Question

How does titration of a strong, monoprotic acid with a strong base differ from titration of a weak, monoprotic acid with a strong base with respect to the following: (a) quantity of base required to reach the equivalence point, (b) \(\mathrm{pH}\) at the beginning of the titration, \((\mathrm{c}) \mathrm{pH}\) at the equivalence point, (d) \(\mathrm{pH}\) after addition of a slight excess of base, (e) choice of indicator for determining the equivalence point?

Short answer

In both strong and weak acid titrations, the quantity of base required to reach the equivalence point is the same. The pH at the beginning of the titration is lower for a strong acid as it is completely ionized, whereas a weak acid is partially ionized. At the equivalence point, the pH is neutral (7) for strong acid-strong base titration, and basic (greater than 7) for weak acid-strong base titration. After adding a slight excess of base, the pH is slightly alkaline for both titrations, but lower for a weak acid titration. The choice of indicator depends on the titration: phenolphthalein or bromothymol blue for strong acid-strong base titration, and phenolphthalein for weak acid-strong base titration.

Step-by-step solution

(a) Quantity of base required to reach the equivalence point

In both cases, the amount of base needed to reach the equivalence point is the same. At the equivalence point, the number of moles of the acid and the base must be equal for both strong and weak monoprotic acids. The balanced chemical equation is: \[HA + OH^- \rightarrow H_2O + A^-\]

(b) pH at the beginning of the titration

For a strong monoprotic acid like hydrochloric acid (HCl), the pH at the beginning of the titration is very low (0-2) because strong acids are completely ionized in water, increasing the concentration of hydrogen ions. For a weak monoprotic acid like acetic acid (CH3COOH), the pH at the beginning of the titration is higher (2-6) because the acid does not fully ionize in water.

(c) pH at the equivalence point

At the equivalence point, both the acid and the base have reacted, and the pH of the solution depends on the resulting salt. For strong acid-strong base titration, the resulting salt does not undergo any hydrolysis; hence, the pH will be neutral (7). For weak acid-strong base titration, the resulting salt does undergo hydrolysis, which produces hydroxide ions. This results in a basic pH (greater than 7) at the equivalence point.

(d) pH after addition of a slight excess of base

For both strong and weak acid titrations, the pH after adding a slight excess of base will be slightly alkaline due to the excess hydroxide ions from the strong base. However, the pH will be a bit lower for a weak acid titration compared to a strong acid titration because the weak acid is not completely dissociated.

(e) Choice of indicator for determining the equivalence point

The choice of indicator for a strong acid-strong base titration would be one with a pH transition range of about 4 to 10, such as phenolphthalein (transition range 8.2 - 10) or bromothymol blue (transition range 6-7.6). For a weak acid-strong base titration, the indicator must have a pH transition range above 7, as the equivalence point will have a pH greater than 7. Phenolphthalein (transition range 8.2 - 10) would be a suitable choice in this case.

Key concepts

Equivalence Point

In a titration process, the equivalence point is the stage where the quantity of titrant added is just enough to completely neutralize the analyte solution. Essentially, this means the moles of acid are equal to the moles of base present during the reaction.

For both strong monoprotic acids and weak monoprotic acids, the quantity of base required to reach the equivalence point remains the same. This is because the stoichiometry of the reaction demands equal moles of hydrogen ions (\(H^+\)) and hydroxide ions (\(OH^-\)) to form water (\(H_2O\)).

• This can be represented by the equation: \[ HA + OH^- \rightarrow H_2O + A^- \]
• "HA" is a general representation of any acid, where "A-" stands for the conjugate base following the reaction.

Acid-Base Reaction

An acid-base reaction involves the transfer of protons from the acid to the base. Titration is a method used to determine the unknown concentration of an acid or base by gradually adding a known concentration of another acid or base.

When a strong, monoprotic acid completely ionizes in water, it forms a high concentration of hydrogen ions (\(H^+\)), leading to a significantly low pH at the beginning of the titration. For example, hydrochloric acid (HCl) falls into this category. Conversely, a weak acid, such as acetic acid (\(CH_3COOH\)), only partially ionizes in water. This only leads to a moderate increase in hydrogen ion concentration; therefore, the starting pH is higher compared to strong acids.

• Strong acids are typically entirely dissociated in solution.
• Weak acids remain largely undissociated, indicating \(HA\) is present more than \(H^+\) and \(A^-\).

pH

pH is a measure of the acidity or basicity of a solution. It is an important factor in titration, as the pH changes markedly throughout the process.

At the equivalence point:

• In a titration involving a strong acid and a strong base, the solution will have a pH of 7 because the resulting salt (e.g., NaCl from \(HCl\) and \(NaOH\)) does not undergo hydrolysis.
• In a weak acid-strong base titration, the salt formed (e.g., \(CH_3COONa\) from \(CH_3COOH\) and \(NaOH\)) will hydrolyze to produce hydroxide ions, shifting the pH above 7, making the solution slightly basic.

After the addition of a slight excess of the base, the pH rises even more significantly in both cases due to the presence of excess hydroxide ions. However, the strength and nature of the acid play a crucial role in determining the specific pH at various stages of the titration.

Indicator Choice

The correct choice of indicator is vital for accurately determining the equivalence point in a titration. An indicator is a chemical that changes color at a certain pH range.

For strong acid versus strong base titrations, indicators such as bromothymol blue or phenolphthalein are suitable, as their transition range falls around the neutral pH, where the equivalence point occurs. These indicators change color close to a pH range of 7, making them viable for observing the exact moment of neutralization.

• Bromothymol blue transitions between pH 6 - 7.6.
• Phenolphthalein transitions between pH 8.2 - 10.

In the case of a weak acid-strong base titration, phenolphthalein is preferred, because the pH at the equivalence point is above 7. This indicator would visibly change color at the critical point where equilibrium is achieved during the reaction.