Dayton Electrolytic CellFig. 9.—Dayton Electrolytic Cell.
Fig. 9.—Dayton Electrolytic Cell.
The outer cell is made of soapstone and is approximately 21⁄2feet 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 thepurpose 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 21⁄2pounds 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:
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
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 hypochloritecontains an excess of sodium chloride, which accelerates it (Race[4]) but the ultimate effect is the same with both. This is shown inTable XXIV.
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 salt.
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 increased.
An electrolyser of the non-submerged diaphragm type is the Allen-Moore cell which has been adopted by the MontrealWater and Power Co. This has been described by Pitcher and Meadows.[8]The general lay-out of the installation is shown inFig. 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.
Brine Saturating and Purifying EquipmentFig. 10—Brine Saturating and Purifying Equipment.
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 determinedand 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 203⁄8inches 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.
Sections of Allen-Moore CellFig. 11.—Sections of Allen-Moore Cell.
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:
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 (videpage 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.
[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 (NH2Cl), 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 carbolicacid 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+ (NH4)2C2O4= CaC2O4+ 2NH4OCl.
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. NH4OCl = NH2CL + H2O. 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 inTable XXV.
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 similarconcentration 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 (videp. 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 inTables XXVI,XXVIIandXXVIII. 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.
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.
Sketch of Ottawa Chloramine PlantFig. 12.—Sketch of Ottawa Chloramine Plant.
Fig. 12.—Sketch of Ottawa Chloramine Plant.
The general design of the plant is shown inFig. 12. The bleach is mixed in tankAas a solution containing 0.3 to 0.6 per cent of available chlorine and delivered to tanksBandD, each of which has a twenty-four-hour storage capacity. The ammonia solution is mixed and stored in tankBand contains 0.3-0.5 per cent of anhydrous ammonia. The two solutions are run off into boxesEandFwhich maintain a constant head on valvesVandV′controlling the head on the orifices. Both orifices discharge into a common feedboxGfrom which the mixture is carried by the water injectorJthrough one of duplicate feed pipes and discharged into the suction well through a perforated pipe.
As tankBwas 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 XXIXshows 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.
At the height of the spring floods the raw water contained 80 p.p.m. of turbidity and over 500B. coliper c.cm. but 0.6 p.p.m. of chlorine and 0.13 p.p.m. of ammonia reduced theB. coliindex 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 aB. coliindex 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 theB. coliindex to 2.0 per 100 c.cms. under similar physical and bacteriological conditions.
During the period of nine months covered by the results inTable 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 inDiagram VIII. The percentage of samples of treated water showingB. coliin 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+ 2NH3= 2NH2Cl + Ca(OH)2
would indicate a ratio of 2 : 1; but only one molecule of Ca(OCl)2is produced from two molecules of bleach and the theoretical ratio is therefore 4 : 1 (142 : 34),
The chlorine to ammonia ratio is very important because of its influence on the economics of the process (videp. 124).
DIAGRAM VIIICHLORAMINE TREATMENT, OTTAWAChloramine Treatment, Ottawa
DIAGRAM VIIICHLORAMINE TREATMENT, OTTAWA
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 capacityof 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 (seep. 56) have been entirely eliminated and that theB. colicontent 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 (seeDiagram 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, andthe 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.