Chemical Equilibrium

1. Introduction

Many chemical reactions do not go to completion; instead they reach a state in which reactants and products coexist and their amounts no longer change. This state is called chemical equilibrium. Understanding equilibrium lets us predict how much product a reaction can give and how to shift conditions (temperature, pressure, concentration) to increase the yield, which is central to industrial processes such as the manufacture of ammonia and sulphuric acid.

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2. Reversible and Irreversible Reactions

• An irreversible reaction proceeds in one direction only, essentially to completion; the products do not react back to give the reactants. Examples: combustion, and precipitation such as

$$ \ce{AgNO3(aq) + NaCl(aq) -> AgCl(s) + NaNO3(aq)} $$

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• A reversible reaction proceeds in both directions simultaneously: reactants form products (forward reaction) and products reform reactants (backward reaction). It is written with a double half-arrow ($\ce{<=>}$). Example:

$$\ce{N2(g) + 3H2(g) <=> 2NH3(g)}$$

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3. Chemical Equilibrium

Chemical equilibrium is the state of a reversible reaction, carried out in a closed system, at which the rate of the forward reaction equals the rate of the backward reaction, so the concentrations of all reactants and products remain constant with time.

Characteristics of chemical equilibrium:

1. It is attained only in a closed system (nothing escapes).
2. It is dynamic: both reactions continue at the molecular level, but at equal and opposite rates, so no net change is observed.
3. The observable properties (concentration, colour, pressure) stay constant once equilibrium is reached.
4. It can be approached from either direction (starting from reactants or from products) and gives the same equilibrium state.
5. A catalyst speeds up the attainment of equilibrium but does not change the equilibrium position or the amounts present.

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