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

Step-by-step 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=([C_5H_5NH^+][OH^-])/([C_5H_5N])`

`K_b=(x.x)/((0.135-x))`

`0.135(1.4* 10^(-9))-x(1.4* 10^(-9))=x^2`

`x^2+1.4x* 10^(-9)-0.189* 10^(-9)=0`

On solving the quadratic equation, we get

`x=\pm 1.37* 10^(-5)`

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

Therefore,
`x=1.37* 20^(-5)`

`x=[OH^-]=1.37* 10^(-5)`

pOH can be calculated as:

`pOH=-log[OH^-]`

`pOH=-log(1.37* 10^(-5))`

pOH = 4.863

pH can be calculated by:

`pK_w=pH+pOH`

`pK_w=14`

`pH=14-4.863`

pH = 9.137