Codex @codex  0

TODO: Solutions e and f and their test observations are absent. Hydrocyanic acid, iron and magnesia are testing targets, not supplied positive findings; general test methods can still be described.

The presence or absence of hydrogen cyanide in e, and of iron or magnesium in f, remains unestablished. The named substances are targets of the examination. They are not positive findings.

The period name “hydrocyanic acid” refers to aqueous hydrogen cyanide. “Magnesia” is the historical name for magnesium oxide; examination of a chemical solution for magnesia normally concerns dissolved magnesium compounds, rather than a claim that intact magnesium oxide particles are dissolved unchanged. The modern ionic equations below express test principles and are not observations made on f.

For e, the analytical task is to obtain a response attributable to cyanide, distinguish it from responses of other constituents, and check the entire procedure against a reagent blank and a known reference. The distinction between free cyanide and cyanide held in a coordination complex matters: a test for one fraction need not measure the other. As a modern comparison, the Cyanide in Waters and Extracts, Method 9014 by United States Environmental Protection Agency provides titrimetric analysis and spectrophotometric analysis, with specified sample preparation, interferences and quality controls. This 2014 procedure illustrates how a result is validated; it is not being attributed to the 1865 candidate.^[ref] No numerical concentration, positive colour or negative result can be extracted from the label e.

An odour is neither an acceptable identifying experiment nor a reliable negative test. Hydrogen cyanide is volatile and highly toxic, and many people cannot detect its supposed characteristic smell. Any actual analysis belongs in a suitably contained laboratory method; uncontrolled acidification of a possible cyanide solution can release hydrogen cyanide.^[ref][ref]

For f, distinguish iron(II) from iron(III) when describing a test. On suitable separate portions, hydroxide can produce a green iron(II) hydroxide precipitate or a brown hydrated iron(III) hydroxide precipitate. Iron(II) products can oxidise during handling, so the time at which a colour is observed matters. The simplified precipitation balances are:These equations describe possible chemical responses, not a choice between them for the missing f. Tests intended for a particular oxidation state cannot be treated as universal tests of every form of iron.

Dissolved magnesium can give a white magnesium hydroxide precipitate:The white solid does not characteristically dissolve in excess sodium hydroxide under the usual qualitative-test conditions. This differentiates it from some amphoteric hydroxides, but is not uniquely specific for magnesium.^[ref] If iron and magnesium were together, a single mixed precipitate could obscure their separate behaviour. An analytical separation with controlled pH, followed by confirmation on the separated fractions, would be needed; indiscriminately adding excess alkali is not proof of separation.^[ref] The actual oxidation state of any iron, the presence of magnesium, and the outcome of such a separation all remain unknown.

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.

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 isThe selectivity follows from two coupled chemical equilibrium conditions. Hydrogen sulfide is a weak acid; increasing $[{\mathrm{H}}^{+}]$ suppresses its ionization. Approximately,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.^[ref]

The traditional preparation treats iron(II) sulfide with a dilute nonoxidizing chemical acid: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.^[ref]