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CHOLERA, PLAGUE, SMALLPOX, TYPHUS FEVER, AND YELLOW

FEVER-Continued.

Reports Received from December 30, 1922, to January 19, 1923-Continued.

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CHOLERA, PLAGUE, SMALLPOX, TYPHUS FEVER, AND YELLOW

FEVER-Continued.

Reports Received from December 30, 1922, to January 19, 1923-Continued.

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PUBLIC HEALTH REPORTS

VOL. 38

FEBRUARY 2, 1923

No. 5

An Experimental Study of the Relation of Hydrogen Ion Concentrations to the Formation of Floc in Alum Solutions.

By EMERY J. THERIAULT, Assistant Chemist, and W. MANSFIELD CLARK, Chief of Division of Chemistry, Hygienic Laboratory, United States Public Health Service.

In approaching a scientific analysis of the art of water clarification it seems to us essential to distinguish the several aspects of the subject. These are so integrated in actual plant operation that it is difficult to perceive the true importance of each of the several factors which have to be mastered by the operator under every exigency. The isolation of phenomena, and their exact quantitative evaluation, will alone permit a true appraisal of any factor in relation to the process as a whole.

We have limited our attention to certain laboratory experiments which clarify one distinct aspect of the alum process. Our data doubtless lack the scope desirable for general practical application, but they indicate that, unless factors still to be investigated have an unexpected influence, maximum precipitation of added aluminium will occur within definite and narrow limits of hydrogen ion concentration.

It is well recognized that a precipitate is not formed from alum when the final solution is either too" acid" or too" alkaline." Hitherto the essential degree of "acidity" or "alkalinity" has been sought in the quantity of acid or alkali determined by one or another analytical method. More recently there has been a growing appreciation of the fact that the waterworks operator is dealing with reversible reactions, that his task is to control equilibria, and that all too often the methods of the analyst, devised originally for other purposes, upset an established equilibrium to yield information of dubious value to the case at hand.

There is little need to review here the relations which have been brought to light in studies of equilibria among acids and bases. They have been emphasized frequently and summarized in various treatises. However, we shall preface our contribution by a brief discussion of a set of equilibria pertinent to the subject.

It has frequently been emphasized that the construction of a titration curve may show the salient characteristics of an acid, a base, or an ampholyte. It is therefore interesting to approach the

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problem first by a consideration of the titration curve of an aluminium salt.

In Figure 1 are shown two such curves. These were constructed from the data obtained when known concentrations of potassium alum were titrated with sodium hydroxide solutions and the hydrogen ion concentration was measured after each addition of alkali.

For the determination of hydrogen ion concentration, the electrometric method was used. The potentiometric system consisted of a Leeds and Northrup type K potentiometer and type R galvanometer with properly shielded switchboard and wiring. The Weston cell values were certified by the Bureau of Standards. The calomel half-cells used as working standards were compared with a battery of tenth normal KCl calomel half-cells and the system was brought to the standard recommended in "The Determination of Hydrogen Ions" (Clark, 1920).

The hydrogen used was electrolytic hydrogen supplied in tanks. It was passed over heated platinized asbestos and then over sodalime, and wherever possible was led through copper tubing rather than rubber tubing.

Since very dilute solutions were to be dealt with in some of the experiments, it was considered of more importance to guard the solution from atmospheric contamination during handling than to take advantage of several features in Clark's electrode vessel. Therefore, instead of titrating aliquots and measuring each when separately transferred to the electrode vessel, continuous electrode measurements were carried out in Erlenmeyer vessels, into which were led the electrodes, the gas inlet and outlet, a gooseneck siphon from the saturated calomel half-cell, and a burette tip. The chief objection to this arrangement was in the narrow liquid junction-made narrow to prevent too great contamination of the solution by the saturated KCl of the liquid junction and too great a loss of titrated material by the exchange. However, this inherent error was probably small and constant and displaces our curves in no essential respect. All titrations were made at constant temperature.

The alum solutions used in the titrations illustrated in Figure 1 were prepared from conductivity water, and a sample of

K2SO ̧∙Al1⁄2 (SO,),24H2O

which had been recrystallized several times. Analysis by the procedure of Blum (1916) indicated 99.6 per cent, 100.1 per cent, and 100.1 per cent, or an average of 99.93 per cent of the theoretical amount of aluminium. Once prepared, this sample was carefully protected against change in moisture content.

The sodium hydroxide solution and the calcium hydroxide solutions used in other measurements were prepared from electrolytic

amalgams which were decomposed with boiled-out "conductivity" water. All operations were made with adequate protection against contamination by atmospheric CO,.

In accordance with custom, we shall express hydrogen ion concentrations in terms of Sørensen's pH, which is defined by the relation

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With the type of vessel used in these titrations, considerable time was necessary at the start to thoroughly displace oxygen and to establish equilibrium stable within a small part of a millivolt. Thereafter there was found evidence that the potentials obtained after each addition of alkali rapidly became indicative of a true electrode equilibrium. This conclusion rests on the following evidence: The same equipment rapidly gave equilibrium potentials with systems which do not change as does the aluminium system, and it furnished smooth titration curves on rapid titration of aluminium salts.

On the other hand, with mixtures of aluminium salts and alkali, there was noted a slow drift in potential which was not in all cases entirely orderly, but which had the characteristics indicated below.

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This, the smoothest relation found, shows, on graphic extrapolation, that if the drift had continued until the pH of the solution was altered 0.05 pH units it would have required about 166 hours. Generally, however, the drift was not so linear in relation to log t, but continued. to decline at a continuously lower rate than that indicated above. Evidently such drifts would be overlooked by one using a cruder method of measurement. We believe them to be due to a slow shift in equilibria due to the properties of the solid phase.

In the more alkaline regions the drift was often of a serious nature, and the plotted curves can not be taken to indicate that complete equilibrium was approached in the solutions more alkaline than pH 7. The drift was of uncertain nature, and it is probably to be accounted for by the complexity of the processes occurring.

The "aluminium hydroxide" precipitated at the higher acidities is slowly redissolved on the addition of more alkali, and at the same time there is a tendency for the separation of the crystalline form of the hydroxide as noted by various observers. From this complex,

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