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142 Cards in this Set

  • Front
  • Back
energy change from electron transitions
angular momentum quantum number
l, takes any values from 0 to n-1, a given n shell contains all the orbitals from l=0 to l=n-1
magnetic quantum number
describes orbital orientation, each can hold up to 2 electrons, each orbital has 2l+1 orientations, magnetic quantum number takes on values from -l to +l (and zero)
spin quantum number
±1/2, electrons in each orbital with a particular orientation don't repel each other because they have opposite spins, the magnetic fields they create causes a magnetic attraction
Pauli exclusion principle
every electron can be described by a unique set of quantum numbers n, l, ml, ms
maximum electrons in a shell
max = 2n^2
Aufbau principle
electron filling of orbitals based on energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s...
Hund's rule
when filling orbitals an electron is always added to an empty orbital before pairing up and has the same spin as other unpaired electrons
stability of very large orbitals
especially stable when full or half filled (d or f), sometimes s electrons from a higher n value than that larger orbital can donate an electron to give a more stable large orbital
representative elements
elements with either s or p orbitals for valence orbital (IA through VIIIA)
lust, malleable, conduct electricity/heat, lose electrons easily, most bonding is ionic
lower molar masses tend to be gases in elemental state, groups V, VI, VII can gain electrons to form anions, most bonding is covalent
silicon, germanium, arsenic, antimony, tellurium, polonium, astatine; down a column (groups III to VII) increase metallic character (used as semiconductors)
alkali metal group
hydrogen is an exception to the properties, generally soft, low-melting, lustrous metals, react violently with water (form hydroxide salt and H2 gas), more reactive as atomic number increases
alkaline earth metals
meryllium has different properties (nonmetal), generally fairly soft with low density, readily form oxides and hydroxides, oxides and hydroxides are generally insoluble (barium hydroxide is a soluble strong base)
oxygen group
group VIA have their number of possible oxidation states increase with atomic number, oxygen reacts well with most metals but reactivty decrease down the column, selenium/tellurium are semiconductors, sulfur is an electrical insulator
elementally they exist as diatomic molecules
transition metals
lose electrons easily and form ionic compounds, ionic states can vary from +1 to +8, they lose electrons from a higher n value shell first, cations with half or full d orbital are very stable, many have more than one stable oxide (oxide with higher % oxygen forms most easily)
ionization energy trends
decreases down a group, increases from left to right
electron affinity
energy change that occurs when electrons are added to the valence orbital (the greater the Ea, the more stable the anion that formed is), increases across period to VII then drops at the noble gases and then increases again; each group has approximately equal Ea's, Ea is higher if an electron is needed for a half or full d orbital
atom's ability to pull electron density toward itself
effective nuclear charge
Zeff increases going across a row and up a column, the attractive force on the electrons increases and atomic radius decreases
atomic radius
size decreases across period (electron number increases but Zeff does too), Zeff decreases down a group so radius increases
ionic radius
the more electrons an ion has gained, the larger its radius; isoelectronic (ions with same number of electrons) have decreased radius across a row and increased down a column
lattice energy
energy required to break up the lattice into individual ions, energy given off when lattice is created is -U where U=kz1z2/d (d is avg. distance between nuclei, z1 and z2 are ion charges)
sp3d2 geometry
octahedral, 90 degree bond angle
sp3d geometry
trigonal bipyramidal, 120 and 90 degree bond angles
lone pairs and bond angles
as the number of lone pairs increase the bond angle between atoms on the central atom decreases
trigonal pyramidal geometry
three bonds and a lone pair
bent geometry
two bonds and two lone pairs
kinetic molecular theory
1. a gas consists of very small particles that move randomly
2. the volume of each gas particle is negligable compared to the spaces between particles
3. there are no intermolecular attractive forces between the gas particles
4. when gas particles collide with each other or with the walls of the container, there is no net gain or loss of kinetic energy
5. the average kinetic energy of each gas particle is proportional to the temperature
kinetic energy of gas particle
velocity of gas particle
real gas conditions
high pressure and/or small volume can cause intermolecular attractive forces, when pressure increases greatly then repulsive forces become important (makes actual volume slightly above that predicted by ideal gas law)
real gas equation

(1st part is correction for pressure/intermolecular forces; 2nd part is correction for volume/size of gas particles)
typical phase diagram
solid/liquid line has a positive linear slope; liquid/gas line has a concave up curve; solid/gas line has a slighter concave up
vapor pressure equation
(vapor pressure is constant at any given temperature, boiling point is T1)
mole solute/kg solvent
mole fraction
Xsolute=mole solute/mole total
vapor pressure depression of solution
(Pº=vapor pressure of pure solvent)
actual vapor pressure is found by using P=Pº-∆P
boiling and freezing point elevation and depression in solution
∆Tb=Kb x molality
∆Tf=Kf x molality
(can be multiplied by the number of particles per molecule when dissolved)
osmotic pressure
πV=nRT (π=osmotic pressure)
Henry's law for dissolved gases
amount of gas dissolved is directly proportional to the pressure of the gas above the solution (P=kX, X is mole fraction of gas and P is pressure above solution), law works best when gas molecules remain intact in solvent and don't react
suspension of small particles in a medium, distinguished from solids by Tyndall effect where they scatter light
strong acid with weak base reaction
ionic compound that forms is acidic in solution
weak acid with strong base reaction
ionic compound that forms is basic in solution
enthalpy of reaction
enthalpy of formation
∆Hºf=∑BE bonds broken-∑BE bonds formed
Hess's law
the sum of the ∆H's for each individual equation sum to the ∆Hrxn for the net equation (reverse the sign of ∆H if reverse reaction, multiply entire reaction and ∆H by scalar)
half life of 1st order reaction
Arrhenius equation

(k1 is rate constant at T1, k2 is rate constant at T2, greater at higher T)
equilibrium constant factors
temperature affects K, amount of starting material and products affects K
problems involving small K
starting material change is negligable, product amounts are small
Le Chatelier and volume
increase V (decrease P), make more gas molecules; decrease V (increase P), remove gas molecules
effect of pH on basic compound solubiltiy
lower pH increases solubility of basic ionic compounds
six strong acids
HCl, HBr, HI, HNO3, H2SO4, and HClO4
common strong bases
LiOH, NaOH, KOH, Ba(OH)2 (for barium hydroxide the base concentration is half the hydroxide concentration, pH greater)
solution that resists change in pH and has approximately equal concentrations of an acid and its conjugate base
Henderson-Hasselbach equation
enthalpy and internal energy and work
H=U+PV, at atmospheric pressure ∆H=q
second law of thermodynamics
for a sponataneous process ∆S must be > q/T, at equilibrium ∆S=q/T=∆H/T
third law of thermodynamics
a substance that is perfectly crystalline at 0K has an entropy of zero
processes with increasing entropy
more gas produced than used up, solid converted to liquid, solid dissolved in a solvent
processes with decreasing entropy
less gas produced than used up, liquid converted to solid, solid precipitation from a solution
standard free energy
work and ∆G
maximum work for a spontaneous reaction is equivalent to ∆G, the free energy change is the maximum energy available to do work
K and ∆G
not at equilibrium: ∆G=∆Gº+RTlnQ
at equilibrium: ∆Gº=-RTlnK (∆G=0)
relationship between equilibrium constant and temperature
balancing redox
start with main atoms, add H2O to balance oxygen, add protons to balance hydrogen
voltaic/galvanic cells
the redox reactions in two compartments produce an electric current
anode half-cell
oxidation, electrode wears away becoming ion in solution, gives up electrons
cathode half-cell
reduction, electrode builds up with metal ions in solution plating out as pure metal, receives electrons and uses them
cell notation for galvanic cells
anode|anode's ion||cathode's ion|cathode

(single line means phase change, double line means salt bridge)
work produced by a cell
work is negative when work is produced
cell potential uner standard conditions of 25 C, 1M concentrations of solutions, 1 atm pressure of gases
spontaneous reaction listing order
anode reaction is placed above cathode reaction
sign of standard potential and reaction manipulation
only sign changes, the value never changes with differening numbers of electrons
free energy and reduction potential
reduction potential and equilibrium
Nernst equation
calculate nonstandard cell potential when concentrations are other than 1M: Ecell=Eºcell-(RT/nF)lnQ
electrolytic cells
constructed for nonspontaneous redox reactions that require electricity to make them run
calculating current needed to deposit a given mass of metal by electroplating
need Faraday constant, electrons (n), time in seconds, molar mass of metal
radioactive decay equation
collision theory
rxn rate = fZ

(f=fraction of effective collisions with enough KE, z=total collisions)
factors that affect rxn rate
1. reactant concentrations (higher z in r=fZ)
2. higher temp (greater KE, higher f)
3. medium polarity/phase
4. catalyst addition
equilibrium and reaction rates
forward rate equals reverse rate
Keq is equal to the ratio of the rate constants (multiply if mroe than one step)
isolated, closed, open
isothermal, adiabatic, isobaric
state functions
pressure, volume, temperature, enthalpy, entropy, free energy, internal energy
heat of formation
enthalpy change if one mole of a compound is formed directly from its elements in their standard states
STP conditions
0 C (273K), 1 atm
density of ideal gas
d=(MM x P)/RT
molecular speeds
avg. KE=(3/2)kT
avg speed=(3RT/MM)^(1/2)
at higher temp, peak shifts toward higher KE/speed and down at peak
liquid attractive forces
the forces overcome the KE that can keep them apart as in the gas phase
strong attractive forces that only have vibrational motion of atoms
ionic solids
high MP and BP, poor solid state electrical conductivity
metallic solids
metal atoms are packed together with high MP, BP (covalent attractions), layer of spheres shape
crystal unit cells
simple cubic (cube), body-centered cubic (simple plus 1 in middle), face-centered cubic (simple plus one on each face)
gas-liquid equilibrium
gas forms when high KE particles leave liquid and pressure over gas can force some back into liquid, vapor pressure is pressure gas exerts over liquid when condensation-evaporation equilibrium occurs
vapor pressure multiple components
gas pressure exert by A and B over a combined solution are the same as each exerts over its pure solution
Raoult's law
solute B added to solvent A causes vapor pressure of A to decrease where the change is ∆P=Paº-Pa=XbPaº (pure solvent is when Pa=XaPaº), attraction of A and B molecules are the same as with themselves
solubility: always soluble
salts of alkali metals, ammonium salts
halide solubility
soluble except with Ag+, Pb+, Hg2(2+)
sulfate solubility
soluble except Ca2+, Sr2+, Ba2+, Pb2+
metal oxide solubility
insoluble except with alkali metals, ammonium, CaO, SrO, BaO (hydrolyze to soluble hydroxides)
hydroxide solubility
insoluble except with alkali metals, ammonium, Ca2+, Sr2+, Ba2+
carbonates, phosphatess, sulfides, sulfites
insoluble except with alkali metals and ammonium
Arrhenius acid/base
acid produces H+, bases produce OH- (only true in aqueous media)
Bronsted-Lowry acid/base
acid is a proton donor, base in a proton acceptor (works for nonaqueous)
Lewis acid/base
acid is electron pair acceptor, base is electron pair donor
angular momentum of electron
ang. mom.=nh/2π
Lyman, Balmer series
n>1 to n=1, n>2 to n=2
paramagnetic atoms
have unpaired electrons, a magnet will align the spins and weakly attract the atom
have no unpaired electrons, a magnet will slightly repel it (antimagnetic)
complex ions
transition metals form them with water or nonmetals, d orbitals split and d electrons can absorb light
octet exceptions
hydrogen (2), lithium (2), beryllium (4), boron (6), 3rd row (+8)
formal charge
van der waals forces
1. dipole-dipole (solid, liquid, not gas), higher BPs
2. H-bonds, high BPs
3. London dispersion (electron density shifts around randomly, close molecules can interact)
reaction rate (aA+bB->cC+dD)
molar volume of gas at STP
1 atm=22.4mol/L
Dalton's laws of partial pressure
partial pressure of gas equals mole fraction times total pressure, total pressure equals all gas components partial pressure added up
disproportionation reactions
reaction where an element is both oxidized and reduced at the same time
zeroth law of thermodynamics
if two thermodynamic systems are in thermal equilibrium with a third, then they are also in thermal equilibrium with each other
energy equivalence
same units for mechanical (KE/PE), chemical, electrical, and thermal energy
temperature scales
PV work
pressure versus volume curve where work done is area under or enclosed by curve
thermodynamic control
depends on stability of the product, occurs either in vigorous reaction conditions (heat) or over long time when equilibrium is allowed to settle out
kinetic control
occurs best in non-equilibrium situations when present with a thermodynamically stable reaction because its low Ea will be overcome more easily even if its product is less stable
Law of mass action
can explain thermodynamic or kinetic situations (the Keq for thermodynamics and the rxn rate for kinetics)
common ion effect and separation
addition of soluble salt with common ion of slightly soluble salt will push the slightly soluble salt further towards precipitation
complex ion formation
electrophilic metal can bind to electron-donating groups like water and nonmetals, metals become more soluble in aqueous phase as complexes
uses current to separate bonded elements and compounds
electrolytic electrolyte function
because water is not a good enough conductor, an electrolyte must be added to solution
electrolytic cell electron pathway
inert electrodes are used, electrons flow from positive anode to negative cathode, negative cathode precipitates a positive ion when it collides with negative electrode and picks up an extra electron, positive anode takes an electron from an anion to precipitate it and complete the circuit
Faraday's laws of electrolysis
the mass of a substance produced at an electrode during electrolysis is proportional to the number of moles of electrons transferred at that electrode; the number of faradays of electric charge required to discharge one mole of substance at an electrode is equal to the number of excess elementary charges on thta ion
mass of substance deposited or gas liberated at electrolytic electrode
(m=mass, F=constant, Q=current, M=molar mass, n=number of electrons on ion)
electron flow in galvanic cells
anode is where oxidation takes place (solid to ion) and electrons are produced, moved in wire to positive cathode where reduction takes place (ion to solid), salt bridge prevents charge build-up
Henry's law
C=k x Pgas
(C=gas solubility in solvent, k=constant, Pgas=partial gas pressure of that gas of itnerest
Higher pressure of that gas over the liquid phase causes greater solubiltiy
Raoult's law
dissolved solute lowers vapor pressure of a solvent, the partial pressure of each component equals the mole fraction times the pressure of that liquid only
osmotic pressure qualitative
the pressure necessary to stop osmotic flow, depends only on the concentration (not nature) of the solute (the concentration in the osmotic equation is based on the solution concentration)
∆H and bond energy
redox titrations
determination of the concentration of an oxidant or reductant based on titration with opposite redox partner