SOLUBILITY O F OXYGEX I N SALT SOLUTIONS A S D T H E HYDRATES O F THESE SALTS BY C. G. hlACilRTHUR
Introduction In another investigation, data on the solubility of oxygen in solutions of several of the ordinary salts were needed. The literature contained determinations for a few acids and bases and one salt, sodium ch1oride.l It, therefore, became necessary to find out how soluble oxygen (in the presence of the other gases of the air) was in different molecular concentrations of several typical salts. Incidentally, it will be possible to use these data to calculate the degree of hydration of these salts in solution.
Experimental The salt was made up to a definite concentration in tall cylinders and placed in a thermostat at 2 5 O C. The cylinders were shaken by hand several times a day for about four days. The solution was then poured into a 2 5 0 cc rneasuring flask that was also graduated to 252 cc and had a neck of such length that the bottom of a rubber stopper came just to the 252 cc mark. To the 250 cc of solution in the flask was added simultaneously I cc of the alkaline potasI O grams KI dissolved in sium iodide (33 grams KaOH water and diluted to IOO cc) and I cc of the manganous chloride solution (40 grams MnC12.4H20made up to IOO cc with H20). The flask was stoppered to the 2 j 2 cc mark, thoroughly shaken, and the precipitate allowed to settle. 3 cc of strong hydrochloric acid was added. The flask was again stoppered and shaken. The contents were then poured into a 7 0 0 cc flask and titrated with S IOO sodium thiosulfate solution, using starch as indicator. By the above method the amount of oxygen in distilled water a t 25' C and a t a pressure of 760 rnm of mercury was found to be j . 7 8 cc. Dupli-
+
1
Geffcken: Zeit. phys. Chem., 49, z j 7 (1904).
496
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Salt and concentration
Solubility of Density a t
2j
I
I 1
j.6j 5.49
8.8 5.3 6.0 8.9
I ,056 I . 116
5.20 4.75
1.23
3.77
8.1 7 .0 6.4
1.46
I
.8r
1.j
Solubilitj~o j Oxygen in Salt Solutiom, Etc.
49 7
cate experiments were not run except on the tn 8 concentration o€ each salt. Controls with distilled water were run along with each series of salts. Therefore, barometric pressure could be neglected because it would have practically the same effect on the control as on the solution. The densities of the solutions were taken or calculated from the density tables in the International Tables of physicochemical constants.
C. G.,lIacArthur
498
The degree of hydration was calculated from the following formula a-b 100-c nz XX 7= Hydration a C yn a = cc of 0 in liter of HzO. h = cc of 0 in 1000 grams of the water in the solution = cc oxygen per liter of solution density -grams of salt per cc of solution' c = per cent of salt in solution. YM = molecular weight of the salt. tn' = molecular weight of water.
Discussion of Data on Solubility of Oxygen From the above data one will notice that in all cases the solubility of oxygen decreases with an increase in the concentration of the salt. The decrease per molecule of salt in solubility is greatest in the most dilute solutions. If one constructs a curve for the solubility of oxygen in increasing concentrations of salt, the curve flattens and becomes parallel with high concentration, indicating that all the oxygen would be removed only from concentrations that are much higher than are possible. The solubility of oxygen in m, 8 solutions of the chlorides does not decrease regularly from that of caesium to lithium with a decrease in molecular weight as one would expect. The figures are caesium chloride 5.67, rubidium chloride 5.65, potassium chloride 5.5 2 , sodium chloride 5.5 2 , lithium chloride 5.63. This irregularity can probably best be explained by the difference in potential of the ions in solution. It can not be due simply to the fact that there are ions present because the difference in number of ions in m 8 solutions does not thus vary with the different salts. It must be due to specific differences between the ions. Another fact to be noted is the marked effect of ammonium chloride on the solubility of oxygen. It varies with the concentration, but the decrease in solubility is extremely large. At first it was thought to be due to a combination of the am1
Philip: Trans. Faraday SOC.,3, 1 2 3 .
Solubilitj' o j Oxygen in Salt Solutions, Etc.
499
monium chloride with the manganous chloride, thus preventing the manganous compound from being oxidized by the oxygen in solution. This was shown not to be true by adding the ammonium chloride to water containing the usual amount of the alkaline iodide and manganous chloride solutions and finishing the determination as usual. This gave the amount of oxygen that distilled water usually contains. Though it is more likely that the ammonium chloride so influences the oxygen that it is not available to oxidize the manganous hydroxide, it is still possible that the ammonium chloride markedly decreases the amount of oxygen in solution. By adding the manganous chloride solution and the alkaline iodide solution simultaneously and allowing the precipitate to stand about two hours before adding the hydrochloric acid, it is believed the insolubility of the hydroxides of barium, calcium and magnesium did not affect the completion of the oxidation of the manganous hydroxide. In determining the amount of oxygen in sucrose solutions a difficulty was encountered that renders the results given in the table somewhat low. I t was noticed that the sugar slowly uses the oxygen in the solution to oxidize itself, even at 25' C. The error in the figures of the table is not large because it was found that if the solutions were shaken vigorously immediately before adding the manganous chloride and alkaline iodide solutions rather constant maximum results were obtained that agreed fairly well with each other and with the results from the salts.
Discussion of Hydration Data If one considers that a substance dissolves in water because it forms a hydrate, it would seem that the force causing the ionization is the one causing the combination between solute and water. When the pull of the water on the molecule is great enough to overcome the force holding the parts of the molecule together, i t will split into parts. Thus when NaCl is placed in contact with HzO the attraction of the water for the salt to form a hydrate will cause a separation of the salt
500
C.
G.M a d rthur
crystals into molecules which are surrounded by water of hydration. The amount of this water of hydration is continually changing. TT’hen enough of this water pulls on the salt molecule it will cause its ionization. According to the kinetic theory, some of these hydrated ions will lose part of these loosely combined water molecules ; then the ions will recombine to form molecules. The larger the dilution the greater the percent of molecules that will be separated into ions, because the chances are then greater, kinetically, that enough water will be influencing the molecule to ionize it. Data on the degree of hydration of ions or molecules will indicate their average hydration. Any particular ion or molecule may be hydrated considerably more or less than the average. It is the moleciiles that become largely hydrated that will part into ions. It is very probable that the molecules in a dilute solution have a hydration nearly as great as the sum of the hydration of their ions. When oxygen dissolves in water it combines in a similar way to form a hydrate. The water will be saturated with oxygen molecules moving a t a given rate at a given temperature and pressure, when the attraction of the water for the oxygen molecules is just large enough to hold that number (not the same molecules continually, of course) from flying back into space above the water. Other molecules or ions, like those of sodium chloride, which also form hydrates will decrease the amount of water that will influence the oxygen molecules ; therefore, the number of dissolved oxygen molecules will decrease. It is possible that the number of water molecules attached to the salt molecule determines simply and definitely the amount-of oxygen that can remain in the water. In this case a definite proportion would exist between the amount of oxygen in a solution and the degree of hydration of the salt added. By consulting the data it will be seen that in the case of potassium chloride this definite proportion exists. This is also to a certain extent true of potassium bromide, sodium
sulfate, potassium sulfate, magnesium chloride, calcium chloride, and probably barium chloride. I n all the other salts studied the degree of hydration does not increase with a decrease in concentration, a t least not consistently. It is probable then that in these salts that do not give a proportional increase in hydration with an increase in ionization, the ions that are present influence the amount of oxygen dissolved in the solution. In the higher concentrations, where the percentage of ionization is small, there is fairly good agreement between the hydration data obtained for these concentrations and results calculated from data obtained for dilute solutions by other methods. However, when the ionized particles begin to predominate the expected increase in hydration, as calculated from the oxygen solubility, does not appear. It is conceivable that the oxygen in the solution does decrease the hydration of the ions more than it does that of the molecules. But it is probable that the sum of the hydration of the ions is somewhat larger than the hydration of the molecules that are made up of these ions. It is easier to believe that the difference in potential between the ions (because some ions do not have this effect) cause an increase in the solubility of the oxygen by attracting it, thus giving data for low hydration values. If the data are studied in the light of what has just been stated and the assumption is made that the hydration of the chlorine ions calculated from the diffusion coefficient is correct (9.6 a t infinite dilution)’ the hydration of a potassium ion in an w i 8 solution becomes 8. If 9.6 H 2 0 is correct for potassium, the data give the bromine ion 6.6 H?O in w 4 solution. If chlorine is 9.6 H 2 0 , calcium equals 6.1, barium 6.7 and magnesium 1 2 . 7 . These are minimal values, of course. nTth potassium hydration 9 6, the sulfate radicle is calculated t o be 28.9. These values are not far from those calculated from the diffusion coefficient, nor from those obtained by the more reliable experimental methods. Smith. Jour. Am. Chem. SOC., 37, Smith Ibid., 37, 7 2 2 (1915).
722
(1915).
By comparing the hydration of a particular concentration of one salt with the same concentration of another, it will be noticed that sodium hydration is greater than potassium, potassium than rubidium, rubidium than caesium. The chlorides are more hydrated than the bromides; the bromides than the iodides. Summary The solubility of oxygen in various concentrations of the following substances was determined by the manganous hydroxide method : the chlorides of lithium, sodium, potassium, rubidium, caesium, ammonium, barium, calcium, magnesium ; the bromides of sodium and potassium; the iodide of potassium; the nitrate of potassium; the sulfates of sodium and potassium, and sucrose. From the data for the solubility of oxygen the hydration of the above named salts was calculated. In all cases it was found that an increase in concentration of the substance decreased the solubility of oxygen in a regular manner. Ammonium chloride has a very large effect on the solubility of oxygen. Sucrose is slowly oxidized by the dissolved oxygen. This interferes with the determination of oxygen in sugar solutions. Potassium chloride, potassium bromide, potassium sulfate, barium chloride, calcium chloride, and magnesium chloride seem to give hydration values that are approximately correct. The other salts studied give low values for the highly ionized solutions, indicating that the particular ions increase the solubility of oxygen to a definite extent which is specific for that particular ion; or that it is the potential difference between the two ions present that causes this increase in oxygen solubility . Baochemzstry Department T h e Cnzz'erszty of Illznoas
