High Yield MCAT Chem - Johan

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

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

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

P_theoretical > P_obs, since particles have IMFs with each other

V_theoretical < V_obs, since particles have size

(P_obs + an^2/V^2)(V_obs - 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

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

reaction rate

factors that increase it:
1. increased temperature
2. decreased activation energy
3. not affected by concentration of reactants!

rate law

an equation that relates the reaction rate to the concentration of reactants, typically in the form: rate = k[A]^m[B]^n.

solubility equilibrium

the state in which the rate of dissolution of a solute is equal to the rate of crystallization, resulting in a constant concentration of the solute in solution.