Basics of Acids and Bases

Identifying Acids and Bases

• Acids: Typically have a hydrogen (H) in front (e.g., HCl, HF, HC2H3O2).
  • When hydrogen has a positive charge, it's an acid.
  • Acids are positively charged.
• Bases: Typically have a hydroxide ion (OH) (e.g., NaOH, KOH).
  • When hydrogen is next to a metal with a negative charge, it's a base.
  • Bases are usually negatively charged.

Definitions

• Arrhenius Definition
  • Acids release H⁺ ions in solution, forming hydronium (H3O⁺).
  • Bases release OH⁻ ions in solution.
• Bronsted-Lowry Definition
  • Acids are proton donors.
  • Bases are proton acceptors.

Acid-Base Reactions

• Conjugate Acid-Base Pairs:
  • Acid turns into its conjugate base after donating a proton.
  • Base turns into its conjugate acid after accepting a proton.
• Example: HCl in water:
  • HCl (acid) donates H⁺ to H2O (base), forming Cl⁻ (conjugate base) and H3O⁺ (conjugate acid).

Writing Conjugate Acids and Bases

• Conjugate Acid: Add H⁺ and increase charge by 1.
• Conjugate Base: Remove H⁺ and decrease charge by 1.
• Example:
  • H2O -> Conjugate Acid: H3O⁺, Conjugate Base: OH⁻.
  • NH3 -> Conjugate Acid: NH4⁺, Conjugate Base: NH2⁻.

pH and pOH

• pH Scale: Typically 0 to 14.
  • 7 is neutral.
  • <7 is acidic.

  • 7 is basic.

• Calculating pH:
  • pH = -log[H3O⁺].
  • pOH = -log[OH⁻].
  • pH + pOH = 14 at 25°C.
  • [H3O⁺] = 10^(-pH).
  • [OH⁻] = 10^(-pOH).

Strong vs. Weak Acids/Bases

• Strong Acids: Ionize completely (e.g., HCl, HBr, HI, HNO3, H2SO4, HClO4).
• Weak Acids: Partial ionization (e.g., HF, acetic acid).
• Strong Bases: Soluble ionic compounds that ionize completely (e.g., NaOH, KOH).
• Weak Bases: Partially ionize and are often insoluble.

Chemical Reactions

• Strong Acids in Water: Use a single arrow (→).
• Weak Acids in Water: Use a reversible arrow (⇌).

Properties of Acids and Bases

• Acids: Taste sour, turn blue litmus paper red.
• Bases: Taste bitter, feel slippery, turn red litmus paper blue.
• Both can conduct electricity, strong ones are strong electrolytes.
• Acids react with active metals to produce hydrogen gas.

Definitions Revisited

• Lewis Acid-Base Theory:
  • Lewis Acids: Electron pair acceptors.
  • Lewis Bases: Electron pair donors.

Calculations and Practice Problems

• Ka and Kb: Acid and base ionization constants.
  • Stronger acids have larger Ka values and smaller pKa values.
• Practice Problems: Calculate pH, pOH, [H3O⁺], and [OH⁻] given concentrations.

Summary

• Understand how to identify and differentiate between acids and bases.
• Know how to calculate pH and related concentrations.
• Be able to write chemical equations and identify conjugate acids and bases.
• Recognize the properties and definitions associated with various types of acids and bases.