Let us consider a general reversible reaction:
`color{red}(A+B ⇌ C+ D)`
where A and B are the reactants, C and D are the products in the balanced chemical equation. On the basis of experimental studies of many reversible reactions, the Norwegian chemists Cato Maximillian Guldberg and Peter Waage proposed in 1864 that the concentrations in an equilibrium mixture are related by the following equilibrium equation,
`color{red}(K_c = ([ C] [D])/([A] [B]))` ...........(7.1)
where `color{red}(K_c)` is the equilibrium constant and the expression on the right side is called the equilibrium constant expression.
The equilibrium equation is also known as the law of mass action because in the early days of chemistry, concentration was called “active mass”. In order to appreciate their work better, let us consider reaction between gaseous `color{red}(H_2)` and `color{red}(I_2)` carried out in a sealed vessel at 731K.
`color{red}(underset(1 mol)(H_2 (g)) +underset(1 mol)(I_2 (g)) ⇌ underset(2 mol)(2HI (g)))`
Six sets of experiments with varying initial conditions were performed, starting with only gaseous `color{red}(H_2)` and `color{red}(I_2)` in a sealed reaction vessel in first four experiments (1, 2, 3 and 4) and only `color{red}(HI)` in other two experiments (5 and 6). Experiment 1, 2, 3 and 4 were performed taking different concentrations of `color{red}(H_2)` and / or `color{red}(I_2)`, and with time it was observed that intensity of the purple colour remained constant and equilibrium was attained. Similarly, for experiments 5 and 6, the equilibrium was attained from the opposite direction.
Data obtained from all six sets of experiments are given in Table 7.2.
It is evident from the experiments 1, 2, 3 and 4 that number of moles of dihydrogen reacted = number of moles of iodine reacted =
½ (number of moles of HI formed). Also, experiments 5 and 6 indicate that,
`color{red}([H_2 (g)]_(eq) = [I_2 (g) ]_(eq))`
Let us consider the simple expression
`color{red}([HI (g)]_(eq) // [H_2 (g) ]_(eq) [ I_2 (g)]_(eq))`
It can be seen from Table 7.3 that if we put the equilibrium concentrations of the reactants and products, the above expression is far from constant. However, if we consider the expression,
`color{red}([HI (g)]_(eq)^2 // H_2 [H_2 (g)]_(eq) [I_2 (g) ]_(eq))`
we find that this expression gives constant value (as shown in Table 7.3) in all the six cases. It can be seen that in this expression the power of the concentration for reactants and products are actually the stoichiometric coefficients in the equation for the chemical reaction. Thus, for the reaction `color{red}(H_2(g) + I_2(g) ⇌ 2HI(g))`, following equation 7.1, the equilibrium constant `color{red}(K_c)` is written as,
`color{red}(K_c = [HI (g) ]_(eq)^2 // [H_2 (g)]_(eq) [I_2 (g) ]_(eq))` .............(7.2)
Generally the subscript ‘eq’ (used for equilibrium) is omitted from the concentration terms. It is taken for granted that the
concentrations in the expression for `K_c` are equilibrium values. We, therefore, write.
`color{red}(K_c = [HI (g) ]^2 // [ H_2 (g) ] [I_2 (g) ])` ...........(7.3)
The subscript `color{red}(‘c’)` indicates that `color{red}(K_c)` is expressed in concentrations of `color{red}(mol L^(–1))`.
`color{purple}(✓✓)color{purple} " DEFINITION ALERT"`
At a given temperature, the product of concentrations of the reaction products raised to the respective stoichiometric coefficient in the balanced chemical equation divided by the product of concentrations of the reactants raised to their individual stoichiometric coefficients has a constant value. This is known as the Equilibrium Law or Law of Chemical Equilibrium.
The equilibrium constant for a general reaction,
` color{red}(a A + b B ⇌ c C + d D)`
is expressed as,
`color{red}(K_c = [C]^c [D]^d // [A]^a [B]^b)` ...........(7.4)
where [A], [B], [C] and [D] are the equilibrium concentrations of the reactants and products. Equilibrium constant for the reaction,
`color{red}(4NH_3 (g) +5O_2 (g) ⇌ 4NO(g) +6H_2O (g)) ` is written as
`color{red}(K_c = [NO]^4 [H_2O]^6 // [NH_3]^4 [O_2]^5)`
Molar concentration of different species is indicated by enclosing these in square bracket and, as mentioned above, it is implied that these are equilibrium concentrations. While writing expression for equilibrium constant, symbol for phases (s, l, g) are generally ignored
Let us write equilibrium constant for the reaction, `color{red}(H_2(g) + I_2(g) ⇌ 2HI(g))` .......................... (7.5)
as `color{red}(K_c = [HI]^2 // [H_2] [I_2] = x)` ....................(7.6)
The equilibrium constant for the reverse reaction, `color{red}(2HI(g) ⇌ H_2(g) + I_2(g)),` at the same temperature is,
`color{red}(K_c = [H_2] [I_2] // [HI]^2 = 1/x = 1/K_c)` ...................(7.7)
Thus `color{red}(K_c = 1/K_c)` .....................(7.8)
Equilibrium constant for the reverse reaction is the inverse of the equilibrium constant for the reaction in the forward
direction.
If we change the stoichiometric coefficients in a chemical equation by multiplying throughout by a factor then we must make
sure that the expression for equilibrium constant also reflects that change. For example, if the reaction (7.5) is written as,
`color{red}(1/2 H_2 (g) +1/2 I_2 (g) ⇌ HI (g))` ............(7.9)
the equilibrium constant for the above reaction is given by
`color{red}(K_c^('') = [HI] // [H_2]^(1/2) [I_2]^(1/2) = { [ HI]^2 // [ H_2] [I_2]}^(1/2))`
` =color{red}( x^(1/2) = K_c^(1/2))` ...............(7.10)
On multiplying the equation (7.5) by n, we get
`color{red}(nH_2 (g) + n I_2 (g) ⇌ 2 n HI (g)) ` ........(7.11)
Therefore, equilibrium constant for the reaction is equal to `color{red}(K_c^n)`. These findings are summarised in Table 7.4. It should be noted
that because the equilibrium constants `color{red}(K_c)` and `color{red}(K_c^('))` have different numerical values, it is important to specify the form of the balanced chemical equation when quoting the value of an equilibrium constant.
Let us consider a general reversible reaction:
`color{red}(A+B ⇌ C+ D)`
where A and B are the reactants, C and D are the products in the balanced chemical equation. On the basis of experimental studies of many reversible reactions, the Norwegian chemists Cato Maximillian Guldberg and Peter Waage proposed in 1864 that the concentrations in an equilibrium mixture are related by the following equilibrium equation,
`color{red}(K_c = ([ C] [D])/([A] [B]))` ...........(7.1)
where `color{red}(K_c)` is the equilibrium constant and the expression on the right side is called the equilibrium constant expression.
The equilibrium equation is also known as the law of mass action because in the early days of chemistry, concentration was called “active mass”. In order to appreciate their work better, let us consider reaction between gaseous `color{red}(H_2)` and `color{red}(I_2)` carried out in a sealed vessel at 731K.
`color{red}(underset(1 mol)(H_2 (g)) +underset(1 mol)(I_2 (g)) ⇌ underset(2 mol)(2HI (g)))`
Six sets of experiments with varying initial conditions were performed, starting with only gaseous `color{red}(H_2)` and `color{red}(I_2)` in a sealed reaction vessel in first four experiments (1, 2, 3 and 4) and only `color{red}(HI)` in other two experiments (5 and 6). Experiment 1, 2, 3 and 4 were performed taking different concentrations of `color{red}(H_2)` and / or `color{red}(I_2)`, and with time it was observed that intensity of the purple colour remained constant and equilibrium was attained. Similarly, for experiments 5 and 6, the equilibrium was attained from the opposite direction.
Data obtained from all six sets of experiments are given in Table 7.2.
It is evident from the experiments 1, 2, 3 and 4 that number of moles of dihydrogen reacted = number of moles of iodine reacted =
½ (number of moles of HI formed). Also, experiments 5 and 6 indicate that,
`color{red}([H_2 (g)]_(eq) = [I_2 (g) ]_(eq))`
Let us consider the simple expression
`color{red}([HI (g)]_(eq) // [H_2 (g) ]_(eq) [ I_2 (g)]_(eq))`
It can be seen from Table 7.3 that if we put the equilibrium concentrations of the reactants and products, the above expression is far from constant. However, if we consider the expression,
`color{red}([HI (g)]_(eq)^2 // H_2 [H_2 (g)]_(eq) [I_2 (g) ]_(eq))`
we find that this expression gives constant value (as shown in Table 7.3) in all the six cases. It can be seen that in this expression the power of the concentration for reactants and products are actually the stoichiometric coefficients in the equation for the chemical reaction. Thus, for the reaction `color{red}(H_2(g) + I_2(g) ⇌ 2HI(g))`, following equation 7.1, the equilibrium constant `color{red}(K_c)` is written as,
`color{red}(K_c = [HI (g) ]_(eq)^2 // [H_2 (g)]_(eq) [I_2 (g) ]_(eq))` .............(7.2)
Generally the subscript ‘eq’ (used for equilibrium) is omitted from the concentration terms. It is taken for granted that the
concentrations in the expression for `K_c` are equilibrium values. We, therefore, write.
`color{red}(K_c = [HI (g) ]^2 // [ H_2 (g) ] [I_2 (g) ])` ...........(7.3)
The subscript `color{red}(‘c’)` indicates that `color{red}(K_c)` is expressed in concentrations of `color{red}(mol L^(–1))`.
`color{purple}(✓✓)color{purple} " DEFINITION ALERT"`
At a given temperature, the product of concentrations of the reaction products raised to the respective stoichiometric coefficient in the balanced chemical equation divided by the product of concentrations of the reactants raised to their individual stoichiometric coefficients has a constant value. This is known as the Equilibrium Law or Law of Chemical Equilibrium.
The equilibrium constant for a general reaction,
` color{red}(a A + b B ⇌ c C + d D)`
is expressed as,
`color{red}(K_c = [C]^c [D]^d // [A]^a [B]^b)` ...........(7.4)
where [A], [B], [C] and [D] are the equilibrium concentrations of the reactants and products. Equilibrium constant for the reaction,
`color{red}(4NH_3 (g) +5O_2 (g) ⇌ 4NO(g) +6H_2O (g)) ` is written as
`color{red}(K_c = [NO]^4 [H_2O]^6 // [NH_3]^4 [O_2]^5)`
Molar concentration of different species is indicated by enclosing these in square bracket and, as mentioned above, it is implied that these are equilibrium concentrations. While writing expression for equilibrium constant, symbol for phases (s, l, g) are generally ignored
Let us write equilibrium constant for the reaction, `color{red}(H_2(g) + I_2(g) ⇌ 2HI(g))` .......................... (7.5)
as `color{red}(K_c = [HI]^2 // [H_2] [I_2] = x)` ....................(7.6)
The equilibrium constant for the reverse reaction, `color{red}(2HI(g) ⇌ H_2(g) + I_2(g)),` at the same temperature is,
`color{red}(K_c = [H_2] [I_2] // [HI]^2 = 1/x = 1/K_c)` ...................(7.7)
Thus `color{red}(K_c = 1/K_c)` .....................(7.8)
Equilibrium constant for the reverse reaction is the inverse of the equilibrium constant for the reaction in the forward
direction.
If we change the stoichiometric coefficients in a chemical equation by multiplying throughout by a factor then we must make
sure that the expression for equilibrium constant also reflects that change. For example, if the reaction (7.5) is written as,
`color{red}(1/2 H_2 (g) +1/2 I_2 (g) ⇌ HI (g))` ............(7.9)
the equilibrium constant for the above reaction is given by
`color{red}(K_c^('') = [HI] // [H_2]^(1/2) [I_2]^(1/2) = { [ HI]^2 // [ H_2] [I_2]}^(1/2))`
` =color{red}( x^(1/2) = K_c^(1/2))` ...............(7.10)
On multiplying the equation (7.5) by n, we get
`color{red}(nH_2 (g) + n I_2 (g) ⇌ 2 n HI (g)) ` ........(7.11)
Therefore, equilibrium constant for the reaction is equal to `color{red}(K_c^n)`. These findings are summarised in Table 7.4. It should be noted
that because the equilibrium constants `color{red}(K_c)` and `color{red}(K_c^('))` have different numerical values, it is important to specify the form of the balanced chemical equation when quoting the value of an equilibrium constant.