Hello, and welcome to today's chemistry lesson. We are going to explore one of the most important compounds in the chemical industry: sulphuric acid. Often called the King of Chemicals, this remarkable substance touches nearly every aspect of modern manufacturing. By the end of this lesson, you will understand how sulphuric acid is produced on a massive scale, why it behaves so differently when dilute versus concentrated, and what makes it such a versatile reagent.
Let us begin with some historical context. Sulphuric acid has been known for centuries. In medieval times, it was obtained by heating green vitriol crystals, which are FeSO₄·7H₂O, also known as hydrated iron two sulphate or ferrous sulphate. When heated, this compound decomposes to produce iron three oxide, sulphur dioxide, sulphur trioxide, and water. The reaction is: two molecules of FeSO₄·7H₂O yield Fe₂O₃, sulphur dioxide, sulphur trioxide, and fourteen water molecules. The sulphur trioxide then dissolves in water to form sulphuric acid. Because of this origin, sulphuric acid was historically known as oil of vitriol. Today, we represent this important acid by the formula H₂SO₄, with a relative molecular mass of 98.
Now, let us examine how sulphuric acid is manufactured industrially. The modern method is called the Contact Process, and it operates on a truly massive scale.
The first step is producing sulphur dioxide. This can be done by burning sulphur in purified air, giving sulphur dioxide directly. Alternatively, iron pyrites, which is FeS₂, can be roasted in air to produce iron three oxide and sulphur dioxide: four FeS₂ plus eleven oxygen molecules yield two Fe₂O₃ and eight sulphur dioxide molecules. Many plants now prefer burning pure sulphur because it avoids the impurities that come with mineral sources.
The second step is absolutely critical: purification of the gases. The mixture of sulphur dioxide and air must be freed from dust, moisture, and especially arsenic compounds. Arsenic oxide is particularly dangerous because it poisons the catalyst, destroying its efficiency. The gases pass through an electric precipitator to remove dust particles, then through a water scrubber for further cleaning. They are dried by spraying with concentrated sulphuric acid itself, and finally passed through an arsenic purifier containing ferric hydroxide.
The third step is the heart of the process: catalytic oxidation. The clean, dried mixture of sulphur dioxide and air passes through a contact tower packed with V₂O₅ catalyst. Pt can also be used, but it is more expensive and easily poisoned by impurities. The reaction converts sulphur dioxide to sulphur trioxide. This reaction is highly exothermic, meaning it releases substantial heat. The catalyst is initially heated to about 450 degrees Celsius, but once the reaction begins, the heat it generates maintains the temperature automatically.
Here, we must understand the conditions that favour this conversion. Since the reaction is exothermic, lower temperatures theoretically give better yields. However, the practical optimum is 410 to 450 degrees Celsius. High pressure would also favour the reaction because the product has less volume than the reactants, but building acid-resistant towers that withstand high pressure is extremely difficult. Therefore, only one to two atmospheres of pressure are used. An excess of oxygen is maintained to push the reaction forward, and the catalyst enables the reaction to proceed at a reasonable rate.
The fourth step involves absorbing the sulphur trioxide. You might think we could simply dissolve sulphur trioxide in water to make sulphuric acid, but this does not work well in practice. The reaction is extremely vigorous, producing so much heat that a mist of acid droplets forms instead of a clean liquid. Instead, the sulphur trioxide is absorbed into concentrated sulphuric acid, 98 percent strength, forming a compound called oleum or pyrosulphuric acid, H₂S₂O₇. The reaction is: sulphur trioxide plus H₂SO₄ yields H₂S₂O₇.
The final step is dilution. Calculated amounts of water are carefully added to the oleum to produce sulphuric acid of the desired concentration: H₂S₂O₇ plus water yields two molecules of H₂SO₄. This controlled approach avoids the dangerous mist formation that would occur if sulphur trioxide contacted water directly.
Now let us turn to the physical properties of sulphuric acid. Pure sulphuric acid is a colourless, odourless, dense, oily liquid. It is extremely hygroscopic, meaning it absorbs moisture from the air so aggressively that it must always be kept in tightly stoppered bottles. The pure acid has a density of 1.85 grams per cubic centimetre and boils at 338 degrees Celsius. Interestingly, pure sulphuric acid freezes at 10.4 degrees Celsius into colourless crystals. It mixes with water in all proportions, but this mixing releases tremendous heat. Pure sulphuric acid is almost a non-conductor of electricity, but when diluted, it becomes a good conductor because it ionises.
The chemical properties of sulphuric acid depend dramatically on whether it is dilute or concentrated. Let us examine each case carefully.
When dilute, sulphuric acid behaves as a typical strong acid. It ionises in water to form hydronium ions and hydrogen sulphate ions. It ionises in two stages, first forming hydronium ion and hydrogen sulphate ion, then further ionising to produce more hydronium ion and sulphate ion. The reactions are: H₂SO₄ plus water forms H₃O⁺ and HSO₄⁻; then HSO₄⁻ plus water forms more H₃O⁺ and SO₄²⁻. Note that pure sulphuric acid, without water, does not ionise and therefore shows no acidic properties. This two-stage ionisation makes sulphuric acid dibasic, meaning it can form two types of salts. With limited base below 200 degrees Celsius, it forms acid salts like sodium hydrogen sulphate or sodium bisulphate, NaHSO₄. With excess base above 200 degrees Celsius, it forms normal salts like sodium sulphate, Na₂SO₄. For example, NaCl with H₂SO₄ below 200 degrees Celsius gives NaHSO₄ and HCl. Above 200 degrees Celsius, two molecules of NaCl with H₂SO₄ give Na₂SO₄ and two molecules of HCl.
Dilute sulphuric acid reacts with active metals, those above hydrogen in the activity series, to produce metal sulphates and hydrogen gas.
Magnesium, zinc, and iron all react this way.
Noble metals like gold and platinum remain unaffected.
It neutralises bases, both metal oxides and metal hydroxides, forming salts and water.
Copper oxide, iron oxide, zinc hydroxide, and sodium hydroxide all react readily.
With carbonates and bicarbonates, dilute sulphuric acid liberates carbon dioxide gas along with salt and water.
Sodium carbonate, copper carbonate, sodium bicarbonate, and potassium bicarbonate all show this effervescence.
Metal sulphides react to produce hydrogen sulphide gas, which has a characteristic rotten egg smell.
Sodium sulphide, iron sulphide, and zinc sulphide all behave this way.
Finally, sulphites and bisulphites release sulphur dioxide gas when treated with dilute sulphuric acid. Sulphur dioxide is a poisonous gas that contributes to air pollution, adversely affecting human health and vegetation, but it also has interesting properties as both an oxidising and reducing agent. Its temporary bleaching action is worth noting: sulphur dioxide with water forms sulphuric acid and nascent hydrogen, which reduces coloured materials to colourless products that can regain their original colour when exposed to air.
Concentrated sulphuric acid is an entirely different reagent. Its behaviour is dominated by four key characteristics that we must understand thoroughly.
First, concentrated sulphuric acid is non-volatile. Its very high boiling point of 338 degrees Celsius means it does not evaporate easily. This property makes it invaluable for preparing volatile acids from their salts. When concentrated sulphuric acid reacts with sodium chloride or potassium chloride, hydrogen chloride gas is liberated. With sodium nitrate or potassium nitrate, nitric acid is produced. Even acetic acid can be prepared from sodium acetate. In each case, the volatile acid escapes while the non-volatile sulphuric acid remains behind.
Second, concentrated sulphuric acid is a powerful oxidising agent. When heated, it decomposes to release nascent oxygen. This oxygen can oxidise non-metals like carbon and sulphur themselves. Carbon is oxidised to carbon dioxide. The reaction is: carbon plus two molecules of H₂SO₄ yields carbon dioxide, two water molecules, and two sulphur dioxide molecules. Sulphur is oxidised to sulphur dioxide. The reaction is: sulphur plus two molecules of H₂SO₄ yields three sulphur dioxide molecules and two water molecules. Metals like copper and zinc, which do not react with dilute acid, are oxidised by concentrated sulphuric acid to form their sulphates, with sulphur dioxide and water as byproducts. Copper forms copper sulphate: copper plus two molecules of H₂SO₄ yields CuSO₄, two water molecules, and sulphur dioxide. Zinc forms zinc sulphate: zinc plus two molecules of H₂SO₄ yields ZnSO₄, two water molecules, and sulphur dioxide. Even hydrogen bromide is oxidised to bromine: two molecules of HBr plus H₂SO₄ yields bromine, two water molecules, and sulphur dioxide. Hydrogen sulphide is oxidised to sulphur: H₂S plus H₂SO₄ yields sulphur, two water molecules, and sulphur dioxide. These reactions demonstrate the powerful oxidising nature of hot concentrated sulphuric acid.
Third, and perhaps most dramatically, concentrated sulphuric acid is a potent dehydrating agent. It has an extraordinary affinity for water, stripping water molecules away from compounds with which they are chemically combined.
When a few drops of concentrated sulphuric acid are added to blue copper two sulphate pentahydrate crystals, CuSO₄·5H₂O, the colour fades to white as the water of crystallisation is removed, leaving anhydrous copper sulphate, CuSO₄.
More spectacular is the reaction with carbohydrates. Sugar, glucose, and cellulose in paper, cotton, and wood are all carbohydrates containing hydrogen and oxygen in the same ratio as water. Concentrated sulphuric acid removes these elements as water, leaving behind a black spongy mass of pure carbon called sugar charcoal. For glucose, C₆H₁₂O₆ yields six carbon atoms and six water molecules. For cane sugar, C₁₂H₂₂O₁₁ yields twelve carbon atoms and eleven water molecules. The carbon rises up as a black spongy mass. The reaction is highly exothermic, so the mass becomes hot and steam is evolved. This same destructive power explains why concentrated sulphuric acid causes terrible burns, charring skin black by dehydration.
Organic acids like formic acid and oxalic acid are dehydrated by concentrated sulphuric acid. Formic acid, HCOOH, yields carbon monoxide and water. Oxalic acid, (COOH)₂ or C₂H₂O₄, yields carbon monoxide, carbon dioxide, and water. Ethanol, C₂H₅OH, can be dehydrated to ethylene, C₂H₄, at 170 degrees Celsius.
Fourth, concentrated sulphuric acid precipitates insoluble sulphates from solution. Pb(NO₃)₂ forms insoluble lead sulphate, BaCl₂ forms insoluble barium sulphate, and calcium salts form insoluble calcium sulphate. These insoluble sulphates are important in qualitative analysis. Barium sulphate and lead sulphate are particularly important as they are white precipitates used in qualitative analysis.
Before we conclude, let us consider how to distinguish between dilute and concentrated sulphuric acid. Dilute acid ionises extensively, making it a strong acid and good conductor. Concentrated acid ionises poorly, behaving as a weak acid and weak electrolyte. Only concentrated acid can act as an oxidising agent, dehydrating agent, or non-volatile acid. Dilute acid cannot dry gases or dehydrate substances.
One crucial safety note: when diluting sulphuric acid, never add water to the acid. The heat released is so intense that the water instantly turns to steam, causing the acid to spurt violently. Always add acid slowly to water with continuous stirring, allowing the heat to dissipate safely.
Let us recap the essential points of this lesson.
First, sulphuric acid is manufactured by the Contact Process. This involves production of sulphur dioxide, thorough purification including removal of arsenic oxide which poisons the catalyst, catalytic oxidation to sulphur trioxide using V₂O₅ or Pt at 410 to 450 degrees Celsius and one to two atmospheres pressure, absorption into oleum, and careful dilution.
Second, dilute sulphuric acid behaves as a typical strong dibasic acid, reacting with metals, bases, carbonates, sulphides, and sulphites to produce characteristic products.
Third, concentrated sulphuric acid shows four distinctive properties: it is non-volatile, a powerful oxidising agent, a strong dehydrating agent, and capable of precipitating insoluble sulphates.
Fourth, the dehydration of sugar to sugar charcoal and the conversion of blue copper two sulphate pentahydrate crystals to white anhydrous copper sulphate demonstrate the remarkable affinity of concentrated sulphuric acid for water.
Fifth, the non-volatile nature enables preparation of volatile acids like hydrochloric acid from NaCl or KCl, and nitric acid from NaNO₃ or KNO₃.
Finally, always remember the safety rule: add acid to water, never water to acid.
Sulphuric acid truly deserves its title as the King of Chemicals. From fertilisers to dyes, from explosives to pharmaceuticals, from metal processing to petroleum refining, its applications are virtually limitless. Understanding its dual nature, as both a simple acid when dilute and a powerful reactive agent when concentrated, gives you insight into one of chemistry's most important substances. Keep exploring, keep questioning, and you will continue to uncover the remarkable patterns that govern our material world. Until next time, stay curious and stay safe in your chemical journey.