03/28 Chem Lecture - Strong and Weak Acids And Bases

  • Exam Information

    • Chapter 17 content won't be included in the upcoming exam on Monday.
    • Office hours extended today from 11:30 to 1:00 in Conference Room 302.
    • Be prepared to ask questions related to practice exams, practice problems, etc.
    • There is a CAl (Collaborative Learning) session at 2 PM today.
    • Homework due at 5 PM today.
    • No PCA (Periodic Check Assessment) due on Monday; expected next week.
    • No class next Friday.

  • Kinetics Review

    • An exam on reaction mechanisms (part of kinetics) will be on Monday.
    • Need to read and understand reaction coordinate diagrams.
    • Count the number of peaks to determine the number of steps and transition states in a reaction.
    • Determining exothermic vs. endothermic based on energy comparison between reactants and products.
    • Identifying the rate determining step through activation energy comparisons.

  • Introduction to Acids and Bases

    • Review of pH and pOH:
    • Aqueous solutions contain H₃O⁺ (hydronium) and OH⁻ (hydroxide) ions.
    • Neutral solutions have equal concentrations of H₃O⁺ and OH⁻.
    • Acidic solutions have more H₃O⁺ than OH⁻; Basic solutions have more OH⁻ than H₃O⁺.
    • pH = -log[H₃O⁺] allows for easier comparison of acid concentrations on a manageable scale.
    • Common pH scale ranges from 0 (most acidic) to 14 (most basic).

  • Calculating pH from Concentrations

    • If you know [H₃O⁺], you can calculate pH and vice versa (H₃O⁺ concentration = 10^(-pH)).
    • Example: If [H₃O⁺] = 4.5 x 10^(-6) M, then pH = -log[4.5 x 10^(-6)] = 5.35.
    • Follow significant figures when calculating to report pH values correctly.

  • Acid and Base Relationships

    • pKₗ = -log(Kₗ) relates to the concentration of ions in full equations and can be utilized to derive pH from ion concentrations.
    • pH + pOH = 14 at 25°C due to Kw (ion product of water).
    • If [H₃O⁺] increases, then [OH⁻] decreases (inverse relationships).
    • Example: If pH = 5.35, then pOH = 14 - 5.35 = 8.65.

  • Common Strong Acids

    • Know the six strong acids:
    1. Hydrochloric acid (HCl)
    2. Hydrobromic acid (HBr)
    3. Hydroiodic acid (HI)
    4. Nitric acid (HNO₃)
    5. Sulfuric acid (H₂SO₄)
    6. Perchloric acid (HClO₄)
    • Strong acids fully dissociate, meaning no molecules remain in the solution.

  • Calculating pH of Strong Acids

    • For strong acids, the concentration of H₃O⁺ equals the initial concentration of the acid.
    • Example: 0.75 M HNO₃ results in a pH of 0.12 since all HNO₃ fully dissociates into H₃O⁺ and NO₃⁻.

  • Common Strong Bases

    • Know common strong bases:
    1. Sodium hydroxide (NaOH)
    2. Potassium hydroxide (KOH)
    3. Calcium hydroxide (Ca(OH)₂)
    4. Strontium hydroxide (Sr(OH)₂)
    5. Barium hydroxide (Ba(OH)₂)
    • Strong bases fully dissociate in solution.

  • Calculating pH of Strong Bases

    • A strong base like Ca(OH)₂ produces two hydroxide ions per molecule.
    • Example: For 0.25 M Ca(OH)₂, OH⁻ concentration is 0.50 M, leading to a pOH = -log(0.50), and therefore pH = 14 - pOH.
    • For NaOH, with the same molarity, OH⁻ concentration equals the initial concentration of NaOH.

  • Summary

    • Understand the relationships between strong acids, strong bases, their dissociation, and calculations for pH/pOH in solution.
    • Practice conversions between hydrogen and hydroxide concentrations through log scale calculations and ensure adherence to significant figures in all calculations.