Chemistry (No. III.)

Original question.

What other substances besides Oxygen and Nitrogen may be found in Atmospheric Air ? By what tests is their presence or absence indicated ?

Ordinary air contains water vapour, carbon dioxide, variable trace gases and suspended particles in addition to nitrogen and oxygen. Which traces are detected depends on location, sampling volume and test sensitivity. A negative small-sample test is not proof of complete absence.

Water vapour is shown by condensation on a sufficiently cold surface or by the colour change of a suitable dry indicator: white anhydrous copper(II) sulfate becomes blue on hydration. Passing a known air stream through a drying tube and observing a controlled mass increase gives a more quantitative principle, provided other absorbable vapours are excluded.

Carbon dioxide gives a calcium carbonate turbidity when a suitable air sample is passed into clear limewater:

CO_{2} + Ca (OH)_{2} ⟶ CaCO_{3} ↓ + H_{2} O .

Large excess of carbon dioxide can redissolve the deposit as dissolved bicarbonate species, so excessive exposure can obscure the initial result. Absorption into an alkaline solution, followed by an appropriate carbonate test, concentrates the trace component.

Ammonia and ammonium-bearing particles may be collected by an acid wash and examined after suitable neutralization or liberation of the Ammonia. Nessler's reagent supplies a sensitive colour test, with a blank to exclude contaminated reagents. Simple moist litmus is much less sensitive. Acidic gases and nitrogen oxides may instead be collected in appropriate wash solutions and examined with tests for the resulting anions; their identity cannot be inferred from acidity alone.

Near relevant sources, hydrogen sulfide can blacken moist lead acetate paper by forming lead sulfide. Ozone-sensitive starch–iodide paper can indicate an oxidizing trace constituent through liberation of iodine, but other oxidants can give the same response: the historical observation is not a specific proof of ozone. Dust and droplets can be collected on a filter or surface and examined by microscopy, with chemical tests applied to separated material as needed.

Modern analysis also recognizes argon and other noble gases, as well as many low-concentration pollutants, using instrumental separation and spectroscopy. Their modern recognition should not be silently included among discoveries available to an 1865 candidate. The useful answer pairs each proposed constituent with a collection method, a specific reaction or observation, and the limits of that test.

2

Original question.

Why is rain water always ‘soft’ ? Describe and explain the method of estimating quantitatively the ‘hardness’ of water.

Solution

Fresh rainwater is ordinarily soft because evaporation leaves most nonvolatile dissolved salts behind; “always” is an idealization, not an exceptionless law. Water hardness is principally caused by dissolved calcium and magnesium ions. Water vapour rising from a hard-water source does not carry these ions with it as vapour. Condensation consequently supplies low-mineral water before contact with dust, surfaces or soil. Mineral particles, sea spray and a collecting surface can add salts, so actual collected rainwater need not have zero water hardness. Nor does softness establish purity or freedom from microorganisms.^[ref]

The historically appropriate quantitative principle is soap titration of water hardness. Soluble soap reacts with hardness-producing ions to form poorly soluble salts, using up soap before a persistent lather can form. For a representative carboxylate soap,

Ca^{2 +} + 2 RCOO^{-} ⟶ Ca (RCOO)_{2} ↓ .

Magnesium behaves similarly. Once the dissolved hardness-producing ions have substantially reacted, additional soap can sustain the specified lather. The endpoint is therefore an operational observation under standardized shaking, standing time, temperature and sample volume; it is not the first transient bubble.

Use a measured volume of sample in a stoppered vessel with enough headspace for shaking. Add a soap solution of known strength incrementally from a burette, shaking in a reproducible way, until the chosen persistent-lather endpoint appears. Determine a blank with a corresponding amount of soft or distilled water, because some soap is needed for the lather even when no hardness salts are present. Standardize the soap against a reference solution containing a known calcium amount, expressed as an equivalent mass of calcium carbonate. This standardization is essential: the apparent strength cannot safely be inferred from the mass of an incompletely characterized commercial soap.

Let

V_{s}

be the sample volume in millilitres,

V

the soap titre in millilitres and

B

the corresponding blank. If

f

is the standardized capacity in milligrams of calcium carbonate equivalent per millilitre of soap solution, then

H = \frac{1000 f ( V - B )}{V _{s}} mg L^{- 1} as CaCO_{3} .

For a standard containing

m_{0}

milligrams of calcium carbonate equivalent and requiring blank-corrected titre

V_{0} - B_{0}

f = m_{0} / (V_{0} - B_{0})

. This makes the measurement quantitative without assuming an unreported soap concentration or inventing a sample titre. Historical results could instead be stated as grains of calcium carbonate equivalent per gallon, provided the gallon convention was specified.

Repeating the examination after controlled boiling, cooling, removal of the precipitate and restoration of the original volume estimates the residual or permanent hardness. The reduction estimates the boiling-removable fraction, often called temporary hardness. A representative reaction is

Ca^{2 +} + 2 HCO_{3}^{-} ⟶ CaCO_{3} ↓ + CO_{2} + H_{2} O .

Carbonate equilibria and magnesium chemistry make this more subtle than assuming every dissolved salt is destroyed by boiling. Very hard samples may need measured dilution; magnesium, other metal ions and contamination can disturb the lather endpoint. Thus the soap result is an experimentally calibrated hardness equivalent, not an exact identification of every original salt.

3

Original question.

How would you examine a sample of water for (1) Lead, (2) Nitrates, (3) Organic matter ? Under what circumstances is water liable to contain each of these ?

Solution

Examine separate, properly collected aliquots for lead, nitrate and organic matter, and interpret the results with the source and sampling conditions. These are three different analytical questions. No sample observations are supplied, so no actual contaminant concentration or positive finding can be assigned.

For lead, preserve a representative portion in clean apparatus, avoiding lead-containing fittings or contaminated reagents. Distinguish dissolved material from suspended or particulate material if that distinction is wanted; filtering at the outset otherwise changes what is being measured. Historically, a suitable concentration step followed by hydrogen sulfide gives dark lead sulfide:

Pb^{2 +} + H_{2} S ⟶ PbS ↓ + 2 H^{+} .

A separate confirmation uses potassium iodide to give yellow lead(II) iodide. Its characteristic recrystallization behaviour on heating and cooling a suitable aqueous preparation is useful supporting evidence. Dark sulfide alone is insufficient, because other metals give dark precipitates. Establish a reagent blank, compare with known material, and record the sensitivity; a negative visual test cannot exclude all low-level lead.

The principal circumstance is contact with lead-bearing pipes, cisterns, solder or other plumbing materials. Stagnation, acidity and water chemistry influence release; newly exposed metal and disturbed corrosion deposits can also matter. Low hardness does not by itself predict a universally safe or unsafe supply: pH, alkalinity, dissolved gases, protective films and the actual materials must be considered together. Current CDC guidance confirms the importance of corrosion of plumbing and of contact time and water chemistry.^[ref]

For nitrate, the historical brown-ring test uses fresh iron(II) sulfate and a carefully formed strongly acidic interface. Reduction of nitrate produces nitrogen monoxide, which forms the brown nitrosyl complex with excess iron reagent:

NO_{3}^{-} + 3 Fe^{2 +} + 4 H^{+} ⟶ NO + 3 Fe^{3 +} + 2 H_{2} O .

The observable ring is associated with the resulting iron–nitrosyl species, not with free nitrate having a brown colour. Use nitrate-free reagents and, if necessary, a suitably concentrated sample. Nitrite can interfere and must be distinguished independently; adding nitric acid during an earlier preparation would invalidate this test for nitrate originally in the water. Strong acid and heating can disturb the ring, so the observation depends on the procedure as well as the constituents.

Nitrate may occur after oxidation of nitrogenous wastes in soil, from manure, sewage, cesspit leakage, fertilizer leaching or natural nitrogen cycling. It can persist after much readily oxidizable organic material has disappeared. Its presence alone therefore neither proves recent sewage entry nor identifies a particular source. Current EPA source-water guidance includes agricultural runoff, septic leakage, sewage and atmospheric inputs among nitrate sources.^[ref]

For organic matter, inspect colour, turbidity and sediment, and examine a suitable concentrate by microscopy. Evaporating a measured portion and examining the residue for charring or change on ignition supplies further evidence, but loss on heating is not specific: salts and retained water can also change. A controlled potassium permanganate consumption experiment compares a sample with a reagent blank under stated acidity, reaction time and temperature. It detects oxidizable material, which can include inorganic reducing substances, and may incompletely oxidize organic material. The amount consumed is consequently an operational observation, not a universal conversion into a mass of “sewage”.

Sources include decaying vegetation, peat and soil-derived material as well as animal wastes, sewage and some industrial discharges. Natural organic colour is not synonymous with sewage, and clear water may contain dissolved organic contaminants or microorganisms. A modern safety assessment therefore adds validated microbiological and chemical measurements appropriate to the source; these historical qualitative tests alone cannot establish potability.^[ref]

4

Original question.

Give an account of some method by which Arsenic may be detected in the presence of organic substances.

Solution

The historical Marsh test separates arsenic into a volatile hydride and then into a visible deposit, allowing it to be examined apart from much of the original organic matrix. It requires controls and confirmation because a dark deposit is not unique to arsenic. The test is associated with James Marsh; its primary account explicitly addresses mixtures containing animal and vegetable matter.^[ref]

Retain representative material and a control portion before destructive preparation. A complex organic sample can be clarified and, where needed, subjected to a validated wet oxidation that destroys interfering organic matter while retaining the arsenic in the analytical preparation. Avoid uncontrolled heating to dryness or conditions that volatilize an arsenic compound; successful mineralization cannot simply be assumed. The reagent sequence must leave the arsenic in a reducible form and must not leave enough oxidant to defeat the subsequent reduction. A recovery experiment with a known addition is a check on the entire preparation, not merely on the last colour reaction.

In the hydrogen-generating stage, arsenic-free zinc reacts with a suitable dilute acid. The sample's reducible arsenic is converted to arsine, which passes with the carrier gas out of the generating vessel. A representative net reduction balance for arsenic(III) is

H_{3} AsO_{3} + 3 Zn + 6 H^{+} ⟶ AsH_{3} + 3 Zn^{2 +} + 3 H_{2} O .

This is a reaction balance, not a claim that the old phrase “nascent hydrogen” identifies one uniquely established microscopic mechanism.

The apparatus needs controlled reagent entry, an unobstructed gas path and a means of excluding droplets carried from the sample. In the later heated-tube form, the hydride decomposes at a hot section:

2 AsH_{3} ⟶ 2 As + 3 H_{2} .

Arsenic collects as a mirror on the cooler glass beyond it. Marsh's original flame method deposited material on a cold surface held in relation to the burning jet. The original arrangement and the heated-tube improvement should not be treated as identical apparatus. In either case the useful analytical separation is transfer of the arsenic away from the nonvolatile matrix, not just darkening of that matrix.

Run the complete reagent blank first: contaminated zinc or acid can otherwise make an arsenic-free sample appear positive. A positive known control checks that hydride production and collection work. Antimony can form stibine and a similar dark deposit, so mirror formation alone is not conclusive. Confirm with independent chemistry, for example controlled oxidation of the collected deposit followed by appropriate arsenic tests; the greater susceptibility of an arsenic mirror to suitable hypochlorite treatment supplies an additional historical distinction from an antimony deposit. No single visual resemblance should replace these discriminations.

A negative test is meaningful only relative to the validated preparation, detection limit and controls. Modern arsenic analysis uses instrumental methods with defined recovery and sensitivity; the historical demonstration is not itself a contemporary forensic certification. Arsine is acutely toxic, and the hydrogen-bearing gas stream introduces further apparatus hazards: this account explains the analytical chemistry, not a household procedure.^[ref]

5

Original question.

What reactions do the solutions exhibit which are obtained by dissolving (1) Iron and (2) Tin in Hydrochloric and Nitrohydrochloric acids respectively ?

Solution

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,

Fe + 2 HCl ⟶ FeCl_{2} + H_{2} .

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:

MnO_{4}^{-} + 5 Fe^{2 +} + 8 H^{+} ⟶ Mn^{2 +} + 5 Fe^{3 +} + 4 H_{2} O .

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:

Sn + 2 HCl ⟶ SnCl_{2} + H_{2} .

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:

2 HgCl_{2} + SnCl_{2} ⟶ Hg_{2} Cl_{2} ↓ + SnCl_{4},

Hg_{2} Cl_{2} + SnCl_{2} ⟶ 2 Hg ↓ + SnCl_{4} .

(5)

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, is

3 Sn + 4 HNO_{3} + 12 HCl ⟶ 3 SnCl_{4} + 4 NO + 8 H_{2} O .

(6)

Actual 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

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.

6

Original question.

Give some account of the metal Antimony and its combinations.

Solution

Antimony is the brittle, silvery crystalline element $Sb$ , encountered especially in sulfide ores and in compounds of oxidation states $+ 3$ and $+ 5$ . Its historical treatment as a metal is understandable from appearance and alloying behaviour; modern classifications commonly call it a metalloid. It is too brittle to behave like a readily drawn, ductile metal and is used in suitable alloys to alter properties such as hardness.

The important ore is stibnite,

Sb_{2} S_{3}

. A historical reduction can use iron to bind the sulfur:

Sb_{2} S_{3} + 3 Fe ⟶ 2 Sb + 3 FeS .

Another route roasts suitable sulfide material to an oxide and then reduces that oxide with carbon. For an ideal trioxide product,

2 Sb_{2} S_{3} + 9 O_{2} ⟶ 2 Sb_{2} O_{3} + 6 SO_{2},

Sb_{2} O_{3} + 3 C ⟶ 2 Sb + 3 CO .

Real ore processing requires control of impurities and volatile losses, so these balances do not alone constitute a complete extraction flowsheet.^[ref]

Among its principal combinations, antimony(III) oxide is an amphoteric oxide. It dissolves under suitable acidic or alkaline conditions rather than fitting a universal rule that a metal oxide must be solely basic. Antimony(III) chloride, historically “butter of antimony”, hydrolyses on dilution and can deposit a white oxychloride-containing material. A useful simplified balance is

SbCl_{3} + H_{2} O ⇌ SbOCl ↓ + 2 HCl .

Added acid favours dissolution; a more extensively hydrolysed precipitate need not be pure

SbOCl

. Pentavalent compounds include antimony(V) chloride and antimonate species; chloride-rich solutions and their hydrolysis require attention to the medium.

Antimony(III) sulfide can occur as the dark crystalline ore and as an orange precipitate obtained from a suitable solution with hydrogen sulfide:

2 Sb^{3 +} + 3 H_{2} S ⟶ Sb_{2} S_{3} ↓ + 6 H^{+} .

(5)

Its dissolution in suitable alkaline sulfide solutions forms thioantimony species, providing an analytical separation from some other sulfides. Acidification can reverse that operational separation. A volatile hydride, stibine,

SbH_{3}

, can decompose to a dark antimony deposit and thus interfere with the Marsh test for arsenic.

Antimony potassium tartrate, “tartar emetic”, is a historically important complex salt, not merely an arbitrary mixture of an antimony salt and a potassium salt. It was used medicinally, but historical use is not proof of safety or a present indication; antimony-containing substances can cause serious poisoning. No dosage or treatment recommendation follows from describing them.^[ref]

7

Original question.

Describe the modes of preparation and the properties of Carbonic Oxide.

Solution

“Carbonic oxide” means carbon monoxide, $CO$ : a combustible, reducing, colourless gas distinct from carbon dioxide. Several preparations remove part of the oxygen from a carbon–oxygen compound or prevent complete oxidation of carbon.

Passing carbon dioxide over sufficiently hot carbon gives

CO_{2} + C ⇌ 2 CO .

High temperature favours substantial conversion. The reaction is important in furnaces, but gas emerging from such a process need not be pure carbon monoxide; residual carbon dioxide and other gases must be considered. Incomplete combustion of carbonaceous fuel is another source, with the ideal net balance

2 C + O_{2} \to 2 CO

. It is a source principle rather than a selective method guaranteeing a pure sample.

For a historical laboratory preparation, formic acid can be dehydrated by a suitable strongly dehydrating medium:

HCOOH ⟶ CO + H_{2} O .

Oxalic acid under an appropriate dehydrating treatment yields a mixture:

H_{2} C_{2} O_{4} ⟶ CO + CO_{2} + H_{2} O .

An alkaline absorption stage removes much of the carbon dioxide, while carbon monoxide passes on. A generator, droplet trap, purifier and collecting vessel illustrate why preparation and purification are separate operations. The gas is only slightly soluble in water, so collection over water is possible in principle. Concentrated sulfuric acid acting on carbonaceous material can also give mixed carbon oxides and sulfur dioxide; that route requires additional purification and should not be presented as clean CO evolution without by-products.

The relative density of carbon monoxide against dry air is approximately

28.01/28.97 ≃ 0.967

. It has no reliable warning odour. It burns with a blue flame when supplied with suitable oxygen:

2 CO + O_{2} ⟶ 2 CO_{2} .

A CO sample itself does not ordinarily support combustion of a taper. Nor does it immediately give the usual limewater test of carbon dioxide; oxidation of the CO first gives a gas that does. This distinguishes fuel behaviour from the property of supporting combustion of another material.

Its reducing power is illustrated by

CuO + CO ⟶ Cu + CO_{2}

(5)

and by reduction of suitable iron oxides in metallurgy. The products reflect transfer of oxygen from the oxide to carbon monoxide. Its ability to bind metal centres is also chemically important. In biological exposure, binding to haemoglobin and other effects disrupt oxygen delivery and use; it is a serious poison, not merely an innocuous replacement of atmospheric oxygen. This modern explanation must not turn an historical gas-preparation answer into a proposed exposure experiment.^[ref]^[ref]

8

Original question.

How are Ether, Bisulphide of Carbon, and Chloroform prepared ?

Solution

The three preparations use different chemistry: diethyl ether from ethanol by acid-catalysed dehydration, carbon disulfide by combination of carbon and sulfur, and chloroform by chlorination and cleavage of a suitable organic starting material. “Ether” here means diethyl ether, not every member of the modern ether class; “bisulphide of carbon” means carbon disulfide.

For diethyl ether, heating ethanol with a suitable proportion of sulfuric acid favours intermolecular dehydration:

2 C_{2} H_{5} OH ⟶ C_{2} H_{5} O C_{2} H_{5} + H_{2} O .

The historical explanation invokes formation of ethyl hydrogen sulfate and its reaction with a further ethanol molecule, regenerating sulfuric acid. The acid therefore need not be consumed stoichiometrically in the overall balance. The modern mechanism treats a protonated alcohol or related activated species as the substrate for attack by another ethanol molecule.^[ref] Conditions matter: excessive heating favours formation of Ethylene rather than ether, and strongly oxidizing conditions can give additional products. Condensation of the volatile product, removal of acidic and water-soluble impurities, drying and redistillation separate production from purification. A cold receiver is appropriate to the volatile ether; an open flame near the escaping vapour would be incompatible with its ready flammability.

For carbon disulfide, sulfur vapour is brought into contact with sufficiently hot charcoal in a retort:

C + 2 S ⟶ CS_{2} .

Here

S

records sulfur atoms in the balance, not a claim that the sulfur vapour has only monatomic molecules. The outgoing vapours are cooled and the liquid separated from excess sulfur and other impurities, then purified by distillation. The high-temperature reaction zone and the cooled collecting region have different functions. Carbon disulfide is a volatile, highly flammable solvent; its historical use does not establish harmlessness.^[ref]

For chloroform, a historically important route uses ethanol with a slurry containing bleaching powder. Oxidation first supplies acetaldehyde; further chlorination gives chloral, which alkali cleaves to chloroform and formate. An illustrative final step with calcium hydroxide is

2 CCl_{3} CHO + Ca (OH)_{2} ⟶ 2 CHCl_{3} + Ca (HCOO)_{2} .

The bleaching material supplies oxidizing and chlorinating capacity, while the alkaline conditions permit cleavage. It is misleading to describe the entire operation as one simple substitution directly replacing three hydrogens of intact ethanol. It is also unnecessary to pretend that commercial bleaching powder has a single invariant molecular composition.

Acetone supplies an alternative methyl-carbonyl starting material for the same broad haloform chemistry. In ionic notation a clean idealized balance is

CH_{3} COCH_{3} + 3 OCl^{-} ⟶ CHCl_{3} + CH_{3} COO^{-} + 2 OH^{-} .

The volatile product is condensed with water-containing material; a chloroform-rich phase is separated, washed appropriately, dried and redistilled. Identity and purity require more than observing a dense lower layer. These operations describe preparation chemistry rather than a specified pharmacopoeial recipe or a claim of suitability for anaesthesia. The historical anaesthetic importance of chloroform does not remove its toxic effects.^[ref]

9

Original question.

Describe the process by which the percentage composition of a substance consisting of Carbon, Hydrogen, and Oxygen is determined.

Solution

Burn a known mass of pure, suitably dry material completely, collect the resulting carbon dioxide and water separately, calculate the carbon and hydrogen masses, and obtain oxygen by difference under the stated three-element assumption. This is organic elemental analysis, not an identification of molecular structure.

Weigh a representative portion of mass

m

in a clean combustion boat or suitable carrier. In the classical apparatus, a heated combustion tube contains an oxidizing material such as copper(II) oxide, with a controlled stream of clean, dry oxygen-bearing gas where the arrangement requires it. The purpose is complete conversion of carbon into carbon dioxide and hydrogen into water, without unburnt material escaping. An equation for a model substrate is

C_{x} H_{y} O_{z} + (x + \frac{y}{4} - \frac{z}{2}) O_{2} ⟶ x CO_{2} + \frac{y}{2} H_{2} O,

with the oxygen amount understood as the net external requirement when positive.

Downstream, a preweighed water-absorption tube collects the water. A following preweighed alkaline absorber collects the carbon dioxide; the order prevents water from being incorrectly attributed to the carbon dioxide absorber. Guards exclude atmospheric moisture and carbon dioxide during handling, and a final clean-gas sweep transfers products still held in the apparatus. Let the blank-corrected mass increases be

a

for the water collector and

b

for the carbon dioxide collector.

Using the simple historical atomic-mass ratios

C = 12

H = 1

O = 16

m_{C} = \frac{12}{44} b = \frac{3}{11} b, m_{H} = \frac{2}{18} a = \frac{1}{9} a .

Consequently,

% C = 100 \frac{3 b}{11 m}, % H = 100 \frac{a}{9 m},

% O = 100 - % C - % H .

Modern work uses the appropriate modern molar masses in the same ratios. No sample mass or absorption results are supplied, so no numerical percentages are justified.

The oxygen already present in the sample cannot be obtained merely by measuring the oxygen in the combustion products: the oxidant supplies additional oxygen. Difference works here only because the sample is assumed to contain exclusively carbon, hydrogen and oxygen and to be free of unaccounted water, solvent, ash and other elements. If the combustion leaves a mineral residue, that mass cannot silently be reported as oxygen. A separate check of purity and residue therefore tests an assumption of the calculation.

Incomplete combustion, escaped vapour, exhausted absorbers, moisture entering from the carrier gas, and contamination of the carbon dioxide absorber all bias the result. A blank and an independent repeat with a suitable reference material check recovery. Oxygen by difference accumulates the errors in both other determinations; forcing the percentages to sum to 100 is not independent proof that the analysis is correct.

10

Original question.

What other substances besides those named in the preceding question may an organic substance contain, and how is their presence discovered ?

Solution

An organic substance may also contain nitrogen, sulfur, phosphorus, halogens and, in some cases, metals or other elements. “Organic” does not mean “made of only carbon, hydrogen and oxygen”. Detection generally requires converting a covalently bound element into a species that an appropriate inorganic test can recognize.

A historically available general method is the sodium fusion test. Fusion converts suitable nitrogen-containing material into cyanide, sulfur into sulfide, and chlorine-, bromine- or iodine-containing material into the corresponding soluble halides. The cooled reaction material is worked up to an aqueous extract, with separate portions reserved for separate tests. The transformations are analytical conversion principles, not claims that every nitrogen or sulfur compound must give quantitative recovery under any fusion conditions. Blanks and a known positive material are essential when a negative result is to be interpreted.^[ref]

For nitrogen, a suitable cyanide-containing extract is treated with a fresh iron(II) reagent and then brought through the appropriate oxidation and acidification sequence to obtain Prussian blue. Formation of ferrocyanide is represented by

Fe^{2 +} + 6 CN^{-} ⟶ [Fe (CN)_{6}]^{4 -} .

The eventual ferric ferrocyanide material provides the blue observation. It is not sufficient simply to add any iron salt and assume that a positive blue must appear. Another classical principle converts nitrogen in suitable organic materials into Ammonia by heating with alkaline reagents, then recognizes the ammonia in a receiver. That test has substrate-dependent limitations and is not a universal replacement for complete nitrogen analysis. A later specialized nitrogen-digestion protocol must not be silently attributed to 1865.

For sulfur, a sulfide-containing alkaline extract can give black lead sulfide with a suitably introduced lead reagent:

Pb^{2 +} + S^{2 -} ⟶ PbS ↓ .