chem unit 3

interference

interaction between waves

constructive interference

waves are in phase and combine

destructive interference

waves are out of phase and cancel

diffraction

bending and spreading of waves as they hit an object 

photoelectric effect

observation that metals emit electrons when light is shined upon them

n

principle quantum number, determines energy level of electron

indeterminacy

present circumstances do not necessarily determine future events

l

angular momentum quantum number, determines shape of electron

ml

magnetic quantum number, determines orientation of electron

orbitals

probability distribution maps showing where electrons are likely to be found

ms

spin quantum number, determines spin of electron

phase

the sign of a waves amplitude, positive or negative

pauli exclusion principle

no two identical electrons can have the same four quantum numbers, implies there can only be two electrons per orbit

aufbau principle

electrons fill lower energy orbits before moving to higher energy levels

hunds principle

electrons will fill orbits individually before pairing up

madelung principle

explains the order in which atomic orbits are filled

radial distribution plot 

graphical representation of probability of finding electrons 

degenerate orbitals

orbitals with the same energy within the same subshell

ionization energy

energy required to remove an electron in the gaseous state

electron affinities

energy change associated from gaining an electron in the gaseous state

van der waals radius

radius between non bonded atoms

covalent radius 

radius between bonded atoms 

paramagnetic

has unpaired electrons, is attracted to a magnetic field

diamagnetic

does not have unpaired electrons, is slightly repelled by a magnetic field

lattice energy

energy associated with the formation of a crystalline lattice

born-haber cycle

type of energy cycle used to calculate the lattice enthalpy of an ionic compound

hess’s law

∆E = sum of enthalpy changes of the steps

intermolecular forces 

force in between molecules 

intramolecular forces

force within molecules

electronegativity

ability of an atom to attract electrons to itself

dipole moment

vector quantity indicating polarity in a molecule

resonance structures 

2 or more valid lewis structures that together describe the bonding in a molecule 

formal charge

(valence e) - (nonbonding e + ½ bonding e)

free radicals

molecules with an odd number of electrons

coordinate covalent bond

both electrons in the shared pair come from one atom

bond energy 

energy required to break 1 mol of the bonds in gas phase (bond enthalpy)

exothermic

release energy

endothermic

require energy

valence shell electron pair repulsion theory (VSEPR)

model that helps us examine the shape of molecules 

electron geometry

arrangement of electron groups

molecular geometry

arrangement of the atoms

valence bond theory

describes covalent bonds as the overlap of atomic orbitals

hybrid orbitals

combined atomic orbitals that create new, equivalent hybrid orbitals

perturbation theory

a complex system is viewed as a simpler system that is slightly altered by some additional force

bond energy 

the measure of the strength of a chemical bond

pi bond

when orbitals overlap side by side

sigma bond

when orbitals overlap end to end

isomers

same molecular formula but different structures

bonding scheme 

describes how atoms are connected in a molecule 

molecular orbital theory

describes the electronic structure of molecules by combining orbitals to form new orbitals that are delocalized over multiple atoms

bonding orbital

in molecular orbit theory, the bond that is lower in energy than any atomic orbital from which it was formed

antibonding orbital

in molecular orbital theory, orbitals with higher energy that any atomic orbital from which it was formed

bond order 

(bonding electrons - antibonding) / 2 

homonuclear diatomic molecules

molecules made up of 2 atoms of the same kind

nonbonding orbitals

in molecular orbital theory, electrons that remain localized on an atom

intermolecular forces

forces between molecules

intramolecular forces 

forces within a molecule 

dispersion force

present in all atoms, weak temporary attraction between molecules due to instantaneous dipoles

dipole dipole force

exists only in polar molecules, attractive forces between the positive end of one polar molecule and the negative end of another

hydrogen bond

a type of super dipole dipole bond, occurs when hydrogen is bonded to an electronegative ion (O, N, F)

to polarize

to form a dipole moment

miscibility

the ability to mix without separating, determined by polarity

ion dipole force

force between an ionic compound and a polar compound

volatility

tendency to evaporate at a given temperature 

steric number

the electron groups surrounding an atom