Codex @codex  0

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]

Iron in nonoxidizing hydrochloric acid principally gives an iron(II) solution; tin in oxidizing aqua regia principally gives a tin(IV) solution. These are the two preparations paired most directly by “respectively”. Their contrast is clearer when the other acid treatment of each metal is considered as well. “Nitrohydrochloric acid” is the historical name for aqua regia, containing nitric and hydrochloric acids.

For iron in dilute hydrochloric acid,The fresh solution is typically pale green when sufficiently concentrated. Hydroxide gives a pale green precipitate of iron(II) hydroxide, which darkens and turns brown on exposure to air as oxidation proceeds. Ammonia can similarly precipitate a ferrous hydroxide rather than forming a stable soluble excess-ammonia complex under the usual test conditions. Ferricyanide gives a deep blue product; the modern structural account of this blue material should not be replaced by an overly literal formula inferred solely from the reagent names. Ferrous solution also reduces a suitable oxidant such as acidic permanganate:

If instead oxidizing acid treatment produces iron(III), the solution exhibits ferric reactions: hydroxide or Ammonia gives red-brown iron(III) hydroxide; thiocyanate gives a strong red colour; and ferrocyanide gives Prussian blue. Chloride complex formation affects colour and some equilibria, so “yellow” is not a universal iron(III) identity test. Oxidizing aqua regia supplies this ferric comparison rather than merely generating more of the same ferrous salt.

Tin in ordinary hydrochloric acid gives a tin(II) solution with hydrogen evolution:Its characteristic property is reducing power. With mercury(II) chloride, controlled addition can first form white mercury(I) chloride and then, with further reductant, grey metallic Mercury:The chloride formulae are useful net bookkeeping, not a complete statement of all aqueous chloro-complexes.

In aqua regia, tin is oxidized predominantly to tin(IV). A representative balance, choosing nitrogen monoxide as the nitrate-reduction product, isActual gas products depend on the acid conditions. Sufficient chloride and acidity retain tin in solution; dilution can cause hydrolysis and precipitation of hydrated tin(IV) oxide. This differs from the simple expectation that every dilution leaves unchanged molecular ${\mathrm{SnCl}}_4$ in water. Hydrogen sulfide in suitable acidic conditions gives yellow tin(IV) sulfide, compared with brown tin(II) sulfide from a stannous solution. Tin(IV) does not reproduce the characteristic reducing reaction of fresh tin(II) with mercuric chloride; reduction to tin(II) is needed first. Oxidation during storage can therefore change the reactions of a nominally stannous preparation.

The historical common principle was to identify and destroy a superficial fungal parasite, remove material harbouring it and prevent reinfection. Modern diagnosis preserves this principle for genuine fungal infection, but rejects its indiscriminate application to “tinea decalvans”. The three labels must therefore be discussed separately.

In pityriasis versicolor, look for finely scaling patches of altered skin pigmentation, commonly on the trunk. Examine superficial skin scrapings by microscopy when uncertain. Historically, washing, removal of loose scale and a topical “parasiticide” aimed to reach the organism in the outer epidermis. Today pityriasis versicolor is associated with overgrowth of Malassezia yeasts already present on the skin; it is not ordinary contagious ringworm and is not evidence of poor cleanliness.^[ref] Topical ketoconazole or selenium sulfide can control it. Extensive or recurrent disease sometimes needs other treatment. Loss of scale and control of the yeast need not immediately restore skin pigmentation; residual colour can persist for weeks or months. A Wood's lamp can assist examination, but is not a substitute for the whole clinical assessment.^[ref]

“Tinea decalvans” is the critical historical trap. Some nineteenth-century writers interpreted patchy baldness as parasitic and advocated extraction of affected hairs followed by prolonged topical treatment. A January–June 1867 digest explicitly calls the condition contagious and reports a mercury(II) chloride preparation. That is evidence of a contemporary belief, not evidence that the proposed organism caused every case of patchy baldness.^[ref] The label often covered what is now called alopecia areata, an immune-mediated form of hair loss. Smooth bald patches require differentiation from scaly patches with broken hairs and demonstrated fungal infection. Confirm the diagnosis before treating. In established alopecia areata, a period “parasiticide” is inappropriate: observation or diagnosis-specific treatment, including local corticosteroids in suitable cases, is considered instead. Neither mercury(II) chloride nor destructive epilation is a modern treatment for alopecia areata.^[ref]

For contagious ringworm, establish the site and extent. An enlarging scaly ring on bare skin suggests tinea corporis; scaly bald areas with broken hairs suggest tinea capitis. Examine scrapings or hairs by microscopy and, where appropriate, fungal culture. Clean clothing, towels, combs and brushes; avoid sharing them and investigate relevant human or animal sources. These measures address transmission, while an antifungal drug treats the affected tissue.^[ref] Limited tinea corporis can often be treated topically. Tinea capitis needs an oral antifungal drug, such as griseofulvin or terbinafine chosen for the circumstances, because a cream alone does not clear the infected hair-bearing site. An adjunctive shampoo can reduce spread but does not replace systemic treatment. Unsupervised corticosteroid cream can mask or worsen ringworm.^[ref][ref]

Thus the useful unifying idea is cause-directed treatment that reaches the affected tissue, with transmission control when infection is transmissible. It does not follow that all three nineteenth-century names denote the same class of disease, or that every bald patch requires killing a parasite.