Question

Which of the following mixture can act as a buffer solution? (1) I litre of \(0.2 \mathrm{~N} \mathrm{CH}_{3} \mathrm{COOH}+1\) litre of \(0.2 \mathrm{~N} \mathrm{NaOH}\) (2) I litre of \(0.1 \mathrm{~N} \mathrm{CH}_{3} \mathrm{COOH}+1\) litre of \(0.2 \mathrm{~N} \mathrm{NaOH}\) (3) 1 litre of \(0.2 \mathrm{~N} \mathrm{CH}_{3} \mathrm{COOH}+1\) litre of \(0.1 \mathrm{~N} \mathrm{NaOH}\) (4) All the above

Short answer

Options 1 and 3 can act as buffer solutions.

Step-by-step solution

- Define a Buffer Solution

A buffer solution is a solution that can resist significant changes in pH when small amounts of an acid or a base are added. It is typically made of a weak acid and its conjugate base, or a weak base and its conjugate acid.

- Analyze Option 1

For Option 1: Mix 1 litre of 0.2N \(\text{CH}_{3}\text{COOH}\) (acetic acid) with 1 litre of 0.2N NaOH (sodium hydroxide).The NaOH will completely react with \(\text{CH}_{3}\text{COOH}\) to produce \(\text{CH}_{3}\text{COONa}\) (sodium acetate) and water:\[ \text{CH}_{3}\text{COOH} + \text{NaOH} \rightarrow \text{CH}_{3}\text{COONa} + \text{H}_{2}\text{O} \]Since the quantities are equal, the resulting solution will contain the weak acid \(\text{CH}_{3}\text{COOH}\) and its conjugate base \(\text{CH}_{3}\text{COO}^-\). This can function as a buffer.

- Analyze Option 2

For Option 2: Mix 1 litre of 0.1N \(\text{CH}_{3}\text{COOH}\) with 1 litre of 0.2N NaOH.The limiting reactant is \(\text{CH}_{3}\text{COOH}\). Therefore, all of the \(\text{CH}_{3}\text{COOH}\) will react with NaOH to form \(\text{CH}_{3}\text{COONa}\).Since one mole of NaOH reacts with one mole of \(\text{CH}_{3}\text{COOH}\), after reaction, there will be an excess of NaOH in the solution. This will not be a buffer solution.

- Analyze Option 3

For Option 3: Mix 1 litre of 0.2N \(\text{CH}_{3}\text{COOH}\) with 1 litre of 0.1N NaOH.Since NaOH is the limiting reactant, it will be completely consumed in the reaction, forming \(\text{CH}_{3}\text{COONa}\). This will leave excess \(\text{CH}_{3}\text{COOH}\).The resulting solution contains both \(\text{CH}_{3}\text{COOH}\) and \(\text{CH}_{3}\text{COO}^-\), forming a buffer solution.

- Conclusion

From the analysis:- Option 1 forms a buffer solution.- Option 2 does not form a buffer solution (excess NaOH).- Option 3 forms a buffer solution.Therefore, Options 1 and 3 can act as buffer solutions.

Key concepts

Weak Acid

A weak acid is an acid that partially dissociates in solution. This means that not all of its molecules release hydrogen ions (H\textsuperscript{+}) when dissolved in water. Most remain intact. A common example of a weak acid is acetic acid (\text{CH}_{3}\text{COOH}), which is found in household vinegar.
When a weak acid like acetic acid is present in a solution, it establishes an equilibrium between the intact acid molecules and the ions they release:
\[ \text{CH}_{3}\text{COOH} \rightleftharpoons \text{CH}_{3}\text{COO}^{-} + \text{H}^{+} \]
This equilibrium is crucial because it allows the solution to resist changes in pH. When an acid or base is added, the equilibrium shifts to minimize the change in H\textsuperscript{+} concentration.
In the given exercise, the weak acid involved is acetic acid. Its partial dissociation is what helps in maintaining the pH balance in buffer solutions.

Conjugate Base

The conjugate base is what's left after an acid has donated a proton (H\textsuperscript{+}). In the case of a weak acid, its conjugate base is important for buffering. For acetic acid (\text{CH}_{3}\text{COOH}), the conjugate base is acetate ion (\text{CH}_{3}\text{COO}^{-}).
The conjugate base plays a pivotal role in neutralizing added acids. When an acid is added to the buffer solution, the conjugate base reacts with the H\textsuperscript{+} ions released by the acid to form the weak acid, which minimizes the pH change:
\[ \text{CH}_{3}\text{COO}^{-} + \text{H}^{+} \rightarrow \text{CH}_{3}\text{COOH} \]
In the exercise, both Option 1 and Option 3 form buffer solutions because they contain the weak acid acetic acid and its conjugate base, acetate ion. This combination is what makes the solution capable of resisting changes in pH.

pH Resistance

Buffer solutions are superb at resisting changes in pH. This pH resistance is vital, especially in biological systems where enzymes and other biochemical processes require a stable pH to function correctly.
A buffer solution typically contains a mixture of a weak acid and its conjugate base. When small amounts of an acid or base are added, the buffer solution reacts to prevent significant changes in pH.
For example, if an acid (H\textsuperscript{+}) is added to a buffer solution containing acetic acid and acetate, the acetate ions will neutralize the added H\textsuperscript{+}, forming more acetic acid:
\[ \text{CH}_{3}\text{COO}^{-} + \text{H}^{+} \rightarrow \text{CH}_{3}\text{COOH} \]
Conversely, if a base like OH\textsuperscript{-} is added, the acetic acid will react with the OH\textsuperscript{-} to form water and acetate ions:
\[ \text{CH}_{3}\text{COOH} + \text{OH}^{-} \rightarrow \text{CH}_{3}\text{COO}^{-} + \text{H}_{2}\text{O} \]
This ability to counteract added acids and bases is what gives buffer solutions their pH resistance, ensuring stability in various chemical and biological contexts. Options 1 and 3 in the exercise exemplify this property well, as they effectively form buffers that can stabilize pH.