Chapter 11: Study of Compounds: Sulphuric Acid
In 1831 a vinegar manufacturer in Bristol named Peregrine Phillips took out a patent that almost nobody used. His idea was simple to state. Pass sulphur dioxide and air over a hot surface of platinum, he said, and the two gases will join to give sulphur trioxide, which can then be turned into sulphuric acid. The gases met on the surface of the metal, made contact with it, and the process later took its name from that word.
Phillips never made a fortune from it. Platinum was expensive, and the smallest trace of certain impurities in the gas ruined its power to speed the reaction. For half a century, factories went on making the acid the older, dirtier way, in vast lead-lined rooms. Only when chemists learned to purify the gas stream carefully, and later to replace platinum with a cheaper oxide of vanadium, did the contact process take over the world’s acid supply.
That supply matters more than its dull name suggests. The German chemist Justus von Liebig judged a country’s industrial strength by the quantity of sulphuric acid it consumed, and the measure still works surprisingly well. Fertilisers, paints, detergents, dyes, storage batteries and refined petrol all pass through this one oily liquid on their way to you.
Diagram
Fig. 11.0-1 — Sulphuric Acid In Everyday Products A central labelled flask of colourless oily liquid marked sulphuric acid, with arrows radiating to six labelled items: a bag of ammonium sulphate fertiliser, a car battery, a tin of paint, a detergent pack, a length of rayon cloth, and a fuel pump nozzle.
11.1 Introduction
Of all the substances made in chemical plants, one is produced in greater quantity than any other, and it is not a fuel, a plastic or a medicine. It is a colourless, oily, heavy liquid called sulphuric acid. Because so many key industries depend on it — fertilisers, dyes, detergents, paints, batteries, refined petrol — it has long been called the king of chemicals.
The name is not flattery. A factory that makes fertiliser needs the acid; a factory that makes the dye needs it; even the plant that makes another acid usually begins with this one. Nothing else manufactured is called upon by so many different industries at once.
11.1.1 Formula, Structure And Relative Molecular Mass
A molecule of sulphuric acid contains two atoms of hydrogen, one of sulphur and four of oxygen, so its molecular formula is \(\mathrm{H_2SO_4}\).
Adding the atomic masses of these atoms gives the relative molecular mass. Hydrogen contributes \(2 \times 1 = 2\), sulphur contributes \(32\), and the four oxygen atoms contribute \(4 \times 16 = 64\). The three add up to \(98\), so the relative molecular mass of sulphuric acid is \(98\).
Inside the molecule, the sulphur atom sits at the centre. Two of its oxygen atoms each carry a hydrogen atom, giving two \(\mathrm{-O-H}\) groups; the other two oxygen atoms are joined to sulphur alone. It is the two hydrogen atoms of those \(\mathrm{-O-H}\) groups that the acid can release in water, and almost every reaction in this chapter turns on that fact. The arrangement is shown in Fig. 11.1-1.
Diagram
Fig. 11.1-1 — Structure Of The Sulphuric Acid Molecule Central S atom; one O drawn directly above it and one directly below, each joined to S by an arrow pointing away from S (coordinate bond); to the left and right, an O atom each joined to S by a plain line and carrying an H atom on its outer side, giving H–O–S–O–H across the middle. Label the two outer groups as hydroxyl (–O–H) groups and note that these supply the two replaceable hydrogen atoms.
11.1.2 Oil Of Vitriol: The Older Route To The Acid
Sulphuric acid has been known for a very long time. In the later Middle Ages it was made by heating the green crystals of hydrated iron(II) sulphate, known then as green vitriol. What distilled over and condensed was a thick, oily liquid, and the acid took the name oil of vitriol from the crystals it came out of.
Heating drives off the water of crystallisation and breaks the sulphate down as well. In words: hydrated iron(II) sulphate on strong heating gives iron(III) oxide, sulphur dioxide, sulphur trioxide and water vapour.
\[2\,\mathrm{FeSO_4\!\cdot\!7H_2O}\text{(s)} \xrightarrow{\ \Delta\ } \mathrm{Fe_2O_3}\text{(s)} + \mathrm{SO_2}\text{(g)} + \mathrm{SO_3}\text{(g)} + 14\,\mathrm{H_2O}\text{(g)}\]
The useful product here is the sulphur trioxide. Passed into water, it gives the acid itself.
\[\mathrm{SO_3}\text{(g)} + \mathrm{H_2O}\text{(l)} \rightarrow \mathrm{H_2SO_4}\text{(aq)}\]
That second equation is worth remembering, because the modern factory described later in this chapter has to reach the same product by a distinctly less direct route.
11.1.3 Occurrence In The Free And Combined States
Sulphuric acid occurs in the free state — that is, as the acid itself — in certain mineral springs, where it is formed by the action of water on sulphide minerals underground.
Far more of it occurs in the combined state, locked up in sulphates. Three of these minerals are worth naming: barytes, \(\mathrm{BaSO_4}\); gypsum, \(\mathrm{CaSO_4\!\cdot\!2H_2O}\); and kieserite, \(\mathrm{MgSO_4\!\cdot\!H_2O}\).
11.2 Preparation Of Sulphuric Acid
Before looking at how sulphuric acid is made by the tonne in a factory, it is worth seeing the small-scale chemistry that produces it. Two routes are described here. Both work, both are reactions you should be able to write, and neither is used industrially — one gives only a dilute acid, and the other wastes an expensive acid to make a cheaper one.
Notice what the two routes have in common. In each of them, sulphur or one of its compounds is oxidised, that is, made to combine with more oxygen. Sulphuric acid is the most oxygen-rich of the common sulphur compounds, so every road to it runs through oxidation.
11.2.1 Oxidation Of An Aqueous Solution Of Sulphur Dioxide
When sulphur dioxide is passed into water it dissolves and forms sulphurous acid, \(\mathrm{H_2SO_3}\) — an acid with one oxygen atom fewer than sulphuric acid. Supply that missing oxygen, and the solution becomes sulphuric acid. Oxygen, chlorine and bromine will each do the job.
With oxygen, the change in words is: sulphur dioxide + water + oxygen → sulphuric acid.
\[2\,\mathrm{SO_2}\text{(g)} + 2\,\mathrm{H_2O}\text{(l)} + \mathrm{O_2}\text{(g)} \rightarrow 2\,\mathrm{H_2SO_4}\text{(aq)}\]
With chlorine and with bromine there is something to watch as well as something to write, and the observation comes first. Bubble sulphur dioxide through chlorine water and the greenish-yellow colour of the chlorine fades away; do the same with bromine water and the reddish-brown colour fades. The colour belonged to the free halogen, and it disappears because the halogen has been converted into a colourless hydrogen halide in solution.
\[\mathrm{SO_2}\text{(g)} + 2\,\mathrm{H_2O}\text{(l)} + \mathrm{Cl_2}\text{(g)} \rightarrow \mathrm{H_2SO_4}\text{(aq)} + 2\,\mathrm{HCl}\text{(aq)}\]
\[\mathrm{SO_2}\text{(g)} + 2\,\mathrm{H_2O}\text{(l)} + \mathrm{Br_2}\text{(aq)} \rightarrow \mathrm{H_2SO_4}\text{(aq)} + 2\,\mathrm{HBr}\text{(aq)}\]
The fading of the colour is set out in Fig. 11.2-1. It is a useful piece of evidence: since the halogen has been used up and the sulphur has gained oxygen, the sulphur dioxide here is behaving as a reducing agent.
Diagram
Fig. 11.2-1 — Sulphur Dioxide Decolourises Bromine Water Two test tubes side by side. Left tube labelled ‘before’: reddish-brown bromine water with a delivery tube dipping into it carrying sulphur dioxide from a flask. Right tube labelled ‘after’: the same tube now colourless, labelled solution contains sulphuric acid and hydrobromic acid.
11.2.2 Action Of Concentrated Nitric Acid On Sulphur
Sulphur itself can be oxidised straight to sulphuric acid if a strong enough oxidising agent is used. Concentrated nitric acid is such an agent. Warmed with sulphur, it gives sulphuric acid, and dense brown fumes of nitrogen dioxide come off.
In words: sulphur + concentrated nitric acid → sulphuric acid + nitrogen dioxide + water.
\[\mathrm{S}\text{(s)} + 6\,\mathrm{HNO_3}\text{(aq)} \rightarrow \mathrm{H_2SO_4}\text{(aq)} + 6\,\mathrm{NO_2}\text{(g)} + 2\,\mathrm{H_2O}\text{(l)}\]
Nitrogen dioxide is poisonous, so the reaction is carried out only in a fume cupboard, never on an open bench. That, together with the cost of nitric acid, is why the reaction stays in the laboratory and never enters a factory.
11.3 Manufacture Of Sulphuric Acid: The Contact Process
Almost all the sulphuric acid used in the world is made by the contact process. The whole plant exists to carry out one difficult step — persuading sulphur dioxide to take up a further atom of oxygen — and everything before and after that step is preparation or clean-up.
There are five stages in all, and they are set out end to end in FC 11.1. Read it first, then take the stages one at a time.
Flowchart
FC 11.1 — The Contact Process, Stage By Stage Sulphur or iron pyrites burnt in air → sulphur dioxide + air → purification of the gas mixture (dust removed, cooled, scrubbed, dried, arsenic removed, purity tested) → catalytic oxidation over vanadium(V) oxide at about 450°C → sulphur trioxide → absorbed in 98% sulphuric acid in the absorption tower → oleum (H₂S₂O₇) → diluted with a calculated quantity of water in the dilution tank → sulphuric acid of the required strength
11.3.1 Production Of Sulphur Dioxide
The plant begins by burning sulphur into sulphur dioxide, and it can start from either of two raw materials.
The first is the ore iron pyrites, \(\mathrm{FeS_2}\). Heating a sulphide ore strongly in a plentiful supply of air is called roasting, and here it gives iron(III) oxide as a residue and sulphur dioxide as the useful gas.
\[4\,\mathrm{FeS_2}\text{(s)} + 11\,\mathrm{O_2}\text{(g)} \rightarrow 2\,\mathrm{Fe_2O_3}\text{(s)} + 8\,\mathrm{SO_2}\text{(g)}\]
The second material is sulphur itself, mined at a purity of about 99.5 per cent and simply burnt in clean air.
\[\mathrm{S}\text{(s)} + \mathrm{O_2}\text{(g)} \rightarrow \mathrm{SO_2}\text{(g)}\]
Burning sulphur is the tidier route, because the gas that leaves the burner is already fairly clean. Pyrites brings dust and other impurities with it, and those must all be taken out before the gas goes any further.
11.3.2 Purification Of The Gases
The mixture of sulphur dioxide and air leaving the burners carries dust, moisture and traces of arsenic compounds. Every one of these must go, because the catalyst waiting further along the plant is easily spoiled. A substance that destroys a catalyst’s power in this way is said to poison it, and a poisoned catalyst has to be replaced at great expense.
The gas is cleaned in stages, each stage removing one kind of impurity.
Solid particles are removed first in the dust chamber, where a downward blow of steam knocks them out of the stream. The gases are then cooled by passing them through pipes, and led into a scrubbing tower, where a downward spray of water washes out the remaining dust and soluble impurities.
Fine solid particles that survive the washing are taken out in an electrostatic precipitator — a chamber strung with wires held at a high electric potential, which attracts the charged solid particles and drops them out of the stream.
The gas is now clean but wet, and moisture would dilute the acid later on, so it is dried in a drying tower by a downward spray of concentrated sulphuric acid. Finally it passes over ferric hydroxide in the arsenic purifier, which holds back every trace of arsenic oxide, and through a testing box, where a beam of light shone through a darkened chamber shows at once whether any dust remains.
11.3.3 Catalytic Oxidation Of Sulphur Dioxide
The clean, dry mixture of sulphur dioxide and air now enters the contact tower, a tall iron vessel holding vertical pipes packed loosely with the catalyst on perforated shelves. This tower is also called the converter, and it is here that the gases make contact with the catalyst surface — the step that gives the whole process its name.
In words: sulphur dioxide + oxygen → sulphur trioxide, with heat given out.
\[2\,\mathrm{SO_2}\text{(g)} + \mathrm{O_2}\text{(g)} \xrightleftharpoons[\ 450^\circ\text{C}\ ]{\ \mathrm{V_2O_5}\ \text{or Pt}\ } 2\,\mathrm{SO_3}\text{(g)} + \Delta\]
The double arrow says that the reaction does not go entirely to the right; some sulphur trioxide keeps breaking back down into sulphur dioxide and oxygen. The conditions chosen for the tower are all aimed at pushing the balance as far towards sulphur trioxide as possible, and they are dealt with in 11.3.6.
The reaction gives out a great deal of heat. So the catalyst is warmed to about 450 °C only at the start; once the reaction is running, the heat it releases keeps the tower at that temperature and no further heating is needed.
11.3.4 Absorption Of Sulphur Trioxide To Form Oleum
The hot sulphur trioxide is cooled in a heat exchanger and passed into the absorption tower, down which concentrated sulphuric acid of 98 per cent strength is trickling. The gas dissolves in that acid to give a heavy, fuming liquid called oleum, or pyrosulphuric acid.
\[\mathrm{SO_3}\text{(g)} + \mathrm{H_2SO_4}\text{(l)} \rightarrow \mathrm{H_2S_2O_7}\text{(l)}\]
At first sight this is a needless detour, since sulphur trioxide and water give sulphuric acid directly. The trouble is that sulphur trioxide does not dissolve in water readily, and the reaction releases so much heat that the acid formed is thrown up as a dense mist of fine droplets. That mist is difficult to condense and drifts out of the plant, so water is never used as the absorbing liquid.
11.3.5 Dilution Of Oleum To Obtain Sulphuric Acid
The oleum is run into a dilution tank, where a calculated quantity of water is added to it. Each molecule of oleum takes up one molecule of water and splits into two of sulphuric acid.
\[\mathrm{H_2S_2O_7}\text{(l)} + \mathrm{H_2O}\text{(l)} \rightarrow 2\,\mathrm{H_2SO_4}\text{(l)}\]
Because the quantity of water is measured, the plant can turn out acid of whatever strength a customer needs, from the 98 per cent acid sold as concentrated sulphuric acid down to much weaker grades.
11.3.6 Favourable Conditions For The Conversion Of Sulphur Dioxide
The conversion of sulphur dioxide to sulphur trioxide is the step the whole plant is built around, and four conditions decide how much of it is obtained.
Low Temperature
The reaction gives out heat, and a reaction that gives out heat yields more product at a low temperature than at a high one. The temperature should therefore be kept as low as the catalyst will allow, and the yield is found to be greatest at about 410–450 °C.
High Pressure
The two gases on the left occupy more volume than the sulphur trioxide they form, so raising the pressure squeezes the reaction towards the product. In practice, though, towers able to resist a strongly acidic gas at high pressure are difficult to build, so only 1–2 atmospheres are used.
Excess Of Oxygen
Air is supplied in more than the quantity the equation demands. The extra oxygen drives the reaction forward and increases the amount of sulphur trioxide formed.
A Suitable Catalyst
Platinum is the more efficient catalyst of the two. It is also expensive, and it is poisoned very easily by impurities such as arsenic(III) oxide. Vanadium(V) oxide is cheaper and much more resistant to poisoning, and it is the catalyst actually used.
11.3.7 The Contact Plant And The Work Of Its Parts
The complete plant is shown in Fig. 11.3-1, and each of its parts does one of the jobs described in the sections above. Table 11.1 sets them out in the order the gas meets them.
Diagram
Fig. 11.3-1 — Plant For The Manufacture Of Sulphuric Acid By The Contact Process Left to right flow diagram. Air blower feeding pyrite or sulphur burners; outlet labelled SO₂ + air passing into a dust chamber with a steam inlet at the top; then cooling pipes; then a scrubbing tower with a water spray at the top; then a drying tower with a conc. H₂SO₄ spray at the top; then the arsenic purifier; then the testing box; then the contact tower (converter) with vertical pipes packed with V₂O₅ catalyst; then the absorption tower with conc. H₂SO₄ (98%) entering at the top and waste gas leaving at the top; oleum drawn from the base into a dilution tank with a water inlet; final outlet labelled sulphuric acid. Label every vessel.
| Part of the plant | Function |
|---|---|
| Blower | Passes purified air or oxygen into the plant |
| Sulphur or iron pyrite burners | Produce sulphur dioxide by burning sulphur or roasting iron pyrites |
| Dust chamber | Removes dust particles by a downward blow of steam |
| Cooling pipes | Cool the gases as they pass through isolated pipes |
| Scrubbing tower | Removes further impurities and dust by a downward spray of water |
| Drying tower | Dries the moist gases by a downward spray of concentrated sulphuric acid |
| Arsenic purifier | Removes arsenic oxide as the gas passes over ferric hydroxide |
| Testing box | Tests the purity of the gases by passing light through a darkened box |
| Contact tower | Oxidises sulphur dioxide to sulphur trioxide over vanadium(V) oxide, heated initially to 450 °C |
| Absorption tower | Absorbs sulphur trioxide in a descending stream of 98% sulphuric acid |
| Dilution tank | Dilutes the oleum with the required quantity of water |
Table 11.1 — Parts of the contact plant and the work each one does
11.4 Properties Of Sulphuric Acid
Sulphuric acid is unusual in that its behaviour changes completely with how much water is present. The dilute acid behaves like any typical acid. The concentrated acid does things no ordinary acid does — it oxidises, it dehydrates, and it drives other acids out of their salts. Keep that division in mind throughout this section.
11.4.1 Physical Properties
Pure sulphuric acid is a colourless, odourless liquid, noticeably heavier and more oily than water. Its properties are collected in Table 11.2.
| Property | Sulphuric acid |
|---|---|
| Colour | Colourless |
| Odour | Odourless |
| Taste | Slightly sour |
| Nature | Dense, oily, hygroscopic liquid; it absorbs moisture from the air, so it is always kept stoppered |
| Density | Pure acid: 1.85 g/cc |
| Boiling point | 338 °C |
| Melting point | Pure acid freezes to colourless crystals at 10.4 °C |
| Solubility | Soluble in water in all proportions |
| Conductivity | Pure acid: almost a non-conductor. Dilute acid: a good conductor |
| Constant boiling mixture | Boils constantly at 338 °C, containing 98.5% of the acid |
| Physiological nature | The concentrated acid is highly corrosive and chars the skin black |
Table 11.2 — Physical properties of sulphuric acid
Two of these entries deserve a word. The acid is hygroscopic, meaning it takes up moisture from the air around it; a bottle left open gains water and loses strength, which is why the stopper matters. And the sour taste is a property of the substance, not an invitation — no acid in a laboratory is ever tasted.
Notice also the two conductivity entries. The pure acid barely conducts electricity, while the dilute acid conducts well. The next section explains why.
11.4.2 Ionisation And The Dibasic Nature Of The Acid
Pure sulphuric acid, with no water in it at all, does not ionise, and a substance that does not release hydrogen ions cannot behave as an acid. That is why the pure liquid hardly conducts a current.
Add water and everything changes. Each molecule hands a hydrogen ion to a water molecule, producing the hydronium ion, \(\mathrm{H_3O^{+}}\), and it does so in two separate stages.
\[\mathrm{H_2SO_4} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{HSO_4^{-}}\]
\[\mathrm{HSO_4^{-}} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{SO_4^{2-}}\]
Because two hydrogen ions can be released from one molecule, sulphuric acid is a dibasic acid. The two replaceable hydrogen atoms are exactly the two you met in the structure in Fig. 11.1-1.
11.4.2.1 Acid Salts And Normal Salts
A dibasic acid can lose its two hydrogen atoms one at a time, so it forms two kinds of salt. If only one hydrogen atom is replaced, the salt still contains a replaceable hydrogen atom and is called an acid salt. If both are replaced, the salt is a normal salt.
Sodium chloride shows both. Heated gently with the concentrated acid, below 200 °C, it gives sodium hydrogen sulphate, an acid salt.
\[\mathrm{NaCl} + \mathrm{H_2SO_4} \xrightarrow{\ \text{below } 200^\circ\text{C}\ } \mathrm{NaHSO_4} + \mathrm{HCl}\]
Heated more strongly, above 200 °C, with twice as much salt, the second hydrogen atom goes too and sodium sulphate is formed.
\[2\,\mathrm{NaCl} + \mathrm{H_2SO_4} \xrightarrow{\ \text{above } 200^\circ\text{C}\ } \mathrm{Na_2SO_4} + 2\,\mathrm{HCl}\]
11.4.3 Dilution Of Sulphuric Acid
Mixing the concentrated acid with water releases a very large amount of heat, and how the two are brought together decides whether that heat is safe.
Water is never poured on to the acid. The water sits on top, the heat released turns it instantly into steam, and the steam throws burning acid out of the vessel. Instead the acid is added slowly to a larger quantity of water, in a thin stream, with constant stirring. Being denser, the acid sinks and mixes, and the heat is carried away by the whole body of water. The rule to remember is simple: acid into water, and never the reverse. The correct method is shown in Fig. 11.4-1.
Diagram
Fig. 11.4-1 — The Safe Way To Dilute Sulphuric Acid Two panels. Left panel marked with a tick: a beaker of water on a bench, concentrated sulphuric acid being poured slowly down the side from a small measuring cylinder, a glass rod stirring, label ‘acid added slowly to water, stirring throughout’. Right panel marked with a cross: water being poured into a beaker of concentrated acid, with steam and droplets spurting out of the mouth, label ‘never add water to the acid’. Show safety goggles beside both panels.
A related point: the concentrated acid is less volatile than the dilute acid, because it boils at the much higher temperature of 338 °C. That single fact becomes important in 11.4.5.1.
11.4.4 Properties Of Dilute Sulphuric Acid
Dilute sulphuric acid behaves as a typical strong acid. It is fully ionised, and the five reactions below are the ones any acid of that kind will give.
11.4.4.1 Action On Active Metals
Metals placed above hydrogen in the activity series displace hydrogen from the dilute acid at ordinary temperature, forming the metal sulphate. Drop a piece of magnesium into the dilute acid and bubbles rise at once; with zinc and with iron the same thing happens more slowly.
In words: active metal + dilute sulphuric acid → metal sulphate + hydrogen.
\[\mathrm{Mg}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{MgSO_4}\text{(aq)} + \mathrm{H_2}\text{(g)}\]
\[\mathrm{Zn}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{ZnSO_4}\text{(aq)} + \mathrm{H_2}\text{(g)}\]
\[\mathrm{Fe}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{FeSO_4}\text{(aq)} + \mathrm{H_2}\text{(g)}\]
Metals below hydrogen give no hydrogen at all. Copper is one such metal, and the noble metals gold and platinum do not react with sulphuric acid under any conditions.
11.4.4.2 Action On Metal Oxides And Metal Hydroxides
Metal oxides and metal hydroxides are bases, and the acid neutralises them to give a salt and water alone.
\[\mathrm{CuO}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{CuSO_4}\text{(aq)} + \mathrm{H_2O}\text{(l)}\]
\[\mathrm{FeO}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{FeSO_4}\text{(aq)} + \mathrm{H_2O}\text{(l)}\]
\[\mathrm{Zn(OH)_2}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{ZnSO_4}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)}\]
\[2\,\mathrm{NaOH}\text{(aq)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{Na_2SO_4}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)}\]
Look at the last two equations and you will see why two formula units of sodium hydroxide are needed while one of zinc hydroxide is enough: the acid has two hydrogen ions to give away, and each hydroxide ion takes one.
11.4.4.3 Action On Carbonates And Hydrogen Carbonates
Add the dilute acid to a carbonate and there is brisk effervescence. The gas turns lime water milky, and it is therefore carbon dioxide.
\[\mathrm{Na_2CO_3}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{Na_2SO_4}\text{(aq)} + \mathrm{H_2O}\text{(l)} + \mathrm{CO_2}\text{(g)}\]
\[\mathrm{CuCO_3}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{CuSO_4}\text{(aq)} + \mathrm{H_2O}\text{(l)} + \mathrm{CO_2}\text{(g)}\]
Hydrogen carbonates behave in the same way. Since each of them carries only one carbonate group’s worth of hydrogen, two formula units react with one of the acid.
\[2\,\mathrm{NaHCO_3}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{Na_2SO_4}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)} + 2\,\mathrm{CO_2}\text{(g)}\]
\[2\,\mathrm{KHCO_3}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{K_2SO_4}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)} + 2\,\mathrm{CO_2}\text{(g)}\]
11.4.4.4 Action On Metal Sulphides
With a metal sulphide, the dilute acid releases hydrogen sulphide — a gas with the smell of rotten eggs, and a poisonous one, so the test tube is kept in a fume cupboard.
\[\mathrm{Na_2S}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{Na_2SO_4}\text{(aq)} + \mathrm{H_2S}\text{(g)}\]
\[\mathrm{FeS}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{FeSO_4}\text{(aq)} + \mathrm{H_2S}\text{(g)}\]
\[\mathrm{ZnS}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{ZnSO_4}\text{(aq)} + \mathrm{H_2S}\text{(g)}\]
11.4.4.5 Action On Sulphites And Hydrogen Sulphites
Sulphites and hydrogen sulphites give off sulphur dioxide, recognised by its sharp, choking smell.
\[\mathrm{Na_2SO_3}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{Na_2SO_4}\text{(aq)} + \mathrm{H_2O}\text{(l)} + \mathrm{SO_2}\text{(g)}\]
\[2\,\mathrm{NaHSO_3}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{Na_2SO_4}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)} + 2\,\mathrm{SO_2}\text{(g)}\]
Sulphur dioxide is poisonous, and released from chimneys in quantity it damages human health and vegetation, so industrial plants are required to trap it rather than vent it.
The Bleaching Action Of Sulphur Dioxide
Sulphur dioxide can act as an oxidising agent and as a reducing agent. In the presence of moisture it also bleaches, and it does so by its reducing action: with water it produces nascent hydrogen, written \([\mathrm{H}]\), which is hydrogen at the moment of its formation and far more reactive than ordinary hydrogen gas.
\[\mathrm{SO_2} + 2\,\mathrm{H_2O} \rightarrow \mathrm{H_2SO_4} + 2[\mathrm{H}]\]
\[\text{coloured material} + [\mathrm{H}] \rightarrow \text{colourless product}\]
The bleaching is only temporary. Left in air, the colourless material is oxidised by the oxygen of the air and its original colour returns.
11.4.5 Properties Of Concentrated Sulphuric Acid
With little water present, the acid ionises hardly at all, and the properties of the following pages have nothing to do with hydrogen ions. They belong to the \(\mathrm{H_2SO_4}\) molecule itself.
11.4.5.1 Non-Volatile Nature
A liquid that boils at 338 °C does not escape easily as vapour, so sulphuric acid is described as non-volatile. Acids such as hydrochloric acid, nitric acid and acetic acid boil far lower and are volatile.
Warm the concentrated acid with the salt of a volatile acid, and the volatile acid is driven off as a vapour and can be collected, while the sulphate stays behind in the flask. This is a double decomposition: the two compounds simply exchange partners, and the reaction goes to completion because one product keeps leaving. Table 11.3 gives the standard examples.
| Salt of volatile acid | + concentrated acid | → Acid salt | + Volatile acid |
|---|---|---|---|
| NaCl | H₂SO₄ | NaHSO₄ | HCl |
| KCl | H₂SO₄ | KHSO₄ | HCl |
| NaNO₃ | H₂SO₄ | NaHSO₄ | HNO₃ |
| KNO₃ | H₂SO₄ | KHSO₄ | HNO₃ |
| CH₃COONa (sodium acetate) | H₂SO₄ | NaHSO₄ | CH₃COOH (acetic acid) |
Table 11.3 — Volatile acids prepared from their salts by concentrated sulphuric acid
Written out, the preparation of nitric acid runs as follows.
\[\mathrm{KNO_3} + \mathrm{H_2SO_4}\text{(conc.)} \xrightarrow{\ \Delta\ } \mathrm{KHSO_4} + \mathrm{HNO_3}\]
11.4.5.2 Concentrated Sulphuric Acid As An Oxidising Agent
On heating, concentrated sulphuric acid decomposes and yields nascent oxygen, written \([\mathrm{O}]\).
\[\mathrm{H_2SO_4} \rightarrow \mathrm{H_2O} + \mathrm{SO_2} + [\mathrm{O}]\]
That nascent oxygen is what does the oxidising, and it will attack non-metals, metals and certain compounds. In every case the sulphur is reduced and sulphur dioxide comes off, so the smell of sulphur dioxide is the sign that the acid has acted as an oxidising agent.
Non-Metals
Hot concentrated acid oxidises carbon to carbon dioxide, and sulphur to sulphur dioxide.
\[\mathrm{C}\text{(s)} + 2\,\mathrm{H_2SO_4}\text{(conc.)} \rightarrow \mathrm{CO_2}\text{(g)} + 2\,\mathrm{H_2O}\text{(l)} + 2\,\mathrm{SO_2}\text{(g)}\]
\[\mathrm{S}\text{(s)} + 2\,\mathrm{H_2SO_4}\text{(conc.)} \rightarrow 3\,\mathrm{SO_2}\text{(g)} + 2\,\mathrm{H_2O}\text{(l)}\]
The second of these is worth a second look. All three molecules of sulphur dioxide contain sulphur, but only one of those sulphur atoms came from the element; the other two came from the acid.
Metals
Copper does not react with the dilute acid at all, yet with the hot concentrated acid it dissolves to give a blue solution and sulphur dioxide. Zinc, which gives hydrogen with the dilute acid, gives sulphur dioxide with the concentrated acid instead.
\[\mathrm{Cu}\text{(s)} + 2\,\mathrm{H_2SO_4}\text{(conc.)} \rightarrow \mathrm{CuSO_4}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)} + \mathrm{SO_2}\text{(g)}\]
\[\mathrm{Zn}\text{(s)} + 2\,\mathrm{H_2SO_4}\text{(conc.)} \rightarrow \mathrm{ZnSO_4}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)} + \mathrm{SO_2}\text{(g)}\]
Inorganic Compounds
Compounds that are easily oxidised are attacked in the same way. Hydrogen bromide gives free bromine, and hydrogen sulphide gives free sulphur.
\[2\,\mathrm{HBr} + \mathrm{H_2SO_4}\text{(conc.)} \rightarrow \mathrm{Br_2} + 2\,\mathrm{H_2O} + \mathrm{SO_2}\text{(g)}\]
\[\mathrm{H_2S} + \mathrm{H_2SO_4}\text{(conc.)} \rightarrow \mathrm{S} + 2\,\mathrm{H_2O} + \mathrm{SO_2}\text{(g)}\]
11.4.5.3 Concentrated Sulphuric Acid As A Dehydrating Agent
Concentrated sulphuric acid has a great affinity for water. It will pull out not only free moisture but the elements of water — hydrogen and oxygen in the ratio 2 : 1 — that are chemically combined inside another compound. Acting in this way, it is a dehydrating agent.
Removing Water Of Crystallisation
Blue crystals of copper(II) sulphate contain five molecules of water of crystallisation each. Add a few drops of the concentrated acid to them in a test tube and the blue colour is destroyed; what remains is the dirty-white anhydrous salt, the water having been taken away by the acid.
\[\mathrm{CuSO_4\!\cdot\!5H_2O}\text{(s)} \xrightarrow{\ \text{conc. } \mathrm{H_2SO_4}\ } \mathrm{CuSO_4}\text{(s)} + 5\,\mathrm{H_2O}\]
Charring Carbohydrates
Carbohydrates — glucose, cane sugar, and the cellulose of paper, cotton and wood — contain hydrogen and oxygen in exactly the proportion found in water. The acid removes all of it and leaves the carbon behind. Pour the concentrated acid on to sugar crystals in a hard glass test tube and a black, spongy mass of carbon rises out of the tube, steam is given off, and the whole mass grows hot because the reaction is exothermic. The substance is said to be charred. The change is shown in Fig. 11.4-2.
\[\mathrm{C_6H_{12}O_6}\text{(s)} \xrightarrow{\ \text{conc. } \mathrm{H_2SO_4}\ } 6\,\mathrm{C}\text{(s)} + 6\,\mathrm{H_2O}\]
\[\mathrm{C_{12}H_{22}O_{11}}\text{(s)} \xrightarrow{\ \text{conc. } \mathrm{H_2SO_4}\ } 12\,\mathrm{C}\text{(s)} + 11\,\mathrm{H_2O}\]
The carbon left behind is very pure and is called sugar charcoal.
This is also why the acid is so dangerous on skin. It removes water from the flesh, raising blisters and charring the skin black, which is why goggles and gloves are worn whenever the concentrated acid is handled and why a spill is washed at once with plenty of water.
Diagram
Fig. 11.4-2 — Action Of Concentrated Sulphuric Acid On Sugar Two beakers side by side. Left beaker: white sugar crystals with concentrated sulphuric acid being poured on from a small flask, labelled crystals of sugar. Right beaker: a tall black spongy column of carbon risen well above the rim, with wavy arrows at the top labelled steam, and the mass labelled black spongy carbon mass (sugar charcoal).
Dehydrating Organic Acids And Alcohols
Certain organic compounds lose water to the acid and give a gas or a new compound.
\[\mathrm{HCOOH} \xrightarrow{\ \text{conc. } \mathrm{H_2SO_4}\ } \mathrm{CO} + \mathrm{H_2O}\]
\[\mathrm{(COOH)_2} \xrightarrow{\ \text{conc. } \mathrm{H_2SO_4}\ } \mathrm{CO} + \mathrm{CO_2} + \mathrm{H_2O}\]
\[\mathrm{C_2H_5OH} \xrightarrow[\ 170^\circ\text{C}\ ]{\ \text{conc. } \mathrm{H_2SO_4}\ } \mathrm{C_2H_4} + \mathrm{H_2O}\]
The first two are formic acid and oxalic acid; the third is ethyl alcohol, which loses a molecule of water to become the unsaturated gas ethylene.
11.4.5.4 Preparation Of Insoluble Sulphates
Most sulphates dissolve in water, but those of lead, barium and calcium do not. So when sulphuric acid is added to a solution containing one of these metals, the sulphate appears at once as a white precipitate.
\[\mathrm{Pb(NO_3)_2}\text{(aq)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{PbSO_4}\!\downarrow + 2\,\mathrm{HNO_3}\text{(aq)}\]
\[\mathrm{BaCl_2}\text{(aq)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{BaSO_4}\!\downarrow + 2\,\mathrm{HCl}\text{(aq)}\]
This behaviour is the basis of the tests in the next section.
11.4.6 Dilute Sulphuric Acid Compared With Concentrated Sulphuric Acid
The differences that run through this whole section are collected in Table 11.4.
11.5 Tests For Sulphuric Acid And Sulphates
Suppose you are handed a bottle of colourless liquid and told only that it is an acid. Three tests will settle whether it is sulphuric acid, and each of them rests on a property established earlier in this chapter.
11.5.1 The Barium Chloride Test
This is the standard test for sulphuric acid and for any soluble sulphate. Add barium chloride solution to the liquid under test. If a sulphate is present, a white precipitate of barium sulphate appears at once, because barium sulphate is one of the few sulphates that will not dissolve in water.
You met this reaction in 11.4.5.4 as a way of preparing an insoluble sulphate; here the same reaction is being used the other way round, as evidence. A soluble sulphate such as magnesium sulphate gives exactly the same precipitate, which is why the test detects the sulphate ion rather than the acid alone.
\[\mathrm{BaCl_2}\text{(aq)} + \mathrm{MgSO_4}\text{(aq)} \rightarrow \mathrm{BaSO_4}\!\downarrow + \mathrm{MgCl_2}\text{(aq)}\]
The precipitate alone is not quite proof. Barium sulphite is also a white precipitate, so a sulphite solution would look the same at this stage. The two are separated by adding a mineral acid such as dilute hydrochloric or dilute nitric acid: barium sulphite dissolves in it, while barium sulphate does not. A white precipitate that survives the acid is a sulphate. The full test is set out in FC 11.2 and shown in Fig. 11.5-1.
Flowchart
FC 11.2 — Testing A Solution For The Sulphate Ion Solution under test → add barium chloride solution → white precipitate forms → add dilute hydrochloric acid → branch one: precipitate dissolves, so a sulphite is present and not a sulphate; branch two: precipitate does not dissolve, so the sulphate is confirmed. If no precipitate forms with barium chloride, no sulphate is present
Diagram
Fig. 11.5-1 — The Barium Chloride Test For A Sulphate Three test tubes in a rack. First: clear colourless solution under test with a dropper adding barium chloride solution. Second: dense white precipitate settling at the bottom, labelled barium sulphate. Third: dilute hydrochloric acid being added to the second tube, the precipitate unchanged, labelled precipitate insoluble in mineral acid, sulphate confirmed.
11.5.2 Heating With Copper
Warm the liquid under test with copper turnings. If it is concentrated sulphuric acid, the copper dissolves and sulphur dioxide is given off — the reaction met in 11.4.5.2, where the acid was acting as an oxidising agent.
The gas is then identified. Hold a strip of filter paper moistened with acidified potassium dichromate solution at the mouth of the tube; its orange colour turns green. The change happens because sulphur dioxide is a reducing agent, and it reduces the orange dichromate to a green chromium(III) salt.
\[\mathrm{K_2Cr_2O_7} + \mathrm{H_2SO_4} + 3\,\mathrm{SO_2} \rightarrow \mathrm{Cr_2(SO_4)_3} + \mathrm{K_2SO_4} + \mathrm{H_2O}\]
Because dilute sulphuric acid does not attack copper at all, this test does two jobs at once: it identifies the acid, and it tells the concentrated acid from the dilute.
11.5.3 Heating With Sodium Chloride
Warm the liquid with a little sodium chloride. If it is the concentrated acid, its non-volatile nature drives the volatile acid out, and pungent fumes of hydrogen chloride come off — the reaction of 11.4.2.1.
To identify those fumes, dip a glass rod in ammonium hydroxide solution and hold it at the mouth of the tube. Dense white fumes of ammonium chloride appear where the two gases meet.
\[\mathrm{HCl}\text{(g)} + \mathrm{NH_4OH}\text{(aq)} \rightleftharpoons \mathrm{NH_4Cl}\text{(s)} + \mathrm{H_2O}\text{(l)}\]
All three tests are carried out on very small quantities, with the tube pointed away from everyone, and the last two in a fume cupboard, since both hydrogen chloride and sulphur dioxide are choking gases.
11.6 Uses Of Sulphuric Acid
Every property in the last section turns into a use somewhere. The acid that chars sugar dries gases in a laboratory; the acid that will not boil away makes other acids; the acid that dissolves metal oxides cleans steel before it is coated. This section gathers those uses together.
Preparing Gases In The Laboratory
Because concentrated sulphuric acid acts both as a dehydrating agent and as a drying agent, it is used to prepare and to dry a range of gases. Halogens, carbon monoxide, carbon dioxide, hydrogen and sulphur dioxide can all be obtained with its help, and a gas such as hydrogen chloride is passed through the concentrated acid to remove the moisture it carries.
Not every gas can be dried this way. Hydrogen sulphide and hydrogen both react with the acid instead of merely losing their moisture to it, so a different drying agent is chosen for them.
\[\mathrm{H_2}\text{(g)} + \mathrm{H_2SO_4}\text{(conc.)} \rightarrow 2\,\mathrm{H_2O}\text{(l)} + \mathrm{SO_2}\text{(g)}\]
A drying agent must take the water and leave the gas alone. Here the acid oxidises the gas, so it fails the only condition that matters.
In Metallurgy
The acid has two quite separate jobs in the winning and finishing of metals.
In extraction, it converts metallic compounds into their sulphates. These sulphates dissolve, and passing a current through the solution deposits the metal in a pure form.
In pickling, sheets of metal are dipped in the dilute acid before they are galvanised. The acid dissolves away the oxides and carbonates that have formed on the surface, leaving clean metal for the coating to grip. This is simply the neutralisation of 11.4.4.2 put to industrial work.
In Storage Batteries And Oil Refining
The lead accumulator — the rechargeable battery in a car — holds dilute sulphuric acid as its electrolyte, and the acid undergoes electrolysis in the aqueous state as the battery is charged and discharged. Its construction is shown in Fig. 11.6-1.
Oil refineries use the acid to remove harmful impurities while purifying petrol, kerosene and lubricating oils.
Diagram
Fig. 11.6-1 — The Lead Accumulator Sectional view of a single cell of a car battery: a hard plastic case holding two sets of interleaved plates, the negative plate labelled spongy lead and the positive plate labelled lead(IV) oxide, both immersed in a liquid labelled dilute sulphuric acid (electrolyte); terminals marked + and − on the lid.
In Industry
The largest quantities of the acid go into manufacture, and Table 11.5 lists the main products.
| Industry | Products made with the help of sulphuric acid |
|---|---|
| Fertilisers | Ammonium sulphate, \(\mathrm{(NH_4)_2SO_4}\); superphosphate of lime, \(\mathrm{Ca(H_2PO_4)_2} + \mathrm{CaSO_4}\) |
| Fibres, dyes and drugs | Artificial fibres such as rayon and nylon; dyes and drugs from coal tar derivatives |
| Explosives | Tri-nitro toluene (T.N.T.), picric acid, tri-nitro glycerine, \(\mathrm{C_3H_5O_3(NO_2)_3}\) |
| Acids | Nitric acid, hydrochloric acid, phosphoric acid, acetic acid |
| Other compounds | Sodium sulphate, used in the glass industry; ferrous sulphate, used in the ink industry |
Table 11.5 — Industrial products that depend on sulphuric acid
Read the first row again and the opening story of this chapter makes better sense. A country that grows more food needs more fertiliser, a country that makes more fertiliser needs more sulphuric acid, and so the quantity of acid a nation consumes has long been treated as a rough measure of its industry.
11.7 Summary
Sulphuric acid, \(\mathrm{H_2SO_4}\), has a relative molecular mass of 98. Its molecule has two \(\mathrm{-O-H}\) groups joined to a central sulphur atom, and these supply its two replaceable hydrogen atoms.
It is called the king of chemicals because so many key industries depend on it.
In the Middle Ages it was made by heating green vitriol, hydrated iron(II) sulphate, and was therefore called oil of vitriol. The sulphur trioxide evolved dissolves in water to give the acid: \(\mathrm{SO_3} + \mathrm{H_2O} \rightarrow \mathrm{H_2SO_4}\).
It occurs free in some mineral springs, and combined in barytes (\(\mathrm{BaSO_4}\)), gypsum (\(\mathrm{CaSO_4\!\cdot\!2H_2O}\)) and kieserite (\(\mathrm{MgSO_4\!\cdot\!H_2O}\)).
An aqueous solution of sulphur dioxide is oxidised to sulphuric acid by oxygen, chlorine or bromine. The colour of chlorine water or bromine water fades as the halogen is used up.
Sulphur warmed with concentrated nitric acid also gives the acid: \(\mathrm{S} + 6\,\mathrm{HNO_3} \rightarrow \mathrm{H_2SO_4} + 6\,\mathrm{NO_2} + 2\,\mathrm{H_2O}\).
Sulphuric acid is manufactured by the contact process in five stages: production of sulphur dioxide, purification of the gases, catalytic oxidation, absorption in 98% acid to give oleum, and dilution of the oleum.
Sulphur dioxide is obtained by roasting iron pyrites, \(4\,\mathrm{FeS_2} + 11\,\mathrm{O_2} \rightarrow 2\,\mathrm{Fe_2O_3} + 8\,\mathrm{SO_2}\), or by burning sulphur, \(\mathrm{S} + \mathrm{O_2} \rightarrow \mathrm{SO_2}\).
The gases are freed of dust, cooled, scrubbed, dried with concentrated sulphuric acid, passed over ferric hydroxide to remove arsenic oxide, and tested for purity, because impurities poison the catalyst.
Sulphur dioxide is oxidised over vanadium(V) oxide at about 450 °C: \(2\,\mathrm{SO_2} + \mathrm{O_2} \rightleftharpoons 2\,\mathrm{SO_3}\). The reaction is exothermic, so the catalyst is heated only at the beginning.
The conditions favouring the conversion are a low temperature of 410–450 °C, a pressure of 1–2 atmospheres, an excess of oxygen, and vanadium(V) oxide as catalyst — chosen over platinum because it is cheaper and is not poisoned easily.
Sulphur trioxide is absorbed in 98% sulphuric acid, not in water, because with water it forms an unmanageable acid mist. \(\mathrm{SO_3} + \mathrm{H_2SO_4} \rightarrow \mathrm{H_2S_2O_7}\) (oleum).
Oleum is diluted with a calculated quantity of water: \(\mathrm{H_2S_2O_7} + \mathrm{H_2O} \rightarrow 2\,\mathrm{H_2SO_4}\).
Pure sulphuric acid is a colourless, odourless, dense, oily, hygroscopic liquid of density 1.85 g/cc. It boils at 338 °C and freezes at 10.4 °C, and forms a constant boiling mixture at 338 °C containing 98.5% of the acid.
The pure acid does not ionise and is almost a non-conductor; the dilute acid ionises fully and conducts well. The concentrated acid is highly corrosive and chars the skin black.
The acid ionises in two stages and is therefore a dibasic acid. It forms an acid salt, \(\mathrm{NaHSO_4}\), below 200 °C and a normal salt, \(\mathrm{Na_2SO_4}\), above 200 °C.
Acid is always added slowly to water with stirring, never water to acid, because the heat evolved would turn the added water to steam and throw the acid out of the vessel.
Dilute sulphuric acid displaces hydrogen from metals above hydrogen in the activity series, such as magnesium, zinc and iron. Gold and platinum do not react with it under any condition.
Dilute sulphuric acid neutralises metal oxides and hydroxides to a sulphate and water.
Dilute sulphuric acid liberates carbon dioxide from carbonates and hydrogen carbonates, hydrogen sulphide from metal sulphides, and sulphur dioxide from sulphites and hydrogen sulphites.
Sulphur dioxide bleaches in the presence of moisture by producing nascent hydrogen, \(\mathrm{SO_2} + 2\,\mathrm{H_2O} \rightarrow \mathrm{H_2SO_4} + 2[\mathrm{H}]\). The bleaching is temporary, and the colour returns in air.
Being non-volatile, concentrated sulphuric acid displaces volatile acids from their salts: hydrochloric acid from chlorides, nitric acid from nitrates and acetic acid from acetates.
On heating, the concentrated acid yields nascent oxygen, \(\mathrm{H_2SO_4} \rightarrow \mathrm{H_2O} + \mathrm{SO_2} + [\mathrm{O}]\), and so acts as an oxidising agent. It oxidises carbon to carbon dioxide, sulphur to sulphur dioxide, copper to copper sulphate and zinc to zinc sulphate, giving sulphur dioxide each time.
As a dehydrating agent the concentrated acid removes water of crystallisation from blue \(\mathrm{CuSO_4\!\cdot\!5H_2O}\), leaving the white anhydrous salt, and chars carbohydrates: \(\mathrm{C_{12}H_{22}O_{11}} \rightarrow 12\,\mathrm{C} + 11\,\mathrm{H_2O}\), the carbon left behind being sugar charcoal.
It also dehydrates formic acid to carbon monoxide, oxalic acid to carbon monoxide and carbon dioxide, and ethyl alcohol at 170 °C to ethylene.
It precipitates the insoluble sulphates of lead, barium and calcium from solutions of their salts.
Sulphuric acid and soluble sulphates give a white precipitate of barium sulphate with barium chloride solution, and that precipitate does not dissolve in dilute hydrochloric or nitric acid.
Warmed with copper, the concentrated acid gives sulphur dioxide, which turns acidified potassium dichromate solution from orange to green. Warmed with sodium chloride, it gives hydrogen chloride, which gives dense white fumes with a rod dipped in ammonium hydroxide.
Sulphuric acid is used to prepare and dry gases, in metallurgy for extraction and pickling, in lead accumulators, in oil refining, and in making fertilisers, fibres, dyes, drugs, explosives, other acids and sulphate compounds. It cannot dry hydrogen or hydrogen sulphide, which it oxidises.
Flowchart
FC 11.2 — Sulphuric Acid From Sulphur To Product Sulphur or iron pyrites → burnt or roasted in air → sulphur dioxide → gases purified and dried → oxidised over V₂O₅ at about 450°C → sulphur trioxide → absorbed in 98% H₂SO₄ → oleum (H₂S₂O₇) → diluted with a measured quantity of water → sulphuric acid → branches to: dilute acid, showing typical acid reactions with metals, bases, carbonates, sulphides and sulphites; and concentrated acid, acting as a non-volatile acid, an oxidising agent and a dehydrating agent
11.8 Key Terms
Acid salt — a salt formed when only one of the two replaceable hydrogen atoms of sulphuric acid is replaced, so the salt still contains a replaceable hydrogen atom, as in sodium hydrogen sulphate.
Dehydrating agent — a substance that removes the elements of water, hydrogen and oxygen in the ratio 2 : 1, from a compound in which they were chemically combined.
Dibasic acid — an acid whose one molecule contains two replaceable hydrogen atoms and can therefore donate two hydrogen ions in aqueous solution.
Hygroscopic — able to absorb moisture from the air, which is why sulphuric acid is always kept in a stoppered bottle.
King of chemicals — the name given to sulphuric acid because no other manufactured compound is used by so many key industries.
Non-volatile — having a high boiling point and so not escaping easily as vapour; sulphuric acid boils at 338 °C.
Oil of vitriol — the old name for sulphuric acid, from the oily liquid obtained by heating crystals of green vitriol.
Oleum — pyrosulphuric acid, \(\mathrm{H_2S_2O_7}\), formed when sulphur trioxide dissolves in concentrated sulphuric acid.
Oxidised — made to combine with oxygen, or to lose hydrogen or electrons; every route to sulphuric acid oxidises sulphur or one of its compounds.
Roasting — heating a sulphide ore strongly in a plentiful supply of air, as when iron pyrites is converted to iron(III) oxide and sulphur dioxide.
Sugar charcoal — the very pure carbon left behind as a black spongy mass when concentrated sulphuric acid dehydrates cane sugar.
11.9 Points To Ponder
The contact process spends a great deal of money cleaning a gas mixture before it ever reaches the catalyst. If the catalyst could be made completely resistant to poisoning, which parts of the plant could be removed — and which would still be needed for reasons that have nothing to do with the catalyst?
Sulphuric acid is used to prepare hydrochloric acid and nitric acid from their salts. Could either of those two acids be used in the same way to prepare sulphuric acid from a sulphate? Reason your answer from the property that makes the process work.
Zinc gives hydrogen with the dilute acid and sulphur dioxide with the concentrated acid, yet the formula of the acid is \(\mathrm{H_2SO_4}\) in both bottles. What does this suggest about how much of a substance’s chemistry is decided by its formula alone?
The same acid that dries hydrogen chloride destroys hydrogen sulphide. What single question would you ask about any new gas before deciding whether concentrated sulphuric acid is a safe drying agent for it?
Concentrated sulphuric acid chars sugar, paper and cotton, and it also injures skin. Are these three observations separate facts about the acid, or one fact seen in different places?
A country’s consumption of sulphuric acid has long been treated as a measure of its industrial strength. Which of the uses in 11.6 do you think best justifies that claim, and which industries would suffer first if the supply were interrupted?
11.10 Objective Questions
- The gas evolved when dilute sulphuric acid is added to sodium sulphite is: [1]
- Hydrogen
- Carbon dioxide
- Sulphur dioxide
- Hydrogen sulphide
- In the contact process, sulphur trioxide is absorbed in: [1]
- Water
- 98% sulphuric acid
- Dilute sulphuric acid
- Ferric hydroxide
- The white precipitate formed when barium chloride solution is added to a sulphate is: [1]
- Soluble in dilute nitric acid
- Soluble in dilute hydrochloric acid
- Insoluble in both these acids
- Soluble only on warming
- Concentrated sulphuric acid is warmed with potassium nitrate. The property of the acid being used is its: [1]
- Oxidising nature
- Dehydrating nature
- Non-volatile nature
- Hygroscopic nature
- Iron pyrites is converted to sulphur dioxide in the contact process by: [1]
- Calcination
- Roasting
- Electrolysis
- Reduction with carbon
- The gas that cannot be dried by concentrated sulphuric acid is: [1]
- Hydrogen chloride
- Carbon dioxide
- Sulphur dioxide
- Hydrogen sulphide
- Blue crystals of copper(II) sulphate turn dirty white on adding a few drops of concentrated sulphuric acid because the acid removes: [1]
- Sulphate ions
- Water of crystallisation
- Copper ions
- Dissolved oxygen
- The pressure actually used in the contact tower is about: [1]
- 1–2 atmospheres
- 20 atmospheres
- 200 atmospheres
- 400 atmospheres
The following questions are assertion–reason based. Choose the answer from the codes given below. (1) Both A and R are true and R is the correct explanation of A. (2) Both A and R are true but R is not the correct explanation of A. (3) A is true but R is false. (4) A is false but R is true.
- Assertion (A): In the contact tower the catalyst is heated only at the beginning. Reason (R): The oxidation of sulphur dioxide to sulphur trioxide is exothermic. [1]
- Assertion (A): Water is never poured on to concentrated sulphuric acid. Reason (R): Sulphuric acid is less dense than water. [1]
- Assertion (A): Pure sulphuric acid is almost a non-conductor of electricity. Reason (R): Sulphuric acid ionises only when water is present. [1]
- Match each substance in Column I with the product it gives when treated with sulphuric acid under the stated conditions in Column II. [4]
| Column I | Column II |
|---|---|
| a. Zinc sulphide + dilute acid | i. Sulphur dioxide |
| b. Sodium chloride + concentrated acid, below 200 °C | ii. Carbon dioxide |
| c. Potassium hydrogen carbonate + dilute acid | iii. Hydrogen sulphide |
| d. Sulphur + concentrated acid | iv. Hydrogen chloride |
- Complete the following sentences with a suitable word or phrase. [4]
- Sulphur trioxide is dissolved in concentrated sulphuric acid to give …………… .
- The catalyst used in the contact process is …………… .
- Sulphuric acid can donate two hydrogen ions and is therefore a …………… acid.
- The carbon left behind when concentrated sulphuric acid acts on cane sugar is called …………… .
- State whether each statement is true or false. [3]
- Dilute sulphuric acid liberates hydrogen from copper.
- Barium sulphite dissolves in dilute hydrochloric acid.
- Concentrated sulphuric acid is a stronger electrolyte than dilute sulphuric acid.
- Name the following. [4]
- The acid formed when sulphur dioxide dissolves in water.
- The part of the contact plant that removes arsenic oxide.
- The salt formed when one replaceable hydrogen atom of sulphuric acid is replaced by sodium.
- The gas evolved when dilute sulphuric acid acts on iron(II) sulphide.
- Give one word or phrase. [3]
- A liquid that takes up moisture from the air.
- Heating a sulphide ore strongly in a plentiful supply of air.
- A substance that destroys the activity of a catalyst.
- Pick the odd one out from each set and give a reason for your choice. [2]
- Magnesium, zinc, iron, gold.
- Barytes, gypsum, kieserite, iron pyrites.
11.11 Check Your Concepts
Define a dibasic acid, and state the two salts sulphuric acid forms with sodium chloride. [2]
State why sulphuric acid is stored in stoppered bottles. [2]
Give a reason why concentrated sulphuric acid is used to dry hydrogen chloride but not hydrogen. [2]
State what you would observe when concentrated sulphuric acid is poured on to crystals of cane sugar in a hard glass test tube. [2]
Explain why the gases entering the contact tower must be free of dust, moisture and arsenic oxide. [3]
Describe how dilute sulphuric acid is prepared from the concentrated acid, and give the reason for the order in which the two liquids are mixed. [3]
Sulphur trioxide reacts with water to give sulphuric acid, yet water is not used in the absorption tower. Explain, and state what is used instead. [3]
Differentiate between dilute and concentrated sulphuric acid, stating the basis of difference in each case, with respect to (i) electrical conduction and (ii) action on zinc. [4]
Explain, with one example of each, what is meant by calling concentrated sulphuric acid (i) an oxidising agent and (ii) a dehydrating agent. [4]
State the four conditions favourable to the conversion of sulphur dioxide to sulphur trioxide, and give the reason behind each one. [4]
Why does a sample of pure sulphuric acid hardly conduct electricity while its dilute solution conducts well? Support your answer with the two stages of ionisation. [4]
Describe the five stages of the contact process, naming the raw material, the catalyst, the temperature used and the product of each stage. [5]
11.12 Equations, Conversions & Identification
Write balanced equations for the action of dilute sulphuric acid on: [3]
- Magnesium
- Zinc hydroxide
- Sodium hydrogen carbonate
Write balanced equations for the action of hot concentrated sulphuric acid on: [3]
- Carbon
- Copper
- Sulphur
Complete and balance the following: [3]
- \(\mathrm{FeS} + \mathrm{H_2SO_4}\text{(dil.)} \rightarrow\)
- \(\mathrm{H_2S_2O_7} + \mathrm{H_2O} \rightarrow\)
- \(\mathrm{Pb(NO_3)_2} + \mathrm{H_2SO_4} \rightarrow\)
Dilute sulphuric acid is added to sodium sulphite. Name the gas evolved, give one test by which you would identify it, and state what you would see in that test. [3]
State what you would see when: [3]
- Sulphur dioxide is bubbled through bromine water.
- Barium chloride solution is added to magnesium sulphate solution, and dilute nitric acid is then added to the mixture.
- A few drops of concentrated sulphuric acid are added to blue copper(II) sulphate crystals.
How would you obtain, in one or more steps, writing equations for each step: [4]
- Hydrogen chloride from sodium chloride
- Sulphuric acid from sulphur
- Zinc sulphate from zinc oxide
A white solid P reacts with dilute sulphuric acid to give a colourless gas Q with a choking smell. Q turns acidified potassium dichromate solution green and decolourises bromine water. Identify P and Q, and write the equation for the reaction between P and the acid. [4]
Study Fig. 11.3-1. Name the part of the plant labelled as the drying tower, the arsenic purifier and the contact tower, and state the purpose of each. [3]
Study Fig. 11.5-1. Name the reagent being added in the first tube, name the precipitate formed, and state why dilute hydrochloric acid is added in the third tube. [3]