4.1: Precipitation of group II cations

The bases of Group II cations separation

The solubility guideline#1 of insoluble ions states “Hydroxide (\(\ce{OH^{-}}\)) and sulfides (\(\ce{S^{2-}}\)) are insoluble except when the cation is alkali metal, ammonia, or a heavy alkaline earth metal ion, i.e., \(\ce{Ca^{2+}}\), \(\ce{Ba^{2+}}\), and \(\ce{Sr^{2+}}\)”. The sulfide of \(\ce{Cr^{3+}}\) is also in the exceptions list as its sulfide is unstable in water. It is obvious that the number of insoluble sulfides and hydroxides is large. The solution is made acidic to decrease [\(\ce{OH^{-}}\)] to below the level that can cause precipitation of any ion. The [\(\ce{S^{2-}}\)] also remains low due to the common ion effect of \(\ce{H3O^{+}}\) in the acidic medium as explained in the next section. Therefore among the insoluble sulfides, only those that have very low solubility limits are selectively precipitated. These include \(\ce{Bi^{3+}}\), \(\ce{Cd^{2+}}\), \(\ce{Cu^{2+}}\), and \(\ce{Sn^{4+}}\) among the cations selected in this study that are left in the solution after group I cations have been separated. Group II comprise of \(\ce{Bi^{3+}}\), \(\ce{Cd^{2+}}\), \(\ce{Cu^{2+}}\), and \(\ce{Sn^{4+}}\).

Precipitation of group II cations

Among the ions in the initial solution after removal of group I cations, the following ions form insoluble sulfides: \(\ce{Bi^{3+}}\), \(\ce{Cd^{2+}}\), \(\ce{Cu^{2+}}\), \(\ce{Fe^{2+}}\), \(\ce{Fe^{3+}}\), \(\ce{Ni^{2+}}\), and \(\ce{Sn^{4+}}\). Among these, \(\ce{Bi^{3+}}\), \(\ce{Cd^{2+}}\), \(\ce{Cu^{2+}}\), and \(\ce{Sn^{4+}}\) are in group II that form very insoluble sulfides, and \(\ce{Cr^{3+}}\), \(\ce{Fe^{2+}}\), \(\ce{Fe^{3+}}\), and \(\ce{Ni^{2+}}\) are in group III form insoluble hydroxides and sulfides in basic medium, as reflected by their solubility product constants (\(\ce{K_{sp}}\)) listed in Table 1. The minimum concentration of \(\ce{S^{2-}}\) needed to start precipitation of the cation can be calculated from the \(\ce{K_{sp}}\) expressions as shown in Table 1. It can be observed from Table 1 that there is a huge difference in the minimum \(\ce{S^{2-}}\) concentration (1.8 x 10^-20M) needed to precipitate \(\ce{Ni^{2+}}\) -the least soluble sulfide of group III and \(\ce{Cd^{2+}}\) (7.8 x 10^-26) -the most soluble sulfide of group II. If the \(\ce{S^{2-}}\) is kept more than 1.8 x 10^-20 M but less than 7.8 x 10^-26 M group II cations will selectively precipitate while group III cations and the rest of the cations will remain dissolved.

Table 1: Solubility product constants of insoluble sulfides of group II, and group III and minimum sulfide ion concentration needed to start precipitation from 0.1M cation solution*.

Ion	Sulfide	Ksp at 25 ^oC	Minimum [S^-2] needed to precipitate
\(\ce{Fe^{2+}}\)	\(\ce{FeS}\)	\(\mathrm{K}_{\mathrm{sp}}=\left[\mathrm{Fe}^{2+}\right]\left[\mathrm{S}^{2-}\right]=4.9 \times 10^{-18}\)	\(\left[\mathrm{S}^{-2}\right]=\mathrm{K}_{\mathrm{sp}} /\left[\mathrm{Fe}^{2+}\right]=4.9 \times 10^{-17}\)
\(\ce{Ni^{2+}}\)	\(\ce{NiS}\)	\(\mathrm{K}_{\mathrm{sp}}=\left[\mathrm{Ni}^{2+}\right]\left[\mathrm{S}^{2-}\right]=1.8 \times 10^{-21}\)	\(\left[\mathrm{S}^{-2}\right]=\mathrm{K}_{\mathrm{sp}} /\left[\mathrm{Ni}^{2+}\right]=1.8 \times 10^{-20}\)
\(\ce{Cd^{2+}}\)	\(\ce{NiS}\)	\(\mathrm{K}_{\mathrm{sp}}=\left[\mathrm{Cd}^{2+}\right]\left[\mathrm{S}^{2-}\right]=7.8 \times 10^{-27}\)	\(\left[\mathrm{S}^{-2}\right]=\mathrm{K}_{\mathrm{sp}} /\left[\mathrm{Cd}^{2+}\right]=7.8 \times 10^{-26}\)
\(\ce{Bi^{3+}}\)	\(\ce{Ba2S3}\)	\(\mathrm{K}_{\mathrm{sp}}=\left[\mathrm{Bi}^{3+}\right]^{2}\left[\mathrm{~S}^{2-}\right]^{3}=6.8 \times 10^{-97}\)	\(\left[\mathrm{S}^{-2}\right]=\sqrt[3]{\mathrm{K}_{\mathrm{sp}} /\left[\mathrm{Bi}^{3+}\right]^{2}}=4.1 \times 10^{-32}\)
\(\ce{Sn^{4+}}\)	\(\ce{SnS2}\)	\(\mathrm{K}_{\mathrm{sp}}=\left[\mathrm{Sn}^{4+}\right]\left[\mathrm{S}^{2-}\right]^{2}=1.0 \times 10^{-70}\)	\(\left[\mathrm{S}^{-2}\right]=\sqrt[2]{\mathrm{K}_{\mathrm{sp}} /\left[\mathrm{Sn}^{4+}\right]}=3.2 \times 10^{-35}\)
\(\ce{Cu^{2+}}\)	\(\ce{CuS}\)	\(\mathrm{K}_{\mathrm{sp}}=\left[\mathrm{Cu}^{2+}\right]\left[\mathrm{S}^{2-}\right]=8.7 \times 10^{-36}\)	\(\left[\mathrm{S}^{-2}\right]=\mathrm{K}_{\mathrm{sp}} /\left[\mathrm{Fe}^{2+}\right]=8.7 \times 10^{-35}\)

• * Following cations that may be present the initial solution are not listed in this table due to the following reasons: i) group I cations, i.e., \(\ce{Pb^{2+}}\), \(\ce{Hg2^{2+}}\), and \(\ce{Ag^{+}}\) are already removed, ii) \(\ce{Ca^{2+}}\) and \(\ce{Ba^{2+}}\) for soluble sulfides, iii) sulfide of \(\ce{Cr^{3+}}\) is not stable in water, and iv) \(\ce{Fe^{3+}}\) is reduced to \(\ce{Fe^{2+}}\) by \(\ce{H2S}\) in acidic medium: \(\ce{2Fe^{3+}(aq) + S^{2-} <=> 2Fe^{2+}(aq) + S(s)}\). Source of \(\ce{K_{sp}}\) values: chem 202 lab manual, 2008, by Michael Stranz, cengag learning, ISBN 13: 978-0-534-66904-1

Source of \(\ce{S^{2-}}\) is \(\ce{H2S}\) gas -a week diprotic acid that dissociated in water by the following equilibrium reactions:

\[\ce{H2S(g) + H2O(l) <=> H3O^{+}(aq) + HS^{-}(aq)}\quad \mathrm{K}_{\mathrm{a} 1}=\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\left[\mathrm{HS}^{-}\right] /\left[\mathrm{H}_{2} \mathrm{~S}\right]=1.0 \times 10^{-7}\nonumber\]

\[\ce{HS^{-}(aq) + H2O(l) <=> H3O^{+}(aq) + S^{2-}(aq)}\quad \mathrm{K}_{\mathrm{a} 2}=\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\left[\mathrm{S}^{2-}\right] /\left[\mathrm{HS}^{-}\right]=1.3 \times 10^{-13}\nonumber\]

\[\text{Overall reaction: }\ce{H2S(g) + 2H2O(l) <=> 2H3O^{+}(aq) + S^{2-}(aq)}\quad\mathrm{K}_{\mathrm{a}}=\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]^{2}\left[\mathrm{~S}^{2-}\right] /\left[\mathrm{H}_{2} \mathrm{~S}\right]=1.3 \times 10^{-20}\nonumber\]

Extent of \(\ce{H2S}\) dissociation, and, consequently, the concentration of \(\ce{S^{2-}}\) produced is dependent on \(\ce{H3O^{+}}\):

\[K_a = \frac{\ce{[H3O^{+}]^{2}[S^{2-}]}}{\ce{[H2S]}}\quad\quad\text{ rearranges to: }\quad\quad\ce{[S^{2-}]} = \frac{K_{a}\ce{[H2S]}}{\ce{[H3O^{+}]^{2}}}\nonumber\]

It is obvious from the above formula that \(\ce{[S^{2-}]}\) is dependent on \(\ce{[H3O^{+}]}\), which is related to pH ( \(pH = Log\frac{1}{\ce{[H3O^{+}]}} = \text{-Log}\ce{[H3O^{+}]}\). Therefore, \(\ce{[S^{2-}]}\) can be controlled by adjusting the pH.

\(\ce{H2S}\) is a toxic gas. To minimize the exposure, \(\ce{H2S}\) is produced in-situ by decomposition of thioacetamide (\(\ce{CH3CSNH2}\)) in water:

\[\ce{CH3CSNH2(aq) + 2H2O <=> CH3COO^{-} + NH4^{+}(aq) + H2S(aq)}\nonumber\]

The decomposition of thioacetamide is an endothermic reaction, which, according to Le Chatelier's principle, moves in the forward direction upon heating. An aqueous solution of thioacetamide is heated in a boiling water bath in a fume hood producing ~0.01M \(\ce{H2S}\) solution.

Rearranging acid dissociation constant of \(\ce{H2S}\) and plugging in 0.01M \(\ce{H2S}\) in the rearranged formula allows calculating \(\ce{S^{2-}}\) concentration at various concentrations of \(\ce{H3O^{+}}\), i.e., at various pH values:

\[\ce{[S^{2-}]} = \frac{K_{a}\ce{[H2S]}}{\ce{[H3O^{+}]^{2}}} = \frac{1.3\times10^{-20}\times0.01}{\ce{[H3O^{+}]^{2}}} = \frac{1.3\times10^{-22}}{\ce{[H3O^{+}]^{2}}}\nonumber\]

It shows that \(\ce{S^{2-}}\) concentration can be varied by [\(\ce{H3O^{+}}\)], i.e., by varying pH. At pH 1 and 0, \(\ce{H3O^{+}}\) is 0.10 M and 1.0 M, respectively, that produces [\(\ce{S^{2-}}\)] in the range of 1.3 x 10^-20 M S^2- and 1.3 x 10^-22 M S^2-:

\[\ce{[S^{2-}]} = \frac{1.3\times10^{-22}}{(0.10)^{2}} = 1.3\times10^{-20} ~M\quad\quad\text{ and }\quad\quad\ce{[S^{2-}]} = \frac{1.3\times10^{-22}}{(1.0)^{2}} = 1.3\times10^{-22}~M\nonumber\]

This range of [\(\ce{S^{2-}}\)] is less than the solubility limit of \(\ce{Ni^{2+}}\) -the least soluble cation of group III but more than the solubility limit of \(\ce{Cd^{2+}}\) -the most soluble cation of group II. If pH of the test solution is maintained between 0 and 1, group II cations will precipitate and group III and higher group cations will remain dissolved. At pH 0.5, S^2- is 5.2 x 10^-22 M that will precipitate more than 99.99% \(\ce{Cd^{2+}}\):

\[\mathrm{K}_{\mathrm{sp}}=\left[\mathrm{Cd}^{2+}\right]\left[\mathrm{S}^{2-}\right]=7.8 \times 10^{-27}\quad\quad\text{gives:}\quad\quad\ce{[Cd^{2+}]} = \frac{7.8\times 10^{-27}}{\ce{[S^{-2}]}} = \frac{7.8\times 10^{-27}}{5.2\times10^{-22}} = 1.5\times10^{-5}~M\nonumber\]

, which is 0.0002% of the initial [\(\ce{Cd^{2+}}\)].

The supernatant after removal of group I chlorides is usually within the pH range of 0.5 ±0.3, which is the appropriate pH for precipitation of group II cations under the conditions of this study. If the pH of the test sample is outside this range, the pH can be increased to ~0.5 by adding 0.5M \(\ce{NH3(aq)}\) drop by drop under stirring. Determine pH by using a pH paper after each drop of 0.5M \(\ce{NH3(aq)}\)_­ is added and thoroughly mixed. Keep in mind that \(\ce{NH3}\) solution in water is also labeled as \(\ce{NH4OH}\). Similarly, the pH can be decreased to ~0.5 by adding 0.5M \(\ce{HCl(aq)}\) drop by drop under stirring. Determine pH by using a pH paper after each drop of 0.5M \(\ce{HCl(aq)}\)_­ is added and thoroughly mixed.

Thioacetamide reagent is added to the test solution at pH ~0.5 and heated in a boiling water bath to precipitate out group II cations.

The precipitates include \(\ce{SnS2}\) (yellow), \(\ce{CdS}\) (yellow-orange), \(\ce{CuS}\) (Black-brown), \(\ce{Bi2S3}\) (black), formed by the following precipitation reactions:

\[\ce{SnCl6^{2-}(aq) + 2S^{2-}(aq) <=> 6Cl^{-}(aq) + SnS2(s, yellow)}\nonumber\]

\[\ce{Cd^{2+}(aq) + S^{2-}(aq) <=> CdS(s, yellow-orange)}\nonumber\]

\[\ce{Cu^{2+}(aq) + S^{2-}(aq) <=> CuS(s, black-brown)}\nonumber\]

\[\ce{2Bi^{3+}(aq) + 3S^{2-}(aq) <=> Bi2S3(s, black)}\nonumber\]

The overall color of the combined precipitate may vary depending on its composition. Black color dominates, i.e., if all precipitates are present, the color of the mixture will be black as shown in Figure \(\PageIndex{1}\).

Figure \(\PageIndex{1}\): Mixture of Group II precipitates, i.e., \(\ce{SnS2}\) (yellow), \(\ce{CdS}\) (yellow-orange), \(\ce{CuS}\) (Black-brown), and \(\ce{Bi2S3}\) (black).

The solution is cooled to room temperature by using a room temperature water bath. Cooling helps precipitation of \(\ce{CdS}\). A drop of 0.5 M \(\ce{NH3(aq)}\) is added while stirring, which promotes precipitation of \(\ce{CdS}\) and \(\ce{CnS2}\), as both tend to stay dissolved in a supersaturated solution. The mixture is centrifuged and decanted to separate the supernatant that is used for the analysis of group III and higher group cations. The precipitate is washed with 0.1M \(\ce{NH4Cl}\) solution and the washed precipitate is used to separate and confirm individual cations of group II.