Acids and Bases: Properties, Equilibria, and Applications

Les acides et les bases (Acids and Bases)

Introduction

This section introduces the fundamental concepts of acids and bases, focusing on their definitions, properties, and the equilibria they establish in aqueous solutions. The notes are structured to provide a comprehensive overview suitable for college-level General Chemistry.

Les couples acide-base conjugués (Conjugate Acid-Base Pairs)

Definitions and Concepts

• Acide de Brønsted: A substance that donates a proton (H+).

• Base de Brønsted: A substance that accepts a proton.

• Conjugate base: What remains after an acid has donated a proton.

• Conjugate acid: What is formed when a base accepts a proton.

Example:

• $\mathrm{CH_3COOH(aq) + H_2O(l) \rightarrow CH_3COO^-(aq) + H_3O^+(aq)}$

• Here, $\mathrm{CH_3COOH}$ and $\mathrm{CH_3COO^-}$ form a conjugate acid-base pair, as do $\mathrm{H_2O}$ and $\mathrm{H_3O^+}$.

Identifying Conjugate Pairs in Reactions

• In any acid-base reaction, identify the acid and base on each side of the equation to determine the conjugate pairs.

• Example:

  • (a) $\mathrm{NH_4^+(aq) + OH^-(aq) \rightarrow NH_3(aq) + H_2O(l)}$

  • Conjugate pairs: $\mathrm{NH_4^+}$/$\mathrm{NH_3}$ and $\mathrm{H_2O}$/$\mathrm{OH^-}$

  • (b) $\mathrm{H_2O(aq) + CO_3^{2-}(aq) \rightarrow OH^-(aq) + HCO_3^-(aq)}$

  • Conjugate pairs: $\mathrm{H_2O}$/$\mathrm{OH^-}$ and $\mathrm{CO_3^{2-}}$/$\mathrm{HCO_3^-}$

Les propriétés acido-basiques de l'eau (Acid-Base Properties of Water)

Autoionization of Water

• Water can act as both an acid and a base (amphoteric).

• Autoionization reaction: $\mathrm{2H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)}$

• Equilibrium constant: $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$ at 25°C

• In pure water: $[H^+] = [OH^-] = 1.0 \times 10^{-7}$ M

Le pH: une mesure du degré d'acidité (pH: A Measure of Acidity)

Definition and Calculation

• pH is a practical measure of the concentration of $H^+$ ions: $\mathrm{pH = -\log[H^+]}$

• Acidic solution: $\mathrm{pH < 7.00}$

• Neutral solution: $\mathrm{pH = 7.00}$

• Basic solution: $\mathrm{pH > 7.00}$

• pOH is less commonly used: $\mathrm{pOH = -\log[OH^-]}$

• Relationship: $\mathrm{pH + pOH = 14.00}$ at 25°C

Examples

• Calculate the pH of a $0.76$ M $\mathrm{HNO_3}$ solution:

• $\mathrm{pH = -\log(0.76) = 0.12}$

• Given a fruit juice with $\mathrm{pH = 3.33}$, $[H^+] = 10^{-3.33} = 4.7 \times 10^{-4}$ M

Les acides forts et les bases fortes (Strong Acids and Bases)

Strong Acids

• A strong acid is a strong electrolyte that dissociates completely in water.

• Examples: $\mathrm{HCl}$, $\mathrm{HNO_3}$, $\mathrm{H_2SO_4}$

Weak Acids

• A weak acid is a weak electrolyte that only partially dissociates in water.

• Examples: $\mathrm{HF}$, $\mathrm{CH_3COOH}$, $\mathrm{NH_4^+}$

Strong Bases

• A strong base is a strong electrolyte that dissociates completely in water.

• Examples: $\mathrm{NaOH}$, $\mathrm{KOH}$

Weak Bases

• A weak base is a weak electrolyte that only partially reacts with water.

• Example: $\mathrm{NH_3}$

Les acides faibles et les constantes d'ionisation des acides (Weak Acids and Acid Ionization Constants)

Acid Dissociation and $K_a$

• Weak acids do not dissociate completely in water.

• General reaction: $\mathrm{HA(aq) + H_2O(l) \rightleftharpoons H_3O^+(aq) + A^-(aq)}$

• Acid dissociation constant: $K_a = \dfrac{[H_3O^+][A^-]}{[HA]}$

• The larger the $K_a$, the stronger the acid.

Calculating Equilibrium Concentrations

• For weak acids, use the initial concentration and $K_a$ to set up an ICE (Initial, Change, Equilibrium) table.

• Approximations can be made if the degree of dissociation is small (typically if $K_a$ is much less than the initial concentration).

• If the approximation is not valid, solve the quadratic equation for $x$ (the amount dissociated).

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Examples

• Given $0.12$ M HA with $K_a = 5.7 \times 10^{-4}$, solve for $[H^+]$ and pH.

• Set up: $K_a = \dfrac{x^2}{0.12 - x}$, solve for $x$.

• If $x$ is not negligible compared to $0.12$, use the quadratic formula.

Le pourcentage d'ionisation (Percent Ionization)

Definition and Calculation

• Percent ionization measures the fraction of acid molecules that ionize in solution.

• Formula: $\text{Percent ionization} = \dfrac{[H_3O^+]_{eq}}{[HA]_{initial}} \times 100\%$

• For a monoprotic acid: $\text{Percent ionization} = \dfrac{x}{[HA]_{initial}} \times 100\%$

• Percent ionization decreases as the initial acid concentration increases.

Les diacides et les polyacides (Polyprotic Acids)

Stepwise Ionization

• Polyprotic acids can donate more than one proton, with each step having its own $K_a$.

• Each successive ionization has a smaller $K_a$ value.

• Example: $\mathrm{H_2CO_3}$

  • First ionization: $\mathrm{H_2CO_3(aq) + H_2O(l) \rightleftharpoons HCO_3^-(aq) + H_3O^+(aq)}$

  • Second ionization: $\mathrm{HCO_3^-(aq) + H_2O(l) \rightleftharpoons CO_3^{2-}(aq) + H_3O^+(aq)}$

Les bases faibles et les constantes d'ionisation des bases (Weak Bases and Base Ionization Constants)

Base Ionization and $K_b$

• Weak bases react with water to produce $OH^-$ ions.

• General reaction: $\mathrm{B(aq) + H_2O(l) \rightleftharpoons BH^+(aq) + OH^-(aq)}$

• Base ionization constant: $K_b = \dfrac{[BH^+][OH^-]}{[B]}$

• Example: $\mathrm{NH_3(aq) + H_2O(l) \rightleftharpoons NH_4^+(aq) + OH^-(aq)}$

La relation entre les constantes d'ionisation des couples acide-base conjugués (Relationship Between $K_a$ and $K_b$)

• For a conjugate acid-base pair: $K_a \times K_b = K_w$

• $K_w = 1.0 \times 10^{-14}$ at 25°C

• If $K_a$ is large (strong acid), $K_b$ for its conjugate base is small (weak base), and vice versa.

• Example: If $K_a$ for an acid is $1.8 \times 10^{-5}$, $K_b$ for its conjugate base is $K_b = \dfrac{K_w}{K_a} = \dfrac{1.0 \times 10^{-14}}{1.8 \times 10^{-5}} = 5.6 \times 10^{-10}$

Les propriétés acido-basiques des sels (Acid-Base Properties of Salts)

Salts in Water

• The hydrolysis of a salt depends on the acid and base from which it is derived.

• If both the cation and anion are from strong acid and strong base, the solution is neutral.

• If the cation is from a weak base and the anion from a strong acid, the solution is acidic.

• If the cation is from a strong base and the anion from a weak acid, the solution is basic.

Examples

• $\mathrm{CH_3COONa}$ in water: $\mathrm{CH_3COO^-}$ reacts with water to produce $OH^-$, making the solution basic.

• $\mathrm{NH_4Cl}$ in water: $\mathrm{NH_4^+}$ reacts with water to produce $H_3O^+$, making the solution acidic.

Les oxydes acides, basiques, et amphotères (Acidic, Basic, and Amphoteric Oxides)

Reactions with Water

• Acidic oxides (usually nonmetal oxides) react with water to form acids.

• Example: $\mathrm{SO_3 + H_2O \rightarrow H_2SO_4}$

• Basic oxides (usually metal oxides) react with water to form bases.

• Example: $\mathrm{Na_2O + H_2O \rightarrow 2NaOH}$

• Amphoteric oxides can react as either acids or bases, depending on the conditions.

Periodic Trends

• Across a period, oxides become more acidic from left to right.

• Down a group, oxides become more basic.

Summary Table: Strong and Weak Acids and Bases

Type	Definition	Examples
Strong Acid	Completely dissociates in water	HCl, HNO3, H2SO4
Weak Acid	Partially dissociates in water	CH3COOH, HF, HCN
Strong Base	Completely dissociates in water	NaOH, KOH
Weak Base	Partially reacts with water	NH3, CH3NH2

Key Equations

• $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$

• $\mathrm{pH = -\log[H^+]}$

• $\mathrm{pOH = -\log[OH^-]}$

• $\mathrm{pH + pOH = 14.00}$

• $K_a = \dfrac{[H_3O^+][A^-]}{[HA]}$

• $K_b = \dfrac{[BH^+][OH^-]}{[B]}$

• $K_a \times K_b = K_w$

• Percent ionization: $\dfrac{[H_3O^+]_{eq}}{[HA]_{initial}} \times 100\%$

Additional info:

• Some context and explanations have been expanded for clarity and completeness, as the original slides were in French and used brief bullet points.

• Examples and equations have been formatted for clarity and to match standard General Chemistry curriculum.