Thermochemistry and Enthalpy: Key Concepts for Chemistry

Thermochemistry

The study of the heat absorbed or released during chemical and physical changes.

Energy

The capacity to supply heat or do work.

Potential Energy

The energy an object has because of its relative position, composition, or condition.

Kinetic Energy

The energy that an object possesses because of its motion.

Law of Conservation of Energy

During a chemical or physical change, energy can be neither created nor destroyed, although its form can change.

Thermal Energy

A type of kinetic energy (KE) associated with the random motion of atoms and molecules.

Temperature

A quantitative measure of 'hot' or 'cold.'

Heat (q)

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

Heat Flow

The transfer of thermal energy that increases the thermal energy of one body and decreases the thermal energy of the other.

Thermal Equilibrium

The state when two substances reach the same temperature and their molecules have the same average kinetic energy.

Exothermic Process

A process that releases heat.

Endothermic Process

A process that absorbs heat.

Chemical Changes

Changes that involve the transformation of substances into different substances.

Physical Changes

Changes that do not alter the chemical composition of a substance.

Combustion Reaction

A chemical reaction that produces energy in the form of heat and light.

Metabolism of Food

The chemical reactions that occur in the body to convert food into energy.

Fossil Fuels

Natural substances like gasoline, natural gas, and coal that are burned for energy.

Thermal Energy Transfer

The process of thermal energy moving from a hotter object to a cooler one.

Calorimetry

The measurement of heat transfer during chemical reactions or physical changes.

Kinetic Energy Conversion

The process of converting kinetic energy into other forms of energy, such as electrical energy.

Average Kinetic Energy

The mean energy of the motion of particles in a substance.

Calorie

The amount of energy required to raise one gram of water by 1 °C (or 1 kelvin).

Calorie (large)

Commonly used in quantifying food energy content, it is a kilocalorie.

Joule (J)

The SI unit of heat, work, and energy, defined as the amount of energy used when a force of 1 newton moves an object 1 meter.

Heat capacity (C)

The quantity of heat (q) a body of matter absorbs or releases when it experiences a temperature change (ΔT) of 1 °C (or 1 Kelvin).

Specific heat capacity (c)

The quantity of heat required to raise the temperature of 1 gram of a substance by 1 °C (or 1 Kelvin).

Extensive property

A property that depends on the amount of matter present, such as heat capacity.

Intensive property

A property that does not depend on the amount of matter present, such as specific heat capacity.

Heat calculation

The amount of heat, q, entering or leaving a substance can be calculated based on temperature changes.

System

The substance or substances undergoing the chemical or physical change.

Thermal energy gain

If a substance gains thermal energy, Tfinal > Tinitial, then the value of q is positive.

Thermal energy loss

If a substance loses thermal energy, Tfinal < Tinitial, then the value of q is negative.

Specific heat of helium

5.193 J/g °C

Specific heat of water

4.184 J/g °C

Specific heat of ethanol

2.376 J/g °C

Specific heat of ice

2.093 J/g °C (at −10 °C)

Specific heat of water vapor

1.864 J/g °C

Specific heat of nitrogen

1.040 J/g °C

Specific heat of air

1.007 J/g °C

Specific heat of oxygen

0.918 J/g °C

Specific heat of aluminum

0.897 J/g °C

Specific heat of carbon dioxide

0.853 J/g °C

Surroundings

The other components of the measurement apparatus that serve to either provide heat to the system or absorb heat from the system.

Calorimeter

A device used to measure the amount of heat involved in a chemical or physical process.

Exothermic Reaction

A reaction where the heat produced is absorbed by the solution.

Endothermic Reaction

A reaction where the heat required for the reaction to occur is provided by the solution.

Calorimetric Determination

Involves either an exothermic process where heat, q, is negative, or an endothermic process where heat, q, is positive.

Heat Transfer

The process of thermal energy being transferred from the system to its surroundings or vice versa.

Coffee-cup Calorimeter

A calorimeter constructed of polystyrene cups, often used in general chemistry labs.

Commercial Calorimeters

Calorimeters of better design used in industry and for research.

Heat Exchange

The transfer of heat that must take place only between the system and the surroundings for accurate results.

Net Change in Heat

In a calorimetry process, the net change in heat is zero.

Heat Gained and Lost

The heat gained by one substance is equal to the heat lost by another, in magnitude but opposite in sign.

Example 5.3

A calculation involving a 360-g piece of rebar dropped into 425 mL of water at 24.0 °C, resulting in a final temperature of 42.7 °C.

Specific Heat of Steel

Assumed to be approximately the same as that for iron.

Density of Water

1.0 g/mL, so 425 mL of water equals 425 g.

Initial Temperature of Rebar

Calculated to be 248 °C based on heat transfer principles.

Energy Conservation in Reactions

During a chemical reaction, energy is neither created nor destroyed, resulting in no overall energy change.

Heat of Reaction

The heat produced or consumed by the reaction, qreaction, plus the heat absorbed or lost by the solution, qsolution, must add up to zero.

Bomb Calorimeter

A device used to measure the energy produced by reactions that yield large amounts of heat and gaseous products.

Whole-body Calorimeter

A calorimeter large enough to hold an individual human being, used to measure metabolism under various conditions.

Nutritional Calorie (Calorie)

The energy unit used to quantify the amount of energy derived from the metabolism of foods.

Specific Heat

The amount of heat required to raise the temperature of one gram of a substance by one degree Celsius.

Temperature Change

The difference in temperature before and after a reaction or process.

Heat Absorbed

The amount of heat taken in by the solution during a reaction.

Heat Produced

The amount of heat released by a reaction.

Combustion Reactions

Reactions that involve the burning of a substance in the presence of oxygen, producing heat and gaseous products.

Aqueous Solution

A solution in which water is the solvent.

Mass of Solution

The mass of the solution calculated based on its volume and density.

Chemical Hand Warmers

Devices that produce heat through exothermic reactions to warm hands.

Instant Cold Pack

A pack that becomes cold due to the endothermic dissolution of ammonium nitrate.

Energy Reservoirs

Substances that can store energy, allowing it to be added or removed.

Significant Figures

Digits in a number that contribute to its precision.

Maximum Temperature

The highest temperature reached during a reaction, such as 28.9 °C in the example.

Internal energy (U)

The total of all possible kinds of energy present in a substance.

1st Law of Thermodynamics

The relationship between the change in internal energy (DU), heat (q), and work (w).

Heat absorbed by the system

+q indicates that the system absorbs heat from the surroundings.

Work done on the system

+w indicates that the surroundings do work on the system.

Heat released by the system

-q indicates that the system releases heat to the surroundings.

Work done by the system

-w indicates that the system does work on the surroundings.

Change in internal energy (ΔU)

If heat flows into the system or work is done on the system, ΔU > 0; if heat flows out or work is done by the system, ΔU < 0.

Expansion work

A type of work that occurs when a system pushes back the surroundings against a restraining pressure.

Internal combustion engine example

In an internal combustion engine, the reaction is exothermic, releasing heat to the surroundings (-q) and doing work on the surroundings (-w).

State function

A quantity whose value depends only on the state of a system, not on how that state is reached.

Altitude as a state function

Altitude does not depend on the path taken to reach it, similar to internal energy.

Enthalpy (H)

Defined as the sum of a system's internal energy (U) and the product of its pressure (P) and volume (V).

Enthalpy change (ΔH)

The change in enthalpy that occurs during a chemical or physical process, commonly at constant pressure.

Pressure-volume work (w)

The work done by or on a system due to changes in pressure and volume.

Heat of reaction at constant pressure (qp)

The heat flow associated with a reaction occurring at constant pressure.

Thermochemical equation

An equation that represents changes in both matter and energy, with the enthalpy change (ΔH) written to the right of the balanced equation.

Example of a thermochemical equation

H2(g) + ½O2(g) ⟶ H2O(l) ΔH = −286kJ indicates the heat associated with the reaction.

Extensive property of ΔH

If the coefficients of the chemical equation are multiplied, the enthalpy change must also be multiplied by that same factor.

ΔH

The change in enthalpy for a reaction.

Standard state

A commonly accepted set of conditions used as a reference point for the determination of properties under other different conditions.

Standard state conditions

Pressure of 1 bar (or 1 atm; 1 bar = 0.987 atm) and solutions at 1 M concentration.

Standard enthalpy of combustion (ΔH°C)

The enthalpy change when exactly 1 mole of a substance burns under standard state conditions.

Enthalpy change symbol

The symbol used to indicate an enthalpy change for a process occurring under standard state conditions, often superscripted with 'o'.

C12H22O11

The chemical formula for sucrose.

KClO3

The chemical formula for potassium chlorate.

Limiting reactant

The reactant that is completely consumed in a reaction, limiting the amount of product formed.