Equilibrium Test - 4

The reaction quotient Q. For a general reaction: a A + b B ⇌ c C + d D

 Qc = [C]^c [D]^d / [A]^a [B]^b

Then, If Qc > Kc ,

the reaction will proceed in the direction of reactants (reverse reaction).

 If Qc < Kc , the reaction will proceed in the direction of the products (forward reaction).

If Qc = Kc , the reaction mixture is already at equilibrium.

A mathematical expression of this thermodynamic view of equilibrium can be described by the following equation: ∆G = ∆G⁰ + RT lnQ

 where, G⁰ is standard Gibbs energy.

At equilibrium, when ∆G = 0 and Q = Kc ,

 ∆G = ∆G⁰ + RT ln K = 0

∆G⁰ = – RT lnK (7.22)

lnK = – ∆G⁰ / RT

Taking antilog of both sides, we get,

K = e^{– ∆G0 / RT}

Le Chatelier’s principle- It states that a change in any of the factors that determine the equilibrium conditions of a system will cause the system to change in such a manner so as to reduce or to counteract the effect of the change.

A catalyst increases the rate of the chemical reaction by making available a new low energy pathway for the conversion of reactants to products. It increases the rate of forward and reverse reactions that pass through the same transition state and does not affect equilibrium. Catalyst lowers the activation energy for the forward and reverse reactions by exactly the same amount. Catalyst does not affect the equilibrium composition of a reaction mixture. It does not appear in the balanced chemical equation or in the equilibrium constant expression.

Hydrogen ion by itself is a bare proton with very small size (~10–15 m radius) and intense electric field, binds itself with the water molecule at one of the two available lone pairs on it giving H₃O⁺ . This species has been detected in many compounds (e.g., H₃O⁺Cl ^– ) in the solid state.
In aqueous solution the hydronium ion is further hydrated to give species like H₅O⁺₂

Hydronium ion concentration in molarity is more conveniently expressed on a logarithmic scale known as the pH scale. The pH of a solution is defined as the negative logarithm to base 10 of the activity (_H+) of hydrogen ion. In dilute solutions (< 0.01 M), activity of hydrogen ion (H⁺ ) is equal in magnitude to molarity represented by [H⁺ ].
From the definition of pH, the following can be written, pH = – log H⁺= – log {[H⁺] / mol L^–1}
Thus, an acidic solution of HCl (10^-2 M) will have a pH = 2.
Similarly, a basic solution of NaOH having [OH^{​​​​​-} ] =10^-4 M and [H3O⁺ ] = 10^-10 M will have a pH = 10.
At 25 °C, pure water has a concentration of hydrogen ions, [H⁺ ] = 10^-7M. Hence, the pH of pure water is given as: pH = –log(10^–7) = 7

The solutions which resist change in pH on dilution or with the addition of small amounts of acid or alkali are called Buffer Solutions. Buffer solutions of known pH can be prepared from the knowledge of pKa of the acid or pK b of base and by controlling the ratio of the salt and acid or salt and base. A mixture of acetic acid and sodium acetate acts as buffer solution around pH 4.75 and a mixture of ammonium chloride and ammonium hydroxide acts as a buffer around pH 9.25.