9.5.C - Sulfuric Acid


Sulfuric acid is a widely used and important acid in industry used to make paper, pickle iron and steel, manufacture other important industrial chemicals and fertilisers.

Reactions of sulfuric acid include: as an oxidant, in neutralisation reactions and as a dehydrating agent. The dissolutions of sulfuric acid is exothermic so acid must always be added to water. Sulfuric acid is a strong oxidant so it must be handled with care. Protective laboratory clothing and eye protection must always be worn.

Sulfuric acid is most commonly manufactured from naturally occurring sulfur. This sulfur is extracted from underground beds using the Frasch process. The sulfur is oxidised to sulfur trioxide in the Contact process and the sulfur trioxide is converted to oleum and then sulfuric acid in an absorption tower. The yield of acid is about 98% v/v and the final concentration is usually about 18 mol L-1.

In this topic students:

  • Describe the properties, uses and some reactions of sulfuric acid
  • Outline procedures for the safe handling and transport of sulfuric acid
  • Describe the production process for sulfuric acid and explain any equilibrium considerations


Properties, Uses and Reactions

Properties

Sulfuric acid (H2SO4), is a strong mineral acid. It is soluble in water at all concentrations. It was once known as Zayt al-Zaj, or oil of vitriol, coined by the 8th-century Arab Alchemist Jabir ibn Hayyan, the chemical's probable discoverer. Sulfuric acid has many applications, and is produced in greater amounts than any other chemical besides water. World production in 2001 was 165 million tonnes, with an approximate value of $US 8 billion. Principal uses include ore processing, fertilizer manufacturing, oil refining, wastewater processing and chemical synthesis.

Sulfuric acid is a corrosive, oily, colourless liquid. It melts at 10.36°C, boils at 340°C and is soluble in all proportions in water. Sulfuric acid is a strong acid that in aqueous solution is largely changed to hydrogen ions (H+) and sulfate ions (SO42.). Each molecule gives two H+ ions, thus sulfuric acid is diprotic. Dilute solutions of sulfuric acid show all the behavioural characteristics of acids. They taste sour, conduct electricity, neutralize alkalis, and corrode active metals with formation of hydrogen gas. From sulfuric acid one can prepare both normal salts containing the sulfate group, SO4, and acid salts containing the hydrogen sulfate group, HSO4.

Concentrated sulfuric acid, formerly called oil of vitriol, is a valuable dehydrating agent. It acts so vigorously in this respect that it removes water from, and therefore chars, wood, cotton, sugar, and paper. It is used in the manufacture of ether, nitroglycerine, and dyes for its property as a dehydrating agent. When concentrated sulfuric acid is heated, it behaves also as an oxidizing agent (oxidant), capable, for example, of dissolving such relatively unreactive metals as copper, mercury, and lead to produce metal sulfate, sulfur dioxide, and water.

During the 19th century, the German chemist Baron Justus von Liebig discovered that sulfuric acid, when added to the soil, increased the amount of soil phosphorus available to plants. This discovery gave rise to an increase in the commercial production of sulfuric acid and led to improved methods of manufacture.

Uses

Sulfuric acid is a very important commodity chemical, and indeed a nation's sulfuric acid production is a good indicator of its industrial strength. The major use (60% of total worldwide) for sulfuric acid is in the "wet method" for the production of phosphoric acid, used for manufacture of phosphate fertilizers as well as trisodium phosphate for detergents. In this method phosphate rock is used, and more than 100 million tonnes is processed annually. This raw material is shown below as fluorapatite, though the exact composition may vary. This is treated with 93% sulfuric acid to produce calcium sulfate, hydrogen fluoride (HF) and phosphoric acid. The HF is removed as hydrofluoric acid. The overall process can be represented as:

Ca5F(PO4)3 + 5H2SO4 + 10H2O → 5CaSO4.2H2O + HF + 3H3PO4

Sulfuric acid is used in large quantities in iron and steel making principally as pickling-acid used to remove oxidation, rust and scale from rolled sheet and billets prior to sale into the automobile and white-goods business. The used acid is often re-cycled using an acid recovery plant in which the acid is boiled away from the dissolved iron salts, often using a submerged hydrogen-oxygen flame as the heat source.

Ammonium sulfate, an important nitrogen fertilizer is most commonly produced as a by-product from coking plants supplying the iron and steel making plants. Reacting the ammonia produced in the thermal decomposition of coal with waste sulfuric acid allows the ammonia to be crystallized out as a salt (often brown because of iron contamination) and sold into the agro-chemicals industry. The equation for such a reaction is:

NH3 + H2SO4(NH3)2SO4

Another important use for sulfuric acid is for the manufacture of aluminium sulfate, also known as papermaker's alum. This can react with small amounts of soap on paper pulp fibres to give gelatinous aluminium carboxylates, which help to coagulate the pulp fibres into a hard paper surface. It is also used for making aluminium hydroxide, which is used at water treatment plants to filter out impurities, as well as to improve the taste of the water. Aluminium sulfate is made by reacting bauxite with sulfuric acid:

Al2O3 + 3H2SO4 → Al2(SO4)3 + 3H2O

Sulfuric acid is used for a variety of other purposes in the chemical industry. For example, it is the usual acid catalyst for the conversion of cyclohexanoneoxime to caprolactam, used for making nylon. It is used for making hydrochloric acid from salt via the Mannheim process. Much H2SO4 is used in petroleum refining, for example as a catalyst for the reaction of isobutane with isobutylene to give isooctane, a compound that raises the octane rating of gasoline (petrol). Sulfuric acid is also important in the manufacture of dyestuffs.

A mixture of sulfuric acid and water is used as the electrolyte in various types of lead-acid battery where it undergoes a reversible reaction where lead and lead dioxide are converted to lead (II) sulfate. Sulfuric acid is also the principal ingredient in some drain cleaners, used to clear blockages consisting of paper, rags, and other materials not easily dissolved by caustic solutions.

Reactions

Dehydrating Agent

Because the hydration of sulfuric acid is thermodynamically favorable (ΔH = -880 kJ/mol), sulfuric acid is an excellent dehydrating agent, and is used to prepare many dried fruits. The affinity of sulfuric acid for water is sufficiently strong that it will take hydrogen and oxygen atoms out of other compounds; for example, mixing starch (C6H12O6)n and concentrated sulfuric acid will give elemental carbon and water which is absorbed by the sulfuric acid (which becomes slightly diluted):

(C6H12O6)n → 6C + 6H2O

The effect of this can be seen when concentrated sulfuric acid is spilled on paper; the starch reacts to give a burned appearance, the carbon appears as soot would in a fire. A more dramatic illustration occurs when sulfuric acid is added to a tablespoon of white sugar in a cup when a tall rigid column of black porous carbon smelling strongly of caramel emerges from the cup. Concentrated sulfuric acid can also be used to dehydrate ethanol to produce ethene.

C2H5OH (g) → C2H4 (g) + H2O (g)

Reaction with Water

The hydration reaction of sulfuric acid is highly exothermic. If water is added to concentrated sulfuric acid, it can boil and spit dangerously. One should always add the acid to the water rather than the water to the acid. This can be remembered through mnemonics such as "Do as you oughta: add acid to water", "A.A.: Add Acid", or "Drop acid, not water." Note that part of this problem is due to the relative densities of the two liquids. Water is less dense than sulfuric acid and will tend to float above the acid. The reaction is best thought of as forming hydronium ions, by:

H2SO4 + H2O → H3O+ + HSO4-

and then:

HSO4- + H2O → H3O+ + SO42-

Acid-Base Reactions

As an acid, sulfuric acid reacts with most bases to give the corresponding sulfate. For example, copper(II) sulfate, the familiar blue salt of copper used for electroplating and as a fungicide, is prepared by the reaction of copper(II) oxide with sulfuric acid:

CuO + H2SO4 → CuSO4 + H2O

Sulfuric acid can be used to displace weaker acids from their salts, for example sodium acetate gives acetic acid:

H2SO4 + CH3COONa → NaHSO4 + CH3COOH

Oxidising Agent

Sulfuric acid reacts with most metals in a single displacement reaction to produce hydrogen gas and the metal sulfate. Dilute H2SO4 attacks iron, aluminium, zinc, manganese and nickel, but tin and copper require hot concentrated acid. Lead and tungsten are, however, resistant to sulfuric acid. The reaction with iron (shown) is typical for most of these metals, but the reaction with tin is unusual in that it produces sulfur dioxide rather than hydrogen.

Fe (s) + H2SO4 (aq) → H2 (g) + FeSO4 (aq)


Sn (s) + 2H2SO4 (aq) → SnSO4 (aq) + 2H2O (l) + SO2 (g)

In both reactions the metal is oxidised. In the first reaction, the hydrogen is reduced. In the second, the sulfur is reduced.



A video from the University of Nottingham of the properties and reactions of sulfuric acid.

 

Safe Handling and Transport

Laboratory Hazards

The corrosive properties of sulfuric acid are accentuated by its highly exothermic reaction with water. Hence burns from sulfuric acid are potentially more serious than those of comparable strong acids (e.g. hydrochloric acid), as there is additional tissue damage due to dehydration and particularly due to the heat liberated by the reaction with water, i.e. secondary thermal damage. The danger is obviously greater with more concentrated preparations of sulfuric acid, but it should be remembered that even the normal laboratory "dilute" grade (approx. 1 mol L-1, 10%) will char paper by dehydration if left in contact for a sufficient length of time. The standard first aid treatment for acid spills on the skin is, as for other corrosive agents, irrigation with large quantities of water: in the case of sulfuric acid it is important that the acid should be removed before washing, as a further heat burn could result from the exothermic dilution of the acid. Washing should be continued for a sufficient length of time—at least ten to fifteen minutes—in order to cool the tissue surrounding the acid burn and to prevent secondary damage. Contaminated clothing must be removed immediately and the underlying skin washed thoroughly.

Preparation of the diluted acid can also be dangerous due to the heat released in the dilution process. It is essential that the concentrated acid is added to water and not the other way round, to take advantage of the relatively high heat capacity of water. Addition of water to concentrated sulfuric acid leads at best to the dispersal of a sulfuric acid aerosol, at worst to an explosion. Preparation of solutions greater than 6 mol L-1 (35%) in concentration is the most dangerous, as the heat produced can be sufficient to boil the diluted acid. Efficient mechanical stirring and external cooling (e.g. an ice bath) are essential.

Industrial Hazards

Although sulfuric acid is non-flammable, contact with metals in the event of a spillage can lead to the liberation of hydrogen gas. The dispersal of acid aerosols and gaseous sulfur dioxide is an additional hazard of fires involving sulfuric acid. Water should not be used as the extinguishing agent because of the risk of further dispersal of aerosols: carbon dioxide is preferred where possible.

Sulfuric acid is not considered toxic besides its obvious corrosive hazard, and the main occupational risks are skin contact leading to burns (see above) and the inhalation of aerosols. Exposure to aerosols at high concentrations leads to immediate and severe irritation of the eyes, respiratory tract and mucous membranes: this ceases rapidly after exposure, although there is a risk of subsequent pulmonary edema if tissue damage has been more severe. At lower concentrations, the most commonly reported symptom of chronic exposure to sulfuric acid aerosols is erosion of the teeth, found in virtually all studies: indications of possible chronic damage to the respiratory tract are inconclusive as of 1997. In the United States, the permissible exposure limit (PEL) for sulfuric acid is fixed at 1 mg/m3: limits in other countries are similar.

Transport

Concentrated sulfuric acid can be transported and stored in iron or steel containers, as it is molecular and does not readily react with iron. The steel tanks provide strength to prevent leakage into the environment in the event of an accident. Dilute sulfuric acid is stored in glass containers. The lid must be kept tightly sealed, as sulfuric acid absorbs water from the atmosphere.

 

Industrial Production

Two processes for the production of sulfuric acid are in use today. In their initial steps, both require the use of sulfur dioxide, which is produced by burning iron pyrites, FeS2, or sulfur, in air. The first of these methods, the lead-chamber process, employs as reaction vessels large lead-sheathed brick towers. In these towers, sulfur-dioxide gas, air, steam, and oxides of nitrogen react to yield sulfuric acid as fine droplets that fall to the bottom of the chamber. Almost all the nitrogen oxides are recovered from the outflowing gas and are brought back to the chamber to be used again. Sulfuric acid produced in this way, and labelled acid, is only about 62 to 70 percent H2SO4, with the rest being water. The lead-chamber process now makes about 20 percent of all sulfuric acid, but that percentage is diminishing.

The second and more common method of manufacturing sulfuric acid, the contact process, came into commercial use about 1900. This process uses elemental sulfur (obtained from the Frasch process) and depends on oxidation of sulfur dioxide to sulfur trioxide, SO3, under the accelerating influence of a catalyst. Finely divided platinum, the most effective catalyst, has two disadvantages: It is very expensive, and it is vitiated by certain impurities in ordinary sulfur dioxide that reduce its activity. Many sulfuric-acid producers use two catalysts in tandem; first, a more rugged but less effective one like iron oxide or vanadium oxide to bring about the bulk reaction; then, a smaller amount of platinum to finish the job. At 400°C, the conversion of sulfur dioxide to trioxide is nearly complete. The trioxide is dissolved in concentrated sulfuric acid, and at the same time a regulated influx of water maintains the concentration at a selected level usually about 95 percent. By reducing the flow of water, a product with more SO3 than shown in the formula H2SO4 may be made. This product, called fuming sulfuric acid, or oleum, or Nordhausen acid, is needed in some organic chemical reactions.



A video showing the production of sulfuric acid throught the Contact process in the UK.
This process is described below.

 

1. Frasch Process

The Frasch process showing how elemental sulfur is extracted from ground-based
deposits.

Elemental sulfur deposits are found in volcanic or sedimentary areas of Italy, USA, Russia and the Ukraine. These deposits have been extracted from the ground using the Frasch process. Holes are drilled down through the overlying rock into the sulfur deposits. A special series of pipes are then inserted into the drill hole. The pipes are arranged in a circular pattern. The outer pipes will contain superheated steam (usually about 160°C) that is pumped down into the deposit. Superheated steam is steam th at is much hotter than the boiling point of water.

Since the melting point of sulfur is so low, the sulfur is readily melted. As the sulfur becomes molten, it is removed by pumping air down the central pipe which forces the liquid sulfur to the surface under pressure. When the molten sulfur reaches the surface, it is pumped onto wooden blocks where the sulfur again solidifies. The Frasch process is able to produce sulfur of very high purity, often above 99%.

2. Contact Process

In the first part of this process, sulfur obtained from the Frasch process is burned in an exothermic reaction to produce sulfur dioxide.

S (s) + O2 (g) → SO2 (g)

This is then oxidised to sulfur trioxide using oxygen in another exothermic reaction in the presence of a vanadium(V) oxide catalyst.

2SO2 (g) + O2 (g) → 2SO3(g)

The second step is typically the most difficult step in the process and the step where equilibrium considerations are most important. Yield is increased to 99% by using:

  • Low temperature - This is an exothermic reaction so a compromise temperature of 400°C is used.

  • High pressure - Only a pressure of about 2 atm is used because a satisfactory yield is obtained from the lwer temperature. Creating higher pressures would increase the yield but not enough to pay for the cost of the increased pressure.

  • Excess oxygen - The typical input mixture: 10 kPa SO2, 10 kPa O2 and 80 kPa N2. The nitrogen acts as an inert gas to increase the pressure.

During this step the gaseous mixture passes into the first catalyst bed at 400°C and the heat liberated from the exothermic reaction raises the temperature to 550°C. The rate of reaction is high at this temperature and about 70% conversion is produced. High yields are ensured in the next two cycles where the catalyst beds are held at 400°C. The recycling of the reaction mixture over successive catalyst beds produces an overall yield for this step of about 98%.

3. Conversion to Oleum and Sulfuric Acid

An absorption tower in a sulfuric acid
manufacturing plant.

In this step, the sulfur trioxide is treated with sulfuric acid (usually as 97-98% H2SO4 containing 2-3% water) to produce a substance called oleum. In an absorption tower (right) the SO3 is absorbed into H2SO4 to produce oleum (H2S2O7).

H2SO4 (l) + SO3 (g) → H2S2O7 (l)

Note that directly dissolving SO3 in water in this step, as shown in the reaction below, is impractical due to the highly exothermic nature of the reaction. Mists are formed instead of a liquid and this makes is difficult to extract the sulfuric acid.

SO3 (g) + H2O (l) → H2SO4 (l)


4. Conversion to Sulfuric Acid

In the final step, oleum is reacted with water to form concentrated H2SO4.

H2S2O7 (l) + H2O (l) → 2H2SO4 (l)


The yield of acid at the completion of the process is about 98% v/v and the final concentration is usually about 18 mol L-1.



A flowchart summarising the sequences of reactions involved in producing sulfuric acid on a large scale.