Thermochemistry and Chemical Kinetics Lecture Notes

Flashcards covering thermochemistry topics (enthalpy types, Hess's Law, lattice energy, entropy) and chemical kinetics topics (reaction rates, mechanisms, activation energy, and reaction orders).

Thermochemistry

The branch of chemistry studying heat energy changes in terms of absorption and release, accompanying physical transformation and chemical reactions.

Exothermic reactions

Processes that release heat into the surroundings, typically resulting in a negative enthalpy change ($-\Delta H$).

Endothermic reactions

Processes that absorb heat from the surroundings, resulting in a positive enthalpy change ($+\Delta H$).

Enthalpy ($H$)

A thermodynamic quantity representing the total heat content of a system, where the change ($\Delta H$) is equivalent to the heat exchanged at constant pressure.

Enthalpy of formation ($\Delta H_f^\circ$)

The change in enthalpy when one mole of a compound in its standard state is formed from its elements in their standard states.

Enthalpy of reactions ($\Delta H_{rxn}$)

The amount of heat evolved or absorbed when the moles of reactants and products as indicated by a balanced chemical equation react completely, calculated as $\Delta H_{rxn} = \Sigma \Delta H_{\text{products}} - \Sigma \Delta H_{\text{reactants}}$.

Enthalpy of Combustion ($\Delta H_c$)

The heat change when one mole of a substance is burnt completely in excess oxygen or air.

Heat energy ($q$) formula

The formula used to calculate heat energy experimentally: $q = mc\Delta T$, where $m$ is the mass of water, $c$ is the specific heat capacity, and $\Delta T$ is the temperature change.

Specific heat capacity of water

The value defined as $4184\,J/kgK$ or $4.184\,kJ/kgK$.

Enthalpy of Neutralization

The change in enthalpy of the system when $1\,g$ equivalent of an acid is completely neutralized by $1\,g$ equivalent of a base (or vice versa) in dilute solution.

Standard heat of neutralization for strong acid/base

A constant value always equal to $-57.1\,kJ$, representing the heat of formation of $1$ mole of water from H+ and OH- ions.

Enthalpy of Atomization

The energy change required to convert one mole of a substance into its individual constituent gaseous atoms; it is always an endothermic process.

Enthalpy of Ionization energy

The minimum energy required to remove the most loosely bonded electron from one mole of isolated gaseous atoms to form a cation.

Penetration effect

The proximity of an electron in an orbital to the nucleus, which is a factor affecting ionization enthalpy.

Shielding effect

The effect where inner electrons develop a shield for outer shell electrons, preventing the full nuclear charge from reaching them.

Effective nuclear charge ($Z_{\text{effective}}$) formula

$Z_{\text{effective}} = Z - S$, where $Z$ is the actual nuclear charge and $S$ is the screening constant.

Electron Affinity ($\Delta H_{EA}$)

The energy released when an electron is added to an atom to form an anion.

Enthalpy of Lattice Energy

The energy released when $1$ mole of an ionic crystalline compound is formed from gaseous ions, or the energy required to break $1$ mole of a solid ionic compound into gaseous ions.

Born-Haber cycle

An indirect method used to calculate lattice energy by considering sublimation, dissociation, ionization energy, and electron affinity.

Hess’ Law

The total enthalpy change of a chemical reaction is the same regardless of whether the reaction takes place in one step or a series of steps.

Enthalpy of Hydration

The amount of heat energy released when $1$ mole of gaseous ions is completely dissolved in water to form an infinitely dilute solution.

Enthalpy of ligation ($\Delta H_{lig}$)

The energy released during the first step of hydration as the solvent coordinates with the ion.

Energy of dispersion ($\Delta H_{disp}$)

The energy change associated with the second step of hydration, where hydrated ions are dispersed into the solvent.

Entropy ($S$)

A thermodynamic property that measures the randomness or disorder of the molecules in a system, measured in $J/K$.

Spontaneous reactions

Reactions that favor the formation of products at the conditions occurring, releasing free energy and having a negative change in Gibbs free energy ($-\Delta G$).

Gibbs Free Energy equation

$\Delta G^\circ = \Delta H^\circ - T\Delta S^\circ$, where spontaneity is determined by combining changes in enthalpy and entropy.

Chemical Kinetics

The branch of chemistry that studies reaction speeds (rates) and the step-by-step molecular pathways (mechanisms) by which reactions occur.

Reaction rate

The speed at which reactants are converted into products, expressed as the change in concentration over time: $Rate = \frac{\Delta [A]}{\Delta t}$.

Activation Energy ($E_a$)

The minimum amount of energy required to initiate a chemical reaction.

Catalysis

Substances that speed up a reaction by providing an alternative pathway with a lower activation energy without being consumed themselves.

Rate Law expression

$Rate = k[A]^p[B]^q$, where $p$ and $q$ are the orders of reactions for reactants $A$ and $B$, and $k$ is the rate constant.

Zero-order reaction

A reaction where the rate does not depend on the concentration of the reactants, such as the photochemical combination of $H_2$ and $Cl_2$.

Arrhenius equation

$k = Ae^{-\frac{E_a}{RT}}$, used to calculate activation energy where $A$ is the pre-exponential factor, $R$ is the gas constant ($8.314\,J/Kmol$), and $T$ is temperature.

Boltzmann Distribution curve

A graph showing how molecular energy or speed is distributed among sample particles at a given temperature.