Question

Pyridine (C5H5N) is a particularly foul-smelling substance used in manufacturing pesticides and plastic resins. Calculate the pH of a 0.135 M solution of pyridine (Kb = 1.4×10-9 ). 

Answer 1

Answer: pOH of the given 0.135M solution of Pyridine is 9.137

Explanation: Pyridine is a weak base and during its hydrolysis, pyridine forms pyridinium ion and releases hydroxide ions. Reaction follows:

                          $C_5H_5N+H_2O\rightleftharpoons C_5H_5NH^++OH^-$

at $t=0$                0.135M                     0               0

at $t=t_{eq}$               0.135-x                     x                x

$K_b$ can be written as:

$K_b=\frac{[C_5H_5NH^+][OH^-]}{[C_5H_5N]}$

$K_b=\frac{x.x}{(0.135-x)}$

$0.135(1.4\times 10^{-9})-x(1.4\times 10^{-9})=x^2$

$x^2+1.4x\times 10^{-9}-0.189\times 10^{-9}=0$

On solving the quadratic equation, we get

$x=\pm 1.37\times 10^{-5}$

Negative value is neglected, as it is concentration and concentration cannot be in negative.

Therefore, $x=1.37\times 20^{-5}$

$x=[OH^-]=1.37\times 10^{-5}$

pOH can be calculated as:

$pOH=-log[OH^-]$

$pOH=-log(1.37\times 10^{-5})$

pOH = 4.863

pH can be calculated by:

$pK_w=pH+pOH$

$pK_w=14$

$pH=14-4.863$

pH = 9.137