Thermochemistry Lecture Review

A set of practice vocabulary flashcards covering the fundamental concepts of thermochemistry, energy types, calorimetry, and enthalpy transformations based on the provided lecture notes.

Thermochemistry

The study of heat change in chemical reactions.

System

The specific part of the universe that is of interest in a study.

Surroundings

The part of the universe in contact with the system, such as the water inside a calorimeter.

Energy

The capacity to supply heat or do work.

Radiant energy

Energy that comes from the sun and is the earth’s primary energy source.

Thermal energy

The energy associated with the random motion of atoms and molecules.

Chemical energy

The energy stored within the bonds of chemical substances.

Nuclear energy

The energy stored within the collection of neutrons and protons in the atom.

Potential energy

The energy available by virtue of an object’s position.

Kinetic energy

The energy measured in Joules and calculated by the formula $KE = \frac{1}{2}mv^2$, where $m$ is mass in kg and $v$ is velocity in m/sec.

Joule (J)

The SI unit of heat, work, and energy; $1 \text{ Joule} = 1 \text{ kg} \text{ m}^2 / \text{s}^2$.

Law of Conservation of Energy

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

Heat

The transfer of thermal energy between two bodies that are at different temperatures.

Temperature

A measure of the thermal energy.

Exothermic process

Any process that gives off heat and transfers thermal energy from the system to the surroundings.

Endothermic process

Any process in which heat has to be supplied to the system from the surroundings.

calorie (cal)

The historical unit of energy required to raise $1 \text{ gram}$ of water by $1 \text{ degree C}$ (or $1 \text{ K}$).

Calorie (Cal)

Equivalent to $1000 \text{ calories}$ or $1 \text{ kilocalorie (kcal)}$, used in measuring food energy content.

Heat Capacity (C)

The quantity of heat ($q$) that an object absorbs when its temperature increases $1 \text{ } ^{\circ}\text{C}$ (or $1 \text{ K}$).

Specific Heat Capacity (c)

Also known as Specific Heat, it is the quantity of heat required to raise the temperature of $1 \text{ gram}$ of substance by $1 \text{ } ^{\circ}\text{C}$ (or $1 \text{ K}$).

Calorimetry

A technique used to measure the amount of heat released or absorbed during a chemical reaction or physical change.

Calorimeter

An instrument used for measuring the heat changes in a system.

Cup calorimeter

A type of calorimeter used to measure heat changes for reactions occurring in solution.

Bomb calorimeter

A type of calorimeter used specifically to measure heat changes for combustion reactions.

Thermodynamics

The scientific study of the interconversion of heat and other kinds of energy.

State functions

Properties determined by the state of the system, regardless of how that condition was achieved, including energy, pressure, volume, and temperature.

First law of thermodynamics

The principle stating energy can be converted between forms but not created or destroyed, expressed as $\Delta E_{\text{system}} + \Delta E_{\text{surroundings}} = 0$.

Internal Energy Change (\Delta E)

The change in internal energy of a system defined by the formula $\Delta E = q + w$, where $q$ is heat and $w$ is work.

Work (w)

The product of external pressure and change in volume in a system, expressed as $w = -PΔV$ for an expanding gas.

Enthalpy (H)

A thermodynamic quantity used to quantify heat flow into or out of a system in a process occurring at constant pressure.

Enthalpy of reaction (\Delta H)

The heat given off or absorbed during a reaction at constant pressure, defined as $H(\text{products}) - H(\text{reactants})$.

Thermochemical Equation

A balanced chemical equation that includes the specified enthalpy change ($\Delta H$).

Standard Enthalpy of Combustion (\Delta H^{\circ}_{C})

The heat released by the complete combustion of $1 \text{ mole}$ of a compound at $25 \text{ } ^{\circ}\text{C}$ and $1 \text{ atm}$, yielding products also at those conditions.

Standard Enthalpy of Formation (\Delta H^{\circ}_{f})

The heat change that results when one mole of a compound is formed from its elements at a pressure of $1 \text{ atm}$.

Hess’s Law

When reactants are converted to products, the change in enthalpy is the same whether the reaction takes place in one step or a series of steps.

Enthalpy of solution (\Delta H_{\text{soln}})

The heat generated or absorbed when a certain amount of solute dissolves in a certain amount of solvent.

Bond Energy

The energy stored within chemical bonds, changes in which drive the overall enthalpy change of formation for compounds.

Specific Heat of Liquid Water

The specific heat capacity of liquid H2O is $4.184 \text{ J/g} \text{ } ^{\circ}\text{C}$ or $1.0 \text{ cal/g} \text{ } ^{\circ}\text{C}$.

Specific Heat of Iron

The specific heat capacity of iron is valued at $0.45 \text{ J/g} \text{ } ^{\circ}\text{C}$.

Specific Heat of Aluminum

The specific heat capacity of aluminum is valued at $0.897 \text{ J/g} \text{ } ^{\circ}\text{C}$.

Specific Heat of Lead

The specific heat capacity of lead is valued at $0.13 \text{ J/g} \text{ } ^{\circ}\text{C}$.

Heat of neutralization

The enthalpy change for the reaction $HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l)$, which is $-56.2 \text{ kJ/mol}$.

Heat of fusion

The enthalpy change for the physical process of melting, such as $H_2O(s) \rightarrow H_2O(l)$, equal to $6.01 \text{ kJ/mol}$.

Heat of vaporization

The enthalpy change for the process of boiling, such as $H_2O(l) \rightarrow H_2O(g)$, equal to $44.0 \text{ kJ/mol}$ at $25 \text{ } ^{\circ}\text{C}$.

Carbohydrate Fuel Value

The energy content for carbohydrates is $4.0 \text{ kcal/g}$ or $17 \text{ kJ/g}$.

Fat Fuel Value

The energy content for fats is $9.0 \text{ kcal/g}$ or $38 \text{ kJ/g}$.

Protein Fuel Value

The energy content for proteins is $4.0 \text{ kcal/g}$ or $17 \text{ kJ/g}$.

Element Standard Reference Point

The convention that the standard enthalpy of formation of any element in its most stable form is zero.

Standard enthalpy of formation of O3 (g)

For ozone gas at $25 \text{ } ^{\circ}\text{C}$, the $\Delta H^{\circ}_{f}$ is $142.2 \text{ kJ/mol}$.

Standard enthalpy of formation of Carbon (diamond)

For the diamond allotrope of carbon, the $\Delta H^{\circ}_{f}$ is $1.90 \text{ kJ/mol}$.

Standard enthalpy of formation of Carbon (graphite)

For the most stable graphite form of carbon, the $\Delta H^{\circ}_{f}$ is $0 \text{ kJ/mol}$.

Standard enthalpy of formation of H2O (l)

The enthalpy change for creating liquid water from elemental hydrogen and oxygen is $-285.8 \text{ kJ/mol}$.

Standard enthalpy of formation of CO2 (g)

The enthalpy change for creating carbon dioxide gas from its elements is $-393.5 \text{ kJ/mol}$.

Standard enthalpy of reaction formula

Calculated as $\Delta H^{\circ}_{rxn} = \sum n\Delta H^{\circ}_{f}(\text{products}) - \sum m\Delta H^{\circ}_{f}(\text{reactants})$.

Constant-Volume Calorimetry condition

Condition where volume is constant, work is zero, and $\Delta E = q_{rxn}$.

Constant-Pressure Calorimetry condition

Condition where pressure is constant, often in cup calorimeters, and $\Delta H = q_{rxn}$.

Exothermic Sign Convention

Signified by a negative enthalpy change ($\Delta H < 0$) as heat is released by the system.

Endothermic Sign Convention

Signified by a positive enthalpy change ($\Delta H > 0$) as heat is absorbed by the system.

Enthalpy and States of Matter Convention

Enthalpy values differ depending on whether the products are generated in gas, liquid, or solid states.

Mass (m)

The variable in kinetic energy ($KE = \frac{1}{2}mv^2$) measured in kilograms (kg).