Chapter 16: Acids and Bases

Definitions of Acids and Bases

Acids and bases are fundamental chemical species with distinct definitions according to different theories:

• Arrhenius Definition:

  • Acid: Produces H+ ions when dissolved in water.

  • Base: Produces OH- ions when dissolved in water.

• Brønsted–Lowry Definition:

  • Acid: H+ donor.

  • Base: H+ acceptor.

• Lewis Definition:

  • Acid: Electron-pair acceptor.

  • Base: Electron-pair donor.

Properties of Acids and Bases

Acids and bases exhibit characteristic physical and chemical properties:

• Acids:

  • Sour taste

  • Ability to dissolve many metals

  • Neutralize bases

  • Turn blue litmus paper red

• Bases:

  • Bitter taste

  • Slippery feel

  • Turn red litmus paper blue

  • Neutralize acids

Structures of Acids

Acids are classified based on their molecular structure:

• Binary Acids: Acid hydrogens attached to a nonmetal atom (e.g., HCl, HF).

• Oxyacids: Acid hydrogens attached to an oxygen atom (e.g., H2SO4, HNO3).

Problems with Arrhenius Theory

The Arrhenius theory is limited because it cannot explain:

• Basic solutions from substances without OH- (e.g., NH3).

• Acidic solutions from substances without H+ (e.g., CO2).

• Acid–base reactions outside aqueous solutions.

Brønsted–Lowry Acid–Base Theory

This theory defines acids and bases based on proton transfer:

• Acid: H+ donor

• Base: H+ acceptor (must have a lone pair of electrons)

• All Arrhenius acids and bases are also Brønsted–Lowry acids and bases.

Example reaction:

• HCl(aq) + H2O(l) → Cl-(aq) + H3O+(aq)

Conjugate Acid–Base Pairs

In Brønsted–Lowry reactions, each acid and base forms a conjugate pair:

• A base accepts a proton and becomes a conjugate acid.

• An acid donates a proton and becomes a conjugate base.

Strong and Weak Acids and Bases

Acids and bases are classified by their degree of ionization:

• Strong Acid: Ionizes completely in water; strong electrolyte.

• Weak Acid: Ionizes only partially; weak electrolyte.

• Strong Base: Forms OH- ions completely; strong electrolyte.

• Weak Base: Forms OH- ions partially; weak electrolyte.

Examples of Strong and Weak Acids

Acid and Base Strength: Ionic Attraction

The strength of an acid or base depends on the equilibrium position of its reaction with water:

• Strong acids have weak attraction between the base form and H+, leading to complete ionization.

• Weak acids have strong attraction, resulting in partial ionization.

Acid Ionization Constant (Ka)

The acid ionization constant, Ka, measures acid strength:

• Larger Ka = stronger acid

• Ka is defined as:

Acidic, Basic, and Neutral Solutions

All aqueous solutions contain both H3O+ and OH- ions:

• Neutral: [H3O+] = [OH-] = 1.00 × 10-7 M

• Acidic: [H3O+] > [OH-]

• Basic: [OH-] > [H3O+]

Measuring Acidity: pH and pOH

The acidity or basicity of a solution is quantified by pH and pOH:

• pH:

• pOH:

• Relationship:

• pH < 7: acidic; pH > 7: basic; pH = 7: neutral

Example Calculations

pKa and pKb

Acid and base strength can also be expressed as pKa and pKb:

• Smaller pKa or pKb indicates a stronger acid or base.

Finding pH of Strong and Weak Acid Solutions

For strong acids, [H3O+] ≈ [acid]. For weak acids, equilibrium calculations are required:

• Use an ICE (Initial, Change, Equilibrium) table to solve for [H3O+].

Finding the Equilibrium Constant from pH

Percent Ionization

Percent ionization measures the fraction of acid molecules that ionize:

Strong and Weak Bases

• Strong Bases: Completely dissociate in water (e.g., NaOH).

• Weak Bases: Only partially accept H+ (e.g., NH3).

Base Ionization Constant (Kb)

Base strength is measured by the base ionization constant, Kb:

• Larger Kb = stronger base

Finding the [OH-] and pH of a Weak Base Solution

• Use an ICE table to solve for [OH-].

• Calculate [H3O+] using Kw.

• Find pH from [H3O+].

Relationship Between Ka and Kb

Ka and Kb for a conjugate acid–base pair are related:

• For a given acid and its conjugate base:

Classifying Salt Solutions

Salt solutions can be acidic, basic, or neutral depending on the nature of their ions:

• Neutral: Cation from strong base, anion from strong acid (e.g., NaCl).

• Basic: Cation from strong base, anion from weak acid (e.g., NaF).

• Acidic: Cation is conjugate acid of weak base, anion from strong acid (e.g., NH4Cl).

• Acidic: Cation is highly charged metal ion, anion from strong acid (e.g., Al(NO3)3).

• Depends: Both cation and anion are conjugates of weak acid/base (e.g., NH4F).

Strengths of Binary and Oxyacids

• Binary acid strength increases across a period and down a group.

• Oxyacid strength increases with electronegativity and oxidation number of the central atom.

• More oxygens attached to the central atom increases acid strength.

Lewis Acid–Base Theory

Lewis theory focuses on electron pair transfer:

• Lewis Base: Electron pair donor (nucleophile).

• Lewis Acid: Electron pair acceptor (electrophile).

• Does not require H atoms.

Summary Table: pH of Common Substances

Summary Table: Strong and Weak Acids and Bases

Summary Table: Acid Ionization Constants

Summary Table: pH Calculations