5.3.1 Reversible reactions and dynamic equilibrium

Reversible reaction products can turn back into reactants

1. The change from reactants to products is the forward reaction, and the change from products back to reactants is the reverse reaction.
2. A reversible reaction is shown with the equilibrium sign, $\rightleftharpoons$, instead of a single arrow.
3. Heating ammonium chloride decomposes it into ammonia and hydrogen chloride: $${NH}_{4}Cl(s)\rightleftharpoons {NH}_3(g)+HCl(g)$$
   1. Near the cool top of the test tube, the two gases react again to form a white solid of ammonium chloride.
4. Heating blue hydrated copper(II) sulfate drives off water and leaves white anhydrous copper(II) sulfate.
   1. Adding water to the white solid turns it blue again.

Changing conditions can reverse a reaction's direction

1. Which direction a reversible reaction goes depends on the conditions, such as temperature, pressure and the amounts of substances present.
2. For ammonium chloride, heating favours the forward reaction and cooling favours the reverse reaction.
3. For copper(II) sulfate, heating removes water, while adding water reverses the change.
4. When anhydrous copper(II) sulfate turns blue, the mixture gets hot, because the reverse reaction gives out the energy that heating put in.
5. Most reactions, such as burning a fuel, cannot be reversed by changing the conditions.

Reversible reactions reach dynamic equilibrium in closed systems

1. At the start, only reactants are present, so the forward reaction is fast and the reverse reaction has not begun.
2. As products build up, the reverse reaction speeds up, while the forward reaction slows as reactants are used up.
3. Eventually the forward and reverse reactions happen at the same rate.
   1. From then on, the amounts of reactants and products stay constant.
4. Equilibrium can be reached only in a closed system, because an escaping product could never react back.
   1. Heating calcium carbonate in an open container lets the carbon dioxide escape, so the reaction goes to completion instead.

Both reactions continue at dynamic equilibrium

1. The equilibrium is dynamic because particles keep reacting in both directions.
2. No change can be seen, because each product particle formed is balanced by one reacting back.
3. An equilibrium mixture can contain mostly products, mostly reactants or similar amounts of both.
4. How far an equilibrium lies towards the products, and how that can be changed, is covered in the next topic.

5.3.2 Changing the position of equilibrium

Equilibrium shifts to oppose a change in conditions

1. When an equilibrium lies to the right, the mixture contains mostly products.
2. When it lies to the left, the mixture contains mostly reactants.
3. Changing the concentration, temperature or pressure disturbs the equilibrium.
   1. The system responds by shifting in the direction that opposes the change, until a new equilibrium is reached.
4. At the new equilibrium, the forward and reverse rates are equal again, but the amounts of each substance are different.

Adding or removing a substance shifts the equilibrium

1. Adding more reactant shifts the equilibrium to the right, forming more product.
2. Removing product as it forms also shifts the equilibrium to the right.
3. Adding more product shifts the equilibrium to the left, forming more reactant.
4. Removing a reactant shifts the equilibrium to the left, to replace some of it.

Raising temperature favours the endothermic direction

1. If the forward reaction is exothermic, the reverse reaction is endothermic by the same amount.
2. Raising the temperature shifts the equilibrium in the endothermic direction, which takes in energy and so opposes the rise.
3. Lowering the temperature shifts the equilibrium in the exothermic direction, which releases energy.
   1. So for an exothermic forward reaction, a lower temperature gives more product at equilibrium.
4. For an endothermic forward reaction, a higher temperature gives more product.

Raising pressure favours fewer gas molecules

1. To compare the sides, add up the balancing numbers of the gases on each side of the equation.
2. Raising the pressure shifts the equilibrium to the side with fewer gas molecules, which reduces the pressure.
3. Lowering the pressure shifts it to the side with more gas molecules.
   1. In $2{NO}_2(g)\rightleftharpoons N_2{O}_4(g)$, there are 2 gas molecules on the left and 1 on the right, so high pressure shifts the equilibrium to the right.
4. If both sides have the same number of gas molecules, as in $H_2(g)+I_2(g)\rightleftharpoons 2HI(g)$, pressure does not change the position.

Choosing conditions to give the most product

1. First find which side the desired product is on and whether its formation is exothermic or endothermic.
2. Choose a low temperature if the forward reaction is exothermic, or a high temperature if it is endothermic.
3. Choose a high pressure if the product side has fewer gas molecules, or a low pressure if it has more.
4. Keep the reactant concentrations high and remove the product as it forms.
5. A catalyst speeds up the forward and reverse reactions equally, so it does not change the position of equilibrium.
   1. Industry often uses compromise conditions that also give a fast enough rate at a sensible cost, and these are covered with industrial processes.