What is pH of a liter buffer solution containing 0.6 M benzoic acid and 0.07 M sodium benzoate?

step 1 Okay, so we're given a pH of 3 .10 and a concentration of HA, which is equal to .200M. So we're

What is pH of a liter buffer solution containing 0.6 M...

What is $\mathrm{pH}$ of a liter buffer solution containing $0.6 \mathrm{M}$ benzoic acid and $0.07 M$ sodium benzoate? The dissociation constant $\left(\mathrm{K}_{\mathrm{a}}\right)$ for benzoic acid at $25^{\circ} \mathrm{C}$ is $6.3 \times 10^{-5}$.

Key Concepts

pH is a measure of the acidity or basicity of an aqueous solution, represented as the negative logarithm (base 10) of the hydrogen ion concentration. It provides insight into the acid-base balance of a solution and is crucial for understanding reaction equilibria and biochemical processes in the solution.

Buffer Solutions

Buffers are solutions that resist changes in pH upon the addition of small amounts of acid or base. They typically consist of a weak acid and its conjugate base (or a weak base and its conjugate acid) and are important for maintaining a stable pH in various chemical and biological systems.

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation is an approximation that relates the pH of a buffer solution to the pKa of the weak acid and the ratio of the concentrations of the conjugate base to the weak acid. This equation is a key tool for calculating the pH of buffer solutions and designing buffers with desired pH values.

Acid Dissociation Constant (Ka) and pKa

The acid dissociation constant (Ka) quantifies the strength of a weak acid in solution by describing its degree of ionization. The pKa, defined as the negative logarithm of Ka, is more convenient to use in calculations and indicates the pH at which the acid is half dissociated. Understanding Ka and pKa is essential for predicting the behavior of acids in buffer systems and other chemical equilibria.

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