AP Chemistry MASTER REVIEW

all units + memorization

convert moles to grams

multiply by molar mass

convert moles to atoms/particles

multiply by Avogadro's number

Avogadro's number

$$6.022\cdot10^{23}$$

convert grams to moles

divide by molar mass

convert atoms to moles

divide by Avogadro's number

to convert between substances, always use

moles

with answers always round to the

lowest given sig fig

how to calculate average atomic mass

take a weighed average (multiply each isotope by its percentage and add all of them together)

mass spectra graph

peaks represent isotopes

what do the heights of the peaks represent in a mass spectra graph?

higher abundance

molecular formula

describes actual number of atoms

empirical formula

simplified smallest ratio of atoms

can molecular formula = empirical formula

yes!

how to calculate the empirical formula

convert to moles of each element and find mole ratio (divide by smallest moles)

what do you do when given percentages?

assume a 100g total sample

pure substance vs. mixture

mixture consists of 2 or more pure substances with different properties and can be physically separated, while pure substances are just one element/compound

homogenous mixture

uniform composition

heterogenous mixture

different composition, visible differences

ways to separate a mixture

physical separation, filtration, distillation, crystallization

filtration

pour heterogenous mixture with a precipitate on filter paper

distillation

separate 2 liquids based on boiling point

coulomb’s law

describes attractive/repulsive forces at the nuclear level based on distance and charge

electron configuration

arrangement of electrons in an atom

coefficient in electron config

energy level

letter in electron config

type of sublevel

superscript in electron config

number of electrons in each sublevel

types of sublevels

s, p, d, f

how many electrons in s

2

how many electrons in p

6

how many electrons in d

10

how many electrons in f

14

aufbau principle

each electron occupies lowest energy orbital available

exceptions to aufbau principle

Cr, Mo, W, Cu, Ag, and Au pull 1 electron from s to put it in d for stability (ex: Cr electron config is [Ar] 4s¹3d⁵)

pauli exclusion principle

max 2 electrons in an orbital and must have opposite spins

hund’s rule

electrons will spread out over equal-level orbitals before more electrons can occupy the same orbital

how to write electron config

determine the number of electrons and go from there/look at the periodic table

group 1 & 2 (including helium)

s

groups 3-8

p

transition metals

d

inner transition metals

f

d energy levels always

1 less than s & p

f energy levels always

2 less than s & p

noble gas notation

abbreviate electron config by writing the last noble gas in brackets and continuing from there

orbital diagrams

visually show electron config & direction of spin

valence electrons

found in atom’s highest energy level (outermost orbitals); determine chemical properties and reactivity

nuclear charge (Z_eff)

attractive force of the nucleus on valence electrons

trend of nuclear charge

increases as you move from left to right across a period because there are more protons in the nucleus

shielding effect

minimized attraction between protons & valence electrons because of layers of electrons in between

trend of shielding effect

attraction decreases as you move from top to bottom across a group becuase there are more layers—more distance to valence electrons

magnetism

depends on the number of unpaired electrons

paramagnetism

1 or more unpaired electron, attracted to magnetic field

diamagnetism

no unpaired electrons, not attracted to magnetic field

photoelectron spectroscopy (PES)

graph of the binding energy of sublevels

what does the height of a peak in a PES mean

number of electrons

how to figure out the y-axis scale

first peak will always be 1s²

binding energy

amount of energy to add/remove electron

atomic radius

size of an atom

atomic radius trend

decreases as you move left to right, increases as you move down

ionization energy

how much energy required to remove a valence electron

ionization energy trend

increases as you move left to right, decreases as you move down

how to identify the amount of valence electrons from an ionization energy chart

when there’s a big jump in ionization energy, you removed a core electron (i.e. jump between 2nd and 3rd ionization energy means there’s 2 valence electrons)

electron affinity

energy released when an atom accepts an electron

electron affinity trend

increases as you move left to right, decreases as you move down

electronegativity

ability of an atom to attract electrons in a bond

electronegativity trend

increases as you move from left to right, decreases as you move downe

what group isn’t electronegative?

noble gases except those who have an expanded octet

ionic radius

size of an ion

compare cation size to neutral ion

cations smaller than neutral atom because the loss of electrons causes it to lose an energy level

compare anion size to neutral ion

anions larger than neutral atom because the added electrons outnumber the protons, rendering the nuclear charge less effective

in order for an ionic bond to form, the elements must have

a big difference in electronegativity (usually metal to nonmetal)

what charge do ionic compounds have?

neutral

process of ionic bonding

metal gives away electron(s) to nonmetal

metallic bonding

sea of electrons act as a buffer between repelling cations, between atoms with similar electronegativities

properties of metallic bonds

• good conductors

• high MP and BP

• malleable

nonpolar covalent bonding

equal/almost equal sharing of electrons, between atoms with similar high electronegativities (same element or C to H)

polar covalent bonding

unequal sharing of electrons, creates dipoles (partial negative/positive poles), between atoms of different high electronegativities

where do the electrons localize closer to in a polar covalent molecule

the atom with higher electronegativity

properties of covalent bonds

• poor conductors

• when in networks, high MP and BP

• when left as molecules, low MP and BP

• soft or brittle

ionic bonding

transfer electrons, between atoms with very different electronegativities, attraction between cations and anions

where do the electrons transfer to in an ionic bond

the atom with the higher electronegativity (nonmetal)

properties of ionic bonds

• poor conductors when solid

• good conductors when liquid

• form networks - high MP and BP

• brittle

delocalized electrons

electrons that can move around freely

localized electrons

“fixed” electrons restricted to a certain region

intramolecular force

attraction in a bond within the molecule

bond energy/enthalpy

energy released when a bond forms/energy needed to break a bond, refers to covalent bonds

potential energy diagrams

relationship between potential energy (y) and bond length (x)

low potential energy means

more attraction (bond)

what creates higher attraction

smaller atoms, charged atoms, more bonds (double/triple bonds)

lattice energy

amount of energy needed to separate 1 mole of an ionic compound

always mention this when talking about lattice energy

coulomb’s law

coulomb’s law

attraction between two particles is based on charge and distance

which component of attraction is more important

charge

higher lattice energy

higher MP and BP

rxns to form ionic compounds typically

endothermic

rxns to form ionic compounds from gaseous ions are

exothermic because they need to come together to stabilize

lattice structure based on

minimized repulsions and maximized attraction (alternating cation-anion pattern)

alloy

replacing 1 metal atom with another metal atom, usually stronger than metals

types of alloys

substitutional and interstitial

substitutional alloy

consists of 2 metal atoms with similar atomic radii

properties of substitutional alloys

malleable, stronger, retain many metallic properties