Contact Process

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CONTACT PROCESS Sulfuric acid is one of the most important industrial chemicals  Outline three uses of sulfuric acid in industry

1. The major use of sulfuric acid in Australia is in the manufacture of fertilizers such as ammonium sulfate and superphosphate. Superphosphate is produced by reacting sulfuric acid with rock phosphate. Ammonium sulfate is produced by neutralising ammonia with sulfuric acid. 2. Production of titanium (IV) oxide from titanium minerals eg ilmenite. Titanium is an important lightweight metal used to produce strong alloys and white, opaque pigments. H2SO4 is used to leach the titanium from the minerals after mining. 3. Cleaning iron – because very corrosive used to remove the oxide layer from iron or steel before they are galvanised or electroplated.  Describe the processes used to extract sulfur from mineral deposits, identifying the properties of sulfur which allow its extraction and analyzing potential environmental issues that may be associated with its extraction Most sulphur is extracted from mineral deposits using the Frasch process. Superheated stream is pumped down the outer of 3 concentric pipes into the sulphur deposit, and since sulphur has a low melting point (119) it is readily melted. At the same time, compressed air is blown down the inner pipe, and because sulphur has a relatively low density, the air is able to force the molten sulphur up the middle pipe to the surface where it resolidifies. The insolubility of sulphur in water means that it separates from any water, leaving 99.5% pure sulphur. Sulfur is also obtained from hydrogen sulphide in natural gas and petroleum. Incomplete combustion of H2S in a furnace produces SO2 and S. 3H2S(g) + O2(g)  H2S(g) + 3S(g) + SO2(g) The mixture is cooled to condense the sulphur. Sulfur is also released as sulphur dioxide when metal sulphide ores are smelted. Eg. ZnS(s) + O2(g)  Zn(s) + SO2(g)




Environmental Issues: - Sulfur is easily oxidised to sulphur dioxide or reduced to hydrogen sulfide, both of which are serious air pollutants at quite low concentrations. Care is needed to ensure that there is no inadvertent oxidation or reduction of sulphur - It is very difficult to back-fill the underground caverns left by extraction of sulphur  Outline the steps and conditions necessary for the industrial production of H2SO4 from its raw materials

Today most H2SO4 is manufactured by the Contact Process. Step 1 – molten sulphur (or sulphide ore eg pyrite) is combusted to form SO2. S(l) + O2(g)  SO2(g) or 4FeS2(s) + 11O2(g)  4Fe2O3(s) + 8SO2(g) Step 2 – So2 gas is transferred to a catalytic converter where it is oxidised to SO3. 2SO2(g) + O2(g) 2SO3(g) + heat

Conditions necessary include a pressure of 1-2 atmospheres, a small excess of O2, a catalyst of vandium (V) oxide, and temps of 400-500oC. Step 3 – SO3 is dissolved in conc H2SO4 to form oleum, H2S2O7. Water is then added to the oleum to produce H2SO4. H2S2O7(l) + H2O(l)  2H2SO4

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Describe the reaction conditions necessary for the production of SO2 and SO3 The production of SO2 is carried out in a combustion furnace. The exothermic reaction occurs quickly and goes to completion. An excess of dry air is used so that the SO2 produced is already mixed with O2 for the next step. To produce SO3 the reaction conditions necessary include a pressure of 1-2 atmosphere, as small excess of O2 to increase the yield, temperatures of 400500oC and a catalyst of V2O5. The catalyst is needed to increase the reaction rate at moderate temperatures without decreasing the yield. The conditions used are a compromise between reaction rate and equilibrium yield to produce as much SO3 as possible. 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. Temperature A high yield of SO3 could be achieved at low temps but the rate of the reaction would be very slow. A faster...
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