Study guide
This chapter is educational content only, not medical advice, and does not guarantee any exam outcome. It follows the AAMC's content outline for this section, which asks you to apply general chemistry, physics, organic chemistry, and biochemistry to living systems. The MCAT tests this material largely through passages: a stimulus (data, an experiment, or a described system) followed by questions that require you to combine outside knowledge with information given in the text, so the passage below is built the same way the real exam presents this content.
Passage: Enzyme-Catalyzed Rate Studies and a Buffered Equilibrium
A researcher, Dr. Amara Solberg, studies a hypothetical enzyme, ligase-X, that catalyzes the reversible reaction A + B <-> C in aqueous buffer at pH 7.4 and 37 C. She first measures initial reaction rate (v) at fixed enzyme concentration while varying substrate A concentration ([A]), holding B in excess. At low [A], v rises nearly linearly with [A]; at high [A], v plateaus near a maximum rate (Vmax), consistent with saturation kinetics, where every enzyme active site is occupied. From a Lineweaver-Burk (double-reciprocal) plot of 1/v versus 1/[A], she extracts Vmax = 20 micromol/min and Km = 4 mM, where Km is the substrate concentration at which v = Vmax/2. She next studies the reverse reaction and finds that at equilibrium, [C]_eq / ([A]_eq x [B]_eq) = Keq = 5.0 x 10^2 M^-1, indicating the equilibrium favors product C. Because Keq is large and DG-standard = -RT ln(Keq) is therefore negative, the forward reaction is favored under standard conditions, though the enzyme changes only the rate at which equilibrium is reached, not the equilibrium position itself. In a second experiment, Dr. Solberg buffers the reaction with a weak acid HA (pKa = 7.2) and its conjugate base A-, using the Henderson-Hasselbalch equation, pH = pKa + log([A-]/[HA]), to hold pH near the enzyme's optimum. She also runs the reaction in a sealed vessel connected to a piston, noting that when the piston compresses the gas-phase headspace, pressure rises as volume falls, consistent with Boyle's law (P1V1 = P2V2 at constant temperature), and that dissolved gas concentration in the buffer rises with partial pressure per Henry's law. Finally, she adds a fixed concentration of a competitive inhibitor structurally similar to A; the inhibitor raises the apparent Km without changing Vmax, because it competes for the same active site but can be out-competed by enough substrate.
Questions 1-5 (Passage-Based)
1. Based on the passage, at [A] = 4 mM (with B still in saturating excess), the initial reaction rate v is closest to: (A) 5 micromol/min (B) 10 micromol/min (C) 20 micromol/min (D) 40 micromol/min. Reasoning: by definition Km is the [A] at which v = Vmax/2, so v = 20/2 = 10 micromol/min, making (B) correct. 2. If the competitive inhibitor is added and [A] is increased sufficiently, the reaction rate will: (A) plateau below the original Vmax (B) still approach the original Vmax (C) approach zero (D) become independent of [A]. Correct answer (B) — competitive inhibition raises apparent Km but does not lower Vmax, since enough substrate outcompetes the inhibitor. 3. A buffer of HA/A- at pH 7.2 has: (A) mostly HA (B) mostly A- (C) equal HA and A- (D) no buffering capacity. Since pH equals pKa here, log([A-]/[HA]) = 0, so [A-] = [HA], giving answer (C); a buffer resists pH change best within about 1 pH unit of its pKa. 4. If the piston reduces headspace volume to one-third its original value at constant temperature, and initial pressure was 1.0 atm, the new pressure is approximately: (A) 0.33 atm (B) 1.0 atm (C) 3.0 atm (D) 9.0 atm. By Boyle's law, P1V1 = P2V2, so P2 = P1V1/V2 = 1.0 atm x 3 = 3.0 atm, giving answer (C). 5. Because Keq = 5.0 x 10^2 M^-1 favors product formation, DG-standard for the forward reaction A + B -> C is: (A) positive (B) negative (C) zero (D) cannot be determined. Since DG-standard = -RT ln(Keq) and Keq > 1, ln(Keq) is positive, making DG-standard negative, so answer (B) is correct; a negative DG-standard indicates the reaction is thermodynamically favorable (exergonic) under standard-state conditions, though actual spontaneity at any moment also depends on the reaction quotient Q, not DG-standard alone.
Thermodynamics: Energy, Entropy, and Free Energy
The MCAT treats thermodynamics as a set of bookkeeping rules for energy in chemical and biological systems. The first law states that energy is conserved: the change in internal energy DE of a system equals heat added (q) minus work done by the system (w), DE = q - w. Enthalpy (H) tracks heat exchanged at constant pressure, and a reaction is exothermic (releases heat, DH negative) or endothermic (absorbs heat, DH positive). Entropy (S) measures the dispersal of energy and matter, and the second law states that the total entropy of an isolated system tends to increase; biological systems can locally decrease entropy (as in building an ordered protein) only by increasing entropy elsewhere, usually as heat released to the surroundings. Gibbs free energy combines these ideas: DG = DH - TDS, where T is absolute temperature in Kelvin. A negative DG means a process is spontaneous (exergonic) as written, releasing usable energy; a positive DG means the process is nonspontaneous (endergonic) and requires an energy input to proceed, which is why cells couple endergonic reactions (such as biosynthesis) to the hydrolysis of ATP, an exergonic reaction. It matters that spontaneity (sign of DG) says nothing about reaction rate: a reaction can be thermodynamically favorable yet proceed too slowly to matter without a catalyst, which lowers the activation energy (Ea) of the rate-limiting transition state without changing DG or Keq. Hess's law allows DH for a multistep pathway to be calculated by summing the DH of each step, since enthalpy is a state function depending only on initial and final states, not the path taken. Calorimetry experiments, often appearing in passages, measure heat exchanged directly and let a specific heat capacity (q = mcDT) be used to find DH per mole of reactant.
Electrochemistry: Redox Reactions and Electrochemical Cells
Oxidation-reduction (redox) reactions transfer electrons between species, and the MCAT tests both balancing these reactions and predicting whether they proceed spontaneously. Oxidation is loss of electrons (and increase in oxidation number); reduction is gain of electrons (and decrease in oxidation number) — remembered by the mnemonic OIL RIG (oxidation is loss, reduction is gain). In a galvanic (voltaic) cell, a spontaneous redox reaction generates electrical current: oxidation occurs at the anode, reduction occurs at the cathode, and electrons flow through an external wire from anode to cathode while ions flow through a salt bridge to maintain electrical neutrality. Standard reduction potentials (E-standard) rank how readily a species is reduced; the overall cell potential is E-cell = E-cathode - E-anode, and a positive E-cell corresponds to a spontaneous reaction, linked to free energy by DG-standard = -nFE-standard, where n is moles of electrons transferred and F is Faraday's constant. In an electrolytic cell, an external power source drives a nonspontaneous redox reaction, as in electroplating or the recharging of a battery; here the same electrode-naming convention holds (oxidation at the anode, reduction at the cathode), but the required E-cell is negative without the external voltage, which is a common point of confusion worth memorizing directly. The Nernst equation, E = E-standard - (RT/nF) ln(Q), adjusts cell potential for non-standard concentrations, showing that a cell's voltage changes as reactants are consumed and products accumulate, approaching zero as the system nears equilibrium. Biological electron transport, such as the mitochondrial electron transport chain, applies these same principles: electrons move through a series of carriers with increasing reduction potential (increasing affinity for electrons), releasing free energy in steps that pump protons and ultimately drive ATP synthase.
Atomic and Molecular Structure
Atomic structure underlies the periodic trends and bonding behavior tested throughout the chemistry portions of the exam. An atom's electron configuration follows the Aufbau principle (filling lowest-energy orbitals first), the Pauli exclusion principle (no two electrons in an atom share all four quantum numbers), and Hund's rule (electrons fill degenerate orbitals singly before pairing). Periodic trends follow predictably from effective nuclear charge and shielding: atomic radius decreases left to right across a period (increasing nuclear charge pulls electrons in) and increases down a group (added electron shells); ionization energy (energy to remove an electron) and electronegativity (an atom's pull on shared electrons in a bond) both generally increase left to right and decrease down a group. Chemical bonds form to lower overall energy: ionic bonds involve electron transfer between atoms of large electronegativity difference, forming a lattice of oppositely charged ions, while covalent bonds involve electron sharing, which may be nonpolar (similar electronegativity) or polar (unequal sharing, creating partial charges). Molecular geometry follows VSEPR theory (valence shell electron pair repulsion), which predicts shape from the number of bonding and lone electron pairs around a central atom; for example, four bonding pairs and no lone pairs produce a tetrahedral geometry, while two bonding pairs and two lone pairs (as in water) produce a bent shape. Intermolecular forces, weaker than covalent bonds but essential to biological structure, include hydrogen bonding (a strong dipole interaction when H is bonded to N, O, or F), dipole-dipole forces, and London dispersion forces (present in all molecules, arising from transient electron distribution shifts); these forces govern boiling point trends, protein folding, and the double helix structure of DNA, where hydrogen bonds pair complementary bases.
Key terms
- Km (Michaelis constant)
- — The substrate concentration at which reaction velocity equals half of Vmax; a lower Km reflects higher apparent enzyme-substrate affinity.
- Competitive inhibition
- — Enzyme inhibition in which a molecule resembling the substrate binds the active site, raising apparent Km while leaving Vmax unchanged.
- Henderson-Hasselbalch equation
- — pH = pKa + log([conjugate base]/[weak acid]); used to calculate buffer pH and buffering capacity near the pKa.
- Gibbs free energy (DG)
- — DG = DH - TDS; a negative value indicates a spontaneous (exergonic) process, independent of how fast that process occurs.
- Activation energy (Ea)
- — The energy barrier a reaction must overcome to proceed; catalysts lower Ea and speed a reaction without changing DG or Keq.
- Standard reduction potential (E-standard)
- — A measure of a species' tendency to be reduced; cell potential E-cell = E-cathode - E-anode predicts spontaneity of a redox reaction.
- Nernst equation
- — E = E-standard - (RT/nF) ln(Q); adjusts electrochemical cell potential for non-standard concentrations of reactants and products.
- VSEPR theory
- — A model predicting molecular geometry from the repulsion between bonding and lone electron pairs around a central atom.
- Electronegativity
- — An atom's relative pull on shared electrons in a covalent bond; it increases left to right across a period and decreases down a group.
- Boyle's law
- — P1V1 = P2V2 at constant temperature and moles of gas; pressure and volume are inversely related.
- Henry's law
- — The concentration of a dissolved gas is proportional to its partial pressure above the liquid, at constant temperature.
- Hess's law
- — The total enthalpy change for a reaction is the sum of enthalpy changes for its individual steps, since enthalpy is a state function.
Exam tips
- In kinetics passages, locate Km and Vmax from the description or graph before answering — many questions are solved directly from the definition (v = Vmax/2 at [A] = Km) without new calculation.
- Do not confuse thermodynamic favorability (sign of DG, position of Keq) with reaction rate (Ea, catalyst effect) — the exam frequently tests this distinction directly.
- In any electrochemistry item, first identify which electrode is oxidation (anode) and which is reduction (cathode), then apply E-cell = E-cathode - E-anode before worrying about the Nernst equation.
- For gas-law and periodic-trend items, sketch the direction of change (increase/decrease) in words before choosing an answer choice — this catches sign errors from rushing through the algebra.
- When a passage gives you a defined equilibrium constant or rate law, plug the passage's own numbers into the standard equation rather than relying purely on memorized values, since MCAT passages often use invented systems to test transfer of the underlying principle.