How the MCAT tests kinetics

Kinetics falls in AAMC content category 5E — the principles of chemical thermodynamics and kinetics. The MCAT pairs it with enzyme catalysis in biochemistry, so the same vocabulary (Ea, catalyst, rate-determining step) reappears across passages. Expect data tables you must convert into a rate law, and energy diagrams you must interpret.

workspace_premium The key distinction

Thermodynamics answers “does it happen?”; kinetics answers “how fast?” A reaction with a very negative ΔG can still be kinetically trapped behind a large activation energy (think diamond → graphite).

Rate laws and reaction order

The rate law relates reaction rate to reactant concentrations: rate = k[A]ᵐ[B]ⁿ. The exponents (the order) must be found from experiment — usually the method of initial rates — not from stoichiometric coefficients.

rate = k[A]m[B]nRate lawm and n are the orders in A and B; their sum is the overall order. k is the temperature-dependent rate constant.
zero first second [A] → Rate →
Order sets the shape. Zero order: rate is independent of [A]. First order: rate is proportional to [A]. Second order: rate rises with [A]².
functions Worked example — find the rate law from initial rates

Doubling [A] doubles the rate; doubling [B] quadruples the rate. What is the rate law and overall order?

  1. Doubling [A] × 2 → rate × 2 = 2¹, so the order in A is 1.
  2. Doubling [B] × 2 → rate × 4 = 2², so the order in B is 2.
  3. Combine: rate = k[A]¹[B]²; overall order = 1 + 2 = 3.

Answer. rate = k[A][B]²; the reaction is third order overall.

Activation energy and the energy diagram

Every reaction must climb an energy barrier — the activation energy — to reach the transition state before rolling down to products. The reaction energy diagram displays Ea, the transition state, any intermediates, and the overall ΔH.

Ea ΔH < 0 transition state catalyzed (lower Ea) reactants products Reaction progress → Potential energy →
Reaction energy diagram. Ea is the climb from reactants to the transition-state peak; ΔH is the reactant-to-product drop. A catalyst (dashed) offers a lower-Ea path but does not change ΔH.
error Common trap

A catalyst lowers activation energy for both the forward and reverse reactions equally. It speeds the approach to equilibrium but never changes the position of equilibrium, ΔH, or ΔG.

What changes the rate

Anything that increases the frequency of successful, sufficiently energetic collisions speeds a reaction. Temperature is the most dramatic lever because of how molecular energies are distributed.

Ea T₁ (lower) T₂ (higher) can react Molecular kinetic energy → Fraction of molecules →
Maxwell–Boltzmann distribution. Raising the temperature broadens and shifts the curve to the right, so a much larger fraction of molecules exceed Ea and can react (shaded tail).
  • Temperature ↑ → more molecules exceed Ea and collisions are more frequent → rate rises sharply.
  • Concentration (or pressure for gases) ↑ → more frequent collisions → rate rises for any positive-order reactant.
  • Surface area ↑ (for heterogeneous reactions) → more contact → faster.
  • Catalyst → provides a lower-Ea pathway → faster without being consumed.
k = A·e−Ea/RTArrhenius equationk rises with temperature (T) and falls with a larger activation energy (Ea). A is the frequency factor; R is the gas constant.

Reaction mechanisms and the rate-determining step

Most reactions occur as a sequence of elementary steps. The overall rate can go no faster than its slowest step — the rate-determining step — so the experimental rate law reflects that step (and any fast equilibria before it).

  • For an elementary step (and only an elementary step), the order equals the molecularity — the coefficients do give the exponents.
  • An intermediate is produced in one step and consumed in a later one; it never appears in the overall rate law.
  • A valid mechanism must sum to the overall balanced equation and be consistent with the observed rate law.
workspace_premium Link to biochemistry

Enzymes are biological catalysts: they lower Ea, are not consumed, and follow the same kinetic logic — the foundation of Michaelis–Menten kinetics tested in the bio/biochem section.

Integrated rate laws in brief

For single-reactant reactions, the integrated rate laws let you find concentration over time and give a linear plot that identifies the order.

OrderLinear plotHalf-life behavior
Zero[A] vs tShortens as reaction proceeds
Firstln[A] vs tConstant (independent of [A])
Second1/[A] vs tLengthens as [A] falls

A constant half-life is the fingerprint of a first-order process (also true for radioactive decay).

How to study it — and the traps to avoid

Kinetics rewards a few procedures. Practice extracting rate laws from data and reading energy diagrams until both are automatic.

  1. Never read reaction order off the balanced equation — use initial-rate data.
  2. Keep kinetics and thermodynamics separate: fast ≠ favorable.
  3. Know that a catalyst lowers Ea only; ΔH, ΔG, and K are untouched.
  4. Recognize a constant half-life as first order.
error Highest-yield trap

Adding a catalyst or changing temperature changes the rate, but only temperature changes the equilibrium constant K. A catalyst gets you to the same equilibrium faster — it does not shift it.