Thermodynamics, Kinetics, and Redox Review

Vocabulary-style practice flashcards covering basic concepts, laws, and equations of Thermodynamics, Kinetics, and Redox chemistry as presented in the lecture notes.

Thermodynamics

The study of whether a reaction will happen, focusing on variables like Enthalpy ($\Delta H$), Gibbs Free Energy ($\Delta G$), and the Equilibrium Constant ($K$).

Kinetics

The study of how fast a reaction will occur, focusing on reaction rate and Activation Energy ($E_a$).

Redox

The study of electron flow, focusing on Standard Cell Potential ($E^{\circ}$) and the Nernst equation.

Enthalpy ($\Delta H$)

The heat energy change of a system at constant pressure; a "battle" where the system wants lower energy.

Exothermic

A process where $\Delta H$ is negative, heat is released, and products are at a lower energy than reactants.

Endothermic

A process where $\Delta H$ is positive, heat is absorbed, and products are at a higher energy than reactants.

Standard Formation Enthalpy

The enthalpy change for the formation of ONE mole of compound from its elements in their standard states, calculated as $\Delta H = \Sigma \text{products} - \Sigma \text{reactants}$.

Hess's Law

A law stating that if reactions are added together, their enthalpies are added as well; if a reaction is flipped, the sign of $\Delta H$ is changed.

Entropy ($\Delta S$)

The number of possible arrangements in a system; more arrangements (like more gas particles) lead to higher entropy ($\Delta S > 0$).

Gibbs Energy ($\Delta G$)

Determined by the equation $\Delta G = \Delta H - T \Delta S$, where the winner of the battle between enthalpy and entropy determines spontaneity.

Spontaneous

A condition indicated by $\Delta G < 0$, meaning the reaction will proceed without outside intervention.

Equilibrium (Thermodynamic)

The state where $\Delta G = 0$.

Non-spontaneous

A condition indicated by $\Delta G > 0$, meaning the reaction will not proceed spontaneously.

Gibbs Energy and Equilibrium Connection

The relationship expressed by the equation $\Delta G^{\circ} = -RT \ln(K)$, where a large $K$ corresponds to a negative $\Delta G$ and products being favored.

Rate

The change in concentration with time, expressed in units of $\text{mol\,L}^{-1}\text{s}^{-1}$.

Rate Law

The equation $\text{rate} = k[A]^n[B]^m$, where exponents are determined by comparing experiments changing one reactant at a time.

First Order Reaction

A reaction where doubling the concentration doubles the rate ($m = 1$), represented by a straight line on a $\ln[A]$ vs. $t$ graph.

Second Order Reaction

A reaction where doubling the concentration quadruples the rate ($m = 2$), represented by a straight line on a $1/[A]$ vs. $t$ graph.

Zero Order Reaction

A reaction where doubling the concentration results in the rate staying the same ($m = 0$), represented by a straight line on an $[A]$ vs. $t$ graph.

First Order Integrated Rate Equation

$\ln[A]_t = \ln[A]_0 - kt$.

Half Life Equation (First Order)

$t_{1/2} = \frac{0.693}{k}$, where $k$ is the first order rate constant.

Collision Theory

States that for a reaction to occur, there must be a collision, sufficient energy, and correct orientation.

Activation Energy ($E_a$)

The minimum energy needed for a reaction to occur.

Arrhenius Equation (Two Temperatures)

The equation $\ln\left(\frac{k_2}{k_1}\right) = -\frac{E_a}{R}\left(\frac{1}{T_2} - \frac{1}{T_1}\right)$, where temperatures must be converted from $^{\circ}\text{C}$ to $\text{K}$.

Intermediate

A species made in one step of a reaction mechanism and used in another, which cancels out in the overall reaction.

Rate Determining Step

The slowest step in a reaction mechanism that controls the overall rate and from which the rate law usually comes.

Oxidation

The loss of electrons.

Reduction

The gain of electrons.

Oxidising Agent

A substance that causes oxidation in another and is itself reduced.

Reducing Agent

A substance that causes reduction in another and is itself oxidised.

Galvanic Cells (Anode and Cathode)

Electrochemical cells where oxidation occurs at the anode and reduction occurs at the cathode.

Standard Cell Potential ($E^{\circ}_{\text{cell}}$)

Calculated as $E^{\circ}_{\text{cell}} = E^{\circ}_{\text{cathode}} - E^{\circ}_{\text{anode}}$.

Redox Connection to $\Delta G$

The relationship $\Delta G = -nFE^{\circ}$, where a positive $E^{\circ}$ yields a negative $\Delta G$ (spontaneous).

Redox Connection to $K$

The relationship $\Delta E^{\circ} = \frac{RT}{nF} \ln(K)$.

Nernst Equation

$E = E^{\circ} - \frac{RT}{nF} \ln(Q)$, used for non-standard conditions like variable concentrations or pH.

Biological Standard Conditions

The state where standard $[H^+] = 10^{-7}$ because $pH = 7$.

Ligands

Molecules or ions that donate an electron pair to a metal.

Electron Withdrawing Ligand

A ligand that pulls electron density away from the metal, making reduction easier and increasing $E^{\circ}$.

Electron Donating Ligand

A ligand that pushes electrons toward the metal, making reduction harder and decreasing $E^{\circ}$.