain times."[22]

During the next few years the use of hypochlorite in water purification,
both alone and in conjunction with filtration, became very popular and
in 1911 over 800 million gallons per day were treated in this manner.
Amongst the users were some of the largest cities in North America,
including Brooklyn, Albany, and New York City, N. Y., Cincinnati and
Columbus, Ohio, Harrisburg, Philadelphia, Pittsburg, and Erie, Pa.,
Hartford, Conn., Nashville, Tenn., St. Louis and Kansas City, Mo.,
Montreal, P. Q., Toronto and Ottawa, Ont., Baltimore, Md., and
Minneapolis, Minn. At present (1918) over 3,000 million gallons per day
are being chlorinated in North America and more than 1,000 cities and
towns are employing this process.


[1] Schwann. Microskopische Untersuchungen über die Übereinstimmung in
der Textur und dem Wachstum der Tiere und Pflanzen. Berlin. 1839

[2] Schroder and Dusch. Ann. der Chem. u. Pharm., 1854, 89, 232.

[3] Webster. The Engineer. 1889, 67, 261.

[4] McLintock. Brit. Med. Jour., 1890, 11, 498.

[5] Eng. News. 1893, 30, 41.

[6] Eng. Record. 1894, 29, 110.

[7] Proskauer and Elsner. Vierteljahresschr. ger. Med. u. öff.
Sanitätswesen. 1898, 16, Supp. Heft.

[8] Schwartz. Gas. Eng., 1906, 29, 773.

[9] Kranejuhl and Kurjuivut. Mitteilungen aus der Königlichen
Prüfungsanstalt für Wasserversorgung und Abwässerbeseitigung zu Berlin,
1907, 9, 149.

[10] Kellerman, Pratt, and Kimberly. Bull. 115, Bur. Plant Ind., U. S.
Dept. of Agr., 1907.

[11] Phelps. Water Supply Paper 229, Dept. of Int., U. S. Geo. Survey.

[12] Houston and McGowan. 5th Rpt. Royal Commission on Sewage Disposal.

[13] Nesfield. Public Health. 1903, 15, 601.

[14] Moor and Hewlett. Rpt. of M. O. to L. G. B., 1909-10.

[15] Traube. Zeit. f. Hyg., 1894, 16, 149.

[16] Bassenge. Zeit. f. Hyg., 1895, 20, 227.

[17] Lode. Archiv. f. Hyg., 1895, 24, 236.

[18] Berge. Rev. d'Hyg., 1900, 22, 905.

[19] Schumburg. Zeit. f. Hyg., 1903, 45, 125.

[20] Plucker. J. Gasbeleucht., 1911, 54, 385.

[21] Langer. Zeit. f. Hyg., 1916, 81, 296.

[22] Johnson. Jour. Amer. Pub. Health Assoc., 1911, 1, 566.



Before considering the "modus operandi" of chlorine and hypochlorites,
it will be advisable to take up the composition of the latter substances
and particularly that of "bleach." Bleach is manufactured by passing
chlorine gas over slaked lime and the ensuing reactions are often
represented by the equation Ca(OH)_{2} + Cl_{2} = CaOCl_{2} + H_{2}O.
This represents the substance formed as a pure oxychloride of calcium
which contains approximately 50 per cent of chlorine, but the article
commercially produced never contains this amount of chlorine, the usual
percentage being from 35-37. The general composition of commercial
bleach is fairly uniform. This is shown in the following analyses of
which two are of German bleach examined by Lunge and one of Canadian
manufacture analysed by the author.

                                 |     Lunge.     | Race.
                                 |   %    |   %   |   %
  Available chlorine             | 37.00  | 38.30 | 37.50
  Chlorine as chlorides          |  0.35  |  0.59 |  0.52
  Chlorine as chlorates          |  0.25  |  0.08 |  0.18
  Lime                           | 44.49  | 43.34 | 44.12
  Magnesia                       |  0.40  |  0.31 |  1.28
  Iron oxide                     |  0.05  |  0.04 |  0.11
  Alumina                        |  0.43  |  0.41 |  0.46
  Carbon dioxide                 |  0.18  |  0.31 |  0.22
  Silica                         |  0.40  |  0.30 |  0.52
  Water and undetermined         | 16.45  | 16.32 | 15.09

From these analyses the constitutional of commercial bleach might be
represented by the formula


which assumes it to contain:

      68.0 per cent of calcium hypochlorite,
      20.0 per cent of calcium hydroxide,
  and 12.0 per cent of water.

In this formula calcium hypochlorite has been written CaOCl_{2}, but
this substance actually contains one atom of oxygen less than the true
hypochlorite, which has the constitutional formula ClO-Ca-OCl. This
difference led some of the earlier chemists to regard CaOCl_{2} as a
mixture of equal molecules of calcium chloride and calcium hypochlorite
(CaCl_{2} + Ca(OCl)_{2} = 2CaOCl_{2}), but it has been definitely
established that no calcium chloride exists in the free state in dry
commercial bleach.

Since the very earliest days when the process of bleaching was
investigated it was considered to be a process of oxidation and it is
not surprising that Lavoisier and his pupils, who had noted the strong
decolourising action of the gas discovered previously by Scheele, should
regard it as a compound that contained oxygen. They were confirmed in
this view by the fact that an aqueous solution of the gas slowly evolved
oxygen when placed in bright sunlight, and lost its bleaching
properties. Watt disproved this and showed that the evolution of oxygen
was due to the action of the chlorine on water.

  Cl_{2} + H_{2}O = 2HCl + O.

The bleaching action was not due to the chlorine "per se" but to the
nascent oxygen produced in the presence of moisture. Later, when bleach
and other chlorine compounds came into use as deodourisers, their action
was attributed to the oxygen produced and when their germicidal
properties became known it was natural to assume that the destruction of
bacteria was due to the same cause. Some of the earlier experimental
work supported this view. Fischer and Proskauer[1] found that humidity
played an important part in chlorine disinfection, probably because it
favoured oxidation. In air saturated with moisture micro-organisms were
killed by 0.3 per cent of chlorine in three hours but when the air was
dry practically no action occurred. They concluded that chlorine was not
directly toxic. Warouzoff, Winogradoff, and Kolessnikoff[2] were unable
to confirm the results of Fischer and Proskauer and found that a mixture
of chlorine gas and air killed tetanus spores in one minute.

The nascent oxygen hypothesis was clearly and succinctly expressed by
Prof. Leal during the hearing of the Boonton, N. J., case and the
following abstracts have been taken from his evidence:

"... That on the addition of bleach to water the loosely formed
combination forming the bleach splits up into chloride of calcium and
hypochlorite of calcium. The chloride of calcium being inert, the
hypochlorite acted upon by the carbonic acid in the water either free or
half bound, splits up into carbonate of calcium and hypochlorous acid.
The hypochlorous acid in the presence of oxidisable matter gives off its
oxygen; hydrochloric acid being left. The hydrochloric acid then drives
off the weaker carbonic acid and unites with the calcium forming
chloride of calcium.

"That the process was wholly an oxidising one, the work being done
entirely by the oxygen set free from the hypochlorous acids in the
presence of oxidizable matter....

"We have used during our investigations, the term 'potential oxygen' as
expressing its factor of power. When set free, it is really nascent or
atomic oxygen and is, in its most active state, entirely different from
the oxygen normally in water...."

The reactions suggested are expressed in the following equations:

    (i).  2CaOCl_{2} = CaCl_{2} + Ca(OCl)_{2}

   (ii). Ca(OCl)_{2} + CO_{2} + H_{2}O = CaCO_{3} + 2HClO

  (iii). 2HClO = 2HCl + O_{2}

   (iv). CaCO_{3} + 2HCl = CaCl_{2} + CO_{2} + H_{2}O.

Phelps, during the hearing of this case, suggested that hypochlorites
were directly toxic to micro-organisms but this view was not supported
by any definite evidence and the nascent oxygen hypothesis met with
almost universal acceptance. Investigations made by the author in 1915,
1916 and 1917 have produced data which cannot be adequately explained by
the nascent oxygen hypothesis.[3]

The disinfecting action of bleach can be most conveniently considered by
regarding it as a heterogeneous mixture of the reactants and resultants
of the reaction

  CaO + H_{2}O + Cl_{2} <=> CaOCl_{2} + H_{2}O

which is in equilibrium for the temperature and pressure obtaining
during the process of manufacture. Under suitable physical conditions
the chlorine content can be increased to 40-42 per cent but such a
product is not so stable as those represented by the analyses on page 14
and which contain approximately 20 per cent of excess hydrate of lime.
The stability of bleach depends upon this excess of base (Griffen and
Hedallen[4]) and although magnesia can be partially substituted for this
excess of lime, a minimum of 5 per cent of free hydrate of lime is
required to ensure stability.

On dissolving bleach in water the first action is the decomposition of
calcium oxychloride into an equal number of molecules of calcium
hypochlorite and calcium chloride.

  2CaOCl_{2} = Ca(OCl)_{2} + CaCl_{2}.

In dilute solution these salts are dissociated and hydrolysis tends to
occur in accordance with the equations

  2Ca(OCl)_{2} + 4H_{2}O <=> 2Ca(OH)_{2} + HOCl + HCl and

  CaCl_{2} + 2H_{2}O <=> Ca(OH)_{2} + 2HCl.

Calcium hydrate and hydrochloric acid are both practically completely
dissociated, i.e. there is a large and equal quantity of H^{.} and OH',
and the product is much greater than K_{_w_} (ionic product of water),
and hence there is a combination of these ions, leaving the solution
neutral and no undissociated acid or base exists. This statement is only
approximately correct as hydrochloric acid is slightly more dissociated
than calcium hydroxide (ratio 9:8) and the solution is consequently
slightly acid, i.e. the H^{.} concentration is greater than 1 × 10^{-7}.

Hypochlorous acid is only very slightly dissociated, especially in the
presence of the OCl' ion due to the dissociation of the Ca(OCl)_{2}, as
compared with Ca(OH)_{2} and hydrolysis of the Ca(OCl)_{2} proceeds with
increased dilution. The action is best represented by the equation

  2Ca(OCl)_{2} + 2H_{2}O <=> CaCl_{2} + Ca(OH)_{2} + 2HOCl

The hydrolytic constant of hypochlorous acid has apparently not been
determined but as the acid is weaker than carbonic acid, which has a
hydrolytic constant of 1 × 10^{-4}, the value is probably between 1 ×
10^{-3} and 1 × 10^{-4}. From the formula _x_^{2}/(1 - _x_)_v_ =
_k__{_wv_} in which 1 mole of pure Ca(OCl)_{2} is dissolved in _v_
litres, _x_ is the fraction hydrolysed, and _k__{_wv_} is the hydrolytic
constant, complete hydrolysis occurs (_x_ = 1) when _v_ is not greater
than 1 × 10^{4} litres. This is equivalent to a concentration of not
less than 7.1 p.p.m. of available chlorine. Solutions of pure
hypochlorites are alkaline in reaction because of the excess of hydroxyl
ions (minimum concentration 1 × 10^{-4}). In solutions of bleach the
hydrolytic action is retarded by the OH' due to the free base, and
accelerated by the excess of H^{.} caused by the dissociation and
partial hydrolysis of CaCl_{2}; the final result is determined by the
relative proportions and the effect of the free base usually
preponderates. The addition of any substance that reduces the OH'
concentration enables hydrolysis to proceed to completion and affords a
rational explanation of the fact that solutions of bleach, on
distillation with such weak acids as boric acid, yield a solution of
hypochlorous acid. It also explains why the addition of an acid is
necessary in Bunsen's method (_vide_ p. 79) of analysing hypochlorite
solutions. It has been stated that when hydrochloric acid is employed
the increase in the oxidising power is due to the action of the acid
upon calcium chloride but this never occurs under ordinary conditions;
weak acids such as carbonic or acetic will give practically the same
result as hydrochloric acid in solutions of bleach of the strength used
in water treatment. The slightly higher result obtained with strong
acids is due to the decomposition of chlorates.

The effect of dilution alone is shown by the data given below. A 2 per
cent bleach solution, containing very little excess base, was diluted
with distilled water and the various dilutions titrated with
thiosulphate after the addition of potassium iodide. In one series the
solutions were titrated directly, and after acidification in the other.
The results[A] were as follows:


  Strength of Solution. Grams Bleach | Direct Titration × 100
           Per 100 c.cms.            |  --------------------.
                                     |    Acid Titration
               2.0                   |         30.8
               0.2                   |         34.3
               0.1                   |         41.8
               0.02                  |         67.5
               0.002                 |        100.0

  [A] Corrected for the alkali produced by HClO + 2KI = KCl + KOH +

Although every precaution was taken to exclude carbonic acid, a portion
of the hydrolysis was probably due to this acid, which would remove
calcium hydrate from the sphere of action and consequently alter the
equilibrium. The above figures are only applicable to the particular
sample used; other samples containing different excesses of base would
yield different hydrolytic values. The results are in agreement with the
hypothesis presented and confirm the theoretical deduction that very
dilute bleach solutions are completely hydrolysed if no salts are
present that will dissociate and increase the OH' concentration.
Hydrolysis is reduced by caustic alkalies and alkaline carbonates, and
increased by acids and acid carbonates that reduce the OH'

The effect of chlorides is anomalous and no adequate explanation for
their action can be given at present. The addition of small quantities
of sodium chloride (0.1 per cent) increases the hydrolysis of bleach
solutions but much larger quantities tend to the opposite direction.

The effect of these substances upon the velocity of the germicidal
action of bleach solutions is in the same direction as the hydrolysing
effect.[4] Sodium chloride in quantities up to 10 parts per million has
a very limited effect but larger quantities (90 p.p.m.) increase the
velocity of the reaction. Sodium chloride, in the absence of
hypochlorites, was found to have no influence upon the viability of _B.
coli_ in water.

In quantities up to approximately 5 p.p.m., sodium hydroxide has but
little influence; 5-10 p.p.m. reduce the velocity to a marked degree,
but when the quantity of caustic is still further increased the
germicidal action of the alkali commences to be appreciable and may
nullify the retarding action on the hypochlorite. Normal carbonates tend
to reduce the velocity of the germicidal action and bicarbonates to
increase it.

Sulphuric acid, even in very small quantities (5 p.p.m.), has a marked
accelerating effect and the total effect produced is much greater than
can be accounted for by the germicidal activity of the acid alone. Weak
acids such as carbonic acid and acetic acid are also effective
accelerators. In one experiment a 0.01 per cent solution of bleach was
found to be 40 per cent hydrolysed. By passing carbonic acid gas this
was increased to 95 per cent and the velocity of the germicidal action
of this solution was found to be approximately 100 per cent greater than
that of the uncarbonated one. Norton and Hsu[5] have shown that the
germicidal activity of some disinfectants is a function of the hydrogen
ion concentration, but this factor is insufficient to account for the
effect of acids on bleach solutions.

The effect of sodium chloride on the bacteriological results, like that
on the hydrolytic constant, is anomalous. Similar effects have been
observed on the addition of this salt to phenol and other disinfectants.
The _raison d'être_ of the increased activity is obscure but it is
possible that the salt renders the organisms more susceptible to the
action of the germicide.

Ammonia, though decreasing the hydrogen ion concentration of bleach and
other hypochlorite solutions, markedly increases the velocity of the
reaction; chlorinated derivatives of ammonia (chloramines), which have a
specific germicidal action, are formed. These will be discussed at
length in Chapter IX, p. 115.

Rideal[6] has shown that the addition of ammonia to sodium hypochlorite
destroys the bleaching activity in acid solution. This has been found by
the author to be also true for calcium hypochlorite (bleach). If the
bleaching effect is due to oxidation, the oxidising power of
hypochlorites must be considered to be destroyed by the addition of
ammonia. The property of oxidising organic matter in water is also
destroyed; this is well illustrated in Table II which shows the rate of
absorption of chlorine and chloramine by the Ottawa River water. The
water used in this experiment contained 40 p.p.m. of colour and absorbed
9.5 p.p.m. of oxygen (30 mins. at 100° C.).


                    |   ABSORPTION OF AVAILABLE CHLORINE AT 63° F.
   Time of Contact  +----------------------+--------------------------
      Minutes.      | Chlorine as Bleach.  | Chlorine as Chloramine.
          Nil.      |        10.00         |         9.98
          5         |         6.50         |         9.98
         10         |         5.91         |         9.90
         20         |         5.18         |         9.90
         40         |         4.47         |         9.84
         60         |         3.90         |         9.84
         80         |         3.65         |         9.84
         20 hours   |         ....         |         9.68

  [A] Results are parts per million.

From a consideration of these and other experiments made by the author
in January, 1916, it became apparent that the nascent oxygen hypothesis
entirely failed to explain the results obtained, and that they must be
attributed to a direct toxic action of the chlorine or chloramine.

Dakin et al.[7] arrived at a similar conclusion from a consideration of
the results obtained during the use of hypochlorite solutions in the
treatment of wounds by Carrel's method of irrigation. They attributed
the marked beneficial action to the formation of chloramines _in situ_
by the action of hypochlorous acid upon amino acids and proteid bodies.
Compound chloramines (chlorinated aminobenzoic acids) were prepared in
the laboratory and found to give excellent results in reducing wound
infection. Later, other compounds were prepared for the purpose of
sterilising small quantities of water for the use of mobile troops (see
p. 128).

Rideal[6] was the first to note the strong germicidal power of
chloramine and attributed the persistent germicidal activity of
hypochlorites in sewage to the formation of chloramine and chloramine

Further evidence against the nascent oxygen theory of chlorine
disinfection is to be found in the fact that such active oxidising
agents as sodium, potassium, and hydrogen peroxides have a much lower
germicidal activity than chlorine when compared on the basis of their
oxygen equivalents. Table III shows chlorine to be approximately five
times as active as potassium permanganate when compared on this basis.


             |     BLEACH       |
             |Available Chlorine|     POTASSIUM
  Contact    |    0.35 p.p.m.   |   PERMANGANATE.
    Period.  +------------------+-------------------
             |Oxygen Equivalent. Parts Per Million.
             |  0.08  |  0.133  |  0.266  |  0.400
  Nil        |  140   |   ...   |   ...   |   ...
  30 mins    |   90   |   122   |   115   |   110
  1 hour     |   68   |   115   |   100   |    80
  1-1/2 hours|   63   |   108   |    95   |    75
  4 hours    |   50   |    95   |    80   |    50

  [A] Results are _B. coli_ per 10 c.cms.

The germicidal activity of oxidising agents has been shown by Novey and
others to be somewhat proportional to the energy liberated during the
reaction but even when this factor is taken into consideration chlorine
compounds are more active than other oxidising agents. Hypochlorous acid
is far superior to hydrogen peroxide as a germicidal agent and is as
active as ozone, which liberates a greater amount of energy.

  2HClO = 2HCl + O_{2} + 18,770 calories

  2H_{2}O_{2} = 2H_{2}O + O_{2} + 46,120 calories

  2O_{3} = 3O_{2} + 60,000 calories.

Again, solutions of chlorine gas and hypochlorites having the same
oxidising activity, as determined by titration with thiosulphate after
the addition of potassium iodide and acid, i.e. contain equal amounts of
available chlorine, show approximately the same germicidal activity in
water. On the addition of ammonia, the hypochlorite solutions retain
their ability to liberate iodine from potassium iodide (Wagner test) but
the property of oxidising such dyestuffs as indigo is destroyed and the
germicidal activity is increased. Ammonia, when added to solutions of
chlorine gas, diminishes the property of liberating iodine from
potassium iodide, the bleaching effect on dyestuffs, and the germicidal
action. It is often assumed that chlorine forms hypochlorous acid on
solution in water Cl_{2} + H_{2}O = HClO + HCl but the results obtained
on the addition of ammonia indicate that either very little hypochlorous
acid is formed or that ammonia and hypochlorous acid do not form
chloramine in the presence of hydrochloric acid.

When chlorine gas was treated with a 0.5 per cent solution of ammonia in
the proportion of 1 molecule of chlorine to 1.90-1.95 molecules of
ammonia, Noyes and Lyon[8] found that nitrogen and nitrogen-trichloride
were formed in equimolar quantities.

  12NH_{3} + 6Cl_{2} = N_{2} + NCl_{3} + 9NH_{4}Cl.

Bray and Dowell[9] showed that this reaction depended upon the hydrogen
ion concentration and proceeded in accordance with the following

   (i). Acid solution 4NH_{3} + 3Cl_{2} = NCl_{3} + 3NH_{4}Cl

  (ii). Alkaline solution 8NH_{3} + 3Cl_{2} = N_{2} + 6NH_{4}Cl.

In (i) with a ratio of chlorine to ammonia of 3:1 by weight, one-half of
the chlorine is lost as ammonium chloride and one-half forms nitrogen
trichloride, concerning which comparatively little is known; in (ii) the
whole of the chlorine forms ammonium chloride, which has no germicidal

The effect of ammonia on the germicidal action of a solution of chlorine
gas is shown in the Table IV.


_Conditions._ Colour of water 40 p.p.m. Turbidity, 5 p.p.m.

           |Available Chlorine 0.20 p.p.m., Ammonia.
   Contact |           Parts Per Million.
   Period. +---------+---------+----------+---------
           |   Nil.  |   0.05  |   0.10   |   0.20
  Nil.     |   130   |   ...   |   ...    |   ...
  10 mins. |   135   |   140   |   130    |   135
   1 hour  |   130   |   130   |   128    |   120
   4 hours |   120   |   112   |   110    |   105
  24 hours |   120   |   145   |   160    |   170

  [A] Results are _B. coli_ per 10 c.cms.

Even when the ratio of Cl:NH_{3} was 4:1 by weight, practically the same
as was used in the experiments of Noyes and Lyon, and Bray and Dowell,
quoted above, the germicidal action was totally destroyed and the
24-hour results showed aftergrowths which were somewhat proportional to
the amount of ammonia added. This was probably due to the formation of
ammonium chloride, which provided additional nutriment for the

It has often been assumed that hypochlorite solutions are decomposed on
addition to water containing free or half-bound carbonic acid with the
production of free chlorine, but no evidence has been adduced in
support. Free chlorine can be separated from hypochlorous acid in
aqueous solution by extraction with carbon tetrachloride and when this
solvent is shaken with a carbonated hypochlorite solution it is found
that only traces of chlorine are removed.

Hypochlorous acid reacts with hydrochloric acid with the evolution of
free chlorine HClO + HCl = Cl_{2} + H_{2}O but in very dilute solution
the amount of free chlorine formed is exceedingly minute. Jakowkin[10]
has shown that this reaction does not proceed to completion and that the
concentration of free chlorine can be calculated from the equation HClO
× H^{.} × Cl' = 320Cl_{2} in which the reactions are expressed in gram
molecules per litre. The hydrogen ions and chlorine ions are obtained
from the dissociation of carbonic acid (H_{2}CO_{3} <=> H^{.} +
HCO_{3}') and chlorides (NaCl <=> Na^{.} + Cl') and also by the
dissociation of hydrochloric acid produced by the interaction of
hypochlorous acid and organic matter. HClO = O + HCl <=> H^{.} + Cl'. If
the formula of Jakowkin can be correctly applied to solutions containing
fractions of a part per million of hypochlorous acid the free chlorine
liberated by the addition of 1 p.p.m. of bleach to a water low in
chlorides would be of the order 10^{-7}-10^{-8} p.p.m.

_Sodium hypochlorite_ is probably hydrolysed in dilute solution in a
manner similar to that of bleach.

  2NaOCl = NaCl + NaOH + HClO.

For solutions containing equal amounts of available chlorine,
electrolytic sodium hypochlorite is more dissociated than bleach because
of the absence of an excess of base, and this, together with the
presence of sodium chloride, accounts for the slightly higher germicidal
velocity obtained. The experience of pulp mills, with bleach and
electrolytic hypochlorites, confirms this: the latter is a much quicker
bleaching agent than bleach and it is often so rapid as to make it
desirable to reduce the velocity by the addition of soda ash.

Regarding hypochlorite solutions a phenomenon of more scientific
interest than of practical importance has been noted by Breteau[12] who
found that alkaline solutions of sodium hypochlorite containing 0.94 per
cent of available chlorine lost 3.6 per cent of their titer on dilution
with 80 volumes of water; also that this loss was increased by the
addition of small quantities of salt (sodium chloride) and more so by
carbonates and bicarbonates. The author has noted similar losses on
diluting bleach solutions and that the loss increased on standing. The
loss can be explained by the decomposition of hypochlorous acid, in the
presence of light, into hydrochloric acid and oxygen. 2HClO = 2HCl +

CHLORINE WATER. When a solution of chlorine in water is used as a
germicide the chemical reactions that occur differ materially from those
of hypochlorite solutions. On solution in water, hydration or solvation
probably takes place with the production of heat. Cl_{2}·Aq. = 2,600
calories. Chlorine water is comparatively stable but decomposes under
the influence of light in accordance with the equation Cl_{2} + H_{2}O =
2HCl + O; a similar reaction occurs in the presence of organic matter or
any substance capable of oxidation. Chlorine water contains only minute
traces of hypochlorous acid and there is no evidence that the
endothermic reaction

  Cl_{2}·Aq + H_{2}O = HClO·Aq + HCl·Aq
    -2600   - 68,460 = -29,930 - 39,315 - 1815

occurs in a measurable degree.

From thermochemical considerations hypochlorous acid and chlorine water
should be about equally active as oxidising agents.

  2HClO·Aq = 2HCl + O_{2} + 18,770 calories

  2Cl_{2}·Aq + 2H_{2}O = 2HCl + O_{2} + 15,340 calories

  2Cl_{2}· + Aq + 2H_{2}O = 2HCl + O_{2} + 20,540 calories

When a solution of chlorine or hypochlorite is added to water as a
germicidal agent, a variety of reactions occur the character of which is
determined by the nature of the mineral and organic matter in the water
and the type of chlorine compound added. The general reactions are of
three types (1) oxidation of the organic matter, (2) direct chlorination
of the organic matter, and (3) a bactericidal action.

In the treatment of waters that contain appreciable amounts of organic
matter almost all the chlorine is consumed in reaction (1) and even with
filter effluents it is probably true that oxidation accounts for the
greater portion of the chlorine consumed. The author has found that a
dosage of 0.02 part per million of available chlorine was more effective
in destroying _B. coli_ in distilled water than 0.40 p.p.m. in a water
absorbing 9.5 p.p.m. of oxygen (30 mins. at 100° C.).

Reaction (1) can be adequately explained by the nascent oxygen
hypothesis and it is this reaction that determines the dosage required
for effective sterilisation. (See Chap. III.)

Very little information is available regarding reaction (2) but there is
little doubt that a direct chlorination of the organic matter does occur
and it is more than probable that these chlorinated derivatives are
largely responsible for the obnoxious tastes and odours produced in some
waters. It has been suggested that these were due to the formation of
chloramines. This view was formerly supported by the author but the
chloramine treatment at Ottawa and other places has demonstrated the
inadequacy of this explanation. It is true that the odour of chloramine
is stronger and more pungent than that of chlorine, but chloramine in
the Ottawa supply, even with doses as high as 0.5 part per million of
available chlorine, has caused no complaints.

The odour of some of the organo-chloro compounds is more penetrating and
obnoxious than those of chlorine and chloramine, and it is quite
possible that some of the higher homologues of chloramine are in this
class. It should be noted, however, that some of the chloro-amido
compounds prepared by Dakin are white, odourless, crystalline

Practically nothing is known regarding the specific nature of the
mechanism involved in reaction (3). The hypothesis that chlorine, and
chlorine compounds, exert a direct toxic action on the micro-organisms
marks an advance in the science of water treatment but does not indicate
the physiological processes involved. Cross and Bevan[11] have shown
that chloro-amines have a tendency to combine with nitrogenous molecules
and to become fixed on cellulose; it is therefore possible that reaction
is a cytolytic one in which the chlorine attacks and partially or wholly
destroys the membranous envelope of the organisms. A portion of the
chlorine or chlorine-compound may also penetrate the membrane and
produce changes that result in the death of the organism.


[1] Fischer and Proskauer, Rev. d'Hyg., 1884, 6, 515.

[2] Warouzoff, Winogradoff, and Kolessnikoff. Russkaia medicina, 1886,
Nos. 3 and 32.

[3] Race. Jour. Amer. Water Works Assoc., 1918, 5, 63.

[4] Griffen and Hedallen. Jour. Soc. Chem. Ind., 1915, 34, 530.

[5] Norton and Hsu. Jour. Inf. Dis., 1916, 18, 180.

[6] Rideal, S. Jour. Roy. San. Inst., 1910, 31, 33.

[7] Dakin, Cohen, Duafresne, and Kenyon. Proc. Roy. Soc., 1916, 89B,

[8] Noyes and Lyon. Jour. Amer. Chem. Soc., 1901, 23, 460.

[9] Bray and Dowell. Jour. Amer. Chem. Soc., 1917, 39, 905.

[10] Jakowkin. Zeit. f. Phys. Chim., 1899, 19, 613.

[11] Cross and Bevan. Jour. Soc. Chem. Ind., 1898, 28, 260.

[12] Breteau. Jour. Pharm. Chim., 1915, 12, 248.



The amount of chlorine required for efficient treatment is very largely
determined by the amount required to satisfy the oxidisable matter
present in the water. Many experimenters have reported results that
would indicate that appreciable concentrations of chlorine are required
for bactericidal action but the details of the technique, as published,
show that the effect of the organic matter added with the test organism
was not thoroughly appreciated. One cubic centimetre of a culture in
ordinary peptone water, added to one litre of water, would increase the
organic content by approximately 10 parts per million, an amount that
would absorb appreciable amounts of chlorine.

Other conditions also make it very difficult to compare the results
obtained in the past: one of these is the degree of purity set as the
objective. German bacteriologists added enormous numbers of the test
organism and endeavoured to obtain the complete removal of the organism
from such quantities as one litre of water with a contact period often
as short as 10 minutes. Nissen,[1] of the Hygienic Institute of Berlin,
found that a 1: 800 dilution of bleach (420 p.p.m. of chlorine) was
required to destroy _B. typhosus_ in one minute and a 1:1600 dilution
(210 p.p.m. of chlorine) in 10 minutes. Delépine[2] obtained somewhat
similar results by means of the thread method for testing disinfectants.
Phelps,[3] using gelatine plates for enumeration of the bacteria,
obtained a 90 per cent reduction of _B. typhosus_ in twenty minutes
with 5 p.p.m. of available chlorine; over 99 per cent reduction in one
hour, and over 99.99 per cent reduction in 18 hours. Wesbrook,
Whittaker, and Mohler[4] tested bleach solutions with various strains of
_B. typhosus_ by means of the plate method and found that the most
resistant one was reduced from 20,000 per c.cm. to sterility (in 1
c.cm.) by 3 p.p.m. of available chlorine in fifty minutes and that the
least resistant one only required 1.0 p.p.m. with a thirty minutes'

Lederer and Bachmann[5] have reported the following results:



            |               NATURE OF TEST ORGANISM.
  Available |       |   B.  |       |       |       |   B.  |
  Chlorine  |   B.  |fæcalis|  B.   |Proteus|  B.   |lactis | B.
  p.p.m.    |cloacæ.|alkali-|paraty-| mira- |enter- | aero- |choler[oe]-
            |       |genes. |phosus.|bilis. |itidis.| genes |suis.
  0.1       |  .....|  99.98|  .....|  27.3 |  .....|  .....| .....
  0.2       |  99.69|  99.99|  99.97|  45.5 |  99.83|  99.17|  95.8
  0.3       |  99.75| 100.00| 100.00|  63.7 |  99.98|  99.98| 100.0
  0.5       | 100.00|  .....|  .....|  72.7 | 100.00| 100.00| .....
  0.7       |  .....|  .....|  .....|  63.7 |  .....|  .....| .....
  1.0       |  .....|  .....|  .....|  63.7 |  .....|  .....| .....
  3.0       |  .....|  .....|  .....|  90.9 |  .....|  .....| .....
  5.0       |  .....|  .....|  .....|  90.0 |  .....|  .....| .....
  Original }|       |       |       |       |       |       |
  number of}|       |       |       |       |       |       |
  organisms}|160,000|  9,500|  3,000| 8,000 |180,000|180,000|  500
  per c.cm.}|       |       |       |       |       |       |

With the exception of _P. mirabilis_, which forms endospores, all the
organisms were killed (less than 1 per c.cm.) by 0.5 p.p.m. of available
chlorine in fifteen minutes.

All these observers found that _B. coli_, the organism usually employed
as an index of contamination, had approximately the same degree of
resistance to chlorine as _B. typhosus_, though Wesbrook et al. directed
attention to the varying viability of organisms derived from different

These experiments merely indicate the dosage required for exceptional
conditions such as it is inconceivable would ever occur in water-works
practice. No information is available regarding the actual _B. typhosus_
content of waters that have caused epidemics of typhoid fever, but for
the present purpose it may be assumed that the extreme condition would
be a pollution by fresh sewage giving a _B. coli_ content of 1,000 per
c.cm. or 200 times worse than the average condition that can be
satisfactorily purified without overloading a filter plant (500 _B.
coli_ per 100 c.cms.). Experiments made by the author indicate that a
suspension of 1,000 _B. coli_ per c.cm. in water, in the absence of
organic matter, can be reduced to a 2 _B. coli_ per 100 c.cms. standard
(the U.S. Treasury Standard) by 0.1 p.p.m. of available chlorine in ten
minutes at 65° F. This experiment indicates the amount of chlorine that
is required for the bactericidal action only; such a dosage could never
be used in practice to meet a pollution of this degree because of the
accompanying organic matter. In actual practice the author has
experienced the above condition but once, and on that occasion the _B.
coli_ were derived from soil washings and not from fresh sewage.

The amount of chlorine required for germicidal action is small, and the
main factors that determine the dosage necessary to obtain this action
are (1) the content of readily oxidisable organic matter, (2) the
temperature of the water, (3) the method of application of the chlorine
and (4) the contact period.

=Oxidisable Matter.= The oxidisable matter may be divided into two
classes (_a_) inorganic and (_b_) organic. The inorganic constituents
naturally found in water, that are readily oxidisable, are ferrous salts
(usually carbonates), nitrites, and sulphuretted hydrogen, and these
react quantitatively with chlorine until fully oxidised. The oxygen
value of chlorine is approximately one-quarter (actually 16: 71) the
available chlorine content in accordance with the equation Cl_{2}/71 +
H_{2}O = 2HCl + O/16. One part per million of available chlorine will
oxidise 1.58 p.p.m. of ferrous iron; 0.197 p.p.m. of nitrous nitrogen;
and 0.479 p.p.m. of sulphuretted hydrogen.



                  | Water "A" Colour 3 | Water "B" Colour 40
                  | Available Chlorine | Available Chlorine
  Contact Period. |       p.p.m.       |      p.p.m.
                  |        0.2         |  0.2 |  0.4 |  0.5
  Nil             |        194         |  194 |  194 |  194
  5 minutes       |        121         |  165 |  129 |   66
  1 hour          |          7         |   95 |   20 |    1
  5 hours         |          0         |    4 |    0 |    0
  24 hours        |          0         |    1 |    1 |    0
  48 hours        |          0         |    0 |    0 |    0

  [A] Results are _B. coli_ per 10 c.cms. of water.

The organic matter found in water may be derived from various substances
such as urea, amido compounds, and cellulose; humus bodies derived from
soil washings and swamps may also be present. The humus compounds of
swamps and muskeg are usually associated with the characteristic colour
of the water derived from these sources. The effect of this coloured
organic matter upon the chlorine dosage is well illustrated in Table VI.
In this experiment _B. coli_ was used as the test organism and the only
varying factor was the organic matter. To obtain the same result with a
contact period of one hour at 63° F. it was necessary to use about two
and one-half times the amount of chlorine with a water containing 40
p.p.m. of colour as with one practically free from colour. It will be
noted that water "A," in which the colour had been reduced to 3 p.p.m.
by coagulation with aluminium sulphate, required a greater dosage of
chlorine than was necessary for bactericidal action only. This was due
to a residual organic content which produced none or but a trace of
colour, for although the colour had been reduced by 92 per cent the
organic matter, as measured by the oxygen absorbed test, had only been
reduced by 70 per cent.

The results obtained by Harrington[6] at Montreal are in the same
direction. During the greater part of the year the water is obtained
from the St. Lawrence river, which is colourless and low in organic
matter; in the spring months the flood waters of the Ottawa, a highly
coloured river, enter the intake and necessitated a much higher dosage.


               |        |       | Oxygen | Chlorine|          |
   Source of   | Alkali-|Colour.|Absorbed| Required| Bacteria | Per Cent
    Supply.    |  nity. |       |  (30   |  p.p.m. | per c.cm.| Removed.
               |        |       | mins.) |         |          |
  Ottawa river | 15-20  | 50-70 |  14.0  |  1.50   |  3,000   | over 98
  St. Lawrence |        |       |        |         |          |
    river      | 90-100 |  Nil. |   0.30 |  0.30   |    500   | over 99

Ellms[7] obtained similar results and reported "that the rate at which
sterilisation proceeds varies, in a general way, directly with the
concentration of the applied available chlorine and the temperature, and
inversely as the amount of easily oxidisable matter present."

Experience with filter plants shows the same facts, the amount of
chlorine required for the sterilisation of a filter effluent being
invariably less than that necessary to purify the raw water to the same

The effect of coloured organic matter upon the absorption of chlorine,
in the form of hypochlorite, is shown on Diagram I.

[Illustration: DIAGRAM I


  |                               | Value of K calculated from |
  |                               |                            |
  |    Absorption of Chlorine     |     Log(N_{1}/N_{2})       |
  |      by water at 63° F.       | K = -----------------      |
  |                               |     _t__{2} - _t__{1}      |
  |                               |                            |
  +-------+                       +-------+ When _t__{1} = 0   |
  |Time of+-----------------------+Time of+--------------------+
  |Contact|    Colour of Water    |Contact|       Colour       |
  |  in   +-------+-------+-------+  in   +------+------+------+
  |Minutes|   3   |   25  |   40  |Minutes|   3  |  25  |  40  |
  |  Nil  | 10.00 | 10.00 | 10.00 |       |      |      |      |
  |   5   |  9.62 |  7.70 |  6.50 |   5   |0.0033|0.0227|0.0374|
  |  10   |  9.41 |  7.03 |  5.91 |  10   |0.0026|0.0153|0.0228|
  |  20   |  9.17 |  6.40 |  5.18 |  20   |0.0018|0.0096|0.0190|
  |  40   |  8.95 |  5.82 |  4.47 |  40   |0.0012|0.0057|0.0087|
  |  60   |  8.85 |  5.63 |  3.90 |  60   |0.0008|0.0041|0.0068|
  |  80   |  8.80 |  5.58 |  3.65 |  80   |0.0007|0.0032|0.0056|

The shape of the curve obtained with a colour of 40 p.p.m. somewhat
resembled that of a mono-molecular reaction and the results were
calculated accordingly. The mathematical expression of this law is
_dN_/_dt_ = _KN_ where _N_ is the concentration of the available
chlorine in parts per million. Integrating between _t__{1} and _t__{2}
the formula _K_ = log(_N__{1}/_N__{2})/(_t__{2} - _t__{1}) is obtained.
If the compound absorbing the chlorine were simple in character, and the
chlorine were present in large excess, the value of _K_ would be
constant. In the experiments recorded, _K_ constantly decreases, due to
the decreasing concentrations of the reacting substances and the complex
nature of the organic matter.

The results show the effect of organic matter on the reduction of the
chlorine concentration available for germicidal action and also the
importance of avoiding a local excess of chlorine (_vide_ p. 41).

An effort has been made by some observers to find a quantitative
relation between the organic matter, expressed as oxygen absorbed in
parts per million, and the chlorine required for oxidation, but without
definite result. Some of the results obtained are given in Table VII.


                           |        Oxygen Absorbed
        Observer.          | Ratio -----------------.
                           |       Chlorine Absorbed
  Rouquette                |           1
  Bonjean                  |           0.5
  Orticoni                 | Less than 1
  Valeski and Elmanovitsch |           0.4
  Race                     |           0.4
  Theoretical              |           0.22

The value of 0.4 (0.39) obtained by the author is the average of over
one hundred determinations covering a period of two years. The
experiments of Zaleski and Elmanovitsch were made with the water of the
Neva River.

The divergence in the ratios affords additional evidence in favor of
reaction (2) mentioned on page 28 and also shows that the chlorinated
compounds are less readily oxidized than those from which they are
produced. Heise[8] has found that the amount of chlorine consumed is
usually proportional to the concentration in which it is added though
not necessarily a function of the concentration of the organic matter.

=Temperature.= The evidence regarding the effect of temperature upon the
dosage required is somewhat conflicting. Ellms (_vide supra_) found that
the velocity of the germicidal action varied directly with the
temperature and this has also been the author's experience with
laboratory experiments. Typical examples of these are given in Tables
VIII and IX.



                 |Temperature, degrees, Fahrenheit.
  Contact Period.+-----------+----------+----------
                 |     36    |    70    |    98
  Nil            |    424    |   424    |   424
  5 minutes      |    320    |   280    |   240
  1.5 hours      |    148    |    76    |    12
  4.5 hours      |     38    |    14    |     3
  24 hours       |      2    |     0    |     0
  48 hours       |      2    |     0    |     0

  [A] Results are _B. coli_ per 10 c.cms.



                 |Temperature, degrees, Fahrenheit.
  Contact Period.+-----------+----------+----------
                 |     36    |    70    |    98
  Nil            |    240    |   240    |   240
    5 minutes    |    240    |   250    |   235
    1 hour       |    245    |   235    |   195
    4 hours      |    215    |   190    |   170
   24 hours      |    143    |   130    |   115
   48 hours      |    130    |    59    |    19
   72 hours      |    ...    |    28    |   ...
   96 hours      |    ...    |    16    |   ...
  120 hours      |    ...    |     6    |   ...

  [B] Results are _B. coli_ per 10 c.cms.

The reaction velocity of a germicide is proportional to the
temperature[9] and the influence of temperature may be mathematically
expressed by the formula _K__{1}/_K__{2} = _[theta]_(_T__{2} - _T__{1}),
in which _K__{1} and _K__{2} are the constants of the reaction at
temperatures _T__{2} and _T__{1}, respectively, and _[theta]_ is the
temperature coefficient. From the value of _[theta]_, the velocity
constant of a germicide for any temperature may be calculated from the
equation _K__{_T_} = _K__{20}° × _[theta]_^{(_T_ - _T__{20}°)}.
_K__{1} and _K__{2} are obtained from the formula _K__{_T_} =
log(_N__{1}/_N__{2})/(_t__{2} - _t__{1}) in which _N__{1} - _N__{2} is
the number of bacteria destroyed in the interval _t__{2} - _t__{1}.

A reduction of temperature also lowers the oxidizing activity of the
chlorine so that a greater concentration is available for germicidal
action. This is shown by the results plotted in Diagram II.

[Illustration: DIAGRAM II


  |                            |   Value of K calculated from     |
  |                            |      absorption at 63° F.        |
  | Absorption of Chlorine by  |                                  |
  | water containing 40 p.p.m. |         Log(N_{1}/N_{2})         |
  |        of colour           |     K = -----------------        |
  |                            |         _t__{2} - _t__{1}        |
  +-------+--------------------+                                  |
  |Time of|Temperature of Water+-------+--------+--------+--------+
  |Contact+--------------------+_t__{2}|_t__{1} |_t__{1} |_t__{1} |
  |Minutes|32° F.|46° F.|63° F.|minutes|  = 0   |  = 5   |  = 10  |
  |  Nil  |10.00 |10.00 |10.00 |       |        |        |        |
  |   5   | 8.00 | 7.45 | 6.50 |   5   | 0.0374 |  ----  |  ----  |
  |  10   | 7.23 | 7.09 | 5.91 |  10   | 0.0228 | 0.0082 |  ----  |
  |  20   | 7.00 | 6.60 | 5.18 |  20   | 0.0190 | 0.0066 | 0.0057 |
  |  40   | 6.42 | 6.05 | 4.47 |  40   | 0.0087 | 0.0043 | 0.0040 |
  |  60   | 6.22 | 5.60 | 3.90 |  60   | 0.0068 | 0.0040 | 0.0036 |
  |  80   | 6.13 | 5.40 | 3.65 |  80   | 0.0056 | 0.0033 | 0.0029 |

Tables VIII and IX, however, show that the temperature coefficient of
the germicidal action has a greater effect than the reduction in the
amount of chlorine absorbed and removed from the reaction.

The results obtained on the works scale with these waters are very
different to the laboratory ones and show that more chlorine is required
during the summer season than in winter. The results with bleach and
liquid chlorine are in the same direction (_vide_ Diagrams III and IV).
The bleach was regulated so as to maintain a constant purity, whilst in
the other case the dosage was constant with a varying _B. coli_ content.
In Diagram IV the _B. coli_ is plotted; this does not represent all the
factors involved as the _B. coli_ content of the treated water is also a
function of that of the raw water, but in the example given this factor
is of no moment because it was comparatively constant during the period
plotted (extreme variation 80 per cent).

The discrepancies between the laboratory and works results cannot be
easily explained. The only difference in the conditions is the nature of
the containing vessel. Glass is practically inert at all temperatures
but the iron pipes, through which the water passed before the samples
were taken, may exert an absorptive influence on the chlorine at the
higher temperatures experienced during the summer months.

Waters containing organic matter that differs much in quantity from the
examples above may yield very different results and no generalisation
can be made that will cover all cases. An increase of temperature
increases the germicidal velocity and also the rate of absorption of
chlorine by the organic matter; other factors determine which of these
competitive actions predominates.

=Method of Application (admixture).= A thorough admixture of the water
and chlorine is a _sine qua non_ for successful operation. This should,
if possible, be attained by natural means, but if there is any doubt as
to the efficiency of the mixing process, mechanical appliances should be
utilised. Pumps, especially centrifugal pumps, constitute a very
convenient and efficacious method of mixing the germicide and the
water, and the solutions should never be injected into the discharge
pipes when it is possible to make connections with the suctions.

[Illustration: DIAGRAM III


[Illustration: DIAGRAM IV


Inefficient admixture leads to local concentration of the chlorine, a
condition which (_vide_ p. 35), results in a wastage of the
disinfectant. Two practical examples of this effect may be cited. In one
case the water was free from colour and contained very little organic
matter. This water was chlorinated at one plant by allowing the bleach
solution to drop into one vertical limb of a syphon approximately 6,000
feet long, the other vertical limb being used as a suction well for the
pumps which discharged into the distribution mains. At the other plant
the bleach solution was injected into the discharge pipe of a
reciprocating pump through a pipe perforated with a number of small
holes. The results for two typical months are given in Table X.


        | Available | BACTERIA PER C.CM.    |
        |  Chlorine +-----------------------+
        | Parts Per |       |               | B. Coli Index
  Month.|  Million. |  Raw  |Treated Water. | Per 100 c.cms.
        +-----+-----+ Water.+-------+-------+-------+-------
        |  A. |  B. |       |   A.  |   B.  |   A.  |   B.
  July  | 0.20| 0.25|   864 |   27  |   93  |  <0.2 |   8.5
  August| 0.20| 0.27| 1.108 |   12  |  120  |  <0.2 |  10.2
  A = efficient mixing. B = inefficient mixing.

The results with the "B" plant were very irregular. The hypochlorite and
water did not mix thoroughly and, as several suctions pipes were
situated in the suction shaft, there was no subsequent admixture in the
pumps; this also caused complaints regarding taste and odour but the
complaints were localised, and not general as would result from an
overdose of solution due to irregularities at the plant.

The second example deals with a water containing 40-45 p.p.m. of colour.
This supply was taken from the river by low-lift pumps and discharged
into a header which was connected with the high-lift pumps by two intake
pipes about 5,000 feet in length. During 1914 a baffled storage basin of
two hours capacity was constructed and in June the hypochlorite was
added at the inlet to this basin by means of a perforated pipe. The
object was to increase the contact period prior to the delivery of the
water into the header. The results for this month were as follows:


                          | BACTERIA PER C.CM. AGAR.|
                          +------------+------------+  B. Coli.
                          |  3 Days at |  1 day at  |   Index
                          |    20 C.   |    37 C.   | Per c.cm.
  Raw water               |    410     |    104     |   0.280
  Treated water           |     49     |     26     |   0.036
  Percentage purification |     88.2   |     75.0   |  87.5

During August the point of application of the hypochlorite was changed
from the inlet of the basin to the suctions of the pumps and the
solution proportioned to the amount of water pumped by the starch and
iodide test. The average of the daily tests for this month were:


                          | BACTERIA PER C.CM. AGAR.|
                          +------------+------------+  B. Coli.
                          |  3 Days at |  1 day at  |   Index
                          |    20 C.   |    37 C.   | Per c.cm.
  Raw water               |    448     |    100     |   0.600
  Treated water           |     26     |     12     |   0.005
  Percentage purification |     91.9   |     88.0   |  99.2

Here again thorough admixture produced better results than inefficient
admixture plus a longer contact period. Langer[10] has also noted the
effect of local concentration and found that the disinfecting action is
increased by adding the bleach solution in fractions, a cumulative
effect replacing that of concentration.

The importance of the admixture factor was not thoroughly appreciated
during the earlier periods of chlorination but later installations, and
particularly the liquid chlorine ones, have been designed to take full
advantage of it.

The point of application in American water-works practice varies
considerably (Longley[11]). In 57 per cent of those cases in which it is
employed as an adjunct to filtration, it is used in the final treatment;
in 26 per cent it is used after coagulation or sedimentation and before
filtration; in the remaining 17 per cent it is applied before
coagulation and filtration. The report of the committee adds: "The data
at hand do not give any reasons for the application before coagulation.
In general, an effective disinfection may be secured with a smaller
quantity of hypochlorite, if it is applied after rather than before
filtration. It should be noted that the storage of chlorinated water in
coagulating basins, and its passage through filters, tend to lessen
tastes and odors contributed by the treatment and this fact may in some
cases account for its use in this way."

=Contact Period.= Other things being equal, the efficiency of the
treatment will vary directly, within certain limits, with the contact
period. When a chlorinated water has to be pumped to the distribution
mains directly after treatment, the dosage must be high enough to secure
the desired standard of purity within twenty to thirty minutes. The
chlorine is sometimes not completely absorbed in this period and may
cause complaints as to tastes and odours. The examples given above show
that the lack of contact period can be largely compensated by ensuring
proper admixture. Experience has amply demonstrated that there is no
necessity to use heroic doses for water that is delivered for
consumption almost immediately after treatment, and that, with proper
supervision, complaints can be almost entirely prevented.

The general effect of the effect of contact period is shown in Tables
VIII and IX on page 37. Another example of a coloured water is given in
Table XI, whilst Table XII shows the results obtained with a colourless


                   |  CHLORINE, PARTS PER MILLION.
   Contact Period. +-------+-------+-------+-------
                   |  0.30 |  0.40 |  0.55 |  1.21
  Nil              | 3,800 |   ... |  ...  |  ...
   1 minute        | 1,400 |   120 |    0  |    0
  10 minutes       |   720 |     5 |    0  |    0
  20 minutes       |    35 |     0 |    0  |    0

  [A] Results are _B. coli_ per 10 c.cms.


                   |        Sampling Point.         | Bacteria Per c.cm.
  Average of series|  5,000 ft. from pumping station|       300
    of samples     |  6,000  "   "     "       "    |       203
                   |  7,000  "   "     "       "    |       103
                   | 12,000  "   "     "       "    |        86
                   | 14,000  "   "     "       "    |        87

Table XIII is taken from the work of Wesbrook et al.[4]


               |      CONTACT PERIOD. (TEMP. 22°-26° C.).
  Available Cl.+---------+---------+---------+---------+---------
     P.p.m.    |         |  1 Hr.  |         |  6 Hrs. |
               | 30 Mins.| 30 Mins.| 3 Hrs.  | 30 Mins.| 24 Hrs.
      0        | 230,000 | 200,000 | 160,000 | 150,000 | 140,000
      0.5      |  14,000 |   7,400 |   2,000 |   6,000 |  11,000
      1.0      |      20 |      14 |     170 |     450 |  60,000
      1.5      |      10 |       6 |      16 |      45 |  70,000
      2.0      |       7 |       8 |      10 |      97 |  70,000
      2.5      |       7 |      14 |      30 |     116 |  65,000
      3.0      |       6 |      12 |       5 |      12 |  16,500

  [B] Results are bacteria per c.cm.

In Tables VIII, IX, XI, and XII, the bacteria decreased constantly with
increase of contact period, but the results in Table XIII show that no
advantage was to be gained by prolonging the contact beyond three hours;
after this period the bacteria commenced to increase in number and when
twenty-four hours had elapsed the number approached the original. This
increase in the bacteria is technically known as "aftergrowth" and will
be discussed more fully in Chapter IV.

The replies to queries sent out by the Committee on Water Supplies of
the American Public Health Association[11] indicate that the contact
period after treatment varies considerably in American water-works
practice. Forty per cent of the replies indicated no storage after
treatment; 18 per cent less than one hour; 9 per cent from one to three
hours; 5 per cent three to twelve hours; 11 per cent twelve to
twenty-four hours, and 17 per cent a storage of more than twenty-four

=Turbidity= is usually considered to exert an effect upon the dosage
required but no definite evidence has been adduced in support of this
hypothesis. Turbidity is generally caused by the presence of very finely
divided suspended matter, usually silt or clay, which is inert to
hypochlorites. The condition that produces turbidity, however, produces
a concomitant increase in the pollution and some of the organisms are
embedded in mineral or organic material that prevents access of the
chlorine to the organisms which consequently survive treatment. A larger
concentration is required to meet these conditions but it is not
necessitated by the turbidity _per se_.

=Effect of Light.= Light exerts a marked photo-chemical effect on the
germicidal velocity of chlorine and hypochlorites. When chlorinated
water is passed through closed conduits and basins the effect of light
is of course nil but in open conduits and reservoirs this factor is
appreciable and reduces the necessary contact period. The effect of
light on laboratory experiments made with colourless glass bottles is
so marked as to make it impossible to compare the results obtained on
different days under different actinic conditions. The following figures
illustrate the effect of sunlight:


                   |  AVAILABLE CHLORINE 0.35 P.P.M.
   Contact Period. +-------------------+-----------------
                   | Exposed to Bright | Stored in Dark
                   | Sunlight (April)  |    Cupboard.
   Nil             |        215        |      215
   30    minutes   |        130        |      145
   1     hour      |        122        |      136
   2-1/2 hours     |         61        |      130
   3-1/2 hours     |          0        |       32

=Determination of Dosage Required.= The dosage required for the
treatment of a water can only be accurately determined by treating
samples with various amounts of chlorine and estimating the number of
bacteria and _B. coli_ after an interval of time equal to that available
in practice. The temperature of the water during the experiment should
be the same as that of the water at the time of sampling.

In order to limit the range covered by the experiments the approximate
dosage can be ascertained from Diagram V if the amount of oxygen
absorbed by the water is known. This diagram is calculated on the amount
of available chlorine, present as chlorine or hypochlorite, that will
reduce the _B. coli_ content to the U. S. Treasury standard (2 _B. coli_
per 100 c.cms.) in two hours. If the oxygen absorbed values are
determined by the four-hour test at 27° C. they should be multiplied by

Another method which has been generally adopted for military work during
the war, consists in the addition of definite volumes of a standard
chlorine solution to several samples of the water and, after a definite
interval, testing for the presence of free chlorine by the
starch-iodide reaction. The details of the method of Gascard and
Laroche, which is used by the French sanitary service, have been given
by Comte.[12] One hundred c.cms. of the water to be examined are placed
in each of 5 vessels and 1, 2, 3, 4, and 5 drops of dilute Eau de
Javelle (1:100) are added and the contents stirred. After twenty
minutes, 1 c.cm. of potassium iodide-starch reagent (1 gram each of
starch, potassium, iodide and crystallized sodium carbonate to 100
c.cms.) is added and the samples again stirred. The lowest dilution
showing a definite blue colour is regarded as the dose required, and the
number of drops is identical with that required of the undiluted Eau de
Javelle for 10 litres of water when the same dropping instrument is
used. The actual concentration represented by these dilutions depends
necessarily upon the size of the drops and the strength of the undiluted
Eau de Javelle, but one drop per 100 c.cms. usually represents
approximately 1 p.p.m.

[Illustration: DIAGRAM V


In Horrocks's method, as used in the British army, a standard bleach
solution is added and is almost immediately followed by the zinc
iodide-starch reagent. The two methods were compared by Massy,[13] who
found that the French method gave an average result of only 0.06 m.gr.
per litre (0.06 p.p.m.) higher than the English method. Water in the
Gallipoli campaign required from 0.21 to 1.06 p.p.m. as determined by
both methods.

Diénert, Director of the Paris Service for investigating drinking water,
adds 3 p.p.m. of available chlorine and allows the mixture to stand
fifteen minutes after shaking; the residual chlorine is then titrated
with thiosulphate. The amount absorbed is increased by 0.5 p.p.m. and in
the opinion of Diénert this dosage is correct for a contact period of
three hours.

For military camps where a standpipe usually provides a reasonable
contact period, it has been found good practice to add sufficient
chlorine to give a rich blue colour with the starch-iodide reagent and
subsequently reduce the dosage gradually until the water, after standing
one hour, gives but a faint reaction to the test reagent. This method
should be checked up as soon as possible by bacteriological
examinations. An example of this method is given in Table XIV.


     Starch-iodide    |    BACTERIA ON AGAR PER C.CM.     |
   Reaction After One +-----------------+-----------------+ B. Coli Per
         Hour.        | 1 Day at 37 C.  | 2 Days at 20 C. | 100 c.cms.
        000[¤][¤]     |       40        |       15        |      0
        0000[¤]       |       37        |       18        |      8
        00000         |       68        |      268        |     34
        00000         |      115        |      553        |     61
    Raw water         |      114        |      685        |     89
  The number of [¤] signs indicates the intensity of the reaction.


[1] Nissen. Zeit. f. Hyg., 1890, =8=, 62.

[2] Delépine, J. Soc. Chem. Ind., 1911, =29=, 1350.

[3] Phelps. Water Supply Paper No. 220, U. S. Geo. Survey.

[4] Wesbrook, Whittaker, and Mohler, J. Amer. Public Health Assoc.,
1911, =1=, 123.

[5] Lederer and Bachmann. Eng. Rec., 1912, =65=, 360.

[6] Harrington. J. Amer. Waterworks Assoc., 1914, =1=, 438.

[7] Ellms. Eng. Rec., 1911, =63=, 472.

[8] Heise. Philippine Jour. Sci., 1917, =12=, A, 17-34.

[9] Norton and Hsu, Jour. Inf. Dis., 1916, =18=, 180.

[10] Langer. Zeit. f. Hyg., 1916, =81=, 296.

[11] Longley. J. Amer. Public Health Assoc., 1915, =5=, 920.

[12] Comte. J. Pharm. Chim., 1916, =14=, 261.

[13] Massy. J. Pharm. Chim., 1917, =15=, 209.



A disinfectant is usually described as a substance capable of destroying
bacteria and other micro-organisms, and an antiseptic as one that
restrains or retards their growth or reproduction. This distinction is
entirely arbitrary as the ability of a substance to kill organisms or
merely inhibit their growth depends upon the concentration employed.

Chlorine and hypochlorites, even in minute doses, exert a toxic effect
that is sufficient to produce death in organisms but when still smaller
concentrations are employed the toxic effect is transient and the
reproductive faculty may be entirely regained.

The enumeration of bacteria by means of solid media depends upon the
ability of the organism to reproduce at such a rate as to produce a
visible colony within the period of incubation and any substance that
prevents the growth of a visible colony is classified as a disinfectant;
if on further incubation the bacterial count approximates that of the
untreated sample the added substance has acted mainly as an antiseptic.
In practice no substance acts entirely as an antiseptic as the organisms
present have varying degrees of resistance and the less viable ones are
killed by doses that are only antiseptic to the more resistant ones.
An example of an antiseptic effect followed by a mild disinfectant
action, caused by small doses of bleach is shown in Table XV. In this
experiment the water designated as control was from the same source as
the treated water. In order to make the bacterial count in this water
approximately the same as in the treated water, the original count was
reduced by diluting the sample with water from the same source,
sterilised by boiling, and afterwards reaërated with sterile air.


Sample treated with 0.1 part per million of available chlorine.

               |                                       |    RATIO OF
               |                                       |     COUNTS.
    Time. |Day.|    2  |   3   |   4   |   5   |   6   | 2:4 | 2:5 | 2:6
          |    |       |       |       |       |       |Days.|Days.|Days.
   11 a.m.|  1 |    520|    940|  1,350|  2,360|  2,780|1:2.6|1:4.5|1:5.3
   12 noon|  1 |    390|    770|  1,080|  2,040|  2,320|  2.8|  5.2|  5.8
    2 p.m.|  1 |    187|    260|    690|  1,840|  2,080|  3.7|  9.9| 16.4
    4 p.m.|  1 |     91|    130|    280|    760|    840|  3.1|  8.3|  9.2
   10 a.m.|  2 |     42|    120|    670|    920|    ...| 15.9| 22. |  ...
   10 a.m.|  3 |    320|  1,210|  3,500|    ...|    ...| 10.9|  ...|  ...
   10 a.m.|  4 |  8,700| 14,200| 26,000|    ...|    ...|  2.9|  ...|  ...
                         CONTROL. NO CHLORINE ADDED
               |                                       |    RATIO OF
               |                                       |     COUNTS.
    Time. |Day.|    2  |   3   |   4   |   5   |   6   | 2:4 | 2:5 | 2:6
          |    |       |       |       |       |       |Days.|Days.|Days.
   11 a.m.|  1 |    121|    184|    285| liquid|    ...|1:2.4|1:...|  ...
   12 noon|  1 |    115|    171|    223|    380|    392|  1.9|1:3.2|1:3.2
    2 p.m.|  1 |    109|    152|    221|    362|    375|  2.0|  3.3|  3.4
    4 p.m.|  1 |    121|    175|    251|    410|    415|  2.1|  3.4|  3.4
   10 a.m.|  2 |  6,200|  8,500|  8,800|  8,900| liquid|  1.4|  1.4|  ...
   10 a.m.|  3 |425,000|650,000|670,000| liquid|    ...|  1.5|  ...|  ...
   11 a.m.|  1 |    915|  1,410|  1,630|  2,150|  3,200|1:2.2|1:2.8|1:3.5

  [A] Results are bacteria per c.cm.

Table XVI shows the effect of a concentration of 1.0 p.p.m. of chlorine;
the hypochlorite at this concentration acted almost entirely as a
germicide or disinfectant.



    Time.  |Day. |  2  |  3  |  4  |  5  |  6
   11 a.m. |  1  |    2|    5|    7|    8|   10
   12 noon |  1  |    1|    1|    2|    2|    4
    2 p.m. |  1  |    0|    0|    0|    2|    2
    4 p.m. |  1  |    1|    2|    2|    6|    6
   10 a.m. |  2  |    0|    0|    0|    1|   ..
   10 a.m. |  3  |    0|    0|    0|   ..|   ..
   10 a.m. |  4  |    5|   13|   16|   ..|   ..
   10 a.m. |  5  |   79|  166|   ..|   ..|   ..
  Untreated|     |     |     |     |     |
   water   | ..  |  915|1,410|1,680|2,150|3,200

  [A] Results are bacteria per c.cm.

Table XV shows a recovery of the anabolic functions after treatment with
0.1 p.p.m. of chlorine but since this was obtained by plating on such a
suitable medium as nutrient gelatine, it is probable that reproduction
in water having a low organic content would be still further diminished.
This is indicated by the results obtained.

There is no evidence of any marked difference in the resistance of
ordinary water bacteria to chlorine and these are the first to be
affected by the added germicide. The common intestinal organisms are
also very susceptible to destruction by chlorine and there is
considerable evidence that _B. Coli_ is slightly more susceptible than
many of the vegetative forms usually found in water. The specific
organisms causing the water-borne diseases, typhoid fever and cholera,
are, on the average, not more resistant than _B. coli_.

The spore-forming bacteria usually found in water are those of the
subtilis group, derived largely from soil washings, and _B. enteritidis
sporogenes_, from sewage and manure. The spores of these organisms are
very resistant and survive all ordinary concentrations. Wesbrook et
al.[1] found that 3 p.p.m. of available chlorine had little effect on a
spore-forming bacillus isolated from the Mississippi water and the
author has obtained similar results with _B. subtilis_.

Thomas,[2] during the chlorination of the Bethlehem, Pa., supply, found
four organisms that survived a concentration of 2 p.p.m. of available
chlorine: _Bact. ærophilum_, _B. cuticularis_, and _B. subtilis_, all
spore formers and _M. agilis_.

In practice no attempt is made, except in special cases, to destroy the
spore-bearing organisms as they have no sanitary significance and the
concentration of chlorine required for their destruction would cause
complaints as to tastes and odours if the excess of chlorine were not
removed. Such doses are unnecessary and result in waste of material. It
is found that, when the dose is sufficient to eliminate the _B. coli_
group from 25-50 c.cms. of water, the majority of the residual bacteria
are of the spore-bearing type. Smeeton[3] has investigated the bacteria
surviving in the Croton supply of New York City after treatment with 0.5
p.p.m. of available chlorine as bleach. Table XVII gives the results

The organisms of the _B. subtilis_ group outnumbered all the others, 66
(62.8 per cent) belonging to this group alone. This group contained _B.
subtilis_--Cohn (36 strains), _B. tumescens_--Chester (15 strains) _B.
ruminatus_--Chester (13 strains), and _B. simplex_--Chester 1904, (2
strains). Three of the four coccus forms were classified as _M. luteus_.
No intestinal forms were found.

Clark and De Gage[4] in 1910 directed attention to the fact that the
bacterial counts, made at 37° C. on chlorinated samples, were often much
greater than the counts obtained at room temperature. "This phenomenon
of reversed ratios between counts at the two temperatures," they stated,
"has been occasionally observed with natural water, but a study of the
record of many thousands of samples shows that the percentage of such
samples is very small, not over 3-5 per cent.... On the other hand 20-25
per cent. of samples treated with calcium hypochlorite show higher
counts at body temperature than at room temperature." Clark and De Gage
were unable to state the true significance of this phenomenon but were
of the opinion that it was not due to larger percentages of
spore-forming bacteria in the treated samples. Other observers, on the
contrary, have invariably found the spore-formers to be more resistant
to chlorine and thermophylic in type.



                 |  Morphology   | Spore       |   Gelatine  | Reaction in
                 |               | Formation   |Liquefaction | Litmus Milk
                 |Bacilli.|Cocci.| Pos. | Neg. | Pos. | Neg. | Pos. | Neg.
  No. of strains | 100    |  5   | 89   | 16   | 68   | 37   | 98   |  7
  Per cent.      |  95.2  |  4.7 | 84.7 | 15.2 | 64.7 | 35.2 | 93.3 |  6.6
                 |     Indol     |    Acid     |  Reduction  | Inhibition
                 |   Production  | Production  |     of      | by Gentian
                 |               | in Glucose  |  Nitrates   |   Violet
                 |  Pos.  | Neg. | Pos. | Neg. | Pos. | Neg. | Pos. | Neg.
  No. of strains |  75    | 30   | 61   | 44   | 40   | 65   | 98   |  7
  Per cent.      |  71.4  | 28.5 | 58   | 41.9 | 38   | 61.9 | 93.3 |  6.6

The removal of intestinal forms is, of course, merely a relative one and
when large quantities of treated water are tested their presence can be

The author, in 1915, made a number of experiments to ascertain whether
the _B. coli_ found after chlorination were more resistant to chlorine
than the original culture. The strains surviving treatment with
comparatively large doses were fished into lactose broth and subjected
to a second treatment, the process being repeated several times. The
velocity of the germicidal reaction with the strains varied somewhat,
but not always in the same direction, and the variations were not
greater than were found in control experiments on the original culture.
No evidence was obtained that the surviving strains were in any way more
resistant to chlorine than the original strain; in considering the
results it should be borne in mind that the surviving strains were
cultivated twice on media free from chlorine before again being
subjected to chlorination.

A number of the strains that survived several treatments were cultivated
in lactose broth and the acidity determined quantitatively. All the
cultures produced less acid than the original culture, and the average
was materially less than the original. These results point to a
diminution of the bio-chemical activity by action of the chlorine.

A point of perhaps more scientific interest than practical utility is
the relative proportion of the various types of _B. coli_ found before
and after treatment with chlorine. The author, in 1914, commenced the
differentiation of the types by means of dulcite and saccharose and
obtained the results shown in Table XVIII. These figures are calculated
from several hundreds of strains.

Although there is a slight difference in the relative proportions of the
types found at Ottawa and Baltimore, both sets of results show
definitely that there is no difference in the resistance of the various
types to chlorination.

=Aftergrowths.= In Tables XIII (p. 44) and XV (p. 51), it will be
noticed that, after the preliminary germicidal action has subsided, a
second phase occurs in which there is a rapid growth of organisms. This
is usually known as aftergrowth. When the contact period between
chlorination and consumption is short, the reaction does not proceed
beyond the first phase, but when the treated water is stored in service
reservoirs the second phase may ensue. At one purification plant, where
the service reservoirs are of large capacity, the aftergrowths amounted
to 20,000 bacteria per c.cm. although the water left the purification
plant with a bacterial count usually lower than 50 per c.cm.


                    |           PERCENTAGE OF ORGANISMS.
                    |  B. coli  | B. coli   | B. lactis | B. acidi
                    |  communis | communior | aerogenes | lactici
                    |     |Chlo-|     |Chlo-|     |Chlo-|     |Chlo-
                    |Raw. | ri- |Raw. | ri- |Raw. | ri- |Raw. | ri-
                    |     |nated|     |nated|     |nated|     |nated
  Ottawa, 1914      |  5  |  4  | 40  | 48  | 44  | 36  | 11  | 12
  Ottawa, 1915      |  8  |  8  | 50  | 46  | 34  | 31  |  8  | 15
  Baltimore, 1913[A]| 11  | 14  | 33  | 25  | 35  | 31  | 21  | 30

  [A] Thomas and Sandman.[5]

Regarding the nature of this aftergrowth, there has been a considerable
difference of opinion: some regard it as the result of the
multiplication of a resistant minority of practically all the species of
organisms present in the untreated water; others, that it is partially
due to the organisms being merely "slugged" or "doped," i.e. are in a
state of suspended animation, and afterwards resume their anabolic
functions; whilst others believe that with the correct dosage of
chlorine, only spore-forming organisms escape destruction and that the
aftergrowth is the result of these cells again becoming vegetative.

The aftergrowths obtained under the usual working conditions vary
according to the dosage of chlorine employed, and none of the above
hypotheses alone provides an adequate explanation. When the dosage is
small, a small number of active organisms, in addition to the spore
bearers, will escape destruction, and others will suffer a reduction of
reproductive capacity. The flora of the aftergrowth in this case will
only differ from the original flora by the elimination of a majority of
the organisms that are most susceptible to the action of chlorine and
the weaker members of other species of greater average resistance. As
the dose is increased these factors become relatively less important
until a stage is reached when only the most resistant cells, the spores,
remain. The resultant aftergrowth must necessarily be almost entirely
composed of spore-bearing organisms. A small number of the most
resistant members of non-sporulating organisms may also be present but
they will, in the majority of instances, form a very small minority.
This is the condition that usually obtains in practice and it is
necessary to consider whether the aftergrowth may have any sanitary

Concerning the secondary development of _B. coli_, the usual index of
pollution, there is but little information. H. E. Jordon[6] reported
that, of 201 samples, 21 gave a positive _B. coli_ reaction immediately
after treatment, 39 after standing for twenty-four hours, and 42 after
forty-eight hours. These increases were confined to the warm months, the
cold months actually showing a decrease. The following figures, taken
from the author's routine tests for 1913 and 1914, show a similar
tendency, but an analysis of the results by months did not show that
this was confined to the warm season. The sequence of the results from
left to right, in the following Table, is in the same order as the
contact period. Approximately 290 samples were taken at each sampling

At station No. 2 the germicidal action was still proceeding but at No.
5, representing an outlying section of the city, the increase in the _B.
coli_ content is very apparent.

During 1915 and 1916 the author endeavoured to duplicate these results
under laboratory conditions and entirely failed. These experiments,
which were made with the same materials as were in use at the city
chlorination plant, but in glass containers, were usually only carried
to a forty-eight hours contact, as this was the extreme limit for the
city mains; one, however, was prolonged to five days. Many experiments
were made under varying conditions, with similar results. Typical
examples are given in Tables VI, VIII and IX on pages 33 and 37.



       |        SAMPLING POINT NO.
       |   1  |   2  |   3  |  4   |  5
  1913 | 15.2 | 14.4 | 16.3 | 16.8 | 26.8
  1914 |  7.0 |  5.7 |  6.0 | .... | 11.6

In every case there was persistent diminution in the number of _B. coli_
with increase of contact period. Determination of the bacterial count on
nutrient agar showed that, in several experiments, the aftergrowth had
commenced, and in some instances there was evidence that the second
cycle was partially complete i.e. the number had reached a maximum and
then commenced to decline. The time required for the completion of the
two cycles, comprising the first reduction caused by the chlorine, the
increase or aftergrowth, and the final reduction due to lack of suitable
food material, is dependent upon several factors of which the dosage and
temperature are the most important. With a small dosage the germicidal
period is short and the second phase is quickly reached; with large
doses, the second phase is not reached in forty-eight hours; the higher
the temperature the quicker is the action and the development of the
aftergrowth. These statements refer only to the bacteria capable of
development on nutrient agar. The _B. coli_ group behaved differently
and persistently diminished in every case. If _B. typhosus_ acts in a
similar manner to _B. coli_, the laboratory experiments show that
aftergrowths are of no sanitary significance and can safely be ignored,
but as the results obtained in practice are contradictory to the
laboratory ones, the matter must be regarded as _sub judice_ until more
definite evidence is available.

It is common knowledge that samples of water from "dead ends" of
distribution mains show high counts and much larger quantities of _B.
coli_ than the water delivered to the mains. This is another phase of
aftergrowth problem that often causes complaints and can only be
eliminated by "blowing off" the mains frequently or by providing
circulation by connecting up the "dead ends." One extreme case of this
description might be cited. A small service was taken off the main at
the extreme edge of the city to supply a Musketry School two miles away
and was only in use for a few months in the summer season. This service
pipe delivered water containing _B. coli_ in a considerable percentage
of the 10 c.cm. samples and in a few instances in 1 c.cm., although the
water delivered to the city mains never exceeded 2 _B. coli_ per 100
c.cms. and averaged about one-tenth that quantity. No epidemiological
records of the effect of this water are available because it was put
through a Forbes steriliser before consumption.

In some instances the rate of development of the organisms after
chlorination is greater than in the same water stored under similar
conditions. This is especially noticeable in the presence of organic
matter and has been ascribed to the action of the chlorine on the
organic matter with the production of other compounds that are available
as food material for the organisms.

Houston, during the treatment of prefiltered water Lincoln in 1905,
found that although the removal of _B. coli_ and other organisms growing
at 37° C. was satisfactory, there was almost invariably an increase in
the bacteria growing on gelatine at 20° C. This was ascribed to the
action mentioned above and the chemical results supported this view,
more organic matter being found in the filter effluents than in the
prefiltered water. Rideal's experiments with sewage at Guildford
indicate that a similar action may occur in contact beds. The addition
of bleach to the prefiltered water at Yonkers also resulted in an
increased count and in these instances the aftergrowths are due to a
disturbance of the equilibrium by the action of the chlorine on the
zooglea and other organic matter invariably found in ripe filters.
Similar results can be produced by the addition of chlorinated water to
small experimental sand filters. This is shown by the results in Tables
XX and XXI.


                |  Bacteria Per |               |
    Available   |    Gram of    |Typical B. coli| FREE CHLORINE
   Chlorine in  |   Sand After  |After 24 Hours.| AFTER 24 HRS.
   Water p.p.m. +-------+-------+---+---+---+---+-------+--------
                |       |       |100| 10| 1 |0.1|Without| After
                | 3 Hrs.|24 Hrs.|Gr.|Gr.|Gr.|Gr.| Acidification.
   Nil          | 12,000| 21,000| + | + | + | - |   -   |   -
    3.0         |     80|114,000| - | - | - | - |   -   |   -
    5.0         |     50|150,000| - | - | - | - |   -   |   -
    7.0         |     25|214,000| - | - | - | - |   -   |   -
   10.0         |     26|500,000| - | - | - | - |   -   |   -


   Available in +----------+----------+-----------
   Water p.p.m. | 3 Hours. |24 Hours. |48 Hours.
   Nil          |  70,000  |   .....  |   .....
   0.1          |   7,200  |  20,400  |  12,800
   0.3          |   5,240  |   6,400  |  11,200
   0.5          |   5,120  |   4,700  |  10,800
   1.0          |   1,100  |   8,800  |  20,400

It is observable that the effect of small doses was comparatively small
and transient; large doses of bleach reduced the bacteria very
materially but the reduction was not maintained and the subsequent
increase was abnormally rapid.


[1] Wesbrook, Whittaker and Mohler. J. Amer. Pub. Health Assoc., 1911,
1, 123.

[2] Thomas. Jour. Ind. and Eng. Chem., 1914, 6, 548.

[3] Smeeton. Jour. of Bact., 1917, 2, 358.

[4] Clark and De Gage. Rpt. Mass. B. of H., 1910, p. 319.

[5] Thomas and Sandman. J. Ind. and Eng. Chem., 1914, 6, 638.

[6] Jordan, H. E. Eng. Record, 1915, May 17.



The complaints that have been made against chlorinated water since the
practice was commenced have been very diversified in character and can
be numbered by the legion and although some have been justifiable, the
great majority has been unsubstantiated and must be ascribed to

Almost every one who has had charge of chlorination plants has noted the
latter phenomenon, for in some instances complaints have been made
following the publication of the information that chlorination was to be
commenced but antecedent to its actual operation, and in others when for
some reason or another, the chlorination plant has been temporarily
stopped. Similar observations have been made in laboratory experiments
when independent observers have been requested to detect the chlorinated
waters from an equal number of treated and untreated waters. Such
observers are wrong in the majority of the waters which they designate
as treated ones if the dosage is not in excess of that required for
satisfactory purification.

One amusing example of auto-suggestion was experienced by the author
some years ago. During a ceremonial visit to the waterworks, the Mayor
and several civic representatives happened to visit a hypochlorite plant
that was built on a pier over the river and which had no ostensible
connection with the city mains. One of the party expressed a desire for
a drink of good river water without any hypochlorite in it and was
served with water from the plant supply by an assistant engineer of the
waterworks department. The water was consumed by all with great relish
and as it was being finished, the writer entered the plant and was
invited to join them in the enjoyment of this "dopeless" water; on
asking where it had been obtained he was astonished to hear that it was
from a tap which was supplied with the ordinary chlorinated water of the

On many occasions, complaints are justifiable and should be carefully
investigated instead of, as is often the case, being attributed to
auto-suggestion. The time and energy that are often devoted to
endeavouring to persuade water consumers that their complaints are
without foundation, can better be utilised in so improving the
chlorination process as to eliminate tastes and odours. All complaints
should be carefully investigated and a record kept for future reference,
for the cause, although not manifest at the time, may be discovered
later. The records then provide valuable corroborative evidence.

The nature of the complaints against chlorinated water is very
diversified and includes imparting foreign tastes and odours, causing
colic, killing fish and birds, the extraction of abnormal amounts of
tannin from tea, the destruction of plants and flowers, the corrosion of
water pipes, and that horses and other animals refuse to drink it.

_Tastes and Odours._ When an excess of hypochlorite or liquid chlorine
is added to a water it imparts a sharp pungent odour and acid taste,
characteristic of chlorine, that render it offensive to the nose and
palate. In some instances the presence of chlorine compounds is not
obtrusive when the temperature of the water is low but becomes so when
the temperature is raised. It is especially observable when the faucets
of hot water services are first opened and the chlorine is carried off
as a vapour by the other gases liberated by the reduction in pressure.
For this reason the complaints regarding hot water are relatively more
numerous and sometimes constitute the whole of the complaints. In cold
water containing appreciable quantities of mineral salts the
hypochlorites and hypochlorous acid might not be entirely dissociated;
they may become more hydrolysed with an increase in temperature and
finally broken down under the influence of the carbonic acid liberated
from the bicarbonates by heat.

Chlorine also forms chlorinated organic compounds by action on the
organic matter present in water and some of the objectionable tastes and
odours of chlorinated waters have been attributed to this agency. Some
observers have stated that chloramines were amongst the chloro-organo
compounds produced but the author's experience with the Ottawa supply
has demonstrated that simple chloramine (NH_{2}Cl) can be successfully
employed for water treatment without causing complaints. It was
suggested on page 28 that some of the higher chloro-amines might be the
cause of some complaints but at present there is no definite information
regarding the formation of these compounds in water and all such
hypotheses are little more than conjectures. Letton[1] has reported that
at Trenton, in 1911, when the water of the Delaware River was first
treated, the dosage was as high as 1.2 p.p.m. of available chlorine and
although chemical tests showed the absence of free chlorine, the water
had an extremely disagreeable taste which was especially noticeable in
the hot water. The conclusion was reached that "the taste and odour were
not those of chlorine, but were due to some complex chemical change
brought about by the action of the chlorine on the organic matter
present in the water."

The waters that require the most accurate adjustment of chlorine dosage,
if complaints are to be avoided, are those containing very small amounts
of organic matter. The margin between the dosage required for the
attainment of a satisfactory degree of bacteriological purity and that
which may cause complaints is usually very small, often less than 25
per cent, with the waters of the Great Lakes and many filter effluents.
On the other hand, coloured waters containing large amounts of organic
matter can be treated with an excess of chlorine without causing tastes
and odours. The author found that the addition of 1.5 p.p.m. of
available chlorine to the Ottawa River water did not cause complaints
although only 0.8 to 0.9 p.p.m. were usually required for satisfactory
purification. Harrington of Montreal has had a similar experience with
this water.

The presence of traces of foreign substances in water sometimes produces
chlorinated derivatives having repugnant tastes and odours. Creosote and
tar oils have caused an odour somewhat resembling that of iodoform and
industrial wastes have also produced complaints.

The substitution of chlorine gas (liquid chlorine) for bleach solutions
has apparently eliminated tastes and odours in some cases but this may
be due to a more perfect control over the dosage rather than to any
property of the bleach _per se_.

In some instances the sludge from bleach plants has caused complaints by
producing an excessive concentration of chlorine during the period of
its discharge. This occurred in Ottawa on several occasions before it
was discovered and corrected. When the sludge in the storage tanks
reached the discharge valve it was customary to wash out the tank and
discharge the sludge into the river. The operators opened the wash out
valves to the full extent and the sludge and liquor were discharged into
the river about 70 feet away from the inlet to the sedimentation basin
and on the downstream side of it. A portion of the hypochlorite was
almost invariably carried into the basin and increased the dosage. This
condition was remedied by carrying the sludge drain farther down stream
and insisting upon the sludge being discharged at a slower rate.

Kienle[2] has reported similar occurrences at Chicago. The hypochlorite
was applied at the intake cribs situated a considerable distance off
shore. The direction of the wind often necessitated holding the sludge
for a considerable length of time but occasionally it was found
impossible to await favourable conditions with the result that the wind
and wave action carried a portion of the sludge back into the crib and
down into the shaft and tunnel.

The temperature of the water at the time of treatment is another factor
bearing on the production of tastes and odours. When the temperature is
low, water absorbs relatively less chlorine (_vide_ Diagram No. II, page
38) in the same period of time with the consequence that, if the dosage
is kept constant, more chlorine is present in the free condition. At
Milwaukee (Kienle)[2] with a dosage of 0.24 p.p.m. of available chlorine
(as bleach) no complaints were received during the spring, summer, and
autumn seasons but when the temperature reached 40° F., they were
compelled to reduce the chlorine to 0.12 p.p.m. in order to prevent
objectionable tastes and odours in the tap waters.

Abnormal conditions such as freshets, and storms, sometimes cause
complaints regarding tastes and odours. Adams[3] found that the
complaints in Toronto usually accompanied a change in the direction of
the wind, a sustained east wind being the one most productive of
trouble. The exact cause for this could not be ascertained but it was
usually found that there was an accompanying increase in the number of
microscopical organisms (plankton) present in the raw water.

Freshets usually increase the bacterial contamination and necessitate an
increased dosage which may cause complaints.

Complaints as to tastes and odours can be best avoided by ensuring
regularity of dosage, perfect admixture, and storage of the treated
water for a reasonable period. These factors are discussed in detail

_Colic._ Although claims have been made that the consumption of
chlorinated water has produced "colic" no corroborative evidence has
been adduced and the symptoms have probably been due to some other
cause. Dilute solutions of chlorine have been used as intestinal
antiseptics in the treatment of typhoid fever without producing
irritation of the mucous lining and the usual dose for this treatment is
one grain of chlorine. Before taking a _medicinal_ dose of chlorine 140
gallons of water containing 0.1 p.p.m. would have to be consumed, a
quantity greater than is ordinarily drunk in a year.

Chlorine and hypochlorites are destructive and irritant to skin and it
is possible that hot chlorinated water has, in some instances, a similar

It is inconceivable that the addition of minute traces of bleach or
chlorine to water should cause it to extract abnormal amounts of tannin
from tea but it is possible that free chlorine, when present, acts upon
the tea extractives and produces compounds having obnoxious tastes and
odours. Tannin to the ordinary tea drinker represents the disagreeable
portion of the tea and an obnoxious taste in tea brewed with chlorinated
water would consequently be ascribed to the extraction of abnormal
quantities of tannin.

Almost all waterworks departments using chlorination have received
complaints to the effect that the water had killed fish and small birds.
There is usually no evidence that the loss was due to chlorinated water
but it is generally impossible to convince the owners that the process
of water treatment was not the cause. Many continuous physiological
tests have been made of the effect of chlorinated water on small fish
and have shown that the concentration used in water treatment is without
effect. The author kept a tank of minnows in one of the pumping stations
for months without loss although the tank was continuously supplied with
water that had been treated but a few seconds previously. The bleach
solution was discharged into the suction of the pumps and the water for
the fish test was taken from the discharge header.

It has been found on many occasions that fish are extremely susceptible
to chlorine and hypochlorites. This knowledge has been sometimes used
for such nefarious purposes as fish poaching, a few pounds of bleach in
a small stream being a simple and most effective method of killing all
the fish which are then carried down stream into a convenient net.
Chlorinated sewage effluents have also been known to destroy the fish
life of the stream into which they were discharged.

The opinion of fish culturists as to the action of chlorinated waters
upon fish eggs in hatcheries is almost unanimously to the effect that it
is a destructive one. Fish eggs are extremely sensitive to chlorine and
hypochlorous acid and very few will survive in a water containing 0.1
p.p.m. of free chlorine. The Department of Fisheries of the Dominion of
Canada has informed the author that free chlorine in the water had a
marked adverse effect on the hatching of the eggs of Atlantic salmon,
Great Lake trout, pickerel, and whitefish, but no effect was noticed
when free chlorine was absent. The Department has, however, decided to
remove all the hatcheries to localities where water that does not
require chlorination can be obtained.

The effect of chlorinated water upon seeds, plants, and flowers has been
investigated by the Dominion Department of Agriculture and Dr. Gussow
(Dominion Botanist) and Dr. Shutt (Agricultural Chemist) who were in
charge of the work, have reported that water treated with hypochlorite
caused no apparent injury to carnations and hybrid roses. Six varieties
of wheat seed, after soaking in freshly prepared hypochlorite solutions
(0.05 to 10 parts per million of available chlorine) were all sown on
the same day. Germination was found to be uniform throughout and no
effect of the chlorine was observed either as regards the rate of
germination or the development of the young plants. Experiments on
barley and oats produced similar results. Radishes, turnips, cucumbers,
and beans also showed no retardation in development after treatment with
chlorinated water.

These experiments were conducted with solutions of bleach in distilled
water, but identical results were obtained in a later series when the
treated city supply (Ottawa) was used.

The results proved conclusively that statements alleging damage to
plants, flowers, and seeds by the hypochlorite treatment of water are
unfounded and do not merit the slightest consideration.

_Corrosion of Pipes._ Chlorinated water, it has been alleged on many
occasions, causes rapid corrosion of galvanised iron water services and
especially of the water tubes of boilers, water heaters, etc. When
bleach is used for water treatment, a slight increase in the hardness is
produced but as this is mostly due to calcium chloride, there is no
corresponding increase in the salts that form a protective coating. The
presence of traces of calcium chloride and chloro-organic compounds
might tend to increase the corrosive properties of a water but this
increase is probably so small as to be negligible.

If pipe corrosion is considered by the carbonic acid hypothesis, the use
of bleach should tend to reduce it because bleach contains an excess of
base that combines with a portion of the free carbonic acid. The results
of routine tests for free carbonic acid made on the raw and treated
waters at Ottawa are as follows:

         |   CARBONIC ACID.     |
         |  PARTS PER MILLION   |
   Year. +----------+-----------+   Nature of Treatment.
         |Raw Water.|Chlorinated|
         |          |   Water.  |
   1915  |   1.44   |   1.41    | Bleach
   1916  |   0.92   |   0.85    | Bleach
   1917  |   0.84   |   0.81    | Bleach first four months
         |          |           | Chloramine during last
         |          |           | eight months

These figures shown that the hypochlorite treatment produced a small but
definite decrease in the carbonic acid content and should, _cæteris
paribus_, tend to reduce and not increase corrosion.

If the corrosion of pipes is considered according to the electrolytic
theory, a slight increase, due to an increased electrical conductivity,
might be anticipated. The effect of the addition of hypochlorite upon
the electrical conductivity of distilled water and the Ottawa River
water is shown in Diagram VI.

[Illustration: DIAGRAM VI

Effect of Calcium Hypochlorite on Electrical Conductivity]

With the concentrations of hypochlorite ordinarily used in water
treatment it is inconceivable that the slight increase in the electrical
conductivity has any practical significance at low temperatures. The
conductivity increases rapidly, however, with increase of temperature
and any increment due to chlorination might produce a slight appreciable
effect at temperatures approaching the boiling-point of water.

Liquid chlorine does not increase the conductivity to the same extent as
an equivalent quantity of hypochlorite but it increases the carbonic
acid content in proportion to the dosage used.

The author investigated the action of hypochlorite on galvanised pipes
in 1914 and was unable to detect any definite corrosion with normal
concentrations of chlorine. The experiments were made with 2-inch pipes
and an examination of the first consignment received showed that,
although the galvanising on the outside was perfect, the inner coat was
very inferior: in some parts there was an excess of zinc that broke away
on scraping whilst in others the iron pipe was bare.

A committee of the Pittsburg Board of Trade, appointed to investigate
complaints as to pipe corrosion, reported in 1917 that they were largely
due to inferior qualities of pipes and not to the method of water
purification employed (slow sand filtration and chlorination).

The effect of chlorination on the _plumbo-solvency_ of water was
investigated in 1904 by Houston who found that chlorine, as chloros, in
amounts between one and ten parts per million, did not appreciably
increase the plumbo-solvent action of either unfiltered or filtered
water. Similar results were obtained by the author with the Toronto
supply: raw lake water, filtered water, and water treated with 0.25 and
0.50 p.p.m. of chlorine, all dissolved the same quantity of lead in
twenty-four hours. The amount in each case was too small to be of any


[1] Letton. J. Amer. Waterworks Assoc., 1915, 2, 688.

[2] Kienle. J. Amer. Waterworks Assoc., 1915, 2, 690.

[3] Adams. J. Amer. Pub. Health Assoc., 1916, 6, 867.



The treatment of water with bleach alone has been largely supplanted by
the liquid chlorine process but the following details will be of use on
meeting conditions for which liquid chlorine cannot be used and also for
the preparation of the hypochlorite solution required in the chloramine

The essential features of a bleach installation are the solution or
mixing tanks, storage tanks, piping system, discharge orifice or weir,
and sludge drain.

Bleach is usually sent out by the manufacturers in sheet steel drums, 39
inches high and 29-1/2 inches in diameter, which contain about 14 cu.
ft. of bleach and weigh approximately 750 pounds gross and 690 pounds
net. It can be most economically purchased in car lots and if the
consumption warrants this procedure storage should be provided for about
70 drums or rather more than one car load. According to Hooker[1] bleach
loses 1 per cent of available chlorine per month in hot seasons and 0.3
per cent in cold ones so that it is advisable to carry as little stock
as possible during hot weather. Hot weather also causes a further loss
by accelerating the action of the bleach on the drum which rapidly
disintegrates and cannot be handled. Bleach can often be purchased more
cheaply in hot weather but such a policy is a short sighted one unless
it is required for immediate use.

The general design of a hypochlorite plant is largely determined by the
capacity but in all cases an effort should be made to avoid complicated
details which may appear advantageous in the drafting office but do not
stand up in actual practice. Many metals rapidly develop a protective
coating on immersion in bleach solution but if this is removed by
friction, rapid erosion ensues; bearing metallic surfaces should be
reduced to a minimum.

_Mixing Tanks._ All tanks, whether mixing or storage, should be
constructed of concrete and painted with two coats of asphalt.
Experience has shown that wooden tanks are not suitable. The author has
used pine, oak, and cypress tanks but all were rapidly leached by the
hypochlorite and ultimately had to be lined with concrete.

There is a considerable variation in the concentration of bleach
solution made in mixing tanks at various works. Some operators use about
one gallon of water per pound of bleach and mix the two to a cream by
wooden paddles, revolving on a central axis, for 1-2 hours; the paddles
are then stopped and the cream run out into the storage tanks and
diluted to the required strength by passing water through the mixing
tank. There are two objections to this method: (1) the addition of small
quantities of water to bleach tends to gelatinisation which may protect
lumps from the further action of water and (2) a stratification of the
solution occurs in the storage tank unless agitation is used.
Gelatinisation causes loss of available chlorine and stratification
causes irregular dosage unless corrected by agitation, which
necessitates power. Other operators mix the bleach and water to the
final concentration in the mixing tank and discharge the contents into
the storage tank, the intermittent process being repeated until the
storage tank is full. Gelatinisation is avoided by using a low original
concentration and as all batches are of equal density no stratification
is produced.

At Ottawa the bleach is crushed and, after weighing, dumped into a
circular concrete tank provided with a hinged wooden lid. The stirring
arrangement consists of a bronze shaft on which an aluminium impeller is
fixed which revolves in an iron tube set slightly above the bottom of
the tank (see Fig. 1). After the requisite amount of water has been
added the motor connected to the bronze shaft is started and the mixture
pumped for 15-20 minutes; without waiting for the sludge to settle the
contents are discharged into the storage tank and the operation repeated
until the tank is full. The piping between the mixing and storage tanks
is of galvanised iron of generous dimension so as to compensate for
incrustation. The pipes are straight and are provided with crosses at
every change of direction to enable excessive incrustation to be
removed. The valves should be made of hard rubber or special bronze; if
brass valves are used they will probably require renewing every twelve

[Illustration: Fig. 1.--Mixing Tank for Bleach.]

The concentration of solution necessarily depends upon local conditions
but it is usually advisable to keep it below 2.5 per cent of bleach,
which is equivalent to 0.85 per cent of available chlorine.

_Storage Tanks._ These should be built of reinforced concrete and
painted inside with asphalt, which should be periodically renewed to
prevent the solution seeping through to the reinforcement. At least two
tanks should be provided so that one may be filled and allowed to settle
before being put in operation. The hypochlorite discharge pipe is
usually 6-9 inches from the bottom to permit the collection of sludge,
which is run off when it reaches the elevation of the hypochlorite
discharge. The sludge drain, which opens into the bottom of the tank, is
usually a 4- or 6-inch cast-iron pipe, with suitable gate valve, which
discharges into a common drain made of clay pipe.

The storage tanks should be provided with either glass gauges or float
indicators to enable the orifice discharge to be checked up at
periodical intervals.

_Regulation of Dosage._ The discharge of the hypochlorite solution is
usually regulated either by maintaining a constant head on an orifice of
variable dimension or by varying the head on an orifice of fixed
dimension. The weir principle may also be used but it is not so well
adapted for hypochlorite as for other chemicals.

In the constant head method, the head is maintained by a bronze valve
connected to a float made of glass or tinned copper. In many cases the
orifice is a rectangular slot in a brass plate and is adjusted by means
of a brass slide operated by a micrometer screw. Brass plates are not
very suitable as they become corroded and so reduce the size of the
orifice; if the incrustation is removed the orifice will discharge more
than the calibration indicates. Needle valves are unsuitable for similar

An example of an orifice feed box of the constant head type is shown in
Fig. 2. A vertically arranged hard-rubber pipe passes though a hard
rubber stuffing box in the bottom of the tank and has one or more
orifices near its upper end. The area of the submerged portions of the
orifices is controlled by the hand wheel which is connected with the
threaded stem of the pipe. The stem has sixteen threads per inch, and
one revolution of the wheel will submerge the orifices one-sixteenth of
an inch. The extent to which the orifices are submerged is indicated on
the dial fixed to the side of the tank.

[Illustration: FIG. 2.--Dosage Tank.]

Fig. 3 shows the regulating mechanism of another apparatus of the
constant head type. The orifice consists of a circular slot in a hard
rubber disc and is regulated by means of a hand wheel which operates a
hard rubber slide.

[Illustration: FIG. 3.--Orifice Controlling Device.]

The general arrangement of one of the variable head types is shown in
Fig. 4. A constant head is maintained on the valve _V_ by a float and
cock operating in a lead- or porcelain-lined tank. The circular tapered
orifice _O_, cut in glass, is situated in the flanged end of the iron
casting _C_ and the head, indicated on the gauge glass, is regulated by
valve _V_. This arrangement is simple and reasonably accurate. The
orifice may show slight incrustation after being in service for some
time but it can be easily cleaned by means of a test-tube brush or a
small swab moistened with acid; a wire or rod tends to break the edge of
the conical orifice and should not be used.

The volume of solution discharged by orifices of various dimensions is
shown in Diagram XV, page 149. Diagram XVI, page 149, facilitates the
calculation of the number of pounds of bleach required for any dosage.

[Illustration: FIG. 4.--Variable Head Dosage Box.]

The solution discharged from the orifice box is carried to the point of
application either in galvanised iron pipes of generous dimension or in
rubber hose. Pumps may be used for raising the solution to a higher
elevation but unless special material is used in their construction they
corrode rapidly and cannot be kept in service. Whenever possible, a
water injector should be used as it does not corrode and assists in
maintaining the delivery pipes free from sludge. All delivery pipes
should be duplicated and blown out regularly by water under pressure;
they should also be protected from frost.

The adjustment of the hypochlorite dosage can be automatically regulated
in plants where the flow of the water to be treated is measured by a
Venturi meter or other suitable appliance. Various devices have been
suggested and used but, in general, they are not so successful as
automatic regulators for liquid chlorine on account of the presence of
sludge particles which tend to diminish the area of the orifice.

For small plants, barrels have often been used as solution and storage
vessels with, in some instances, fairly successful results. The bleach
process, however, cannot be recommended for small installations because
the chemical control necessary for successful operation is usually not
available. One drum of bleach may suffice for several months operation
and as the powder gradually loses strength, the dosage constantly
diminishes and may jeopardise the safety of the supply. Liquid chlorine
machines are much more suitable than hypochlorite installations for
supplies having no chemical control.

Bleach is being very extensively used for the sterilisation of the water
used by the allied troops in France. The water supplies on the British
front are all more or less subject to pollution and it is consequently
necessary, to ensure adequate protection, to chlorinate all supplies
with bleach. Other forms of chlorine have been tried but have not proved
successful near the firing lines. The details of the technique employed
cannot be given but it may be stated that the concentration of chlorine
employed is always more than sufficient and that residual tastes and
odours are regarded as secondary considerations. Treated water is always
tested by the starch-iodide method and a bacteriological examination is
frequently made by mobile laboratories.

=Control of Hypochlorite Plants.= If efficient operation and regular
dosage is to be obtained, it is necessary that hypochlorite plants
should be controlled by a trained chemist. Good results are occasionally
obtained without such control but in every plant circumstances arise at
some period or another which only a chemist is qualified to deal with.

The points that require consideration are (1) the composition of the
bleach; (2) concentration of available chlorine in the prepared
solutions; and (3) chemical tests for free chlorine in the treated

(1) _Composition of Bleach._ Each drum of bleach should be sampled and
analysed before use. The sample is obtained by cutting out the head of
the drum and removing a vertical section by means of a special sampling
tube or a piece of half-inch iron pipe which is forced to the bottom of
the drum with a boring motion and then removed; the core is then forced
out by means of a rod, mixed, and quartered down to the required size.

For analysis weigh out 5 grms. on a balance sensitive to 0.01 grm. and
grind in a mortar with 50-70 c.cms. of water; wash into a 250 c.cm.
flask and make the volume up to 250 c.cms.; shake. After allowing the
sludge to settle remove 10 c.cms. by means of a pipette and titrate by
one of the following methods:

_Bunsen's Method._ Add 10 c.cms. of a 5 per cent solution of potassium
iodide and 0.5 c.cm. glacial acetic acid and titrate with sodium
thiosulphate (24.8 grms. of the C.P. crystalline salt and 1 c.cm. of
chloroform per litre) using a starch solution as indicator. Each cubic
centimetre of thiosulphate used = 1.755 per cent of available chlorine
(1 c.cm. N/10 sodium thiosulphate = 0.00355 grm. available chlorine).

_Penot's Method._ Dilute the hypochlorite solution with 15 c.cms. of
water and titrate with a solution of N/10 sodium arsenite using
starch-iodide paper as an external indicator. Each c.cm. of solution
used = 1.755 per cent of available chlorine (1 c.cm. = 0.00355 grm.
available chlorine). The use of an external indicator makes this process
a slow one and to overcome this objection Mohr proposed the addition of
an excess of sodium arsenite solution and then titrating with N/10
iodine solution after adding a few drops of starch solution.

Griffen and Hedallen[2] compared these three methods and found that
Penot's method and Mohr's modification of that method gave results
which were 0.6 per cent lower than those obtained by Bunsen's method.

For a separate estimation of the chlorine present as chloride, chlorate,
and hypochlorite the method given in Sutton's Volumetric Analysis, 10th
edition, page 178, should be followed.

_Storage Liquor._ This is tested by any of the above methods. It has
been proposed to determine the strength of the bleach solution by the
use of a hydrometer but the results are not sufficiently accurate and
the method cannot be recommended.

If bleach is properly broken up and thoroughly agitated in the mixing
tank at least 95 per cent of the available chlorine should be extracted.
The efficiency of the extraction process is checked by comparing the
tests of the storage liquor with those of the dry bleach and each batch
of liquor should be tested daily. It is sometimes advisable to take two
samples from each tank, one soon after a tank has been put into
operation, and a second sample at the end of the run. Considerable
differences are occasionally found between these samples and are due,
either to inadequate agitation of the liquor in the storage tank, or
inefficient mixing in the mixing tank. If the results are irregular the
former is the more probable cause but if the second sample is invariably
stronger the mixing tank operations should be investigated. The
increased concentration of the second sample is due to unextracted
bleach passing out of the mixing tank and gradually becoming leached as
the tank contents are run off. If the bleach is lumpy and is not
subsequently broken up, losses are almost inevitable.

Hale[3] found that during the period when the New York City supply was
being treated with bleach it was necessary to constantly check the
operations of the labourers by frequent samples. "During one week about
95 per cent of the chlorine added was actually applied, the second week
it dropped to 85 per cent. and the third week to 75 per cent. Whenever
a poor run is called to the attention of the labourers, results

By taking two samples daily from each tank discharged the author has
been able to obtain an average annual efficiency on the Ottawa plant of
94 per cent., i.e. the solutions contained 94 per cent. of the available
chlorine contained in the bleach. In making such checks it is necessary
to keep a careful account of the stock of bleach to prevent labourers
adding a few extra pounds of bleach to compensate for losses.

Sludge forms an appreciable but unavoidable source of loss of material.
When the sludge reaches the outlet of the hypochlorite pipe the sludge
must be run to waste; otherwise it will pass over and tend to choke the
dosage control apparatus. If the sludge is run into the same body of
water that forms the source of supply, it must be discharged very slowly
to prevent a possibility of over dosage and damage to fish life. With
proper control, sludge losses can easily be kept under 2 per cent. and
often under 1 per cent.

The greatest source of unavoidable loss in hypochlorite plants is from
deterioration of the bleach during storage; in warm climates this loss
may exceed 10 per cent. In Ottawa where high temperatures are only
experienced during the summer months the loss from this cause has
averaged from 7-8 per cent. on the bleach stored during that period.

_Detection and Estimation of Free Chlorine._ The oldest and probably the
best known test for free chlorine in water is the Wagner test, made by
adding a few drops of potassium iodide and starch; the presence of
chlorine is indicated by a deep rich blue colouration that is
proportional in intensity to the quantity of chlorine present. When this
test is used as a colorimetric method for the estimation of chlorine
several difficulties are encountered; the intensity of the colour
produced by the majority of treated waters gradually diminishes and the
loss is usually more rapid than in the standards made up with distilled
water; a different result is obtained if the solutions are acidified and
the results vary with different acids, acetic acid yielding a much lower
result than a mineral acid such as hydrochloric acid; in the presence of
acid the colouration usually intensifies on standing, whereas the
standard intensifies but little. The difference caused by the addition
of acid is imperfectly understood but it is obvious that the chlorine
set free by the acid cannot be present in the "free" state; it is
probably in a semi-labile condition loosely attached to organic
compounds. Whether this semi-labile chlorine is available for germicidal
action is at present not definitely known but it has been noted by
several observers that the germicidal action proceeds after the "free"
chlorine reaction has disappeared.

The method used by the author for the estimation of free chlorine is as
follows: place 500 c.cms. of the sample in a stoppered bottle, add 1
c.cm. of 5 per cent KI solution, 2 drops of conc. HCl and 1 c.cm. of
starch solution and titrate with N/1000 sodium thiosulphate until
colourless. The difficulty introduced by the opalescence of the liquid
is overcome by pouring portions of the liquid into two Nessler tubes and
adding a drop of thiosulphate solution to one and noting if any
reduction of colour occurs on shaking; if the intensity of the colour is
diminished, the contents of both tubes are poured back into the bottle
and titrated until no further colour removal, as shown by the tubes, can
be obtained. One c.cm. of N/1000 sodium thiosulphate = 0.07 p.p.m. of
available chlorine when 500 c.cms. of water are used.

Adams[4] has employed the colorimetric method of estimating the colour
obtained after the addition of dilute H_{2}SO_{4}, KI, and starch but
used standard solutions of dyes for comparison. The standards were
prepared from mixtures of Brilliant Mill Green "S" and Cardinal Red "J"
and were made up weekly.

Phelps found that ortho-tolidine in acetic acid solution produced an
intense yellow colouration with free chlorine and suggested the use of
this reagent as a qualitative test for chlorine. Ellms and Hauser[5]
developed this process into a quantitative one and substituted
hydrochloric acid for acetic acid as a solvent. One c.cm. of the reagent
(1 gram of pure _o_-tolidine dissolved in 1 litre of 10 per cent of
hydrochloric acid) is added to 100 c.cms. of the sample in a Nessler
tube and the colour compared after five minutes with permanent standards
made up with mixtures of potassium bichromate and copper sulphate. This
method was adopted as the official standard method of the American
Public Health Association; the details are given in the Appendix (p.

The author has found that this method gives excellent results except for
coloured waters. The colouring matter in many waters diminishes in
intensity on the addition of acids and is somewhat similar in tint to
that produced by addition of _o_-tolidine. If the reaction is used
qualitatively on coloured treated water and a comparison made with the
untreated sample, a negative result, due to the reduction in colour
produced by the acid being greater than the increase caused by the
reagent, might be obtained when traces of free chlorine are present.
Similar difficulties are encountered when quantitative comparisons are
made against permanent standards.

Benzidine (Wallis[6]) has also been suggested for the detection of free
chlorine. On adding this reagent a blue colouration is produced but on
stirring it rapidly changes to a bright yellow which is proportional in
intensity to the amount of free chlorine present. Ellms and Hauser[5]
investigated benzidine in 1913 and found it to be inferior to
_o_-tolidine as a test reagent for free chlorine.

LeRoy[7] has proposed the use of
hexamethyltri_para_-aminotriphenylmethane for detecting and estimating
free chlorine. On the addition of a hydrochloric acid solution of this
compound to a sample containing free chlorine a violet colouration is
produced that can be matched in the usual way with standards. It is
stated that 0.03 p.p.m. of free chlorine gives a distinct colouration
and that the reagent reacts very slowly with nitrites and is quite
unaffected by hydrogen peroxide.

The starch-iodide and _o_-tolidine reactions are affected by oxidising
agents or reducible substances; nitrites and ferric salts are the
compounds that are most likely to interfere and Ellms and Hauser[5] have
found that these bodies do not affect the _o_-tolidine reaction to the
same extent as the starch-iodide reaction. Very small quantities of
nitrites (0.03 p.p.m. of N) and ferric salts (0.2 p.p.m. Fe) give a blue
colouration with the starch-iodide reagent and for this reason it is
always advisable, whenever possible, to make a control test on the
untreated water. Nitrites are oxidised by free chlorine and consequently
do not interfere with the estimation of it by the thiosulphate method;
the influence of ferric salts can be overcome by substituting 3 c.cms.
of 25 per cent phosphoric acid for hydrochloric acid (Winkler[8]).

An electrical instrument called a "chlorometer" has been devised by E.
K. Rideal and Evans[9] for the estimation of free chlorine. The
diagrammatic sketch, reproduced in Fig. 5, shows the general
construction of the apparatus. When water containing no free chlorine
passes through the copper tube, hydrogen is liberated on the platinum
rod by the electrolytic solution pressure of the copper and an electric
current is generated; a polarizing action follows and the flow of
current ceases. When free chlorine is present it combines with the
hydrogen as produced and so enables more copper to dissolve and produces
a permanent flow of current. The current produced is a function of the
depolarizing action, i.e. of the free chlorine, and is indicated by the
current meter which is graduated in parts per million of available
chlorine. The usual range of instrument is 5 p.p.m. and each division of
the scale is equal to one-tenth of one part per million.

Only strong oxidisers, such as chlorine, ozone, and permanganates, which
have a great affinity for hydrogen, are able to produce a permanent
current; ferric chloride and other weak oxidisers do not affect the

[Illustration: FIG. 5.--Rideal-Evans Chlorometer.]


_Cost of Construction._ According to the replies received by
the Committee on Water Supplies of the American Public Health
Association[10] the total cost of equipment for disinfection varies
widely and bears no apparent relation to the capacity of the equipment.
This is due to the temporary nature of the plants erected in many cities
and the necessity of erecting expensive structures in others. The cost
of construction varies also in different localities. The cost of
equipping hypochlorite plants with standard concrete tanks and dosage
regulators would be more uniform and for capacities between 10 and 50
million gallons per day would approximate $15 to $50 per million

_The operating cost_ of bleach plants shows similar wide variations. In
some cases the labour required for mixing and supervision can be
obtained without extra cost whilst in others the labour charge exceeds
the cost of hypochlorite.

The price of bleach has shown violent fluctuations during the last three
years (see Diagram IX, page 125) but is now (1918) comparatively steady
at $2.25 to $2.75 per 100 pounds. Assuming that 33.3 per cent of
available chlorine can be extracted, each pound of chlorine costs
6.75-7.25 cents as compared with 15-25 cents for liquid chlorine. The
fixed charges on the capital expenditures together with the labour and
incidental charges almost invariably make the total cost of operation of
a straight bleach plant higher than that of a liquid chlorine plant. The
tendency during the last four years has been to substitute liquid
chlorine for hypochlorite and the majority of the plants are now of the
former type.


Substances used for the removal of excess chlorine are usually known as
"antichlors" and those that have been most frequently employed are
sodium bisulphite, NaHSO_{3}, and sodium thiosulphate Na_{2}S_{2}O_{3}.
The reactions with chlorine are:

   (i) NaHSO_{3} + Cl_{2} + H_{2}O = NaHSO_{4} + 2HCl.

  (ii) Na_{2}S_{2}O_{3} + Cl_{2} = Na_{2}S_{4}O_{6} + 2NaCl.

Sodium bisulphite is a very efficient "antichlor," only 1.46 parts being
required to remove 1 part of chlorine, but owing to its instability the
action is uncertain. Sodium thiosulphate is a comparatively stable
cheap salt, containing 5 molecules of water of crystallization,
Na_{2}S_{2}O_{3} · 5H_{2}O but 7 parts are necessary to remove 1 part by
weight of chlorine.

"Antichlors" are used as aqueous solutions and the dosage controlled in
the same manner as for bleach solutions. The action is an instantaneous
one and it is consequently necessary that the germicidal action should
be complete before the "antichlor" is added.

Filters, containing solid materials capable of absorbing free chlorine,
have also been used for removing the excess of the germicidal reagent.
Iron borings and aluminium were used experimentally by Thresh[11] but
the process was not commercially developed. The "De Chlor" filter, in
which carbon is the active substance, has been installed at several
water works in England (Reading, Exeter, Aldershot) with apparently
successful results. The Reading experimental installation, described by
Walker,[12] consisted of a steel drum, 8 feet 3 inches in width, the top
and bottom being domed. In the upper portion, 10 feet 9 inches in depth,
provision was made for thorough admixture of the bleach solution and
water and a subsequent storage of thirty minutes. The lower section of
the filter was divided into three compartments, the first and last of
which contained graded silica; the middle compartment was filled with a
layer (20 inches deep) of specially prepared granulated charcoal or

The filter was operated under pressure and passed an average of 192,000
Imp. gallons per day, the rate being 32,000 Imp. gallons per square yard
per day.

Water from the pre-filters (polarite and sand) was treated with bleach
to give a concentration of 1 p.p.m. of available chlorine and passed
through the De Chlor filter. The average bacteriological results
obtained during the first six months operation were as follows:

                              Bacteria Per  c.cm.       B. coli Index
                           Gelatine 3 Days at 20° C.    Per 100 c.cms.

  Raw river water                    6,775                    600

  Water from pre-filters               579                    119

  Water from De Chlor filter            33                    Nil

Free chlorine could not be detected by chemical tests in the filtered
water which was also free from abnormal tastes and odours. It is stated
that the carbon has to be removed and revivified periodically. The
filter was washed about once per week, the wash water being only
one-tenth of one per cent.

The experimental filter was operated for nearly two years before being
removed to permit the erection of larger units having a total capacity
of one million Imp. gallons per day.


[1] Hooker. Chloride of Lime in Sanitation, New York, 1913.

[2] Griffen and Hedallen. J. Soc. Chem. Ind., 1915, =34=, 530.

[3] Hale. Proc. N. J. San. Assoc., 1914.

[4] Adams. J. Amer. Pub. Health Assoc., 1916, =6=, 867.

[5] Ellms and Hauser. J. Ind. and Eng. Chem., 1913, =5=, 915 and 1030;
_ibid._, 1914, =6=, 553.

[6] Wallis. Ind. Jour. Med. Res., 1917, =4=, 797.

[7] Le Roy. Comptes rend., 1916, =163=, 226.

[8] Winkler. Zeit. angew. Chem., 1915, =28=, 22.

[9]: Rideal, E. K. and Evans. Analyst, 1913, =38=, 353.

[10] J. Amer. Pub. Health Assoc. 1915, =5=, 921.

[11] Thresh. Internat. Congress Appl. Chem., 1908.

[12] Walker. Jour. Roy. Inst. Pub. Health, Jan., 1911.



The use of liquefied chlorine for the disinfection of water was first
proposed by Lieutenant Nesfield[1] of the Indian Medical Service. He
stated that: "It occurred to me that chlorine gas might be found
satisfactory ... if suitable means could be found for using it.... The
next important question was how to render the gas portable. This might
be accomplished in two ways: By liquefying it, and storing it in
lead-lined iron vessels, having a jet with a very fine capillary canal,
and fitted with a tap or a screw cap. The tap is turned on, and the
cylinder placed in the amount of water required. The chlorine bubbles
out, and in ten to fifteen minutes the water is absolutely safe, and has
only to be rendered tasteless by the addition of sodium sulphite made
into a cake or tablet.... The cylinders could, of course, be refilled.
This method would be of use on a large scale, as for service water

The first _practical_ demonstration of the possibilities of this method
was made by Major Darnall[2] of the Medical Corps, United States Army,
in 1910. Chlorine was taken from steel cylinders and passed through
automatic reducing valves which provided a uniform flow of gas for the
water requiring treatment. A uniform flow of water was maintained
through the mixing pipe and so secured a uniform dosage. This apparatus
might be considered as the forerunner of the various commercial types of
machines that were developed later and which are being so extensively
used at the present time.

A working model, having a capacity of 500 gallons per hour, was erected
at Fort Myer, Va., and was operated on water that had been treated with
alum but had received no further purification. Despite the presence of
the flocculated organic matter, satisfactory purification was obtained
with 0.5 to 1.0 p.p.m. of available chlorine and no taste or odour was
imparted to the supply.

From the results obtained at Fort Myer, and Washington, D.C., Darnall
concluded that "In general, it may be said that with an average
unfiltered river water such as that of the Potomac, about one-half of
one part (by weight) of chlorine gas per million of water will be
required. For clear lake waters three-tenths to four-tenths of a part
per million will be sufficient."

A Board of Officers of the War Department examined the results and
reported (June, 1911) "That the apparatus is as efficient as
purification by ozone or hypochlorite and is more reliable in operation
than either.... That it could be installed at a very low cost and that
the cost of operation would be very slight."

In June, 1912, Ornstein experimented with chlorine gas, obtained from
the liquefied gas in cylinders, for sewage and water disinfection but
his method differed from Darnall's in first dissolving the gas in water
and feeding the solution to the liquid to be treated.

Kienle[3] made experiments at Wilmington, Del., in November, 1912, and
obtained a constant flow of gas by means of high- and low-pressure
valves; the gas was dissolved in water in an absorption tower and
afterwards fed to the water to be treated.

Van Loan and Thomas of Philadelphia experimented with liquid chlorine on
a large scale at the Belmont Filter Plant in September, 1912. The
chlorine was fed into the filtered water basin in the gaseous state and
the quantity was regulated by the loss in weight of the containers. The
dosage was approximately 0.14 p.p.m. (West[4]).

Jackson, of Brooklyn, made similar experiments about the same time at
the Ridgewood Reservoir, Brooklyn, and his type of apparatus was shortly
afterwards put on the market as the Leavitt-Jackson Liquid Chlorine
Machine. The regulation of the flow in this machine was determined by
the loss in weight of the gas cylinder which was suspended from a
sensitive scale beam. By moving the counterbalancing weight on the beam
at a constant rate, a uniform flow of gas was obtained, the area of the
orifice being kept constant by the equilibrium in the balance operating
controlling valves through a system of levers.

This type of apparatus was tried at several places but it was found that
the adjustment of the regulating mechanism was too sensitive and
produced considerable irregularities in the flow of gas.

The type used by Ornstein and Kienle were combined and commercially
developed by the Electric Bleaching Gas Co. of New York.[A] In this
combined type the gas was collected from one or more cylinders by means
of a manifold which delivered it to the regulating mechanism at the
pressure indicated by a gauge attached to the inlet pipe. Beyond this
gauge were two pressure-regulating devices, the first being used
primarily to reduce the initial pressure to about 15 pounds per square
inch, and the second for controlling the pressure through a range
sufficient to give the desired discharge of gas. The gas from the second
regulator passed through an orifice in a plate at a pressure indicated
by a suitable gauge which was calibrated in terms of weight of chlorine
per unit of time. The gas, on leaving the regulating apparatus, passed
up an absorption tower of hard rubber, where it met a descending stream
of water. The solution was carried by suitable piping to the point of
application. This type was modified in some cases by the substitution of
a flow meter of the float type for the inferential pressure meter.

  [A] This type has recently been withdrawn from the market.

[Illustration: FIG. 6.--Manual Control Chlorinator, Solution Feed, Type

Another type of apparatus, developed by Wallace and Tiernan,[A] is shown
in Figs. 6 and 7. The gas under the pressure indicated by the tank
pressure gauge (Fig. 6) passes into the pressure compensating chamber,
which maintains a constant drop in pressure across the chlorine control
valve, through the check valve, and into the solution jar after
measurement in the pulsating meter. The water required for dissolving
the chlorine enters the jar through the feed line and check valve and
the solution passes along the feed line after being water sealed in a
special chamber. The meter is a volumetric displacement one and is
regulated by observing the number of pulsations per minute. Each
pulsation corresponds to 100 milligrams or 0.00022 pound of chlorine;
diagrams for converting pulsations per minute into weight per
twenty-four hours are usually provided with the apparatus. This type of
meter is suitable for quantities between 0.1 and 12 pounds per day and
possesses the distinct advantage of enabling the operator to see the
actual delivery of the gas.

  [A] Manufactured by Wallace and Tiernan Co. Inc. N. Y.

[Illustration: FIG. 7.--Manual Control Chlorinator, Solution Feed, Type

The quantities of gas exceeding 12 pounds per day the type shown in Fig.
7 may be used. The gas from the control valve passes through a visible
glass orifice which is connected with the manometer. This manometer, or
chlorine meter, contains carbon tetrachloride and is graduated
empirically in terms of weight of chlorine per unit of time. A suitable
gauge indicates the back pressure thrown by the check valve and
registers the same pressure as the tank gauge when the flow of gas is
stopped. The gas passes into the glass cylinder where it is dissolved in
water and passes out by the feed pipe.

The most accurate range of the orifice type is from 1-6, i.e. if the
minimum graduation on the scale is 10, the maximum is 60. If quantities
less than the minimum graduation are desired, a smaller orifice with its
corresponding scale can be substituted in a few minutes.

These types are manually controlled, but automatic control types, to
meet almost any condition, can be obtained and are in use in many

In some instances (dry-feed types) the chlorine gas is not dissolved in
water prior to addition to the water requiring treatment but is carried
to the point of application as a dry gas and enters the water through a
diffusion plate made of carborundum sponge. The sponge becomes saturated
with water because of the capillary action of the carborundum upon the
water. The pressure of the chlorine in the feed pipe forces the gas
through the diffuser in the form of minute bubbles which become
saturated with moisture. On meeting the water they immediately go into
solution and no gas escapes.

The operation of liquid chlorine machines is exceedingly simple. After
the cylinders have been connected, the cylinder valves are opened and
the joints tested for leakage by holding a swab of absorbent cotton
saturated with strong ammonia under them; a leakage is indicated by the
appearance of white fumes of ammonium chloride. The control valve is
then slightly opened and the auxiliary cylinder valves partially opened;
whilst the pressure in the apparatus is slowly increasing the remainder
of the joints are tested and if found to be tight, the cylinder valves
are fully opened and the control valve opened to the desired amount. In
the solution feed types the water required as solvent is turned on
before the control valve is opened. Once the apparatus is working, no
further attention is required, except for the regulation of the dosage
in the manual control types, until the cylinders are replaced. When the
stock of gas in the cylinders is almost depleted the pressure falls but
it is always preferable to determine the stock by standing the cylinders
on a platform scale and weighing at regular intervals. This also
provides a check on the apparatus and can be utilised to check the

The accumulation of substances that impede the flow of gas is usually
slow and is indicated by a gradual increase in the back pressure. The
orifice is calibrated at 25 pounds back pressure and any deviation from
this figure will show a discrepancy between the actual weight of
chlorine evaporated and the amount calculated from the scale reading.

Liquid chlorine is usually sent out by the manufacturers in steel
cylinders which contain about 1.1 cubic feet of liquid or approximately
100 pounds (1 cu. ft. = 89.75 pounds).[A]

  [A] An effort is now being made to standardise cylinders of 150 lbs.

For small installations only one cylinder is necessary but it is always
preferable to connect more than one. When the flow of gas is rapid the
temperature of the liquid chlorine falls and reduces the pressure. The
effect of the fall in temperature, due to the latent heat of
evaporation, can be partially overcome by using a larger number of
cylinders; in addition a source of external heat should be provided that
will maintain the temperature of the cylinders at a minimum of 80° F.
This is a "sine qua non" for successful operation. The effect of the
temperature upon the pressure in the cylinders is shown in Diagram VII.

[Illustration: DIAGRAM VII


In practice it is found impossible to utilise all the gas contained in
the containers; when the cylinders are almost empty the pressure
necessary for the operation of the regulating device cannot be obtained
and full cylinders must be attached. When sufficient heat is provided
the weight of chlorine in the cylinder can be reduced to 1 - 1-1/2
pounds before the tank pressure becomes too low.

Liquid chlorine machines will operate, with ordinary care, for long
periods. The various parts are made of such metals as experience has
demonstrated to be best able to resist the corrosive action of the dry
gas and the apparatus is designed to prevent the access of moisture
which would otherwise produce corrosion and impede the flow of gas.
Stoppages are sometimes caused by brown deposits derived from
impurities in the liquid chlorine. These are primarily due to variations
in the graphite electrodes used in the electrolytic process for the
manufacture of chlorine from salt.

[Illustration: FIG. 8.--Dunwoodie Chlorinating Plant Treating
400,000,000 Gallons Per Day for New York City.]

To convey the dry gas from the apparatus to the point of application,
copper or iron pipes may be used; for aqueous solutions, flexible rubber
hose must be employed. Chlorine water is exceedingly active, chemically,
and rapidly attacks all the common metals; ordinary galvanised iron pipe
is eroded in a few days and should never be used.

Liquid chlorine, for water disinfection, possesses several marked
advantages over the ordinary bleach process.

(1) The sterilising agent is practically 100 per cent pure, the only
impurities being traces of carbon dioxide and air, and does not
deteriorate on storage; it will, in fact, keep almost indefinitely.

(2) Liquid chlorine practically eliminates all labour costs because of
the simplicity of the apparatus and the concentrated form of the
sterilising agent. The apparatus is so compact that all the cylinders
and regulating apparatus required for delivering 200 pounds of gas per
day can be placed in an area of about 50 square feet and it can
consequently be almost invariably accommodated in locations where the
trifling amount of attention required can be obtained without extra

(3) The sludge problem, inseparable from bleach installations, is

(4) Regulation of the dosage is simpler and consequently usually more
accurate. The dosing apparatus in bleach plants invariably tends to
choke and demands regular attention from intelligent operators; a
similar tendency in liquid chlorine machines is easily detected and
electrical devices can be installed to indicate automatically any
changes in the flow.

(5) The first cost is smaller. The cost of liquid chlorine machines
varies from $400, for the small manual control types, to $1,200, for the
automatic control types. The capital outlay is mainly determined by the
number of machines and accessories required and not, within certain
limits, by the capacity. One machine will deliver up to 200 pounds of
gas per day, an amount sufficient to treat 60,000,000 U. S. A. gallons
(50,000,000 Imp. gals.) at 0.40 p.p.m. of available chlorine. Unless
duplicate machines are installed for the higher rates, the first cost is
inversely proportional, though not directly so, to the volume of water
treated. It is in all cases less than the first cost of a bleach plant
of equal capacity, accuracy, and durability.

(6) Liquid chlorine installations usually tend to produce less
complaints as to tastes and odours. This is probably due, not to any
merit of the chlorine _per se_, but to a more accurate regulation of the
dosage and efficient distribution of the chlorine in the treated water.
The advantages ensuing from thorough admixture had only become partially
appreciated before liquid chlorine machines were fully developed and
they have been more fully utilised in the design of these later

Claims have also been made that liquid chlorine prevents "aftergrowths"
but no evidence can be adduced in support of this statement.
Aftergrowths have occurred at many places where this process is employed
and in this respect it possesses no advantage over hypochlorite

It is also claimed that one pound of liquid chlorine is more efficient,
as a germicide, than an equal weight of chlorine in the form of bleach.
Jackson[5] has stated that 1 pound of chlorine is equal to 9 pounds of
bleach; Kienle (_loc. cit._) that it was equal to 8 pounds of bleach,
whilst Huy claimed to have obtained an efficiency ratio of 1:10 at
Niagara Falls, N. Y. The conditions of the experiment were not
comparable however, in the last mentioned ratio. Catlett, at Wilmington,
N. C. (West[4]) obtained a better bacterial reduction with 1 pound of
liquid chlorine than with 6 pounds of bleach.

The efficiency ratio of chlorine to bleach has been reported upon by
West.[4] From 1910-1913 the mixed filter effluents of the Torresdale
plant at Philadelphia were treated with bleach but in November, 1913 the
liquid chlorine process was substituted. On comparing the results
obtained during the same months of the two periods it was found that, in
general, 1 pound of liquid chlorine gave a slightly higher percentage
purification than 6-7 pounds of bleach. Similar results were obtained at
the other Philadelphia plants. The figures published by West show that
the hypochlorite solutions used were abnormally strong (3.6-10.4 per
cent of available chlorine), a condition that would increase the
difficulty of extracting all the soluble hypochlorite. It was found
indeed, that, under the most advantageous conditions, only 87 per cent
of the available chlorine was extracted. The average chlorine content of
the bleach used during 1912-1913 was 36.1 per cent but the figures given
would indicate that at least 1.5 per cent, a reduction of 4.6 per cent
of the total, was lost during storage. It would seem not improbable that
the total loss under average conditions was not less than 20 per cent,
which would reduce the efficiency ratio to 1:4.8-5.6.

Hale[6] also made a comparison of the relative efficiency of liquid
chlorine and hypochlorite of lime at New York, and the earlier results
agreed with West's ratio of 1:6-7. An investigation showed that large
quantities of chlorine were not extracted from the bleach and when this
condition was rectified the total loss averaged only 4 per cent and the
results obtained were equal to those given by the liquid chlorine
machines. Hale's comparative figures are given in Table XXIII.


    Treatment.    |  Water   | Number of |  Chlorine  |   Reduction
                  | Treated. |  Samples. |   p.p.m.   |  of B. coli.
  Bleach          | Croton   |    84     | 0.27-0.36  |     93%
  Liquid chlorine | Bronx    |    84     | 0.27-0.36  |     93%

Hale concluded that, when efficiently used, the ratio of chlorine to
bleach required to produce equal bacterial purification, approached 1:3.

The results obtained by the author in Ottawa are similar to those of
Hale. During the earlier period of the bleach treatment a dosage of 1.5
p.p.m. of available chlorine was required to obtain satisfactory
purification but various improvements that were subsequently made
enabled the quantity to be reduced to 0.8 p.p.m. The same raw water
usually requires 0.75 to 0.80 p.p.m. of liquid chlorine to obtain the
same purification. The total losses in the Ottawa bleach plant averaged
6-8 per cent and based on these figures the efficiency ratio is
approximately 1:3.5.

Ratios as low as 1:3.5 can only be obtained by the supervision of a
chemist and this analytical control involves additional expense that
must be charged against the bleach process. No chemical analyses are
necessary for the control of liquid chlorine plants.

_Disadvantages of Liquid Chlorine Plants._ The main objection to the use
of liquid chlorine is that the slight leaks of gas occur occasionally
and unless removed by forced ventilation may produce a concentration of
chlorine that will injure the operators.

Pettenkofer and Lehmann[7] found that 0.001-0.005 per cent of chlorine
in air affected the respiratory organs; 0.04-0.06 per cent produced
dangerous symptoms, whilst concentrations exceeding 0.06 per cent
rapidly proved fatal.

The danger of gas leakages can be eliminated by placing the apparatus in
a small separate room provided with a fan and a ventilation duct. By the
liberal use of glass in the construction of the room, the operation of
the plant can be seen at all times without entering the chamber.

A portion of the liquid chlorine apparatus is made of glass and is
consequently easily fractured. Duplicates of the glass parts should be
kept in stock to prevent interrupting the supply of gas; a duplicate
machine is also advisable in large installations.

_Cost of Treatment._ Prior to the outbreak of war in 1914, liquid
chlorine sold at 10-11 cents per pound in small quantities and for 8-9
cents per pound in large shipments. In 1917 the price was 18-20 cents
per pound for small quantities and 15 cents upwards for large contracts.
Canadian prices are 25 per cent higher.

The amount of chlorine required for satisfactory disinfection (see
Chapter III) depends upon the nature of the water and the cost of
treatment varies accordingly. In the majority of plants the cost varies
from 25-90 cents per million gallons.

_Popularity of Process._ Since 1913, when the first commercial liquid
chlorine machines were used, the popularity of this process has
increased in a most remarkable manner. In 1913 over 1,700 million
gallons per day were treated with hypochlorite; in 1915, 1,000 million
gallons per day were treated with liquid chlorine and an approximately
equal amount with hypochlorite; in January 1918, the amounts were 3,500
million gallons per day (liquid chlorine) and 500 million gallons per
day (hypochlorite).

This wonderful development has been largely due to the intrinsic merits
of the process and the reliability of the machines manufactured although
it has been indirectly assisted by the excessive cost of hypochlorite
during 1915-1916.

Liquid chlorine machines are being used for the purification of water on
the Western Front of the European battlefield. The outfit is a mobile
one and consists of a rapid sand filter, liquid chlorine apparatus, a
small storage tank and solution tanks. Owing to the limited contact
period available a large dosage of chlorine is employed and the excess
afterwards removed by the addition of a solution of sodium thiosulphate.

_Chlorine Water._ Marshall[8] has proposed the use of chlorine water for
the sterilisation of water for troops. The solution is contained in
ampoules which are of two sizes, one for water carts and the other for
water bottles of one quart capacity.

The coefficient of solubility of chlorine, from 10°-41° C. is _C_ =
3.0361 - 0.04196_t_ + 0.0001107_t_^{2}; when _t_ = 10° C. 1 c.cm. of
water absorbs 2.58 c.cms. of chlorine or 8.2 m.gr., a quantity
sufficient to give a concentration of 1 p.p.m. in 8 litres of water.
Marshall has stated that, when pure materials are used, chlorine water
is stable but the author is unable to confirm this. A saturated solution
of chlorine in distilled water lost over 50 per cent of its available
chlorine content when stored for five days in the dark at 70° F. The
chlorine present as hypochlorous acid increased slightly but the
quantity never exceeded very small proportions. Chlorine solutions
decompose in accordance with the equation, Cl_{2} + H_{2}O = 2HCl + O.

Although chlorine water appears to be of little value because of its
instability there appears to be no reason why chlorine hydrate should
not be successfully employed. The hydrate was first prepared by
Faraday[9] by passing chlorine into water surrounded by a freezing
mixture. A thick yellow magma resulted from which the crystals of
chlorine hydrate were separated by pressing between filter paper at 0°
C. The hydrate prepared by Faraday was found to have the composition
represented by the formula Cl·5H_{2}O but later investigators have shown
that more concentrated hydrates can be prepared. Roozeboom[10] prepared
a hydrate represented by the formula Cl·4H_{2}O and Forcrand[11] one
containing only 3-1/2 molecules of water (Cl_{2}·7H_{2}O). Chlorine
hydrate separates into chlorine gas and chlorine water at 9.6° C. in
open vessels and at 28.7° C. in closed vessels. Pedler[12] has shown
that when the ratio of Cl_{2}:H_{2}O is 1:64 or greater, the mixture of
chlorine hydrate and water exhibits great stability and can be exposed
to tropical sunlight for several months without decomposition.

Cl_{2}·64H_{2}O contains 5.8 per cent of chlorine and about 8. c.cms.
would be required to give a concentration of 1 p.p.m. in 110 Imp.
gallons of water, the usual capacity of a military water cart.


[1] Nesfield. Public Health, 1903, =15=, 601.

[2] Darnall. J. Amer. Pub. Health Assoc., 1911, =1=, 713.

[3] Kienle. Proc. Amer. Waterworks Assoc., 1913, 274.

[4] West. J. Amer. Waterworks Assoc., 1914, =1=, 400-446.

[5] Jackson. Proc. Amer. Waterworks Assoc., 1913.

[6] Hale. Proc. N. J. San. Assoc., 1914.

[7] Pettenkofer and Lehmann. Munich Acad., 1887.

[8] Marshall. Conv. Amer. Elect. Chem. Soc., 1917. Eng. and Contr.,
1918, =49=, 40.

[9] Faraday. Q. J. S., =15=, 71.

[10] Roozeboom. Rec. Trav. Chim., 1885, =3=, 59.

[11] Forcrand. Comp. rend., 1902, =134=, 991.

[12] Pedler. J. C. S., 1890, =83=, 613.



Since 1889 when Webster first proposed the use of electrolysed sea-water
as a disinfectant, various attempts have been made to introduce
electrolytic hypochlorites for the bactericidal treatment of water and
sewage. Two of these preparations were named Hermite fluid, and
electrozone (c.f. page 5). Sodium hypochlorite, made by passing chlorine
into solutions of caustic soda, or by the decomposition of bleach by
sodium carbonate, has also been used and preparations of this character
have been sold under such names as Eau de Javelle, Labarraque solution,
chloros, and chlorozone. These solutions contain mixtures of sodium
hypochlorite and sodium chloride together with some free alkali.
Chlorozone was the name given by Count Dienheim-Brochoki to a number of
preparations patented in 1876 and subsequently down to 1885. They were
produced by passing air and chlorine into solutions of caustic soda.
Lunge and Landolt[1] have shown that the air introduced is without
effect and that the advantages claimed for chlorozone are illusory.

The earliest electrolytic installation on this continent was operated at
Brewster, N. Y., in 1893 and since that date several plants have been
erected where local conditions conduced to economical operation.

When a uni-directional current of electricity is passed through a
solution of sodium chloride, the salt is dissociated and the components
liberated, NaCl = Na + Cl. If the elements are not separated, the
chlorine combines with the sodium hydrate, formed by the action of the
sodium on the water, to form sodium hypochlorite. The equations 2Na +
2H_{2}O = 2NaOH + H_{2}, and 2NaOH + Cl_{2} = NaOCl + NaCl + H_{2}O show
that only one-half of the chlorine produced is found as hypochlorite;
the other half reforming sodium chloride.

Several types of electrolysers have been used for the production of
hypochlorites and chlorine but only two are suitable for water-works
purposes: in one, the cathodic and anodic products recombine in the main
body of the electrolyte; in the other, the diaphragm process, they are
separated as removed and the final products are chlorine gas and a
solution containing caustic soda and some undecomposed salt.

Until a few years ago the non-diaphragm process was the only one used
for water treatment and it will consequently be discussed first.

_Non-diaphragm Process._ The theoretical voltage required for the
decomposition of sodium chloride is 2.3 but when the products recombine
in the electrolyte, side reactions occur which increase the minimum
voltage to 3.54. On this basis one kilowatt hour gives 272 ampere hours
and as one ampere hour is theoretically capable of producing 1.33 grams
of chlorine, 1.21 kilowatt hours are necessary for the production of 1
pound of chlorine by the decomposition of 1.65 pounds of salt.

Charles Watt (1851) discovered this process and was the first to
recognize the necessary conditions which are (1) insoluble electrodes,
(2) low temperature of electrolyte, and (3) rapid circulation of
electrolyte from the cathode to the anode. The control of the
temperature is very important, for as it increases, side reactions occur
with the formation of chlorates, and the efficiency is decreased.

The non-diaphragm cells used in Europe (Haas and Oettel, Kellner,
Hermite, Vogelsand, and Mather and Platt) have been described by
Kershaw.[2] In the Haas and Oettel electrolyser the electrodes are
composed of carbon but in the other types at least one electrode is made
from platinum or a platinum alloy. The Dayton electrolyser, which is the
cell most familiar in North America, is shown in Fig. 9.

[Illustration: FIG. 9.--Dayton Electrolytic Cell.]

The outer cell is made of soapstone and is approximately 2-1/2 feet long
and 2 feet wide. The main electrodes consist of four pieces of Atcheson
graphite connected together by screws and metal strips to which is
attached a clamp for connecting electrical terminals. Circulation of the
brine is produced by glass baffle plates and secondary electrodes placed
one inch apart between the main electrodes. The cell is intended to be
used at 110-volts pressure but by wiring two cells in series a 220-volt
circuit may be employed. An inlet and outlet are provided at each end of
the tank to permit the direction of the flow to be periodically reversed
for the purpose of removing the lime deposit from the graphite plates.

The salt solution is prepared in wooden tanks from coarse clean salt
(ground rock salt is unsuitable), containing as little iron as possible,
in the proportion of 50 pounds to 100 gallons of water. After passing
through a gravel or other suitable filter the brine solution is carried
by brass pipes to the electrolyser. The rate of flow is adjusted to the
temperature of the hypochlorite solution leaving the cell but under
normal conditions it is stated that the cell described will pass 40
gallons per hour with a consumption of 70 amperes and produce 2-1/2
pounds of chlorine per hour. This is equal to 8 pounds of salt and 3.08
kilowatt hours per pound of chlorine. After the cells have been operated
for several months the efficiency usually falls and 10-11 pounds of salt
and 3.5-3.7 kilowatt hours are required for the production of one pound
of chlorine. The concentration of the hypochlorite solution is usually
about 6 grams per litre.

Rickard[3] stated that by cooling the Dayton cell with ice 1 pound of
chlorine could be produced from 2.65 kilowatt hours and 6.9 pounds of
sodium chloride; without cooling the figures were 3.62 kilowatt hours
and 7.2 pounds of salt. Based on the figures that have been obtained
from mature cells, the efficiency of the Dayton cell as compared with
those described by Kershaw is as follows:

                    |        SALT.         |        POWER.
                    |     PER POUND OF AVAILABLE CHLORINE.
    Type of Cell.   +----------+-----------+----------+-----------
                    |  Pounds. | Per Cent  | Kilowatt | Efficiency
                    |          | Consumed. |  Hours.  | Per Cent.
  Haas and Oettel   |   10.7   |    15.4   |   3.8    |    31.9
  Kellner           |    7.5   |    22.0   |   2.75   |    43.9
  Hermite           |   11.2   |    14.5   |   2.87   |    42.2
  Mather and Platt  |   ....   |    ....   |   2.75   |    43.9
  Dayton            |   10.0   |    16.5   |   3.6    |    33.6
  Theoretical       |    1.65  |   100.0   |   1.21   |   100.0

The cost of production depends upon local conditions: if alternating
current is available at $30 per horse-power per annum, and low-grade
salt can be obtained for $5 per ton the cost of 1 pound of chlorine
would be

                   | COST (CENTS) PER POUND OF
                   |    AVAILABLE CHLORINE.
    Type of Cell.  +--------+----------+--------
                   |  Salt. | Current. | Total.
  Haas and Oettel  |  2.67  |   1.97   |  4.64
  Kellner          |  1.87  |   1.43   |  3.30
  Hermite          |  2.80  |   1.49   |  4.29
  Dayton           |  2.50  |   1.92   |  4.42

The electrical and chemical efficiencies of the Haas and Oettel and
Dayton cells, which contain carbon electrodes, are smaller than those
containing platinum electrodes but for water sterilisation the carbon
cells have been found to be more suitable. To prevent the action of
magnesium salts on the platinum electrodes it is necessary to use a
higher grade of salt or to provide means of purification. Because of the
absence of a base and the presence of chlorides, electrolytic
hypochlorite cannot be stored for more than a few hours without
appreciable loss of titre. Unless used for the treatment of the effluent
of a filter plant operated at a fairly constant rate a small storage
tank is necessary to compensate for the irregular demand and to provide
the head required by orifice feed boxes. Small variations can be made by
regulating the flow through the cells but large ones are not compatible
with efficiency, which is the highest under a constant load.

Claims have been made that electrolytic hypochlorite is more efficient
as a germicide than bleach when compared on the basis of their available
chlorine content but no evidence of it has been produced. Bleach
contains an excess of base, which retards the germicidal action, and
electrolytic hypochlorite contains an excess of sodium chloride, which
accelerates it (Race[4]) but the ultimate effect is the same with both.
This is shown in Table XXIV.


                  |   BLEACH.         ||  ELECTROLYTIC
                  |                   ||  HYPOCHLORITE.
  Contact Period. | Available Chlorine. Parts Per Million.
                  |   0.4   |   0.6   ||   0.4   |   0.6
  Nil             |   182   |   ...   ||   ...   |   ...
  10 minutes      |   130   |    10   ||   120   |    8
   1 hour         |    66   |     1   ||    60   |    0
   2 hours        |     3   |     0   ||     1   |    0
   3-1/2 hours    |     0   |     0   ||     0   |    0

  [A] Results are B. coli per 10 c.cms.

Electrolytic hypochlorite has a greater germicidal velocity than bleach
but the difference is so small as to be of no practical importance.
Rabs[5] experimented with various hypochlorites but was unable to find
any appreciable differences in their germicidal action.

If electrical power can be obtained at a very low cost, or if the cost
is merely nominal, as it is when there is an appreciable difference
between the normal consumption and the peak load upon which the rate is
based, the electrolytic hypochlorite method offers some advantages but
in the great majority of plants it cannot economically compete with
bleach. The instability of the liquor and cell troubles have also
prevented the process being generally utilised. Baltimore and Cincinnati
experimented with this method but did not adopt it. Winslow[6] has
reported that, owing to the difficulty in obtaining bleach since the
outbreak of war, Petrograd has used electrolytic hypochlorite made from

_Diaphragm Process._ The various types of diaphragm cells that have been
commercially operated are of two varieties: (1) cells with submerged
diaphragms and (2) cells in which the electrolyte comes in contact with
one face only of an unsubmerged diaphragm.

The Le Sueur, Gibbs, Crocker, Billiter-Siemens, Nelson, and
Hargreaves-Bird cells are of the submerged diaphragm variety. The Nelson
cell has been operated for some time at the filtration plant at Little
Falls, N. J. The cells are fed with brine solution previously purified
by the addition of soda ash and have given fairly successful results
although the cost of maintenance is comparatively high. Tolman[7] has
reported that several towns in West Virginia use a bleach solution
prepared by absorbing chlorine, manufactured by the Hargreaves-Bird
process, in lime water; the solution contains about 1.95 per cent of
available chlorine.

The diaphragms in both the submerged and unsubmerged types are usually
constructed either with asbestos paper or cloth, placed in such a manner
as to divide the cells into two separate compartments: the anodic, into
which the brine is fed and where the chlorine is produced; and the
cathodic, where caustic soda is formed.

By maintaining the liquor in the anodic compartment at a higher
elevation than in the cathodic one, the direction of flow is towards the
latter, but owing to osmosis and diffusion the separation is not
complete and a portion of the caustic soda passes the diaphragm and
produces hypochlorite with a consequent loss of efficiency and rapid
deterioration of the anodes. With the exception of the Billiter-Siemens
cell, the submerged diaphragm cells operate at not more than 85 per cent
efficiency and the cost of maintenance is usually high.

In the non-submerged diaphragm types the invasion of the anodic
compartment by caustic is much reduced and the efficiency and life

An electrolyser of the non-submerged diaphragm type is the Allen-Moore
cell which has been adopted by the Montreal Water and Power Co. This
has been described by Pitcher and Meadows.[8] The general lay-out of the
installation is shown in Fig. 10, and the essential features are: a salt
storage bin having a capacity of 40 tons; the brine saturating and
purifying apparatus; duplicate 15 horse-power motor-generator sets; four
chlorine cells; and the silver ejectors and distributing lines for
carrying the chlorine solution to the point of application.

[Illustration: FIG. 10--Brine Saturating and Purifying Equipment.]

The brine solution, which is prepared by passing water through the
saturators previously filled with salt, is delivered to the two concrete
reaction tanks where an amount of soda ash and caustic liquor sufficient
to combine with the calcium and magnesium salts is added, and the
mixture filtered through sand and stored in the purified brine tanks. To
prevent the formation of hypochlorites by the interaction of chlorine
and alkali, the alkalinity of the liquor is determined and sufficient
hydrochloric acid added to ensure an acidity of 0.01 per cent. The acid
brine is delivered at one end of the four cells (Fig. 11) each of which
is 7 feet long and 20-3/8 inches wide and consumes 600 amperes at 3.3
volts. The cell box is built of concrete and is provided with a
perforated wrought iron cathode box and graphite anode plates which are
separated by an unsubmerged asbestos paper diaphragm.

[Illustration: FIG. 11.--Sections of Allen-Moore Cell.]

Each cell has a capacity of 32 pounds of chlorine per day and the gas
flow is determined by measuring the volume of caustic soda produced in a
given period of time and calculating the weight from the volume and
concentration as determined by titration with standard acid; each gram
of NaOH is equal to 0.88 gram of chlorine. The efficiency of the cell is
obtained by dividing the number of grams of chlorine produced per hour
by the product of the current volume (in amperes) and the factor 1.33,
the theoretical production of chlorine for one ampere hour. The average
efficiency of the Montreal cells was found to be 93 per cent. The
installation comprises four cells, one being held in reserve, and the
annual cost of producing 90 pounds of chlorine per day is given as
$2,500. The details are:

  Salt at $8.00 per ton, delivered         $500.00
  Power, 15 H.P., at $30.00 flat rate       450.00
  Labour and superintendence                500.00
  Interest at 6 per cent on capital cost    300.00
  Depreciation, 15 per cent                 750.00

cost per pound of chlorine = 7.6 cents.

The diaphragm cells, like the non-diaphragm ones, operate most
efficiently under a constant load; they are consequently suitable for
treating the effluent of filter plants.

Where very cheap electrical power can be obtained, the cost per pound of
available chlorine is less for the electrolytic method just described
than for liquid chlorine or chlorine obtained from bleach; but this
condition obtains in very few places. Mr. J. A. Meadows has suggested to
the author that the cost could be reduced by converting the chlorine gas
into hypochlorite and then adding dilute ammonia as in the chloramine
process (_vide_ page 115). The caustic liquor, usually run to waste from
the cathodic compartment, could be delivered into a feed box from which
it would be drawn off by the water injector used for dissolving the
chlorine gas.


[1] Lunge and Landolt. Jour. Soc. Dyers and Colourists, Nov. 25, 1885.

[2] Kershaw. Jour. Soc. Chem. Ind., 1912, =31=, 54.

[3] Rickard. Quar. Bull. Ohio Board of Health, Oct.-Dec., 1904.

[4] Race. Jour. Amer. Waterworks Assoc., 1918, =5=, 63.

[5] Rabs. Hygienische Rundschau, 1901, 11.

[6] Winslow. Public Health Rpts. U. S. P. H. S., 1917, =32=, 2202.

[7] Tolman. Jour. Amer. Waterworks Assoc., 1917, =4=, 337.

[8] Pitcher and Meadows. Jour. Amer. Waterworks Assoc., 1917, =4=, 337.



Chloramine (NH_{2}Cl), a chemical compound in which one of the hydrogen
atoms of ammonia has been replaced by chlorine, was discovered by
Raschig[1] in 1907. Chloramine was prepared by cooling dilute solutions
of bleach and ammonia and adding the latter to the former contained in a
flask surrounded by a freezing mixture. The proportions were as the
equivalent weights of anhydrous ammonia and available chlorine
(approximately two parts by weight of chlorine to one part by weight of
ammonia). After gas evolution had ceased the mixture was saturated with
zinc chloride and the magma distilled under reduced pressure. The
distillate was a dilute solution of comparatively pure chloramine.

The first to notice the effect of ammonia on the germicidal value of
hypochlorites was S. Rideal[2] who noted that during the chlorination of
sewage, the first rapid consumption of chlorine was succeeded by a
slower action which continued for days in some instances, and was
accompanied by a germicidal action after free chlorine or hypochlorite
had disappeared. Rideal stated that: "It became evident that chlorine,
in supplement to its oxidising action, which had been exhausted, was
acting by substitution for hydrogen in ammonia and organic compounds,
yielding products more or less germicidal." On investigating the effect
of ammonia on hypochlorite it was found that the addition of an
equivalent of ammonia to electrolytic hypochlorite increased the
carbolic acid coefficient of 2.18, for one per cent available chlorine,
to 6.36 (nearly three times the value). Further experimental work showed
that the increase was due to the formation of chloramine.

The author, in 1915, during a series of experiments on the relative
germicidal action of hypochlorites, attempted to prepare the ammonium
salt by double decomposition of bleach and ammonium oxalate solutions.

  Ca(OCl)_{2} + (NH_{4})_{2}C_{2}O_{4} = CaC_{2}O_{4} + 2NH_{4}OCl.

The velocity of the germicidal action of the solution was found to be
about ten times greater than the germicidal velocities of other
hypochlorites of equal concentrations, (Race[3]), and from a
consideration of the chemical formula of ammonium hypochlorite it
appeared probable that it would be very unstable and decompose into
chloramine, which Rideal had previously shown to have an abnormal
germicidal action, and water. NH_{4}OCl = NH_{2}CL + H_{2}O. After these
results have been confirmed, the effect of adding ammonia to bleach
solution was tried and it was found that 0.20 p.p.m. of available
chlorine and 0.10 p.p.m. of ammonia produced equally good results as
0.60 p.p.m. of chlorine only. Similar results were obtained on the
addition of ammonia to electrolytic hypochlorite.

Experiments made with a view to determining the most efficient ratios of
ammonia gave very surprising results: chlorine to ammonia ratios (by
weight) between 8:1 and 1:2 gave approximately the same germicidal
velocity.[3] The action of the ammonia on the oxidising power of bleach,
as measured by the indigo test, was also found to be disproportionate to
the amount added.

The oxidising action of various mixtures of bleach and ammonia as
measured by the rate of absorption of the available by the organic
matter in the Ottawa River water is shown in Table XXV.


         by Weight.         +-----------+-----------+-----------
                            | 10 Mins.  | 4 Hours.  | 20 Hours.
  Infinity (ammonia absent) |   66.8    |    40.0   |    25.1
    8:1                     |   83.2    |    77.8   |    67.3
    4:1                     |   97.2    |    94.7   |    88.5
  2.7:1                     |   98.3    |    96.5   |    92.8
    2:1                     |   99.8    |    98.2   |    96.2

The 8:1 ratio caused a marked reduction in the rate of absorption of the
chlorine whilst a 4:1 ratio was almost as active as the ratios
containing more ammonia.

At the time when the abnormal results were obtained with ammonium
hypochlorite and mixtures of bleach and ammonia, the phenomenon appeared
to be of scientific interest only and especially so as Rideal had
attributed the obnoxious tastes and odours, sometimes produced by
chlorination, to the formation of chloramines. During the winter of
1915-1916 the price of bleach, however, advanced to extraordinary
heights and the author then determined to try out the process on a
practical scale for the purification of water. A subsidiary plant
pumping about 200,000 Imperial gallons per day (240,000 U. S. A.
gallons) was found to be available for this purpose and the chloramine
process was substituted for the bleach method previously in operation.
The process was commenced by the addition of pure ammonia fort, in the
amount required to give a chlorine to ammonia ratio of 2:1, to the
bleach solutions in the barrels. The results were not in accordance with
those obtained in the laboratory and it was found that the samples of
bleach solutions received for analysis were far below the strength
calculated from the amount of dry bleach used. This experience was
repeated on subsequent days and the deficiency was found to increase on
increasing the ammonia dosage. Solutions of similar concentration were
then used in the laboratory with similar losses, and it was observed
that on the addition of ammonia a copious evolution of gas occurred. An
investigation showed that the ammonia and bleach must be mixed as dilute
solutions and prolonged contact avoided (_vide_ p. 127). Alterations
were accordingly made in the plant and the bleach and ammonia were
prepared as dilute solutions in separate vessels and allowed to mix for
only a few seconds before delivery to the suction of the pumps. This
method of application was instantaneously successful and results equal
to those obtained in the laboratory were at once secured. The dosage was
reduced until the bacteriological results were adversely affected and
continued at values slightly in excess of this figure (0.15 p.p.m.) for
a short period to prove that the process was reliable.

From a consideration of the work of Raschig and Rideal, it appeared that
the most efficient proportions of available chlorine and ammonia would
be two parts by weight of the former to one part of the latter and this
ratio was maintained during the run on the experimental plant. Lower
ratios of chlorine to ammonia were contra-indicated by the laboratory
experiments, which showed that the efficiency was not increased thereby
whilst higher ratios were left for future consideration.

The results obtained on the experimental plant, together with those
obtained on the main plant, where 24 million gallons per day were
treated with bleach only, are given in Tables XXVI, XXVII and XXVIII.
The two periods given represent the spring flood condition and that
immediately preceding it; these are respectively the worst and best
water periods. The results in both cases are from samples examined
approximately two hours after the application of the chemicals.

The cost data were calculated on the current New York prices of bleach
and ammonia.



       |                  |   CHLORITE ALONE.    |       AND AMMONIA.
  1916 | Bacteria   |     | Bacteria  |     |    | Bacteria   |     |    |
       | per cubic  |     | per cubic |     |    | per cubic  |     |    |
       |centimeter. | _B. |centimeter.| _B. |    |centimeter. | _B. |    |
       +-----+------+coli_+-----+-----+coli_|    +------------+coli_|    |
       |Agar | Agar |Index|Agar |Agar |Index| (1)|Agar |Agar  |Index| (1)| (2)
       | 1   |  3   | per | 1   | 3   | per |    | 1   | 3    | per |    |
       | day | days | 100 | day |days | 100 |    | day | days | 100 |    |
       | at  |  at  | cc. | at  | at  | cc. |    | at  | at   | cc. |    |
       | 37° |  20° |     | 37° | 20° |     |    | 37° | 20°  |     |    |
       | C.  |  C.  |     | C.  | C.  |     |    | C.  | C.   |     |    |
  Mar. |   44|   238| 35.7|  4  | 12  |<0.14|0.90|   4 | 12   |0.14 |0.22|0.11
  15-31|     |      |     |     |     |     |    |     |      |     |    |
  April|3,099|14,408|195.5| 32  | 56  | 0.50|1.10|  33 |246   |0.74 |0.25|0.13
  1-19 |     |      |     |     |     |     |    |     |      |     |    |
  Legend: (1) Available chlorine parts per million.
          (2) Ammonia, parts per million.


_Percentage Reduction_

         |    Bacteria   |       |Available|    Bacteria  |       |Available
         |   per cubic   |  _B.  |chlorine |   per cubic  |  _B.  |Chlorine
         |  centimeter.  | coli_ |  parts  |  centimeter. | coli_ | Parts
         +------+--------+ Index |   per   +------+-------+ Index |  per
         | Agar |  Agar  |per 100|million. | Agar | Agar  |per 100|Million.
         |1 Day | 3 Days | cubic |         |1 Day |3 Days | cubic |
         |  at  |   at   |centi- |         |  at  |  at   |centi- |
         |37° C.| 20° C. |meters.|         |37° C.| 20° C.|meters.|
  Mar.   | 90.9 |  95.8  | 99.9+ |  0.90   | 90.0 |  95.0 |   99.7|  0.22
    15-31|      |        |       |         |      |       |       |
  April  | 98.9 |  99.6  | 99.7  |  1.10   | 98.3 |  98.9 |   99.6|  0.25
    1-19 |      |        |       |         |      |       |       |


_Cost Per Million Imperial Gallons_[A]

             | Hypochlorite | Hypochlorite
             |   alone.     | and ammonia.
  Mar. 15-31 |    $1.12     |    $0.46
  April      |     1.26     |     0.54

  [A] Calculated as Bleach at $3.80 per 100 pounds and aqua ammonia (26°
      Bé.) at 5-1/2 cents per pound.

The results were so satisfactory that the author recommended the
adoption of the process on the main chlorinating plant but owing to
conditions imposed by the Provincial Board of Health the process was not
operated until February, 1917.

In place of ammonia fort, aqua ammonia (26° Bé.), containing
approximately 29 per cent of anhydrous ammonia, was used. The material
was first examined by the presence of such noxious substance as cyanides
and found to be very satisfactory.

[Illustration: FIG. 12.--Sketch of Ottawa Chloramine Plant.]

The general design of the plant is shown in Fig. 12. The bleach is mixed
in tank _A_ as a solution containing 0.3 to 0.6 per cent of available
chlorine and delivered to tanks _B_ and _D_, each of which has a
twenty-four-hour storage capacity. The ammonia solution is mixed and
stored in tank _B_ and contains 0.3-0.5 per cent of anhydrous ammonia.
The two solutions are run off into boxes _E_ and _F_ which maintain a
constant head on valves _V_ and _V'_ controlling the head on the
orifices. Both orifices discharge into a common feed box _G_ from which
the mixture is carried by the water injector _J_ through one of
duplicate feed pipes and discharged into the suction well through a
perforated pipe.

As tank _B_ was previously used as a bleach storage tank, the change
from hypochlorite alone to chloramine necessitated very little expense.

The treatment was commenced by gradually increasing the quantity of
ammonia, until a dosage of 0.12 p.p.m. was reached, and constantly
increasing the dosage of bleach, which was formerly 0.93 p.p.m. of
available chlorine. Owing to the restrictions imposed by the Provincial
authorities it has not been possible to maintain a dosage as low as that
indicated as sufficient by the experimental plants results, but some
interesting data have been obtained. Table XXIX shows the results
obtained from February to October, 1917, from the chloramine treatment
at Ottawa and also those obtained with liquid chlorine at Hull where the
same raw water is treated with 0.7-0.8 p.p.m. of chlorine.


           |  _B. coli_   |          |        |                  |  Hull
           |PER 100 C.CMS.|          |        |  DOSAGE P.P.M.   |_B. coli_
    1917   +------+-------+Turbidity.| Colour.+---------+--------+Per 100
           | Raw  |  Tap  |          |        |         |        |c.cms.
           |Water.| Water.|          |        |Chlorine.|Ammonia.|
  Feb.     |  268 |  0.88 |     3    |   40   |   0.57  |  0.05  | ....
  Mar. 1-18|  250 |  0.96 |     4    |   40   |   0.32  |  0.11  | ....
  Mar. 1-31|  643 |  0.43 |     4    |   40   |   0.47  |  0.14  | ....
  April    | 5228 |  0.34 |    31    |   32   |   0.56  |  0.10  | ....
  May      |  162 | <0.08 |     3    |   39   |   0.52  |  0.08  | ....
  June     |  114 | <0.08 |     3    |   41   |   0.51  |  0.08  | ....
  July     |  237 |  0.08 |     5    |   41   |   0.51  |  0.08  | 44.4
  Aug.     |  165 |  0.08 |     4    |   42   |   0.51  |  0.10  | 28.0
  Sept.    |   55 | <0.08 |     6    |   42   |   0.50  |  0.09  | 15.2
  Oct.     |   31 |  0.15 |     5    |   42   |   0.42  |  0.08  |  1.1
  Average  |  211 |  0.22 |     7    |   40   |   0.51  |  0.09  |

At the height of the spring floods the raw water contained 80 p.p.m. of
turbidity and over 500 _B. coli_ per c.cm. but 0.6 p.p.m. of chlorine
and 0.13 p.p.m. of ammonia reduced the _B. coli_ index in the tap
samples to 2.5 per 100 c.cms.; samples taken in Hull on the same day
(treated with 0.7-0.8 p.p.m. of liquid chlorine) gave a _B. coli_ index
of 26.7. Previous experiences in Ottawa has shown that a dosage of
approximately 1.5 p.p.m. of available chlorine is required to reduce the
_B. coli_ index to 2.0 per 100 c.cms. under similar physical and
bacteriological conditions.

During the period of nine months covered by the results in Table XXIX,
only five cases of typhoid fever were reported in which the evidence did
not clearly indicate that the infection had occurred outside the city.
The reduction in the bleach consumed during the same period effected a
saving of $3,200.

During one period of operation the hypochlorite dosage was gradually
reduced to ascertain what factor of safety was maintained with a dosage
of 0.5 p.p.m. of available chlorine and 0.06-0.08 p.p.m. of ammonia. The
results are shown in Diagram VIII. The percentage of samples of treated
water showing _B. coli_ in 50 c.cms. was calculated from the results of
the examination of 4-7 samples daily.

The results showed that it was possible to reduce the chlorine dosage to
0.25 p.p.m. with 0.06 p.p.m. of ammonia without adversely affecting the
bacteriological purity of the tap supply and fully confirmed the
experimental results previously obtained.

The lowest ratio of available chlorine to ammonia used during this test
was approximately 4:1. This is the ratio indicated by a consideration of
the theory of the reaction, and not 2:1 as was formerly stated
(Race[4]). If bleach is represented as Ca(OCl)_{2}, the equation

  Ca(OCl)_{2} + 2NH_{3} = 2NH_{2}Cl + Ca(OH)_{2}

would indicate a ratio of 2:1; but only one molecule of Ca(OCl)_{2} is
produced from two molecules of bleach and the theoretical ratio is
therefore 4:1 (142:34),

  2CaOCl_{2} = CaCl_{2} + Ca(OCl)_{2} and Ca(OCl)_{2} + 2NH_{3} =
   Cl = 142                                              34
                                               = 2NH_{2}Cl + Ca(OH)_{2}.

The chlorine to ammonia ratio is very important because of its influence
on the economics of the process (_vide_ p. 124).

[Illustration: DIAGRAM VIII


All the laboratory and works results that have been obtained in Ottawa
indicate the importance of an adequate contact period. The superiority
of chloramine over other processes is due to the non-absorption of the
germicidal agent and to obtain the same degree of efficiency the contact
period must be increased as the concentration is decreased. For this
reason the best results will be obtained by chlorinating at the entrance
to reservoirs or under other conditions that will ensure several hours
contact. At Ottawa the capacity of the pipes connecting the pumping
station (point of chlorination) and the distribution mains provides a
contact period of one and a quarter hours but even better results would
be obtained if the contact period were increased.

The general results obtained during the use of chloramine at Ottawa in
1917 have shown that the aftergrowths noted during the use of
hypochlorite (see p. 56) have been entirely eliminated and that the _B.
coli_ content of the tap samples from outlying districts has been
invariably less than that of samples taken from taps near to the point
of application of the chloramine. At Denver, Col., where the chloramine
process has also been used, similar results were obtained[5]: four days
after the initiation of the chloramine treatment the aftergrowth count
on gelatine of the Capitol Hill reservoir dropped from 15,000 to 10 per
c.cm. The hypochlorite dosage was cut from 0.26-0.13 p.p.m. of available
chlorine and 0.065 p.p.m. of ammonia added.

_Economics of the Chloramine Process._ The chloramine process was
introduced at Ottawa for the purpose of obtaining relief from the effect
of the high price of bleach caused by the cessation of imports from
Europe in 1915. The results obtained with the experimental plant
indicated that, calculated on the prices current at the beginning of
1917, appreciable economies could be made. Although the reduction in the
chlorine dosage has not been as great as was anticipated, due to the
restrictions previously mentioned, the cost of sterilising chemicals in
1917 was $3,200 less than the cost of straight hypochlorite treatment.

During the latter part of 1917 the relative cost of bleach and ammonia
changed (see Diagram IX).

When calculated on the New York prices for January, 1918, the cost of
chloramine treatment in the United States would be greater than
hypochlorite alone unless a large reduction in the dosage could be
secured by very long contact periods. This condition is only temporary,
however, and the price of ammonia will probably gradually decline as
the plants for fixation of atmospheric nitrogen commence operations and
reduce the demand for the ammonia produced from ammoniacal gas liquor.

In Canada, the market conditions are still (1918) favourable to the
chloramine process: bleach is 25 per cent higher than the U.S.A. product
and ammonia can be obtained for one-half the New York prices.

[Illustration: DIAGRAM IX


_Advantages of the Chloramine Process._ Although the market conditions
may, in some instances, be unfavourable to the chloramine process, the
method possesses certain advantages that more than offset a slight
possible increase in the cost of materials. The taste and odour of
chloramine is even more pungent than that of chlorine but since the
introduction of the process in Ottawa no complaints have been received.
Owing to the reduced dosage, slight proportional fluctuations in the
dosage do not produce the same variations in the amount of free chlorine
which is the usual cause of complaints. A public announcement that the
amount of hypochlorite has been reduced also has a psychological
effect upon the consumers and tends to reduce complaints due to

The most important advantage of the process is the elimination of the
aftergrowth problem. At Denver, where the aftergrowth trouble is
possibly more acute than at any other city on the continent, it was
effectively banished by the use of chloramine. At Ottawa, the sanitary
significance of _B. coli_ aftergrowths is no longer of practical
interest because such aftergrowths have ceased to occur. Whatever may be
their opinion as to the sanitary significance of aftergrowths, all water
sanitarians will agree that the better policy is to prevent their

_Operation of Chloramine Process._ For the successful operation of the
chloramine process, the essential factors are low concentrations of the
hypochlorite and ammonia solutions. The author has found that
hypochlorite containing 0.3-0.5 per cent of available chlorine and
ammonia containing 0.3-0.5 per cent of anhydrous ammonia can be mixed in
a 4:1 or 8:1 ratio without appreciable loss in titre. Solutions of these
concentrations mixed in 4:1 ratio lost only 2-3 per cent of available
chlorine in fifteen minutes and less than 10 per cent in five hours. The
effect of mixing solutions containing 4.35 per cent of available
chlorine and 2.2 per cent of ammonia is shown in Table XXX.

The stability of chloramine is a function of the concentration and the
temperature and in practice it will be found advisable to determine in
the laboratory the maximum concentrations that can be used at the
maximum temperature attained by the water to be treated (cf. Muspratt
and Smith[6]).

According to Raschig[1] two competing reactions occur when ammonia is in

      (1) NH_{2}Cl + NH_{3} = N_{2}H_{4}HCl hydrazine hydrochloride
  and (2) 3NH_{2}Cl + 2NH_{3} = N_{2} + 3NH_{4}Cl.

When the excess of ammonia is large, as on the addition of ammonia fort,
the second reaction predominates and the yield of nitrogen gas is almost
quantitatively proportional to the quantity of available chlorine
present. As ammonium chloride has no germicidal action, and hydrazine a
carbolic coefficient of only 0.24 (Rideal), the formation of these
compounds should be avoided.


Hypochlorite containing 4.35 per cent available chlorine. Ammonia
contained 2.2 per cent NH_{3}

                               |   LOSS OF AVAILABLE CHLORINE AFTER
  Ratio Chlorine to Ammonia by +--------------+----------+-----------
          Weight.              | Few Minutes. |  1 Hour. | 24 Hours.
                               |  Per cent    | Per cent | Per cent
            6:1                |     19       |    19    |    19
            4:1                |     24       |    25    |    25
            2:1                |     45       |    47    |    47
            1:1                |     91       |    91    |    92
            1:2                |     20       |    28    |    65

The dosage of chloramine can be checked by titration of the available
chlorine (see p. 82) immediately after treatment or by the estimation of
the increment in the total ammonia (free and albuminoid). Routine
determinations of the latter made in Ottawa show that practically the
whole (90-95 per cent) of the added ammonia can be recovered by
distillation with alkaline permanganate and that 85-90 per cent is in
the "free" condition.

In operating the chloramine process it is important that the pipes used
for conveying the chloramine solution should be of ample dimensions and
provided with facilities for blowing out the lime that deposits from the

  Ca(OCl)_{2} + 2NH_{3} = 2NH_{2}Cl + Ca(OH)_{2}.

The marked activity of chloramine as a chlorinating agent could be
predicated from its heat of formation, which is 8,230 calories. The
other possible chloramines should be even more active as the heat of
formation of these compounds are:

  Dichloramine          NHCl_{2} -- 36,780 calories.
  Nitrogen trichloride  NCl_{3}  -- 65,330 calories.

Dichloramine is unknown but nitrogen chloride has been prepared and is a
highly explosive yellow oil that decomposes slowly when kept under water
in the ice box. NCl_{3} can be easily prepared by passing chlorine gas
into a solution of ammonium chloride and this process would suggest that
a method might be found of utilising chlorine and ammonia as gases for
the production of nitrogen trichloride as a germicide for water
chlorination. NH_{4}Cl + 3Cl_{2} = NCl_{3} + 4HCl.

The "available" chlorine content of the chloramines is double the actual
chlorine content as each atom of chlorine will liberate two atoms of
iodine from hydriodic acid.

  NH_{2}Cl + 2HI = I_{2} + NH_{4}Cl.

  NCl_{3} + 6HI  = 3I_{2} + NH_{4}Cl + 2HCl.


For the sterilisation of small individual quantities of water such as
are required by cavalry and other mobile troops bleach and acid sulphate
tablets have been usually employed. Such tablets have given fairly
satisfactory results but certain difficulties inherent to these
chemicals have made it desirable to seek other methods.

The subject was investigated by Dakin and Dunham,[7] who first tried
chloramine-T (sodium toluene-_p_-sulphochloramide). It was found that
heavily contaminated waters, and particularly those containing much
carbonates, required a comparatively high concentration of the
disinfectant: 40 parts per million of chloramine-T were necessary in
some cases and such an amount was distinctly unpalatable. By adding
tartaric acid or citric acid the effective concentration could be
reduced to 4 p.p.m. but the mixture could not be made into a tablet
without decomposition and a two-tablet system was deemed undesirable.

Toluene sulphodichloramines were next tried. Excellent bacteriological
results were obtained but the manufacture of tablets again presented
difficulties. When the necessary quantity of dichloramine was mixed with
what were assumed to be inert salts--sodium chloride for example--the
normal slow rate of decomposition was accelerated. The dichloramine, in
tablet form, was also found to be too insoluble to effect prompt

The most suitable substance found by Dakin and Dunham was "halazone" or
_p_-sulphodichloraminobenzoic acid (Cl_{2}N·O_{2}S·C_{6}H_{4}·COOH).
This compound is easily prepared from cheap readily available materials
and was found to be effective and reasonably stable.

The starting point in the preparation of halazone is
_p_-toluenesulphonic chloride, a cheap waste product in the manufacture
of saccharine. By the action of ammonia, _p_-toluene sulphonamide is
produced and is subsequently oxidised by bichromate and sulphuric acid
to _p_-sulphonamidobenzoic acid. This acid, on chlorination at low
temperatures, yields _p_-sulphondichloraminobenzoic acid (halazone). The
reactions may be expressed as follows:

   CH_{3}              COOH                COOH
  / \                 / \                 / \
  | |           -->   | |           -->   | |
  \ /                 \ /                 \ /
   SO_{2}·NH_{3}       SO_{2}·NH_{2}       SO_{2}·NCl_{2}

Halazone is a white crystalline solid, sparingly soluble in water and
chloroform, and insoluble in petroleum. It readily dissolves in glacial
acetic acid from which it crystallizes in prisms (M.P. 213° C.).

The purity of the compound can be ascertained by dissolving in glacial
acetic acid, adding potassium iodide, and titrating with thiosulphate;
0.1 gram should require 14.8 to 14.9 c.cms. of N/10 sodium thiosulphate.
Each chlorine atom in halazone is equivalent to 1 molecule of
hypochlorous acid and the "available" chlorine content is consequently
52.5 per cent or double the actual chlorine content.

  >SO_{2}·NCl_{2} + 4HI = >SO_{2}·NH_{2} + 2HCl + 2I_{2}.

From the bacteriological results given by Dakin and Dunham it would
appear that 3 parts per million of halazone (1.5 p.p.m. available
chlorine) are sufficient to sterilise heavily polluted waters in thirty
minutes and that this concentration can be relied upon to remove
pathogenic organisms.

The formula recommended for the preparation of tablets is halazone 4 per
cent, sodium carbonate, 4 per cent (or dried borax 8 per cent), and
sodium chloride (pure) 92 per cent.

Halazone and halazone tablets, when tested in the author's laboratory on
the coloured Ottawa River water seeded with _B. coli_, have given rather
inferior results. With 1 tablet per quart, over six hours were required
to reduce a _B. coli_ content of 100 per 10 c.cms. to less than 1 per 10
c.cms. Clear well waters gave excellent results and large numbers of _B.
coli_ were reduced to less than 1 per 10 c.cms. in less than thirty
minutes. McCrady[A] has also obtained excellent results with various
strains of _B. coli_ seeded into the colourless St. Lawrence water.

  [A] Private communication.


[1] Raschig. Chem. Zeit., 1907, =31=, 926.

[2] Rideal. S. J. Roy. San. Inst., 1910, =31=, 33-45.

[3] Race. J. Amer. Waterworks Assoc., 1918, =5=, 63.

[4] Race. Eng. and Contr., 1917, =47=, 251.

[5] Contract Record. Aug. 15, 1917, 696.

[6] Muspratt and Smith. J. Soc. Chem. Ind., 1898, =17=, 529.

[7] Dakin and Dunham. Brit. Med. Jour., 1917, No. 2943, 682.



The object of adding chlorine or chlorine compounds to water is for the
purpose of destroying any pathogenic organisms that may be present. In a
few instances some collateral advantages are also obtained but, in
general, no other object is aimed at or secured.

Chlorination does not change the physical appearance of water; it does
not reduce or increase the turbidity nor does it decrease the colour in
an appreciable degree.

The chemical composition is also practically unaltered. When bleach is
used there is a proportionate increase in the hardness but the amount is
usually trifling and is without significance. During 1916 when the
Ottawa supply was entirely treated with bleach at the rate of 2.7 parts
per million (0.92 p.p.m. of available chlorine) the average increase in
the total hardness as determined by the soap method was 2.5 parts per

When chlorine is added to prefiltered water, as an adjunct to
filtration, an increase in the number of gallons filtered per run has
been noted at some plants. This increase is not so great with rapid as
with slow sand filters but in some instances it has led to appreciable

Walden and Powell[1] of Baltimore, found that the addition of a quantity
of bleach equal to approximately 0.50 p.p.m. of available chlorine
enabled the alum to be reduced from 0.87 to 0.58 grain per gallon. The
percentage of water used in washing the filters was also reduced, from
4.1 per cent to 2.9 per cent, whilst the filter runs were increased on
the average by one hour and ten minutes. The net saving in coagulant
alone amounted to 30 cents per million gallons.

Clark and De Gage[2] found that the use of smaller amounts of coagulant
during the period of combined disinfection and coagulation resulted in
an increase of nearly 25 per cent in the quantity of water passed
through the filter between washings, and also in a material reduction of
the cost of chemicals, which averaged $2.62 per million gallons for
combined disinfection and coagulation as against $4.86 for coagulation
alone. The water used in these experiments was obtained from the
Merrimac River at Lawrence.

The effect of hypochlorite on the reduction of algæ growths on slow sand
filters was first noticed by Houston during the treatment of the Lincoln
supply in 1905. Two open service reservoirs were fed with treated water
and were themselves dosed from time to time. "Previous to 1905 they
developed seasonally most abundant growths, but during the hypochlorite
treatment it was noticed that they remained bright, clear, and
remarkably free from growths" (Houston[3]).

Ellms,[4] of Cincinnati, has also noted the effect of hypochlorite on
algæ. When the bleach was added to the coagulated water the destruction
of the plankton was not as satisfactory as had been anticipated and it
was found that large doses destroyed the coating of the sand particles
and rendered the filters less efficient. The use of bleach in the
filtered water basin was more successful and cleared it of troublesome

In 1916, during the treatment of the London Supply with bleach (dosage
0.5 p.p.m. of available chlorine), Houston made further observations on
this point. The Thames water, taken at Staines, had previously been
stored for considerable periods in reservoirs, but this necessitated
lifting the water by pumps which consumed large quantities of coal that
were urgently needed for national purposes. As a war measure, the
storage was eliminated and the water treated with hypochlorite at
Staines and allowed to flow by gravitation to the various works where
the slow sand filters are situated. The treatment resulted in a marked
reduction in the growths of algæ, the reduction in the area of filters
cleaned in 1916 (June to September) as compared with 1915 being as

      Filter Works.             Reduction
  Grand Junction (Hampton)          6
  Grand Junction (Kew)             43
  East London (Sunbury)            30
  Kempton Park                     33
  West Middlesex                   56

A portion of this reduction can probably be attributed to the
elimination of storage.

Chlorination, by decreasing the load on filter beds, has enabled the
rate of filtration to be increased in some cases. This increased
capacity, which would otherwise have necessitated additional filter
units, has been obtained without any further capital outlay. At
Pittsburg (Johnson[5]) the rate of filtration, after cleaning, was
increased 250,000 gallons each hour until the normal rate was reached;
restored beds were maintained at a 250,000 gallon rate for one week.
After the introduction of chlorination it was found possible to increase
the rates more rapidly without adversely affecting the purity of the
mixed filter affluents.

_Hygienic Results._ Evidence as to the actual reduction of the number of
such pathogenic germs as _B. typhosus_ in water supplies by chlorination
is most readily found in the death rates from typhoid fever in cities
that have no other means of water purification. In some cases this
evidence is necessarily of a circumstantial nature; in others it is
definite and conclusive.

Some of the earlier results of the effect of chlorination on typhoid
morbidity and mortality rates were compiled by Jennings[6] and others
have been published by Longley.[7] These data have been brought up to
date in Table XXXI and other statistics added.



      City.       |  Commenced  | BEFORE USING.| AFTER USING. |
                  |Chlorination.+-------+------+-------+------+ Percentage
                  |             |Period.|Rate. |Period.|Rate. | Reduction.
  Baltimore       |June 1911    |1900-10| 35.2 |1912-15| 22.2 |    36
  Cleveland       |Sept. 1911   |1900-10| 35.5 |1912-16|  8.2 |    77
  Des Moines      |Dec. 1910    |1905-10| 22.7 |1911-13| 13.4 |    41
  Erie            |Mar. 1911    |1906-10| 50.6 |1912-14| 15.0 |    70
  Evanston, Ill.  |Dec. 1911    |1908-11| 29.0 |1912-13| 14.5 |    50
  Jersey City     |Sept. 1908   |1900-17| 18.7 |1909-16|  8.4 |    55
  Kansas City, Mo.|Jan. 1911    |1900-10| 42.5 |1911-16| 14.2 |    66
  Omaha, Neb.     |May 1910     |1900-09| 22.5 |1911-16| 10.6 |    53
  Trenton         |Dec. 1911    |1907-11| 46.0 |1911-14| 28.7 |    35
  Montreal        |Feb. 1910    |1906-10| 40.0 |1911-16| 25.0 |    37
  Toronto         |Apr. 1911    |1906-10| 31.2 |1912-16|  7.8 |    75
  Ottawa          |Sept. 1912   |1906-10| 34.0 |1913-17| 17.0 |    50

The figures given in this table show the effect of chlorination only; no
other form of purification was used during the periods given, except at
Toronto where a portion of the supply has been subjected to filtration.

It will be seen that since chlorination was adopted the typhoid death
rates have been reduced by approximately 50 per cent and that the
averages for the period after treatment are almost invariably less than
20 per 100,000, a figure that a few years ago was regarded as
satisfactory. The average death rate for the last available year is 11
per 100,000, a result that is even more satisfactory and exceeds the
anticipations of the most optimistic of sanitarians.

A portion of the reduction in the typhoid rates is no doubt due to
improvements in general sanitary conditions but the reduction is much
greater than can be accounted for in that manner alone and in many
cases there was a sharp decline immediately following the commencement
of chlorination.

In a few instances there is evidence that chlorination has reduced the
typhoid rates of cities previously supplied with filtered water. Diagram
X, drawn from data supplied by Dr. West, of the Torresdale Filtration
Plant, shows the effect of disinfecting the filter effluents at

[Illustration: DIAGRAM X


During the years 1909-10-11, when practically the whole of the city
supply was filtered, the average typhoid death rate was 18, but when the
water was also chlorinated, in 1914-15-16, the rate was only 7, a
reduction of 61 per cent.

The figures in Table XXXII show that the Torresdale filters, during
1915-16 were unable to adequately purify the water and that chlorination
was necessary.



      |         |         |          | BACTERIA PER CUBIC CENTIMETER.
      | Oxygen  |         |          +-----------------+----------------
      |Consumed.| Colour. |Turbidity.|    Untreated.   |    Treated.
      |         |         |          +---------+-------+---------+------
      |         |         |          |Gelatine.| Agar. |Gelatine.| Agar.
  1915|  1.70   |   12    |   0.6    |   141   |  30   |   28    |  14
  1916|  1.90   |   12    |   Nil.   |    88   |  23   |   38    |  11
      |               _B. coli communis_               |
      |            Per Cent Positive Tests.            |
      |    Untreated.        |         Treated.        | Added Chlorine
      +----------+-----------+------------+------------+   Parts Per
      |10 c.cms. |  1 c.cm.  | 10 c.cms.  |  1 c.cm.   |   Million.
  1915|    66    |     24    |     5      |    0.3     |      0.18
  1916|    49    |     16    |     7.4    |    1.9     |      0.15

In Diagram XI the typhoid death rates of Columbus, Ohio, and New Orleans
are shown to exemplify conditions that have not been improved by
chlorination. The endemic condition of typhoid in Columbus was brought
to an abrupt conclusion by the installation and operation of the
softening and filter plant in September, 1908, and no further reduction
followed the introduction of chlorination in December, 1909.

In New Orleans the typhoid rate decreased on the inception of the new
water works system in 1909 and again after the installation of the
Carrollton filters in 1912. The product of the filtration plants has
always been above suspicion but aftergrowths occasionally developed and
the bacterial count then exceeded the United States Treasury standard.
To overcome this difficulty, hypochlorite was used in 1915, but, as was
anticipated, it had no effect on the typhoid rate. The high rate in New
Orleans is largely due to outside cases received for hospital treatment
and to other circumstances beyond the control of the water and sewerage

In all the examples previously cited, the evidence as to the effect of
chlorination on typhoid mortality rates is circumstantial but, taken as
a whole, it is fairly conclusive. In the examples to be considered next
the evidence is more direct.

[Illustration: DIAGRAM XI


One of the most conclusive experiments as to the beneficial effect of
chlorination is that reported by Young[8] of Chicago. The water supply
of Chicago was obtained from Lake Michigan by means of intake pipes and
pumped to various parts of the city. The distribution system was divided
into four districts and, although there was a certain amount of mixing
along the borders, the water supplied to each district was substantially
separate. The rapid and progressive decline in the typhoid rate of
Chicago (from 19 in 1900 to 10.8 in 1911) subsequent to the diversion
of the city sewage from the lake, led to the assumption that water-borne
typhoid had ceased to be of any moment. Early in 1912, however,
permission was secured to chlorinate the supply of one district (No. 1)
and the treatment was continued until December when the solutions
commenced to freeze. Diagram XII shows the effect of the treatment on
the autumnal increase in District No. 1 as compared with the other three
districts. The autumnal increase was calculated from the excess of
typhoid incidence for July to November inclusive, over that for February
to June inclusive.

[Illustration: DIAGRAM XII


These results demonstrate in a most striking manner the beneficial
effect of chlorination. The general conditions, with the exception of
the raw water supply, were approximately the same in all four districts.
Diagram XIII shows that the raw water supply of District No. 1 was
slightly worse than any of the others, 21.8 per cent of the samples from
District No. 1 containing _B. coli_ in 1 c.cm. as compared with 21.0 per
cent in the most polluted supply of the other districts.

[Illustration: DIAGRAM XIII


The results obtained at Ottawa are also conclusive. Following two
epidemics of typhoid fever in 1911 and 1912, caused by breaks in the
intake pipe, hypochlorite treatment was commenced and has been in
continuous operation until February, 1917, when chloramine treatment was
substituted. The dosage has been so regulated as to assure a high degree
of purity at all times in the water delivered to the mains and as
evidence of this it might be mentioned that the average _B. coli_ index
(calculated by Phelps' method) for the years 1916 and 1917 was only 0.27
per 100 c.cms. The typhoid rates for the five years preceding the
epidemic years and for a similar subsequent period are given in Diagram

[Illustration: DIAGRAM XIV


The diagram shows that there has been a constant reduction in the city
typhoid rate since the last severe epidemic with the exception of the
year 1915. The high rate of that year was caused by a localised epidemic
started by polluted well water and spread by flies from an unsewered
area. This outbreak was the cause of about seven deaths registered
during that year (population 100,000).

The objection might be raised that if the reduction of the typhoid rate
were due to the water treatment, the decline should have been abrupt and
not a gradual one. It is probable that there has been practically no
water-borne typhoid in the city since chlorination was commenced but
this fact is masked by cases from other sources. During 1911 and 1912
over 3,500 cases of typhoid were reported, of which an appreciable
number would become carriers for various periods of time. As these
carriers decreased the number of cases infected by them would also
decrease and so account for a gradually declining death rate.

It might be further objected that the reduced typhoid rate is due to a
general improvement in the sanitary conditions. If the death rate from
causes other than typhoid can be regarded as a measure of the general
sanitary conditions it is obvious from the data in Table XXXIII that the
improvement in the typhoid rate is immeasurably greater than can be
ascribed to that cause.


                            |  RATE PER 100,000 |      PERCENTAGE
          Cause.            +---------+---------+-----------+---------
                            | 1908-12 | 1913-17 | Reduction | Increase
  Total[A]                  |  14.90  |  14.78  |    1.2    |   ...
  Typhoid, total            |  34[B]  |  17     |   50.0    |   ...
  Typhoid, city             |  26[B]  |   8     |   69.2    |   ...
  Pneumonia                 | 100     | 107     |    ...    |   7.0
  Tuberculosis              | 133     | 138     |    ...    |   3.7
  Diarrh[oe]a and Enteritis | 139     | 128     |    7.9    |   ...
    under 2 years           |         |         |           |

  [A] Rate per 1,000.

  [B] 1906-10, epidemic years 1911-12 excluded.

One further objection might be made: that the raw water was not infected
during 1913-17 or infected to a smaller extent than during the previous
period. Attempts to isolate _B. typhosus_ from the raw water have
invariably been futile but their presence in 1914 might be inferred from
the fact that during the latter part of the summer of that year an
epidemic of typhoid fever occurred at Aylmer, a village that discharges
its sewage into the Ottawa River about six miles above the Ottawa
intake. Hull, situated on the opposite bank of the river and having a
population of 20,000, takes its water supply from the same channel that
supplies Ottawa but at a point a few hundred feet further down stream.
During November and December, 1914, some 200 cases of typhoid fever
(incidence 1,000 per 100,000) occurred in Hull as compared with 28 in
Ottawa. As the Ottawa intake is situated between the Hull intake and the
outlet of the Aylmer sewer it is incredible that the Ottawa raw water
was not also infected.

In 1916 a liquid chlorine plant was installed in Hull, but in 1917,
owing to an accident, it was out of commission for a short period and at
least 100 cases of fever developed during the following month. During
the same period only two cases were reported in Ottawa and of these one
was obviously contracted outside the city.

In view of the preceding facts it must be granted that the improvement
in the typhoid rate of Ottawa can be definitely attributed to an
improvement in the water supply caused by chlorination.

The efficacy of chlorination to prevent and check epidemics of
water-borne typhoid has never been doubted. Innumerable instances could
be cited in which the prompt treatment of large public supplies has
promptly checked outbreaks that threatened to assume serious proportions
and there is no doubt that the extremely low typhoid morbidity rate on
the Western Front of the European battlefield is partially due to the
extensive and rigorous chlorination measures that have been instigated.
Prophylactic vaccination and the prompt isolation of typhoid carriers
have largely contributed to the wonderful results obtained but due
credit must also be given to the systematic purification and treatment
of water supplies. Similar results have been obtained at training camps
in Canada and in other countries by effective treatment with either
liquid chlorine or hypochlorite.

Since the inception of water chlorination in America in 1908, the merit
of the method has been very generally recognized throughout the
Continent but was regarded with scepticism in Europe, except as a
temporary expedient, until the results obtained by the military forces
compelled more general recognition. Before the war, chlorination of
water supplies in England was only practised in a few isolated and
relatively unimportant instances; in 1917, practically the whole supply
of London was chlorinated and at Worcester a similar treatment has been
recommended to enable the slow sand filters to be operated at higher
rates without reducing the quality of the water supplied to the

_Use and Abuse of Chlorine._ Inasmuch as chlorination has no beneficial
effect on water except the reduction of the bacterial content it should
be used for this purpose only and under such conditions as permit the
operations to be under full control at all times. The supplies that can
be most efficiently and safely treated are those that are relatively
constant in chemical composition and bacterial pollution. Changes in
volume can be dealt with by automatic apparatus but sudden changes in
organic and bacterial content require a change of dosage that cannot be
made by any mechanical appliance. Long experience and accurate
meteorological records may in some cases enable those in charge of
chlorination plants to anticipate changes in the conditions of the water
supply, but it is always preferable to provide a positive method of
preventing sudden changes by using chlorination merely as an adjunct to
other processes of purification. Unpurified waters that are
objectionable on account of their bacterial content only are very rare,
as the cause that produces the bacterial pollution usually produces
other conditions that are equally objectionable though not so dangerous
to health. Sudden storms in summer, or sudden thaws in winter, usually
cause large increments in turbidity accompanied by soil washings that
often carry appreciable quantities of fæcal matter into surface water
supplies. Lake supplies often suffer in the same manner and sewage,
which during normal conditions is carried safely away from water
intakes, obtains access to the supply. If the dosage is maintained at a
level sufficiently high to meet these abnormal conditions, complaints as
to taste and odour would ensue, and in general, such a practice is
impossible. Some supplies have been chlorinated successfully for years
but the principle of using chlorination as the first and last line of
defence cannot be recommended. Success can only be obtained by eternal
vigilance and the responsibility for results is more than water works
officials should be called upon to assume.

Chlorination is an invaluable adjunct to other forms of water
purification and it is not improbable that, in the future, filter plants
will be designed to remove æsthetic objections at the lowest possible
cost and that chlorination will be relied upon for bacterial reduction.
Chlorination is the simplest, most economical, and efficient process by
which the removal of bacteria can be accomplished and there is no valid
reason why it should not be used for that purpose.

The popularity of this process has suffered through the efforts of over
zealous enthusiasts who have been unable either to recognize its
limitations or to appreciate the fact that a domestic water supply
should be something more than a palatable liquid that does not contain
pathogenic organisms. Every system of water purification has its limited
sphere of utility and chlorination is no exception to the rule.


[1] Weldon and Powell. Eng. Rec., 1910, =61=, 621.

[2] Clark and De Gage, 41st Annual Rpt. Mass. State B. of H. 1910.

[3] Houston. 12th Research Rpt. Metropolitan Water Board, London.

[4] Ellms. Eng. Rec., 1911, =63=, 388.

[5] Johnson. Eng. Rec., 1911, =64=, No. 16.

[6] Jennings. 8th Inter. Congr. Appl. Chem., =26=, 215.

[7] Longley. J. Amer. Waterworks Assoc., 1915, =2=, 679.

[8] Young. J. Amer. Public Health Assoc., 1914, =4=, 310.



REAGENTS. 1. Tolidine solution. One gram of _o_-tolidine, purified by
recrystallization from alcohol, is dissolved in 1 litre of 10 per cent
hydrochloric acid.

2. Copper sulphate solution. Dissolve 1.5 grams of copper sulphate and 1
c.cm. of concentrated sulphuric acid in distilled water and dilute the
solution to 100 c.cms.

3. Potassium bichromate solution. Dissolve 0.025 gram of potassium
bichromate and 0.1 c.cm. of concentrated sulphuric acid in distilled
water and dilute the solution to 100 c.cms.

PROCEDURE. Mix 1 c.cm. of the tolidine reagent with 100 c.cms. of the
sample in a Nessler tube and allow the solution to stand at least five
minutes. Small amounts of free chlorine give a yellow and larger amounts
an orange colour.

For quantitative determination compare the colour with that of standards
in similar tubes prepared from the solutions of copper sulphate and
potassium bichromate. The amounts of solution for various standards are
indicated in the following table:


      Chlorine.       |    Solution of    |     Solution of
                      |  Copper Sulphate. | Potassium Bichromate.
  Parts per million.  |       c.cms.      |       c.cms.
       0.01           |        0.0        |        0.8
        .02           |        0.0        |        2.1
        .03           |        0.0        |        3.2
        .04           |        0.0        |        4.3
        .05           |        0.4        |        5.5
        .06           |        0.8        |        6.6
        .07           |        1.2        |        7.5
        .08           |        1.5        |        8.7
        .09           |        1.7        |        9.0
        .10           |        1.8        |       10.0
        .20           |        1.9        |       20.0
        .30           |        1.9        |       30.0
        .40           |        2.0        |       38.0
        .50           |        2.0        |       45.0

[Illustration: DIAGRAM XV]

[Illustration: DIAGRAM XVI]


  Adams, 66, 82

  Bassenge, 9
  Baxter, 4
  Berge, 9
  Berthollet, 1
  Bevan, 29
  Bonjean, 36
  Bray, 24
  Breteau, 26
  Bucholtz, 5

  Catlett, 99
  Clark, 53, 133
  Comte, 47
  Cross, 29
  Cruikshank, 3

  Dakin, 22, 28, 129
  Darnall, 89
  Davy, 1
  DeGage, 53, 133
  DeMorveau, 3
  Dibden, 6
  Diénert, 48
  Dienheim-Brochoki, 105
  Dowell, 24
  Dunbar, 6
  Dunham, 129
  Dupré, 5
  Dusch, 4

  Ellms, 34, 83, 84, 133
  Elmanovitsch, 36
  Elsner, 6
  Evans, 84

  Faraday, 103
  Fischer, 16
  Forcrand, 103
  Fuller, G. W., 11

  Gascard, 47
  Griffen, 17, 79

  Haberkorn, 5
  Hale, 80, 100
  Harrington, 34, 65
  Hauser, 83, 84
  Hedallen, 17, 79
  Heise, 36
  Henry, 2
  Hermite, 5
  Hewlett, 9
  Hooker, 72
  Horrocks, 48
  Houston, 8, 59, 71, 133
  Hsu, 21

  Jackson, 91, 99
  Jakowkin, 26
  Jennings, 135
  Johnson, 11, 134
  Jordan, H. E., 57

  Kanthack, 6
  Kauffman, 9
  Kellerman, 7
  Kershaw, 107
  Kienle, 65, 66, 90, 99
  Kimberly, 7
  Klein, 5
  Koch, 4
  Kolessnikoff, 16
  Kranejuhl, 7
  Kuhn, 5
  Kurpjuivat, 7

  Landolt, 105
  Langar, 10
  Laroche, 47
  Lavoisier, 1, 15
  Leal, 16
  Lehmann, 101
  LeRoy, 83
  Letton, 64
  Longley, 43, 135
  Lunge, 105
  Lyon, 24

  Marshall, 102
  Massy, 48
  Meadows, 112, 114
  McCrady, 130
  McGowan, 8
  McLintock, 5
  Mohler, 31
  Mohr, 79
  Moor, 9
  Muspratt, 126

  Nesfield, 8, 89
  Nissen, 30
  Norton, 21
  Novey, 23
  Noyes, 24

  Ornstein, 90
  Orticoni, 36

  Pedler, 103
  Percy, 3
  Pettenkofer, 101
  Phelps, 7, 17, 82
  Pitcher, 112
  Plucker, 10
  Powell, 132
  Pratt, 7
  Proskauer, 6, 16

  Rabs, 110
  Race, 36, 110, 116
  Raschig, 115
  Rickard, 108
  Rideal, E. K., 84
  Rideal, S., 6, 9, 21, 22, 60, 115, 116
  Roscoe, 5
  Roozeboom, 103
  Rouquette, 36
  Ruffer, 5

  Sandman, 56
  Scheele, 1, 15
  Schroder, 4
  Schuder, 10
  Schumacher, 7
  Schumburg, 10
  Schwann, 4
  Schwartz, 7
  Semmelweiss, 4
  Sickenberger, 9
  Smeeton, 53
  Smith, 126

  Tennant, 2
  Thomas, 53, 56
  Thresh, 87
  Tiernan, 92
  Tolman, 111
  Traube, 9

  Valeski, 36
  Von Loan, 90

  Walden, 132
  Walker, 87
  Wallace, 92
  Wallis, 83
  Warouzoff, 16
  Watt, 2, 3, 15, 106
  Webster, 5, 105
  Wesbrook, 31, 44, 53
  West, 91, 99, 136
  Whittaker, 31
  Winkler, 84
  Winogradoff, 16
  Winslow, 110
  Woodhead, 7
  Woolf, 5

  Young, 138

  Zirn, 6


  Absorption of chlorine by water, 35
  Abuse of chlorination, 144
  Acids, effect of, 19, 21
  Action of chlorine, 16
  Admixture, effect of, 39
  Aftergrowths, 55
    accelerated growth, 58
    _B. coli_ in, 57
    effect of liquid chlorine, 99
    views as to nature of, 56
  Algæ, effect of chlorine on, 133
  Alkalies, effect of, 19, 20
  Allen-Moore cell, 111
  Ammonia, and chlorine, 24
    and sodium hypochlorite, 114
    effect on bleach, 21
    effect on oxidising action, 21
    soda process, 2
  Antichlors, 86
  Antiseptics, early work on, 3
    chlorine as an, 50
  Application of chlorine, point of, 43
  Auto-suggestion, 62

  _B. choleræ_ suis, 31
  _B. cloacæ_, 31
  _B. coli_, aftergrowths, 57
    in sewage, 6, 7
    in water, 9, 28, 31
    standard, 46
    viability of, 52, 55
  _B. cuticularis_, 53
  _B. fæcalis alkaligenes_, 31
  _B. enteritidis_, 31
  _B. enteritidis sporogenes_, 53
  _B. lactis ærogenes_, 31
  _B. subtilis_, 53
  _B. tetani_, 9
  _B. typhosus_, 9, 10, 30, 31
  Bacteria surviving chlorination, 50
    aftergrowths, 55
    nature of, 53
    spores, 57
  Benzidine, 83
  Bleach, analysis of solution, 79
    as deodourant, 3, 6
    as sewage disinfectant, 6, 7
    at Adrian, 11
    at Boonton, 11, 16
    at Bubbly Creek, 11
    composition, 14
    decomposition of, 25
    discovery, 2
    germicidal velocity, 20, 21
    hydrolysis, 18, 19
    production, 3
    stability of, 17
    toxic action, 22
    treatment, 72
      control of, 78
      cost, 86
      dosage regulation, 75
      in France, 78
      losses in, 81
      mixing tank, 73
      plant design, 72
      storage tank, 75
  Brest experiments, 5

  Carnallite, 1
  Chicago, typhoid rate, 138
  Chloramine, 114
    at Denver, 124, 126
    at Ottawa, 28, 116
    contact period, 123
    cost of, 124
    decomposition of, 126
    experimental results, 119
    germicidal power, 116
    operation of process, 126
    plant design, 120
    preparation of, 115
    ratio of chlorine and ammonia, 116, 122
    tastes and odours, 28, 64, 117
    toxic action, 22, 29
  Chlorides, effect of, 20
  Chlorine, and ammonia, 24, 25
    discovery of, 1
    disinfection, effect of pabulum, 4
      general reactions, 28
    hydrate, 103
    detection of, 81
    effect on flowers, 68
    estimation of, 81
    in sanitary work, 4
    medicinal dose, 67
    oxygen equivalent, 23
    liquid, 89
      advantages of, 97
      cost of treatment, 101
      disadvantages of, 101
      germicidal efficiency, 99
      machines, 89
    peroxide, 9
    water, 102
      corrosion of pipes, 69
      damage to seeds, 68
      decomposition of, 15
      heat of formation, 27
  Chlorometer, 84
  Chloros, 8
  Chlorozone, 105
  Colour, effect on dosage, 33
  Columbus, typhoid rates, 137
  Complaints, 62
  Contact period, effect on dosage, 44
    effect on taste, 43
    usual practice, 45
  Cost of bleach plant, 85
    bleach treatment, 86
    liquid chlorine treatment, 101
  Crossness experiments, 5

  Dayton cell, 107
  DeChlor filters, 87
  Denver, chloramine treatment, 124, 126
  Dichloramine, 128
  Disinfectants, 50
  Disinfection, early views of, 3
  Dosage, 30
    determination of, 46
    effect of, admixture, 39
      colour, 33
      contact period, 43
      initial contamination, 32
      light, 45
      oxidisable matter, 32
      standard of purity, 30, 32
      temperature, 34, 36
      turbidity, 45
    for military work, 48
    regulation of bleach, 75
    relation to oxygen absorbed, 36
    tanks, 75

  Eau de Javelle, 3, 47
  Electrical conductivity of treated water, 70
  Electrolysed sea water, 5
  Electrolytic hypochlorite, 2, 104
    Bradford, 5
    Brest, 5
    Brewster, 6, 105
    cost of, 113
  Electrolytic hydrochlorite, Crossness, 5
    discovery of, 3
    diaphragm cells, 110
    early use of, 5
    efficiency of, 109
    Havre, 5
    non-diaphragm cells, 106
  Electrozone, Brewster, 6
    Maidenhead, 6
    Tonetta Creek, 6

  Filter effluents, chlorination of, 34
  Filters, effect on beds, 60
    effect on runs, 132
  Fish, effect on, 8, 67, 68

  Germicidal velocity, effect of acids, 21
    alkalies, 20
    ammonia, 21
    chlorides, 20
  Guildford, chlorination at, 9

  Haas and Oettel cell, 108
  Halazone, 128
  Hardness, effect of chlorine on, 132
  Havre experiments, 5
  Hermite fluid, 5
  Hexamethyl-_p_-aminotriphenylmethane, 83
  Historical, 1
  Hooghly River, 7
  Hydrazine, 126
  Hydrogen peroxide, 24
  Hydrolysis of hypochlorites, effect of, acids, 19
    alkalies, 19
    chlorides, 20
  Hygienic results, 134
  Hypochlorous acid, 17
    decomposition of, 24, 25, 26
    hydrolytic constant, 18

  Initial contamination, effect on dosage, 32
  Intestinal organisms, viability of, 52
  Iodoform taste, 65
  Iron salts, effect on dosage, 33

  Jersey City, court case, 11, 16

  Kellner cell, 108

  Labarraque solution, 105
  Leavitt-Jackson machine, 91
  Leblanc process, 2
  Light, effect on dosage, 45
  Lincoln, chlorination at, 8, 59
  Liquid chlorine, advantages of, 97
    and tastes, 65
    effect of temperature on, 95
    machines, 89
      dry feed, 94
      E. B. G. Co., 91
      Leavitt-Jackson, 91
      operation of, 95
      Wallace and Tiernan, 92
  L'Orient, experiments at, 5

  M. agilis, 53
  Maidstone, use of bleach at, 8
  Margin of safety for taste and odour, 64
  Material for bleach plants, 74
  Military work, bleach method for, 78
    chlorine water, 103
    dosage for, 47, 48, 78
    early European, 10
    liquid chlorine, 102
    typhoid reduction, 143
    use of chlorine in, 8
  Mixing tank for bleach, 73
  Moisture, effect on chlorine gas, 16
  Montreal, dosage at, 34
    electrolytic cells, 112

  Nascent oxygen hypothesis, 17
  Nelson cell, 111
  Neva River, 36
  New Orleans, typhoid rates, 137
  New York, bacteria surviving treatment, 53
    bleach efficiency, 100
    liquid chlorine plant, 97
  Nitrites, effect on dosage, 33
  Nitrogen trichloride, 24, 128

  Odours, effect of contact period on, 43
    nature of, 63
  Ottawa, aftergrowths at, 57
    bleach plant efficiency, 100
    chloramine plant, 120
    chloramine results, 121
    sludge trouble, 65
    typhoid rates, 140
  Oxidisable matter, effect on dosage, 32, 36
  Oxychloride, Guildford, 9
    Middlekerke, 9
    Ostend, 9
  Ozone, 24

  Philadelphia and chlorination, 136
  Pipe corrosion, 69
    Pittsburg report, 71
  Plumbo solvency, 71
  _P. mirabilis_, 31
  Potassium permanganate, 23
  Puerperal fever in Vienna, 4
  Pumps, for admixture, 41

  Red Bank, sewage disinfection at, 7
  Reversed ratio of counts, 54

  Sewage disinfection at Baltimore,
    Berlin, 7
    Boston, 7
    Brewster, 6
    Hamburg, 6
    Maidenhead, 6
  Sludge, as cause of complaints, 65
  Sodium bisulphite, 86
  Sodium chloride, deposits, 1
    decomposition of, 106
  Sodium hypochlorite, 105
    decomposition of, 26
    effect of ammonia on, 21
    hydrolysis of, 26
  Sodium thiosulphate, 87
  Standard of purity, 30
  Storage tanks, 75
  Sulphuretted hydrogen, 33
  Sylvine, 1

  Tannin, 67
  Tastes, effect of contact period on nature of, 63
  Temperature, effect on absorption of chlorine, 35, 38
    bleach deterioration, 72
    dosage, 34, 36
    germicidal velocity, 38
    pressure of liquid chlorine, 96
    tastes and odours, 66
  Thermophylic organisms, 54
  Tolidine, 82
  Toxic action of chlorine, 22, 29
  Turbidity, effect on dosage, 45
    effect of chlorine on, 132

  Use of chlorination, 144

  Water mains, disinfection of, 8
  Well water, 7
  Worcester, chlorination at, 11
  Worthing experiments, 5

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  |                   Additional Transcriber's Notes:                   |
  |                                                                     |
  | * The original symbol in Table XIV (a circled +) has been changed   |
  |   to [¤].                                                           |
  | * Some formulas have been spaced out for better readability.        |
  | * Some minor typographical errors have been corrected (including    |
  |   indicators for references and missing diacritical marks from      |
  |   German words).                                                    |
  | * In-line multi-line formulas have been changed to in-line single-  |
  |   line formulas, if necessary with the addition of brackets.        |
  | * Inconsistencies in spelling, hyphenation, lay-out or formatting   |
  |   have not been corrected, except in the following cases:           |
  |   * Bassenege, Schemmelweiss, Langar and Kanthdack in the name index|
  |     have been changed to Bassenge, Semmelweiss, Langer and Kanthack |
  |     as in the text.                                                 |
  |   * Heisse, Jordon, Tonnetta Creek and Horrock's have been changed  |
  |     to Heise, Jordan, Tonetta Creek and Horrocks's, respectively, as|
  |     elsewhere in the text.                                          |
  |   * Page 35: N^{1} and N^{2} in formula changed to N_{1} and N_{2}  |
  |     as elsewhere.                                                   |
  |   * Page 7: Hadallen changed to Hedallen as elsewhere in the text.  |
  | * Changes made to the text:                                         |
  |   * Page 17: --> changed to <=> in chemical formula as described in |
  |     the text.                                                       |
  |   * Page 26: H^{.} + HCO_{3} changed to H^{.} + HCO_{3}'.           |
  |   * Page 26: chlor-ions changed to chlorine ions.                   |
  |   * Page 54: Gention Violet changed to Gentian Violet.              |
  |   * Page 103: Footnote marker [11] inserted (missing in original).  |
  | * The author called Kurpjuivut, Kurjuivut and Kurpjuivat in various |
  |   places in the text is probably called Kurpjuweit. The author      |
  |   called Schumburg and Schumberg in the text is called Schumberg.   |
  |   The book contains references to both Zaleski and Elmanovitsch     |
  |   and Valeski and Elmanovitsch; Zaleksi is probably correct.        |
  | * Other remarks:                                                    |
  |   * Footnote on Page 119: fraction unclear in the original,         |
  |     presented here as 5-1/2.                                        |
  |   * Page 134: affluents should probably be effluents.               |
  |   * In the original work, there is no TABLE XXII between TABLE XXI  |
  |     and XXIII.                                                      |

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