5.1: Separation of group III cations
Group II cations form sulfides that have very low solubility. After group II cations are removed under a low concentration of \(\ce{S^{2-}}\) in an acidic medium, the solution is made alkaline. Remember that like sulfides, hydroxides are also insoluble according to insoluble ions rule#1 of solubility guidelines described in chapter 1 states “Hydroxide (\(\ce{OH^{-}}\)) and sulfides (\(\ce{S^{2-}}\)) are insoluble except when the cation is a heavy alkaline earth metal ion: \(\ce{Ca^{2+}}\), \(\ce{Ba^{2+}}\), and \(\ce{Sr^{2+}}\), alkali metal ions, and ammonium ion.”
Table 1 lists solubility product constants of hydroxides of group III & IV cations at 25 o C, maximum hydroxide (\(\ce{OH^{-}}\)) concentration, and the maximum pH that can exist in a saturated solution containing 0.1M cation solutions that may be present in the test solution at this stage. It can be observed that the ions listed in table 1 will not precipitate as hydroxides during the precipitation of group II cations under the acidic pH range of 0.5 to 1.
\(\ce{Fe^{3+}}\) forms the most insoluble hydroxide, but it is reduced to \(\ce{Fe^{2+}}\) by \(\ce{H2S}\) during precipitation of group II cations:
\[\ce{2Fe^{3+}(aq) + S^{2-}(aq) <=> 2Fe^{2+}(aq) + S(s)}\nonumber\]
\(\ce{Fe^{3+}}\) may be present only if precipitation of group III starts from a fresh sample that has not been subjected to group II separation.
It can be observed from Table 1 that if the pH of the sample solution is increased to a range of 7 to 10, \(\ce{Fe^{3+}}\), \(\ce{Cr^{3+}}\), \(\ce{Ni^{2+}}\), and \(\ce{Fe^{2+}}\) will precipitate as \(\ce{Fe(OH)3(s, rusty)}\), \(\ce{Cr(OH)3(s, gray-green)}\), \(\ce{Ni(OH)2(s, green)}\), and \(\ce{Fe(OH)2(s, green)}\), leaving behind in the solution rest of the ions that may still be present at this stage. Group III comprise of , \(\ce{Fe^{3+}}\), \(\ce{Cr^{3+}}\), \(\ce{Ni^{2+}}\), and \(\ce{Fe^{2+}}\) ions.
|
Ion |
Salt |
K sp at 25 o C |
Minimum [OH - ] and pH needed to precipitate |
|---|---|---|---|
| \(\ce{Fe^{3+}}\) |
\(\ce{Fe(OH)3}\) |
\(\mathrm{K}_{\mathrm{sp}}=\left[\mathrm{Fe}^{3+}\right]\left[\mathrm{OH}^{-}\right]^{3}=2.8 \times 10^{-39}\) |
\(\left[\mathrm{OH}^{-}\right]=3.0 \times 10^{-13}~M=\mathrm{pH} ~1.5\) |
| \(\ce{Cr^{3+}}\) |
\(\ce{Cr(OH)3}\) |
\(\mathrm{K}_{\mathrm{sp}}=\left[\mathrm{Cr}^{3+}\right]\left[\mathrm{OH}^{-}\right]^{3}=1.0 \times 10^{-30}\) |
\(\left[\mathrm{OH}^{-}\right]=2.2 \times 10^{-10}~M=\mathrm{pH} ~4.3 \) |
| \(\ce{Ni^{2+}}\) |
\(\ce{Ni(OH)2}\) |
\( \mathrm{K}_{\mathrm{sp}}=\left[\mathrm{Ni}^{2+}\right]\left[\mathrm{OH}^{-}\right]^{2}=5.5 \times 10^{-16}\) |
\(\left[\mathrm{OH}^{-}\right]==7.4 \times 10^{-9}~M=\mathrm{pH} ~4.3 \) |
| \(\ce{Fe^{2+}}\) |
\(\ce{Fe(OH)2}\) |
\( \mathrm{K}_{\mathrm{sp}}=\left[\mathrm{Fe}^{2+}\right]\left[\mathrm{OH}^{-}\right]^{2}=4.9 \times 10^{-17}\) |
\( \left[\mathrm{OH}^{-}\right]=2.2 \times 10^{-9}~M=\mathrm{pH} ~5.6\) |
|
\(\ce{Ca^{2+}}\) |
\(\ce{Ca(OH)2}\) |
\( \mathrm{K}_{\mathrm{sp}}=\left[\mathrm{Ca}^{2+}\right]\left[\mathrm{OH}^{-}\right]^{2}=5.0 \times 10^{-6}\) |
\(\left[\mathrm{OH}^{-}\right]=7.1 \times 10^{-4}~M=\mathrm{pH} \mathrm{} ~10.9 \) |
|
\(\ce{Ba^{2+}}\) |
\(\ce{Ba(OH)2}\) |
\( \mathrm{K}_{\mathrm{sp}}=\left[\mathrm{Ba}^{2+}\right]\left[\mathrm{OH}^{-}\right]^{2}=2.6 \times 10^{-4}\) |
\(\left[\mathrm{OH}^{-}\right]=5.1 \times 10^{-3}~M=\mathrm{pH} \mathrm{} 11.7 \) |
- * Following cations that may be present in the initial solution are not listed in this table due to the reason: i) \(\ce{Pb^{2+}}\), \(\ce{Hg2^{2+}}\), and \(\ce{Ag^{+}}\) are already removed as chloride precipitates of group I cations, ii) \(\ce{Sn^{4+}}\), \(\ce{Cd^{2+}}\), \(\ce{Cu^{2+}}\), and \(\ce{Bi^{3+}}\) has been removed as group II sulfides under pH 0.5 to 1, iii) \(\ce{Na^{+}}\) and \(\ce{K^{+}}\) form soluble compounds with all anions according to rule#1 of solubility described in chapter 1. Source: Engineering ToolBox, (2017). Solubility product constants. [online] Available at: https://www.engineeringtoolbox.com/s...sp-d_1952.html [Accessed Feb. 5 th , 2022]
Buffers, that resist change in pH are employed in such a situation where pH needs to be maintained in a narrow range. Buffers are a mixture of a weak acid and its conjugate base or a mixture of a weak base and its conjugate acid. Ammonia (\(\ce{NH3}\)), i.e., a week base and ammonium ion (\(\ce{NH4^{+}}\)) is its conjugate acid.
The \(\ce{NH3}\)/\(\ce{NH4^{+}}\) is a suitable buffer that can maintain pH of around 9. The buffer is prepared by adding 2 drops of 6M \(\ce{HCl}\) into 15 drops of the sample and then adding 6M \(\ce{NH3}\) drop by drop to neutralize the acid.
\[\ce{HCl(aq) + H2O(l) -> H3O^{+}(aq) + Cl^{-}(aq)}\nonumber\]
\[\ce{NH3(aq) + H3O^{+}(aq) -> NH4^{+}(aq) + H2O(l)}\nonumber\]
\[\text{Overall reaction:} \ce{~HCl(aq) + NH3(aq) -> NH4^{+}(aq) + Cl^{-}(aq)}\nonumber\]
Then 5 drops more of 6M \(\ce{NH3}\) are added after the \(\ce{HCl}\) has been neutralized to make a mixture of \(\ce{NH3}\) and \(\ce{NH4^{+}}\) that maintains pH ~9 and OH - at around 1 x 10 -5 M.
The group III cations precipitate at this stage as hydroxides, as shown in Figure \(\PageIndex{1}\), except \(\ce{Ni^{2+}}\):
\[\ce{Fe^{3+}(aq) + 3OH^{-}(aq) -> Fe(OH)3(s, reddish-brown ~or ~rusty)(v),}\nonumber\]
\[\ce{Cr^{3+}(aq) + 3OH^{-}(aq) -> Cr(OH)3(s, gray-green)(v),}\nonumber\]
\[\ce{Fe^{2+}(aq) + 2OH^{-}(aq) -> Fe(OH)2(s, green)(v).}\nonumber\]
The concentration of \(\ce{Fe^{2+}}\), i.e., the most soluble hydroxide of group III cations, is reduced by more than 99.99%, i.e., from 0.1M to 4.9 x 10 -7 M when pH is increased to 9 and \(\ce{OH^{-}}\) concentration is increased to 1 x 10 -5 M:
\[\mathrm{Fe}^{2+}=\frac{\mathrm{K}_{\mathrm{sp}}}{\left[\mathrm{OH}^{-}\right]^{2}}=\frac{4.9 \times 10^{-17}}{\left(1 \times 10^{-5}\right)^{2}}=4.9 \times 10^{-7} \mathrm{~M}\nonumber\]
Nickle ion is not precipitated at this stage as it forms soluble coordination cation \(\ce{[Ni(NH3)6]^{2+}}\) with ammonia:
\[\ce{Ni^{2+}(aq, green) + 6NH3(aq) <=> Ni(NH3)6(aq, blue)}\nonumber\]
Therefore, \(\ce{S^{2-}}\) is introduced by adding thioacetamide and heating the mixture in a boiling water bath. Decomposition of thioacetamide produces ~0.01M \(\ce{H2S}\):
\[\ce{CH3CSNH2(aq) + 2H2O(l) <=> CH3COO^{-}(aq) + NH4^{+}(aq) + H2S(aq)}\nonumber\]
Nearly all of the \(\ce{H2S}\) dissociates to form ~0.01M \(\ce{S^{2-}}\) at pH ~9:
\[\ce{H2S(aq) + 2H2O(l) <=>2H3O^{+}(aq) + S^{2-}(aq)}\quad K_a = \frac{\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]^{2}\left[\mathrm{~S}^{2-}\right]}{\left[\mathrm{H}_{2} \mathrm{~S}\right]}=1.3 \times 10^{-20}\nonumber\]
The ammonia complex of nickel, i.e., \(\ce{[Ni(NH3)6]^{2+}}\) precipitates out as \(\ce{NiS}\), and, at the same time, \(\ce{Fe(OH)3}\) and \(\ce{Fe(OH)2}\) also convert to \(\ce{Fe2S3}\) and \(\ce{FeS}\):
\[\ce{Ni(NH3)6^{2+}(aq, blue) + S^{2-}(aq) <=> NiS(s, black) + 6NH3(aq)}\nonumber\]
\[\ce{2Fe(OH)3(s, reddish-brown) + 3S^{2-}(aq) <=> Fe2S3(s, yellow-green) + 6OH^{-}(aq)}\nonumber\]
\[\ce{Fe(OH)2(s, geen) + S^{2-}(aq) <=> FeS(s, black) + 2OH^{-}(aq)}\nonumber\]
Chromium remains as \(\ce{Cr(OH)3}\) precipitate because chromium sulfide is unstable in water.
Group III precipitates, i.e., \(\ce{Cr(OH)3(s, gray-green)}\), \(\ce{NiS(s, black)}\), \(\ce{Fe2S3(s, yellow-green)}\), and \(\ce{FeS(s, black)}\) in the mixture are separated as precipitates, and the rest of the ions, i.e, \(\ce{Ca^{2+}}\), \(\ce{Ba^{2+}}\), \(\ce{Na^{+}}\) and \(\ce{K^{+}}\), etc. remain dissolved in the supernatant, as shown in Figure \(\PageIndex{2}\). The color of the precipitate does not give a clear indication of what ions are present at this stage as several species of different colors may be mixed at this stage.