-
Attraction, Dipole-Dipole
Electrostatic (coulombic) attraction between charges of opposite sign of adjacent molecules with permanent dipole moments.
-
Calorimetry
The technique of measuring heat changes.
The are two basic types of calorimetry: constant pressure and constant volume.
-
Calorimetry, Constant Pressure
The technique of measuring heat change while disallowing changes in pressure.
This technique is often conducted using a "coffee cup" calorimeter. It measures changes in enthalpy.
-
Calorimetry, Constant Volume
The technique of measuring heat change while disallowing changes in volume.
This technique is conducted using a "bomb" calorimeter. It measures changes in internal energy.
-
Conditions, Standard
For gases: 1 atm pressure.
For solutions: 1 atm pressure, 1.00 M concentration.
For pure substances in a condensed state (solid or liquid): the standard state is the pure liquid or solid.
For elements: the standard state is the form in which the element exists (is most stable) under 1 atm pressure and the temperature of interest (usually 298 K).
-
Energy
The capacity to do work or to produce heat.
There are two broad categories of energy: potential energy and kinetic energy.
-
Energy, Bond Dissociation
The energy required to break 1 mole of bonds of a particular type.
Usually given the symbol D. Always has a positive sign.
-
Energy, Internal
The sum of the kinetic and potential energies of all the particles in the system.
The change in internal energy of a system can be expressed as the sum of heat flow into (q) and work performed on (w) the system.
-
Energy, Kinetic
The energy associated with bodies in motion.
-
Energy, Potential
The "built-in" future capacity to do work or to produce heat as a result of position or composition.
-
Enthalpy
Heat flow at constant pressure.
Because chemists most often conduct reactions at a constant pressure of 1 atm, enthalpy and heat of reaction are usually synonymous.
-
Enthalpy, Standard of Formation
The standard enthalpy of formation is the enthalpy change that results from the formation of 1 mol of a compound from its elements, with all components in their standard states.
The enthalpy of formation of an element in its standard state is defined to be zero.
-
Entropy
A measure of the randomness or disorder of a system
See also:
Thermodynamics, First Law Thermodynamics, Second Law
-
Forces, Intermolecular
Electrostatic attractions and repulsions between molecules.
Intermolecular attractive forces take the form of either dipole-dipole attractions or London dispersion forces.
-
Forces, London Dispersion
Attraction between fleeting dipoles induced in neighbouring atoms or molecules by momentary polarisations (distortions) of an electron cloud.
London dispersion forces exist between all molecules.
-
Gas
A fluid that expands to fill its container completely and uniformly, exerts pressure on its container, and mixes completely with any other gas.
-
Gas, Ideal
A gas for which the volume of the particles is tiny compared to the volume of the gas (i.e., to the volume of the container), and there are no forces of attraction or repulsion between the particles.
-
Gas, Real
A gas whose particles have significant volume and which experience appreciable inter-particulate forces.
Contrast with ideal gases.
-
Heat
The transfer of thermal energy from a body at higher temperature to one at lower temperature.
-
Heat Capacity, Specific
The energy required to raise the temperature of 1 kg of a substance by 1 K.
-
Hess's Law
The enthalpy of a reaction is independent of the number and order of steps taken to carry out the reaction.
-
Liquid
A fluid that is incompressible (has a fixed volume) and assumes the shape of the bottom of its container.
See also:
Gas
-
Moment, Dipole
The separation of positive and negative charge. Exists in any molecule in which there is a distribution of positive and negative charges that is not self-canceling by symmetry.
-
Pressure, Partial
The partial pressure of a gas in a mixture of gases is the pressure that the gas would exert if it alone occupied the entire volume.
-
State, Standard
The state (solid, liquid or gas) in which a substance is found under standard conditions.
-
State Function
A function (or property) that defines the present state of a system. It does not depend on past conditions, i.e., is path-independent.
-
Surroundings
That portion of the universe that is not the system.
-
System
The portion of the universe chosen for study.
There are three different types of systems: open, closed and isolated. Contrast with surroundings.
-
System, Closed
A system that allows exchange of energy with its surroundings, but not matter.
-
System, Isolated
A system that allows exchange of neither energy nor matter with its surroundings.
-
System, Open
A system that allows exchange of matter and energy with its surroundings.
-
Theory, Kinetic Molecular (KMT)
A model to describe the behaviour of gases based on the assumption that they are made up of microscopic particles that are in constant random motion.
KMT relies on the following assumptions: (1) The gas is ideal. (2) Collisions between particles are perfectly elastic (i.e., no energy loss in collision) and collisions of particles with the walls of the container are responsible for the observed pressure of the gas. (3) The average kinetic energy of the particles is directly proportional to the Kelvin temperature of the gas.
-
Thermodynamics
The study of the transformations of energy.
See also:
Energy Thermodynamics, First Law Energy, Internal Heat Work Enthalpy Entropy Thermodynamics, Second Law
-
Thermodynamics, First Law
Energy can neither be created nor destroyed; it can only be converted between one form and another.
-
Thermodynamics, Second Law
The entropy of the universe is constantly increaseing
See also:
Thermodynamics, First Law
-
Work
Work is force acting over distance.
In chemical reactions, work is often performed in expanding a gas against an external pressure.
