AP Chemistry Unit 6 Thermodynamics: Understanding Energy and Heat Transfer

System

The part of the universe being studied (e.g., the reacting chemicals in a beaker).

Surroundings

Everything outside the system (e.g., solution, cup, air, water bath) that can exchange energy with it.

Heat ($\text{q}$)

Energy transferred due to a temperature difference; symbolized by q.

Heat sign convention ($\text{q}$)

q > 0 means the system absorbs heat; q < 0 means the system releases heat.

Endothermic process

A process in which the system absorbs heat from the surroundings ($q > 0$; typically $\Delta H > 0$ at constant pressure).

Exothermic process

A process in which the system releases heat to the surroundings ($q < 0$; typically $\Delta H < 0$ at constant pressure).

Coffee-cup calorimeter

A constant-pressure calorimetry setup (often insulated) where temperature change of the solution is used to infer heat of the reaction.

Calorimetry

Experimental measurement of heat transfer by tracking temperature changes and using energy conservation.

Enthalpy ($\text{H}$)

A thermodynamic state function related to heat flow at constant pressure.

Enthalpy change ($\triangle H$)

Change in enthalpy between products and reactants; for many constant-pressure processes, $\Delta H$ corresponds to heat transferred.

Constant-pressure heat ($\text{q}_{\text{p}}$)

Heat transferred at constant pressure; in many AP Chemistry cases, $q_{\text{p}} = \triangle H$.

Energy diagram

Graph (often potential energy vs. reaction progress) showing reactant energy, product energy, activation barrier, and overall $\triangle H$.

Activation energy ($\text{E}_{\text{a}}$)

Energy barrier that must be overcome for a reaction to proceed; $E_{\text{a}}(\text{forward}) = E_{\text{TS}} - E_{\text{reactants}}$.

Transition state (TS)

Highest-energy point (peak) on an energy diagram; the unstable configuration at the top of the activation barrier.

Catalyst

Substance that provides an alternative pathway with lower activation energy (lower peak) without changing reactant/product energies or ΔH.

Thermal equilibrium

Condition when objects in contact reach the same final temperature; heat flow stops.

Calorimetry energy balance

In an ideal insulated setup, heat gained and lost sum to zero: $q_{\text{system}} + q_{\text{surroundings}} = 0$ (so $q_{\text{system}} = -q_{\text{surroundings}}$).

Specific heat capacity ($\text{c}$)

Heat required to raise the temperature of 1 g of a substance by 1°C (or 1 K); units often J/(g·°C).

Molar heat capacity ($\text{C}_{\text{m}}$)

Heat required to raise the temperature of 1 mol of a substance by 1°C (or 1 K).

Calorimeter constant ($\text{C}_{\text{cal}}$)

Heat absorbed by the calorimeter per degree of temperature change; $q_{\text{cal}} = \text{C}_{\text{cal}} \times \Delta T$ (units J/°C).

q = mcΔT

Equation for heat associated with a temperature change (no phase change), where $\Delta T = T_f - T_i$.

Latent heat ($\text{L}$)

Mass-based constant for phase-change energy (J/g) used in q = mL.

Enthalpy of fusion (\text{H}_{\text{fus}})

Heat required to melt 1 mol of a solid at its melting point; melting is endothermic.

Enthalpy of vaporization ($\triangle H_{\text{vap}}$)

Heat required to vaporize 1 mol of a liquid at its boiling point; vaporization is endothermic.

Heating curve

Plot of temperature vs. heat added showing sloped segments (use q = mcΔT) and plateaus during phase changes (use q = nΔH or q = mL).