Chemistry 2B - Summer Session II 2026 Lecture Flashcards

General chemistry flashcards covering thermochemistry basics, laws of thermodynamics, calorimetry, and enthalpy from the Chem 2B Summer Session II 2026 course.

Thermochemistry

The study of the energy, heat, and work associated with chemical reactions and physical changes.

Internal Energy ($U$)

A microscopic energy contained in a substance, composed of thermal energy and chemical energy ($Internal Energy = Thermal Energy + Chemical Energy$).

Thermal Energy

The energy associated with random molecular motion (kinetic energy) that results in temperature and increases as temperature rises.

Chemical Energy

The potential energy associated with chemical bonds and intermolecular attractions; it changes during chemical reactions or physical changes (e.g., it increases when a bond breaks).

System

The specific part of the universe chosen for study where chemical reactions usually occur.

Surroundings

The part of the universe outside of the system, connected to the system via matter, heat, and work.

Open System

A system that can exchange both matter and energy (as heat or work) with its surroundings.

Closed System

A system that does not allow the flow of matter with surroundings but allows the transfer of energy as heat or work.

Isolated System

A system that cannot exchange any heat, work, or matter with its surroundings.

Heat ($q$)

Internal energy transferred between a system and its surroundings as a result of a temperature difference, flowing from hotter to colder bodies.

calorie ($cal$)

The quantity of heat required to change the temperature of one gram of water by one degree Celsius; equal to $4.184\,J$.

Joule ($J$)

The SI unit for heat ($q$).

Law of Conservation of Energy (Heat)

The principle that energy is neither created nor destroyed; heat gained by a system is lost by its surroundings ($q_{sys} + q_{surr} = 0$).

Heat Capacity ($C$)

An extensive property representing the quantity of heat required to change the temperature of a substance by $1\,^{\circ}C$; calculated as $C = \frac{q}{\Delta T}$.

Extensive Property

A property proportional to the system size, such as volume, mass, internal energy ($U$), or heat capacity.

Intensive Property

A property that does not depend on the size of the system, such as temperature, pressure, or density.

Specific Heat Capacity

An intensive property representing the heat capacity per gram of substance ($J/g\,^{\circ}C$).

Molar Heat Capacity

An intensive property representing the heat capacity per mole of substance ($J/mol\,^{\circ}C$).

Exothermic Reaction

A reaction that releases heat to the surroundings ($q_{rxn}$ is negative), converting chemical energy to thermal energy and increasing the system temperature.

Endothermic Reaction

A reaction that absorbs heat from the surroundings ($q_{rxn}$ is positive), converting thermal energy into chemical energy and decreasing the system temperature.

Bomb Calorimeter

A constant-volume calorimeter used to measure the heat of combustion reactions by observing the temperature change of the calorimeter.

Coffee-Cup Calorimeter

A calorimeter used in measuring the heat of solution reactions at constant pressure, typically using nested Styrofoam cups for insulation.

Pressure-Volume Work

Work done by external pressure through a volume change; calculated at constant pressure as $w = -P \times \Delta V$.

State Function

Any property that has a unique value for a specified state of a system and does not depend on how that state was established (e.g., $P$, $T$, $V$, $U$).

Path Dependent Function

A property which depends on how the system changes from the initial state to the final state (e.g., heat and work).

The First Law of Thermodynamics

States that the change in internal energy of a closed system ($\Delta U$) is equal to the heat supplied to the system ($q$) plus the work done on the system ($w$) ($\Delta U = q + w$).

Enthalpy ($H$)

An extensive state function defined as $H = U + PV$. At constant pressure, the change in enthalpy ($\Delta H$) is equal to the heat gained or lost ($q_p$).

Enthalpy of Fusion ($\Delta H_{fus}$)

The enthalpy change of one mole of a substance during melting; it is an endothermic process where $q_p > 0$.

Enthalpy of Vaporization ($\Delta H_{vap}$)

The enthalpy change of one mole of a substance during vaporization; it is an endothermic process where $q_p > 0$.

Standard State (Gas)

The state of a pure gas at exactly $1\,atm$ pressure.

Standard Enthalpy of Reaction ($\Delta H^{\circ}$)

The enthalpy change of a reaction in which all reactants and products are in their standard states (usually at $298.15\,K$).

Hess’s Law

States that if a process occurs in steps, the enthalpy change for the overall process is the sum of the enthalpy changes for the individual steps.

Standard Enthalpy of Formation ($\Delta H_f^{\circ}$)

The enthalpy change in the formation of one mole of a substance in its standard state from its component elements in their standard reference forms.