Reaction rate describes how quickly a reaction occurs. Equilibrium describes the relative amounts of reactants and products in a reversible reaction.
Rate is how quickly a reactant is used up or a product is formed: \(\text{rate} = \frac{\text{quantity used or formed}}{\text{time}}\), in g/s or cm3/s, or mol/s.
It can be followed by measuring the mass lost as a gas escapes, the volume of gas collected, or the time for a solution to turn cloudy enough to hide a cross. On a graph of product against time, the rate at any moment is the gradient, and it is usually steepest at the start, when the reactant concentrations are highest.
Reactions happen when particles collide with enough energy. The minimum needed is the activation energy.
So the rate depends on how often particles collide and how much energy they have when they do.
Each factor changes the rate by changing the frequency of collisions or the proportion of collisions with enough energy.
A catalyst speeds up a reaction without being used up by it.
It works by providing an alternative reaction pathway with a lower activation energy. It does not change the position of the equilibrium or the equilibrium yield, only how quickly equilibrium is reached. Enzymes are biological catalysts.
In a reversible reaction the products can react to re-form the reactants, shown by the ⇌ symbol.
If the forward reaction is exothermic then the reverse is endothermic, and by exactly the same amount of energy. Hydrated copper sulfate provides an example: heating it drives water off and turns it white, and adding water turns it blue again while giving out heat.
In a closed system a reversible reaction reaches equilibrium when the forward and reverse reactions are happening at the same rate.
It is a dynamic equilibrium: both reactions continue, but the amounts of reactants and products stop changing. A closed system is required for the dynamic equilibrium considered here.
Le Chatelier's principle: if a change is made to a system at equilibrium, the position of equilibrium shifts so as to partly oppose that change.
Increasing the concentration of a reactant shifts the equilibrium towards the products until equilibrium is restored; removing a product does the same. Increasing the concentration of a product shifts it back towards the reactants.
Raising the temperature shifts the equilibrium in the endothermic direction, because that absorbs the extra energy. Lowering it shifts the equilibrium in the exothermic direction.
So for a reaction that is exothermic forwards, a lower temperature gives a higher yield, but also a slower reaction, which is the compromise behind industrial conditions.
For reactions involving gases, increasing the pressure shifts the equilibrium towards the side with fewer moles of gas, as shown by the balancing coefficients. This partly opposes the increase in pressure.
Count the moles of gas on each side of the equation first. If both sides have the same number, changing the pressure has no effect on the position of equilibrium.
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