General Chemistry

1. Acids and Bases
   1. Definitions and Properties
      1. Arrhenius Definition
         1. Acids produce H⁺ ions in water
         2. Bases produce OH⁻ ions in water
         3. Example of acids: HCl, H₂SO₄
         4. Example of bases: NaOH, KOH
      2. Brønsted-Lowry Definition
         1. Acids are proton donors
         2. Bases are proton acceptors
         3. Conjugate acid-base pairs
         4. Example of acid-base reaction: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
      3. Lewis Definition
         1. Acids are electron pair acceptors
         2. Bases are electron pair donors
         3. Includes broader range of substances
         4. Example: BF₃ as a Lewis acid
   2. Properties of Acids
      1. Sour taste
      2. Turn blue litmus paper red
      3. React with metals to produce hydrogen gas
      4. Conduct electricity in solution
   3. Properties of Bases
      1. Bitter taste
      2. Turn red litmus paper blue
      3. Slippery feel
      4. Conduct electricity in solution
   4. pH and pOH
      1. Definition of pH
         1. Scale from 0 to 14
         2. Neutral pH = 7
         3. Acidic pH < 7
         4. Basic pH > 7
      2. Calculating pH
         1. pH = -log[H⁺]
         2. Relationship with pOH: pH + pOH = 14
      3. Definition of pOH
         1. Measure of hydroxide ion concentration
         2. pOH = -log[OH⁻]
      4. Chemical significance of pH
         1. Impact on biological systems
         2. Influence on chemical reactivity
   5. Acid-Base Titration
      1. Purpose and applications
         1. Determining concentration of unknown solution
         2. Neutralization reaction as the basis
      2. Indicators
         1. Phenolphthalein, methyl orange, bromothymol blue
         2. End point vs. equivalence point
      3. Titration curve
         1. Strong acid with strong base
         2. Weak acid with strong base
         3. Strong acid with weak base
      4. Calculating unknown concentrations
         1. Using titration data: \( M_1V_1 = M_2V_2 \)
   6. Buffers
      1. Definition and function
         1. Resist changes in pH upon addition of small amounts of acid or base
         2. Composed of weak acid and its conjugate base or weak base and its conjugate acid
      2. Buffer capacity
         1. The amount of acid or base the buffer can neutralize
         2. Depends on concentrations of the buffer components
      3. Henderson-Hasselbalch Equation
         1. \( pH = pK_a + \log \left(\frac{[A⁻]}{[HA]}\right) \)
         2. Useful for calculating pH of a buffer solution
      4. Biological significance of buffers
         1. Blood buffering systems (e.g., bicarbonate buffer)
         2. Importance in maintaining pH stability in physiological processes
   7. Strengths of Acids and Bases
      1. Strong acids and bases
         1. Strong acids fully dissociate in water (e.g., HCl, HNO₃)
         2. Strong bases fully dissociate in water (e.g., NaOH, KOH)
      2. Weak acids and bases
         1. Weak acids partially dissociate in water (e.g., acetic acid)
         2. Weak bases partially dissociate in water (e.g., ammonia)
      3. Measuring strength
         1. Acid dissociation constant (Kₐ)
         2. Base dissociation constant (K_b)
         3. Relating Kₐ and K_b to acid/base strength
   8. Applications and Importance
      1. Industrial applications
         1. Acids and bases in manufacturing (e.g., sulfuric acid in fertilizers)
         2. Importance in chemical synthesis and processing
      2. Environmental impact
         1. Acid rain and its effects on ecosystems
         2. Role in remediation strategies
      3. Everyday uses
         1. Acids in food and beverages (e.g., citric acid, vinegar)
         2. Bases in cleaning products (e.g., baking soda)