AP CHEMISTRY SEMESTER 1 STUDY GUIDE
MEASUREMENTS, METRIC SYSTEM, & MATTER:
 Know the 7 fundamental measurements and their S.I. base units.
 Know the numerical meaning of the metric prefixes.
 Give the abbreviation of metric units given the name and vice versa. Also know what an angstrom is, its abbreviation
and numerical value.
 Convert from one metric unit to another.
 Determine the number of significant digits in a quantity.
 Distinguish between the different forms of matter (pure substance, mixture; element, compound, homogeneous and
heterogeneous mixtures).
 Know the 5 evidences of a chemical reaction.
 Distinguish between chemical and physical changes.
NOMENCLATURE:
 Give the formula of a compound given the name and vice versa. This includes all types of compounds—ionic,
molecular, acids and coordination compounds.
 Know the names and symbols of the elements and the name and formula of the common polyatomic ions.
STOICHIOMETRY:
 Perform mass – moles calculations.
 Determine the empirical formula of a compound given:
 the molecular formula
 mass percents of the elements in the compound
 given the experimental data of a reaction
 Determine the mass percent of the elements in a compound.
 Determine the mass of an element given the mass of the compound.
 Determine the molecular formula of a compound given the empirical formula and the molar mass of the compound.
 Perform the following stoichiometric calculations:
 finding the mass of a substance given the mass of another
 problems involving molarity of substances in a reaction
 determining the amount of a product in a limiting reactant situation
 determining the theoretical yield of a reaction
 determining the percent yield of a reaction
 Distinguish between hydrates and anhydrates.
REACTION TYPES:
 Identifying the type of reaction taking place (redox, precipitation, acid/base)
 Indicate whether or not an acid is a strong or weak acid.
 Explain the difference between strong, weak and nonelectrolytes; strong and weak acids.
 Determine the oxidation state of an element in a compound or polyatomic ion.
 Write the complete dissociation equations for an ionic salt and acids.
 Determine the element undergoing reduction and oxidation in a redox reaction; determine the oxidizing and reducing
agents in a redox reaction.
 Balance chemical equations.
 Complete a chemical equation given the reactants—this includes all the different types we have had.
 Write the ionic and net ionic equations for a reaction.
 Balance redox equations for acidic and basic solutions.
ATOMIC STRUCTURE:
 Determine the frequency of an electromagnetic wave given the wavelength and vice versa.
 Describe the significance of the line spectrum of elements and how it indicates that electrons have quantized
energies.
 List the evidences for supporting:
 the particle properties of light
 the wave properties of particles
 What is meant by the dual nature of light?
 Determine the wavelength of a particle using de Broglie’s formula.
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Describe Bohr’s model of the atom and what he used to support this theory.
Give the two reasons for Bohr’s model not working.
Calculate ΔE for Bohr’s hydrogen electron moving from one energy level to another.
Calculate the wavelength and frequency of a photon given off by a hydrogen electron moving from a higher energy
level to a lower one.
Explain what Schrödinger assumed about an electron in coming up with the quantum theory.
Explain how wave functions are quantized and how they are related to orbitals.
Explain what quantum mechanics is based upon.
Give the set of 4 quantum numbers, what they represent, and how they are determined numerically.
Write the set of quantum numbers for given electrons.
Know the relationship between the quantum numbers, the orbitals, and the number of electrons.
Identify the shapes of the orbitals of different sublevels.
Describe how degenerate orbitals are alike and different.
Describe how orbitals, such as a 3p and a 4p, are alike and different.
Explain why we cannot account exactly for the repulsion among electrons in polyelectronic atoms.
Explain what a 4s electron being more penetrating than a 3d electron means.
Explain the emphasis placed upon valence electrons in discussing atomic properties.
Explain the relationship between valence electrons and elements in the same group on the Periodic table.
Give the orbital diagram, electron configuration, and valence electron configuration for elements.
Identify elements based upon electron configurations and quantum numbers of the electrons.
Identify whether or not a configuration is for an excited electron or one in its ground state.
Determine the number of unpaired electrons in an atom.
Explain why the first ionization energy tends to increase as one proceeds across from left to right across a period; why
the first ionization energy of Al is lower than Mg and that of S is lower than P.
Explain why successive ionization energies of an atom always increases.
Explain why Li has a larger second ionization energy than Be.
Arrange elements in order of increasing: atomic size, ionization energy, electron affinity.
Arrange neutral atoms and their ions in order of increasing ionic size.
Explain why atomic size tends to decrease from left to right across a period and increases as one goes down a group.
Explain why the most reactive metals are on the left side of the periodic table and why the most reactive nonmetals
are to the right side of the periodic table (not counting the Noble gases).
Know the different parts of the periodic table.
Explain why metals are reducing agents and nonmetals are oxidizing agents.
Determine which element would be an oxidizing agent with elements of opposite charges.
GENERAL BONDING THEORIES:
 Determine the type of bonds between two atoms.
 Determine the polarity of: a bond, a molecule
 Determine the direction of the dipole moment of a polar covalent bond.
 Determine whether or not species are isoelectronic.
 Compare the bond length and bond energy of bonds based on bond order.
 Know which elements can have multiple bonds.
 Know the steps of the formation of ionic solids and whether or not these steps are exothermic or endothermic.
 Recognize and draw the Lewis Structure of a molecule.
 Explain why CaS releases more energy than NaF when forming solids. Be able to compare compounds in terms of
the amount of lattice energy.
 Recognize when resonance would exist.
 Know the bond angles and hybridization for different shapes.
 Explain why elements form compounds.
 Know why some elements can end up with more than 8 electrons while others cannot go beyond 8. Why can some
atoms have the same number of bonds as valence electrons and others cannot.
 Determine the lattice energy of an ionic compound.
 Determine the shape of a molecule and draw it.
 Define and recognize sigma and pi bonds.
 Know how hybrid orbitals form and determine the hybridization on an atom.
 Determine the formal charges on atoms in a molecule and determine the best structure based on formal charges.
 Determine the intermolecular force (hydrogen bonding, dipole-dipole, London Forces) between particles.
THERMOCHEMISTRY:
 Declare whether or not a measurement is a state function.
 Know the signs related to the thermochemical quantities in regards to exothermic and endothermic processes. This
includes work (even though it is not endothermic or exothermic) for the expansion and compression of gases. This
will be done from the system’s point of view.
 Given a system containing a gas, explain when work will occur. What is the relationship between work, pressure and
volume? What is the significance of the negative sign in the equation w = PΔV?
 Describe what happens in endothermic and exothermic reactions concerning potential energy. Explain why the
overall reaction is endothermic or exothermic.
 Know how ΔH and q are related. When are they equal?
 Determine which substance is a better choice to use in making certain equipment given specific heats of the
substances (i.e. What metal would be better in dissipating the large amount of heat produced by hi-fi power amplifiers
in order to prevent damage to the electrical components, Al or Fe?)
 Explain why we can use Hess’s Law to determine the amount of energy involved in a reaction.
 List the 2 ways in which we can manipulate thermochemical equations when doing Hess’s Law problems and explain
the resulting effect it has on ΔH.
 List the conventional definitions of standard states for substances.
 There are 4 points to consider concerning H f when doing standard enthalpy of formation problems. What are they?
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What is the S.I. base unit for energy? What is it equal to?
What is true about the amount of energy in the universe?
Perform ΔE calculations for systems involving the expansion or compression of gases.
Calculate the specific heat of a substance given calorimetric data.
Perform ΔH problems concerning:
 the amount of heat involved given a thermochemical equation and the mass of a substance in the equation
 the amount of heat involved for a system given ΔE, pressure and volume changes
 using Hess’s Law to determine ΔH for a reaction
 using H f for substances in a reaction to determine ΔH for the reaction
GAS LAWS:
 Know the standard pressures and temperature of gases.
 List the postulates of the kinetic molecular theory for gases.
 State the relationship between the volume, Kelvin temperature, pressure and moles of a sample of gas.
 Compare the velocity of different gas particles.
 Determine the changes in the volume, temperature, moles and pressure of a gas sample if the situation changes—
know for both a closed and open system.
 State how real gases differ from ideal gases.
 State the relationship between kinetic energy and temperature.
 Compare the rate of effusion and diffusion of different gases.
 Explain what corrections van der Waals did to the KMT model of ideal gases to make it work for real gases.
 Determine the pressure using a manometer.
 Use Boyle’s, Charles’, Combined, Ideal Gas, Graham’s, van der Waals’ equation, and Dalton’s laws in various gas
law calculations.
 Perform stoichiometry calculations for reactions involving gases.
 Calculate the average kinetic energy and velocity of a gas.
SOLIDS & LIQUIDS:
 Identify the type of solid a substance is.
 Distinguish between the properties of the different types of solids.
 Know what is meant by “normal boiling point” for a liquid.
 Perform calculations involving the vapor pressure and temperature of liquids at different situations—includes normal
boiling points.
 Distinguish between substances in terms of properties dependent on IMF’s (vapor pressure, boiling point, surface
tension, viscosity, and capillary action.
 Interpret a phase diagram.
 Perform calculations involving heats of fusion and vaporization and q to find the heat involved with changing physical
states of a substance.
VOCABULARY:
st
1 Law of Thermodynamics
accuracy
anhydrate
anion
bond energy
bond length
bond order
cation
chemical change
compound
continuous spectrum
coordination number
covalent bonds
critical point
dative bonds
degenerate orbital
diffraction
diffusion
dipole moment
effusion
electrolyte
electron affinity
electronegativity
element
empirical formula
endothermic
endpoint
energy
enthalpy
exothermic
formal charge
frequency
ground state
heat
Hess’s law
heterogeneous mixture
homogeneous mixture
Hund’s rule
hybrid orbital
hydrate
immiscible
ionic bonds
ionization energy
isoelectronic
kinetic energy
lattice
lattice energy
Law of Conservation of Energy
limiting reactant
line spectrum
matter
miscible
EQUATIONS GIVEN:
CmmmΔTm = [CwmwΔTw + (Ccal)(ΔTw)]
q = CcalΔT
 1
1 
18
ΔE = 2.178 x 10ˉ J  2 - 2 
n
n i 
 f
rate1

rate 2
MM2
MM1
molar mass
mole
molecular formula
molecule
node
nonelectrolyte
nonpolar covalent bonds
orbital
oxidation
oxidation state
oxidizing agent
partial pressure
Pauli exclusion principle
photon
physical change
pi bonds
polar covalent bonds
potential energy
precision
pressure
quantized
quantum
reducing agent
reduction
resonance
shielding effect
Q Q
E  2.31  10 19 J  nm 1 2
 r
(KE)avg =
urms =
Pobs =
(
INFO GIVEN:
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Periodic table
electronegativity chart
all constants needed
activity series
list of complex ions
3
2

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
=
h
mv
c = 
RT
E = h
3RT
M
E=
nRT
n
 a 
V  nb
V
)
sigma bonds
significant digits
solute
solution
solvent
specific heat capacity
spectator ions
standard enthalpy of formation
standard states
standing wave
state function
surface tension
surroundings
system
temperature
theoretical yield
valence electrons
viscosity
wave function
wavelength
work
(
ΔE = q + w
q = mCΔT
w = PΔV
PV = nRT
hc

2
P1 P2

n1 n2
)
P1 V1 P2 V2

T1
T2