Thermodynamics and Enthalpy Review

Comprehensive vocabulary flashcards covering basic thermodynamics, various enthalpies of reaction, physical laws, system types, and process definitions based on the lecture notes.

Ionization Energy (Absolute Zero)

Energy required to extract an electron from atom/ion in gaseous state.

Electron Affinity

Energy released when a neutral gaseous atom accepts an extra electron to form a negative ion.

Born-Haber cycle

A sequence of steps used to calculate enthalpy where the importance is that the sum of the enthalpy changes round a cycle is zero.

Lattice enthalpy

The enthalpy change which occurs when one mole of an ionic compound dissociates into its ions in gaseous state, or vice-versa.

Enthalpy of hydration (Gaseous Ion)

The enthalpy change when 1 mole of the gaseous ion is hydrated in large amount of water to form aqueous ion, such as $Cl^-(g) + aq \rightarrow Cl^-(aq)$, $\Delta_{hyd}H^\circ$.

Enthalpy of solution ($\Delta_{sol}H^\circ$)

The enthalpy change determined by the selective values of the lattice enthalpy and enthalpy of hydration of ions as $\Delta_{sol}H^\circ = \Delta_{lattice}H^\circ + \Delta_{hyd}H^\circ$.

Bond enthalpy

The average of enthalpies required to dissociate a specific bond present in different gaseous compounds into free atoms or radicals in the gaseous state.

Bond dissociation enthalpy

The enthalpy required to dissociate a given bond of some specific compound, which depends on the nature of molecular species from which the atom is being separated.

Enthalpy of dilution

The amount of heat evolved or absorbed when a solution containing one mole of solute is diluted from one concentration to another.

Enthalpy of neutralization

The amount of heat released when one gram equivalent of an acid is neutralised by one gram equivalent of a base, or the formation of one mole of water; for strong acid/base it is always $-57.1\,kJ/mole$. or -13.7 Kcal/mol.

Enthalpy of ionization

Heat consumed in ionizing weak acids or weak bases that are not completely ionized in solution, defined per mole of the substance.

Greater the enthalpy of ionization of any weak acid/base …

WEAKER will be said acid/base

Enthalpy of transition

The enthalpy change when one mole of one allotropic form changes to another, such as $C(\text{graphite}) \rightarrow C(\text{diamond})$, $\Delta H = 1.90\,kJ\,mol^{-1}$.

Standard enthalpy of formation (Elements)

By convention, assigned a zero value for every element in its stable state of aggregation at $1\,bar$ pressure and specified temperature (usually $25^\circ\text{C}$).

Calorific value

The amount of heat produced in calorie or joule when one gram of a substance (food or fuel) is completely burnt or oxidised.

Lavosier and Laplace's law

Enthalpy of decomposition of a compound is equal to (-ve) the formation of the compound and only the sign changes.

Hess's Law

The law stating the total amount of heat evolved or absorbed in a chemical reaction is the same whether the reaction is carried out in one or several steps.

Kirchoff's equation

Equations used to find heat change of reaction at a temperature when it is known at another: $(\Delta H_2 - \Delta H_1) = \Delta C_p(T_2 - T_1)$ and $(\Delta E_2 - \Delta E_1) = \Delta C_v(T_2 - T_1)$.

Enthalpy (H)

A state function defined by the relation $H = U + PV$

Internal Energy (U)

The definite amount of energy associated with every system, comprising kinetic, potential, electronic, and nuclear energies; it is an extensive property and a state function.

First Law of Thermodynamics

Based on conservation of energy, states that energy can neither be created nor destroyed, and is expressed as $\Delta U = q + w$.

Heat Capacity [C]

The amount of heat required to raise the temperature of a system by $1^\circ\text{C}$, defined as $C = \frac{dq}{dT}$; it is an extensive property.

Molar Heat Capacity (C)

The amount of heat required to raise the temperature of 1 mole of substance by $1^\circ\text{C}$; it is an intensive property.

Specific Heat Capacity (c)

The amount of heat required to raise the temperature of $1\,g$ of substance by $1^\circ\text{C}$; it is an intensive property.

Calorie

The amount of heat required to raise the temperature of $1\,g$ of water by $1^\circ\text{C}$ ($14.5^\circ\text{C}$ to $15.5^\circ\text{C}$).

Work (Mechanical or PV Work)

Form of energy appearing when there is a change in the boundary of the system; calculated as $w = - \int P_{ext} dV$.

Open system

A system which can exchange both energy and matter with the surroundings, such as living organisms.

Closed system

A system which can exchange only energy but cannot exchange matter with the surroundings, such as a closed container.

Isolated system

A system which cannot exchange both energy and matter with the surroundings, such as water in a thermos flask.

Homogeneous system

A system having uniform nature throughout, made up of one phase only.

Heterogeneous system

A system not uniform throughout, consisting of more than one phase.

State Function

A property of a system which is dependent only on the initial and final state, and independent of the path followed (e.g., Temperature, Pressure, Volume, Enthalpy).

Path Function

Quantities dependent on the path or way the system has achieved a particular state, such as Heat, Work, and Heat Capacities.

Extensive Properties

Properties dependent on the mass or size of the system, such as Mass, Volume, Enthalpy ($H$), and Entropy ($S$).

Intensive Properties

Properties which are mass or size independent and not additive, such as Temperature, Pressure, Density, and Concentration.

Isothermal process

A process in which the temperature of the system remains constant ($\Delta T = 0$).

Isobaric process

A process in which the pressure of the system remains constant ($\Delta P = 0$).

Adiabatic process

A process in which no heat exchange takes place between the system and surroundings ($q = 0$).

Isochoric process

A process for which the volume of the system remains constant ($\Delta V = 0$).

Cyclic process

A combination of processes where the final state of the system is identical to the initial state; the net change in all thermodynamic properties is zero.

Mayer's formula

Refers to the relation $C_p - C_v = R$ for 1 mole of an ideal gas.

Poisson's ratio (\gamma)

The ratio of molar heat capacity at constant pressure to that at constant volume, defined as $\gamma = \frac{C_p}{C_v}$.

Resonance Energy

Difference between Enthalpy of formation/combustion/hydrogenation of an actual molecule V.S. of actual molecule.

Enthalpy of hydrogenation

One mole of given compound is hydrogenated (given $H$ +)

Heat of dilution depends on

1. original solution concentration.

2. amount of solvent added.

$\Delta H$ for Bond Enthalpies

$\Delta H$ reactants - $\Delta H$ products

Dissolution of salt solutions increase with temperature. Why?

Most ionic compounds are endothermic thus, $\Delta H$ sol +ve .

If lattice enthalpy is very high, then (dissolution is easier/harder)

Harder. Lattice enthalpy = strength between ions.

Spontaneous Reaction (with example)

Takes place on its own. Irreversible in nature. Eg. Natural processes.

Disorder

Randomness. Any macroscopic property of a system in the average of many microscopic forces which are random.

Second Law of Thermodynamics

$\Delta S$ _universe = $\Delta S$ _system + $\Delta S$ _surrounding > 0

$\Delta S$ _universe during reversible process

0

Entropy

Measure of disorder. State function. $dS=\frac{dq}{T}$ where $dq$ _rev = heat exchanged by system in a very small step (reversibly).