Acids and Bases - Basic Introduction - Chemistry

Basics of Acids and Bases

Identifying Acids and Bases

• Acids:

  • Typically have a hydrogen in front of their chemical formula.

  • Examples include:

    • HCl (hydrochloric acid)

    • HF (hydrofluoric acid)

    • HC₂H₃O₂ (acetic acid)

  • If hydrogen is attached to a nonmetal, it is likely an acid.

• Bases:

  • Typically contain a hydroxide ion (OH⁻).

  • Examples include:

    • NaOH (sodium hydroxide)

    • KOH (potassium hydroxide)

  • If hydrogen is next to a metal (e.g., sodium hydride), it indicates a base.

Charge Considerations

• Hydrogen Ion Charges:

  • Hydrogen with a positive charge (H⁺) indicates an acid.

  • Hydrogen with a negative charge indicates a base.

Definitions of Acids and Bases

Arrhenius Definition

• Acids: Substances that release H⁺ ions in solution (hydronium ions H₃O⁺ in water).

• Bases: Substances that release OH⁻ ions in solution.

Brønsted-Lowry Definition

• Acids: Proton donors.

• Bases: Proton acceptors.

  • Example: HCl acts as a Brønsted-Lowry acid in water where H⁺ is transferred to H₂O, forming H₃O⁺.

Conjugate Acid-Base Pairs

• The acid turns into its conjugate base after donating a proton.

• The base turns into its conjugate acid after accepting a proton.

  • Example: HCl (acid) → Cl⁻ (conjugate base) + H⁺; H₂O (base) → H₃O⁺ (conjugate acid).

pH and Ion Concentration

pH Scale

• Ranges from 0 to 14 (though it can extend beyond these values).

• Neutral pH: 7

• Acidic pH: < 7 (e.g., pH -2 is very acidic)

• Basic pH: > 7

pH Calculations

• pH = -log[H₃O⁺]

• pOH = -log[OH⁻]

• Relationship: pH + pOH = 14 at 25°C

• H₃O⁺ concentration = 10^(-pH)

• OH⁻ concentration = 10^(-pOH)

Strong Acids vs. Weak Acids

• Strong Acids: Ionize completely in solution (strong electrolytes); e.g., HCl, HNO₃, H₂SO₄.

• Weak Acids: Partially ionize (<5%); e.g., acetic acid, HF.

  • Common Strong Acids: HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄.

  • Notice that the number of oxygen atoms in oxyacids relates to strength.

Acid-Base Reactions

Strong Acids

• Reaction with water (single arrow indicates complete ionization):

  • HCl + H₂O → Cl⁻ + H₃O⁺

Weak Acids

• Reaction with water (double arrow indicates equilibrium):

  • HF ⇌ H⁺ + F⁻

Identifying Strong and Weak Bases

• Strong Bases: Soluble ionic compounds (e.g., NaOH, KOH) that ionize completely in water.

• Weak Bases: Insoluble compounds or those that do not ionize completely (e.g., NH₃).

Properties of Acids and Bases

General Properties

• Acids Taste: Sour.

• Bases Taste: Bitter; feel slippery.

• Acids turn blue litmus paper red, bases turn red litmus paper blue.

• Conductivity: Strong acids/bases are strong electrolytes; weak acids/bases are weak electrolytes.

Chemical Reactions

• Acid reacting with active metals (like zinc) produces hydrogen gas.

  • E.g., Zn + HCl → ZnCl₂ + H₂.

• Non-reactive metals (e.g., copper) do not react with acids.

Definitions Recap

• Arrhenius Acids: Release H⁺ ions.

• Arrhenius Bases: Release OH⁻ ions.

• Brønsted-Lowry Acids: Proton donors.

• Brønsted-Lowry Bases: Proton acceptors.

• Lewis Acids/Bases: Electron pair acceptors/donors.

Amphoteric Substances

• Substances that can act as either an acid or a base depending on the reaction context.

• Examples include water (H₂O) and H₂PO₄⁻.

Acid-Base Constants

• Kₐ (Acid Dissociation Constant): Measures strength; higher Kₐ means stronger acid.

• Kₑ (Base Dissociation Constant): Similarly measures base strength.

• Relationship: Kₐ * Kb = Kₕ (1 x 10^(-14) at 25°C).

Practice Problems Recap

• Use prior concepts to calculate pH, pOH, and ion concentrations.

• Understand how to derive Kₐ and Kb values based on given information.