• HCl: acid (donates H⁺)
• H₂O: base (accepts H⁺)

Conjugate acid-base pairs:

$HA + B \rightleftharpoons A^- + HB^+$

• HA/A⁻: conjugate pair (differ by H⁺)
• B/HB⁺: conjugate pair

Example: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

• NH₃/NH₄⁺: conjugate pair
• H₂O/OH⁻: conjugate pair

Lewis Theory

Most general:

Acid: Electron pair acceptor Base: Electron pair donor

Example: BF₃ + NH₃ → F₃B-NH₃

• BF₃: Lewis acid (accepts electron pair)
• NH₃: Lewis base (donates electron pair)

All Brønsted acids/bases are Lewis, but not vice versa

Autoionization of Water

Water self-ionizes:

$2H2O \rightleftharpoons H3O^+ + OH^-$

Simplified: H₂O ⇌ H⁺ + OH⁻

Ion product constant (K_w):

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

Key relationships:

• Pure water: [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ M
• Acidic: [H⁺] > [OH⁻]
• Basic: [H⁺] < [OH⁻]
• Always: [H⁺][OH⁻] = 1.0 × 10⁻¹⁴

pH Scale

pH definition:

$\text{pH} = -\log[H^+]$

pOH definition:

$\text{pOH} = -\log[OH^-]$

Relationship:

$\text{pH} + \text{pOH} = 14.00 \text{ at 25°C}$

pH Scale interpretation:

Ranges:

• pH < 7: Acidic
• pH = 7: Neutral
• pH > 7: Basic

Calculating pH

From [H⁺]:

$\text{pH} = -\log[H^+]$

Example: [H⁺] = 1.0 × 10⁻³ M

pH = -log(1.0 × 10⁻³) = 3.00

From [OH⁻]:

1. Calculate pOH = -log[OH⁻]
2. pH = 14.00 - pOH

Example: [OH⁻] = 1.0 × 10⁻⁴ M

pOH = -log(1.0 × 10⁻⁴) = 4.00 pH = 14.00 - 4.00 = 10.00

Calculating [H⁺] from pH

Inverse relationship:

$[H^+] = 10^{-\text{pH}}$

Example: pH = 5.00

[H⁺] = 10⁻⁵·⁰⁰ = 1.0 × 10⁻⁵ M

Strong Acids and Bases

Strong acids (completely ionize):

• HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄

For strong monoprotic acid:

• [H⁺] = [acid]

Example: 0.010 M HCl

• [H⁺] = 0.010 M
• pH = -log(0.010) = 2.00

Strong bases (completely dissociate):

• Group 1A hydroxides: LiOH, NaOH, KOH
• Group 2A hydroxides: Ca(OH)₂, Ba(OH)₂, Sr(OH)₂

For strong base:

• [OH⁻] = [base] × (number of OH⁻)

Example: 0.010 M NaOH

• [OH⁻] = 0.010 M
• pOH = 2.00
• pH = 12.00

Example: 0.010 M Ca(OH)₂

• [OH⁻] = 2 × 0.010 = 0.020 M
• pOH = -log(0.020) = 1.70
• pH = 14.00 - 1.70 = 12.30

Significant Figures in pH

pH has decimal places = sig figs in [H⁺]

Example: [H⁺] = 2.5 × 10⁻³ M (2 sig figs)

• pH = 2.60 (2 decimal places)

The digits before decimal point come from exponent

Conjugate pairs differ by one H⁺:

Pair 1: HF and F⁻

• HF → F⁻ + H⁺
• HF is acid, F⁻ is its conjugate base
• HF/F⁻ conjugate acid-base pair

Pair 2: NH₃ and NH₄⁺

• NH₃ + H⁺ → NH₄⁺
• NH₃ is base, NH₄⁺ is its conjugate acid
• NH₃/NH₄⁺ conjugate acid-base pair

Summary:

Answers:

• Pair 1: HF/F⁻
• Pair 2: NH₃/NH₄⁺

Pattern: Conjugate pairs always differ by exactly one proton (H⁺)