Principles of Bioenergetics and Thermodynamics (1D) - MCAT Biological and Biochemical Foundations of Living Systems

$\Delta G = 0$. At equilibrium, the system's free energy is minimized, resulting in no net driving force for forward or reverse processes.

$\Delta G > 0$. Endergonic processes require free energy input from surroundings to proceed, characterized by a positive change in Gibbs free energy.

$\Delta G < 0$ at all $T$. With negative enthalpy and positive entropy, both terms in the Gibbs equation contribute to spontaneity regardless of temperature.

Spontaneous does not imply fast. Thermodynamic spontaneity assesses feasibility based on energy changes, independent of kinetic factors that determine reaction speed.

$K$ unchanged. Catalysts equally accelerate forward and reverse rates, maintaining the ratio that defines the equilibrium constant unchanged.

$\Delta G^\circ > 0$. An equilibrium constant less than one signifies reactant-favored reactions, resulting in positive standard free energy change.

$\Delta G < 0$. Exergonic processes release free energy to surroundings, driven by a negative change in Gibbs free energy under constant conditions.

Heat transferred at constant pressure. Enthalpy change measures heat exchange during reactions or phase changes when pressure remains constant, as per thermodynamic definitions.

$\Delta H < 0$. Exothermic processes release heat to surroundings, reducing the system's internal energy and thus yielding a negative enthalpy change.

$\Delta G$ unchanged; $E_a$ decreased. Catalysts lower the activation energy by stabilizing the transition state, but thermodynamics dictate that free energy difference remains constant.

$\Delta G = \Delta H - T\Delta S$. This equation quantifies the free energy available for work by balancing enthalpy and entropy contributions at constant temperature and pressure.

$\Delta G^\circ = -RT\ln K$. This relation connects thermodynamic favorability under standard conditions to the equilibrium constant via temperature and the gas constant.

$\Delta G = \Delta G^\circ + RT\ln Q$. This equation adjusts standard free energy for non-equilibrium conditions by incorporating the reaction quotient and thermal energy.

$\Delta S > 0$. Positive entropy change signifies increased molecular disorder or more accessible microstates, aligning with the second law of thermodynamics.

Energy barrier to reach the transition state. Activation energy quantifies the kinetic barrier reactants must overcome to form the high-energy transition state in a reaction pathway.

$\Delta G^\circ < 0$. An equilibrium constant greater than one indicates product-favored reactions, corresponding to negative standard free energy change.

$Q < K$. When the reaction quotient is below the equilibrium constant, the system shifts forward to minimize free energy.

$\Delta G = -w_{\max}$. Gibbs free energy change represents the maximum reversible work extractable, excluding pressure-volume work, under isothermal-isobaric conditions.

$\Delta G = 4\ \text{kJ/mol}$. Convert $\Delta S$ to kJ/mol·K, compute $T\Delta S = 6$ kJ/mol, then subtract from $\Delta H$ using the Gibbs free energy equation.

Negative at high $T$. At high temperatures, the large negative $-T\Delta S$ term overcomes the positive enthalpy, making free energy negative.

Negative at low $T$. At low temperatures, the negative enthalpy term dominates over the smaller positive $-T\Delta S$ contribution in the Gibbs equation.

$Q > K$. When the reaction quotient exceeds the equilibrium constant, the system favors the reverse direction to reach equilibrium.

$\Delta G > 0$ at all $T$. Positive enthalpy and negative entropy ensure the Gibbs free energy remains positive, preventing spontaneity at any temperature.