Question

Which is not a buffer solution?
A. $N{H_4}Cl$ +$N{H_4}OH$
B. $C{H_3}COOH$ +$C{H_3}COONa$
C. $C{H_3}COON{H_4}$
D. $N{H_4}N{O_3}$

Answer 1

Hint: A buffer solution either is a mixture of a weak acid and its salt with strong base or a mixture of a weak base and its salt with strong acid. So the solution which is not satisfying this condition is not a buffer solution.

Step by step answer: Let us first look at the buffer solution.
Buffer Solution: So buffer solution is an aqueous solution consisting of a mixture of a weak acid and its conjugate base, or vice versa. Its pH changes very little when a small amount of strong base or acid is added to it.
So according to the above definition analyzing all the options
-1st is $N{H_4}Cl$ +$N{H_4}OH$ , in which $N{H_4}Cl$acts as a strong electrolyte and $N{H_4}OH$acts as a weak electrolyte. Its dissociation is further suppressed by common ion $NH_4^ + $ provided by $N{H_4}Cl$in the solution. This solution acts as a basic buffer and maintains its pH around 9.25. It is capable of resisting the change in pH in addition to a small amount of acid or alkali. So it is a buffer solution.
-2nd is $C{H_3}COOH$ +$C{H_3}COONa$, which also acts as a buffer solution in the same way.
-3rd is $C{H_3}COON{H_4}$which is not satisfying the above conditions so it is not a buffer solution so that is why it is a correct option.
-4th is$N{H_4}N{O_3}$, When in solution, ammonium nitrate will dissociate very well into $NH_4^ + $ and $NO_3^ - $ . The ammonium ion is a Bronsted-Lowry acid in that it will donate a ${H^ + }$ to solution. In turn, this forms ammonia. Ammonia being a weak base will accept hydrogen ions from the solution, but not all of them. In conclusion, the solution will form equilibrium where the pH is slightly acidic due to ${H^ + }$ ions being loose in solution. So according to the definition it is a buffer solution.

So the correct option is C.

Note: Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications. In nature, there are many systems that use buffering for pH regulation. For example, the bicarbonate buffering system is used to regulate the pH of blood.