Chemistry (MCAT)

atomic notation

atomic number (Z)- number of protons, determines element

mass number (A)- number of protons and neutrons, determines isotope

charge (C)- number of protons minus electrons

Bohr model

electrons orbit at fixed distances, different between orbits decreases the larger they get

principal quantum number (n) describes the energy level

excitation- electrons are allowed to absorb specific photon energies to jump energy levels, each atom has specific absorption spectrum

emission/relaxation- electrons in excited state can return to lower energy state by emitting a photon, each atom has specific emission spectrum

emission energy is less than excitation energy because there is loss of vibrational energy in electron

Rutherford experiment

gold foil experiment with alpha particles

most of volume of atoms is empty space

nucleus must be positively charged to deflect alpha particle

electromagnetic spectrum

energy = planks*frequency, frequency = c/wavelength

increasing wavelength:

1. gamma rays
2. X-rays
3. UV
4. violet (300 nm)
5. red (700 nm)
6. IR
7. microwaves
8. radiowaves

gay X-men using viagra in my room

quantum numbers

principal quantum number (n)- energy level, n = 1, 2, 3, etc.
azimuthal, orbital angular momentum (l)- shape of orbital, l = 0 is s, l = 1 is p, etc.
magnetic (ml)- orientation of orbital, ml ranges from -l to +l
electron spin (ms)- direction of electron spin, ms = 1/2, -1/2

rules for filling orbitals

Pauli exclusion principle- no two electrons in the same atom can have the same quantum numbers, same spin

aufbau principle- electrons fill in lowest energy orbitals first, two exceptions where s electrons are promoted: copper (and silver, gold) (d9 becomes s1d10) and chromium (and molybdenum) (d4 becomes s1d5)

hund's rule- electrons fill each orbital in a subshell before doubling up

when removing electrons, 4s electrons are removed before 3d to create ions

Fe = 4s2 3d6
Fe2+ = 3d6

isoelectronic

two atoms have same number of electrons

magnetism

paramagnetic- at least one unpaired electron, attracted to an external magnetic field

diamagnetic- all electrons are paired, repelled from an external magnetic field

shielding

valence electrons have lower effective nuclear charge because they are further away from nucleus and are repelled by core electrons

more protons means greater effective nuclear charge

more orbital rings means lower effective nuclear charge

more electrons (greater negative charge) means lower effective nuclear charge

electrostatic attraction

increases from lower left to upper right

1. down a group, valence electrons are farther away means weaker pull
2. across a period, more protons means stronger pull
3. becoming more negatively charged, more shielding means weaker pull

atomic radius

increases from upper right to lower left

1. down a group, more orbital rings means larger atom
2. across a period, more protons means smaller atom

element groups

alkali metals (1), alkaline earth metals (2), halogens (7), nobel gases (8), transition metals (d block), rare earth metals (f block)

alkali metals and alkaline earth metals often get oxidized to become metal cations

H+ often gets reduced to become H2, and leaves solution
halogen anions often get oxidized to become diatomic

ionic radius

increases with increasing negative charge

anions (-) are larger than neutral, cations (+) are smaller

free radicals

atoms with unpaired valence electrons, can be charged or uncharged, damage DNA to cause cancer

reactive oxygen species (ROS) are free radicals

H2O2 is an ROS

antioxidant enzymes remove ROS

ionization energy

increases from lower left to upper right

1. down a group, more rings means valence electrons are less tightly held
2. across a period, more protons means valence electrons are more tightly held

for neutral atoms, second ionization energy always greater than first

closed-shell requires more energy to remove electron

ionization energy of boron slightly lower than beryllium because electron is in alone in p-orbital

ionization energy of oxygen slightly lower than nitrogen because electron has to occupy a p orbital with an electron already in it, calculate energy difference by comparing oxygen if hypothetically spin pairing didn't occur with when it does occur

electron affinity

increases from lower left to upper right (value becomes more negative, but magnitude increases)

1. down a group, more rings means less energy released
2. across a period, more protons means more energy released

defined as energy change when adding an electron to valence shell in gas state

high electron affinity means high reduction potential

electronegativity

increases from bottom right to upper left

1. down a group, more rings means less pull
2. across a period, more protons means more pull

metals lose electrons in presence of non-metals

F > O > N > Cl > Br > I > S > C = H

acidity

increases from upper left to lower right (different!)

1. down a group, size explains: larger atom means conjugate base (cation) is more stable
2. across a period, electronegativity explains: more protons means conjugate base (cation) is more stable

Lewis structures and formal charge

1. place least electronegative atom is center, always C
2. start with all single bonds, deduct 2 electrons for each
3. start with most electronegative atoms and make octets
4. any remaining electrons add to central atom
5. form bonds to fill any atoms without octets
6. assign formal charge = valence electrons - sticks - dots

positive formal charge should be on less electronegative atoms
try to reduce set of formal charges
check that formal charges add up to overall charge

hybridization

sp for 2 groups, sp2 for 3 groups, sp3 for 4 groups

lone pairs determine molecular geometry: NH3 is trigonal pyramidal (sp3), H2O is bent (sp3), XeF4 is square planar (sp3d2), SF6 is octahedral (sp3d2)

octahedral is 6 bonds, 90 degrees (sp3d2)

sp3 to sp2 conversion to attain planarity and aromaticity (4n+2 electrons resonating in a ring)

amide N can convert from sp3 to sp2 to get resonance stabilized by carbonyl O

orbitals with more s character are more stable

bond length

depends on atomic radius (which increases to lower left)

depends on bond order, more electrons shared is a stronger bond and closer bond

breaking bonds

breaking a bond is endothermic (positive dH) but also exergonic! (negative dG)

ATP ->ADP + H2O requires heat but produces net energy

stronger bonds:

1. have more electrons shared
2. shorter distance between atoms
3. have higher dissociation energy

ATP phosphates

alpha- phosphate closest to sugar
beta- middle one
gamma- phosphate on tip

covalent bonds

metal and nonmetal share electrons

electronegativity differences creates dipoles at each bond

molecular dipole found by adding up all the bond dipoles

metallic bonds

sea of electrons delocalized between metal ions

usually S and D block elements

these compounds are conductors and malleable

coordinate covalent bond

formed between atoms with lone pairs and atoms that are electron deficient

usually between transition metals and organic compounds, where compounds donate electron pairs to coordinate with the metal

coordinate number- number of compounds coordinating with metal ions

Fe2+ can form 6 coordinate covalent bonds (that completely fill up its valence shells)

Fe2+ forms coordinate covalent bonds with hemoglobin

ionic bonds

formed between cations and anions

ions dissociate in aqueous solution, become conductor

as solids, these compounds are insulators and brittle

insulator/conductor

insulator- valence electrons tightly bound to atom

conductor- delocalized electrons, metallic bonds

intermolecular forces

pulling apart atoms is always endothermic!

4 types of IMFs:

1. ion-dipole forces- ions and polar molecule
2. dipole-dipole forces- two polar molecules, align along the molecular dipoles, H-bonding is special case
3. dipole-induced dipole forces- polar, nonpolar molecules
4. london dispersion forces- Van der Waals, temporary
5. Hydrogen bonds- align along the bond dipoles, require a donor and an acceptor, only N, O, and F can do hydrogen bonds

solvation shell

A cagelike network of solvent molecules that forms around a solute in a solution

decrease in entropy

active site

acidic and basic amino acids can undergo H-bonding and ionic interactions with the substrate

ATP has negatively charged phosphates that interact well with H, R, and K

entropy

disorder, always increasing in universe

what increases entropy (S):

1. increasing number of particles
2. increasing volume
3. increasing temperature

formation of a more organized compound or state would decrease entropy

enthalpy

breaking bonds is endothermic (dH > 0)
forming bonds is exothermic (dH < 0)

heat of formation (dHf) of elements in standard state is 0

heat of reaction (dH) = Hf products - Hf reactants
multiply by number of moles

positive dH means endothermic, heat is reactant
negative dH means exothermic, heat is product

Gibbs free energy and spontaneous reactions

dG = dH - TdS
dGo = - RTlnK
dG = dGo + RTlnQ

free energy of formation (dGf) of standard state elements is 0

spontaneous process is exergonic (dG < 0) nonspontaneous process is endergonic (dG > 0)

examples:
combustion (-dH, +dS) is spontaneous at all temperatures
freezing (-dH, -dS) is spontaneous at low temperatures
ATP -> ADP (+dH, +dS) is spontaneous at high temperatures

bond and IMF strengths

order of bond strengths:

1. covalent
2. ionic
3. metalic
4. coordinate covalent

order of IMF strengths:

1. ion-dipole
2. dipole-dipole (H-bonds)
3. dipole-induced dipole
4. london dispersion forces

phase diagram

as pressure and temperature increase, phase change from solid to liquid to gas

triple point- all three phases coexist

critical point- above point liquid and gas are no longer distinct, becomes supercritical fluid

sublimation and evaporation line up on the diagonal

melting line is positive sloping, but for water it is negative sloping (MP decreases when pressure increases, since ice has more volume than water)

density is directly proportional to external P, indirectly proportional to external T

calorimetry

measure changes to determine heat transfer

constant pressure calorimeter- coffee cups
bomb calorimeter- volume constant, pressure changes

heating curve

heat for phase change
q = ndH
n = moles
dH = molar enthalpy of phase change

as IMFs break, potential energy of substance increases but no temperature changes occurs

heat for temp increase
q = mcdT
c = specific heat (J/kgK)
C = mc = heat capacity (J/K)

vapor pressure explains boiling point

force exerted by gas particles

Patm has no effect, external T is directly proportional

boiling point is when Pvap reaches Patm

boiling point is directly proportional to:

1. IMFs- the weaker the interactions, the higher the Pvap, the higher the boiling point
2. Patm- the lower the atmospheric pressure, the easier it is for Pvap to reach Patm

fp depression/bp elevation

3 factors cause this widening of fp and bp:

1. impurities (during mp determination)
2. ions (colligative property in solvent)
3. cholesterol (in lipid membrane)

surface tension

factors that decrease surface tension:

1. increased temperature
2. increased surface area

barometer

air pressure inside tube is less than atmospheric pressure, so liquid level rises

increased liquid pressure exerts force downward to compensate for decreased air pressure, until equilibrium

air pressure inside does not change!

solution and electrolytes

homogenous mixture, solute in solvent (usually aqueous)

strong electrolyte- complete dissociation, strong acid/base
weak electrolyte- partial dissociation- weak acid/base
non-electrolyte- no dissociation, covalent bonds

van't Hoff factor (i)- number of particles per mol

solubility rules

solid/liquid solubility in water:

1. directly proportional to temperature
2. not affected by pressure

gas solubility in water:

1. indirectly proportional to temperature
2. directly proportional to pressure

always soluble:

1. alkali metal ions, H+, NH4+
2. NO3-, CH3COO-, ClO4-

usually insoluble:

1. Ag+, Pb2+, Pb4+, Hg2+
2. CO3(2-), PO4(3-), S2-

kinetic-molecular theory of gases

an ideal gas has:

1. no IMFs
2. particles have no volume
3. temperature is average KE
4. elastic collisions with container

high temperature, low pressure creates an ideal gas since interactions are minimized

basic gas laws

Avogadro's law- n proportional to V, so 1 mol is 22.4 L at STP

Boyle's law- P inverse to V, PV = PV

Charles' law- T proportional to V, V/T = V/T

Gay-Lussac's law- T proportional to P, P/T = P/T

Dalton's law- partial pressure proportional to mole fraction, Ptotal = P1 + P2

PdV curve

x-axis is volume
y-axis is pressure

work = PdV = area under the curve

filling up more volume under constant pressure does more work

units of pressure

1 atm = 100 kPa = 760 torr = 760 mmHg

ideal gas law

PV = nRT
R = .08 Latm/Kmol

but you don't need this, just use 22.4 L/mol at 0 C and 1 atm

real gas law and non-ideality

Ptheoretical > Pobs, since particles have IMFs with each other
Vtheoretical < Vobs, since particles have size

(Pobs + an^2/V^2)(Vobs - nb) = nRT

1. a increases as IMF increase
2. b increases as particle size increases

non-ideal conditions are high pressures and low temperatures, which maximize IMFs:

1. V_obs is greater than predicted
2. P_obs is smaller than predicted

Graham's law of diffusion/effusion

rate1/rate2 = sqrt(molar mass2/molar mass1)

heavier particles diffuse slower

reaction coordinate graph

x-axis is reaction progression
y-axis is free energy

peaks are transition states
valleys are intermediates

the highest peak is the rate limiting step
difference between reactant energy and highest peak is Ea

catalysts lower peaks, does not affect equilibrium

reaction rate

factors that increase it:

1. higher concentration of reactants
2. proper orientation, catalysts help
3. minimum energy to overcome Ea

rate constant

factors that increase it:

1. increased temperature
2. decreased activation energy
3. not affected by concentration of reactants!

rate law

rate = k[A]^x[B]^y

order is sum of exponents

determine rate law:

1. coefficients of reaction become the exponents of rate law
2. solids are not included
3. rate law determined by rate limiting step

determine rate law from data:

1. determine rate constant from any trial
2. 0th order if rate is same between two trials with diff. concentrations of a reactant
3. 1st order is rate increase is same as concentration increase of reactant
4. everything else must be 2nd order

catalytic efficiency

vmax = k_cat*[total enzyme]

catalytic efficiency = k_cat/Km

equilibrium constants

K is conc. of reactants over products, K > 1 favors products

K only changes when temperature changes

Keq > 1 means dG < 0, so spontaneous Keq < 1 means dG > 0, so non-spontaneous

Q is reaction quotient using present concentrations state:

1. Q > K favors shift to reactants
2. Q = K is in equilibrium
3. Q < K favors shift products

why? because of the equation below:

1. dG = dGo + RTlnQ
2. dGo = -RTlnK
3. dG = 0 at equilibrium

factors that affect equilibrium

Le Chatelier's principle says system in equilibrium shifts to minimize stress:

1. change in concentration/partial pressure
2. change in pressure/volume
3. change in temperature

change in concentration or partial pressure- causes shift to reduce the change, solids do not affect equilibrium!

increase in pressure- causes shift towards side with less moles of gas

decrease in volume- causes shift towards side with less moles of gas

change in temperature- treat heat as a reactant/product

multiple equilibria

when one reaction in a system shifts, all reactions shift in the same direction

H2O + CO2

combining two equations requires multiplying the K values

reversing an equation requires inversing the K value

solubility equilibrium

Ksp- concentrations of products since the reactant is solid

molar solubility- moles of solid dissolved, you can plug this into Ksp equation

Qsp > Ksp means that oversaturated, precipitate forms (shift towards reactants)

common ion effect- salt solubility decreases when mixed in solution with a common ion (shift towards reactants)

opposite effect is true, salt solubility increases when mixed with solution that reacts with an ion, like an acid/base (shift towards products)

association/dissociation equilibrium

Ka = [AB]/[A][B]
Kd = [A][B]/[AB]

kinetics vs. thermodynamics

kinetics- rate, intermediate, activation energy, catalyst

faster product

thermodynamics- stability, equilibrium, spontaneity, entropy, enthalpy, free energy

more stable product

for example, free energy of ion transport does not tell you about the kinetics of a channel protein

acid/base trends

acids have more electronegative atoms
atoms without H can be acids if electron deficient

increased acidity:

1. more positive charge
2. more electronegative atom (only within a period)
3. larger atom (only within a group)

bases have less electronegative atoms
atoms without lone pairs are not basic

increased basicity:

1. more negative charge
2. less electronegative atom
3. smaller atom

acidity/basicity of salts

salts will completely dissociate in water!

acidic salt contains ion that is a weak acid
cations in group I and II are not acidic (spectators!)

basic salt contains weak base
anions like Cl, Br, or I are not basic (spectators!)

determine pH of solution

1. The acidity of an element increases as one moves to the right or down the periodic table.

2. The conjugate base of a strong acid is a very weak base, forming pH neutral solutions.

3. The conjugate base of a weak acid is a kinda weak base, pH will be higher.

amphoteric

can be both acid or base

examples:

1. amino acids
2. water
3. bicarbonate

strong acids (6) and calculating pH

HCl, HBr, HI, H2SO4, HClO4 (perchloric acid), HNO3

how to calculate pH:

1. write the concentration in scientific notation
2. take the negative of the 10 exponent
3. round down (because you ignored the actual value) and say it's somewhere in between those two values

for diprotic acids, just double the concentration, you can get a negative pH

large Ka, weak conjugate base has small Kb

weak acids and calculating pH

acetic acid, H3PO4 (phosphoric acid), H2CO3 (carbonic acid), NH4, HF

Ka = [H+]^2/[HA] from ICE table

pH = -1/2log(Ka[HA])

use same estimation technique

small Ka, strong conjugate base has large Kb

pH in solution is a bit under 7

weak acids can fully dissociate just like strong acids if another reaction removes protons to shift equilibrium to the right (like bicarbonate buffer)

strong bases (6) and calculating pH

O2-, OH-, OR-, H-, R-, NH2-

how to calculate pOH:

1. write the concentration in scientific notation
2. take the negative of the 10 exponent
3. round down (because you ignored the actual value) and say it's somewhere in between those two values

for diprotic bases, just double the concentration, can get a negative pOH

large Kb, weak conjugate acid has

weak bases and calculating pH

NH3

Kb = [OH-]^2/[B] from ICE table
pOH = -0.5log(Kb[B])

small Kb, strong conjugate acid was large Ka

pH in solution is a bit over 7

differences in pH

pH 4 has 100 times H+ concentration as pH 6

buffers

conjugate acid/base pair, minimize pH change

weak acid or base (acetic acid, phosphoric acid, carbonic acid are weak acids, NH3 is weak base etc.)

both acid and base must be available to push equilibrium to account for changes in conc.

pH = pKa + log([A-]/[HA])
HH equation shows that conc. of acid/base dependent on pH

if pH < pKa, then buffer is more protonated, each difference of 1 in pH is 10-fold increase in concentration

adding water dilutes your buffer, reduces buffering capacity, but does not affect equilibrium

choose a buffer within 1 of the pH you want to maintain!

henderson-hasselbalch equation

pH = pKa + log [base]/[acid]

conc. of acid/base are dependent on pH
calculate pH of buffer solution given conc.

when [base] = [acid], pH = pKa at half-equivalence point

auto-ionization of water, pH/pOH

at standard conditions:
Kw = KaKb = 1e-14

pKa = -logKa, pKb = -logKb

pKa + pKb = 14 (pH + pOH = 14)

pH = 0 is most acidic, pOH = 0 is most basic, neutral is pH = 7

titration

sigmoidal curve, indicator changes color at equivalence point

used to determine concentration of solution, where moles added equals moles in solution to reach each equivalence point

at equivalence point:

1. compound becomes deprotonated
2. moles of acid = moles of base, meaning everything is neutralized
3. center of vertical segments

at half-equivalence point:

1. pKa = pH
2. [acid] = [base], moles of acid = 1/2 moles of base, meaning half of the acid is neutralized
3. center of horizontal segments

polyprotic titration can be done with amino acids, double the moles added to reach the second equivalence point

for titrating a strong acid, the same volume of strong or weak bases is required to neutralize, the strength will only affect the pH of equivalence point

indicators

choose indicator with pKa within 1 of the pH you want to measure!

equivalence point is the pKa of the indicator

redox reactions

OIL RIG

reduction is gain of electrons, decreased charge

reduction potentials are for the reactants of reduction rxn

easily reduced: neutral nonmetals, anything ending with O, halogens

oxidation is loss of electrons, increased charge

oxidation potentials are for the reactants of oxidation rxn

easily oxidized: H2, neutral metals, anything ending with H

redox reactions can be coupled so that a non-spontaneous reaction can occur, check the reduction potentials!

the ETC is a series of reductions, so each step has a higher standard reduction potential, with O2 -> H2O having the highest

oxidation states

cell potential

Ecell = Ered + Eox

free gibbs energy is inversely proportional with E:

1. dG = -nFEcell
2. Ecell > 0, cell is spontaneous (galvanic)
3. Ecell < 0, cell is nonspontaneous (electrolytic)

Faraday's constant

charge per 1 mol of electrons
F = 100,000 C/mol

Faraday's Law and electroplating

over time, electrons can be used to do work like electroplating

I = Q/t
Q = nF
Faraday's constant is charge per mole

calculate moles of metal plated given current and time:

1. determine charge: Q = It
2. determine moles of electrons: ne = Q/F
3. determine moles of metal: nmetal = ne/ionic charge (electrons per metal)

galvanic cell

cell must have two electrodes, an electrolyte bridge, and a wire with resistance (galvanic) or power source (electrolytic)

spontaneous, positive Ecell

oxidation at the anode (an ox), reduction at the cathode (red cat)

electron travel from anode (-) to cathode (+) through wire, discharging battery

cathode is plated, anode loses metal ions

Nerst equation

cell in equilibrium has an actual cell potential of 0 and a standard cell potential of not 0

E = E0 - (RT/nF)lnQ
E0 = (RT/nF)lnK

-as temp increases, Ecell decreases

-as battery approaches equilibrium, Ecell approaches 0

-E0 is constant

salt bridge

anion will always migrate towards anode, cation will always migrate toward cathode

in galvanic cell, salt bridge balances out movement of electrons from anode to cathode

in electrolytic cell, anion just is attracted to positively-charged anode

without a salt bridge, current immediately stops

electrolytic cell

opposite of galvanic cell

recharging battery using external energy source reverses flow of electrons

electrons still going from anode to cathode, but are being forced to cathode is negative and anode is positive

increasing current will increase rate of electroplating

drives nonspontaneous reaction, negative Ecell

electrolysis

in electrolytic cell with water, hydrolysis will occur creating H2 at the cathode and O2 at the anode

in electrolytic cell with NaCl, hydrolysis will occur creating H2 at cathode and Cl2 at anode

redox titration

cerimetry- Ce4+ + e- => Ce3+

adding Ce4+, observing increase in pH on y-axis

alpha decay

4/2He2+ emitted from nucleus, subtract 2 protons and 2 neutrons

larger nuclei more likely to undergo, less dangerous

beta decay

beta decay- electron emitted from nucleus, neutron becomes a proton and electron

positron emission- beta positive decay, positron (positive charged electron) is emitted, proton becomes a neutron and positron

electron capture- another type of beta decay, electron absorbed by nucleus, electron and proton become a neutron

high proton to neutron ratio makes this more likely

gamma decay

photons are emitted

most dangerous type of radiation

half-life

time it takes for a substance to decay to half of its original amount

a radioactive isotope will have a half-life as it decays to a different daughter isotope (which may or may not be radioactive)

atomic mass would be negligibly affected, while atomic number would be affected

nuclear reactions

are always exothermic!

decay, fission, fusion

mass defect- protons and neutrons separately have more mass than when together in nucleus, binding energy between nucleons

binding releases energy, breaking bonds takes energy