Thermochemistry Lecture Notes

This set of vocabulary flashcards covers essential thermochemistry definitions including thermodynamic laws, energy types, work, enthalpy, and standard state measurement techniques based on the lecture material.

Thermodynamics

The study of energy and its changes or transformations.

Thermochemistry

The study of the relationships between chemical reactions and energy changes involving heat.

System

A portion of the universe on which we wish to focus our attention, separated from the surroundings by a boundary.

Surroundings

Everything in the universe that is not part of the defined system.

Boundary

The interface between a system and its surroundings across which energy exchange occurs.

Adiabatic Process

A process in which heat cannot be transferred across the interface between the system and its surroundings.

Isothermal Process

A process where thermal contact between the system and its surroundings is kept at constant temperature while a change occurs.

State of system

A particular set of conditions like pressure, temperature, number of moles, and physical forms that precisely define a system’s properties.

State function (or state variable)

Quantities like $P$, $V$, and $T$ whose values depend only on the current state, not on the prior history of the sample.

Equation of state

An equation form expressing the interrelations between state functions, such as the ideal gas equation $PV = nRT$.

Heat capacity

The quantity of heat energy required to raise the temperature of a given quantity of a substance by one degree Celsius (units: $J ext{ } ^ ext{o}C^{-1}$).

Specific heat

The heat capacity per gram; the quantity of heat necessary to raise the temperature of $1 ext{ g}$ of a substance by $1 ext{ } ^ ext{o}C$.

Molar specific heat

The heat needed to raise the temperature of $1 ext{ mole}$ of a substance by $1 ext{ } ^ ext{o}C$.

Specific heat of water

The quantity equal to $4.184 ext{ J } g^{-1} ext{ } ^ ext{o}C^{-1}$.

First Law of Thermodynamics

The law of conservation of energy stating that in any process, energy is neither created nor destroyed.

Internal Energy ($E$)

The total energy of a system resulting from the kinetic energy of its particles plus all potential energy from binding forces.

$$ΔE$$

The change in internal energy, defined as $E_{ ext{final}} - E_{ ext{initial}}$.

Heat ($q$)

The energy absorbed from the surroundings by a system during a change.

Work ($w$)

The energy removed from or added to a system through mechanical processes, such as expansion or contraction.

Internal Energy Equation

The form of the first law given as $ΔE = q + w$, where $q$ is heat and $w$ is work.

Work (Mechanical Definition)

Defined as $ext{force} imes ext{distance}$.

Positive $q$ (q > 0)

Indicates that heat is added to the system.

Negative $q$ (q < 0)

Indicates that heat is evolved (removed from) the system.

Positive $w$ (w > 0)

Indicates that the surroundings perform work on the system, causing it to contract.

Negative $w$ (w < 0)

Indicates that the system performs work on the surroundings, causing it to expand.

Path-dependent functions

Functions such as $q$ and $w$ whose values depend on the specific route taken between states.

Pressure ($P$)

Defined as force per unit area ($F/A$).

Pascal ($Pa$)

The SI unit of pressure, equal to one Newton per square metre ($1 ext{ N/m}^2$).

Reversible process

A process through which the maximum work available from any change is obtained.

Endothermic reaction

A reaction where energy is absorbed; $q$ and $ΔE$ are positive.

Exothermic reaction

A reaction where energy is evolved; $q$ and $ΔE$ are negative.

Bomb calorimeter

A device used to measure the heat of reaction under conditions of constant volume.

$$q_v$$

The heat of reaction at constant volume, which is equal to $ΔE$.

Extensive quantity

A property that depends on the amount of matter reacted, such as the total energy change $ΔE$.

Intensive quantity

A property characteristic of a reaction regardless of amount, such as heat evolved per mole.

Enthalpy ($H$)

A thermodynamic function defined as $H = E + PV$, also called heat content.

$$q_p$$

The heat absorbed or evolved at constant pressure, which is equal to the enthalpy change $ΔH$.

Pressure-volume work ($PV$ work)

Work calculated by the expression $(Δn)RT$ for chemical reactions involving changes in the number of moles of gas.

$$Δn$$

The change in the number of moles of gas on going from reactants to products.

Hess’s Law of Heat Summation

States that the net value of $ΔH$ for an overall process is the sum of the enthalpy changes of its individual steps.

Thermochemical equation

A chemical equation that includes the associated energy change and is interpreted on a mole basis.

Enthalpy diagram

A graphical illustration used to describe the nature of thermochemical changes.

Enthalpy of formation ($ΔH_f$)

The enthalpy change associated with the formation of one mole of a substance from its elements.

Standard states

A standard set of conditions chosen for recording data, usually $25 ext{ } ^ ext{o}C$ and $1 ext{ atm}$ pressure.

Standard heat of formation ($ΔH_f^o$)

The enthalpy of formation for a substance in its standard state.

$ΔH_f^o$ of elements

For an element in its natural, most stable form at $25 ext{ } ^ ext{o}C$ and $1 ext{ atm}$, this value is arbitrarily taken to be zero.

Heat of combustion

The heat liberated when one mole of a compound is burned, often measured in a bomb calorimeter to indirectly determine $ΔH_f^o$.

Bond energy

The energy required to break a chemical bond to produce neutral fragments.

Atomization energy

The energy needed to break all bonds in a molecule and reduce gaseous molecules to neutral gaseous atoms.

Average bond energy

The total atomization energy of a molecule divided by the number of bonds broken (e.g., $1662 / 4 = 415 ext{ kJ mol}^{-1}$ for $C-H$ in $CH_4$).