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Gen Chem (11)
| Question | Answer |
|---|---|
| What is dynamic equilibrium? | When the rate of the forward reaction equals the rate of the reverse reaction in a reversible reaction. Concentrations of reactants and products remain constant (no net change), but both reactions continue occurring simultaneously. |
| What are the two common misconceptions about equilibrium? | 1) That forward and reverse reactions stop at equilibrium — they continue at equal rates. 2) That conc of reactants and products are equal at eq — they don't need to be equal; the position depends on the nature of the reaction and external conditions. |
| What is the Law of Mass Action? | The rate of forward and reverse reactions is proportional to the product of the concentrations of reactants/products, each raised to the power of its stoichiometric coefficient in the balanced equation. |
| Forward Reaction Example | kf [A]^a [B]^b |
| Reverse Reaction Example | kr[C]^c[D]^d |
| What is the equilibrium constant expression (Keq) for: aA + bB ⇌ cC + dD? | Keq = [C]^c[D]^d / [A]^a[B]^b. Important: pure solids (s) and pure liquids (l) are EXCLUDED. Only aqueous (aq) and gaseous (g) species are included. |
| What does Keq < 1 mean? Keq = 1? Keq > 1? | Keq < 1 → [reactants] > [products] at equilibrium; reactants favored. Keq = 1 → [reactants] = [products] at equilibrium. Keq > 1 → [products] > [reactants] at equilibrium; products favored. |
| What is Kc vs. Kp? | Kc = equilibrium constant expressed in molar concentrations (for aqueous or gaseous systems). Kp = equilibrium constant expressed in partial pressures (for gaseous systems only). |
| What is the relationship between Gibbs free energy (ΔG°) and Keq? | ΔG° = −RT ln(Keq). When Keq < 1 → ΔG° > 0 (non-spontaneous, reactants favored). When Keq = 1 → ΔG° = 0 (at equilibrium). When Keq > 1 → ΔG° < 0 (spontaneous, products favored). |
| What is the Reaction Quotient (Q) and how does it differ from Keq? | Q uses the same formula as Keq but is calculated at ANY point during a reaction, not just at equilibrium. Keq only uses equilibrium concentrations. |
| What does comparing Q to Keq tell you? | Q < Keq → reaction proceeds forward (more products needed). Q = Keq → reaction is at equilibrium. Q > Keq → reaction proceeds in reverse (more reactants needed). |
| What is Le Chatelier's Principle? | If a system at equilibrium experiences a disturbance (change in concentration, pressure, or temperature), the reaction will shift to counteract the disturbance and restore a new equilibrium. |
| How does changing concentration affect equilibrium? | Adding reactants → shifts forward. Removing reactants → shifts reverse. Adding products → shifts reverse. Removing products → shifts forward. |
| How does changing pressure/volume affect a gaseous equilibrium? | Increasing pressure (or decreasing volume) → shifts toward the side with FEWER moles of gas. Decreasing pressure (or increasing volume) → shifts toward the side with MORE moles of gas. |
| How does temperature change affect equilibrium? | Endothermic (ΔH > 0): heat is a reactant → increasing temp shifts forward; decreasing temp shifts reverse. Exothermic (ΔH < 0): heat is a product → increasing temp shifts reverse; decreasing temp shifts forward. |
| Do catalysts or inert gases affect equilibrium? | No. Catalysts lower activation energy and speed up both forward and reverse reactions equally — equilibrium position is unchanged. Inert gases do not react with other species — equilibrium position is unchanged. |
| What is the solubility product constant (Ksp)? | Ksp describes the extent to which an ionic compound dissolves in water. Solids and liquids are excluded from the expression. Higher Ksp = more soluble; Lower Ksp = less soluble. |
| What is the solubility product quotient (Qsp), also called the ion product (IP)? | Qsp describes the current state of a solution at any moment. It uses the same expression as Ksp but with current (non-equilibrium) ion concentrations. |
| What do Qsp vs. Ksp comparisons tell you about a solution? | Qsp < Ksp → unsaturated (can dissolve more solute). Qsp = Ksp → saturated (no more solute can dissolve). Qsp > Ksp → supersaturated (precipitate forms). |
| Equations of Qsp and Ksp | C -> B + A [B]^b[A]^a |
| How do you calculate molar solubility using an ICE table? | Set initial concentrations of ions to 0 M, let x = moles dissolved at equilibrium, write equilibrium concentrations in terms of x, substitute into Ksp expression, and solve for x. x = molar solubility. |
| What is the common-ion effect? | When a solution already contains an ion that is also produced by a dissolving compound, the solubility of that compound is reduced. Example: AgCl dissolves less in a solution already containing Cl⁻. |
| What are amphoteric species? | Molecules or ions that can both donate AND accept a proton (act as both acid and base). Water (H₂O) is the classic example — it donates a proton to NH₃ (acts as acid) and accepts a proton from HCl (acts as base). |
| What is autoionization of water and the water dissociation constant (Kw)? | Water reacts with itself, transferring a proton to form H₃O⁺ (hydronium) and OH⁻ (hydroxide). Kw = [H₃O⁺][OH⁻] = 1×10⁻¹⁴ at 25°C. |
| What is the acid dissociation constant (Ka) and how is it expressed? | For HA + H₂O ⇌ H₃O⁺ + A⁻: Ka = [H₃O⁺][A⁻] / [HA]. pKa = −log(Ka). Higher Ka = stronger acid; lower Ka = weaker acid. |
| What is the base dissociation constant (Kb) and how is it expressed? | For B + H₂O ⇌ OH⁻ + HB⁺: Kb = [HB⁺][OH⁻] / [B]. pKb = −log(Kb). Higher Kb = stronger base; lower Kb = weaker base. |
| What is the relationship between Ka, Kb, and Kw? | Kw = Ka × Kb = 1×10⁻¹⁴ at 25°C. Therefore pKa + pKb = 14 at 25°C. |
| How does Le Chatelier's Principle apply to acid-base equilibrium (5 scenarios)? | 1) Adding strong acid → [H₃O⁺] increases → shifts LEFT. 2) Adding strong base → [H₃O⁺] decreases → shifts RIGHT. 3) Increasing [HA] → shifts RIGHT. 4) Decreasing [A⁻] → shifts RIGHT. 5) Increasing pH → [H₃O⁺] decreases → shifts RIGHT. |
| What is a conjugate acid-base pair? | They differ by exactly one H⁺. A conjugate base has one fewer H⁺ than the acid; a conjugate acid has one more H⁺ than the base. Example: OH⁻ → conjugate base is O²⁻ (remove one H⁺). |
| What are precipitation reactions? | Reactions where ions in solution combine to form an insoluble ionic solid (the precipitate). Example: Pb²⁺ + 2I⁻ → PbI₂(s). |
| What is the ICE table method for weak acid pH problems? | Set up Initial concentrations, track Change (−x for HA, +x for products), find Equilibrium expressions, substitute into Ka = x²/(initial concentration − x). If Ka is small, approximate (initial − x) ≈ initial, then solve for x = [H⁺], and pH = −log[H⁺]. |
Created by:
smurtab