MCAT Thermodynamics and Kinetics
MCAT Thermodynamics and Kinetics covers how energy flows through reactions and how fast those reactions proceed — a foundational MCAT chemistry topic tested heavily on the Chem/Phys section. Expect to use ΔG = ΔH - TΔS to judge spontaneity, link Keq to ΔG°, apply Arrhenius to temperature-rate relationships, and use Le Chatelier's principle in industrial and biological contexts. Both standalone calculations and passage-embedded data tables show up regularly.
Passages give you reaction-coordinate diagrams, calorimetry setups, or concentration-versus-time graphs and ask you to extract meaning, not just recall formulas. Integrated rate law questions require identifying reaction order from which plot linearizes, and initial-rate tables require a systematic approach to pull out rate constants. The math is moderate, but the conceptual traps are where MCAT thermodynamics and kinetics questions claim most of their points.
The misconception that costs students the most points is conflating ΔG° with ΔG — they are not the same thing, and the exam builds questions around that confusion. Students also consistently confuse kinetics with thermodynamics (a catalyst lowers Ea but does not shift K) and mix up Q versus K when predicting reaction direction. Sign conventions for work differ between physics and chemistry conventions, and the MCAT exploits that difference. If your MCAT general chemistry review does not drill these distinctions, this area will hurt you.
First Law of Thermodynamics and Internal Energy
Sign conventions for heat and work (q, w) differ by discipline — know the chemistry convention cold.
- Confuses physics sign convention (w > 0 for work by system) with chemistry convention (w > 0 for work on system)
- Confuses adiabatic (q = 0, T changes) with isothermal (ΔT = 0, T constant)
Enthalpy and Hess's Law
Hess's law, bond energies, and ΔH = q at constant pressure are the three calculation routes tested.
- Forgets to flip the sign of ΔH when reversing a step reaction in Hess's law
- Inverts the bond-energy formula, subtracting broken from formed instead of formed from broken
Calorimetry and Heat Capacity
Bomb versus coffee-cup setups measure different quantities; heating-curve plateaus signal phase transitions, not missing heat.
- Confuses bomb calorimeter (constant V, measures ΔU) with coffee-cup calorimeter (constant P, measures ΔH)
- Fails to negate q when transferring heat between reaction and calorimeter
Entropy and the Second Law
Predicting ΔS sign for phase changes and dissolution requires thinking about the universe, not just the system.
- Applies the second law to the system alone rather than to the universe
- Assumes dissolution always yields positive ΔS without considering hydration-shell ordering
Gibbs Free Energy and Spontaneity
Four ΔH/ΔS sign combinations determine when temperature controls spontaneity — and ΔG° ≠ ΔG.
- Anchors spontaneity judgment on ΔH sign alone without computing the TΔS term
- Inverts the effect of low product concentration on ΔG, saying it makes ΔG more positive
Reaction Rates and Rate Laws
Reaction orders come from experimental data only; stoichiometric coefficients are irrelevant to the rate law.
- Reads reaction orders directly from stoichiometric coefficients instead of experimental data
- Ignores stoichiometric coefficients when relating individual species rates to the overall reaction rate
Integrated Rate Laws and Reaction Order
Which concentration plot gives a straight line tells you the order — and only first-order has a concentration-independent half-life.
- Assigns first-order kinetics to a linear [A] vs. time plot instead of zero-order
- Generalizes the concentration-independent half-life property of first-order reactions to all orders
Arrhenius Equation and Activation Energy
Temperature raises rate exponentially, not linearly; two-temperature Arrhenius form lets you compute rate constant ratios.
- Assumes a linear relationship between temperature and reaction rate instead of an exponential one
- Conflates kinetic activation energy with thermodynamic equilibrium position
Catalysis (Homogeneous, Heterogeneous, Enzymatic)
Alternate pathways lower Ea equally for forward and reverse; equilibrium position and K are untouched.
- Thinks catalysts shift equilibrium toward products by selectively lowering forward activation energy
- Believes catalysts are consumed in the reaction because they appear in mechanistic steps
Chemical Equilibrium and Keq
ICE tables, the Kp-Kc conversion, and excluding pure solids/liquids from K expressions are the core mechanics.
- Includes pure solids and liquids in the Keq expression
- Treats Kp and Kc as interchangeable without accounting for the Δn gas correction
Le Chatelier's Principle
Stress direction, endo/exo character of the reaction, and Δn(gas) all determine the equilibrium shift — K itself never changes.
- Confuses temperature increase with a universal forward shift, ignoring endo/exo directionality
- Applies pressure-shift rule even when Δn(gas) = 0, predicting a shift where none occurs
Reaction Quotient (Q) vs Equilibrium Constant (K)
Comparing Q to K predicts shift direction; ΔG = ΔG° + RT ln Q connects non-standard conditions to spontaneity.
- Thinks Q has a different formula than K rather than the same formula applied at non-equilibrium conditions
- Reverses the direction of shift when Q > K, predicting a forward shift instead of reverse
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