Chemistry (No. III.)

Original question.

In what relation do the chemical elements stand to the other kinds of matter on the earth's surface ? Why are air and water, respectively, no longer regarded as elements ?

Chemical elements are the constituents from which ordinary chemical compounds are built; many familiar materials are mixtures of those chemical compounds and sometimes free chemical elements. In the historical operational definition, a chemical element could not be resolved into simpler substances by chemical means. In the modern definition, its chemical atoms all have the same atomic number. Ordinary chemical reactions rearrange those chemical atoms rather than creating a new chemical element.

Air is a mixture, chiefly of nitrogen and oxygen, whose components can be separated or selectively consumed. For example, burning phosphorus in a suitable enclosed sample removes most of the oxygen and leaves chiefly nitrogen; the residual gas has different properties from the original air. Air also contains carbon dioxide, variable water vapour and smaller amounts of other gases. Its variable composition is incompatible with treating it as one elementary substance.

Water is a chemical compound: electrolysis produces molecular hydrogen and molecular oxygen, approximately in the volume ratio

2 : 1

at the same temperature and pressure. Their recombination produces water again:

2 H_{2} + O_{2} ⟶ 2 H_{2} O .

Thus water has fixed composition but is chemically decomposable, whereas air has no fixed single chemical formula. Neither belongs with the chemical elements merely because it is common and widespread on the Earth's surface.

2

Original question.

By what methods may oxygen be prepared ?

Solution

Oxygen can be obtained by decomposing an oxygen-rich chemical compound, or by separating it from air. For the historical laboratory, heating potassium chlorate, heating certain metal oxides, and electrolysis of water are useful examples. The substance collected is molecular oxygen,

O_{2}

A common preparation heats potassium chlorate, often with manganese dioxide as a catalyst:

2 KClO_{3} Δ 2 KCl + 3 O_{2} .

A heat-resistant vessel and delivery tube conduct the evolved gas to an inverted collecting vessel over water, after the initial displaced air has been excluded. Oxygen is only slightly soluble in water, making this collection principle useful. The manganese dioxide changes the reaction rate and required heating conditions; it is not the net source of the oxygen in this catalysed balance. Without the catalyst, potassium perchlorate can occur as an intermediate.

Two older preparations use heated metal oxides:

2 HgO Δ 2 Hg + O_{2},

4 MnO_{2} strong heating 2 Mn_{2} O_{3} + O_{2} .

The first explains the historical decomposition of mercury(II) oxide, with metallic Mercury deposited on cooler parts of the apparatus. It is historical chemistry, not a preferred present-day teaching preparation. The second is one stage in the reduction of manganese dioxide; at still higher temperature further oxygen loss can give

Mn_{3} O_{4}

, so the oxide residue depends on the conditions. Strong heating of potassium nitrate provides another net balance:

2 KNO_{3} ⟶ 2 KNO_{2} + O_{2} .

(4)

In electrolysis, suitable separated electrodes collect oxygen at the anode and hydrogen at the cathode. With a suitable electrolyte and electrodes, the net chemical equation is

2 H_{2} O ⟶ 2 H_{2} + O_{2} .

(5)

The gases must be kept in separate collecting spaces if the experiment is to demonstrate their separate identities and the approximately

2 : 1

volume ratio. Choice of electrolyte matters: a chloride-containing electrolyte can give chlorine instead of clean oxygen evolution.

For a modern small-scale comparison, catalytic decomposition of hydrogen peroxide gives

2 H_{2} O_{2} ⟶ 2 H_{2} O + O_{2} .

(6)

This is a different feedstock from potassium chlorate. Large-scale separation of oxygen from air by cryogenic fractionation belongs to later industrial practice and should not be substituted for an 1866 laboratory answer. A collected sample can be tested by its vigorous support of combustion, conventionally the relighting of a glowing splint; this observation identifies a property, while the preparation and exclusion of interfering gases establish the sample's provenance.

3

Original question.

What are the sources and properties of ammonia ?

Solution

Ammonia is the volatile nitrogen–hydrogen chemical compound $NH_{3}$ , obtained historically from ammoniacal salts and from the decomposition of nitrogen-containing organic matter. Sources include the decay of animal and vegetable material, the conversion of urea in stale urine, and the destructive distillation of coal, bones or horn. In gasworks, much of the Ammonia was captured in the aqueous ammoniacal liquor during cooling and washing of crude coal gas. These sources do not mean that nitrogen-free constituents of coal spontaneously produce nitrogen.

An ammonium salt releases Ammonia when heated with a suitable chemical base. For ammonium chloride and slaked lime,

2 NH_{4} Cl + Ca (OH)_{2} ⟶ 2 NH_{3} + CaCl_{2} + 2 H_{2} O .

The ionic principle is

NH_{4}^{+} + OH^{-} \to NH_{3} + H_{2} O

. A generating vessel supplies a drying stage and then an inverted dry collecting vessel in which the lighter Ammonia displaces air. It is unsuitable for ordinary collection over water because it dissolves so readily. Quicklime can serve as a drying agent; a strongly acidic drying material would instead consume the Ammonia.

Ammonia is a colourless gas with a penetrating, irritating odour. Its approximate relative density against air is

17.03/28.97 ≃ 0.59

. It is readily liquefied under suitable cooling or pressure, and its dissolution in water releases heat. It is toxic and corrosive at sufficient exposure; direct smelling is not an analytical requirement.^[ref]

The aqueous chemical equilibrium

NH_{3} + H_{2} O ⇌ NH_{4}^{+} + OH^{-}

explains the alkaline reaction and the change of moist red litmus to blue. The historical name “ammonium hydroxide” for the solution should not imply that every dissolved molecule has become a discrete molecule of

NH_{4} OH

. With hydrochloric acid, Ammonia forms ammonium chloride:

NH_{3} + HCl ⟶ NH_{4} Cl .

The familiar white cloud where the gases meet consists of fine solid salt particles, rather than liquid water alone. In solution, Ammonia also acts as a ligand; an excess can produce a deep-blue copper–ammine coordination complex from suitable copper(II) solutions. Although Ammonia does not readily sustain an ordinary flame in air, it is not intrinsically incapable of combustion; sufficiently favourable mixtures burn.

4

Original question.

In what metallic solutions does sulphuretted hydrogen produce precipitates in presence of an acid ? How is this gas prepared ?

Solution

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,

H_{2} 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)	$CuS$	Black
Lead(II)	$PbS$	Black
Mercury(II)	$HgS$	Black when freshly precipitated
Bismuth(III)	$Bi_{2} S_{3}$	Brown-black
Cadmium(II)	$CdS$	Yellow
Arsenic(III)	$As_{2} S_{3}$	Yellow
Antimony(III)	$Sb_{2} S_{3}$	Orange
Tin(II)	$SnS$	Brown or brown-black
Tin(IV)	$SnS_{2}$	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

Ag_{2} 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

M^{2 +} + H_{2} S ⟶ MS ↓ + 2 H^{+} .

The selectivity follows from two coupled chemical equilibrium conditions. Hydrogen sulfide is a weak acid; increasing

[H^{+}]

suppresses its ionization. Approximately,

[S^{2 -}] = \frac{K _{a 1} K _{a 2} [ H _{2} S ]}{[ 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.^[ref]

The traditional preparation treats iron(II) sulfide with a dilute nonoxidizing chemical acid:

FeS + 2 HCl ⟶ FeCl_{2} + H_{2} 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.^[ref]

5

Original question.

What are the principal constituents of coal gas ?

Solution

Purified coal gas is mainly molecular hydrogen, methane and carbon monoxide, with smaller quantities of other hydrocarbons and diluent gases. It is made by heating coal in retorts with little or no admitted air, not by completely burning the coal.

The hot products include permanent gases, condensable organic vapours and water vapour, leaving coke in the retort. Cooling removes much of the coal tar and condenses water, which captures Ammonia as ammoniacal liquor. Further washing and purification remove much of the remaining Ammonia, hydrogen sulfide and other objectionable constituents before the gas is distributed. A crude retort product and purified household coal gas therefore have different compositions.

Molecular hydrogen is commonly the largest component by volume. Methane is another major combustible constituent, while carbon monoxide contributes fuel value and serious toxicity. Ethylene and other unsaturated hydrocarbons, together with vapours such as benzene, contribute strongly to the luminous flame. Their thermal breakdown can produce incandescent soot, explaining why a small fraction of heavier hydrocarbons can disproportionately affect lighting quality. Smaller or variable quantities of carbon dioxide and molecular nitrogen dilute the combustible mixture. Molecular oxygen is not an intended major constituent; air ingress can change the mixture.

There is no single balanced chemical equation for making coal gas: coal is a complex feed and produces many products simultaneously. Nor is one exact percentage composition universal; it depends on the coal, retort temperature, operating conditions and purification. The main distinction is between the light combustible gases, the smaller illuminating hydrocarbon fraction, and contaminants or diluents. Modern natural gas, usually rich in methane, should not be used as a synonym for historical coal gas.

6

Original question.

Describe the mode of preparation and properties of iodide of potassium.

Solution

Potassium iodide is prepared by converting iodine into iodide, then obtaining a pure crystalline potassium salt. One chemical route uses potassium hydroxide; an alternative first forms iron(II) iodide. The account here describes the chemical operations and their balances, without assigning a later pharmacopoeial formula to 1866.

In the alkaline route, iodine reacts with hot potassium hydroxide to give a mixture of potassium iodide and potassium iodate:

3 I_{2} + 6 KOH ⟶ 5 KI + KIO_{3} + 3 H_{2} O .

This is disproportionation: iodine goes from oxidation state

0

to both

- 1

and

+ 5

. Merely evaporating the solution would leave the unwanted iodate with the iodide. Heating the dried salts with charcoal reduces the iodate to iodide. A limiting balance with carbon monoxide as the carbon oxide product is

KIO_{3} + 3 C ⟶ KI + 3 CO .

The carbon oxides can depend on the reaction conditions; the indispensable chemical point is removal of the iodate's oxygen without losing its iodine. Extraction with water, filtration of insoluble residue, concentration and crystallization then recover potassium iodide. A modern USDA technical evaluation explicitly describes this same reaction sequence; it is corroboration of the chemistry, not evidence that its modern document was the 1866 official recipe.^[ref]

In the alternative iron route,

Fe + I_{2} ⟶ FeI_{2},

followed by

FeI_{2} + K_{2} CO_{3} ⟶ 2 KI + FeCO_{3} ↓ .

(4)

The insoluble iron-containing material is separated and the potassium iodide recovered from the liquid. Oxidation of the ferrous precipitate during working can alter its final form, so the displayed chemical equation is the ideal metathesis balance rather than a guarantee that every collected particle remains pure iron(II) carbonate.

Pure potassium iodide is a colourless or white crystalline ionic solid, commonly in cubic crystals, readily soluble in water and less soluble in ethanol. It has no intrinsic brown iodine colour. Its aqueous solution contains potassium ions and iodide ions. Free iodine dissolves in the presence of iodide through triiodide formation,

I_{2} + I^{-} ⇌ I_{3}^{-} .

(5)

With silver nitrate, iodide gives yellow silver iodide, sparingly soluble in water and resistant to ordinary aqueous Ammonia:

Ag^{+} + I^{-} ⟶ AgI ↓ .

(6)

A little chlorine liberates iodine, which gives the characteristic colour with starch:

Cl_{2} + 2 I^{-} ⟶ 2 Cl^{-} + I_{2} .

(7)

By contrast, iodate contamination can liberate iodine simply on acidification in the presence of iodide:

IO_{3}^{-} + 5 I^{-} + 6 H^{+} ⟶ 3 I_{2} + 3 H_{2} O .

(8)

That distinction explains why removal of the iodate is part of preparing a well-defined product, rather than an optional cosmetic step. These are chemical properties and recognition principles, not advice to administer the salt.

7

Original question.

How is phosphorus obtained ?

Solution

The relevant historical route starts with phosphate-rich bone ash, converts it into a reducible phosphate preparation, and heats that preparation with carbon to obtain elemental phosphorus. Bone contains phosphorus in phosphate, not as free pieces of phosphorus awaiting extraction.

Burning cleaned bones removes most organic matter and leaves bone ash, rich in calcium phosphates. Although natural bone mineral is not pure

Ca_{3} (PO_{4})_{2}

, that formula is a useful historical reaction model. Treatment with sulfuric acid forms sparingly soluble calcium sulfate and a soluble acid phosphate fraction. With partial removal of calcium, an idealized balance is

Ca_{3} (PO_{4})_{2} + 2 H_{2} SO_{4} ⟶ Ca (H_{2} PO_{4})_{2} + 2 CaSO_{4} .

Separate the sulfate-rich residue, concentrate the phosphate-bearing liquor, and heat the resulting material with charcoal in a retort. Dehydration and high-temperature reduction produce phosphorus vapour together with carbon oxides, while a calcium phosphate residue remains. One overall balance for the acid-phosphate route is

3 Ca (H_{2} PO_{4})_{2} + 10 C ⟶ Ca_{3} (PO_{4})_{2} + P_{4} + 10 CO + 6 H_{2} O .

This is a summed process balance; it does not imply that the hydrated phosphate survives unchanged until one elementary reaction suddenly occurs. It also makes clear that the partial-acid route leaves some of the original phosphorus in the calcium phosphate residue. Contemporary descriptions distinguish this route from treating the ash with enough sulfuric acid to give a more fully acidified phosphorus-bearing liquid.^[ref]

The retort outlet leads to a water-covered receiver: phosphorus vapour condenses and is protected from air. The collected product is principally white phosphorus, with molecular form

P_{4}

. Its readiness to oxidize and ignite is why collection under water belongs to the explanation of the apparatus. This historical product must be distinguished from red phosphorus, obtained by a subsequent controlled transformation rather than directly identified with the freshly condensed product.

For comparison, the well-known silica-assisted phosphate-reduction balance is

2 Ca_{3} (PO_{4})_{2} + 6 SiO_{2} + 10 C ⟶ P_{4} + 6 CaSiO_{3} + 10 CO .

Here silica binds calcium into a silicate-rich residue. Modern electric-furnace manufacture uses this principle; describing an electric furnace as the ordinary 1866 apparatus would be anachronistic. The historical answer is the sequence bone ash → acid phosphate preparation → carbon reduction in a retort → condensation of phosphorus under water.

8

Original question.

Explain briefly the method of spectrum analysis, and state some of the discoveries which have resulted from its application.

Solution

Spectrum analysis identifies substances by the positions and patterns of characteristic spectral lines, rather than by the overall colour of a flame alone. A small amount of a suitable salt is introduced into a hot, weakly luminous flame. In the historical prism spectroscope, light enters a narrow slit, a collimator makes the emerging rays nearly parallel, a prism separates them by wavelength, and an observing telescope shows images of the slit at different spectral positions. A reference scale or comparison spectrum makes positions reproducible.

A clean platinum wire and reagent blank help distinguish the sample's signal from contamination. A reference sample establishes the expected positions, while repeated cleaning and testing check whether an unfamiliar line belongs to the sample. Several substances can be recognized in one mixture because their characteristic line patterns can coexist, although overlapping lines and changing excitation conditions complicate identification. Robert Bunsen and Gustav Kirchhoff's original account explicitly emphasizes the faint background of the flame, the prism instrument and the characteristic patterns of alkali and alkaline-earth salts.^[ref]

In modern terms, emission lines arise when excited chemical atoms or ions release photons at energies corresponding to allowed differences between energy levels:

h ν = E_{upper} - E_{lower}, λ = c / ν .

The quantum mechanics behind this equation is a later explanation, not the theory available to an 1866 candidate. Similarly, absorption lines appear when a cooler foreground gas removes particular wavelengths from a brighter background. Agreement between laboratory line positions and dark lines in the solar spectrum permits identification of constituents of the solar atmosphere without bringing a solar sample into the laboratory. It does not establish the detailed chemical constitution of every layer of the Sun.

By December 1866, striking examples included caesium (1860), rubidium (1861), thallium (1861) and indium (1863). The first two were recognized in the work of Robert Bunsen and Gustav Kirchhoff, using conspicuous blue and deep-red spectral features respectively. William Crookes recognized the characteristic green line associated with thallium, while Ferdinand Reich and Hieronymus Theodor Richter identified indium through its indigo-blue spectral signature. These observations motivated chemical separation and further characterization; an unexpected line by itself is a clue to a new chemical element, not a complete proof.^[ref]^[ref]^[ref]^[ref]

The corresponding recognition of familiar chemical elements in the solar spectrum opened stellar chemical analysis. Later discoveries such as solar helium should not be credited to the 1866 state of the subject.

9

Original question.

How would you ascertain qualitatively the chemical composition of a tubercle ?

Solution

Use a qualitative separation into aqueous extract, lipid-soluble material, protein-rich residue and mineral matter, and test each fraction independently. A tubercle is a lesion containing several kinds of material, not a pure chemical compound with one diagnostic formula. The question asks how to investigate its composition; no observations or numerical composition are supplied.

First record the lesion's physical appearance and take a portion for microscopy before destructive treatment. Separate the lesion from surrounding tissue as far as practicable, while recording that contamination by blood, connective tissue or adherent salts can affect the analysis. Different stages of a lesion need not have the same proportions of organic material and calcified matter. A small portion gently dried in suitable apparatus shows whether much water is lost; collecting condensation supports identification of the volatile fraction as containing water, although mass loss alone is not a chemically specific water test.

On separate portions, or by a documented extraction sequence, examine the main fractions:

The water extract contains soluble salts and some soluble organic material. Test its reaction with litmus; heating a suitable clear portion can disclose coagulable protein. Acidification and heating must be controlled because changes of pH and high salt content can alter coagulation. The absence of a heat coagulum does not prove the absence of every protein.
An ethanol or diethyl ether extraction can remove lipids and other solvent-soluble material. Evaporating a separated solvent extract leaves material whose solubility, oily or crystalline appearance, and behaviour on heating help characterize it. Hydrolysable fats can undergo saponification; cholesterol belongs to the unsaponifiable fraction. A greasy residue alone does not establish one particular lipid.
Examine the largely insoluble organic residue for protein. A controlled alkaline copper reaction gives the violet biuret test with sufficient peptide bonds; the yellow xanthoproteic reaction with nitric acid supplies corroborating evidence for suitable aromatic residues. Neither colour establishes that all the material is specifically albumin, casein or a hypothetical unique “tuberculin” substance. Record the reaction conditions and use a reagent blank and known positive material.
Ignite a separate portion carefully to remove combustible organic material, then examine the ash. Dissolve suitable portions in a noninterfering chemical acid. Calcium can be recognized by oxalate precipitation under appropriate conditions; phosphate by an established phosphate test such as the yellow ammonium phosphomolybdate reaction. Carbonate can be recognized by acid-induced carbon dioxide evolution and a separate limewater test. These are separate aliquots: acid treatment used to dissolve the ash can destroy evidence for the original carbonate.

Test soluble salts using a fresh portion or a suitable water extract: for example, silver nitrate can detect chloride, while barium under controlled acidic conditions can detect sulfate. A chloride test should not be performed on material already contaminated with hydrochloric acid. The results identify constituents or operational fractions, not necessarily the precise salts originally present in intact tissue; ashing can also change their chemical forms.

If the intended composition is elemental rather than proximate, investigate an isolated organic fraction separately: oxidation with heated copper(II) oxide can convert its carbon into carbon dioxide and its hydrogen into water, which are trapped and identified independently. A suitable nitrogen determination can convert organic nitrogen into Ammonia for identification in a separate receiver. Mineral carbonate and pre-existing moisture must be controlled before attributing these products to the organic material. Such tests answer a different question from identifying the particular proteins or lipids.

This approach has a historical basis. Original analyses by Robert Mortimer Glover, reproduced in an 1847 review, separately extracted material with ethanol, diethyl ether and water, then examined protein-rich residues and mineral salts. They reported variation between lesions and did not consistently find casein. This evidence supports fractionation and caution about specificity, not one universal numerical recipe for all tubercles.^[ref]

A modern conclusion must separate composition from cause. A protein-, lipid- and phosphate-containing lesion is not thereby proved to be tuberculosis. Histological interpretation and appropriate microbiological examination are different investigations. Current confirmation of Mycobacterium tuberculosis involves validated tests such as nucleic acid amplification and culture, rather than the extraction reactions above; these are modern qualifications and are not attributed to the 1866 examination.^[ref]

10

Original question.

Write a short account of the principal chemical changes to which the name ‘fermentation’ has been given.

Solution

The historical name fermentation covered several transformations of organic matter: alcoholic, lactic, butyric and acetous changes, as well as some processes now classified separately. In modern biochemistry, fermentation more narrowly means energy-yielding metabolism that regenerates oxidized cofactors without requiring an external respiratory electron acceptor. The older use of the word should not be forced into that modern definition.

In alcoholic fermentation, suitable yeast converts a fermentable sugar into ethanol and carbon dioxide. For glucose, the principal net balance is

C_{6} H_{12} O_{6} ⟶ 2 C_{2} H_{5} OH + 2 CO_{2} .

A vessel with an outlet for the evolved gas and a separate receiver illustrates the change: test the carbon dioxide with limewater and identify the volatile alcohol in a separated distillate. Gas evolution alone does not prove that the liquid contains ethanol. Actual yeast cultures also make biomass and minor products; the chemical equation describes the principal conversion. In modern terms, glycolysis supplies ATP, and the reduction leading to ethanol regenerates NAD+ from NADH.

In lactic acid fermentation, an ideal homolactic balance is

C_{6} H_{12} O_{6} ⟶ 2 CH_{3} CH (OH) COOH .

The souring of suitable carbohydrate-containing materials, including milk, supplied a familiar historical example. Hydrolysis of lactose supplies monosaccharides, while acid production lowers the pH and can precipitate casein. Casein coagulation is a consequence of the changed conditions, not evidence that the protein has itself been converted into lactic acid. Some organisms instead make additional products, so not every lactic fermentation obeys only this homolactic balance.

In butyric acid fermentation, carbohydrates or lactate can give butyric acid, carbon dioxide and molecular hydrogen. An idealized balance from lactic acid is

2 C_{3} H_{6} O_{3} ⟶ C_{4} H_{8} O_{2} + 2 CO_{2} + 2 H_{2} .