CHEM1010 W2 L1

Introduction to Acidity and pH

• Focus of the Week: Understanding acids, bases, and how to measure acidity.

• Two key areas:

  • Acidity scale (pH levels from acidic < 7 to alkaline > 7).

  • Consequences of poor acidity management (e.g., acid mine tailings).

• Chapter 16 of the textbook will guide the lectures.

• Format: Three shorter lectures throughout the week.

Definitions

Acid

• Generally defined as a substance that produces hydrogen ions (H⁺) or hydronium ions (H₃O⁺) in water.

• Example: Hydrochloric acid (HCl) completely dissociates in water, forming H⁺ and Cl⁻ ions.

• Strong acids dissociate completely (e.g., HCl, nitric acid (HNO₃)).

Base

• Defined as a substance that accepts a proton (H⁺) and usually produces hydroxide ions (OH⁻) in water.

• Example: Sodium hydroxide (NaOH) completely dissociates into Na⁺ and OH⁻ ions.

• Strong bases dissociate completely (e.g., NaOH, KOH).

Theories of Acids and Bases

Bronsted-Lowry Theory

• Proposes acids as proton donors and bases as proton acceptors.

• An acid must have a removable proton, while a base must be able to accept a proton.

Amphoteric Compounds

• Compounds that can act as both acids and bases (e.g., bicarbonate ion, HCO₃⁻).

• Water (H₂O) is amphoteric and can act both as an acid and a base in different contexts.

Reactions Involving Water

• When HCl dissolves in water:

  • HCl + H₂O → H₃O⁺ + Cl⁻

• Water acts as a Bronsted-Lowry base, accepting the proton to form H₃O⁺.

• The chloride ion (Cl⁻) is the conjugate base of HCl.

Conjugate Acid-Base Pairs

• Definition: Acid-base pairs differ by one H⁺ ion.

  • Example: HCl (acid) and Cl⁻ (conjugate base).

  • Example: H₂O (base) and H₃O⁺ (conjugate acid).

Equilibrium and Dissociation

• Strong acids and bases completely dissociate in solution; no equilibrium constant needed as they go to completion.

• Examples of strong acids: HCl, HNO₃, H₂SO₄ (sulfuric acid).

• Examples of strong bases: Compounds containing alkali metal hydroxides (e.g., NaOH, KOH).

Water as a Special Case

• Water can dissociate slightly to produce H₃O⁺ and OH⁻:

  • 2 H₂O ⇌ H₃O⁺ + OH⁻ (hydrolysis).

• The ion product constant, K_w = [H₃O⁺][OH⁻] = 1 x 10⁻¹⁴ at 25°C.

• In pure water, [H₃O⁺] = [OH⁻] = 1 x 10⁻⁷ M; pH = 7.

pH Measurement

• Definition of pH: Negative logarithm of the H₃O⁺ concentration:

  • pH = -log[H₃O⁺].

• pH scales:

• Acidic: pH < 7

• Neutral: pH = 7

• Basic: pH > 7.

• Common substances and their pH values:

  • Stomach acid: pH ∼ 1

  • Soft drinks: pH ∼ 2-3

  • Blood: Slightly basic (pH ∼ 7.4).

Indicators and pH Meters

• Litmus paper: Used for rough pH measurements (red in acid, blue in base).

• Indicators: Change color over a specific pH range; typically a weak acid or base.

• pH meters: Provide precise pH measurements, usually to two decimal places.

Strong Electrolytes

• Strong acids are strong electrolytes as they exist entirely as ions in solution (e.g., HCl → H₃O⁺ + Cl⁻).

• Monoprotic acids (except sulfuric acid) provide one mole of H⁺ per mole of acid.

Calculating pH

• Attention to acid and base dissociation helps in understanding calculations for pH.

• Formula for pH calculations:

  • Given concentration [H₃O⁺], calculate pH using pH = -log[H₃O⁺].

• For bases, similar calculations can apply using [OH⁻] and pOH relationships.

• Relationship: pH + pOH = 14, which is useful for calculations.