Study Notes on Acids, Bases, Salts, and Redox Reactions

Introduction to Acids, Bases, and Salts

• Key Concepts:

  • Characteristics of acids, bases, and salts.

  • Quantitative relationships in acid-base chemistry.

  • Applications of redox reactions in electrochemistry.

Characteristics of Acids, Bases, and Salts

Definitions and Characteristics

• Acids:

  • Taste: Sour.

  • Litmus Test: Changes blue litmus paper to red.

  • Typical Examples: Hydrochloric acid (HCl), sulfuric acid (H₂SO₄).

• Bases:

  • Taste: Bitter.

  • Feel: Slippery/Soapy.

  • Litmus Test: Changes red litmus paper to blue.

  • Typical Examples: Sodium hydroxide (NaOH).

• Salts:

  • Formed by the reaction of acids with bases.

  • Generally neutral in pH when from strong acids/bases, but can be acidic/basic when derived from weak acids/bases.

Theories of Acids and Bases

Arrhenius Theory (Svante Arrhenius, 1883)

• Arrhenius Acid: A compound that produces H⁺ ions in solution.

  • Example:
    $HCl(aq) + H_2O (l) <br>ightarrow H_3O^+(aq) + Cl^-(aq)$

• Arrhenius Base: A compound that produces OH⁻ ions in solution.

  • Example:
    $NaOH(aq) <br>ightarrow Na^+(aq) + OH^-(aq)$

Limitations of Arrhenius Theory

• Does not explain the behavior of weak bases like ammonia (NH₃) in non-ionic forms.

• Examples of ammonia neutralizing HCl:

  • $NH_3(aq) + HCl(aq) <br>ightarrow NH_4Cl(aq)$

  • $NH_3(aq) + H_2O(l) <br>ightleftharpoons NH_4^+(aq) + OH^-(aq)$

• In gaseous reactions like:

  • $NH_3(g) + HCl(g) <br>ightarrow NH_4Cl(s)$, Arrhenius theory fails.

Brønsted-Lowry Theory

• Brønsted-Lowry Acid: A proton (H⁺) donor.

• Brønsted-Lowry Base: A proton acceptor.

  • Example:

  • $F^-(aq) + H_2O(l) <br>ightleftharpoons HF(aq) + OH^-(aq)$

• The acid dissociation constant:

  • $K_a = rac{[A^-][H^+]}{[HA]}$

• The base dissociation constant:

  • $K_b = rac{[HB^+][OH^-]}{[B]}$

Conjugate Acids and Bases

• Each acid (HA) has a conjugate base (A⁻) following the loss of a proton.

• Each base (B) has a conjugate acid (HB⁺) after gaining a proton.

• Strong acids have weak conjugate bases. For example, the conjugate base of HCl is Cl⁻ (very weak).

pH Concept

Definition of pH

• pH is the negative logarithm of the hydrogen ion concentration:

  • $pH = - ext{log}([H^+]) = - ext{log}([H_3O^+])$

• Acidic Solutions: $[H^+] > 1.0 imes 10^{-7} M$, $pH < 7.00$

• Neutral Solutions: $[H^+] = 1.0 imes 10^{-7} M$, $pH = 7.00$

• Basic Solutions: $[H^+] < 1.0 imes 10^{-7} M$, $pH > 7.00$

Calculating pH and pOH

• If given pH, the [H⁺] can be calculated by:

  • $[H^+] = 10^{-pH}$

• Similarly,

  • $pOH = - ext{log}([OH^-])$

  • Relationships: $pH + pOH = 14.00$

Characteristics of Water

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

• Autoionization of Water:

  • $H_2O(l) + H_2O(l) <br>ightleftharpoons H_3O^+(aq) + OH^-(aq)$

• Water has a dissociation constant:

  • $K_w = [H_3O^+][OH^-] = 1.00 imes 10^{-14}$ at 25°C.

Strong and Weak Acids/Bases

Strong Acids

• Completely ionize in solution, examples include HCl, HNO₃, H₂SO₄.

• Dissociation Example:

  • For HCl:

  • $HCl(aq) <br>ightarrow H^+(aq) + Cl^-(aq)$

Weak Acids

• Incompletely ionized in solution, examples include acetic acid (CH₃COOH), hydrocyanic acid (HCN).

• Conjugate Bases produced are weaker compared to strong acids.

Strong Bases

• Include hydroxides of Group I and II metals: LiOH, NaOH, KOH, etc.

Weak Bases

• Examples are ammonia (NH₃) and its derivatives.

Distribution and Properties of Salts

Physical Properties

• Crystalline structure; soluble in water; conduct electricity in molten state.

Classifications of Salts

1. Salt of strong acid and strong base: Neutral.

2. Salt of weak acid and strong base: Basic.

3. Salt of strong acid and weak base: Acidic.

4. Salt of weak acid and weak base: Dependent on relative strengths.

Buffer Solutions

• Combinations of weak acid/base and their salts, resisting changes in pH upon addition of acids or bases.