Chapter 16 Thermochemistry

energy

the capacity to do work or produce heat.

chemical potential energy

the energy stored in the bonds of chemical compounds, which can be released or absorbed during a chemical reaction.

heat

thermal energy transferred between objects at different temperatures

thermochemistry

The study of the transfers of energy that accompany chemical reactions and physical changes

law of conservation of energy

The principle stating that energy can not be created or destroyed, only transformed/changed

specific heat capacity

The amount of energy required to raise the temperature of one gram of a substance by one °C or K

endothermic

• Energy is taken in from the surroundings

• Gains energy

• Feels cooler

• ∆H = +

exothermic

• Energy is given off to the surroundings

• Loses energy

• Feels warmer

• ∆H = -

system

(specific part of the universe being studied)

Open- Mass and Energy are Exchanged

Closed- Only energy is exchanged

Isolated- neither mass nor energy are Exchanged

eg) Chemical Reaction

surroundings

Everything outside of the system

eg) Beaker, Test Tube, Goggles

calorimeter

A device used to measure the temperature changes of a system in order to determine heat changes or specific heat

enthalpy

Amount of heat taken in or released by a system at constant pressure

heat of reaction

Amount of heat released/absorbed when a chemical reaction occurs usually at constant pressure.

+ —> endothermic; reactant

- —> exothermic; product

heat of combustion

the amount of heat released during complete combustion of one mole of substance

heat of fusion

Amount of energy per mole required to change a substance from the solid state to the liquid state at the solid’s melting point (no change in temperature)

units: kJ/mol

heat of solidification

Amount of heat released when 1 mole of a liquid freezes into a solid at its freezing point.

heat of vaporization

Amount of energy per mole required to change a substance from the liquid state to the vapor state at the liquid’s boiling point (no change in temperature)

units: kJ/mol

heat of condensation

amount of heat released when one mole of a gas changes into a liquid at it’s condensation point.

heat of solution

amount of heat absorbed or released when one mole of substance dissolves in a solvent

standard heat of formation

enthalpy change that occurs when one mole of a compound is formed from its elements in their standard state at 25°C and 1atm.

Hess’s law

The overall enthalpy change in a reaction is equal to the sum of enthalpy changes for the individual steps in the process.

specific heat

the amount of energy required to increase the temperature of one gram of a substance by one degree celsius. Written as Cp because it is measured at (Constant Pressure).