AP Chemistry Study Guide

molecule

2 or more of the same elements bonded together

compound

2 or more different elements bonded together

law of definite proportions

all samples of a compound contain the same elements in the same ratio

heterogeneous

non-uniform mixture

homogeneous

uniform mixture

physical change

change observed without altering identity or composition

chemical change

change where a substance transforms into another substance

intensive property

does not depend on the amount of substance

extensive property

depends on the amount of substance

density

mass divided by volume

significant figures

all non-zero digits are significant

sigfig rule 1

zeros between non-zero digits are significant

sigfig rule 2

leading zeros are not significant

sigfig rule 3

trailing zeros are only significant if there's a decimal point

sigfigs for addition/subtraction

result has same number of decimal places as the number with the fewest decimal places

sigfigs for multiplication/division

result has same number of sigfigs as the number with the fewest sigfigs

sigfigs for logs

number of sigfigs in original value equals number of decimal places in the result

-log(X.XX)

result will have three decimal places (X.XXX)

10^(-x.xx)

result will have 3 sigfigs total

dalton’s atomic theory

each element is made of atoms that are identical for that element but different from other elements

law of multiple proportions

if 2 elements form multiple compounds

j.j. thomson

discovered the electron using the cathode ray tube

robert millikan

measured charge of electron using oil drop experiment

henri becquerel

discovered radioactivity

ernest rutherford

discovered nucleus via gold foil experiment

james chadwick

discovered neutron

alpha radiation

2+ charge

beta radiation

1- charge

gamma radiation

high energy

atomic number

number of protons in an atom

mass number

number of protons plus neutrons

isotopes

same number of protons but different number of neutrons

atomic weight

weighted average of isotope masses based on abundance

molecular formula

shows exact number of atoms of each element in a compound

empirical formula

simplest whole-number ratio of atoms in a compound

ion

atom with extra or missing electrons

cation

positively charged ion

anion

negatively charged ion

ionic compound

compound made from a cation and an anion (metal + nonmetal)

isomers

compounds with same molecular formula but different arrangements

law of conservation of mass

matter is not created or destroyed in chemical reactions

combination reaction

A + B → AB

decomposition reaction

AB → A + B

combustion reaction

hydrocarbon + O2 → CO2 + H2O

formula weight

sum of atomic weights of all atoms in a chemical formula

molecular weight

formula weight of a molecule

percentage composition

percent mass of each element in a compound

avogadro’s number

6.02 × 10^23 particles/mol

molar mass

formula weight expressed in grams/mol

grams to moles

divide by molar mass

moles to formula units

multiply by Avogadro’s number

empirical formula steps

find mass → convert to moles → divide by smallest → multiply to get whole numbers

molecular formula steps

find empirical formula mass → divide molar mass by empirical → multiply subscripts

limiting reactant

reactant that produces the least amount of product

percent yield

(experimental yield / theoretical yield) × 100

solvent

substance present in the greatest quantity in a solution

solute

substance that dissolves in the solvent

strong electrolytes

soluble ionic compounds that completely dissociate in water

solubility

maximum amount of a substance that can dissolve in a solvent at a given temperature

exchange reaction

AX + BY → AY + BX (one product must be insoluble)

molecular equation

shows all compounds as neutral substances

complete ionic equation

shows all strong electrolytes as ions

spectator ions

ions that don’t participate in the actual reaction

net ionic equation

complete ionic equation without spectator ions

redox reaction

reaction involving transfer of electrons

oxidation

loss of electrons or increase in oxidation number

reduction

gain of electrons or decrease in oxidation number

oxidation number of elements

0 for pure elements

oxidation number of monatomic ions

equal to charge of ion

oxidation number of oxygen

-2 normally

oxidation number of hydrogen

+1 with nonmetals

oxidation number of fluorine

-1 always

oxidation number of compound

sum of oxidation numbers = 0

oxidation number of polyatomic ion

sum of oxidation numbers = ion’s charge

group 1A oxidation number

+1 in ionic compounds

group 2A oxidation number

+2 in ionic compounds

group 3A oxidation number

+3 in ionic compounds

displacement reaction

A + BX → AX + B

activity series

list of metals ranked by ease of oxidation

molarity (M)

moles of solute / liters of solution

standard solution

solution with a known concentration

reaction prediction based on activity series

metal must be higher than the compound metal in the activity series to react

ionic compound in water

dissociates into ions and breaks intermolecular forces

Ek

½mv² (kinetic energy

Work

Force * distance

ΔE

Efinal - Einitial

ΔE = q + w

q is heat added to or liberated from system

Endothermic

heat absorbed

Exothermic

heat released

1st law of thermodynamics

energy is conserved

Enthalpy (H)

accounts for heat flow in processes occurring at constant pressure

relationship/equation between enthalpy, moles, and heat/energy transfer

enthalpy per mole

ΔH > 0

system gains heat (endothermic)

ΔH < 0

system loses heat (exothermic)

Enthalpy change rules

What happens to delta ΔH when coefficients are multiplied by x

ΔH * x

what happens to enthalpy change when you reverse a reaction

delta H switches sign

– Phase/state change affects ΔH

Hess’s Law

sum of ΔH for multiple steps = ΔH for overall reaction

Heat capacity (c)

amount of heat required to raise temp by 1K