Acids, Bases and Salts — Revision Notes

⚡ 30-Second Revision

Arrhenius: — Acid ($H^+$), Base ($OH^-$) in water.
Brønsted-Lowry: — Acid (Proton Donor), Base (Proton Acceptor).
Lewis: — Acid (Electron Pair Acceptor), Base (Electron Pair Donor).
pH: — $pH = -log[H^+]$. Neutral $pH=7$ at $25^circ C$.
pOH: — $pOH = -log[OH^-]$.
Relationship: — $pH + pOH = 14$ (at $25^circ C$).
Ionic Product of Water: — $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$ (at $25^circ C$).
Acid Dissociation Constant: — $K_a = \frac{[H^+][A^-]}{[HA]}$.
Base Dissociation Constant: — $K_b = \frac{[BH^+][OH^-]}{[B]}$.
Conjugate Pair Relationship: — $K_a \times K_b = K_w$.
Henderson-Hasselbalch (Acidic Buffer): — $pH = pK_a + log\frac{[Salt]}{[Acid]}$.
Henderson-Hasselbalch (Basic Buffer): — $pOH = pK_b + log\frac{[Salt]}{[Base]}$.
Salt Hydrolysis:

- SA+SB: Neutral - SA+WB: Acidic ($K_h = K_w/K_b$) - WA+SB: Basic ($K_h = K_w/K_a$) - WA+WB: Depends on $K_a$ vs $K_b$

2-Minute Revision

Acids, bases, and salts are fundamental to chemistry. Acids donate protons (Brønsted-Lowry) or accept electron pairs (Lewis), while bases accept protons or donate electron pairs. The Arrhenius theory is a simpler, water-specific definition.

The pH scale, $pH = -log[H^+]$, quantifies acidity, with 7 being neutral at $25^circ C$. Strong acids/bases dissociate completely, while weak ones establish equilibrium, characterized by $K_a$ or $K_b$.

For conjugate acid-base pairs, $K_a \times K_b = K_w = 10^{-14}$. Salts are formed from acid-base neutralization and can hydrolyze in water, making solutions acidic (strong acid + weak base salt), basic (weak acid + strong base salt), or neutral (strong acid + strong base salt).

Buffer solutions, containing a weak acid/base and its conjugate, resist pH changes, and their pH can be calculated using the Henderson-Hasselbalch equation. Remember to consider water's autoionization for very dilute acid/base solutions.

5-Minute Revision

Start by solidifying the three acid-base theories: Arrhenius (H+/OH- in water), Brønsted-Lowry (proton donor/acceptor, leading to conjugate pairs), and Lewis (electron-pair acceptor/donor, the broadest definition).

Practice identifying species under each definition. Next, master the pH scale and its calculations. For strong acids/bases, $[H^+]$ or $[OH^-]$ is directly related to concentration. For weak acids/bases, use $K_a$ or $K_b$ and equilibrium expressions, often simplifying to $[H^+] = \sqrt{K_a C_a}$.

Crucially, for very dilute solutions ($<10^{-7}M$), remember to include the $H^+$ or $OH^-$ from water's autoionization. Understand the relationship $K_a \times K_b = K_w$ for conjugate pairs. Salt hydrolysis is vital: a salt of a strong acid and weak base yields an acidic solution (e.

g., $NH_4Cl$), a weak acid and strong base yields a basic solution (e.g., $CH_3COONa$), and strong acid/strong base yields a neutral solution (e.g., $NaCl$). Finally, grasp buffer solutions – their composition (weak acid + conjugate base or weak base + conjugate acid), how they resist pH changes, and apply the Henderson-Hasselbalch equation: $pH = pK_a + log\frac{[Salt]}{[Acid]}$.

Always check units and significant figures in calculations.

Prelims Revision Notes

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Acid-Base Theories:

* Arrhenius: Acid ($H^+$ producer), Base ($OH^-$ producer) in water. Limited. * Brønsted-Lowry: Acid (Proton donor), Base (Proton acceptor). Broader. Conjugate acid-base pairs differ by one $H^+$. * Lewis: Acid (Electron pair acceptor), Base (Electron pair donor). Broadest. Explains reactions without $H^+$ transfer (e.g., $BF_3 + NH_3$).

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pH and pOH:

* $pH = -log[H^+]$, $pOH = -log[OH^-]$. * $pH + pOH = 14$ (at $25^circ C$). * $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$ (at $25^circ C$). * Neutral solution: $pH=7$, $[H^+]=[OH^-]=10^{-7}M$. * Acidic: $pH<7$, Basic: $pH>7$.

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Strength of Acids/Bases:

* Strong: Complete dissociation (e.g., $HCl, NaOH$). $[H^+]$ or $[OH^-]$ directly from concentration. * Weak: Partial dissociation, equilibrium established. $K_a$ (acid dissociation constant), $K_b$ (base dissociation constant).

* For weak acid $HA$: $K_a = \frac{[H^+][A^-]}{[HA]}$. For weak base $B$: $K_b = \frac{[BH^+][OH^-]}{[B]}$. * Approximation for weak acid: $[H^+] = \sqrt{K_a C_a}$ (if $C_a/K_a > 100$). * Conjugate Pair: Strong acid $\rightarrow$ weak conjugate base.

Weak acid $\rightarrow$ strong conjugate base. $K_a \times K_b = K_w$.

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Salt Hydrolysis: — Reaction of salt ions with water to produce $H^+$ or $OH^-$.

* **Strong Acid + Strong Base (e.g., $NaCl$):** No hydrolysis, neutral solution ($pH=7$). * **Strong Acid + Weak Base (e.g., $NH_4Cl$):** Cation hydrolyzes ($NH_4^+ + H_2O \rightleftharpoons NH_3 + H_3O^+$), acidic solution ($pH<7$).

$K_h = K_w/K_b$. * **Weak Acid + Strong Base (e.g., $CH_3COONa$):** Anion hydrolyzes ($CH_3COO^- + H_2O \rightleftharpoons CH_3COOH + OH^-$), basic solution ($pH>7$). $K_h = K_w/K_a$. * **Weak Acid + Weak Base (e.

g., $CH_3COONH_4$):** Both hydrolyze. pH depends on relative $K_a$ and $K_b$. If $K_a > K_b$, acidic; if $K_b > K_a$, basic; if $K_a = K_b$, neutral.

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Buffer Solutions: — Resist pH change.

* Composition: Weak acid + its conjugate base (e.g., $CH_3COOH/CH_3COONa$) OR Weak base + its conjugate acid (e.g., $NH_3/NH_4Cl$). * Henderson-Hasselbalch Equation: $pH = pK_a + log\frac{[Salt]}{[Acid]}$ (for acidic buffer) or $pOH = pK_b + log\frac{[Salt]}{[Base]}$ (for basic buffer).

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Important Note: — For very dilute strong acids/bases ($C \le 10^{-7}M$), always consider the $H^+$ or $OH^-$ contribution from water's autoionization.

Vyyuha Quick Recall

All Boys Love Protons, Electrons, Hydroxides!

Arrhenius: Hydroxides ($OH^-$) for bases.
Brønsted-Lowry: Protons ($H^+$) for acids/bases.
Lewis: Electrons (pairs) for acids/bases.