A Level Chemistry Revision: Physical Chemistry – Chemical Equilibria, Le Chatelier’s Principle & Kc

by Lucy Bell-Young

Chemical equilibrium is achieved when the rate of the forward reaction is equal to the rate of the backward reaction. In other words, it’s the state of a system where the concentration of the reactants and products are constant, i.e. they do not change with time.

All dynamic systems, including chemical systems, tend to settle at certain equilibrium points under certain conditions. Introducing stress to a system in equilibrium will cause the system to shift its balance to counteract the stress and regain a new equilibrium at a different point. 

For example, flowing water tends to go from a high point to the lowest possible level that is in equilibrium in a given terrain. You can also think of it as something similar to a simple level machine that shifts the fulcrum depending on the loads on either end of the lever.

What is Chemical Equilibrium?

Chemical equilibrium is when a reversible chemical reaction doesn’t change the quantity of reactants or products. But, while the concentration of reactants and products are constant when equilibrium is reached, that doesn’t mean they’re equal. 

All chemical reactions are technically reversible under certain conditions. However, when it comes to standard conditions, reversible chemical reactions are those where the products produce the original reactants as soon as they are formed.

The rates of the two opposing reactions are constant once they achieve the equilibrium point. This means no additional amounts of reactants or products are produced. At equilibrium point, the reaction is considered complete because the maximum possible amount of reactants is converted into a new product.

Chemical equilibrium doesn’t only refer to chemical reactions where bonds are broken and created, leading to the formation of new products. It could also refer to a stable state of a particular substance. In fact, you can find many examples of chemical equilibrium in everyday life that you probably never noticed before.

Illustration of chemical equilibrium

What Are Some Examples of Chemical Equilibrium?

One of the most common examples of chemical equilibrium is in a can of fizzy drink. Soft drinks are pressurised because of the carbon dioxide trapped inside the can. Since most of the carbon dioxide molecules are dissolved in the liquid, you won’t initially see bubbles in the drink – but once you open the can, the gas will come rushing out of the top.

As long as the can is tightly closed, the gaseous and liquid carbon dioxide forms are in equilibrium with each other. A constant movement of the carbon dioxide from liquid to gas phase, and vice versa, is continuous but in balance. As long as you do not shake or open the can, you won’t be able to observe any change. This chemical equilibrium can be written as a balanced chemical equation:

CO2(g)+H2O(l) ⇌ H2CO3(aq)

Similarly, many biological functions would not be possible without dynamic equilibria, the best example of which is found within our bodies, namely, the haemoglobin in our blood. Haemoglobin is responsible for transporting oxygen to the different parts of our bodies. Without the chemical equilibrium in haemoglobin, we wouldn’t be able to survive. This macromolecule is in equilibrium in terms of absorbing and releasing the oxygen it carries. It is represented by this chemical equation:

haemoglobin(aq) + 4O2(g) ⇌ haemoglobin(O2)4(aq)

Haemoglobin chemical formula
Many chemical reactions, however, do not seem to achieve dynamic equilibrium. Instead they seem to go only in one direction until one or more of the available reactants are fully consumed. Consider the chemical equation representing the formation of rust as a result of oxidation:

4Fe + 6H2O + 3O2 → 4Fe(OH)3

What’s the Difference Between Static and Dynamic Equilibrium?

As previously mentioned, all chemical reactions are technically reversible under specific conditions. However, many chemical reactions are practically irreversible because of high entropy. For example, a safety match that’s completely burned is virtually impossible to put back together again, even if all the atoms and energy are contained in a closed system. This is known as static equilibrium. 

Chemical static equilibrium is a state of a system wherein the chemical reaction has completely stopped. The rate of reaction is zero and there is no movement between the reactants and the products. Here are the key differences between static equilibrium and dynamic equilibrium:


A table showing static vs dynamic equilibriumWhat is Le Chatelier’s Principle?

All systems that are in a state of dynamic equilibrium are only stable as long as the factors are maintained at a reasonably stable range. Case in point, the reaction between nitrogen and hydrogen to form ammonia:

N2(g) + 3H2(g) ⇔ 2NH3(g)

To produce this reaction, a pressure of between 2100-3600 psi, and a temperature of between 300-550°C are needed. Under these conditions, the ammonia gas decomposes into its constituents. Therefore, the forward and backward reactions are in equilibrium.

In this example, and others similar to it, Le Chatelier’s principle can be observed and applied. The principle states that in a chemical system in equilibrium, the changes in one or more of the factors or conditions will result in the predictable shift in reactions to achieve a new equilibrium state.

The factors that can shift the equilibrium state are the following:

  • Temperature
  • Pressure
  • Concentration
  • Volume

For example, consider the following exothermic neutralisation reaction between a base and an acid:

NaOH + HCl → H2O + NaCl + energy

Under normal conditions, sodium hydroxide readily reacts with hydrogen chloride to form water and sodium chloride (table salt) without the need for input energy or a catalyst. The reaction is exothermic, producing a significant amount of heat. However, the reverse reaction does not readily occur.

The ionic bond between sodium and chlorine is too strong, making the salt highly stable. Similarly, the covalent bond between hydrogen and oxygen in a water molecule is also strong. You need a significant amount of energy input in the form of an electric current to split the hydrogen from oxygen and break the ionic bond in the salt molecules.

In an exothermic reaction, adding energy to the right side of the equation shifts the equilibrium to the left. Conversely, in an endothermic reaction, adding heat or other forms of energy to the reactant shifts the equilibrium to the right side of the equation.

How to Compute the Equilibrium Constant

Reversible chemical reactions in dynamic equilibrium have specific equilibrium constants. The generalised form of reversible chemical reactions can be written as:


Generalised form of reversible chemical reactionsMeanwhile, the equilibrium constant for every reaction can be written as:
The equilibrium constant formula

Here, A and B represent the reactants, and C and D represent the products. Meanwhile, w, x, y, and z represent the coefficients of the reactants and products. The constant of the equation (Keq) represents the concentration constant, which can be written as Kc. Although the units are in moles per litre (mol/L), the K is unitless since it’s just a ratio. The value of Kc provides insight on the proportional concentrations of the reactants and products under ideal conditions:

  • If the value of Kc is ~1000 or more, the products are proportionally greater than the reactants at equilibrium state
  • If the value of Kc is ~0.001 or less, the reactants are proportionally greater than the products
  • If the value of Kc is between 0.001 and 1000, the amounts of reactants and products are about the same

Changing any of the ideal condition parameters, such as temperature and pressure, will shift the equilibrium state either to the left or the right side of the equation.


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