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3.1: Separation of group I cations

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    Selective precipitation of a set of group I, i.e., lead(II) (\(\ce{Pb^{2+}}\)), mercury(I) (\(\ce{Hg2^{2+}}\)), and silver(I) (\(\ce{Ag^{+}}\)) is based on soluble ions rule#3 in the solubility guidelines in section 1.1 which states "Salts of chloride (\(\ce{Cl^{-}}\)), bromide ( \(\ce{Br^{-}}\)), and Iodide ( \(\ce{I^{-}}\)) are soluble, except when the cation is Lead ( \(\ce{Pb^{2+}}\)), Mercury ( \(\ce{Hg2^{2+}}\)), or Silver ( \(\ce{Ag^{+}}\)). The best source of \(\ce{Cl^{-}}\) for precipitating group 1 cations from a test solution is \(\ce{HCl}\), because it is a strong acid that completely dissociates in water producing \(\ce{Cl^{-}}\) and \(\ce{H3O^{+}}\) ions, both do not get involved in any undesired reactions under the conditions.

    The \(\ce{K_{sp}}\) expression is used to calculate \(\ce{Cl^{-}}\) that will be sufficient to precipitate group 1 cations. The molar concentration of chloride ions i.e., [\(\ce{Cl^{-}}\)], in moles/liter in a saturated solution of the ionic compound can be calculated by rearranging their respective \(\ce{K_{sp}}\) expression. For example, for \(\ce{AgCl}\) solution, \(\mathrm{K}_{\mathrm{sp}}=\left[\mathrm{Ag}^{+}\right]\left[\mathrm{Cl}^{-}\right]\) rearranges to:

    \[\left[\mathrm{Cl}^{-}\right]=K_{s p} /\left[\mathrm{Ag}^{+}\right]\nonumber\]

    and for \(\ce{PbCl2}\) solution, Ksp = [Pb2+][Cl-]2 rearranges to:

    \[\left[C l^{-}\right]=\sqrt{K_{s p} /\left[P b^{2+}\right]}\nonumber\]

    The concentration of ions in the unknown sample are ~0.1 M. Plugging in 0.1M value for \(\ce{Pb^{2+}}\) in the above equation shows that [\(\ce{Cl^{-}}\)] in a saturated solution having 0.1M \(\ce{Pb^{2+}}\) is 1.3 x 10-2M:

    \[\left[C l^{-}\right]=\sqrt{K_{s p} /\left[P b^{2+}\right]}=\sqrt[2]{1.6 \times 10^{-5} / 0.1}=1.3 \times 10^{-2} \mathrm{M}\nonumber\]

    It means \(\ce{Cl^{-}}\) concentration up to 1.3 x 10-2M will not cause precipitation from 0.1M \(\ce{Pb^{2+}}\) solution. Increasing \(\ce{Cl^{-}}\) above 0.013M will remove \(\ce{Pb^{2+}}\) from the solution as a \(\ce{PbCl2}\) precipitate. If 99.9% removal is desired, then 1.0 x 10-4 M \(\ce{Pb^{2+}}\) will be left in the solution and the [\(\ce{Cl^{-}}\)] have to be raised to 0.40 M:

    \[\left[C l^{-}\right]=\sqrt[2]{K_{s p} /\left[P b^{2+}\right]}=\sqrt[2]{1.6 \times 10^{-5} / 1.0 \times 10^{-4}}=0.40 \mathrm{M}\nonumber\]

    The solubility of \(\ce{Hg2Cl2}\) and \(\ce{AgCl}\) is less than that of \(\ce{PbCl2}\). So, a 0.40M \(\ce{Cl^{-}}\) will remove more than 99.9% of \(\ce{Hg2^{2+}}\) and \(\ce{Ag^{+}}\) from the solution.

    A sample of 20 drops of the aqueous solution is about 1 mL. In these experiments, ~15 drops of the test solution are collected in a test tube and 3 to 4 drops of 6M \(\ce{HCl}\) are mixed with the solution. This results in about 0.9 mL total solution containing 1 to 1.3 M \(\ce{Cl^{-}}\), which is more than twice the concentration needed to precipitate out 99.9% of group 1 cations.

    A concentrated reagent (6M \(\ce{HCl}\)) is used to minimize the dilution of the test sample because the solution is centrifuged and the supernatant that is separated by decantation is used to analyze the remaining cations. A 12M \(\ce{HCl}\) is available, but it is not used because it is a more hazardous reagent due to being more concentrated strong acid and also because if \(\ce{Cl^{-}}\) concentration is raised to 5M or higher in the test solution, it can re-dissolve \(\ce{AgCl}\), by forming water-soluble [\(\ce{AgCl2}\)]- complex ion.

    The addition of \(\ce{HCl}\) causes precipitation of group 1 cation as milky white suspension as shown in Figure \(\PageIndex{1}\) and by chemical reaction equations below. The precipitates can be separated by gravity filtration, but more effective separation can be achieved by subjecting the suspension to centrifuge in a test tube. Centrifugal force forces the solid suspension to settle and pack at the bottom of the test tube from which the clear solution, called supernatant, can be poured out -a process called decantation. The precipitate is resuspended in pure water by stirring with a clean glass rod, centrifuged, and decanted again to wash out any residual impurities. The washed precipitate is used to separate and confirm the group 1 cations and the supernatant is saved for analysis of group 2, 3, 4, and 5 cations.

    \[\ce{ Pb^{2+}(aq) + 2Cl^{-}(aq) <=> PbCl2(s)(v)}\nonumber\]

    \[\ce{Hg2^{2+}(aq) + 2Cl^{-}(aq) <=> Hg2Cl2(s)(v)}\nonumber\]

    \[\ce{ Ag^{+}(aq) + Cl^{-}(aq) <=> AgCl(s)(v)}\nonumber\]

    White precipitates of group I cations
    The precipitate separated by centrifugation from the supernatant
    The precipitate resuspended in an inert solvent for washing
    Figure \(\PageIndex{1}\): Precipitates of group I cations formed as a suspension upon addition of \(\ce{HCl}\) drops to the test sample (left), the precipitate separated as a pellet at the bottom of the test tube by centrifugation with a clear supernatant at the top (middle), and the precipitate resuspended in washing liquid (right).

    This page titled 3.1: Separation of group I cations is shared under a Public Domain license and was authored, remixed, and/or curated by Muhammad Arif Malik.

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