| Contents for this page | Related topics |  |
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Chemical Systems Le Chatelier's principle Applications in industry Additional questions |
Reaction rates Catalysis Chemical equilibrium |
 Data Glossary |
| Learning Outcomes | ||
| After studying this section, you will (a) know what is meant by a closed and open chemical system, (b) understand Le Chatelier's principle and be able to apply it to chemical reactions at equilibrium. | ||
A chemical system is an assembly of objects that are separated form their surroundings by a boundary. For example, a solution contained in a flask may constitute a system. We can distinguish between three types of systems:
| Le Chatelier's Principle |
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| "If a system in equilibrium is subjected to a stress the equilibrium will shift in the direction which tends to relieve that stress." |
This generalisation enables us to predict the way reversible chemical reactions will be affected if a change in conditions is applied at equilibrium.
Consider the formation of products, P, from reactants R, according to the exothermic reaction R  P, ΔH < 0, in a closed system. Increasing the temperature of the reaction will increase the rate at which equilibrium is attained, which is therefore obtained after a shorter time, and this will shift the equilibrium position in favour of the reactants. This is because increasing the temperature, that is, increasing the heat in the system, is a "stress" imposed on the equilibrium. Le Chatelier′s principle tells us that the system will shift its equilibrium position in order to remove some of that heat. Since heat is given out in the forward reaction, the reverse reaction is favoured, and there will be more reactants and less products at equilibrium than were present at a lower temperature (diagram on the right). (Click here to see an animation showing the changes in concentration at a higher temperature.)(Start the animation.) |
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| If the reaction is carried out at a lower temperature, then, since ΔH < 0, the forward reaction is favoured, and while the time taken to attain equilibrium is longer, the proportion of products over reactants will be higher once that equilibrium is reached. (Start the animation). | |
For endothermic reactions, the situation is reversed: higher temperatures shift the equilibrium in the direction of product formation, while lower temperatures will reduce the proportion of products at equilibrium.
| For a given temperature, a catalyst will speed up the attainment of the equilibrium, but will not alter the proportions of reactants and products that are fixed by the value of Kc. (Start the animation) | |
| Suppose that when the reaction mixture had reached equilibrium, some more reactants are added? This is a stress on the system, which can be relieved by using up these reactants to form products, in order to form a new equilibrium, which again must conform to the value of Kc at that temperature. (Start the animation) |  |
Several important industrial processes rely on the application of Le Chatelier's principle. Here are just three of them:
Ammonia is produced industrially by the Haber-Bosch process which involves the catalytic reduction of nitrogen by hydrogen at temperatures of 450-500 ºC and pressures of 35-40 MPa.
The process provides an excellent illustration of Le Chatelier's Principle, and is a favourite of examiners! If we examine the above equation, we see that 4 volumes of the reacting gases form 2 volumes of ammonia. Le Chatelier's Principle predicts that the forward reaction will be favoured by applying a stress on the system such as to reduce the overall volume. This is achieved by increasing the pressure on the reacting system. In practice, the pressure that is used is around 40 MPa (about 400 times normal atmospheric pressure).
We note also that heat is evolved in the process (ΔH is negative - the reaction is exothermic). The forward reaction will therefore be promoted by a reduction in temperature, thus improving the yield of ammonia at equilibrium. However, at low temperatures the rate of attaining the equilibrium is so slow as to make the process impracticable. So, in order to carry out the process, a temperature of about 500 ºC is used, and even so, it has to be speeded up by the use of an iron oxide catalyst. The equilibrium mixture contains about 25-30% ammonia. This gas is easily liquefied, thus enabling its separation from the unreacted nitrogen and hydrogen, which are recycled.
Sulphuric acid is produced in industry by first oxidising sulphur dioxide gas (SO2) to sulphur trioxide (SO3):
The reaction is exothermic, and results in a reduction in volume. In practice, a temperature of about 450 °C and a pressure of 100-200 kPa are used. A catalyst, vanadium pentoxide, V2O5, is used, thus increasing the rate of the reaction at the relatively high temperature used.
Nitric acid is produced industrially by the Ostwald process, which involves three steps, the first of which is the catalytic oxidation of ammonia:
The reaction involves a decrease in volume, which by Le Chatelier's principle will be favoured by an increase in pressure. In practice, pressures of 400-1000 kPa are used. The reaction is very strongly exothermic, so as low a temperature as is practicable is used, 900 °C. A Platinum-rhodium catalyst speeds the reaction up.
Notice that all three of the above examples involve exothermic reactions, which should be carried out at as low a temperature as possible. The reaction rates are too low to make this a practical consideration, so high temperatures are used, and appropriate catalysts speed the reactions up.