Solubility of AgBrO3 in AgBrO3-MeNO3-H2O Systems
Abstract Generated abstract
The study examines the solubility of silver bromate in aqueous lithium, sodium, and potassium nitrate solutions at 25 degrees C over nitrate concentrations from 0.001 to 1 mol/l, in order to assess electrolyte effects without common ions. Silver bromate solubility was measured analytically, and the data were used to calculate ionic strengths, solubility products, activity products, and activity coefficients using Debye Hückel relations. The solubility and formal solubility product increased regularly with nitrate concentration, with the effect largest for potassium nitrate and smallest for lithium nitrate. The work reports an approximately constant activity product of 5.5 x 10^-5 and proposes a practical calculation method for estimating silver bromate solubility in alkali metal nitrate solutions.
Full Text
SOLUBILITY OF AgBrO₃ IN THE SYSTEMS
AgBrO₃—MeNO₃—H₂O
Yu. L. Lelchuk and S. M. Sasonko
Studies of the solubility of precipitates in solutions of various electrolytes, both those containing ions common with the precipitate and those not containing such ions, are of great importance in the theory and practice of modern analytical chemistry.
In questions of the solubility of precipitates of analytical significance in solutions containing ions common with the precipitate, an essential role was long attributed to the so-called solubility product. The rule concerning the solubility product \((L_p)\) was first formulated by Nernst as early as 1889. Nernst, proceeding from the theory of electrolytic dissociation developed by Svante Arrhenius in 1887, and applying the classical law of mass action to the heterogeneous system precipitate of an electrolyte—saturated solution, arrived at the conclusion that the product of the concentrations of the ions in a saturated solution in contact with the precipitate, at constant temperature, is a constant quantity.
However, as a result of numerous experimental verifications of Nernst’s basic proposition on the constancy of \(L_p\), carried out by many investigators in this country and abroad in recent decades, considerable deviations of experimental data from this proposition became known. On the other hand, the development of views on the theory of solutions and the law of mass action, on which Nernst based his conclusions, led to substantial corrections of the latter. One of the first to point out a number of very substantial discrepancies between the experimental data obtained in the study of solutions of strong electrolytes and Arrhenius’s theory of electrolytic dissociation, and the degrees of dissociation derived from it, was our great compatriot D. I. Mendeleev [1]. Further studies of solutions of strong electrolytes not only justified D. I. Mendeleev’s criticism, but also made it possible to accumulate a number of new facts, as a result of which it became necessary to revise and substantially change the basic propositions of Arrhenius’s theory of electrolytic dissociation.
The new theories of solutions of strong electrolytes that have arisen in recent decades are based on the assumptions that all strong electrolytes in aqueous solutions are completely dissociated; that in the behavior of solutions an essential role is played not by formal concentrations, but by the so-called “active concentrations”; and that the general properties of solutions are substantially influenced by the charges of the ions and by interionic electrical forces, which are not at all taken into account by Arrhenius’s classical theory. Among the theories that arose on this basis, the electrostatic theory of solutions of strong electrolytes of Debye and Hückel has received the widest acceptance.
A substantial change was also introduced into the law of mass action. Significant deviations of experimental data from the law of mass action in its original formulation demonstrated the need to replace, in the mathematical expression of this law, the formal concentrations of the reacting substances by activities.
In this connection it became necessary to clarify and refine in greater detail the theoretical meaning of the solubility-product rule and to determine the limits of its applicability. Among the most important works devoted to this question, a special place is occupied by the work of the Soviet scientist A. F. Kapustinskii [2]. In his article “The Solubility Product and Its Significance in General and Analytical Chemistry,” he showed that the solubility-product rule, in its strict modern formulation, is a consequence of one of the most general laws of nature concerning energy transformations—namely, the second law of thermodynamics: if, in the expression for the solubility product, formal concentrations are replaced by activities, then the solubility-product rule, formulated as a rule for the product of activities, becomes a law admitting no exceptions.
This law, in A. F. Kapustinskii’s formulation, states: “In a saturated solution at constant temperature, the product of the active concentrations of the ions of the dissolved substance is a constant quantity” [2].
In the practice of analytical chemistry, the solutions under investigation, as a rule, contain, in addition to the ion being precipitated, other salts that have no ions in common with the precipitate. In this connection, the influence of like ions on the solubility of a given precipitate must be considered against the background of the influence of other ions. As follows from the Debye–Hückel theory, the conclusions of Kapustinskii, and the investigations of Soviet scientists I. V. Tananaev, Mizetskaya, Kapustinskii, and many others, electrolytes that do not contain ions in common with the precipitate, as a rule, increase the solubility of the precipitate. The degree of increase in solubility depends on the valence type of the electrolyte and on its chemical nature. In certain cases, however, the addition of any electrolyte causes an increase in the activity coefficient of the precipitate and a decrease in its solubility. Only the use of the activity product instead of Nernst’s solubility-product rule makes it possible to take into account the combined influence of the ionic medium on the solubility of precipitates of analytical significance. However, the broad application of the activity product for practical purposes in analytical chemistry is still hindered by the fact that the values of the activity coefficients of salts are not always known at various values of the ionic strength of the solution.
This most important question of the theory and practice of modern analytical chemistry has still received little attention from researchers; nevertheless, the investigations of Soviet scientists in this field already compare favorably with analogous investigations by foreign scientists. These investigations are distinguished by the broad application of the methods of physicochemical analysis developed by the numerous school of N. S. Kurnakov. It is precisely with such an approach that studies of the solubility of precipitates make it possible to encompass both a wide range of concentrations of foreign electrolytes and a large group of electrolytes themselves. Only with such an approach to the study of solubility can one expect the most complete and correct identification of the general regularities in which both physical and chemical factors influencing the behavior of precipitates in solutions of different electrolytes and different concentrations can be taken into account.
Among the most important works in this field, one should first of all name the work of I. V. Tananaev and I. B. Mizetskaya [3], who studied the solubility of lead sulfate in aqueous solutions of nitrates: lithium, sod—
...barium, potassium, calcium, aluminum, and lead, and the sulfates of hydrogen, lithium, sodium, and potassium.
The present investigation is devoted to the study of the solubility of silver bromate in aqueous solutions of alkali-metal nitrates. The experimental data have been used by us to determine the activity coefficients of silver bromate in aqueous salt solutions at various values of the ionic strength of the solutions. This work is the next communication in a series of works devoted to the study of the solubility of silver bromate in solutions of various electrolytes, both containing and not containing ions common with the precipitate. The investigation as a whole has the aim of studying the regularities occurring in ternary and corresponding quaternary systems with the participation of ions of different valence and chemical nature.
Starting Materials and Experimental Procedure
A pure preparation of $\mathrm{AgBrO_3}$ was obtained by us by adding, in the cold and dropwise, an equivalent amount of a 5% solution of potassium bromate to a 10% solution of silver nitrate, with continuous motor stirring. After repeated washing, the precipitate was recrystallized, washed with bidistillate, and pressed on filter paper to an air-dry state. The preparation was stored in a jar protected from the action of light by a cap of black paper. Lithium nitrate was obtained by us from lithium carbonate by treatment with chemically pure nitric acid; sodium and potassium nitrates—by double recrystallization of domestic commercial preparations. All salt solutions were prepared by appropriate dilution of one-molar solutions. Preparation and dilution of the solutions were carried out with bidistillate kept in vessels paraffined on the inside.
Experiments to determine the solubility of silver bromate were carried out in paraffined flasks of 400 ml capacity. A flask closed with a stopper (with an opening for the stirrer), containing 1 g of $\mathrm{AgBrO_3}$ and 250 ml of an alkali-metal nitrate solution of the given concentration, was placed in a water thermostat at a temperature of $25 \pm 0.1^\circ$, where vigorous motor stirring was carried out for 8 hours.
Before taking a sample of the liquid phase for analysis, the solution with the precipitate was allowed to stand in the thermostat for 40–45 minutes. The concentration of $\mathrm{AgBrO_3}$ in the saturated solution was determined both from the content of silver ions and from the bromate ions. Silver was determined gravimetrically as $\mathrm{AgCl}$, and the bromate ion iodometrically. In separate samples, the content of silver bromate in solution was additionally determined gravimetrically as $\mathrm{AgBr}$ after reduction of silver bromate with acid potassium nitrite. Precipitation of $\mathrm{AgCl}$ and $\mathrm{AgBr}$ was carried out in beakers protected from the action of light by caps of black paper.
For each point of the system, two parallel experiments were carried out, each of which was analyzed twice. For calculating the solubility of silver bromate, the ionic strength of the solutions, the activity coefficients, and the activity product, the mean values of these determinations were taken.
Solubility of $\mathrm{AgBrO_3}$ in Water
With regard to the solubility of silver bromate in water, there are contradictory data in the literature. Ricci and Aleshnik [4] explain the large discrepancies in the solubility data obtained by different authors by the difficulty of obtaining a chemically pure preparation of $\mathrm{AgBrO_3}$ from sodium bromate and silver nitrate. As the basis of all our calculations we...
the value for the solubility of silver bromate in water, \(8,24\cdot10^{-3}\ \text{mol/l}\), found by I. V. Tananaev, Lel’chuk, and Petrovitskaya [5], was used. Approximately the same data were obtained by Ridi [7] and Hill [6].
Systems \(\mathrm{MeNO_3—AgBrO_3—H_2O}\), where \(\mathrm{Me=Li, Na}\) and \(\mathrm{K}\)
The results of determining the solubility of silver bromate in aqueous solutions of alkali-metal nitrates are given in Tables 1, 2, and 3. The first column of the tables gives the concentration of the alkali-metal nitrate; the second, the molarity of \(\mathrm{AgBrO_3}\) in the saturated solution; the third, the ionic strength of the solution; the fourth, the activity coefficient of \(\mathrm{AgBrO_3}\); the fifth, the solubility product of \(\mathrm{AgBrO_3}\); and the sixth, the activity product of \(\mathrm{AgBrO_3}\).
Table 1
Solubility of \(\mathrm{AgBrO_3}\) in \(\mathrm{LiNO_3}\) solutions
| \(\mathrm{LiNO_3}\) (mol/l) | \(\mathrm{AgBrO_3}\) (mol/l) | \(\mu\) | \(\gamma\) | \(L_p\) | \(L_a\) | \(a\) |
|---|---|---|---|---|---|---|
| 0,000 | \(8,24\cdot10^{-3}\) | 0,00824 | 0,900 | \(6,80\cdot10^{-5}\) | \(5,50\cdot10^{-5}\) | 3,12 |
| 0,001 | \(8,31\cdot10^{-3}\) | 0,00931 | 0,903 | \(6,91\cdot10^{-5}\) | \(5,63\cdot10^{-5}\) | 3,12 |
| 0,003 | \(8,43\cdot10^{-3}\) | 0,01148 | 0,894 | \(7,10\cdot10^{-5}\) | \(5,67\cdot10^{-5}\) | 3,12 |
| 0,010 | \(8,47\cdot10^{-3}\) | 0,01847 | 0,870 | \(7,18\cdot10^{-5}\) | \(5,44\cdot10^{-5}\) | 3,12 |
| 0,030 | \(9,01\cdot10^{-3}\) | 0,0390 | 0,826 | \(8,11\cdot10^{-5}\) | \(5,53\cdot10^{-5}\) | 3,12 |
| 0,100 | \(9,73\cdot10^{-3}\) | 0,1097 | 0,751 | \(9,47\cdot10^{-5}\) | \(5,34\cdot10^{-5}\) | 3,12 |
| 0,300 | \(1,094\cdot10^{-2}\) | 0,3109 | 0,662 | \(1,20\cdot10^{-4}\) | \(5,23\cdot10^{-5}\) | 3,12 |
| 1,000 | \(1,317\cdot10^{-2}\) | 1,0132 | 0,564 | \(1,73\cdot10^{-4}\) | \(5,50\cdot10^{-5}\) | 3,12 |
Table 2
Solubility of \(\mathrm{AgBrO_3}\) in \(\mathrm{NaNO_3}\) solutions
| \(\mathrm{NaNO_3}\) (mol/l) | \(\mathrm{AgBrO_3}\) mol/l | \(\mu\) | \(\gamma\) | \(L_p\) | \(L_a\) | \(a\) |
|---|---|---|---|---|---|---|
| 0,000 | \(8,24\cdot10^{-3}\) | 0,00824 | 0,900 | \(6,80\cdot10^{-5}\) | \(5,50\cdot10^{-5}\) | 2,09 |
| 0,001 | \(8,22\cdot10^{-3}\) | 0,00922 | 0,901 | \(6,76\cdot10^{-5}\) | \(6,08\cdot10^{-5}\) | 2,09 |
| 0,003 | \(8,31\cdot10^{-3}\) | 0,01131 | 0,891 | \(6,91\cdot10^{-5}\) | \(5,49\cdot10^{-5}\) | 2,09 |
| 0,010 | \(8,50\cdot10^{-3}\) | 0,0185 | 0,865 | \(7,23\cdot10^{-5}\) | \(5,41\cdot10^{-5}\) | 2,09 |
| 0,030 | \(9,27\cdot10^{-3}\) | 0,03927 | 0,817 | \(8,59\cdot10^{-5}\) | \(5,72\cdot10^{-5}\) | 2,09 |
| 0,100 | \(1,008\cdot10^{-2}\) | 0,11008 | 0,731 | \(1,02\cdot10^{-4}\) | \(5,44\cdot10^{-5}\) | 2,09 |
| 0,300 | \(1,138\cdot10^{-2}\) | 0,3114 | 0,626 | \(1,30\cdot10^{-4}\) | \(5,09\cdot10^{-5}\) | 2,09 |
| 1,000 | \(1,48\cdot10^{-2}\) | 1,0148 | 0,501 | \(2,19\cdot10^{-4}\) | \(5,50\cdot10^{-5}\) | 2,09 |
Table 3
Solubility of \(\mathrm{AgBrO_3}\) in \(\mathrm{KNO_3}\) solutions
| \(\mathrm{KNO_3}\) (mol/l) | \(\mathrm{AgBrO_3}\) (mol/l) | \(\mu\) | \(\gamma\) | \(L_p\) | \(L_a\) | \(a\) |
|---|---|---|---|---|---|---|
| 0,000 | \(8,24\cdot10^{-3}\) | 0,00824 | 0,900 | \(6,80\cdot10^{-5}\) | \(5,50\cdot10^{-5}\) | 1,47 |
| 0,001 | \(8,26\cdot10^{-3}\) | 0,00924 | 0,899 | \(6,82\cdot10^{-5}\) | \(5,50\cdot10^{-5}\) | 1,47 |
| 0,003 | \(8,48\cdot10^{-3}\) | 0,01148 | 0,888 | \(7,19\cdot10^{-5}\) | \(5,67\cdot10^{-5}\) | 1,47 |
| 0,010 | \(8,61\cdot10^{-3}\) | 0,01861 | 0,862 | \(7,41\cdot10^{-5}\) | \(5,50\cdot10^{-5}\) | 1,47 |
| 0,030 | \(9,30\cdot10^{-3}\) | 0,0393 | 0,810 | \(8,65\cdot10^{-5}\) | \(5,68\cdot10^{-5}\) | 1,47 |
| 0,100 | \(1,024\cdot10^{-2}\) | 0,1102 | 0,717 | \(1,05\cdot10^{-4}\) | \(5,40\cdot10^{-5}\) | 1,47 |
| 0,300 | \(1,24\cdot10^{-2}\) | 0,3124 | 0,599 | \(1,54\cdot10^{-4}\) | \(5,52\cdot10^{-5}\) | 1,47 |
| 1,000 | \(1,63\cdot10^{-2}\) | 1,0163 | 0,455 | \(2,66\cdot10^{-4}\) | \(5,50\cdot10^{-5}\) | 1,47 |
The activity coefficient of \(AgBrO_3\) was calculated by the method developed by I. V. Tananaev and I. B. Mizetskaya [3], on the basis of the constancy of the activity product of \(AgBrO_3\) in all solutions saturated with it.
For each system separately, the ionic strength of the saturated aqueous solution at \(25^\circ\) was first determined from the formula:
\[ \mu=\frac{1}{2}\sum mz^2 . \tag{1} \]
Then the activity coefficient of \(AgBrO_3\) in the saturated aqueous solution at \(25^\circ\) was determined by the Debye–Hückel formula for solutions whose concentrations do not exceed \(0.02\ \mathrm{mol/l}\).
\[ -\log \gamma = 0.505\, z_1 z_2 \sqrt{\mu}. \tag{2} \]
After this, the activity product was calculated from the equation:
\[ La = Lp\cdot \gamma^2, \tag{3} \]
the activity coefficient of \(AgBrO_3\) in solutions with the highest concentration of the foreign electrolyte from the formula:
\[ \gamma=\sqrt{\frac{La}{Lp}}, \]
the ionic strengths of all solutions from formula (1), the coefficient \(a\) in the Debye–Hückel formula for calculating the activity coefficient in solutions whose concentrations exceed \(0.02\ \mathrm{mol/l}\) from the formula:
\[ -\log \gamma=\frac{0.505\cdot Z_1\cdot Z_2\sqrt{\mu}}{1+0.33\,a\sqrt{\mu}}, \tag{4} \]
and the activity coefficients of \(AgBrO_3\) in all solutions saturated with it from formula (4) with known \(a\).
Discussion of the Results
As can be seen from the tables presented and the corresponding curves in Figs. 1, 2, 3, and 4, the solubility of \(AgBrO_3\) in solutions of all alkali-metal nitrates increases regularly with increasing concentration of the latter. Correspondingly, the solubility product of \(AgBrO_3\) also increases continuously. The latter is clearly illustrated by Fig. 5.
Of the three alkali-metal cations considered, the greatest increase in the solubility of the \(AgBrO_3\) precipitate is produced by the potassium cation, and the smallest by the lithium cation. Thus, the increase in solubility of the precipitate depends not only on the valence type of the foreign electrolyte, but also on the position of the element in the periodic system of D. I. Mendeleev. The value of \(a\) in formula (4) is not the same for all three systems. It varies from 1.47 for the \(AgBrO_3—KNO_3—H_2O\) system, to 2.09 for the \(AgBrO_3—NaNO_3—H_2O\) system, and to 3.12 for the \(AgBrO_3—LiNO_3—H_2O\) system.
The solubility product of \(AgBrO_3\) in all three systems increases appreciably with increasing concentration of the alkali-metal nitrate. In the \(AgBrO_3—KNO_3—H_2O\) system, the solubility product of \(AgBrO_3\) in a one-molar \(KNO_3\) solution is 3.91 times greater than in water; in the \(AgBrO_3—NaNO_3—H_2O\) system it is 3.22 times greater; and in the \(AgBrO_3—LiNO_3—H_2O\) system, 2.54 times.
Fig. 1. Solubility of AgBrO₃ in LiNO₃ solutions
Fig. 2. Solubility of AgBrO₃ in NaNO₃ solutions
Fig. 3. Solubility of \(\mathrm{AgBrO_3}\) in \(\mathrm{KNO_3}\) solutions
Fig. 4. Solubility of \(\mathrm{AgBrO_3}\) in \(\mathrm{MeNO_3}\) solutions.
I—\(\mathrm{KNO_3}\); II—\(\mathrm{NaNO_3}\); III—\(\mathrm{LiNO_3}\)
As I. V. Tananaev and Mizetskaya [3] showed, for less soluble precipitates under analogous conditions the solubility product increases by a factor of 100 or more.
Fig. 5. Change in \(L_p\) of \(\mathrm{AgBrO_3}\) in \(\mathrm{MeNO_3}\) solutions.
I—\(\mathrm{KNO_3}\); II—\(\mathrm{NaNO_3}\); III—\(\mathrm{LiNO_3}\).
As is evident from Fig. 6, there is a definite dependence of the value \(a\) on the ionic radius of the corresponding cation of the external electrolyte. As I. V. Tananaev, Yu. L. Lel’chuk, and Petrovitskaya [5] showed, for ca-
Fig. 6. Dependence of the value \(a\) on the ionic radius of the \(\mathrm{MeNO_3}\) cation
… of alkaline-earth metals, the values of \(a\), as a function of the ionic radii, fall on a straight line on an analogous graph.
Table 6 and Fig. 7 give an idea of the change in the activity coefficient of \(\mathrm{AgBrO_3}\) in solutions as a function of the concentration of the alkali-metal nitrate. To construct the curve, the activity-coefficient values found for individual electrolytes of the corresponding concentrations were averaged.
Fig. 7. Mean value of \(\gamma_{\mathrm{AgBrO_3}}\) as a function of
\(C_{\mathrm{MeONO_3}}\)
These data are of interest because they can be used to calculate the solubility of silver bromate in solutions of alkali-metal nitrates of any concentration (not exceeding \(1\ \mathrm{mol/l}\)) with an accuracy sufficient for practical purposes.
The calculation should be carried out by the formula:
\[ C_{\mathrm{AgBrO_3}}=\sqrt{\frac{L_a}{\gamma^2}}, \]
where \(L_a=5.5\cdot 10^{-5}\), and \(\gamma\) is found with the aid of the curve (Fig. 7).
Table 4
Behavior of \(\mathrm{AgBrO_3}\) in molar solutions of \(\mathrm{MeNO_3}\)
| Electrolyte | \(C_{\mathrm{AgBrO_3}}\cdot 10^2\) | \(L_p' : L_p\)¹ | \(a\) |
|---|---|---|---|
| \(\mathrm{LiNO_3}\) | 1,32 | 2,54 | 3,12 |
| \(\mathrm{NaNO_3}\) | 1,48 | 3,22 | 2,09 |
| \(\mathrm{KNO_3}\) | 1,63 | 3,91 | 1,47 |
¹ \(L_p\) and \(L_p\) are the solubility products in salt and aqueous solutions, respectively.
Table 5
Change in the solubility product (\(L_p\)) of \(\mathrm{AgBrO_3}\) in \(\mathrm{MeNO_3}\) solutions
| \(C_{\mathrm{MeNO_3}}\), (mol/l) | \(\log C_{\mathrm{MeNO_3}}\) | \(\mathrm{LiNO_3}\): \(L_p \cdot 10^5\) | \(\mathrm{LiNO_3}\): \(\log L_p\) | \(\mathrm{NaNO_3}\): \(L_p \cdot 10^5\) | \(\mathrm{NaNO_3}\): \(\log L_p\) | \(\mathrm{KNO_3}\): \(L_p \cdot 10^5\) | \(\mathrm{KNO_3}\): \(\log L_p\) |
|---|---|---|---|---|---|---|---|
| 0.001 | −3.0 | 6.91 | −4.16 | 6.76 | −4.17 | 6.80 | −4.17 |
| 0.003 | −2.5 | 7.10 | −4.14 | 6.91 | −4.16 | 7.19 | −4.14 |
| 0.010 | −2.0 | 7.18 | −4.14 | 7.23 | −4.14 | 7.41 | −4.13 |
| 0.030 | −1.5 | 8.11 | −4.09 | 8.59 | −4.06 | 8.65 | −4.06 |
| 0.100 | −1.0 | 9.47 | −4.02 | 10.20 | −3.99 | 10.50 | −3.98 |
| 0.300 | −0.5 | 12.00 | −3.92 | 12.95 | −3.89 | 15.40 | −3.81 |
| 1.000 | −0.0 | 17.30 | −3.76 | 21.90 | −3.66 | 26.60 | −3.57 |
Table 6
Activity coefficient of \(\mathrm{AgBrO_3}\) in \(\mathrm{MeNO_3}\) solutions of various concentrations
| \(C_{\mathrm{MeNO_3}}\), (mol/l) | \(\gamma_{\mathrm{AgBrO_3}}\) in \(\mathrm{LiNO_3}\) solutions | \(\gamma_{\mathrm{AgBrO_3}}\) in \(\mathrm{NaNO_3}\) solutions | \(\gamma_{\mathrm{AgBrO_3}}\) in \(\mathrm{KNO_3}\) solutions | Mean value \(\gamma_{\mathrm{AgBrO_3}}\) |
|---|---|---|---|---|
| 0.000 | 0.900 | 0.900 | 0.900 | 0.900 |
| 0.001 | 0.903 | 0.901 | 0.899 | 0.901 |
| 0.003 | 0.894 | 0.891 | 0.888 | 0.891 |
| 0.010 | 0.870 | 0.865 | 0.862 | 0.866 |
| 0.030 | 0.826 | 0.817 | 0.810 | 0.818 |
| 0.100 | 0.751 | 0.730 | 0.717 | 0.733 |
| 0.300 | 0.662 | 0.626 | 0.599 | 0.629 |
| 1.000 | 0.564 | 0.501 | 0.455 | 0.507 |
With such a somewhat simplified approach, the solubility values found will not be very accurate; nevertheless, this approach is attractive because of its simplicity for a satisfactory solution of the question of the solubility of a precipitate under the action of electrolytes that have no ions in common with the precipitate. As is known, with the aid of the solubility-product rule, calculation of the solubility of a precipitate in such cases is generally impossible.
We consider it our pleasant duty to express our gratitude to Corresponding Member of the Academy of Sciences of the USSR, Prof. I. V. Tananaev, for providing the topic and for valuable advice in carrying out this work.
Conclusions
The solubility of silver bromate in solutions of alkali-metal nitrates—lithium, sodium, and potassium—has been investigated at concentrations from 0.001 to 1 mol/l at 25°.
It has been established that the solubility of the \(\mathrm{AgBrO_3}\) precipitate depends not only on the valence type of the external electrolyte, but also on the position of the corresponding elements in D. I. Mendeleev’s periodic system of elements.
The experimental data have been used to calculate the value \(a\), by means of which the activity coefficients of silver bromate in aqueous solutions of alkali-metal nitrates were determined according to the second variant of the Debye–Hückel formula.
The solubility product and the activity product of silver bromate in water and in aqueous solutions of alkali-metal nitrates have been determined.
lithium, sodium, and potassium, and the scales of variation of the solubility product of silver bromate as a function of the ionic strength of the solution are shown.
A formula has been proposed for calculating, with accuracy sufficient for practical purposes, the solubility of silver bromate in aqueous solutions of alkali-metal nitrates by means of the activity product of $\mathrm{AgBrO_3}$, found to be equal to $5.5 \cdot 10^{-5}$.
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