A 0.1 M acetic acid solution is titrated against 0.1 M-NaOH solution. What would be the differen

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A 0.1 M acetic acid solution is titrated against 0.1 M-NaOH...

A $0.1 \mathrm{M}$ acetic acid solution is titrated against $0.1 \mathrm{M}-\mathrm{NaOH}$ solution. What would be the difference in $\mathrm{pH}$ between $1 / 4$ and $3 / 4$ stages of neutralization of the acid?
(a) $2 \log (0.75)$
(b) $2 \log (0.25)$
(c) $\log 3$
(d) $2 \log 3$

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Key Concepts

Logarithmic pH Scale

The pH scale is logarithmic, meaning it represents the ratio of hydrogen ion concentration on a logarithmic scale. This property is essential for understanding how small changes in ion concentration can result in significant pH shifts and why expressions involving logarithms, like those in the Henderson-Hasselbalch equation, are used in acid-base chemistry.

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation provides a practical means to calculate the pH of a buffer solution by relating it to the acid dissociation constant (pKa) and the ratio of the concentrations of the conjugate base to the weak acid. This equation is crucial in titration scenarios where partial neutralization creates a buffer, allowing for the prediction of pH at various stages of the titration process.

Acid-Base Titration

Acid-base titration is a technique in analytical chemistry where a solution of known concentration (the titrant) is added to a solution of unknown concentration until the reaction reaches its endpoint. This method is fundamental for determining the concentration of an analyte and involves monitoring changes in pH as the titrant is added, especially in reactions involving weak acids or weak bases and their strong counterparts.

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