Polyprotic and Diprotic Acids

Definition of Polyprotic Acids: Acids that can donate more than one proton (hydrogen ion) in a solution.
Example: Phosphoric acid
- Chemical formula: H₃PO₄
- Contains three acidic protons that can be donated.

Reactions of polyprotic acids occur in multiple steps with each proton donation being its own equilibrium reaction.
Each step for phosphoric acid is labeled as:
- Step 1: H₃PO₄ → H₂PO₄⁻ + H⁺ (equilibrium constant Kₐ₁)
- Step 2: H₂PO₄⁻ → HPO₄²⁻ + H⁺ (equilibrium constant Kₐ₂)
- Step 3: HPO₄²⁻ → PO₄³⁻ + H⁺ (equilibrium constant Kₐ₃)

The stability of anions affects the tendency to donate protons.
Observation: Kₐ values decrease through each proton donation step:
- Kₐ₁ > Kₐ₂ > Kₐ₃
Functional implied behavior: If the change in Kₐ is greater than 10³, the subsequent stages can often be ignored for pH calculations,
- Phosphoric acid specific Kₐ values:
- Kₐ₁ = 10⁻³ (large)
- Kₐ₂ = 10⁻⁸ (smaller)
- Kₐ₃ = 10⁻¹³ (much smaller)
Practical Application: For phosphoric acid, only Kₐ₁ is relevant for calculating pH in a usual scenario because later steps contribute negligibly to the hydrogen ion concentration.

Characteristic: Logarithmic scale (base 10)
A change from pH 1 to pH 2 corresponds to a tenfold change in the concentration of H⁺ ions.
Small contributions from later equilibria do not significantly shift the pH value.

For neutralizing phosphoric acid:
- Required: 3 moles of sodium hydroxide (NaOH) to neutralize 1 mole of H₃PO₄.
- The process consumes hydrogen ions directly with hydroxide ions, not considering equilibrium.

Example of a Functionally Diprotic Acid: Sulfuric acid (H₂SO₄)
- First proton: Strong acid (100% dissociation)
- Second proton: Weak acid (equilibrium present)
Key Reactions:
- H₂SO₄ → HSO₄⁻ + H⁺ (strong acid, no equilibrium)
- HSO₄⁻ → SO₄²⁻ + H⁺ (weak acid, K ≈ 1.2 × 10⁻²)
Consequently, both equilibrium constants must be considered for accurate pH calculation.

Definition: Every acid has a conjugate base; every base has a conjugate acid.
Importance: Evaluate the functional roles of species in acid-base reactions.
- Example: Sodium chloride (NaCl) consists of sodium ions (Na⁺) and chloride ions (Cl⁻), which do not participate in these acid-base reactions.
Contrasting Example: Ammonium chloride (NH₄Cl) contains:
- Ammonium (NH₄⁺): functioning as a weak acid (conjugate of ammonia)
- Chloride (Cl⁻): not a conjugate, non-active in proton exchange.
Expected pH Changes:
- NaCl: pH ≈ 7 (neutral)
- NH₄Cl: pH < 7 (acidic)

Dissolution yields sodium cations (Na⁺) and fluoride anions (F⁻).
- Sodium ions do not affect pH.
- Fluoride ions act as a base, coming from the weak acid HF.
Result: pH > 7 (basic solution). The extent depends on the amount of fluoride.

Ammonium: weak acid, conjugate of ammonia (NH₃)
Fluoride: weak base, conjugate of hydrofluoric acid (HF)
Analysis of the strengths of these ions determines solution pH:
- Ammonia's Kb: 1.8 × 10⁻⁵.
- Hydrofluoric acid's Ka: 6.8 × 10⁻⁴.
Direct comparison indicates that HF > NH₃: HF’s equilibrium lies farther right than NH₃.
Conclusion: Ammonium is a stronger acid compared to fluoride's base, leading to an expected acidic solution from their combination.

pK values: convert from equilibrium constants for more intuitive comparisons (pK = -log(K)).
- pKa of ammonium: 9.25
- pKb of fluoride: 10.82
Comparison of the relative strengths implicates:
- Ammonium (stronger acid) exhibits lower pKa than fluoride (weaker base; larger pKb).

Strong acids fully ionize in solutions versus weak acids which reach equilibrium.

Example: Nitric acid (HNO₃) as a strong acid compared to nitrous acid (HNO₂) - a weak acid.
Importance of electronegativity and bond strength in defining acid strength.
Electronegativity and Stability: Enhanced stability in anions correlates with electronegative atoms (like oxygen in oxyacids).
- The presence of more oxygen atoms within oxyacids typically stabilizes the conjugate anion (explanation for their relative strengths).
- Chloric acid (HClO₃) and perchloric acid (HClO₄) are stronger acids; chlorous acid (HClO₂) and hypochlorous acid (HClO) are weaker.
- Electronegativity increases with more oxygen atoms influence acid strength significantly.

Careful distinction between strong acids and weak acids based on their dissociation behavior and the consequences for pH in solutions should always be noted, emphasizing the impact of conjugate base and acid strengths (both K and pK values).