The reaction between iron and steam is a reversible reaction.
If steam is passed over heated iron in an open vessel, the forward reaction takes place and iron oxide and hydrogen are produced. The steam displaces the hydrogen from the system, thereby limiting the reverse reaction. When hydrogen is passed over iron oxide, iron and steam are produced.
An equilibrium is established, however, if iron and steam are heated in a closed container so the products cannot escape. In equilibrium, the concentrations of the two reactants and the two products remain constant provided external conditions are unchanged. This doesn't mean that the reactions have stopped, however. What happens is that the rate of the forward reaction becomes equal to the rate of the backward reaction so the concentrations are unchanged. It is better to use the term dynamic equilibrium to emphasise that the two reactions are continuing.
Consider the general chemical reaction:
A, B, C and D represent the chemical substances; a, b, c and d represent the number of moles of each in the equation.
If A and B are mixed a reaction occurs and C and D are produced. Eventually equilibrium is established. If the equilibrium concentrations of A, B, C and D are determined, it is found that:
Kc is known as the equilibrium constant (subscript c refers to concentration).
Consider the reaction:
Concentration is normally expressed in mol dm-3. Kc has no units in reactions with equal numbers of moles on both sides of the chemical equation. Where the numbers of moles of reactants and products are not equal Kc will have units, e.g.:
N2(g) + 3H2(g) 2NH3
When a system in equilibrium is suddenly disturbed, the system will respond in some way until equilibrium is eventually re-established.
Le Chatelier's Principle states that if a system in equilibrium is subjected to a change, processes occur which tend to counteract the change imposed.
If the concentration of a reactant in an equilibrium system is suddenly changed, the rate of the forward reaction alters. Similarly, if a product concentration is suddenly altered, the rate of the backward reaction alters. In either of these two cases, equilibrium is disrupted because the rates of the two opposing processes are no longer equal. The concentration of the reactants and products change until a new equilibrium position is reached. Addition of a reactant to an equilibrium results in a shift in equilibrium position to the right. Removal of a product results in a shift in equilibrium position also to the right. Addition of a product results in a shift to the left. In these cases the value of Kc remains unchanged.
Pressure changes can only affect those equilibria that involve gaseous reactants or products. If the total pressure of a gas mixture is increased by compressing it, then Le Chatelier's principle tells us the system will respond by reducing the pressure again.
The pressure exerted by a gas depends on the number of gas molecules present, so the equilibrium shifts to the side with fewer gas molecules.
Consider the reaction:
N2(g) + 3H2(g) 2NH3
There are four moles of gas on the reactant side and two moles on the product side. If pressure is increased, the system reacts to reduce the pressure by reducing the total number of moles of gas. The reaction shifts to the right and produces more ammonia.
If there are the same number of moles of gas on each side of the chemical equation, increasing the pressure has no effect on the equilibrium position. Kc remains unchanged.
In the ammonia reaction, the forward reaction is exothermic:
N2(g) + 3H2(g) 2NH3 H = -92kJ mol-1
Hence the backward reaction is endothermic, so increasing the temperature will favour the backward reaction as this is the reaction which needs heat.
Changing the temperature of a system in equilibrium results in the establishment of a new equilibrium constant. An increase in temperature of an equilibrium results in an increase in the equilibrium constant if the reaction involved is endothermic; and a decrease in the value of the equilibrium constant if the reaction is exothermic.
Catalysts do not affect the position of equilibrium - they only affect the rate at which equilibrium is attained.
In homogeneous systems, all reactants and products are in the same phase, either all gases, in aqueous solution, or all liquids.
In heterogeneous systems, the participating substances are present in more than one phase, e.g.
CaCO3(s) CaO(s) + CO2(g)
The concentrations of CaO and CaCO3 are constant, so we can write a new equilibrium constant:
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