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A chemical compound whose solid composition includes water of crystallization.

In modern notation the formulas are HCN, HCOOH, ${\mathrm{H}}_2{\mathrm{C}}_2{\mathrm{O}}_4$, $\mathrm{H}{\mathrm{g}}_2\mathrm{C}{\mathrm{l}}_2$, $\mathrm{HgC}{\mathrm{l}}_2$, $\mathrm{A}{\mathrm{s}}_2{\mathrm{O}}_3$, a hydrated potassium antimony tartrate, and $\mathrm{KCl}{\mathrm{O}}_3$, respectively. Several historical names need qualification.

The chemical formulas above use modern atomic weights and distinguish molecular formulas, formula units and hydrates. Historical equivalent notation sometimes halves the subscripts attached to divalent chemical elements or writes water as $\mathrm{HO}$. Such formulas must be read together with their contemporary atomic weights, rather than treated as inconsistent modern molecular formulas.

A suitable historical answer gives both the chemical transformation and the means of isolating a reasonably pure product. The following are preparative principles, with modern chemical formulas; they are not instructions for self-preparation of medicines. In particular, hydrogen cyanide, mercury compounds and arsenic compounds present serious toxicity hazards. Hydrogen cyanide is extremely volatile, can be absorbed by several routes, and its odour is an unreliable warning.^[ref]

For hydrocyanic acid, a customary early nineteenth-century route is the controlled distillation of potassium ferrocyanide with dilute sulfuric acid and water. Hydrogen cyanide passes into the distillate; nonvolatile salts remain behind. A generating flask, efficient laboratory condenser and chilled receiver are needed because hydrogen cyanide is readily lost. The contemporary account of hydrocyanic acid preparation in Pharmacopoeia Castrensis Ruthenica in The British and Foreign Medical Review (1842) specifically describes the ferrocyanide–sulfuric acid method and its losses on imperfect condensation. The concentration must be determined analytically before any defined dilution: a nominal starting mixture does not establish the final strength.^[ref] At the level of the released cyanide, the essential acid-base reaction is $\mathrm{C}{\mathrm{N}}^{-}+{\mathrm{H}}^{+}\to \mathrm{HCN}$. This equation alone is not a complete equation for decomposing the stable ferrocyanide coordination complex. Concentrated sulfuric acid can lead to different decomposition chemistry, so the specified dilution is meaningful.

For formic acid, gently heat oxalic acid with glycerol and some water, then separate the aqueous formic acid by distillation. The useful net transformation isThe glycerol participates through intermediate esters and is regenerated rather than consumed in this net equation. Marcellin Berthelot's 1856 account describes the carbon dioxide evolution, subsequent water addition and distillation, and reuse of the residual glycerol. Overheating is disadvantageous because the formic acid can decompose further to carbon monoxide and water. The method therefore separates a volatile acid from a much less volatile carrier; it is not merely the dry destructive heating of oxalic acid.^[ref]

For oxalic acid, oxidize a carbohydrate, such as sucrose or starch, with nitric acid; concentrate the resulting chemical solution appropriately and cool it to obtain crystals, followed by recrystallization. Nitrogen oxides evolve, and excessive oxidation can carry carbon onward to carbon dioxide, so an idealized single-product equation would conceal losses. The use of starch and nitric acid for oxalic acid production is explicitly documented in an 1850 account of nitric oxide recovery.^[ref] An alternative already established before 1864 is alkaline treatment of sawdust, producing an oxalate which can be purified through sparingly soluble calcium oxalate.^[ref] Decomposing the isolated calcium oxalate with sulfuric acid separates much of the calcium as calcium sulfate:Filter off the calcium sulfate and crystallize the oxalic acid from the filtrate. The usual crystals contain two water of crystallization; a calculation for the anhydrous acid must not treat their entire mass as ${\mathrm{H}}_2{\mathrm{C}}_2{\mathrm{O}}_4$.

For calomel, intimately combine mercury(II) sulfate, metallic Mercury and dry sodium chloride, then collect the volatile mercury(I) chloride in a cooler receiving chamber. The net change isThe extra metallic Mercury lowers the average oxidation state from +2 to +1. The product is purified by appropriate sublimation and washing to remove soluble contamination, particularly mercury(II) chloride. Both the chemical composition and the fine subdivision of the product matter: condensation through a chamber can give a powder rather than an adherent crystalline crust. This physical part of calomel preparation is described in the 1843 primary account in Polytechnisches Journal volume 87 (1843); the account's historical therapeutic claims are not adopted here.^[ref]

For corrosive sublimate, heat dry mercury(II) sulfate with sodium chloride and collect the mercury(II) chloride in the cooler part of a suitable sublimation apparatus:Here no metallic Mercury is added. Nonvolatile sodium sulfate stays behind, while mercury(II) chloride is transferred and condensed. Further sublimation or recrystallization separates impurities. ^[ref] The crucial distinction from calomel preparation is whether Mercury remains in the +2 oxidation state or is converted to +1.

For the solid called arsenious acid, roast an arsenic-bearing ore in air and collect the volatile arsenic trioxide in cooler chambers; purify by resublimation. The essential oxidation is represented byActual ores also contain other chemical elements, so their roasting equations and gaseous by-products depend on composition. A contemporary 1860 report describes roasting arseniferous ores, collecting the volatile arsenic trioxide in chambers and refining it by another sublimation. It also demonstrates that apparently good material could contain antimony(III) oxide: physical appearance alone does not establish purity.^[ref] The product intended here is dry $\mathrm{A}{\mathrm{s}}_2{\mathrm{O}}_3$; dissolving it does not make the bottled solid a pure crystalline hydrate ${\mathrm{H}}_3\mathrm{As}{\mathrm{O}}_3$.

For tartar emetic, react antimony(III) oxide with potassium hydrogen tartrate in water. A small amount of water first gives intimate contact in a paste; subsequent heating with more water completes formation of the soluble potassium antimony tartrate. Filter away unreacted insoluble matter and cool the clear filtrate to crystallize the product. This sequence is directly described in a January 1864 discussion of the newly issued British Pharmacopoeia, rather than inferred from its later editions.^[ref] In the traditional antimonyl notation the underlying reaction, before assigning crystal water of crystallization, isThis notation expresses the overall composition; tartrate coordination complexes and water of crystallization need the qualification given in 4. The historical pharmaceutical use of tartar emetic does not make it a safe contemporary emetic.^[ref]

For potassium chlorate, pass chlorine into hot concentrated potassium hydroxide. The disproportionation gives potassium chloride together with potassium chlorate:Concentration and cooling permit fractional crystallization of the much less soluble cold-water potassium chlorate, leaving much of the potassium chloride in the mother liquor; recrystallization improves purity. Cold dilute alkali instead favours hypochlorite, so the heating condition is essential. An economical contemporary manufacturing alternative uses hot calcium hydroxide and potassium chloride; the overall accounting isThe 1850 account in Polytechnisches Journal volume 118 (1850) describes a steam-heated stirred vessel and recovery by cooling, rather than an electrical manufacturing process introduced retrospectively.^[ref]

The main distinction is between metallic iron, iron(II) preparations, iron(III) preparations, and mixed iron oxides. The old prefixes “proto-” and “per-” ordinarily distinguish the lower and higher oxidation states here. “Peroxide of iron” means the historical iron(III) oxide preparation, not a modern peroxide containing an oxygen–oxygen bond.

The principal medicinal preparations can be grouped by the chemical form supplied:

These identifications follow the historical monographs, rather than the names of present commercial supplements.^[ref] A green iron(II) sulfate solution gives the characteristic blue reaction with potassium ferricyanide, whereas the higher iron solutions give blue with potassium ferrocyanide. Those diagnostic reactions explain why the official tests distinguish the two oxidation states; the historical colour test alone does not establish every detail of a complex preparation's composition.

Prepare quinine sulfate by acidic extraction of cinchona bark, alkaline precipitation of the quinine base, conversion into the sulfate, and purification by crystallization. “Quinia” is the historical name for quinine.

The 1867 process starts with one pound of yellow cinchona bark in coarse powder, three fluid ounces of hydrochloric acid, four pints of solution of sodium hydroxide, and sufficient distilled water and diluted sulfuric acid. The hydrochloric acid is diluted with ten pints of distilled water. The bark is first moistened with diluted sulfuric acid, macerated for twenty-four hours with occasional stirring, and then exhausted in a percolator with the diluted hydrochloric acid. The historical endpoint is disappearance of the bitter character of the emerging liquid. This describes the period's test, not an instruction to taste an unidentified medicinal material.^[ref]

The soluble alkaloid salts enter the percolate. Adding solution of sodium hydroxide liberates the much less water-soluble quinine base. After settling, the liquid is decanted and the precipitate collected on a filter and washed with cold distilled water until the washings cease to be coloured. The precipitate is then treated warm with distilled water and diluted sulfuric acid, added gradually to obtain a neutral liquid with nearly all the precipitate dissolved. The hot solution is filtered through paper, the filter washed with distilled water, and the filtrate concentrated until a surface film appears. It is then set aside for crystallization; the crystals are dried on filter paper without applied heat. This is a separation of a naturally occurring alkaloid, not a synthesis of the quinine skeleton.

Writing $B$ for the quinine base, the central changes are:The successive changes in solubility explain the apparatus: a percolator washes soluble salts out of the powder; a filter retains the liberated base; a heated porcelain vessel dissolves the crude base as its sulfate; cooling then recovers the less soluble crystals. Other cinchona alkaloids can accompany the product, which is why appearance and bitterness are insufficient evidence of purity.

The 1867 “Characters and Tests” combine identity and purity checks. They must be read as a group:

The historical quantity and endpoint requirements belong to the 1867 monograph.^[ref] Its formula represents the then-described heptahydrate; the current “Quinine sulfate” instead specifies a dihydrate and modern tests for related cinchona alkaloids. The modern specification explains that distinction; the exam's historical tests remain those listed above.^[ref]