Introduction to chemical equilibria, Chemistry tutorial


According to the theory, all chemical reactions are in fact double reactions: for each and every forward reaction, there is a successive reverse reaction. The idea can be described as follows:

reactants ↔ products

For abundance of reactions, though, the forward reaction is so favored, and the reverse reaction is so negligible, that the reactions are written simply in terms of solid forward arrow, A → B. Though, we will now consider forward or reverse reaction pairs which exist in the chemical equilibrium with one other.

Concept of Chemical Equilibrium:

In a chemical reaction, the chemical equilibrium is a state in which the forward reaction rate and the reverse reaction rate are equivalent. The outcome of this equilibrium is that the concentrations of the reactants and the products don't change. Though, just as the concentrations are not changing doesn't signify that all the chemical reaction has ceased. Just the opposite is true; the chemical equilibrium is a dynamic state in which reactants are being transformed into products at all times, however at the exact rate that products are being transformed back into reactants.

Define: There is equilibrium whenever the concentrations of reactants and products are in an unchanging ratio. The other way of saying this is that a system is in equilibrium if the forward and reverse reactions take place at equivalent rates.

Characteristics of Chemical Equilibrium:

1) At the equilibrium state, rates of forward and backward reactions are equivalent.

2) The observable properties like pressure, density, concentration, color, viscosity and so on, of the system remain unchanged by time.

3) The chemical equilibrium is a dynamic equilibrium, as both the forward and backward reactions continue to take place even although it appears static externally.

The concentrations of reactants and products don't change by time however their inter conversion continue to take place.

4) The chemical equilibrium can be reached via beginning the reaction either from the reactants side or from the products side.

5) Both the pressure and concentration influence the state of equilibrium however don't influence the equilibrium constant.

6) Though, temperature can influence both the state of equilibrium and also the equilibrium constant.

7) A positive catalyst can raise the rates of both forward and backward reactions and therefore helping the system to accomplish the equilibrium faster. However it doesn't influence the state of equilibrium and the equilibrium constant.

Reaching Equilibrium and the Equilibrium Position:

Whenever a system comprises of competing forward and reverse reaction rates, the reaction will carry on till the chemical equilibrium is reached. Consider the given general reaction between A and B:

A ↔ B

Let's state that A and B are introduced to a reaction vessel, and a reaction starts to carry on. The concentrations of A and B will change over time, till they reach equilibrium.

454_concentration versus time for reversible reaction.jpg

Fig: concentration versus time for reversible reaction

It will be noted that over time, the curves level out, or plateau, and the concentrations of A and B are no longer changing. This is the point at which the system has reached the chemical equilibrium. As there are different factors which can increase or decrease the amount of time it takes for a given system to reach the equilibrium, the equilibrium position itself is not affected by such factors. For example, if a catalyst is added to the system, the reaction will continue more quickly, and equilibrium will be reached faster, however the concentrations of both A and B will be similar at equilibrium for both the catalyzed and the uncatalyzed reaction.

The Equilibrium Conditions:

All the chemical reactions are preceded via molecular collisions, however only such collisions having proper orientation and adequate energy to form the transition state complex would next lead to the formation of products. Consider the given reversible reaction:

H2 (g) + I2 (g) ↔ 2 HI (g)

that comprises of forward and reverse processes:

Forward reaction:  H2 (g) + I2 (g) → 2 HI (g).... (i)   

Reverse reaction:  2 HI (g) → H2 (g) + I2 (g).... (ii)

A) The rate of forward reaction (equation-i) based on the frequency of collisions that in turn based on the concentrations of H2 and I2.

B) As HI molecules are made up, they as well collide by one other and some of them dissociate to form H2 and I2 according to equation-ii - the reverse process.

C) As both the opposing reactions progress, the concentration of products rises and thus the rate of the reverse reaction.  At similar time, the reactant concentrations and the rate of forward reaction decrease.

D) Finally, the rates of the two opposing processes become equivalent and there is no total change in the concentrations of the reacting species.

E) The state of dynamic chemical equilibrium is established, like,

H2 (g) + I2 (g) ↔ 2HI (g);

Equilibrium is the precise balancing of two opposing methods. This is a dynamic state, where both the forward and reverse processes continue by equivalent rates. At equilibrium, the concentrations of all the species in system remain constant. 

Irreversible and Reversible Reactions:

Irreversible reaction:

A reaction which takes place only in one direction is termed as an irreversible reaction that is, only the reactants are converted to products and the conversion of products to reactants is not possible.

The single headed arrow (→) is utilized to point out the irreversible reactions.


a) The combustion of methane is an irreversible reaction as it is not possible to convert the products (like carbon-dioxide and water) back to the reactants (like methane and oxygen).

CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l)

b) The decomposition of potassium chlorate is as well an irreversible reaction. This is not possible to make potassium chlorate directly from the KCl and O2.

2KClO3 (s) → 2KCl (s) + 3O2 (g)

Reversible reaction:

A reaction which takes place in both the forward and backward directions is termed as reversible reaction. In a reversible reaction, the reactants are converted into products and the products can as well be converted back to the reactants.

The half headed arrows (↔) are utilized to point out the reversible reactions.

Example: The given reactions are reversible reactions as they take place in both directions.

a) H2 (g) + I2 (g) ↔ 2 HI (g)

b) CaCO3 (s) ↔ CaO (s) + CO2 (g)

c) N2 (g) + 3H2 (g) ↔ 2 NH3 (g)

Chemical equilibrium is possible only in the reversible reactions.

Types of Chemical Equilibria:

The chemical equilibria are categorized into two types: (1) Homogeneous equilibrium and (2) Heterogeneous equilibrium.

1) Homogeneous equilibrium:

A chemical equilibrium is stated to be homogeneous if the entire substance (that is, reactants and products) at equilibrium are in the similar phase.


a) H2 (g) + I2 (g) ↔ 2 HI (g)

b) N2 (g) + 3H2 (g) ↔ 2 NH3 (g)

2) Heterogeneous equilibrium:

A chemical equilibrium is stated to be heterogeneous the entire substance at equilibrium are not in the similar phase.


a) CaCO3 (s) ↔ CaO (s) + CO2 (g)

b) NH4HS (s) ↔ NH3 (g) + H2S (g)

Equilibrium Constant, Kc (In Terms Of Molar Concentration):

Consider the given reaction at equilibrium:

aA + bB ↔ cC + dD

By applying the law of mass action, the rate of forward reaction can be represented as:

Vf α [A]a [B]b

Vf = Kf [A]a [B]b

The rate of backward reaction can be represented as:

Vb α [C]c [D]d

Vb = Kb [C]c [D]d


[A], [B], [C] and [D] are the equilibrium concentrations of A, B, C and D correspondingly.

a, b, c and d represent the stoichiometric coefficients of A, B, C and D correspondingly.

Kf and Kb are the rate constants of forward and backward reactions correspondingly.

However at equilibrium,

Rate of forward reaction = Rate of backward reaction

That is, Vf = Vb

i.e. Kf [A]a [B]b = kb [C]c [D]d

Or Kf/kb = [C]c [D]d/[A]a [B]b

Or KC = [C]c [D]d/[A]a [B]b

Here, KC = Kf/kb

KC is termed as the equilibrium constant deduced in terms of molar concentrations. This is the ratio of products of equilibrium concentrations of products to the product of equilibrium concentrations of reactants. In common, it can be written as:

KC = Product of concentrations of products at equilibrium/ Product of concentrations of reactants at equilibrium

Units of KC:

Most of the times, KC is deduced without units. Though the units of KC can be represented as (mol.L-1)Δn.

Here Δn = (c + d) - (a + b) = (total number of moles of products) - (total number of moles of reactants).

Factors influencing the equilibrium constant:

a) There is no effect of the concentration, pressure and catalyst on the value of equilibrium constant.

b) Though, the equilibrium constant based on the temperature.

Generally, in the exothermic reactions, increase in the temperature decreases the equilibrium constant, KC.

As, in endothermic reactions, increase in temperature increases the KC value.

Equilibrium Constant, KP (In Terms Of Partial Pressures):

KP is the equilibrium constant in terms of partial pressures. Assume that A, B, C and D are gases in the given reaction.

aA (g) + bB (g) ↔ cC (g) + dD (g)

Then for the above reaction, the KP can be represented as:

KP = PCc PDd/PAa PBb

Here PA, PB, PC and PD are the partial pressures of A, B, C and D at equilibrium.

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