MCAT Chemistry

nucleus

where are protons and neutrons located?

where are e- located?

e- cloud

atomic mass

• isotopic mass

• specific mass of certain isotopes

atomic weight

• avg of all masses of naturally occurring isotopes

• abundance x atomic mass

atomic number

number of protons

mass number

• sum of amt of protons + neutrons in an atom

• isotopes always have the same atomic #

principle shell

• where e- are

• the one closest to nucleus is most stable = ground state

electron shells

• the farther the e- is from nucleus, the greater the E

• each e- in same shell has same E

• # e- in shell = 2n²

spectra

range of wavelengths/frequencies emitted/absorbed

excitation

e- absorbs E + moves to a level that can handle the E

de-excitation

• e- releases E as heat/light

• color it can reflect depends on element

• emission of photon/E = wavelength

energy level

where e- might be, done by determining principal quantum #

principle quantum number

• describes shell of atom

• high n = high distance from nucleus

subshell

• area w/in shell that describes shapes of spaces where e- might be

• 3D

• each E level has subshells

azimuthal quantum numbers (I)

• subshell of orbital (shape)

• values only range from 0 to (n-1)

magnetic quantum number

• s: 1s orbital/orientation

• p: 3p orbitals/orientations

• d: 5d orbitals/orientations

• f: 7f “

electron spin number

• orientation which e- spinning

• if final e- arrow points up or down

• ½ = up -½ = down

Aufbau principle

1. e- fill lowest level first

2. order: 1s, 2s, 2p, 3s, 3p, 4s, 3d

3. each e- will hold 2e- w opposite spins

4. e- will fill up orbitals of same E first

Pauli Exclusion principle

no e- can share same quantum #s

Hund’s rule

when assigning e- it must be individually placed b4 pairing up

ions

elemental atoms that have diff # of e-

isotope

diff # of neutrons

photons

main units of light that has dual particle-wave characteristic

metals

• malleable, ↑conductivity bc of willingness to give up e-

• have ↓ionization E + ↓e- affinity

• can attain multiple oxidations states

  • good reducing agents

nonmetals

• ↑ionization E, ↑e- affinity, ↑electronegativity

• non-malleable, brittle, soft solids

• many are gases @ room temp

metalloids

btwn metals + nonmetals (staircase ones)

semiconductors

ionization energy

E needed to remove e- from a natural gaseous atom

electron affinity

E emitted/absorbed when e'- is added to atom

oxidation states

shows # of e- lost/gained to form chem bond

electronegativity

attraction of other molecules to each other

semiconductor

molecules that can conduct + insulate E

valence electrons

e- in outermost shell + participate in chem bonding

effective nuclear charge

• inward/pulling force of pos nuclear charge on valence e-

• increases going left to right on periodic table

• Z_eff = Z — S

  • S: all e- except valence

• 1st E level will have greatest nuclear charge

• Z_eff decreases while principle quantum # increases

  • bc inner core e- shield some (+) charge

periodic table trends

alkaline earth metals

• group 2

• ↓e- affinity/negativity + ionization E

  • bc only have 1-2 e- in valence

• prone to form cations bc they lose e- easily

• moving down, reactivity + atomic density ↑

• melting + boiling pnt ↓

Be, Mg, Ca, Sr, Ba, Ra

alkali metals

• group 1

• soft, ↓density + melting pnts

• melting + boiling pnts ↓ going down

• reactivity ↑moving down

Li, Na, K, Rb, Cs, Fr

chalcogens

• group 16

• nonmetals, metalloid + metal

• metallic character ↑ as you move down

• 6 valence e-

O, S, Se, Te, Po, Lv

halogens

• groups 17

• 7 valence e- + readily accept/share an e-

• form ionic bonds w alkali + covalent w other nonmetals

• going down grp:

  • melting/boiling pnt ↓

  • reactivity ↑

F, Cl, Br, I, At, Ts

noble gases

• group 18

• don’t need to bond w other atoms

• inert/inactive

• color-less + non-flammable

• boiling point ↑

He, Ne, Ar, Kr, Xe, Rn, Og

transition metals

• groups 3-12

• have diff oxidation states

• valence e- in d orbital so Z_eff is low so e- move easier

• ↑melting/boiling points

• good heat + electricity conductors

• elements found in center need more oxidation states

ionic bonds

• transfer of e- from 1 element’s valence shell to another

• forms ions + bonds w other ions via electrostatic interactions

• bonds constructed via interactions btwn transfer of e-

• higher bond strength than covalent

covalent bonds

atoms share e- w one another to form bonds + fulfill octet/duet rule

octet rule

• chem bonds form where valence shell of all atoms will have 8e-

• mainly applies to s + p orbitals

• peroxides and molecules w transition metals don’t follow this

duet rule

• only H+ + He

• want to have 2e- in valence shell

bonding electrons

2e- shared in bonds

nonbinding electrons

• don’t participate in bonding interactions

• free e- bc they remain free in orbitals not participating in bonding

formal charges

• way to assign hypothetical charges on molecules

• formal charge = valence e- — counted e-

formal charge rules

• sum of formal charges should always = charge of molecules

  • neutral molecule = 0

  • ion = -3

• try to have all formal charges = 0

• put negative formal charges on more electronegative atom

electron geometry

arrangement of all e- pairs

bonding + nonbonding

molecular geometry

arrangement of bonding e-

Van der Waals

• weakest

• rely on weak dipole-dipole interactions

• occurs when there’s a temporary polarization w/in molecules bc of distribution of e- → generates temporary dipole

• hydrophobic + nonpolar molecules only do this one

dipole-dipole interactions

• permanent dipoles formed

• bc of differences in distribution of e-

• ↑electroneg = ↑e- density

• more electronegative atom has (-) dipole + other one has (+) dipole

Hydrogen Bonding

• strongest of 3

• H-bond donor + H-bond acceptor

  • donor = electroneg atom w H

  • acceptor = electroneg atom

• occur w O,F,N

molecular weight

sum of all individual atoms in molecule

mole

6.022×10²³ = 1mole

molar mass

• mass in grams of 1 mole of substance

• # of moles x molar mass

• g/mol

molecular formula

elemental composition of compound

empirical formula

ratio btwn all atoms in 1 compound (simplifying)

synthesis reaction

combine 2+ reactants to form product

2H₂ + O₂ → 2H₂O

decomposition reaction

breaking down reactant to simplest form

2NaCl → 2NA + Cl₂

combustion reaction

• substance (usually hydrocarbon) reacts w O to produce heat, CO₂ + water as byproduct

• CH₄ + 2O₂ → CO₂ + 2H₂O + heat

single displacement reaction

displacement of reactant to form new product

Zn + 2HCl → H₂ + ZnCl₂

double displacement reaction

exchange of ions in compound to form new product

NA₂SO₄ + SrCl₂ → 2NaCl + SrSO₄

law of conservation of mass

• in chem rxn, matter cannot be created/destroyed

• balancing equations so #reactants = #products

activation energy

energy needed for rxn to proceed

rate determining step

• slow step

• determines rate of rxn

catalyst

• increase rate of rxn

• consumed 1st + produced later

• don’t show in equation bc they’re reproduced

intermediates

• chem species formed in 1 step of rxn + consumed in next

• don’t show in equation

multistep chemical reaction

collision theory

• examine collisions btwn reactants that form products

• ↑temp = ↑collision = ↑products

• ↑[reactants] = ↑collision = ↑products

transition state theory

↑activation E → ↓rxn rate

(the humps on energy graph are the states)

reaction rate

rate law

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

finding constant “k” in rate law

either [A] or [B] cancels out

dynamic equilibrium

• reactants turn to products + products to reactant @ same rate

• concentrations remain constant

equilibrium constant

Le Chatelier’s principle

• if system @ equilibrium is disturbed, system will adjust to counteract

• concentration

• pressure

• temp (exothermic = ↑temp = ↑reactants)

ICE Table (initial, change, equilibrium

to find equilibrium concentrations

Haber process

• agriculture

• produces ammonia for fertilizers

• N₂(g) + 3H₂ ←→ 2NH₃(g)

contact process

produces sulfuric acid

2SO₂(g) + O₂(g) ←→ 2SO₃(g)

kinetic product

• form faster

• ↓E_a needed

• favored at ↓temp

• less stable

thermodynamic product

• more stable

• takes longer

• favored @ ↑temp

• needs ↑E_a

Diels-Alder Reaction

diene + dienophile = kinetic + thermodynamic product

enolate formation

• kinetic: fast + uses strong, bulky base (LDA)

• thermodynamic: @ ↑temp + uses smaller base → ↑stable enolate

system (thermodynamics)

environment/universe being studied

surrounding (thermodynamics)

everything outside system

open system

can exchange E + matter w surrounding

closed system

can exchange only E w surrounding

isolated system

cannot exchange anything

state functions (thermodynamics)

• depends only on state of system, not how it got there

• ex: temp, P, V

internal energy (U)

• total E w/in system

• all kinetic + potential E of particles

• ΔU = q + w

• q = heat w = work

enthalpy (H)

• heat content of a system @ constant P

• heat absorbed/released @ constant temp

• ΔH = ΔU + PΔV

• ΔH = ∑H_{products ⁺} ∑H_reactants

• -ΔH = exothermic

entropy (S)

• disorder/randomness of system

• ΔS = q_rev/T

• high S = high disorder

• gas has high entropy

• ΔS = ∑S_{products ⁺} ∑S_reactants

Gibbs free E (G)

• determines if rxn if spontaneous

• ΔG = ΔH - TΔS

• ΔG<0: spontaneous

• ΔG>0: non-spontaneous

• ΔG=0: equilibrium

isothermal process

• occurs @ constant temp

• ΔU = 0

adiabatic process

• occurs w/o heat exchange

• all E Δs come from work done on/by system

isochoric process

• occurs @ constant V

• work = 0

• added heat changes U

isobaric process

• occurs @ constant P

• heat added/removed Δs H

heat capacity (C)

• amt of heat needed to Δ temp of substance by 1^oC

• C = q/ΔT

standard enthalpy of formation (ΔH_f^o)

Δq when 1 mole of compound is formed from its elements in standard states