DAT Bootcamp General Chemistry

atom

smallest unit of an element

molecule

atom bonded to at least one other atom

ion

atom that has lost or gained an electron (charged)

allotrope

elemental molecule with different formulas,

example: O2 and O3

ionic compound

metal + nonmetal. high MP, high BP, brittle, hard.

held together by ionic interactions (lattice energy). ex) NaCl or MgO

molecular compound

2+ nonmetals. low MP, do not conduct electricity. held together by intermolecular forces. ex) H2O or Cl2 or CH2

SA and SB

Percent error

actual - theoretical / theoretical

yield of reaction

experimental yield/theoretical yield

Beer-Lambert Law

absorbance = (molar absorptivity)(c)(l)

mass number

protons + neutrons

atomic number

number of protons

Bohr Model

model of an atom that shows electrons in circular orbits around the nucleus

Paramagnetic

unpaired electrons, attracted to magnets ex. O2

Diamagnetic

paired electrons, slightly repelled by magnets ex. N2

principal quantum number

n

shell (distance from nucleus)

(l, infinity)

azimuthal quantum number

l

subshell

l=0 s, l=1 p, l=2 d, l=3 f

magnetic quantum number

ml

specific orbital

-l,...,+l

spin quantum number

ms

electrons up or down

-1/2, +1/2

Aufbau principle

electrons fill the lowest energy orbitals first

Hund's rule

don't pair electrons until needed

Pauli exclusion principle

no 2 electrons in the same atom can have the same 4 quantum numbers

Heisenberg uncertainty

impossible to determine particles position and momentum with perfect accuracy

ionic bonds

transfer electrons between elements with large difference in electronegativity

covalent bonds

sharing of electrons between 2 elements with similar electronegativity

octet rule

every atom wants 8 electrons

octet rule exceptions

H = 2e-, Be = 4e-, B & Al = 6e-, 3rd row and lower can have more than 8

Lattice energy

energy required to completely separate an ionic compounds cations from its anions. larger charge = larger _______ ______. shorter bond distance = larger _______ ______.

single bond

1 sigma bond

double bond

1 sigma and 1 pi bond

triple bond

1 sigma and 2 pi bonds

bond order

The average number of bonds per atom covalently bonded to a central atom

sp

180º bond angle

linear

sp2

120º bond angle

trigonal planar

sp3

109.5º bond angles

tetrahedral

sp3d

90º and 120º bond angles

trigonal bipyramidal

sp3d2

90º bond angle

octahedral

Alkali metals

group 1A

low ionization energy

form ionic bonds

v reactive w water

Alkaline earth metals

group 2A

low ionization energy

reactive w water (increasing reactivity down a column)

Halogens

highly electronegative

high electron affinity

highly reactive with metals

good oxidizing agent

Noble Gases

inert gases

Transition metals

partially filled d-orbitals

several oxidation states

Have high melting and boiling points

hard and have high densities

Can form compounds that are paramagnetic

Can form compounds that are excellent catalysts

Can form colored compounds

Electronegativity

how much something wants to accept an electron.

increases up and to the right

Electron Affinity

energy given off when atom gains an electron

increases up and to the right

Ionization energy

energy required to remove an electron from an atom.

increases up and to the right

Ideal Gas Assumptions

1) volume or size of each individual gas molecule is insignificant

2) gas molecules collisions with each other are perfectly elastic, no IM forces

3) average KE of a gas depends only on the system temperature

gases like what temp and what pressure

high temp low pressure

Boyle's Law

P1V1=P2V2

Charle's Law

V1/T1=V2/T2

Avogadro's Law

V1/n1=V2/n2

Ideal Gas Law

PV=nRT

R

0.0821 L atm / mol K

STP

1atm, 273 K

Gas density equation

Dalton's Law of Partial Pressure

Pt=Pa+Pb+Pc+....

Pa=XaPt

Graham's Law of diffusion

rate a/rate b=square root of molar mass a/ square root of molar mass b

Ion-dipole bond

intermolecular force between ions and polar substances

Hydrogen bonding

N-H, F-H, O-H

Dipole-dipole bond

IM force between polar molecules

London-dispersion force

between all molecules

greater weight = greater force

intermolecular forces (strongest to weakest)

1. Ion dipole

2. hydrogen bonding

3. dipole-dipole

4. dipole induced dipole

5. london disperson/vanderwaals

Intramolecular forces (strongest to weakest)

1. ionic

2 polar covalent bonds

3. metallic

4. non polar covalent bonds

higher IM forces lead to...

higher BP

higher heat of vaporization

higher viscosity

higher surface tension

lower vapor pressure

simple cubic

A unit cell that consists of a cube with one atom at each corner.

body-centered

atoms at corners and one in the center.

face-centered

there are atoms or ions at the corners and the center of each face of the imaginary cube, 4 atoms

Sublimination

solid to gas

Fusion

solid to liquid

Vaporization

liquid to gas

Deposition

gas to solid

Condensation

gas to liquid

Crystallization

liquid to solid

solution

uniform mixture of 2+ substances

solvent

substance present in a larger amount

solute

substance present in a smaller amount

unsaturated solution

less than max amount of solute is dissolved in solvent

saturated solution

max amount of solute is dissolved in solvent

Molarity

mol solute/L solution

molality

mol solute/kg solvent

Freezing point depression

ΔTf= -iKfm

Boiling point elevation

ΔTb= iKbm

osmotic pressure

=iMRT

k unit 0 order

M1s-1

k unit 1st order

s-1

k unit 2nd order

M-1 s-1

k unit 3rd order

M-2s-1

Collision theory

both molecules must collide

both must collide with enough energy

both must collide in correct 3D orientation

catalyst

speed up rxn without being consumed

lowers activation energy

provides alternate pathway

does NOT shift equilibrium

Arrhenius equation

Kc

[products]/[reactants]

Kp

P products/P reactants

Keq

Kforward/Kreverse

Keq>1

products favored

Keq

reactants favored

Q

shift towards product, no ppt

Q>K

shift towards reactants, ppt formed

Q=K

at equilibrium

Arrhenius acid

H+ donor in water

Arrhenius base

OH- donor in water

Bronsted Lowry acid

H+ donor