Buffer Solution MCQ - Practice Questions & Answers

Solve by difficulty

Which of the following statement(s) is/are correct?
(A) The pH of $\mathrm{1\times 10^{-8} \ M \ HCl}$ solution is 8.
(B) The conjugate hase of $\mathrm{H}_2 \mathrm{PH}_4^{-}$ is $\mathrm{HPO}_4^{2-}$
(C) $\mathrm{K}_{\mathrm{w}}$ increases with increase in temperature.
(D) When a solution of a weak monoprotic acid is titrated against a strong base at half neutralisation point.
$\mathrm{pH}=\frac{1}{2} \mathrm{pK}_{\mathrm{a}}$
Choose the correct answer from the options given below.

(A), (B). (C)

(A), (D)

(B), (C)

(B), (C). (D)

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Concepts Covered - 7

Buffer Solution

A solution whose pH does not change very much when H⁺(H₃O⁺) or OH^- are added to it is referred to as a buffer solution.
A buffer solution is prepared by mixing a weak and its salt having common anion(i.e HA + HB forms an acidic buffer) or a weak base and its salt having common cation(i.e BOH + BA forms a basic buffer).
It can be prepared to have a desired value of pH by controlling the amounts of acids and their salts in case of acidic buffer and of bases and their salts in basic buffer.

$\mathrm{Acidic\: buffer:\: \: \: \: \mathrm{CH}_{3} \mathrm{COOH}+\mathrm{CH}_{3} \mathrm{COONa}}$
$\mathrm{Basic\: buffer:\: \: \: \: \: \: \mathrm{NH}_{4} \mathrm{OH}+\mathrm{NH}_{4} \mathrm{Cl}}$

Consider an acidic buffer containing an acid HA and say common ions A^-. Now any H⁺added to this solution within certain limits are neutralized by A^- ions as:
$\mathrm{H}^{+}+\mathrm{A}^{-} \rightleftharpoons \mathrm{HA}$
While the addition of OH^-ions externally (within certain limits) are neutralised by acid HA as:
$\mathrm{HA}\: +\, \mathrm{OH}^{-} \rightleftharpoons \mathrm{H}_{2} \mathrm{O}+\mathrm{A}^{-}$
Hence in both the cases, effect of addition of H⁺or OH^- is almost compensated for (i.e. pH almost remains constant).

Such a system (may be acidic or basic) finds enormous use not only in industrial processes but also most importantly in biological reactions. Like the pH of normal blood is 7.4 and for good health and even for the survival, it should not change below 7.1 or greater than 7.7, the body maintains it through a buffer system made of carbonate and bicarbonate ions and H₂PO₄^- and HPO₄²^-. Similarly, the pH of gastric juice is kept constant in order to operate good digestive functions.

Types of Buffers

Buffer solutions are obtainned if the acids and bases are mixed in different amounts (equivalents).

Buffer solutions are those, which resist a change in pH upon addition of small amount of acid or base. this does not mean that the pH will not change, and all it means is that the pH change would be less than the change that would have occurred had it not been a buffer

Buffer solutions can be classified into three types:

(1) Acidic Buffer Solutions

Acidic buffer solutions are the solutions that are made from a weak acid and one of its salt with a strong base.

For example: Solution of $\mathrm{CH_3COOH}$ and $\mathrm{CH_3COONa}$

It is to be noted that the pH of an acidic buffer may not be always less than 7. It depends upon the Ka values of the acid and also the concentration of the acid and the salt

(2) Basic Buffer Solutions

Basic buffer solutions are the solutions that are made from a weak base and one of its salt with a strong acid.

For example: Solution of $\mathrm{NH_4OH}$ and $\mathrm{NH_4Cl}$

It is to be noted that the pH of an basic buffer may not be more less than 7. It depends upon the Kb values of the base and also the concentration of the salt and base

(3) Simple Buffer Solutions

Simple buffer solutions are the solutions that are made from the salt of a weak acid and weak base.

For example: Solution of $\mathrm{CH_3COONH_4}$

It is to be noted that the pH of simple buffer may be less than, greater than or equal to 7. It depends upon the Ka and Kb values of the acid and the base.

Buffer Action:

A buffer solution resists a change in its pH on addition of small amount of acid or base. This is because there is one component which can neutralise the acid and the other component can neutralise the base

e.g $\mathrm{CH_3COOH\ and\ CH_3COONa}$

When small amount of base is added, then it is the acid which neutralises it

$\mathrm{OH^- +\ CH_3COOH \rightleftharpoons H_2O + CH_3COO^- }$

When small amount of acid is added, then it is the acetate ion which neutralises it

$\mathrm{HCl +\ CH_3COO^- \rightleftharpoons CH_3COOH + Cl^-}$

as neutralisation occurs, the $\mathrm{[H^+] \ or \ [OH^-]}$ does not alter much in the solution and pH change is almost negligible

Cases which are not a buffer solution

(1) Solutions of Strong Acid and its salt e.g. $\mathrm{H_2SO_4\ and\ KHSO_4}$

(2) Solutions of Strong Base and its salt e.g. $\mathrm{NaOH \ and\ NaCl}$

For a solution to be classified as a buffer solution, there must be one weak acid or base and its respective conjugate base or acid

Calculating pH of a Buffer Solution(acidic)

When a solution contains CH₃COOH and CH₃COONa, then the following equilibrium will be established:
$\\\mathrm{CH_{3}COOH\: \rightleftharpoons \: CH_{3}COO^{-}\: +\: H^{+}\: \: \: \: \: \: \: \: \: \: \: \: \: }$

The equilibrium equation for the given system can be calculated using the following equation:
$\\\mathrm{K_{a}\: =\: \frac{[CH_{3}COO^{-}][H^{+}]}{[CH_{3}COOH]}= \frac{[Salt][H^{+}]}{[Acid]}} \\\\\mathrm{[CH_{3}COO^{-}]\: is\: the \: concentration \: of\: salt\:}\\\mathrm{[CH_{3}COOH]\: is\: the \: initial\: concentration \: of\: acid}$

Rearranging the above equation, we get

$\\\mathrm{[H^{+}]\: =\: K_{a}\frac{[Acid]}{[Salt]}}\\\\\mathrm{-log_{10}[H^{+}]\: =\: -log_{10}K_{a}\: -\: log_{10}[Acid]\: +\: log_{10}[Salt]}\\\\\mathrm{pH\: =\: pK_{a}\: +\: log_{10}\frac{[Salt]}{Acid}}$

This equation is also known as the Henderson-Hasselbalch equation.

Some examples

Find the pH of a solution having 0.1M CH₃COOH(K_a = 10^-5) and 0.2M CH₃COONa.

We know that pH of a solution is given as:

$\\\mathrm{pH\: =\: pK_{a}\: +\: log_{10}\frac{[Salt]}{Acid}}\\\\\mathrm{Thus,\: pH\: =\: -log_{10}K_{a}\: +\: log_{10}\frac{[0.2]}{[0.1]}}\\\\\mathrm{\Rightarrow\: pH\: = -log_{10}10^{-5}\: +\: log_{10}2}\\\\\mathrm{\Rightarrow\: pH\: =\: 5\: +\: 0.30\: =\: 5.30}$
Find the pH of a solution containing 0.25 moles of HCN(K_a = 10^-5) and 0.10 moles of NaCN present in 1 litre solution.

We know that pH of a solution is given as:
$\\\mathrm{pH\: =\: pK_{a}\: +\: log_{10}\frac{[Salt]}{Acid}}\\\\\mathrm{Thus,\: pH\: =\: -log_{10}K_{a}\: +\: log_{10}\frac{[Salt]}{[Acid]}}\\\\\mathrm{\Rightarrow\: pH\: = -log_{10}10^{-5}\: +\: log_{10}\frac{0.10}{0.25}}\\\\\mathrm{\Rightarrow\: pH\: = 5\: +\: log_{10}\frac{2}{5}}\\\\\mathrm{\Rightarrow\: pH\: =\: 5\: -\: 0.39\: =\: 4.6}$

Working of Acidic Buffer

Acidic buffer solutions are the solutions that are made from a weak acid and one of its salt mainly sodium salt.

$\\\mathrm{CH_{3}COOH\: \rightleftharpoons \: CH_{3}COO^{-}\: +\: H^{+}}\\\\\mathrm{CH_{3}COONa\: \rightleftharpoons \: CH_{3}COO^{-}\: +\: Na^{+}}$

On addition of acid

$\mathrm{H^{+}\: +\: CH_{3}COO^{-}(solution)\: \rightarrow CH_{3}COOH}$

Although on the addition of acid concentration of CH₃COOH increases, so it wants to go in forward direction but due to common ion effect CH₃COOH cannot dissociate back.
CH₃COO^- concentration decreases but we have abundant amount of CH₃COO^-. So decrease is negligible.
On addition of base

$\mathrm{OH^{-}\: +\: H^{+}(from\: solution)\: \rightarrow \: H_{2}O}$
In this case, [H⁺] concentration decreases and CH₃COOH goes in forward direction to dissociate into H⁺ so as to restore the concentration of [H⁺]

Buffer Capacity

The property of a buffer solution to resist a change in pH is known as buffer capacity. It is defined as the number of moles of acids or bases added in one litre of solution to change the pH by unity, i.e. Thus, buffer capacity is given as:

$\\\mathrm{Buffer\: capacity\: =\: \frac{Mole\: of\: acid\: or\: base\: added\: to\: 1\: litre \: of \: buffer}{Change \: in \: pH}}\\\\\mathrm{\: \: \: \: \: \: \: \: \: \: \: \: \: \: \: \: \: \: \: \: \: \: \: \: \: \: \: \: \: \: \: =\: \frac{n}{\Delta pH}}$

Note: The greater is the buffer capacity, the greater is its capacity to resist change in pH

Salient Features of Buffer Solutions

It has definite pH.
Its pH does not change on standing.
Its pH does not change on dilution.
Its pH does not change significantly on the addition of the small amount of acid or base.
The pH of the buffer solution depends upon pK_a and on the relative molar amount of weak acid and its conjugate base.
Buffer solutions are used in:

Qualitative analysis of mixture, for example, removal of phosphate.
Quantitative analysis of estimations.
Industrial processes such as manufacture of paper, dyes, inks, paints, drugs, etc.
Digestion of food.
Preservation of foods and fruits.
Agriculture and dairy products preservations.

Basic Buffers

Basic buffer solution contains a weak base and its salt with strong acid. Some examples of basic buffers are:

NH₄OH + NH₄Cl
NH₄OH + (NH₄)₂SO₄
CH₃-NH₂ + [CH₃-NH₃⁺]Cl^-

The pH of the basic buffer is given as:

$\mathrm{pOH\: =\: pK_{b}\: +\: log_{10}\frac{[Salt]}{[Base]}}$

We already know that pH = 14 - pOH. Thus can be calculated using this equation.

For example: basic buffer we have:

$\\\mathrm{NH_{4}OH\: \rightleftharpoons \: NH_{4}^{+}\: +\: Cl^{-}}\\\mathrm{NH_{4}Cl\: \rightarrow \: NH_{4}^{+}\: Cl^{-}\: \: \: \: \: \: (Strong\: electrolyte)}\\\\\mathrm{Thus, K_{b}\: =\: \frac{[NH_{4}^{+}][OH^{-}]}{[NH_{4}OH]}}$

In this system:

[NH₄OH]: Initial concentration of [NH₄OH] is taken as at equilibrium negligible dissociation of NH₄OH is there because of common-ion effect.
[NH₄⁺]: The concentration of NH₄OH is mostly from 100% dissociation of NH₄Cl.

Again, as we know:

$\\\mathrm{K_{b}\: =\: \frac{[Salt][OH^{-}]}{Base}}\\\\\mathrm{Thus,\: [OH^{-}]\: =\: K_{b}\frac{[Base]}{[Salt]}}\\\\\mathrm{Using\: log\: on\: both\: sides,\: we\: get:}\\\\\mathrm{-log_{10}[OH^{-}]\: =\: -log_{10}K_{b}\: +\: log_{10}[Salt]\: -\: log_{10}[Base]}\\\\\mathrm{\mathbf{Hence},\: pOH\: =\: pK_{b}\: +\: log_{10}\frac{[Salt]}{[Base]}}$

Action of Basic Buffer

Basic buffer solution contains equimolar quantities of a weak base and its salt with strong acid. For example: ammonium hydroxide i.e. NH₄OH and ammonium chloride i.e NH₄Cl.

On Adding Acid: H⁺release and combines with OH^- of base.

On Adding Base: OH^-releases and combines with NH₄⁺ of salt.

On adding acid to the basic buffer, its H⁺ ions react with OH^- ions of the base and forms H₂O. Thus, in this case, solution feels that its [OH^-] has decreased, thus to neutralize this effect, NH₄OH dissociate in small amounts and gives [OH^-] so as to restore concentration of [OH^-]
On adding base to the basic buffer, its [OH^-] ions react with NH₄⁺ ions and forms NH₄OH. In this case, the solution feels that its NH₄OH concentration is increased. Thus, in this case, the reaction will not proceed forward because of common ion effect.