Aim to answer every question before checking. Missed questions point you to the sections you need most.
1. (5.1) The rate of a chemical reaction is defined as:
A. Rate = -Δ[reactant]/Δt = +Δ[product]/Δt. Stoichiometric coefficients normalize rates across different species.
2. (5.1) Which factors typically increase reaction rate?
B. More energetic collisions, more frequent collisions, and lower activation energy all speed reactions.
3. (5.2) A rate law expresses:
C. Rate = k[A]ᵐ[B]ⁿ. Orders (m, n) come from experiment and need not match stoichiometric coefficients.
4. (5.2) For rate = k[A]²[B], the overall reaction order is:
D. Overall order = sum of individual orders. Units of k shift with order; here k has units M⁻²·s⁻¹.
5. (5.3) Rate laws are determined:
A. Typically via the method of initial rates: vary one reactant at a time and see how the rate changes.
6. (5.3) If doubling [A] quadruples the rate, the reaction order with respect to A is:
B. Factor by which rate changes = (factor on [A])ᵐ. 4 = 2ᵐ, so m = 2.
7. (5.4) For a zero-order reaction in A:
C. Example: saturated enzymatic reactions, where the enzyme is saturated and rate is capped. [A]ₜ = [A]₀ - kt.
8. (5.4) The linear plot for a zero-order reaction is:
D. Slope = -k. First-order gives linear ln[A] vs. t; second-order gives linear 1/[A] vs. t.
9. (5.5) For a first-order reaction:
A. Integrated law: [A]ₜ = [A]₀·e^(-kt). Radioactive decay and many enzymatic and drug-elimination processes follow first-order kinetics.
10. (5.5) The half-life of a first-order reaction:
B. This is why radioactive half-lives are constants. Clinical pharmacology reaches steady state after about 4-5 half-lives of a drug.
11. (5.6) A second-order reaction in A has:
C. Integrated law: 1/[A]ₜ - 1/[A]₀ = kt. A hallmark plot of second-order kinetics.
12. (5.6) The half-life for a second-order reaction:
D. Subsequent half-lives are longer as [A] drops. This is a key distinguishing feature between first- and second-order kinetics.
13. (5.7) The Arrhenius equation k = A·exp(-Ea/RT) implies:
A. A 10 °C rise roughly doubles many reaction rates. Ea is the temperature sensitivity: higher Ea reactions are more accelerated by heat.
14. (5.7) A catalyst speeds up a reaction by:
B. Catalysts do not change ΔG or the equilibrium. They provide a lower-energy pathway and are regenerated at the end.
15. (5.8) Collision theory requires:
C. Rate = (collision frequency) × (fraction with Ea) × (steric factor). All three terms matter.
16. (5.8) The steric factor in collision theory represents:
D. Even energetic collisions are unproductive if the molecules bump the wrong sides together. Bulky or complex reactants tend to have small steric factors.
17. (5.9) Transition state theory introduces:
A. The transition state sits at a saddle point on the potential energy surface and lives only femtoseconds.
18. (5.9) On a reaction coordinate diagram, the transition state is at:
B. The height of the peak above reactants is Ea. An intermediate would sit at a local minimum, not a peak.
19. (5.10) A homogeneous catalyst:
C. Heterogeneous catalysts are in a separate phase, like a Pt surface catalyzing a gas-phase reaction.
20. (5.10) Enzymes are biological catalysts that:
D. Enzymes are catalysts: they obey the same rules as inorganic catalysts but use active-site shape and chemistry to stabilize transition states.
21. (5.11) A reaction mechanism is:
A. For elementary steps (unlike the overall reaction), stoichiometric coefficients can be used directly as reaction orders.
22. (5.11) The rate-determining step in a multi-step mechanism:
B. Think of it as the narrowest point in a pipeline: no matter how fast the other steps are, flow is limited here.
23. (5.12) The steady-state approximation assumes:
C. The Michaelis-Menten derivation uses this approximation for the ES complex.
24. (5.12) The steady-state approximation is most useful when:
D. Formation rate ≈ destruction rate for that intermediate. This lets you solve for [intermediate] in terms of reactants.
