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9.5 Industrial Chemistry: 3. Sulfuric acid

Syllabus reference (October 2002 version)
3. Sulfuric acid is one of the most important industrial chemicals	Students learn to: outline three uses of sulfuric acid in industry describe the processes used to extract sulfur from mineral deposits, identifying the properties of sulfur which allow its extraction and analysing potential environmental issues associated with its extraction outline the steps and conditions necessary for the industrial production of H2SO4 from its raw materials describe the reaction conditions for the production of SO2 and SO3 apply the relationship between rates of reaction and equilibrium conditions to the production of SO2 and SO3 describe, using examples, the reactions of sulfuric acid acting as: an oxidising agent a dehydrating agent describe and explain the exothermic nature of sulfuric acid ionisation identify and describe safety precautions that must be taken when using and diluting concentrated sulfuric acid	Students: gather, process and present information from secondary sources to describe the steps and chemistry involved in the industrial production of H2SO4 and use available evidence to analyse the process to predict ways in which the output of sulfuric acid can be maximised perform first-hand investigations to observe the reactions of sulfuric acid acting as: an oxidising agent a dehydrating agent use available evidence to relate the properties of sulfuric acid to safety precautions necessary for its transport and storage

Extract from Chemistry Stage 6 Syllabus (Amended October 2002).© Board of Studies, NSW.
[Edit: 9 Jul 2009]

Prior Learning: Preliminary module 8.3.5, 8.4.4, HSC module 9.3.2

outline three uses of sulfuric acid in industry

Uses of sulfuric acid include:

• Manufacture of ammonium sulfate fertiliser and phosphate fertiliser
  
  Sulfuric acid removes ammonia from the mixture of gases produced in a coke oven.
  
  2NH₃(g)   +        H₂SO₄(aq)      (NH₄)₂SO₄(aq)
  
  Sulfuric acid converts insoluble calcium phosphate, in phosphate rock, to mixtures that are soluble in water and therefore available for plants.  The mixtures are crushed and used as superphosphate fertilisers.
  
  
• Dehydrating agent
  
  Sulfuric acid is used to dry the chlorine gas produced by the electrolysis of sodium chloride solution. It is also used as a drying agent in the manufacture of explosives, dyes and detergents and it brings about condensation reactions in the production of polymers, and esters.
  
  
• Cleaning iron and steel
  
  To galvanise or plate iron or steel any oxide that has formed on its surface and any grease or dirt must first be removed. This is done by treating with acid such as sulfuric acid.

describe the processes used to extract sulfur from mineral deposits, identifying the properties of sulfur which allow its extraction and analysing potential environmental issues associated with its extraction

Extraction of Sulfur

1. Some sulfur is recovered from underground deposits of the element by the Frasch Process.
   • Three concentric pipes are drilled down into the sulfur deposit.
   • Superheated water (180°C, under pressure) is pumped down the outside pipe. This melts the sulfur (melting point 113°C).
   • Air under pressure is pumped down the inside pipe, pushing the molten sulfur and steam up the middle pipe to the surface.

   

2. Sulfur is also obtained from hydrogen sulfide in natural gas and petroleum. Incomplete combustion of H₂S in a furnace produces SO₂ and S.

   3H₂S(g)   +   O₂(g)      2H₂S(g)   +   3S(g)   +   SO₂(g)  

   This mixture is cooled to condense the sulfur. The gases are then passed over a heated catalyst.

    2H₂S(g)   +   SO₂(g)      2H₂O(g)   +   3S(g)  

   Cooling condenses the remaining sulfur (boiling point 445°C).

3. Sulfur is also released as sulfur dioxide when metal sulfide ores are smelted. A general equation for this reaction, using M to represent a metal (such as copper, zinc or iron), can be written as:

   MS   +   O₂(g)      M(s)   +   SO₂(g)

   The metal M often forms metal oxide MO.

Environmental issues

• Release of sulfur dioxide from metal sulfide smelters.  Most SO₂ gas emitted is used to make sulfuric acid.
• SO₂ emission into the atmosphere by industry is now strictly controlled by government regulations.

  The National Pollutant Inventory , Department of Environment and Heritage, Canberra database monitors industrial sources of SO₂ in Australia, see:

• SO₂ contributes to the formation of acid rain, as you learnt in 9.3.2.

  For more information about acid rain see US Environmental Protection Agency .

  Read about acid rain in Australia and then see and hear a video clip from the CSIRO on acid rain in Australia, CSIRO Publishing 2007

gather, process and present information from secondary sources to describe the steps and chemistry involved in the industrial production of H₂SO₄ and use available evidence to analyse the process to predict ways in which the output of sulfuric acid can be maximised

• Gather information on industrial production of sulfuric acid from text books and the Internet.
• Process the information from these sources by organising it according to the points listed:
  • describe the steps in the production of H₂SO₄
  • write chemical equations to show the reactions involved in each step of the production
  • describe the conditions needed for each step to take place.

  For each step, analyse the information by considering how you could alter conditions to get the maximum yield. Think about and predict:

  • how to increase rates of reactions
  • factors that affect equilibrium reactions.

  You may have other information you want to include as well as the points above.

• Present the information you have gathered in a way you have chosen. It might be in a table or in text format. You could use overhead transparencies if you are giving an oral presentation.

outline the steps and conditions necessary for the industrial production of H₂SO₄ from its raw materials

• The industrial production of sulfuric acid involves the following steps, carried out under the conditions described so as to optimise yield:
  1. Combustion of sulfur (or metal sulfide ore e.g. PbS) to form sulfur dioxide.

     S(s)   +   O₂(g)      SO₂(g)

     2PbS(s)   +   3O₂ (g)      PbO(s)   +   2SO₂(g)

  2. Sulfur dioxide is passed over vanadium pentoxide or platinum catalyst, at 450°C to produce sulfur trioxide. This process takes place in a catalytic converter.

     

  3. Sulfur trioxide is dissolved in water forming sulfuric acid. This process takes place in absorption towers.

     SO₃(g)  +  H₂O(l)      H₂SO₄(aq)  

     Gaseous sulfur trioxide cannot be added directly to water as the reaction is very exothermic and would cause the acid to vaporise and form a dangerous mist.  To avoid this problem, the sulfur trioxide is added to a flowing solution of concentrated sulfuric acid rather than to pure water. Water is added in small amounts, with stirring, to react with the SO₃ and form concentrated H₂SO₄ of the desired concentration.

describe the reaction conditions for the production of SO₂ and SO₃

1. The combustion of sulfur or metal sulfides takes place in a combustion furnace. This is a rapid exothermic reaction and goes to completion. Conditions are those that favour a high rate of reaction and high yield e.g. high temperature, high surface area (crushed rock).
   • Liquid sulfur obtained by the Frasch process is heated in dry air and sent to the catalytic converter.

     

   • SO₂ obtained from metal ores must be separated out, cleaned and dried before it can go to the catalytic converter to produce SO_3.

     

2. The production of SO₃ from SO₂ takes place in a catalytic converter. It is an equilibrium reaction and  involves a compromise between reaction rate, equilibrium yield and economic factors.
   • At room temperature, the yield would be very high, but the reaction would occur at an uneconomically slow rate. Increasing the temperature increases the rate of reaction, however, the forward reaction is exothermic, so increasing the temperature pushes the equilibrium to the left to absorb the heat, thus decreasing the yield. A high temperature could also damage the catalyst, making it less efficient.

     450 - 600°C allows a fairly fast reaction rate plus good yield.

   • A catalyst, vanadium pentoxide, is used to increase the reaction rate. This reaction is called the Contact Process because sulfur dioxide and oxygen molecules react in contact with the surface of the catalyst, which is arranged in layers in towers.
   • Increasing pressure pushes the equilibrium to the right (fewer particles), but the equipment required is expensive, so a low pressure of only 1-2 atmospheres is used. This pressure is sufficient to move gases through the catalyst chamber.
   • Excess oxygen is also used to push the equilibrium to the right and increase yield. The stoichiometric mole ratio for the reaction shows the O₂:SO₂ ratio needed is 1:2. In the industrial process, twice as much oxygen is used, the O₂:SO₂ ratio used is 1:1.

These conditions produce a yield of about 99% sulfur trioxide.

The energy released from these exothermic reactions is used in the plant for melting the sulfur or producing steam to generate electricity.

apply the relationship between rates of reaction and equilibrium conditions to the production of SO₂ and SO₃

• The reaction to produce SO₂ goes to completion.

  However, the reaction to produce SO₃ is an equilibrium reaction and  involves a compromise between reaction rate, equilibrium yield and economic factors. You should be able to show how conditions will affect the rate of reaction and the equilibrium yield and the necessity for compromise in order to obtain an economic yield. Try using a table like the one below to summarise these ideas.

  Conditions	Increase rate of reaction	Increase equilibrium yield	Economic factors	Conditions used
  temperature	high temperature
  pressure			high pressure production is more expensive & dangerous
  concentration of reactants
  other		removal of product

perform first-hand investigations to observe the reactions of sulfuric acid acting as:

• an oxidising agent
• a dehydrating agent

• Sulfuric acid as an oxidising agent

  To perform the investigation

  1. Place a small piece of magnesium, zinc and iron (e.g. a nail) in separate beakers or test tubes.
  2. Add 3 drops of the sulfuric acid to each metal and observe.
  3. Record all observations in a table.
  4. Write equations to show any reactions observed.
• Sulfuric acid as a dehydrating agent

  To perform the investigation

  1. Place about 10 sugar crystals on a piece of flat, dry wood. (A paddle pop stick will do.)
  2. Add 1 drop of sulfuric acid to the sugar and another to the wood.
  3. Observe and record any reactions.
  4. Write an equation for the reaction.

     C₁₂H₂₂O₁₁      12C   +   11H₂O

describe, using examples, the reactions of sulfuric acid acting as:

• an oxidising agent
• a dehydrating agent

• Acts as a moderately strong oxidising agent

  Hot, concentrated sulfuric acid oxidises:

  -bromide and iodide ions to the elements bromine and iodine, while the sulfuric acid is reduced to sulfur dioxide. Sulfur in sulfuric acid has an oxidation number of +6.

  2I^–(aq)  +   3H₂SO₄(aq)    I₂(aq)   +   SO₂(aq)   +   2H₂O(l)   +  2HSO₄^–(aq)

  -unreactive metals such as copper, mercury and lead to produce the metal sulfate, sulfur dioxide and water.

  Cu(s)  +  2H₂SO₄(aq)    Cu ²⁺(aq)  +  SO₄^2–(aq)  +  2H₂O(l)  +  SO₂(aq)

• Acts as a dehydrating agent

  Removes water from carbohydrates and other organic substances - leaving black carbon. It will char wood, cotton, sugar or paper.

  C₁₂H₂₂O₁₁   +   11H₂SO₄ (aq)      12C   +   11H₂SO₄.H₂O
  sucrose

  Removes water from alkanols e.g. converts ethanol to ethene.

  2CH₃-CH₂-OH(l)  +  H₂SO₄(aq)    2CH₂=CH₂(g)  +  2H₃O⁺(aq)  +  SO₄^2–(aq)

  Dries gases that do not react with it e.g. O₂, N₂, Cl₂ and CO₂.

  Helps produce esters by removing a molecule of water when an alkanol reacts with an alkanoic acid (esterification).

  

describe  and explain the exothermic nature of sulfuric acid ionisation

• The ionisation of sulfuric acid is exothermic, releasing lots of heat.

  H₂SO₄(aq)    H⁺(aq)    +      HSO₄^–(aq)   +   heat

• Sulfuric acid dissociates in two steps.

  Notice that sulfuric acid is a strong acid in its first dissociation, but the HSO₄^– ion is a weak acid and only dissociates slightly.

  H₂SO₄(aq)    H⁺(aq)    +      H SO₄^–(aq)           K is very large

  HSO₄^– (aq)   H⁺(aq)      +   SO₄^2–(aq)           K= 1.2 x 10^-2

identify and describe safety precautions that must be taken when using and diluting concentrated sulfuric acid

• Wear protective clothing, including safety glasses.
• ALWAYS ADD ACID TO WATER.  Add a small amount at a time, constantly stirring. This will produce a dilute solution, releasing only a small amount of heat and any splatters that do occur are more likely to be water or dilute acid rather than concentrated acid. If acid is spilt, wash with lots of water and use sodium bicarbonate to neutralise the acid.
• If water is accidentally added to concentrated sulfuric acid, the heat released will make the water boil violently, often splattering droplets of concentrated acid out of the container. There may be enough heat generated to crack the container.

use available evidence to relate  the properties of sulfuric acid to safety precautions necessary for its transport and storage

• Evidence can be obtained from the dot point above or from chemistry practical manuals or the Internet. Concentrated sulfuric acid can be transported and stored in iron containers, as it is molecular and does not readily react with iron.
• Dilute sulfuric acid is stored in glass containers. The lid must be kept tightly sealed, as sulfuric acid absorbs water from the atmosphere.