List five common uses of acids and bases in our daily lives.

Step 1: Lewis structure for the sulfite ion $\mathrm{SO_3^{2-}}$.

Total valence electrons: S contributes 6, 3 × O contribute $3 \times 6 = 18$, +2 electrons from charge = $6 + 18 + 2 = 26$ electrons.

Place S in center bonded to 3 O atoms with single bonds (3 bonds = 6 electrons used).

Distribute remaining electrons: each O gets 3 lone pairs ($3 \times 6 = 18$ electrons), S gets 1 lone pair (2 electrons). Total: $6 + 18 + 2 = 26$.

Formal charges: S: $6 - 2 - \frac{6}{2} = +1$; each terminal O: $6 - 6 - \frac{2}{2} = -1$.

To minimize formal charges, convert one S–O single bond to double bond.

Double-bonded O: 2 lone pairs (4 electrons), FC = $6 - 4 - \frac{4}{2} = 0$.

Single-bonded O: FC = $-1$ each.

S: FC = $6 - 2 - \frac{8}{2} = 0$.

One resonance structure: $[\, \ce{^{-}O-S(=O)-O^{-}} \,]^{2-}$ (with lone pairs on O atoms, S has 1 lone pair).

Resonance delocalizes the double bond among 3 positions.

Step 2: Definition of Lewis acid.

A Lewis acid is a species that accepts an electron pair to form a coordinate covalent bond.

Examples: $\mathrm{HCl}$, $\mathrm{H_2SO_4}$, $\mathrm{BF_3}$.

Step 3: Definition of Lewis base.

A Lewis base is a species that donates an electron pair to form a coordinate covalent bond.

Examples: $\mathrm{NH_3}$, $\mathrm{OH^-}$.

Step 4: Common Lewis acids and bases.

Step 5: Formula for the conjugate base of $\mathrm{HSO_3^-}$.

$\mathrm{HSO_3^- \rightleftharpoons H^+ + SO_3^{2-}}$

Conjugate base: $\mathbf{SO_3^{2-}}$.

Step 6: Formula for the conjugate base of $\mathrm{HSO_4^-}$.

$\mathrm{HSO_4^- \rightleftharpoons H^+ + SO_4^{2-}}$

Conjugate base: $\mathbf{SO_4^{2-}}$.

Step 7: Formula for the conjugate acid of $\mathrm{SO_3^{2-}}$.

$\mathrm{SO_3^{2-} + H^+ \rightleftharpoons HSO_3^-}$

Conjugate acid: $\mathbf{HSO_3^-}$.

Step 8: Formula for the conjugate acid of $\mathrm{SO_4^{2-}}$.

$\mathrm{SO_4^{2-} + H^+ \rightleftharpoons HSO_4^-}$

Conjugate acid: $\mathbf{HSO_4^-}$.

Step 9: Amphiprotic substances.

Amphiprotic (amphoteric in Brønsted-Lowry sense) substances act as both acids (donate $\mathrm{H^+}$) and bases (accept $\mathrm{H^+}$).

$\mathrm{HSO_3^-}$:

As acid: $\mathrm{HSO_3^- \rightleftharpoons H^+ + SO_3^{2-}}$

As base: $\mathrm{HSO_3^- + H^+ \rightleftharpoons H_2SO_3}$

$\mathrm{HSO_4^-}$:

As acid: $\mathrm{HSO_4^- \rightleftharpoons H^+ + SO_4^{2-}}$

As base: $\mathrm{HSO_4^- + H^+ \rightleftharpoons H_2SO_4}$

Step 10: Verify different acid strengths (of $\mathrm{HSO_3^-}$ and $\mathrm{HSO_4^-}$).

$\mathrm{HSO_4^-}$ is the stronger acid ($\mathrm{p}K_a \approx 1.9$) compared to $\mathrm{HSO_3^-}$ ($\mathrm{p}K_a \approx 7.2$), as $\mathrm{SO_4^{2-}}$ is a weaker base than $\mathrm{SO_3^{2-}}$ due to higher oxidation state of S (+6 vs. +4) and more resonance stabilization.

$\mathrm{HSO_4^-}$ is stronger acid than $\mathrm{HSO_3^-}$.