Gen Chem MCAT

atomic number
number of protons and electrons in an atom (Z)
mass number
number of protons plus neutrons (A)
Avogadro’s number
6.022 x 1023 particles

  • same number of protons and electrons, different number of neutrons
  • same or similar chemical properties (because they have the same valence electrons)
  • different atomic masses (because they have different numbers of neutrons)

Quantum theory
energy comes in discrete bundles
Planck’s constant
h = 6.626 x 10 -34 Js
angular momentum of an electron
Rydberg constant

RH = 2.18 x 10 -18 J/electron


E = -RH/n2

Relate the energy level of an electron with the orbital radius
Smaller radius = lower energy state

Discrete energy bundle


E = hν


E = hc/λ


Emission gives rise to fluorescence

Atomic emission spectra

Unique for each element


E = -RH ((1/ni)2 – (1/(nf)2)

Heisenberg Uncertainty Principle
cannot determine precise momentum and position of an electron simultaneously
Pauli Exclusion Principle
no two electrons can have the same four quantum numbers
principle quantum number
n: describes size
angular quantum number
l describes shape, ranges from 0 to n-1
describes orientation, ranges from -l to l
Hund’s rule
electrons prefer to fill empty orbitals before pairing (within same energy level)
has unpaired electrons, is attracted by magnetic field
have no unpaired electrons; repelled by magnetic field
Valence electrons
outer electrons that are most available for bonding
Atomic radius

  • half the distance between centers of two atoms of an element that are just touching
  • decreases left to right
  • increases going down

Ionization energy

  • energy required to remove a valence electron
  • increases going left to right
  • decreases going down


  • attraction atom has for electrons in a bond
  • increases going left to right
  • decreases going down

electron affinity

energy change that occurs when an electron is added to a gaseous atom


(highest in halogens, zero in noble gases)

Exceptions to the octet rule

  • hydrogen (2)
  • lithium (2)
  • beryllium (4)
  • boron (6)
  • phosphorous, sulfur (expanded)

bond length

decreases as number of bonds increases


distance between two bonded atoms

bond energy

increases as number of bonds increases


energy required to break a bond

formal charge

valence electrons – 1/2bonding – nonbonding


sum of charges is charge on ion


less charges on structure means it is more stable

molecular geometry (2 domains)
AX2 linear
molecular geometry (3)

AX3 trigonal planar

AX2E bent

molecular geometry (4)

AX4 tetrahedral

AX3E trigonal pyramidal

AX2E2 bent

molecular geometry (5)

AX5 trigonal bipyramidal

AX4E see-saw

AX3E2 T-shaped

AX2E3 linear

molecular geometry (6)

AX6 octahedral

AX5E square pyramidal

AX4E2 square planar

AX3E3 T-shaped

AX2E4 linear

sigma bond

single bond

head to head overlap

pi bond


in multiple bonds

parallel overlap

relative strengths of intermolecular forces

ion dipole

hydrogen bonding

dipole dipole

dispersion forces

empirical formula
simplest whole number ratio of elements
molecular formula
exact number of atoms present
combination reactions
2 or more reactants combine to form one product
decomposition reactions
a compound breaks into two or more substances
single displacement reactions
an atom replaces an atom in another compound
double displacement reactions
atoms from two different compounds switch to form two new compounds
limiting reactant
reactant consumed first (least number of moles)
percent yield
actual/theoretical x 100%
rate law

aA + bB -; cC + dD


rate = k [A]x[B]y


Difference between rate constant and equillibrium constant

stoichiometric coeffeicients don’t equal orders of reaction


stoichiometric coeffecients do equal superscript in equilibrium 

steps in determining rate law

1. look for 2 trials where all but one substance concentration is held constant

2. repeat for all reactants

3. plug concentrations in to determine rate constnat

zero order reactions

rate is independent of concentration

k units: M/sec

rate only changes with temperature

first order reactions

most common example is radioactive decay

k units: 1/sec

rate is proportional to concentration of one reactant

Second order reaction
k units: 1/Msec
factors that affect reaction rate

  • reactant concentrations (greater concentrations lead to more collisions)
  • temperature (higher temperature leads to greater kinetic energy, which increases number of collisions)
  • medium
  • catalysts

equilibrium constant

aA + bB -> cC + dD


Kc = [C]c[D]d


What does Keq tell us about products and reactants?

Keq >> 1      products > reactants

Keq << 1      reactants > products

Keq ~ 1        reactants ~ products

LeChatelier’s Principle
determines direction reaction will proceed when subjected to stress
effects of concentration on reaction direction

A + B ; C + D


A increases, shifts to products

D decreases, shifts to products

effect of pressure and volume on reaction direction

Increase in pressure shifts equilibrium to side with fewer moles


Reduction in volume shifts equilibrium towards products

What will shift equilibrium towards products?

reactants added

products taken away

pressure applied

volume reduced

temperature reduced (if heat is a product)

What will shift equilibrium towards reactants?

product added

reactants taken away

pressure reduced

volume increased

temperature increased (if heat is a product)

part of universe being studied
everything outside of the system
isolated system
can’t exchange matter or energy
closed system
can exchange energy but not matter
open system
can exchange both matter and energy
isothermal process
temperature of system remains constant
adiabatic process
no heat exchange occurs
isobaric process
pressure of system remains constant

ΔHrxn = Hproducts – Hreactants


Bond formation is always exothermic (releases heat)

Bond dissociation is always endothermic (requires energy)


ΔS = Sfinal – Sinitial


ΔS = qrev/T


ΔSuniverse = ΔSsystem + ΔSsurroundings


ΔSuniverse > 0 in spontaneous reactions

Gibb’s free energy

;G = ;H – T;S


;G ; negative ; spontaneous

;G ; positive ; nonspontaneous

;G ; zero ; equilibrium

Standard Gibb’s free energy

ΔGº = -RT ln Keq


ΔG = ΔGº + RT ln Q



standard temperature pressure

T = 0ºC

Standard state

T = 25ºC

 used in standard enthalpy/entropy problems

Boyle’s law

pressure and volume are inversely related

P1V1 = P2V2

Charles’ Law

volume and temperature are directly proportional


V1/T1 = V2/T2

Ideal Gas Law

PV = nRT


d = m/V = P(MW)/RT

Partial pressures

Ptot = PA + PB + PC



XA = nA/nT

When do real gases deviate from the ideal gas law?

at high pressure, low temperature, and temperatures close to the boiling point


V will be less than predicted

Assumptions of the Kinetic Molecular Theory

  1. particle volume is negligible when compared to container volume
  2. gases have no intermolecular forces
  3. gases particles are in continuous, random motion
  4. collisions are elastic, so there is no overall gain or loss of energy
  5. average kinetic energy of gas is proportional to the absolute temperature

diffusion of gases

heavier gases diffuse more slowly than lighter ones


r1/r2 = √((MW2)/(MM1))

transition between liquid and gas

evaporation: liquid to gas

condensation: gas to liquid

boiling point

vapor pressure of liquid is the same as the external pressure


transition between solid and liquid

melting (fusion): solid to liquid

solidification (crystallization): liquid to solid

solid to gas direct transitions

sublimation: solid to gas

deposition: gas to solid

osmotic pressure

∏ = MRT


water will move towards greater molarity or higher temperature

Solubility: salts of alkali metals
always soluble
Solubility: salts of ammonium ion
always soluble
solubility: chlorides, bromides and iodides
soluble unless Ca, Sr, Ba,  Pb
Solubility: metal oxides
insoluble except CaO, SrO, BaO
Solubility: hydroxides
insoluble except alkali metals an Ca, Sr, Ba
Solubility: carbonates, phosphates, sulfides, and sulfites
insoluble except alkali metals and ammonium ion
solutes whose solutions are conductive
percent composition by mass
mass solute/mass solution x 100%
mole fraction
moles compund / total number moles
molarity (M)
mol solute/L solution
molality (m)
mol solute/kg solution
normality (N)
g solute/L solution
solubility constant
Ksp = [An+]m[Bm-]n
solubility constant vs. reaction quotient

Ksp > Q   solute will continue to dissolve

Q > Ksp   precipitation will occur

Q = Ksp   equilibrium

common ion effect
if a salt is added to a solution already containing one of the ions, the equilibrium will shift to favor the solid salt
conjugate acid-base pairs
related by the transfer of a proton
acid: __________ide
hydro_____ic acid
acid: _____ite
_____ous acid
acid: _____ate
_____ic acid



= 10-14


pH + pOH = 14

dissociation of strong acids and bases
completely dissociate into component ions

Ka = [H3O+][A]/[HA]


measures degree to which acid dissociates

stength of acid compared to Ka
weaker acids have smaller Ka‘s
amphoteric (amphiprotic)
acts as both an acid and a base
titration: strong acid + strong base
quivalence point is at 7
titration: weak acid + strong base
equivalence point is in the basic range
mixture of a weak acid or base with its salt
Hendersen-Hasselbach equation
pH = pka + log[A]/[HA]
polyprotic acid titrations

have more than one equivalence point


(each equivalence point corresponds to the loss/gain of one electron)

oxidizing agent
causes atom to undergo oxidation
reducing agent
causes atom to be reduced as the agent itself is oxidized
LEO the lion says GER

lose electrons = oxidized

gain electrons = reduced

oxidation number of free elements
oxidation number of monatomic ions
equal to the charge state
oxidation state of group IA and IIA elements
+1 and +2 respectively
oxidation state of halides
-1 unless attached to a more electronegative atom
oxidation state of hydrogen
+1 unless attached to a less electronegative atom
oxidation state of oxygen
usually -2
What is the sum of the oxidation numbers equal to?
0 in a neutral compound
Steps in balancing a redox reaction

  1. separate two half reactions
  2. balance all atoms except H and O
  3. add water to balance O
  4. add H+ to balance H
  5. add electrons to balance charge
  6. multiply each half reaction so that number of electrons gained/lost is equal
  7. add half reactions to cancel electrons

Galvanic vs. electrolytic cells (in terms of Gibb’s free energy)

Galvanic cells have spontaneous reactions so -ΔG

Electrolytic cells have nonspontaneous reactions so +ΔG

where reaction occurs
where oxidation occurs (AN OX)
where reduction occurs (RED CAT)
salt bridge
allows for exchange of cations and anions
conventional representation of a cell
anode | anode solution || cathode solution | cathode
anode charge in galvanic and electrolytic cells

positive in electrolytic cells

negative in galvanic cells

reduction potential

tendency of a species to aquire electrons and be reduced


a more positive Eº means greater tendency for reduction to occur

electromotive force (EMF)

difference in potential between two cells


positive in galvanic cells


negative in electrolytic cells

Nernst equation
nFE°cell = RTlnKeq
what information does a positive E°cell give?
K is postive so product formation is favored

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