= Solution
\b[In a suitably dilute, nonoxidizing <chemical acid>, <hydrogen sulfide> precipitates the sufficiently insoluble sulfides, including those of <copper>, <lead>, <Mercury>, <bismuth>, <cadmium>, <arsenic>, <antimony> and <tin>.] “Sulphuretted hydrogen” is the historical name for <hydrogen sulfide>, $\mathrm H_2\mathrm S$. Oxidation state and acidity must be specified: the statement is not independent of every possible <chemical solution> condition.
|| Representative dissolved species
|| Principal precipitate
|| Usual appearance
| <Copper(II)>
| $\mathrm{CuS}$
| Black
| <Lead(II)>
| $\mathrm{PbS}$
| Black
| <Mercury(II)>
| $\mathrm{HgS}$
| Black when freshly precipitated
| <Bismuth(III)>
| $\mathrm{Bi}_2\mathrm S_3$
| Brown-black
| <Cadmium(II)>
| $\mathrm{CdS}$
| Yellow
| <Arsenic(III)>
| $\mathrm{As}_2\mathrm S_3$
| Yellow
| <Antimony(III)>
| $\mathrm{Sb}_2\mathrm S_3$
| Orange
| <Tin(II)>
| $\mathrm{SnS}$
| Brown or brown-black
| <Tin(IV)>
| $\mathrm{SnS}_2$
| Yellow
<Arsenic> is now classed as a <metalloid>; the historical analytical grouping was not a rigorous modern division into metals and nonmetals. Dissolved <silver> also gives black $\mathrm{Ag}_2\mathrm S$ in acidic conditions. It is often absent from the conventional second analytical group because a preceding <hydrochloric acid> treatment has already removed it as <silver chloride>. Suitable <gold> and <platinum> solutions can likewise yield dark sulfide precipitates, although their speciation and reduction chemistry make a simple universal colour rule inappropriate. Thus an analytical group is a separation convention, not a list of every ion that can possibly react.
For a divalent <metal ion>, a useful net balance is
$$
\boxed{\mathrm M^{2+}+\mathrm H_2\mathrm S\longrightarrow\mathrm{MS}\downarrow+2\mathrm H^+.}
$$
The selectivity follows from two coupled <chemical equilibrium> conditions. <Hydrogen sulfide> is a <weak acid>; increasing $[\mathrm H^+]$ suppresses its ionization. Approximately,
$$
[\mathrm S^{2-}]=\frac{K_{a1}K_{a2}[\mathrm H_2\mathrm S]}{[\mathrm H^+]^2}.
$$
Even this small sulfide activity can exceed the <solubility product> of very insoluble sulfides. More soluble sulfides, such as those of <zinc> and <manganese>, generally require less acidic or alkaline conditions. Consequently an excessively strong <chemical acid>, complex formation or an oxidizing medium can change the result. Colours alone are preliminary indications, not final identification.<Qualitative Analysis of Metal Ions in Solution by Williams College>[Experiment 2–3, Introduction and Group II discussion, printed II & III-1–3 and II & III-9–10]{ref}
The traditional preparation treats <iron(II) sulfide> with a dilute nonoxidizing <chemical acid>:
$$
\boxed{\mathrm{FeS}+2\mathrm{HCl}\longrightarrow\mathrm{FeCl}_2+\mathrm H_2\mathrm S.}
$$
A gas-generating flask with controlled acid entry and a delivery outlet conducts the product to the test solution. The apparatus must permit gas flow and contain the effluent rather than accumulating <pressure> in a sealed vessel. <Sulfuric acid> in dilute form can replace <hydrochloric acid>; concentrated oxidizing acids are unsuitable for this simple preparation. <Hydrogen sulfide> is acutely toxic, and loss of odour perception makes smell an unreliable warning: the historical apparatus describes a laboratory principle, not a household test.<Hydrogen sulfide by the National Institute for Occupational Safety and Health>[Physical Description, odour fatigue, Symptoms and Target Organs]{ref}
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