REVERSIBLE AND IRREVERSIBLE REACTIONS

A reversible reaction tends to proceed in both directions, i.e., forward direction and backward direction, and is never completed. It means that a reversible reaction can never be 100% completed. The ultimate goal of a reversible reaction is to achieve the equilibrium position, which is a stable position. The yield obtained in a reversible reaction is comparatively lower, and these reactions are represented by a double-headed arrow (⇋). A reversible reaction can be carried out only in a closed container, especially when gases are involved. The chemical equation of a reversible reaction is shown as: A + B ⇌ C + D.

Whereas the irreversible reactions are those that are unidirectional and tend to completion. The yield produced in an irreversible reaction is better, and these are represented by a single arrow (→). Irreversible reactions can be carried out both in closed and open containers. For example, the following reaction is an irreversible reaction. A + B → C + D.

STATE OF CHEMICAL EQUILIBRIUM OR DYNAMIC EQUILIBRIUM

There are two main types of chemical equilibrium

1. Static equilibrium: In static equilibrium, nothing is happening internally; the forward and reverse processes have ceased.
2. Dynamic equilibrium, in dynamic equilibrium, the individual forward and reverse processes continue at equal rates, so the net reaction is zero.

In chemistry, when we talk about equilibrium, it always means the dynamic equilibrium in a reversible reaction.

            Thus, the chemical reaction in which the rate of forward reaction is exactly equal to the rate of backward reaction, or the reaction in which the concentration of reactants and products becomes constant at a given set of conditions, is called an equilibrium reaction, and this state is called the state or position of equilibrium.

For example, here is the equation of a chemical equilibrium reaction. A + B ⇌ C + D

• As the reaction proceeds by reacting A and B, the products C and D start to form, and the reaction proceeds to the forward direction
• When some moles of products are formed, then they react, and the back reaction (i.e., the formation of A and B) starts.
• At the initial stage, the rate of forward reaction increases, then slows down, and after that, the rate of backward reaction starts to increase.
• At equilibrium, the rate of forward reaction and the rate of backward reaction are exactly equal. i.e., rate of forward reaction = rate of backward reaction, or the concentration of reactants and products remains constant. These concentrations may or may not be equal.

CHEMICAL EQUILIBRIUM CONSTANTS (Kc, Kp, Kn, and Kx)

Kc is the ratio of the rate constant for the forward reaction (K_f) and the rate constant for the reverse reaction (K_r), i.e., Kc = K_f/K_r. Kc is also defined as, it is the ratio of the product of products concentration to the product of reactants concentration. i.e. Kc =[C][D]/[A][B]. The different constants used in equilibrium reactions are Kc, Kp, Kx and Kn.

• The constant Kc is used when the concentration unit for the reactants and products are used is mol/dm³. Generally, we express the molar concentration in square brackets ([]).
• The constant Kp is used as an equilibrium constant when the concentration is expressed in pressure units (according to the Henery’s law the concentration of gas is directly proportional to the applied pressure).
• Kx is the equilibrium constant when the concentration of reactants and products are expressed in moles fractions.
• The constant Kn is the equilibrium constant when the concentration of reactants and products are expressed in the number of moles.

For example, in the following reaction,

N₂ + 3H₂   ⇋ 2NH₃

For this reaction the expressions of Kc = [NH₃]²/[N₂][H₂]³, Kp = P(NH₃)²/P(N₂) X P(H₂)³       

Kn =   n(NH₃)²/n(N₂) X n(H₂)³ and Kx = x (NH₃)²/x (N₂) X x(H₂)³.

These values of equilibrium constants (i.e. Kc, Kn, Kp and Kx) change only by changing the temperature.       

UNITS FOR EQUILIBRIUM CONSTANT (Kc)

In the case of writing the units of equilibrium constant Kc, firstly, write down the expression for Kc and then put the units in mol/ dm³ as shown in the following equation.

N₂ + 3H₂   ⇋ 2NH₃

 Kc = [NH₃]²/[N₂] [H₂]³ → Kc = [mol.dm^-3]²/[mol.dm^-3] [mol.dm^-3]3 = mol^-2.dm⁺⁶              

Shortcut formula for unit

There is also a shortcut formula for expressing units. This formula is,[mol.dm^-3]^Δn, where Δn is the no of moles of products – no of moles of reactants. So, in the above equation, the Δn is -2, putting this value in the expression of Kc, Kc = [mol.dm^-3]^-2 = mol^-2.dm⁺⁶.    

If the number of moles of reactants and products are equal, then Kc of this reaction will not contain any unit. For example, in the reaction, N₂ + O₂ ⇋ 2NO, Kc = [NO]²/[N₂] [O₂] → Kc = [mol.dm^-3]⁰, thus, Kc has no unit.

Sometimes, the expression of the chemical equilibrium has to be written in some mathematical form to determine the units. Suppose we have the following equation,

                                       2N₂ + 6H₂ ⇋ 4NH₃

Initially, when t = 0,       a   +     b    ⇋       0

At equilibrium,              a-2x/V,        b-6x/V    ⇋     4x/V

Now, the Kc value is, Kc = [4x/V]⁴/[a-2x/V]²x[b-6x/V]⁶   → Kc = [mol.dm^-3]^-4 → mol^-4. dm⁺¹²

RELATIONS BETWEEN DIFFERENT EQUILIBRIUM CONSTANTS

• Kp = Kc (RT)^Δn
• Kp = Kx (P)^Δn
• Kp = Kn (P/N)^Δn

When Δn of any reaction is 0, then the power of 0 is equal to 1. For this case, Kp = Kc =Kn = Kx.

When Δn of any reaction is greater than 0 (i.e. Δn ˃ 0), (it means product is greater than reactant, as Δn = no of moles of products – no of moles of reactants), then Kp (contains the alphabet p) will be greater than Kc.

When Δn is less than 0, or has a negative value, then Kp will be less than Kc.   

CHARACTERISTICS OF EQUILIBRIUM STATE OR POSITION

• It is a macroscopic property
• It is a dynamic state, not a static state
• The equilibrium state can be achieved from either direction. For example, in the following reaction, H₂ + I₂ ⇋ 2HI, to obtain the equilibrium position, we can take either HI or H₂ and I₂.  
• A change in pressure/volume, concentration, or temperature changes the equilibrium stage.
• A catalyst does not disturb the equilibrium state, but it helps to achieve the equilibrium state earlier.

CHARACTERISTICS OF EQUILIBRIUM CONSTANT (Kc)

• Kc depends upon the temperature; by changing the temperature, Kc changes.
• Kc is independent of concentration, pressure-volume, or catalyst.
• Kc does not depend upon the initial concentration of reactants, but it depends upon the concentration of products and reactants at the equilibrium stage.
• Kc is the reciprocal of Kc′. Kc = [P]/[R], Kc′ = [R]/[P]
• Kc is just a numerical value of a constant.

APPLICATIONS OF EQUILIBRIUM CONSTANT (Kc)

The equilibrium constant is widely used for the determination of

1. Direction of reaction
2. Extent of reaction

Direction of reaction

By comparing the value of Kc with Q (reaction quotient), we can determine the direction of a reaction. A reaction quotient Q =[P]/[R] before equilibrium. Whereas Kc = [P]/[R] at equilibrium.

• When Q < Kc, then the reaction will proceed in the forward direction
• When Q ˃ Kc, the reaction will proceed in a backward direction
• When Q = Kc, the reaction will be at equilibrium

Extent of reaction

• If the value of Kc is very large, then the reaction is almost complete. As in the expression of Kc, we write Kc = [P]/[R], as Kc∝ [P]. so larger value of Kc means greater amount of product or the product is very stable.    
• If the value of Kc is very small, then the reaction does not go to completion. It means that in this reaction product is less stable and reactant is very stable.
• If the value of Kc is not very large nor very small. It means that both reactants and products are stable. It means the reaction is at equilibrium.

LE CHATELIER’S PRINCIPLE

According to this principle, if stress is applied to a system at equilibrium, the system will act in such a way to nullify this effect as soon as possible. This stress can be applied in the form of factors like change in concentration, pressure/volume, or by change in temperature. To nullify all this stress, the system will gain another equilibrium position because the equilibrium position is the most stable position for a reversible reaction.

EFFECT OF A CHANGE IN CONCENTRATION ON CHEMICAL EQUILIBRIUM

Suppose we have the following reversible reaction,

A + B⇋ C + D

In this reaction, suppose Rf = 0.2 mol.dm^-3.s^-1 and Rr = 0.2 mol.dm^-3.s^-1 . If we increase the concentration of A and B or remove the C and D, then the reaction will move in the forward direction. On the other hand, when we increase the concentration of C and D or remove A and B, then the reaction will move to the backward direction. For example, in the reaction, BiCl₃ + H₂O ⇋ BiOCl + HCl (BiOCl is insoluble in H₂O and is called artificial milk). By adding the water, the milkiness of solution increases, because the reaction moves in the forward direction and more BiOCl is formed. On the other hand, by adding HCl reaction will become clear solution or milkiness disappears. Thus, reaction will move in the backward direction.

EFFECT OF PRESSURE/ VOLUME ON CHEMICAL EQUILIBRIUM

1. In those chemical reactions in which the number of moles of products = number of moles of reactants i.e Δn = 0, in these reactions the change in P/V does not affect the equilibrium position. For example, H₂ (g) + I₂ (g)  ⇋ 2HI (g), nP = nR
2. For gaseous phase reaction in which Δn ≠0, a change in pressure/volume affect the equilibrium state. For example, N_2(g) + 3H_{2 (g)}  ⇋ 2NH_3(g) (Δn = -2) or 2SO₂ + O₂ ⇋ 2SO₃ (Δn = -1), PCl₅ ⇋ PCl₃ + Cl₂ (Δn =1). In these reactions, if we increase the pressure or decrease the volume, the reaction will move in that direction, where the reactants contain lesser number of moles and vice versa.

Similarly, in the relation of pressure-volume, on increasing the pressure the reaction will move in that directions where there is lesser volume or on decreasing the pressure the reaction will move in those directions where there is greater volume. For example, when liquid ⇋ vapours, then by increasing the pressure more liquids will form because the volume of liquid is lesser compared to vapours. In the same way, when ice and liquid are in equilibrium with each other as shown by the following equilibrium, liquid ⇋ ice. As the volume of ice is greater (9%) than the water. So, by increasing the pressure more water will be formed. For any other solid instead of ice by increasing the pressure the amount of solid increase because in any liquid ⇋ solid equilibrium the volume of solid is lesser compared to liquid.   

EFFECT OF TEMPERATURE ON CHEMICAL EQUILIBRIUM

            The change in temperature also affects the state of equilibrium. An increased in temperature shows positive effect (reaction moves in forward direction) in endothermic reactions and negative affect on exothermic reactions (the reaction will move in the reverse reaction).

When we dissolve some solute in the solvent and form a solution. If the solution/beaker becomes hot, the reaction will be exothermic. If the solution/beaker become cold, then the reaction will be endothermic. For example, in the following reaction,

Solid + solvent   ⇋ soln                      ΔH = -KJ/mol

It means that this reaction is exothermic, so by decreasing the temperature solute will dissolve and form solution. However, in the reverse direction, this reaction is endothermic, that is by increasing the temperature solute will be precipitated out.

EFFECT OF THE CATALYST ON  CHEMICAL EQUILIBRIUM

A catalyst neither affects the equilibrium position nor the Kc. However, a catalyst helps to achieve the equilibrium stage earlier.