`=>` The equilibrium constant helps in predicting the direction in which a given reaction will proceed at any stage.
`=>` For this purpose, we calculate the reaction quotient `color{red}(Q)`. The reaction quotient, `color{red}(Q)` (`color{red}(Q_c)` with molar concentrations and `color{red}(Q_P)` with partial pressures) is defined in the same way as the equilibrium constant `color{red}(K_c)` except that the concentrations in `color{red}(Q_c)` are not necessarily equilibrium values. For a general reaction:
`color{red}(a A + b B ⇌ c C + d D)` ........(7.19)
`color{red}(Q_c = [C]^c [D]^d // [A]^a [B]^b)` ......(7.20)
Then, If `color{red}(Q_c > K_c)`, the reaction will proceed in the direction of reactants (reverse reaction). If `color{red}(Q_c < K_c),` the reaction will proceed in the
direction of the products (forward reaction).
If `color{red}(Q_c = K_c),` the reaction mixture is already at equilibrium.
`=>` Consider the gaseous reaction of `color{red}(H_2)` with `color{red}(I_2)`,
`color{red}(H_2(g) + I_2(g) ⇌ 2HI(g); K_c = 57.0)` at `700 K.`
`=>` Suppose we have molar concentrations `color{red}([H_2]t=0.10M, [I2]t = 0.20 M)` and `color{red}([HI]t = 0.40 M).` (the subscript `color{red}(t)` on the concentration symbols means that the concentrations were measured at some arbitrary time `color{red}(t)`, not necessarily at equilibrium).
Thus, the reaction quotient, `color{red}(Q_c)` at this stage of the reaction is given by,
`color{red}(Q_c = [HI]_t^(2) // [H_2]_t [I_2]_t = (0.40)^2// (0.10)×(0.20))`
` color{red}(= 8.0)`
Now, in this case, `color{red}(Q_c (8.0))` does not equal `color{red}(K_c (57.0)),` so the mixture of `color{red}(H_2(g), I_2(g))` and `color{red}(HI(g))` is not at equilibrium; that is, more `color{red}(H_2(g))` and `color{red}(I_2(g))` will react to form more `color{red}(HI(g))` and their concentrations will decrease till `color{red}(Q_c = K_c)`
`=>` The reaction quotient, `color{red}(Q_c)` is useful in predicting the direction of reaction by comparing the values of `color{red}(Q_c)` and `color{red}(K_c)`.
`=>` Thus, we can make the following generalisations concerning the direction of the reaction (Fig. 7.7)
• If `color{red}(Q_c < K_c)`, net reaction goes from left to right
• If `color{red}(Q_c > K_c),` net reaction goes from right to left.
• If `color{red}(Q_c = K_c),` no net reaction occurs.
`=>` The equilibrium constant helps in predicting the direction in which a given reaction will proceed at any stage.
`=>` For this purpose, we calculate the reaction quotient `color{red}(Q)`. The reaction quotient, `color{red}(Q)` (`color{red}(Q_c)` with molar concentrations and `color{red}(Q_P)` with partial pressures) is defined in the same way as the equilibrium constant `color{red}(K_c)` except that the concentrations in `color{red}(Q_c)` are not necessarily equilibrium values. For a general reaction:
`color{red}(a A + b B ⇌ c C + d D)` ........(7.19)
`color{red}(Q_c = [C]^c [D]^d // [A]^a [B]^b)` ......(7.20)
Then, If `color{red}(Q_c > K_c)`, the reaction will proceed in the direction of reactants (reverse reaction). If `color{red}(Q_c < K_c),` the reaction will proceed in the
direction of the products (forward reaction).
If `color{red}(Q_c = K_c),` the reaction mixture is already at equilibrium.
`=>` Consider the gaseous reaction of `color{red}(H_2)` with `color{red}(I_2)`,
`color{red}(H_2(g) + I_2(g) ⇌ 2HI(g); K_c = 57.0)` at `700 K.`
`=>` Suppose we have molar concentrations `color{red}([H_2]t=0.10M, [I2]t = 0.20 M)` and `color{red}([HI]t = 0.40 M).` (the subscript `color{red}(t)` on the concentration symbols means that the concentrations were measured at some arbitrary time `color{red}(t)`, not necessarily at equilibrium).
Thus, the reaction quotient, `color{red}(Q_c)` at this stage of the reaction is given by,
`color{red}(Q_c = [HI]_t^(2) // [H_2]_t [I_2]_t = (0.40)^2// (0.10)×(0.20))`
` color{red}(= 8.0)`
Now, in this case, `color{red}(Q_c (8.0))` does not equal `color{red}(K_c (57.0)),` so the mixture of `color{red}(H_2(g), I_2(g))` and `color{red}(HI(g))` is not at equilibrium; that is, more `color{red}(H_2(g))` and `color{red}(I_2(g))` will react to form more `color{red}(HI(g))` and their concentrations will decrease till `color{red}(Q_c = K_c)`
`=>` The reaction quotient, `color{red}(Q_c)` is useful in predicting the direction of reaction by comparing the values of `color{red}(Q_c)` and `color{red}(K_c)`.
`=>` Thus, we can make the following generalisations concerning the direction of the reaction (Fig. 7.7)
• If `color{red}(Q_c < K_c)`, net reaction goes from left to right
• If `color{red}(Q_c > K_c),` net reaction goes from right to left.
• If `color{red}(Q_c = K_c),` no net reaction occurs.