Chem Test

Chemistry Eighth Edition Chapter 16 Study Notes

Overview

• Authors: Jill K. Robinson, John E. McMurry, Robert C. Fay

• License: Copyright © 2020, 2016, 2012 Pearson Education, Inc., All Rights Reserved.

• Course: Chem 1212, Principles of Chemistry II

• Topics Covered:

  • Acid/Base Chemistry

  • Thermodynamics

  • Reaction Rates

Acid–Base Concepts: The Brønsted–Lowry Theory (1 of 4)

• Arrhenius Acid:

  • Definition: A substance that dissociates in water to produce hydrogen ions, H⁺.

• Arrhenius Base:

  • Definition: A substance that dissociates in water to produce hydroxide ions, OH⁻.

Acid–Base Concepts: The Brønsted–Lowry Theory (2 of 4)

• Brønsted-Lowry Acid:

  • Definition: A substance that can transfer hydrogen ions, H⁺. In other words, a proton donor.

• Brønsted-Lowry Base:

  • Definition: A substance that can accept hydrogen ions, H⁺. In other words, a proton acceptor.

• Conjugate Acid-Base Pairs:

  • Definition: Chemical species whose formulas differ only by one hydrogen ion, H⁺.

Acid–Base Concepts: The Brønsted–Lowry Theory (3 of 4)

• Acid-Dissociation Equilibrium:

  • Reaction:
    HA(aq) + H₂O(l) ⇌ H₃O^+(aq) + A^−(aq)

  • Where HA is the acid, H₂O is the base, H₃O⁺ is the conjugate acid, and A⁻ is the conjugate base.

Acid–Base Concepts: The Brønsted–Lowry Theory (4 of 4)

• Base-Dissociation Equilibrium:

  • Reaction:
    B(aq) + H₂O(l) ⇌ HB^+(aq) + OH⁻(aq)

  • Where B is the base, H₂O is the acid, HB⁺ is the conjugate acid, and OH⁻ is the conjugate base.

Hydrated Protons and Hydronium Ions

• Due to the high reactivity of the hydrogen ion, it is actually hydrated by one or more water molecules.

• For practical purposes, H⁺ is considered equivalent to H₃O⁺.

Acid Strength and Base Strength (1 of 3)

• With equal concentrations of reactants and products, the direction of reaction favored is:
  ext{Stronger Acid} + ext{Stronger Base}
  ightarrow ext{Weaker Acid} + ext{Weaker Base}

Acid Strength and Base Strength (2 of 3)

• Weak Acid:

  • Definition: An acid that is only partially dissociated in water; it acts as a weak electrolyte.

Acid Strength and Base Strength (3 of 3)

• 
  

• Table 16.1: Relative Strengths of Conjugate Acid-Base Pairs
  

  Acid (HA)	Base (A⁻)	Strength
  HClO₄	Cl⁻	Strong Acid
  H₂SO₄	HSO₄⁻	Strong Acid
  HNO₃	NO₃⁻	Strong Acid
  H₃O⁺	H₂O	Weak Acid
  H₂S	HS⁻	Weak Acid
  NH₄⁺	NH₃	Weak Base
  Factors That Affect Acid Strength (1 of 4)

  • Bond Polarity:

    • The strength of an acid is affected by the polarity of the H-A bond, which in turn is influenced by the electronegativity of element A.

  • Electronegativity Values:

    • CH₄: 2.5

    • NH₃: 3.0

    • H₂O: 3.5

    • HF: 4.0

  Factors That Affect Acid Strength (2 of 4)

  • Bond Strength:

    • The strength of the H-A bond determines the acidity such that:
      ext{Acid Strength: HF < HCl < HBr < HI}

  • Relative bond strengths in kcal/mol:

    • HF: 570

    • HCl: 432

    • HBr: 366

    • HI: 298

  Factors That Affect Acid Strength (3 of 4)

  • Oxoacids:

    • Acid strength is influenced by the electronegativity of Y in H-O-Y.

    • General trend in oxoacid strength:
      ext{H-O-I < H-O-Br < H-O-Cl}

  Factors That Affect Acid Strength (4 of 4)

  • Oxoacids:

    • Number of O atoms also impacts acid strength:
      ext{H-O-Cl < H-O-αl < H-O-αo < H-O-α}

  • Oxidation numbers for Cl:

    • Hypochlorous: +1

    • Chlorous: +3

    • Chloric: +5

    • Perchloric: +7

  Dissociation of Water (1 of 4)

  • General reaction for water dissociation:
    2 H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq)

  • Ion-Product Constant for Water:

    • At 25 °C:
      K_w = [H₃O^+][OH^-] = 1.0 imes 10^{-14}

    • Both [H₃O⁺] and [OH⁻] are equal to 1.0 imes 10^{-7} M .

  Dissociation of Water (2 of 4)

  • Equilibrium expression:
    K_w = [H₂O⁺][OH^-]

  Dissociation of Water (3 of 4)

  • Concentration scenarios:

    • Acidic solution:
      [H₃O^+] > [OH^-]

    • Neutral solution:
      [H₂O⁺] = [OH^-]

    • Basic solution:
      [H₃O^+] < [OH^-]

  Dissociation of Water (4 of 4)

  • pH scale:

    • Acidic: pH < 7

    • Neutral: pH = 7

    • Basic: pH > 7

  The pH Scale (1 of 4)

  • 
    

  • The pH is calculated using:
    ext{pH} = - ext{log}([H₃O^+])

  • 
    

  • Concentration levels of hydronium and corresponding pH:
    

    Concentration (M)	pH
    1.0	0
    1.0×10^-1	1
    1.0×10^-2	2
    1.0×10^-3	3
    …	…
    1.0×10^-7	7
    The pH Scale (2 of 4)

    • Example problem:

      • Given [H₃O⁺] = 0.0025 M, calculate pH:
        ext{pH} = - ext{log}(0.0025) ext{ gives } pH ext{ value.}

    The pH Scale (3 of 4)

    • Example problem:

      • Calculate pH for an aqueous ammonia solution with [OH⁻] = 0.0019 M:
        ext{pH} = 14 - ext{pOH} ext{ (use pOH to find pH).}

    The pH Scale (4 of 4)

    • Acid rain is a significant environmental concern as most aquatic life dies in waters with pH lower than 4.5-5.0.

    • Example problem: Calculate [H₃O⁺] for a lake with pH 4.5:
      [H₃O^+] = 10^{-4.5} M

    The pH in Solutions of Strong Acids and Strong Bases (1 of 2)

    • Example: Calculate pH of a 0.025 M HNO₃ solution:

      • Since HNO₃ is a strong acid,
        [H₃O^+] = [HNO₃] = 0.025 M

    • Resulting pH can be obtained using the pH formula.

    The pH in Solutions of Strong Acids and Strong Bases (2 of 2)

    • Example: Calculate pH of a solution made from 500 mg NaOH in 500 mL of water:
      ext{[OH⁻]} = rac{0.5g}{39.99 g/mol imes 0.5 L} = 0.025 M

    Equilibria in Solutions of Weak Acids

    • General reaction for weak acids:
      HA(aq) + H₂O(l) ⇌ H₃O⁺(aq) + A⁻(aq)

    • Acid-Dissociation Constant (Kₐ):
      Kₐ = rac{[H₃O^+][A^-]}{[HA]}

    Acid-Dissociation Constants at 25 °C (Table 16.2)

    • The proton transferred to water during acid dissociation is shown in red.

    • For weak acids, pKₐ is calculated by:
      pKₐ = - ext{log}Kₐ

    Calculating Equilibrium Concentrations of Weak Acids (1 of 2)

    • For example, to find pH of a 0.10 M HCN solution at 25 °C:

      • Use ICE table method to determine equilibrium concentrations.

      • Use the formula to calculate [H₃O⁺] and pH.

    Calculating Equilibrium Concentrations of Weak Acids (2 of 2)

    • Example Calculation: HCN dissociated as follows:
      HCN(aq) + H₂O(l) ⇌ H₃O⁺(aq) + CN⁻(aq)

    • Given:

      • Kₐ = 4.9 imes 10^{-10}

      • Solving for x gives the concentrations leading to a pH calculation.

    Equilibria in Solutions of Weak Acids

    • Example: Hypothetical 0.250 M HF solution has a pH of 2.036:

      • Use pH to find [H₃O⁺] and backtrack to find Kₐ and pKₐ.

    Percent Dissociation in Solutions of Weak Acids

    • Definition:
      ext{Percent Dissociation} = rac{[ ext{HA}]}{[ ext{HA}]_{initial}} imes 100 ext{%}

    • Example calculation illustrates how the percent dissociation increases as concentration decreases.

    Equilibria in Solutions of Weak Bases (1 of 4)

    • Example reaction for weak bases:
      B(aq) + H₂O(l) ⇌ BH⁺(aq) + OH⁻(aq)

    • Base-Dissociation Constant (Kb):
      Kb = rac{[BH^+][OH^-]}{[B]}

    Equilibria in Solutions of Weak Bases (2 of 4)

    • 
      

    • Table 16.4: Kb Values for Some Weak Bases
      

      Base	Formula	Kb	Conjugate Acid	Ka
      Ammonia	NH₃	1.8×10⁻⁵	NH₄⁺	5.6×10⁻¹⁰
      Equilibria in Solutions of Weak Bases (3 of 4)

      • Example: Calculate pH of a 0.40 M NH₃ solution at 25 °C.

      • Using ICE approach and equilibrium expression helps find pH.

      Equilibria in Solutions of Weak Bases (4 of 4)

      • Using approximation method for Kb allows finding of hydroxide concentration:
        [OH^-] = x, ext{ then } pH = - ext{log}([H₃O^+])

      Relation Between Ka and Kb (1 of 2)

      • Example involving NH₄⁺ dissociation provides a relationship between Kₐ and Kb: Kw = Ka imes Kb

      Relation Between Ka and Kb (2 of 2)

      • The relationship holds true for conjugate acid-base pairs, allowing interconversion between pKₐ and pKb values: pKₐ + pKb = 14.00

      Polyprotic Acids (1 of 4)

      • Example of polyprotic acid dissociation with carbonic acid: H₂CO₃(aq) + H₂O(l) ⇌ H₃O⁺(aq) + HCO₃⁻(aq)

        • Ka1 = 4.3×10⁻⁷

      Polyprotic Acids (2 of 4)

      • 
        

      • Table 16.3: Stepwise Dissociation Constants for Polyprotic Acids at 25 °C
        

        Acid	Formula	Ka1	Ka2	Ka3
        Acid–Base Properties of Salts (1 of 9)

        • Salts Yielding Neutral Solutions:

          • Cations from strong bases (e.g., Li⁺, Na⁺).

          • Anions from strong monoprotic acids (e.g., Cl⁻, NO₃⁻).

        Acid–Base Properties of Salts (2 of 9)

        • Salts Yielding Acidic Solutions:

          • Derived from weak bases and strong acids (e.g., NH₄Cl).

        Acid–Base Properties of Salts (3 of 9)

        • Small, highly charged metal cations can produce acidity.

        Acid–Base Properties of Salts (4 of 9)

        • Reaction with water can produce H₃O⁺, leading to an acidic solution.

        Acid–Base Properties of Salts (5 of 9)

        • Salts Yielding Basic Solutions:

          • Derived from strong bases and weak acids (e.g., NaCN).

        Acid–Base Properties of Salts (6 of 9)

        • Solutions depend on relative acid and base strength (e.g., ammonium carbonate).

        Acid–Base Properties of Salts (7 of 9)

        • Distinguishing acidic and basic properties based on strength ratios (Ka and Kb).

        Acid–Base Properties of Salts (8 of 9)

        • Theoretical calculations for salts can result in various pH outcomes.

        Acid–Base Properties of Salts (9 of 9)

        • 
          

• Table 16.5: Acid-Base Properties of Salts
  

  Type of Salt	Examples	Ions That React with Water	pH of Solution
  Lewis Acids and Bases (1 of 3)

  • Lewis Acid: An electron-pair acceptor.

  • Lewis Base: An electron-pair donor.

  Lewis Acids and Bases (2 of 3)

  • Example reactions involving copper ions and ammonia show formation of complexes.

  Lewis Acids and Bases (3 of 3)

  • Sulfuric acid as a Lewis acid accepts electron pairs from bases.

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