3.2.3 Equilibria - Qualitative effects of changes of pressure, temperature and concentration on a system in equilibrium
Students should:
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The equilibrium constant
Kc is a constant which represents how far the reaction will proceed at a given temperature.
When Kc is greater than 1, products exceed reactants (at equilibrium). When much greater than 1, the reaction goes almost to completion. When Kc is less than 1, reactants exceed products. When much less than 1 (Kc can never be negative...so when it is close to zero) the reaction hardly occurs at all.
The only thing which can change the value of Kc for a given reaction is a change in temperature. The position of equilibrium, however, can change without a change in the value of Kc.
Effect of Temperature
The effect of a change of temperature on a reaction will depend on whether the reaction is exothermic or endothermic. When the temperature increases, Le Chatelier's principle says the reaction will proceed in such a way as to counteract this change, ie lower the temperature. Therefore, endothermic reactions will move forward, and exothermic reactions will move backwards (thus becoming endothermic). The reverse is true for a lowering of temperature.
Reactions are either exothermic or endothermic i.e. when a reaction proceeds it does so accompanied by either a release of energy or an absorption of energy. This can be shown on an energy coordinate graph.
You can see that the activation energy needed for an endothermic change is much greater than for the corresponding reverse exothermic change.
If a reaction is exothermic in the forward direction in an equilibrium it will be endothermic in the reverse direction and vice-versa (law of conservation of energy)
Increasing the temperature of an equilibrium mixture will always favour the endothermic direction over the exothermic process. The rate of the endothermic process will increase more than the exothermic direction of change and the equilibrium will be re-established with new concentrations based on more of the endothermic product (i.e. the product of the endothermic direction of change)
To express this in non-scientific terms, it can be considered that the endothermic direction is absorbing heat therefore giving it more heat by increasing the temperature will favour it.
The reverse argument is true for the direction of exothermic change. The exothermic reaction is giving out heat and therefore applying more will hinder this process. Raising the temperature retards the reaction in the exothermic direction.
Effect of Concentration
When the concentration of a product is increased, the reaction proceeds in reverse to decrease the concentration of the products. When the concentration of a reactant is increased, the reaction proceeds forward to decrease the concentration of reactants.
For a system at equilibrium the concentrations of both reactants and products are constant (but not the same). The value of Kc depends on these concentrations and Kc is constant unless the temperature is changed.
If a further quantity of reactant is added to the mixture already at equilibrium then the value of [products]/[reactants] no longer equals the value of Kc and the equilibrium must make adjustments to reestablish the equilibrium concentrations.
In this example the value of the ratio [products]/[reactants] is too small as more reactants have been added. The only way to adjust this is by making more product (and at the same time using up some of the reactants)
The equilibrium responds by making more product until the equilibrium ratio is reestablished at the Kc value.
In summary the effect of adding reactants has been to make more product.
This gives some general rules:
- If we add to the left hand side of an equilibrium we make more of the right hand side
- If we add to the right hand side of an equilibrium we make more of the left hand side
- If we remove from the left hand side the equilibrium makes more of the left hand side component
- If we remove from the right hand side the equilibrium makes more of the left hand side component
It is consequently easy to see how reactions can be driven one way or another by adding or removing components. This is very useful in many reactions.
| Example: The formation of ammonia in the Haber process only results in about 15% of ammonia being present at equilibrium. However, by removing ammonia from the gaseous mixture by liquefaction, the gases can then be recycled through the catalyst to make more ammonia. |
Explanation
All of this can be easily explained by considering the rates of the forwards and back reactions.
At equilibrium the rates of the forward and back reactions are equal:
rate forwards = rate backwards
When we study rates of reaction one of the first conclusions drawn is that the rate of a chemical reaction depends on the concentrations of the reactants. The greater the concentration the more collisions occur and the faster the rate of the reaction.
When we add more reactant the forward rate will now be greater than the back rate. The reaction is now not at equilibrium and the forward reaction proceeds faster than the back reaction until the equilibrium conditions are reestablished.
Similiarly removal of a component from one side will reduce the rate of its reaction and case the equilibrium to make more of it.
In all of these changes the value of Kc remains unchanged. The equilibrium is temporarily disturbed and then the equilibrium concentrations are re-established.
Effect of Pressure
In reactions where gases are produced (or there are more mols of gas on the left), and increase in pressure will force the reaction to move to the left (in reverse). If pressure is decreased, the reaction will proceed forward to increase pressure. If there are more mols of gas on the left of the equation, this is all reversed.
Pressure change only affects gaseous equilibria and only then when there are an unequal number of moles of gas on either side of the equilibrium.
When the pressure is increased on a gaseous mixture at equilibrium the equilibrium will respond so as to release the applied pressure. In other words the equilibrium will move towards the side of the fewer number of moles.
It should be remembered that the pressure of a gas is directly proportional to the number of moles of gas and that this does not depend on the nature of the gas in question - all gases behave in the same way in terms of their physical properties.
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Example: For the Haber Process N2 + 3H2 four moles gas There are four moles of gas on the left hand side and only two moles of gas on the right hand side. Increasing the pressure will move the equilibrium to the right hand side and have the effect of releasing the pressure. |
2NH3 This is rather a difficult concept to provide a satisfactory simple answer. However it is enough to know that the equilibrium responds to the change in conditions by counteracting the change in conditions (i.e. increase the pressure and it responds releasing the pressure by making a fewer number of moles)
The actual explanantion involves considering the consequences of changing the pressure of a gas.
If the number of moles of gas is constant and the temperature doesn't change then the only way to change the pressure is to change the volume of the container.
If the volume of the container changes then the concentration (moles/volume) must also change. Hence the effect of changing the pressure is a corresponding change in the concentration - double the pressure and the concentration doubles.
If the equilibrium expression has a different number of moles on the top and the bottom then doubling the concentration of all the components will change the ratio of the equilibrium expression. The equilibrium expression no longer equals Kc and so the system responds to re-establish the ratio to be equal to Kc again.
Effect of catalysts on equilibrium
A catalyst does not effect either Kc or the position of equilibrium, it only effects the rate of reaction. As the rate of forward reaction and reverse reaction is affected equally then the equilibrium cannot be affected.