Chemical Energy and Thermochemistry (Chapter 7)

A set of vocabulary flashcards covering fundamental terms and concepts from Chapter 7 on chemical energy, thermodynamics, enthalpy, calorimetry, Hess’s law, standard enthalpies, and bond energies.

Energy

The capacity to do work or produce heat.

Law of Conservation of Energy

Principle stating that energy can be converted from one form to another but cannot be created or destroyed.

Potential Energy (PE)

Stored energy due to position or composition; calculated mechanically as mgh.

Kinetic Energy (KE)

Energy of motion, calculated as ½ m v².

Work (w)

Energy transfer that occurs when a force moves an object through a distance.

Heat (q)

Energy transferred between a system and its surroundings because of a temperature difference.

Elastic Collision

Collision in which total kinetic energy is conserved (e.g., gas‐phase molecular impacts).

Inelastic Collision

Collision where some kinetic energy converts to other forms such as heat, sound, or deformation.

Chemical Potential Energy

Energy stored in the relative positions and bonds of atoms within molecules.

Thermodynamics

The study of energy and its interconversions.

Thermochemistry

Branch of thermodynamics focusing on heat involved in chemical processes.

System (thermodynamics)

The part of the universe chosen for study; everything else is the surroundings.

Surroundings

All matter and energy outside the system that can exchange energy or matter with it.

Open System

System that can exchange both energy and matter with its surroundings.

Closed System

System that exchanges energy but not matter with its surroundings.

Isolated System

System that exchanges neither energy nor matter with its surroundings.

Exothermic Reaction

Chemical process that releases heat to the surroundings (ΔH < 0, q < 0).

Endothermic Reaction

Chemical process that absorbs heat from the surroundings (ΔH > 0, q > 0).

Heat of Reaction (q_rxn)

Quantity of heat exchanged between a reacting system and its surroundings.

PV Work

Work associated with volume change against an external pressure; calculated as −P ΔV.

Internal Energy (E)

Total energy contained within a system, including thermal, chemical, and electronic forms.

State Function

Property that depends only on the current state of a system, not on the path taken (e.g., E, H, P, V).

Path Function

Property that depends on the specific process or route taken (e.g., q, w).

First Law of Thermodynamics

Energy of the universe is constant; for a system ΔE = q + w.

Sign Convention for Heat

q > 0 when heat enters the system; q < 0 when heat leaves the system.

Sign Convention for Work

w > 0 when work is done on the system; w < 0 when work is done by the system.

Enthalpy (H)

Thermodynamic quantity defined as H = E + P V, useful at constant pressure.

Enthalpy Change (ΔH)

Heat exchanged at constant pressure; ΔH = Hproducts − Hreactants.

Extensive Property

Property that depends on the amount of substance present (e.g., enthalpy, mass).

Heat Capacity (C)

Quantity of heat required to raise an object’s temperature by 1 °C (units J °C⁻¹).

Specific Heat Capacity (C_s)

Heat required to raise the temperature of 1 g of a substance by 1 °C (J g⁻¹ °C⁻¹).

Molar Heat Capacity (C_m)

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

Calorimetry

Experimental technique for measuring heat changes in physical or chemical processes.

Coffee-Cup Calorimeter

Simple constant-pressure calorimeter typically made from nested Styrofoam cups; used for solution reactions.

Bomb Calorimeter

Sealed, constant-volume calorimeter used to measure heats of combustion.

Constant-Pressure Calorimetry

Calorimetry performed at fixed pressure; heat measured equals ΔH of the process.

Constant-Volume Calorimetry

Calorimetry performed at fixed volume; heat measured equals ΔE of the process.

Hess’s Law

Total enthalpy change for a reaction is the same whether it occurs in one step or several steps.

Standard State

Reference condition of 1 bar pressure, 25 °C, and for solutions 1 M concentration.

Standard Enthalpy of Formation (ΔH_f°)

Enthalpy change for forming 1 mol of a compound from its elements in their standard states.

Standard Enthalpy of Reaction (ΔH_rxn°)

Enthalpy change for a reaction with all reactants and products in their standard states.

Bond-Dissociation Energy (D)

Energy required to break one mole of a specific covalent bond in the gas phase.

Average Bond Energy

Mean of bond-dissociation energies for a given bond across different molecules.

First-Law Equation

Mathematical statement ΔE = q + w expressing energy conservation for a system.

Bond Energy Equation

ΔE = Σ nD (bonds broken) − Σ nD (bonds formed); estimates reaction energy from bond energies.