Question

Which statements are true about the titration represented by the following equation? $$\mathrm{H}^{+}(a q)+\mathrm{F}^{-}(a q) \rightleftharpoons \mathrm{HF}(a q)$$ (a) The net ionic equation could represent the titration of \(\mathrm{NaF}\) by \(\mathrm{H}_{2} \mathrm{SO}_{4}\) (b) The reaction could be considered to be that of a strong acid with a strong base. (c) At the equivalence point, the resulting solution can act as a buffer. (d) Diluting the solution obtained at half-neutralization with water does not change the \(\mathrm{pH}\) of the solution. (e) Phenolphthalein (pH range: 9-11) would be an effective indicator in this titration.

Short answer

a) The net ionic equation could represent the titration of NaF by H2SO4. b) The reaction could be considered that of a strong acid with a strong base. c) At the equivalence point, the resulting solution can act as a buffer. d) Diluting the solution obtained at half-neutralization with water does not change the pH of the solution. e) Phenolphthalein would be an effective indicator in this titration. Answer: a, c, and d are true.

Step-by-step solution

Statement (a)

Check if the net ionic equation could represent the titration of NaF by H2SO4. When NaF reacts with H2SO4, we have: $$\mathrm{NaF + H}_{2} \mathrm{SO}_{4} \rightarrow \mathrm{NaHSO}_{4} + \mathrm{HF}$$ The net ionic equation will be: $$\mathrm{H}^{+}(a q)+\mathrm{F}^{-}(a q) \rightleftharpoons \mathrm{HF}(a q)$$ Thus, statement (a) is true.

Statement (b)

Check if the reaction could be considered that of a strong acid with a strong base. The equation involves the reaction between hydrogen ions (H+), which usually come from strong acids, and fluoride ions (F-), which usually come from strong bases. However, the resulting product, HF, is a weak acid. Therefore, this reaction cannot be considered as a strong acid with a strong base. This statement (b) is false.

Statement (c)

Check if at the equivalence point, the resulting solution can act as a buffer. A buffer solution contains appreciable amounts of both a weak acid and its conjugate base, allowing it to resist changes in pH. Since the titration involves a weak acid (HF), at the equivalence point, we will have a mixture of the weak acid and its conjugate base (F-), which can indeed act as a buffer. Hence, statement (c) is true.

Statement (d)

Check if diluting the solution obtained at half-neutralization with water does not change the pH of the solution. At half-neutralization, the concentration of the weak acid is equal to the concentration of its conjugate base, making the buffered solution most effective at resisting changes in pH. Therefore, diluting with water does not significantly change the pH of the buffered solution. Hence, statement (d) is true.

Statement (e)

Check if phenolphthalein (pH range: 9-11) would be an effective indicator in this titration. Phenolphthalein changes color around the pH range of 9-11, while the pH of HF (a weak acid) will be less than 7. Therefore, phenolphthalein would not be an effective indicator in this titration, as it would only indicate a pH change above its range and not the actual endpoint of the titration. Statement (e) is false.

Key concepts

Titration of NaF by H2SO4

Understanding the titration of sodium fluoride (NaF) by sulfuric acid (H2SO4) is integral to grasping the basics of acid-base reactions. This specific titration can be represented by the net ionic equation:
\[\mathrm{H}^{+}(aq) + \mathrm{F}^{-}(aq) \rightleftharpoons \mathrm{HF}(aq)\]
The reaction showcases the process of adding H2SO4 to NaF. In this reaction, H2SO4 acts as a proton donor (acid), and NaF serves as a proton acceptor (base), leading to the formation of hydrogen fluoride (HF). It is essential to remember that despite H2SO4 being a strong acid, the product HF is a weak acid. This mismatch influences the titration curve, the choice of indicators, and the buffering capacity of the resulting solution.

Buffer Solution

When studying buffering systems, it's important to understand that a buffer solution maintains a relatively stable pH even when small amounts of acid or base are added. Such solutions consist of a weak acid and its conjugate base. In the given exercise, the equivalence point of the titration leads to the formation of a weak acid, HF, and its conjugate base, F-.
This combination can resist pH changes because if an acid is added, the base component (F-) can neutralize it; similarly, if a base is added, the acid component (HF) can counteract the change. Thus, at the equivalence point, the resulting solution forms an effective buffer. This property is exceptionally relevant in the biological and chemical systems where pH stability is crucial. The buffering capacity is greatest at the half-neutralization point, where the concentrations of the weak acid and its conjugate base are equal.

Acid-Base Indicators

Acid-base indicators are compounds that change color at a particular pH range, providing a visible signal for the acidity or alkalinity of a solution. They are crucial in determining the endpoint of a titration, which is ideally close to the equivalence point where equal amounts of acid and base have been mixed.
For the titration of NaF by H2SO4, phenolphthalein was considered as an indicator but was deemed ineffective because its color change occurs at a pH range of 9-11. The endpoint of this titration, however, will have a pH lower than 7 due to the presence of the weak acid HF. Therefore, an appropriate indicator for this titration would be one that changes color at a lower pH range, closer to the acidic pH levels expected at the endpoint. Indicators like methyl orange or bromothymol blue could be more suitable for this type of titration.