Endothermic
endothermic reaction; endothermic process. .
A process that absorbs heat. The enthalpy change for an endothermic process has a positive sign.
Enthalpy change
(H) enthalpy change.
Enthalpy (H) is
defined so that changes in enthalpy (
H)
are equal to the heat absorbed or released by a process running at
constant pressure. While changes in enthalpy can be measured using
calorimetry, absolute values of enthalpy usually cannot be determined.
Enthalpy is formally defined as H = U + PV, where U is the internal
energy, P is the pressure, and V is the volume.
Enthalpy of combustion
(Hc)
heat of combustion.
The change in
enthalpy
when one mole of compound is completely combusted. All carbon in the
compound is converted to CO2(g), all hydrogen to H2O(
),
all sulfur to SO2(g), and all nitrogen to N2(g).
Enthalpy of reaction
(Hrxn)
heat of reaction.
The heat absorbed or released by a chemical reaction running at constant pressure.
Entropy
(S)
.
Entropy is a measure of energy dispersal. Any spontaneous change disperses energy and increases entropy overall. For example, when water evaporates, the internal energy of the water is dispersed with the water vapor produced, corresponding to an increase in entropy
Exothermic
exothermic reaction; exothermic process.
A process that releases heat. The enthalpy change for an exothermic process is negative. Examples of exothermic processes are combustion reactions and neutralization reactions
Free energy
Energy that is actually available to do useful work. A decrease in free energy accompanies any spontaneous process. Free energy does not change for systems that are at equilibrium.
Free-energy change
a change in the "Gibbs free energy", the capacity of a system to do work.
Gibb’s free energy
(G) Gibbs' free energy.
A thermodynamic property devised by Josiah Willard Gibbs in 1876 to predict whether a process will occur spontaneously at constant pressure and temperature. Gibbs free energy G is defined as G = H - TS where H, T and S are the enthalpy, temperature, and entropy. Changes in G correspond to changes in free energy for processes occuring at constant temperature and pressure; the Gibbs free energy change corresponds to the maximum nonexpansion work that can be obtained under these conditions. The sign of DeltaG is negative for all spontaneous processes and zero for processes at equilibrium.
Heat
Heat is a transfer of energy that occurs when objects with different temperatures are placed into contact. Heat is a process, not a property of a material.
Hess’s law
law of constant heat summation; Hess's law of heat summation.
The heat released or
absorbed by a process is the same no matter how many steps the process
takes. For example, given a reaction A 
B,
Hess's law says that H
for the reaction is the same whether the reaction is written as A C B
or as A B.
This is the same as writing that H(A B)
= H(A C)
+ H(C B).
Joule
(J)
The SI unit of energy, equal to the work required to move a 1 kg mass against an opposing force of 1 newton. 1 J = 1 kg m2 s-2 = 4.184 calories.
Kelvin
(K)
The SI base unit of temperature, defined by assigning 273.16 K to the temperature at which steam, ice, and water are at equilibrium (called the triple point of water). The freezing point of water is 273.15 K.
Kinetic energy
The energy an object possesses by virtue of its motion. An object of mass m moving at velocity v has a kinetic energy of ½mv2.
Latent heat of phase change
Heat that is absorbed without causing a rise in temperature. For example, "latent heat of vaporization" refers to the amount of heat required to convert a liquid to vapor at a particular temperature.
Molar enthalpy of formation
Molar enthalpy of fusion
(Hfus)
heat of fusion; molar heat of fusion; molar enthalpy of fusion.
The change in enthalpy when one mole of solid melts to form one mole of liquid. Enthalpies of fusion are always positive because melting involves overcoming some of the intermolecular attractions in the solid.
Molar enthalpy of vaporization
(Hvap)
heat of vaporization
The change in enthalpy when one mole of liquid evaporates to form one mole of gas. Enthalpies of vaporization are always positive because vaporization involves overcoming most of the intermolecular attractions in the liquid.
Nonspontaneous
Specific heat
The heat required to raise the temperature of 1 g of a substance by 1°C is called the specific heat of the substance. Specific heat is an intensive property with units of J g-1 K-1.
Spontaneous
spontaneity; spontaneous process; spontaneous reaction.
A spontaneous process occurs because of internal forces; no external forces are required to keep the process going, although external forces may be required to get the process started. For example, the burning of wood is spontaneous once the fire is started. The combination of water and carbon dioxide to reform the wood and oxygen is NOT spontaneous!
Surroundings
In thermodynamics, the surroundings refer to the universe outside the system.
System
In thermodynamics, the system is the part of the universe that is of interest.
Temperature
Temperature is an intensive property associated with the hotness or coldness of an object. It determines the direction of spontaneous heat flow (always from hot to cold).
Thermochemical equation
An compact equation
representing a chemical reaction that describes both the stoichiometry
and the energetics of the reaction. For example, the thermochemical
equation CH4(g) + 2 O2(g) CO2(g)
+ 2 H2O(g), H
= -2220 kJ means "When one mole of gaseous CH4 is burned in
two moles of oxygen gas, one mole of CO2 gas and 2 moles of
steam are produced, and 2220 kilojoules of heat are released."
Thermochemistry
The study of heat absorbed or released during chemical changes.
Universe
Surroundings