BI1014 - acid base reactions

Bronsted-lowry theory

Acid: Proton donor (H⁺)

Base: Proton acceptor

equilibrium

Governed by:

Rate constants (Kₒᵣ, K₆ₐ꜀ₖ)

Le Chatelier’s Principle: System shifts to oppose change (concentration, pressure, temperature)

Ka

Ka (acid constant):

Large Ka → strong acid (complete dissociation)

Kb

(base constant)

Large Kb → strong base

Kw

ionic product of water

strong acids and bases

Ka or Kb -> >1

Dissociation -> Full

Examples (Acid) HCl, HNO₃, H₂SO₄

Examples (Base) NaOH, KOH

weak acids and bases

Ka or Kb -> <1

Dissociation -> Partial

Examples (Acid) CH₃COOH, citric acid

Examples (Base) NH₃, amines

conjugate base pairs

Every acid has a conjugate base

Every base has a conjugate acid

example of a conjugate pair

CH₃COOH ⇌ CH₃COO⁻ + H⁺(Acid ⇌ Conjugate Base + H⁺)

neutral salts

Strong acid + strong base → pH ~7

acidic salts

Strong acid + weak base → pH <7

basic salts

Weak acid + strong base → pH >7

ampholytes

Can act as both acid & base

Examples: Amino acids, hydrogen phosphate

buffers

Solution that resists pH change

components of a buffer

Weak acid + conjugate base (e.g. CH₃COOH + CH₃COONa)

Weak base + conjugate acid (e.g. NH₃ + NH₄Cl)

buffer action

Add acid (H⁺) → base component neutralizes it

Add base (OH⁻) → acid component donates H⁺

Henderson-Hasselbalch equation

Change of 1 pH = 10× concentration ratio shift

Not valid if [HA] or [A⁻] = 0

titration curve

four stages

1. Initial pH: pH of analyte before titration begins.

2. Buffer Region: Gradual change in pH as titrant is added (especially with weak acids/bases)

3. Equivalence Point:

4. After Equivalence: Excess titrant determines the pH

equivalence point

Chemically, when moles of acid = moles of base

Acid = Base (stoichiometrically)

All acid/base is neutralised

Sharp pH change

end point

When the indicator changes colour (should be close to equivalence point).

indicators

An indicator is a weak acid or base that changes colour depending on the pH of the solution.

how do indicators work

Indicator exists in two forms: HInd (acid) and Ind⁻ (base)

Colour changes depending on pH and which form dominates

types of indicators

Strong acid + strong base

Phenolphthalein or Methyl orange

Colourless → pink or red → yellow

Weak acid + strong base

Phenolphthalein

Colourless → pink

Strong acid + weak base

Methyl orange

Red → yellow

Weak acid + weak base

No good sharp change; use pH meter

pH equation

pH = -log[H+]