Quick questions on Acid-base equilibria - WJEC A-Level Chemistry

A weak acid such as ethanoic acid only partly dissociates, so its $[\text{H}^+]$ is much smaller than its concentration. The standard approximation is $[\text{H}^+] = \sqrt{K_a \times c}$, valid because dissociation is small enough that $[\text{HA}]$ is roughly the original concentration and $[\text{H}^+] = [\text{A}^-]$. This is why a $0.1$ mol dm$^{-3}$ solution of a weak acid has a pH around $3$, while the same concentration of a strong acid has a pH of $1$. The smaller the $K_a$ (the larger the $pK_a$), the weaker the acid and the higher the pH at a given concentration.

A pH titration curve has four features to identify: the starting pH (set by the acid strength), the gentle buffer region (for a weak acid plus base), the steep vertical jump at the equivalence point, and the plateau toward the pH of excess titrant. The midpoint of the buffer region gives $\text{pH} = pK_a$, a quick way to read off the acid's strength. The indicator must change colour entirely within the steep region: methyl orange for strong-acid against weak-base titrations, phenolphthalein for weak-acid against strong-base.

The hydrogencarbonate buffer keeps blood pH near $7.4$; a shift of even $0.1$ unit is dangerous, illustrating the biological importance of buffering. Indicator choice in titrations. A strong acid with a weak base has its equivalence point below pH $7$, so methyl orange (not phenolphthalein) is the correct indicator, read directly from the titration curve.

Calculate the pH of $0.0100$ mol dm$^{-3}$ hydrochloric acid. [1 mark]

State the expression for the acid dissociation constant $K_a$ of a weak acid HA. [1 mark]

Name the two components needed to make an acidic buffer. [1 mark]

State the pH at the midpoint of the buffer region of a weak-acid titration. [1 mark]

Calculate the pH of a $0.100$ mol dm$^{-3}$ weak acid with $K_a = 1.0 \times 10^{-5}$ mol dm$^{-3}$. [2 marks]