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WORKS OF

JOSEPH RACE

PUBLISHED BY

JOHN WILEY & SONS, Inc.

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, 51⁄4 × 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, 51⁄4 × 8, 12 figures and 16 diagrams. Cloth, $1.50 net.

Chlorination of Water

BY

JOSEPH RACE, F.I.C.

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

FIRST EDITION

NEW YORK
JOHN WILEY & SONS, Inc.
London: CHAPMAN & HALL, Limited
1918

Copyright, 1918
BY
JOSEPH RACE, F.I.C.

PRESS OF
BRAUNWORTH & CO.
BOOK MANUFACTURERS
BROOKLYN, N. Y.

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

PREFACE

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 America.

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.

Joseph Race.

Ottawa, Ont.,
April, 1918.

CONTENTS

CHLORINATION OF WATER

CHAPTER I

HISTORICAL

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 χλωρός = 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 success.

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.

TABLE I.—BLEACH STATISTICS.
NORTH AMERICA

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 Scheele.

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 considered.

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 current.

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 organisms.

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 negligible.

“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.

BIBLIOGRAPHY

[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.

CHAPTER II

MODUS OPERANDI

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 + Cl2 = CaOCl2 + H2O. 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.

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

4CaOCl2·2Ca(OH)2·5H2O

which assumes it to contain:

In this formula calcium hypochlorite has been written CaOCl2, 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 CaOCl2 as a mixture of equal molecules of calcium chloride and calcium hypochlorite (CaCl2 + Ca(OCl)2 = 2CaOCl2), 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.

Cl2 + H2O = 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:

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 + H2O + Cl2 → CaOCl2 + H2O

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.

2CaOCl2 = Ca(OCl)2 + CaCl2.

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

2Ca(OCl)2 + 4H2O ⇄ 2Ca(OH)2 + HOCl + HCl and

CaCl2 + 2H2O ⇄ 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 Kw (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 + 2H2O ⇄ CaCl2 + 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 x2⁄(1 - x)v = kwv in which 1 mole of pure Ca(OCl)2 is dissolved in v litres, x is the fraction hydrolysed, and kwv is the hydrolytic constant, complete hydrolysis occurs (x = 1) when v is not greater than 1 × 104 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 CaCl2; 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:

HYDROLYSIS OF BLEACH SOLUTION

[A] Corrected for the alkali produced by HClO + 2KI = KCl + KOH + I2.

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′ concentration.

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.).

TABLE II.[B]

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 derivatives.

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.

TABLE III.[C]—COMPARISON OF BLEACH AND
POTASSIUM PERMANGANATE

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 + O2 + 18,770 calories

2H2O2 = 2H2O + O2 + 46,120 calories

2O3 = 3O2 + 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 Cl2 + H2O = 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.

12NH3 + 6Cl2 = N2 + NCl3 + 9NH4Cl.

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

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 value.

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

TABLE IV.[D]—EFFECT OF AMMONIA ON
CHLORINE GAS SOLUTION

Even when the ratio of Cl : NH3 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 organisms.

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 = Cl2 + H2O 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′ = 320Cl2 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 (H2CO3 ⇄ H· + HCO3′) 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 + O2

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. Cl2·Aq. = 2,600 calories. Chlorine water is comparatively stable but decomposes under the influence of light in accordance with the equation Cl2 + H2O = 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

Cl2·Aq + H2O = 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 + O2 + 18,770 calories

2Cl2·Aq + 2H2O = 2HCl + O2 + 15,340 calories

2Cl2· + Aq + 2H2O = 2HCl + O2 + 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 substances.

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.

BIBLIOGRAPHY

[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, 232.

[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.

CHAPTER III

DOSAGE

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’ contact.

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

TABLE V

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 sources.

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 Cl2/71 + H2O = 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.

TABLE VI.[A]—EFFECT OF COLOUR

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.

Chlorine Treatment at Montreal

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 extent.

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

DIAGRAM I
EFFECT OF COLOUR ON ABSORPTION OF CHLORINE BY WATER

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 t1 and t2 the formula K = log(N1/N2)/(t2 - t1) 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].

TABLE VII.—OXYGEN TO CHLORINE RATIO

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].

TABLE VIII.[B]—EFFECT OF TEMPERATURE

TABLE IX.[C]— EFFECT OF TEMPERATURE

The reaction velocity of a germicide is proportional to the temperature[9] and the influence of temperature may be mathematically expressed by the formula K1/K2 = θ(T2 - T1), in which K1 and K2 are the constants of the reaction at temperatures T2 and T1, respectively, and θ is the temperature coefficient. From the value of θ, the velocity constant of a germicide for any temperature may be calculated from the equation KT = K20° × θ(T - T20°). K1 and K2 are obtained from the formula KT = log(N1/N2)/(t2 - t1) in which N1 - N2 is the number of bacteria destroyed in the interval t2 - t1.

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].

DIAGRAM II
EFFECT OF TEMPERATURE ON ABSORPTION OF CHLORINE BY WATER

[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.

DIAGRAM III
EFFECT OF TEMPERATURE

DIAGRAM IV
EFFECT OF TEMPERATURE

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].

TABLE X.—EFFECT OF EFFICIENT 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:

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:

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 water.

TABLE XI.[D]—EFFECT OF CONTACT PERIOD

TABLE XII.—EFFECT OF CONTACT PERIOD

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

TABLE XIII.[E]—TREATMENT OF MISSISSIPPI RIVER WATER

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 hours.

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:

EFFECT OF SUNLIGHT

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 two.

DIAGRAM V
RELATION OF DOSAGE TO OXYGEN ABSORBED

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.

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].