CHEM 4312 - Chapter 2: Physical Biochemistry Flashcards

Vocabulary flashcards covering bioenergetics, thermodynamics, noncovalent interactions, water properties, and acid-base chemistry.

Bioenergetics

The study of energy conversion in biological systems, such as the transformation of solar energy into chemical energy during photosynthesis.

Homeostasis

The maintenance of a highly ordered steady state in living organisms regarding temperature, concentration of biomolecules, and other parameters, which requires energy and delays reaching equilibrium.

Biological Equilibrium

The state reached when homeostasis is no longer maintained and macromolecules equilibrate with their surroundings.

Osmotic Work

A form of biological work required to maintain differential solute concentrations across biological membranes.

Chemical Work

A form of biological work involved in the biosynthesis and degradation of organic molecules within living organisms.

Mechanical Work

A form of biological work exemplified by muscle contraction in animals.

Reduction

A chemical process characterized by the gain of electrons, an increase in the number of hydrogens, or a decrease in the number of bonds to oxygen.

Oxidation

A chemical process characterized by the loss of electrons, a decrease in the number of hydrogens, or an increase in the number of bonds to oxygen.

Open System

A thermodynamic system in which both matter and energy are freely exchanged with its surroundings.

Closed System

A thermodynamic system in which only energy is exchanged with its surroundings, while matter is not.

Isolated System

A thermodynamic system in which neither matter nor energy is exchanged with its surroundings.

Zeroth Law of Thermodynamics

The thermodynamic law stating that if two bodies are in thermal equilibrium with a third body, they are also in thermal equilibrium with each other.

First Law of Thermodynamics

The principle of conservation of energy stating that energy cannot be created or destroyed, only converted from one form to another ($\Delta E = q - w$).

Second Law of Thermodynamics

The thermodynamic law stating that in all spontaneous processes, the total entropy of the universe increases ($\Delta S_{\text{universe}} = \Delta S_{\text{system}} + \Delta S_{\text{surroundings}} > 0$).

Third Law of Thermodynamics

The thermodynamic law stating that the entropy of a system approaches a constant value as the temperature approaches absolute zero.

Enthalpy ($H$)

A measure of the total heat content of a thermodynamic system, defined as $\Delta H = \Delta E + \Delta PV$.

Exothermic Reaction

A chemical reaction that releases heat into its surroundings and has a negative enthalpy change ($\Delta H < 0$).

Endothermic Reaction

A chemical reaction that absorbs heat from its surroundings and has a positive enthalpy change ($\Delta H > 0$).

Hess's Law

The law stating that the total enthalpy change for a chemical reaction is equal to the sum of the enthalpy changes for its individual steps, regardless of the pathway.

Entropy ($S$)

A measure of the degree of disorder or dispersal of energy within a thermodynamic system.

Gibbs Free Energy ($G$)

A thermodynamic potential defined as $G = H - TS$, representing the energy available to do work at constant temperature and pressure.

Exergonic Reaction

A reaction with a negative free energy change ($\Delta G < 0$) that releases energy and occurs spontaneously.

Endergonic Reaction

A reaction with a positive free energy change ($\Delta G > 0$) that absorbs energy and is non-spontaneous.

Standard Free Energy Change ($\Delta G^\circ$)

The free energy change of a reaction measured under standard conditions: constant pressure of $1\,\text{atm}$, room temperature of $298\,\text{K}$, and reactant/product concentrations of $1\,\text{M}$.

Biochemical Standard Free Energy Change (\Delta G^\circ')

The free energy change measured under standard biochemical state conditions: $pH = 7$, $[H_2O] = 55.5\,\text{M}$, and $[Mg^{2+}] = 1\,\text{mM}$.

Coupled Reactions

The process of pairing an endergonic reaction with an exergonic reaction so that the combined overall reaction has a negative free energy change ($\Delta G < 0$).

ATP Hydrolysis

A highly exergonic reaction (\Delta G^\circ' = -30.5\,kJ\,mol^{-1}) that releases energy due to reduced electrostatic repulsion between phosphoryl groups, resonance stabilization of inorganic phosphate, and greater solvation of ADP and phosphate.

Hydrogen Bond

A noncovalent interaction between an electronegative atom donor bonded to a hydrogen atom and another electronegative acceptor atom, with a bond strength of $12\text{--}30\,kJ\,mol^{-1}$.

Electrostatic Interactions

Attractive or repulsive noncovalent forces occurring between oppositely or similarly charged atoms/dipoles, with a typical strength of $20\,kJ\,mol^{-1}$.

Van der Waals Interactions

Weak, distance-dependent noncovalent forces that temporarily occur between dipoles of nearby electrically neutral molecules, with a strength of $0.4\text{--}4.0\,kJ\,mol^{-1}$.

Hydrophobic Effect

The tendency of nonpolar molecules to aggregate in aqueous solutions to minimize their exposed surface area, reducing ordered water in the hydration layer and increasing universe entropy.

Hydration Layer

A structured shell of ordered water molecules that surrounds nonpolar or polar solutes in an aqueous environment.

Grotthuss Mechanism

The process by which protons rapidly transfer through 'water wires'—chains of hydrogen-bonded water molecules—such as in cytochrome c oxidase.

Ion Product Constant of Water ($K_w$)

The equilibrium constant for the autoionization of water, equal to $1.0 \times 10^{-14}\,M^2 = [H^+][OH^-]$ at $25\,^\circ\text{C}$.

pH

A logarithmic scale representing hydrogen ion concentration, defined by the equation $pH = -\log[H^+]$.

Buffer

An aqueous solution consisting of a weak acid and its conjugate base that resists small changes in pH upon the addition of acid or base.

Henderson-Hasselbalch Equation

An equation relating the pH of a buffered solution to the $pK_a$ of the weak acid and concentrations of conjugate base and weak acid: $pH = pK_a + \log\left(\frac{[A^-]}{[HA]}\right)$.

pK_a

The negative logarithm of the acid dissociation constant ($K_a$), representing the pH at which a weak acid is 50% protonated and 50% deprotonated.

Titration Curve

A graph plotting pH versus added equivalents of titrant (e.g., base), used to determine $pK_a$ values and buffering regions of ionizable groups.

Midpoint of Buffering Region

The point on a titration curve where the concentration of a weak acid equals its conjugate base ($[HA] = [A^-]$), where $pH = pK_a$ and buffering capacity is highest.

Equivalence Point

The point in a titration where the exact stoichiometric quantity of titrant needed to react completely with the target acid or base has been added.

Bicarbonate Buffer System

The primary physiological buffer in human blood that regulates pH through the reversible equilibrium $CO_2 + H_2O \rightleftharpoons H_2CO_3 \rightleftharpoons HCO_3^- + H^+$.

Cytochrome c Oxidase

Complex IV of the mitochondrial electron transport chain that utilizes hydrogen-bonded water wires to relay protons across the inner mitochondrial membrane.

Glutamine Synthetase Reaction

An enzyme-catalyzed pathway that couples the synthesis of glutamine from glutamate and ammonium to the exergonic hydrolysis of ATP, achieving an overall \Delta G^\circ' = -16.3\,kJ\,mol^{-1}.

Selective Hydrophobic Barrier

A fundamental functional property of biological cell membranes formed by lipid bilayers that separates aqueous compartments with different entropy and solute concentrations.