Section One: Terms
work (w) – causing matter to move against an opposing force
potential energy – energy derived from the composition, position, or condition of a particle or
system
kinetic energy - energy of a moving body, in joules, equal to ½mv2 (where m = mass and v =
velocity)
thermal energy – kinetic energy associated with the random motion of atoms and molecules
temperature – quantitative measure of “hot” or “cold”
heat (q) - transfer of thermal energy between two bodies at different temperatures
endothermic - reaction or physical change that absorbs heat
exothermic - reaction or physical change that releases heat
heat capacity (C) - quantity of heat (q) a body of matter absorbs or releases when it experiences a
temperature change (ΔT) of 1 degree Celsius
specific heat capacity (c) - quantity of heat required to raise the temperature of 1 gram of a
substance by 1 degree Celsius
First law of Thermodynamics - internal energy of a system (U) changes due to heat flow (q) in or
out of the system or work (w) done on or by the system
law of conservation of energy - during a chemical or physical change, energy can be neither
created nor destroyed, although its form can be changed
enthalpy (H) - sum of a system’s internal energy (U) and the mathematical product of its pressure
(P) and volume (V)
Standard enthalpy of combustion (ΔHC°) - the enthalpy change when 1 mole of a substance burns
(combines vigorously with oxygen) under standard state conditions
Standard enthalpy of formation (ΔHf °) - the enthalpy change when 1 mole of a substance is
formed from free elements in their most stable states under standard conditions
Hess’s Law - If a process can be written as the sum of several stepwise processes, the enthalpy
change of the total process equals the sum of the enthalpy changes of the various steps
bond energy - energy required to break a specific covalent bond in one mole of gaseous
molecules
viscosity - a measure of a liquid’s resistance to flow
surface tension - energy required to increase the surface area of a liquid
capillary action - when a liquid flows within a porous material due to the attraction of the liquid
molecules to the surface of the material and to other liquid molecules
cohesive force - force of attraction between identical molecules
adhesive force - force of attraction between molecules of different chemical identities
vapor pressure – pressure exerted by vapor in equilibrium with a liquid in a closed container at a
given temperature
, boiling point – temperature at which its equilibrium vapor pressure is equal to the pressure
exerted on the liquid by its gaseous surroundings (atmospheric pressure for open containers)
melting or freezing point – temperature at which liquid and solid of a given substance are in
equilibrium
triple point – temperature at which all 3 states can coexist in equilibrium
critical point – temperature and pressure above which a gas cannot be condensed into a liquid
supercritical fluid - substance at a temperature and pressure higher than its critical point; exhibits
properties intermediate between those of gaseous and liquid states
precipitation reactions – two aqueous ionic compounds react to form a solid precipitate
acid-base reactions – H+ ion is transferred between reactants
redox reactions – reactions where oxidation number of two elements change over the course of
the reaction
pressure – force exerted on a given area P = F/A
Dalton’s law of partial pressures - The total pressure of a mixture of ideal gases is equal to the
sum of the partial pressures of the component gases
diffusion - movement of an atom or molecule from a region of relatively high concentration to
one of relatively low concentration
effusion - transfer of gaseous atoms or molecules from a container to a vacuum through very
small openings
valence bond theory - describes a covalent bond as the overlap of half-filled atomic orbitals
(each containing a single electron) that yield a pair of electrons shared between the two bonded
atoms
orbital overlap – portion of one orbital occupies same space as a second orbital
sigma (σ) bonds - result from the overlap between two s orbitals, 1 s and 1 p orbital, or two p
orbitals
pi (π) bonds - result from the overlap between two side by side p orbitals
atomic orbital theory – describes distribution of electrons in atoms
molecular orbital theory – describes distribution of electrons in molecules
bonding orbitals – result of in-phase combination of orbitals where electron density if between
nuclei and hold atoms together
antibonding orbitals – result of out of phase combination of orbitals where electron density is
opposite of internuclear space
molecular orbital energy diagrams – show relative energy levels of atomic & molecular orbitals
bond order - number of pairs of electrons between two atoms (# bonds between two atoms in a
Lewis structure)
s-p mixing - change that causes σ p orbitals to be less stable than πp orbitals due to the mixing of
s and p-based molecular orbitals of similar energies
formula mass – sum of average atomic masses of all atoms in a chemical formula
aqueous solution – a homogeneous mixture of a substance mixed with water
solute – solution component present in a concentration less than that of the solvent
work (w) – causing matter to move against an opposing force
potential energy – energy derived from the composition, position, or condition of a particle or
system
kinetic energy - energy of a moving body, in joules, equal to ½mv2 (where m = mass and v =
velocity)
thermal energy – kinetic energy associated with the random motion of atoms and molecules
temperature – quantitative measure of “hot” or “cold”
heat (q) - transfer of thermal energy between two bodies at different temperatures
endothermic - reaction or physical change that absorbs heat
exothermic - reaction or physical change that releases heat
heat capacity (C) - quantity of heat (q) a body of matter absorbs or releases when it experiences a
temperature change (ΔT) of 1 degree Celsius
specific heat capacity (c) - quantity of heat required to raise the temperature of 1 gram of a
substance by 1 degree Celsius
First law of Thermodynamics - internal energy of a system (U) changes due to heat flow (q) in or
out of the system or work (w) done on or by the system
law of conservation of energy - during a chemical or physical change, energy can be neither
created nor destroyed, although its form can be changed
enthalpy (H) - sum of a system’s internal energy (U) and the mathematical product of its pressure
(P) and volume (V)
Standard enthalpy of combustion (ΔHC°) - the enthalpy change when 1 mole of a substance burns
(combines vigorously with oxygen) under standard state conditions
Standard enthalpy of formation (ΔHf °) - the enthalpy change when 1 mole of a substance is
formed from free elements in their most stable states under standard conditions
Hess’s Law - If a process can be written as the sum of several stepwise processes, the enthalpy
change of the total process equals the sum of the enthalpy changes of the various steps
bond energy - energy required to break a specific covalent bond in one mole of gaseous
molecules
viscosity - a measure of a liquid’s resistance to flow
surface tension - energy required to increase the surface area of a liquid
capillary action - when a liquid flows within a porous material due to the attraction of the liquid
molecules to the surface of the material and to other liquid molecules
cohesive force - force of attraction between identical molecules
adhesive force - force of attraction between molecules of different chemical identities
vapor pressure – pressure exerted by vapor in equilibrium with a liquid in a closed container at a
given temperature
, boiling point – temperature at which its equilibrium vapor pressure is equal to the pressure
exerted on the liquid by its gaseous surroundings (atmospheric pressure for open containers)
melting or freezing point – temperature at which liquid and solid of a given substance are in
equilibrium
triple point – temperature at which all 3 states can coexist in equilibrium
critical point – temperature and pressure above which a gas cannot be condensed into a liquid
supercritical fluid - substance at a temperature and pressure higher than its critical point; exhibits
properties intermediate between those of gaseous and liquid states
precipitation reactions – two aqueous ionic compounds react to form a solid precipitate
acid-base reactions – H+ ion is transferred between reactants
redox reactions – reactions where oxidation number of two elements change over the course of
the reaction
pressure – force exerted on a given area P = F/A
Dalton’s law of partial pressures - The total pressure of a mixture of ideal gases is equal to the
sum of the partial pressures of the component gases
diffusion - movement of an atom or molecule from a region of relatively high concentration to
one of relatively low concentration
effusion - transfer of gaseous atoms or molecules from a container to a vacuum through very
small openings
valence bond theory - describes a covalent bond as the overlap of half-filled atomic orbitals
(each containing a single electron) that yield a pair of electrons shared between the two bonded
atoms
orbital overlap – portion of one orbital occupies same space as a second orbital
sigma (σ) bonds - result from the overlap between two s orbitals, 1 s and 1 p orbital, or two p
orbitals
pi (π) bonds - result from the overlap between two side by side p orbitals
atomic orbital theory – describes distribution of electrons in atoms
molecular orbital theory – describes distribution of electrons in molecules
bonding orbitals – result of in-phase combination of orbitals where electron density if between
nuclei and hold atoms together
antibonding orbitals – result of out of phase combination of orbitals where electron density is
opposite of internuclear space
molecular orbital energy diagrams – show relative energy levels of atomic & molecular orbitals
bond order - number of pairs of electrons between two atoms (# bonds between two atoms in a
Lewis structure)
s-p mixing - change that causes σ p orbitals to be less stable than πp orbitals due to the mixing of
s and p-based molecular orbitals of similar energies
formula mass – sum of average atomic masses of all atoms in a chemical formula
aqueous solution – a homogeneous mixture of a substance mixed with water
solute – solution component present in a concentration less than that of the solvent
