Chem 104 Test 2 Study Guide

Chem 104 - Test 2 Study Guide

Chapter 14: Acids and Bases Concepts

  • Acid Ionization: The process by which an acid donates a proton (H⁺) to a base.

  • Acid Ionization Constant (Ka): A measure of the strength of an acid in solution. It is defined as:
    Ka=rac[H3O+][A−][HA]K_a = rac{[H_3O^+][A^-]}{[HA]}

  • Acidic: Describes a solution with a pH less than 7.

  • Amphiprotic/Amphoteric: A substance that can either donate or accept a proton; for example, water (H2OH_2O) can act as both an acid and a base.

  • Autoionization: The process in which water ionizes to form hydronium (H3O+H_3O^+) and hydroxide ions (OH−OH^-):
    2H2O<br>ightleftharpoonsH3O++OH−2 H_2O <br>ightleftharpoons H_3O^+ + OH^-

  • Base Ionization: The process by which a base accepts a proton.

  • Base Ionization Constant (Kb): A measure of the strength of a base in solution, defined as:
    Kb=rac[HB+][OH−][B]K_b = rac{[HB^+][OH^-]}{[B]}

  • Basic: A solution with a pH greater than 7.

  • Brønsted-Lowry Acid: A substance that donates a proton.

  • Brønsted-Lowry Base: A substance that accepts a proton.

  • Buffer: A solution that resists changes in pH upon the addition of small amounts of acid or base.

  • Buffer Capacity: The ability of a buffer to maintain pH, dependent on the concentrations of the buffering components.

  • Conjugate Acid: The species formed when a Brønsted-Lowry base accepts a proton.

  • Conjugate Base: The species formed when a Brønsted-Lowry acid donates a proton.

  • Diprotic Acid: An acid that can donate two protons.

  • Diprotic Base: A base that can accept two protons.

  • Ion-Product Constant for Water (Kw): At 25°C, it is defined as:
    Kw=[H3O+][OH−]=1.0imes10−14K_w = [H_3O^+][OH^-] = 1.0 imes 10^{-14}

  • Leveling Effect of Water: Water is a weak acid and base, causing strong acids and bases to appear equally strong in aqueous solution.

  • Monoprotic Acid: An acid that can donate only one proton.

  • Neutral: A solution with a pH of 7, where [H3O+]=[OH−][H_3O^+] = [OH^-].

  • Oxyacid: An acid that contains oxygen, hydrogen, and another element (e.g., H2SO4H_2SO_4).

  • Percent Ionization: The percentage of a given acid that ionizes in solution, calculated as:
    ext{% ionization} = rac{[H_3O^+]_{eq}}{[HA]_0} imes 100

  • 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^-]

  • Stepwise Ionization: The process where polyprotic acids ionize stepwise, losing one proton at a time.

  • Triprotic Acid: An acid that can donate three protons.

  • Henderson-Hasselbalch Equation: Used to calculate the pH of a buffer solution:
    pH=pKa+extlograc[A−][HA]pH = pK_a + ext{log} rac{[A^-]}{[HA]}

Chapter 15: Solubility and Complexes

  • Solubility Product Constant (Ksp): An equilibrium constant for the dissolution of a sparingly soluble ionic compound:
    Ksp=[C]c[D]dK_{sp} = [C]^c[D]^d for a reaction aA+bB<br>ightleftharpoonscC+dDaA + bB <br>ightleftharpoons cC + dD

  • Common Ion Effect: The decrease in solubility of a salt when a common ion is added to the solution.

  • Complex Ion: A species formed from a central metal atom/ion bonded to one or more ligands.

  • Coordinate Covalent Bond: A bond formed when one atom donates both electrons for a bond.

  • Ligand: A molecule or ion that binds to a central metal atom to form a complex.

  • Formation (Stability) Constant (Kf): The equilibrium constant for the formation of a complex ion from its components.

  • Dissociation Constant (Kd): The equilibrium constant for the reverse reaction of the formation of a complex ion.

  • Amphoteric Hydroxide: Hydroxides that can react as either acids or bases, similar to amphiprotic substances.

Chapter 12: Thermodynamics Concepts

  • Spontaneous Process: A physical or chemical change that occurs without outside intervention.

  • Nonspontaneous Process: A process that requires energy input to occur.

  • Standard Entropy (S°): A measure of the disorder of a system at 1 atm and a specified temperature (usually 298 K).

  • Standard Entropy Change (ΔS°): The change in entropy when a reaction occurs under standard conditions.

  • Entropy (S): A measure of the dispersal of energy in a system.

  • 1st Law of Thermodynamics: Energy cannot be created or destroyed, only transformed.

  • 2nd Law of Thermodynamics: The total entropy of an isolated system can never decrease over time.

  • 3rd Law of Thermodynamics: As the temperature of a system approaches absolute zero, the entropy approaches a constant minimum.

  • Standard (Absolute) Entropy: The entropy of a substance at a reference state.

  • Free Energy (G): The energy in a system that can be used to do work. It is defined as:
    G=H−TSG = H - TS where H is enthalpy, T is absolute temperature, and S is entropy.

  • Standard Free Energy of Formation (ΔGf°): The change in free energy when one mole of a compound is formed from its elements in their standard states.

  • Thermodynamic Equilibrium Constant (K): Ratio of the concentrations of products to reactants at equilibrium.

  • Microstate (W): A distinct configuration of a system at the molecular level.

Important Equations to Memorize

  1. Equilibrium Constant:
    Kc=rac[C]c[D]d[A]a[B]bK_c = rac{[C]^c[D]^d}{[A]^a[B]^b}

  2. Ion-Product Constant for Water (at 25°C):
    Kw=[H3O+][OH−]=1.0⋅10−14K_{w}=[H_3O^{+}][OH^{-}]=1.0\cdot10^{-14}

  3. pH and pOH Relations:

    • pH=−log[H3O+]pH=-{log}[H_3O^{+}]

    • pH+pOH=14.00pH + pOH = 14.00

    • pOH=−log[OH−]pOH=-l{og}[OH^{-}]

  4. Acid and Base Ionization Constants:

    • Ka=rac[H3O+][A−][HA]K_a = rac{[H_3O^+][A^-]}{[HA]}

    • Kb=rac[HB+][OH−][B]K_b = rac{[HB^+][OH^-]}{[B]}

  5. Percent Ionization:
    ext{% Ionization} = rac{[H_3O^+]_{eq}}{[HA]_0} imes 100

  6. Relationship between Ka and Kb:
    KaKb=KwK_a K_b = K_w

  7. Henderson-Hasselbalch Equation:
    pH=pKa+extlograc[A−][HA]pH = pK_a + ext{log} rac{[A^-]}{[HA]}

  8. Second Law of Thermodynamics:
    extΔS=racqText{ΔS} = rac{q}{T}

  9. Free Energy and Equilibrium:

    • ΔGrxn°=ΣmΔGf°(prod)−ΣnΔGf°(reac)ΔG_{rxn}° = Σ_m ΔG_f° (prod) - Σ_n ΔG_f° (reac)

    • ΔG=ΔG°+RTextlnQΔG = ΔG° + RT ext{ln} Q

    • ΔG°=−RTextlnKΔG° = -RT ext{ln} K

Problem-Solving Focus Areas

Chapter 14

  • Understanding Brønsted-Lowry acids and bases

  • Calculating pH and pOH

  • Comparing relative strengths of acids and bases

  • Hydrolysis of salt solutions

  • Analyzing polyprotic acids

  • Buffers:

    • Key problems: 9, 10, 12, 13, 18, 19, 20, 21, 35, 36, 40, 41, 53, 54, 65, 66, 69, 70, 74, 75, 80, 81, 82, 83, 90, 91

Chapter 15

  • Precipitation and dissolution processes

  • Analyzing Lewis acids and bases

  • Understanding multiple equilibria

  • Key problems: 2, 6-11, 16, 17, 27, 28, 30, 31, 38, 39, 44, 64, 66-69, 94, 95, 98, 99, 100

Chapter 12

  • Understanding Spontaneity

  • Analyzing entropy

  • Applying 2nd and 3rd laws of thermodynamics

  • Understanding free energy calculations

  • Key problems: 3, 4, 12-13, 16-19, 25, 26, 33, 34, 36, 37, 39-42, 47-50