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Section 14.1: Bronsted-Lowry Acids and Bases

Definitions

  • Acids: Proton (H⁺) donors. According to the Bronsted-Lowry theory, an acid is a substance that donates a proton to another substance.
  • Bases: Proton (H⁺) acceptors. According to the Bronsted-Lowry theory, a base is a substance that accepts a proton from another substance.

Key Relationships

  • Conjugate Acid-Base Pairs:
      - A conjugate base differs from its conjugate acid by one proton.
      - A conjugate acid differs from its conjugate base by one proton.

Reactions

  • Acid Ionization: The process by which an acid donates a proton.
  • Base Ionization: The process by which a base accepts a proton.

Section 14.2: Properties of Water

Amphoteric Compounds

  • Definition: Amphoteric compounds can function as either an acid or a base depending on the environment.
  • Example: Water (H₂O) is an example of an amphoteric compound.

Amphiprotic Compounds

  • Definition: Amphiprotic compounds can act as both a Bronsted-Lowry acid and a Bronsted-Lowry base.
  • Example: Water (H₂O) also qualifies as an amphiprotic compound due to its ability to both donate and accept protons.

Autoionization of Water

  • Equation: Water self-ionizes as follows:
       H_2O
    ightleftharpoons H^+ + OH^-
  • Ion Product Constant of Water at 25 °C:
    Kw=[H3O+][OH−]=1.0imes10−14K_w = [H_3O^+][OH^-] = 1.0 imes 10^{-14}

Learning Check 14.1

  1. Given: Ion product constant of water at 80 °C is Kw=2.4imes10−13K_w = 2.4 imes 10^{-13}.
  2. Calculate: Hydronium ion concentration [H₃O⁺] and hydroxide ion concentration [OH⁻].

Learning Check 14.2

  1. Task: Calculate the hydronium ion concentration in an aqueous solution where [OH⁻] = 0.001 M at 25 °C.

Learning Check 14.3

  1. Task: Write equations representing H₂PO₄⁻ acting as a base with HBr and as an acid with OH⁻.

Section 14.3: pH and pOH

Definitions

  • pH: A measure of the acidity of a solution, defined as:
    pH=−extlog[H3O+]pH = - ext{log}[H_3O^+]
  • pOH: A measure of the basicity of a solution, defined as:
    pOH=−extlog[OH−]pOH = - ext{log}[OH^-]

Learning Check 14.4

  1. Task: Calculate the pH of a solution where [H₃O⁺] = 2.0 x 10⁻⁶ M.

Learning Check 14.5

  1. Task: Calculate [H₃O⁺] for a solution with pH = -1.07.

Learning Check 14.6

  1. Task: Calculate both pH and pOH for a solution where [H₃O⁺] = 4 x 10⁻³ M.

Section 14.4: Relative Strengths of Acids and Bases

Increasing Acid Strength

  • Strong Acids: Acids that ionize completely in solution include:
      - Perchloric acid (HClO₄)
      - Sulfuric acid (H₂SO₄)
      - Hydrogen iodide (HI)
  • Weak Acids: These do not completely ionize in water.
  • Examples of Weak Acids (Ionization in water for selected acids):
      - Acetic acid (HC₂H₃O₂)
      - Hydrofluoric acid (HF)

Acid and Base Ionization Constants

  • Acid Ionization Constant (Kₐ): For a weak acid, defined by the formula:
    Ka=[H3O+][A−][HA]K_a = \frac{[H_3O^+][A^-]}{[HA]}
  • Example: For acetic acid (CH₃COOH), the ionization constant is: Ka=1.8imes10−5K_a = 1.8 imes 10^{-5}

Section 14.5: Polyprotic Acids

Definitions

  • Polyprotic Acids: Acids that can donate more than one proton. Examples include carbonic acid (H₂CO₃) and phosphoric acid (H₃PO₄).

Learning Check 14.19

  1. Task: Calculate [H₃O⁺], [HS⁻], and [S²⁻] in a 0.10 M solution of H₂S.

Section 14.6: Buffers

Definitions

  • Buffer: A solution composed of either a weak acid and its conjugate base or a weak base and its conjugate acid that resists changes in pH when small amounts of strong acid or base are added.
  • Buffer Capacity: The amount of acid or base that can be added before the pH changes significantly.
  • Henderson-Hasselbalch Equation:
    pH=pKa+extlog[A−][HA]pH = pK_a + ext{log} \frac{[A^-]}{[HA]}

Section 14.7: Acid-Base Titration

Definitions

  • Titration Curve: A plot of a solution property (such as pH) vs. the amount of added titrant.

Learning Check 14.21

  1. Task: Calculate the pH for the strong acid/strong base titration between 50.0 mL of 0.100 M HNO₃(aq) and 0.200 M NaOH(aq) at given volumes of added base.

Learning Check 14.22

  1. Task: Calculate the pH for the weak acid/strong base titration for a system described.

Section 15.2: Lewis Acids and Bases

Definitions

  • Lewis Acid: An electron pair acceptor.
  • Lewis Base: An electron pair donor.

Formation Constant Expression

  • Formula:
    Kf=[Ag(NH3)2+][Ag+][NH3]2K_f = \frac{[Ag(NH₃)_2^+] }{[Ag^+][NH₃]^2}

Dissociation Constant Expression

  • Formula:
    Kd=[Ag+][NH3]2[Ag(NH3)2+]K_d = \frac{[Ag^+][NH₃]^2}{[Ag(NH₃)_2^+]}