List topics that need to be taught in order to cover each the content in each knowledge statement thoroughly.
SPQ-1.A.1
One cannot count particles directly while
performing laboratory work. Thus, there must
be a connection between the masses of
substances reacting and the actual number of
particles undergoing chemical changes.
SPQ-1.A.2
Avogadro’s number (N = 6.022 × 1023 mol−1
A )
provides the connection between the number
of moles in a pure sample of a substance and
the number of constituent particles (or formula
units) of that substance.
SPQ-1.A.3
Expressing the mass of an individual atom
or molecule in atomic mass units (amu) is
useful because the average mass in amu of
one particle (atom or molecule) or formula
unit of a substance will always be numerically
equal to the molar mass of that substance
in grams. Thus, there is a quantitative
connection between the mass of a substance
and the number of particles that the
substance contains.
EQN: n = m/M
SPQ-1.B.1
The mass spectrum of a sample containing a
single element can be used to determine the
identity of the isotopes of that element and the
relative abundance of each isotope in nature.
SPQ-1.B.2
The average atomic mass of an element can
be estimated from the weighted average of
the isotopic masses using the mass of each
isotope and its relative abundance.
X INTERPRETING MASS SPECTRA
Interpreting mass spectra of samples
containing multiple elements or peaks arising
from species other than singly charged
monatomic ions will not be assessed on
the AP Exam.
SPQ-2.A.1
Some pure substances are composed of
individual molecules, while others consist
of atoms or ions held together in fixed
proportions as described by a formula unit.
SPQ-2.A.2
According to the law of definite proportions,
the ratio of the masses of the constituent
elements in any pure sample of that compound
is always the same.
SPQ-2.A.3
The chemical formula that lists the lowest
whole number ratio of atoms of the elements in
a compound is the empirical formula.
SPQ-2.B.1
While pure substances contain molecules or
formula units of a single type, mixtures contain
molecules or formula units of two or more
types, whose relative proportions can vary.
SPQ-2.B.2
Elemental analysis can be used to determine
the relative numbers of atoms in a substance
and to determine its purity.
SAP-1.A.1
The atom is composed of negatively charged
electrons and a positively charged nucleus that
is made of protons and neutrons.
SAP-1.A.2
Coulomb’s law is used to calculate the force
between two charged particles.
EQN: Fcoulombic ∝ q1
q2
r2
SAP-1.A.3
In atoms and ions, the electrons can be thought
of as being in “shells (energy levels)” and
“subshells (sublevels),” as described by the
electron configuration. Inner electrons are called
core electrons, and outer electrons are called
valence electrons. The electron configuration is
explained by quantum mechanics, as delineated
in the Aufbau principle and exemplified in the
periodic table of the elements.
X THE ASSIGNMENT OF QUANTUM NUMBERS
TO ELECTRONS IN SUBSHELLS OF AN
ATOM WILL NOT BE ASSESSED ON THE
AP EXAM.
Rationale: Assignment of quantum numbers
to electrons in specific subshells does not
increase students’ understanding of the
structure of the atom.
SAP-1.A.4
The relative energy required to remove an
electron from different subshells of an atom or
ion or from the same subshell in different atoms
or ions (ionization energy) can be estimated
through a qualitative application of Coulomb’s
law. This energy is related to the distance from
the nucleus and the effective (shield) charge of
the nucleus. 
SAP-1.B.1
The energies of the electrons in a given
shell can be measured experimentally with
photoelectron spectroscopy (PES). The
position of each peak in the PES spectrum is
related to the energy required to remove an
electron from the corresponding subshell, and
the height of each peak is (ideally) proportional
to the number of electrons in that subshell.
SAP-2.A.1
The organization of the periodic table is based
on the recurring properties of the elements and
explained by the pattern of electron configurations
and the presence of completely or partially filled
shells (and subshells) of electrons in atoms.
SAP-2.A.2
Trends in atomic properties within the
periodic table (periodicity) can be qualitatively
understood through the position of the element
in the periodic table, Coulomb’s law, the shell
model, and the concept of shielding/effective
nuclear charge. These properties include:
a. Ionization energy
b. Atomic and ionic radii
c. Electron affinity
d. Electronegativity.
SAP-2.A.3
The periodicity (in SAP-2.A.2) is useful to
predict /estimate values of properties in the
absence of data.
SAP-2.B.1
The likelihood that two elements will form
a chemical bond is determined by the
interactions between the valence electrons
and nuclei of elements.
SAP-2.B.2
Elements in the same column of the periodic
table tend to form analogous compounds.
SAP-2.B.3
Typical charges of atoms in ionic compounds
are governed by their location on the periodic
table and the number of valence electrons.
SAP-3.A.1
Electronegativity values for the representative
elements increase going from left to right
across a period and decrease going down
a group. These trends can be understood
qualitatively through the electronic structure of
the atoms, the shell model, and Coulomb’s law.
SAP-3.A.2
Valence electrons shared between atoms
of similar electronegativity constitute a
nonpolar covalent bond. For example, bonds
between carbon and hydrogen are effectively
nonpolar even though carbon is slightly more
electronegative than hydrogen.
SAP-3.A.3
Valence electrons shared between atoms of
unequal electronegativity constitute a polar
covalent bond.
a. The atom with a higher electronegativity will
develop a partial negative charge relative to
the other atom in the bond.
b. In single bonds, greater differences in
electronegativity lead to greater bond dipoles.
c. All polar bonds have some ionic character,
and the difference between ionic and
covalent bonding is not distinct but rather
a continuum.
SAP-3.A.4
The difference in electronegativity is not the
only factor in determining if a bond should
be designated as ionic or covalent. Generally,
bonds between a metal and nonmetal are
ionic, and bonds between two nonmetals are
covalent. Examination of the properties of a
compound is the best way to characterize the
type of bonding.
SAP-3.A.5
In a metallic solid, the valence electrons
from the metal atoms are considered to be
delocalized and not associated with any
individual atom.
SAP-3.B.1
A graph of potential energy versus the distance
between atoms is a useful representation for
describing the interactions between atoms.
Such graphs illustrate both the equilibrium
bond length (the separation between atoms at
which the potential energy is lowest) and the
bond energy (the energy required to separate
the atoms).
SAP-3.B.2
In a covalent bond, the bond length is
influenced by both the size of the atom’s core
and the bond order (i.e., single, double, triple).
Bonds with a higher order are shorter and have
larger bond energies.
SAP-3.B.3
Coulomb’s law can be used to understand the
strength of interactions between cations
and anions.
a. Because the interaction strength is
proportional to the charge on each ion, larger
charges lead to stronger interactions.
b. Because the interaction strength increases
as the distance between the centers of the
ions (nuclei) decreases, smaller ions lead to
stronger interactions.
SAP-3.C.1
The cations and anions in an ionic crystal are
arranged in a systematic, periodic 3-D array
that maximizes the attractive forces among
cations and anions while minimizing the
repulsive forces.
X KNOWLEDGE OF SPECIFIC TYPES OF
CRYSTAL STRUCTURES WILL NOT BE
ASSESSED ON THE AP EXAM.
Rationale: Study of specific crystal structures is
not essential to an understanding of the big ideas.
SAP-3.D.1
Metallic bonding can be represented as an
array of positive metal ions surrounded by
delocalized valence electrons (i.e., a “sea
of electrons”).
SAP-3.D.2
Interstitial alloys form between atoms of
different radii, where the smaller atoms fill the
interstitial spaces between the larger atoms
(e.g., with steel in which carbon occupies the
interstices in iron).
SAP-3.D.3
Substitutional alloys form between atoms
of comparable radius, where one atom
substitutes for the other in the lattice. (In
certain brass alloys, other elements, usually
zinc, substitute for copper.)
Lewis diagrams can be constructed according
to an established set of principles.
SAP-4.B.1
In cases where more than one equivalent Lewis
structure can be constructed, resonance
must be included as a refinement to the Lewis
structure. In many such cases, this refinement
is needed to provide qualitatively accurate
predictions of molecular structure
and properties.
SAP-4.B.2
The octet rule and formal charge can be used
as criteria for determining which of several
possible valid Lewis diagrams provides the
best model for predicting molecular structure
and properties.
SAP-4.B.3
As with any model, there are limitations to
the use of the Lewis structure model,
particularly in cases with an odd number of
valence electrons
SAP-4.C.1
VSEPR theory uses the Coulombic repulsion
between electrons as a basis for predicting
the arrangement of electron pairs around a
central atom.
SAP-4.C.2
Both Lewis diagrams and VSEPR theory must
be used for predicting electronic and structural
properties of many covalently bonded
molecules and polyatomic ions, including
the following:
a. Molecular geometry
b. Bond angles
c. Relative bond energies based on bond order
d. Relative bond lengths (multiple bonds,
effects of atomic radius)
e. Presence of a dipole moment
f. Hybridization of valence orbitals of
the molecule
SAP-4.C.3
The terms “hybridization” and “hybrid atomic
orbital” are used to describe the arrangement
of electrons around a central atom. When the
central atom is sp hybridized, its ideal bond
angles are 180°; for sp2 hybridized atoms the
bond angles are 120°; and for sp3 hybridized
atoms the bond angles are 109.5°.
X AN UNDERSTANDING OF THE DERIVATION
AND DEPICTION OF HYBRID ORBITALS
WILL NOT BE ASSESSED ON THE AP EXAM.
Rationale: The course includes the distinction
between sigma and pi bonding, the use of VSEPR
to explain the shapes of molecules, and the sp,
sp2
, and sp3
 nomenclature. Additional aspects
related to hybridization are both controversial and
do not substantially enhance understanding of
molecular structure.
X HYBRIDIZATION INVOLVING D ORBITALS
WILL NOT BE ASSESSED ON THE AP EXAM.
WHEN AN ATOM HAS MORE THAN FOUR
PAIRS OF ELECTRONS SURROUNDING THE
CENTRAL ATOM, STUDENTS ARE ONLY
RESPONSIBLE FOR THE SHAPE OF THE
RESULTING MOLECULE.
Rationale: Current evidence suggests that
main-group hybridization involving d orbitals
does not exist, and there is controversy about the
need to teach any hybridization. Until agreement
is reached in the chemistry community, we
will continue to include only sp, sp2
, and sp3
hybridization on the AP Exam.
SAP-4.C.4
Bond formation is associated with overlap
between atomic orbitals. In multiple bonds, such
overlap leads to the formation of both sigma
and pi bonds. The overlap is stronger in sigma
than pi bonds, which is reflected in sigma bonds
having greater bond energy than pi bonds. The
presence of a pi bond also prevents the rotation
of the bond and leads to structural isomers. 
SAP-5.A.1
London dispersion forces are a result of the
Coulombic interactions between temporary,
fluctuating dipoles. London dispersion forces
are often the strongest net intermolecular
force between large molecules.
a. Dispersion forces increase with increasing
contact area between molecules and with
increasing polarizability of the molecules.
b. The polarizability of a molecule increases
with an increasing number of electrons in the
molecule; and the size of the electron cloud.
It is enhanced by the presence of pi bonding.
c. The term “London dispersion forces” should
not be used synonymously with the term
“van der Waals forces.”
SAP-5.A.2
The dipole moment of a polar molecule
leads to additional interactions with other
chemical species.
a. Dipole-induced dipole interactions are
present between a polar and nonpolar
molecule. These forces are always attractive.
The strength of these forces increases with
the magnitude of the dipole of the polar
molecule and with the polarizability of the
nonpolar molecule.
b. Dipole-dipole interactions are present
between polar molecules. The interaction
strength depends on the magnitudes of
the dipoles and their relative orientation.
Interactions between polar molecules
are typically greater than those between
nonpolar molecules of comparable size
because these interactions act in addition to
London dispersion forces.
c. Ion-dipole forces of attraction are present
between ions and polar molecules.
These tend to be stronger than dipoledipole forces.
SAP-5.A.3
The relative strength and orientation
dependence of dipole-dipole and ion-dipole
forces can be understood qualitatively by
considering the sign of the partial charges
responsible for the molecular dipole moment,
and how these partial charges interact with an
ion or with an adjacent dipole.
SAP-5.A.4
Hydrogen bonding is a strong type of
intermolecular interaction that exists when
hydrogen atoms covalently bonded to the
highly electronegative atoms (N, O, and F)
are attracted to the negative end of a dipole
formed by the electronegative atom (N, O,
and F) in a different molecule, or a different part
of the same molecule.
SAP-5.A.5
In large biomolecules, noncovalent interactions
may occur between different molecules
or between different regions of the same
large biomolecule.
SAP-5.B.1
Many properties of liquids and solids are
determined by the strengths and types of
intermolecular forces present. Because
intermolecular interactions are broken when a
substance vaporizes, the vapor pressure and
boiling point are directly related to the strength
of those interactions. Melting points also tend
to correlate with interaction strength, but
because the interactions are only rearranged,
in melting, the relations can be more subtle.
SAP-5.B.2
Particulate-level representations, showing
multiple interacting chemical species, are a
useful means to communicate or understand
how intermolecular interactions help to
establish macroscopic properties.
SAP-5.B.3
Due to strong interactions between ions,
ionic solids tend to have low vapor pressures,
high melting points, and high boiling points.
They tend to be brittle due to the repulsion
of like charges caused when one layer slides
across another layer. They conduct electricity
only when the ions are mobile, as when the
ionic solid is melted or dissolved in water or
another solvent.
SAP-5.B.4
In covalent network solids, the atoms are
covalently bonded together into a threedimensional network (e.g., diamond) or layers
of two-dimensional networks (e.g., graphite).
These are only formed from nonmetals:
elemental (e.g., diamond, graphite) or binary
compounds of two nonmetals (e.g., silicon
dioxide and silicon carbide). Due to the strong
covalent interactions, covalent solids have high
melting points. Three-dimensional network
solids are also rigid and hard, because the
covalent bond angles are fixed. However,
graphite is soft because adjacent layers can
slide past each other relatively easily.
SAP-5.B.5
Molecular solids are composed of distinct,
individual units of covalently-bonded
molecules attracted to each other through
relatively weak intermolecular forces. Molecular
solids generally have a low melting point
because of the relatively weak intermolecular
forces present between the molecules. They
do not conduct electricity because their
valence electrons are tightly held within the
covalent bonds and the lone pairs of each
constituent molecule. Molecular solids are
sometimes composed of very large molecules
or polymers.
SAP-5.B.6
Metallic solids are good conductors of
electricity and heat, due to the presence of
free valence electrons. They also tend to be
malleable and ductile, due to the ease with
which the metal cores can rearrange their
structure. In an interstitial alloy, interstitial
atoms tend to make the lattice more rigid,
decreasing malleability and ductility. Alloys
typically retain a sea of mobile electrons and so
remain conducting.
SAP-5.B.7
In large biomolecules or polymers, noncovalent
interactions may occur between different
molecules or between different regions of the
same large biomolecule. The functionality and
properties of such molecules depend strongly
on the shape of the molecule, which is largely
dictated by noncovalent interactions.
SAP-6.A.1
Solids can be crystalline, where the particles
are arranged in a regular three-dimensional
structure, or they can be amorphous, where
the particles do not have a regular, orderly
arrangement. In both cases, the motion of the
individual particles is limited, and the particles
do not undergo overall translation with respect
to each other. The structure of the solid is
influenced by interparticle interactions and the
ability of the particles to pack together.
SAP-6.A.2
The constituent particles in liquids are in
close contact with each other, and they
are continually moving and colliding. The
arrangement and movement of particles are
influenced by the nature and strength of the
forces (e.g., polarity, hydrogen bonding, and
temperature) between the particles.
SAP-6.A.3
The solid and liquid phases for a particular
substance typically have similar molar volume
because, in both phases, the constituent
particles are in close contact at all times.
SAP-6.A.4
In the gas phase, the particles are in constant
motion. Their frequencies of collision and the
average spacing between them are dependent
on temperature, pressure, and volume. Because
of this constant motion, and minimal effects of
forces between particles, a gas has neither a
definite volume nor a definite shape.
X UNDERSTANDING/INTERPRETING PHASE
DIAGRAMS WILL NOT BE ASSESSED ON
THE AP EXAM.
Rationale: Phase diagrams of pure substances
are considered prior knowledge.
SAP-7.A.1
The macroscopic properties of ideal gases are
related through the ideal gas law:
EQN: PV = nRT.
SAP-7.A.2
In a sample containing a mixture of ideal gases,
the pressure exerted by each component (the
partial pressure) is independent of the other
components. Therefore, the total pressure of
the sample is the sum of the partial pressures.
EQN: PA = Ptotal × XA,
where XA = moles A/total moles;
EQN: Ptotal = PA + PB + PC + …
SAP-7.A.3
Graphical representations of the relationships
between P, V, T, and n are useful to describe
gas behavior.
SAP-7.B.1
The kinetic molecular theory (KMT) relates
the macroscopic properties of gases to
motions of the particles in the gas. The
Maxwell-Boltzmann distribution describes the
distribution of the kinetic energies of particles
at a given temperature.
SAP-7.B.2
All the particles in a sample of matter are in
continuous, random motion. The average
kinetic energy of a particle is related to its
average velocity by the equation:
EQN: KE = ½ mv2
.
SAP-7.B.3
The Kelvin temperature of a sample of matter
is proportional to the average kinetic energy of
the particles in the sample.
SAP-7.B.4
The Maxwell-Boltzmann distribution provides
a graphical representation of the energies/
velocities of particles at a given temperature.
SAP-7.C.1
The ideal gas law does not explain the actual
behavior of real gases. Deviations from the
ideal gas law may result from interparticle
attractions among gas molecules, particularly
at conditions that are close to those resulting in
condensation. Deviations may also arise from
particle volumes, particularly at extremely
high pressures. 
SPQ-3.A.1
Solutions, also sometimes called
homogeneous mixtures, can be solids, liquids,
or gases. In a solution, the macroscopic
properties do not vary throughout the sample.
In a heterogeneous mixture, the macroscopic
properties depend on location in the mixture.
SPQ-3.A.2
Solution composition can be expressed in a
variety of ways; molarity is the most common
method used in the laboratory.
EQN: M = nsolute /Lsolution
SPQ-3.B.1
Particulate representations of solutions
communicate the structure and properties
of solutions, by illustration of the relative
concentrations of the components in the
solution and drawings that show interactions
among the components.
COLLIGATIVE PROPERTIES WILL NOT BE
ASSESSED ON THE AP EXAM.
X CALCULATIONS OF MOLALITY, PERCENT
BY MASS, AND PERCENT BY VOLUME WILL
NOT BE ASSESSED ON THE AP EXAM.
The components of a liquid solution cannot
be separated by filtration. They can, however,
be separated using processes that take
advantage of differences in the intermolecular
interactions of the components.
a. Chromatography (paper, thin-layer, and
column) separates chemical species by
taking advantage of the differential strength
of intermolecular interactions between and
among the components of the solution
(the mobile phase) and with the surface
components of the stationary phase.
b. Distillation separates chemical species by
taking advantage of the differential strength
of intermolecular interactions between and
among the components and the effects
these interactions have on the vapor
pressures of the components in the mixture.
Substances with similar intermolecular
interactions tend to be miscible or soluble in
one another.
SAP-8.A.1
Differences in absorption or emission of
photons in different spectral regions are
related to the different types of molecular
motion or electronic transition:
a. Microwave radiation is associated with
transitions in molecular rotational levels.
b. Infrared radiation is associated with
transitions in molecular vibrational levels.
c. Ultraviolet/visible radiation is associated with
transitions in electronic energy levels.
SAP-8.B.1
When a photon is absorbed (or emitted) by an
atom or molecule, the energy of the species is
increased (or decreased) by an amount equal
to the energy of the photon.
SAP-8.B.2
The wavelength of the electromagnetic wave is
related to its frequency and the speed of light
by the equation:
EQN: c = λν.
The energy of a photon is related to the
frequency of the electromagnetic wave
through Planck’s equation (E = ℎν). 
SAP-8.C.1
The Beer-Lambert law relates the absorption of
light by a solution to three variables according
to the equation:
EQN: A = Ebc.
The molar absorptivity E describes how
intensely a sample of molecules or ions
absorbs light of a specific wavelength.
The path length b and concentration
c are proportional to the number of
absorbing species.
SAP-8.C.2
In most experiments the path length and
wavelength of light are held constant. In
such cases, the absorbance is proportional
only to the concentration of absorbing
molecules or ions.
TRA-1.A.1
A physical change occurs when a substance
undergoes a change in properties but not a
change in composition. Changes in the phase
of a substance (solid, liquid, gas) or formation/
separation of mixtures of substances are
common physical changes.
TRA-1.A.2
A chemical change occurs when substances
are transformed into new substances, typically
with different compositions. Production of
heat or light, formation of a gas, formation
of a precipitate, and/or color change provide
possible evidence that a chemical change
has occurred.
TRA-1.B.1
All physical and chemical processes can
be represented symbolically by balanced
equations.
TRA-1.B.2
Chemical equations represent chemical
changes. These changes are the result
of a rearrangement of atoms into new
combinations; thus, any representation of a
chemical change must contain equal numbers
of atoms of every element before and after the
change occurred. Equations thus demonstrate
that mass is conserved in chemical reactions.
TRA-1.B.3
Balanced molecular, complete ionic, and net
ionic equations are differing symbolic forms
used to represent a chemical reaction. The
form used to represent the reaction depends
on the context in which it is to be used.
TRA-1.C.1
Balanced chemical equations in their various
forms can be translated into symbolic
particulate representations.
TRA-1.D.1
Processes that involve the breaking and/or
formation of chemical bonds are typically
classified as chemical processes. Processes
that involve only changes in intermolecular
interactions, such as phase changes, are
typically classified as physical processes.
TRA-1.D.2
Sometimes physical processes involve the
breaking of chemical bonds. For example,
plausible arguments could be made for
the dissolution of a salt in water, as either
a physical or chemical process, involves
breaking of ionic bonds, and the formation
of ion-dipole interactions between ions
and solvent.
SPQ-4.A.1
Because atoms must be conserved during a
chemical process, it is possible to calculate
product amounts by using known reactant
amounts, or to calculate reactant amounts
given known product amounts.
SPQ-4.A.2
Coefficients of balanced chemical
equations contain information regarding the
proportionality of the amounts of substances
involved in the reaction. These values can be
used in chemical calculations involving the
mole concept.
SPQ-4.A.3
Stoichiometric calculations can be combined
with the ideal gas law and calculations involving
molarity to quantitatively study gases
and solutions.
SPQ-4.B.1
Titrations may be used to determine the
concentration of an analyte in solution. The
titrant has a known concentration of a species
that reacts specifically and quantitatively
with the analyte. The equivalence point of the
titration occurs when the analyte is totally
consumed by the reacting species in the
titrant. The equivalence point is often indicated
by a change in a property (such as color) that
occurs when the equivalence point is reached.
This observable event is called the endpoint of
the titration.
TRA-2.A.1
Acid-base reactions involve transfer of one or
more protons between chemical species.
TRA-2.A.2
Oxidation-reduction reactions involve transfer
of one or more electrons between chemical
species, as indicated by changes in oxidation
numbers of the involved species. Combustion
is an important subclass of oxidation-reduction
reactions, in which a species reacts with
oxygen gas. In the case of hydrocarbons,
carbon dioxide and water are products of
complete combustion.
TRA-2.A.3
In a redox reaction, electrons are transferred
from the species that is oxidized to the species
that is reduced.
X THE MEANING OF THE TERMS “REDUCING
AGENT” AND “OXIDIZING AGENT” WILL
NOT BE ASSESSED ON THE AP EXAM.
Rationale: Understanding this terminology is not
necessary for reasoning about redox chemistry
TRA-2.A.4
Oxidation numbers may be assigned to each of
the atoms in the reactants and products; this is
often an effective way to identify the oxidized
and reduced species in a redox reaction.
TRA-2.A.5
Precipitation reactions frequently involve
mixing ions in aqueous solution to produce an
insoluble or sparingly soluble ionic compound.
All sodium, potassium, ammonium, and nitrate
salts are soluble in water.
X ROTE MEMORIZATION OF “SOLUBILITY
RULES” OTHER THAN THOSE IMPLIED IN
TRA-2.A.5 WILL NOT BE ASSESSED ON THE
AP EXAM.
TRA-2.B.1
By definition, a Brønsted-Lowry acid is a
proton donor and a Brønsted-Lowry base is a
proton acceptor.
TRA-2.B.2
Only in aqueous solutions, water plays an
important role in many acid-base reactions,
as its molecular structure allows it to accept
protons from and donate protons to
dissolved species.
TRA-2.B.3
When an acid or base ionizes in water, the
conjugate acid-base pairs can be identified and
their relative strengths compared.
X LEWIS ACID-BASE CONCEPTS WILL NOT
BE ASSESSED ON THE AP EXAM.
Rationale: Lewis acid-base concepts are
important ideas for organic chemistry. However,
as the emphasis in AP Chemistry is on reactions
in aqueous solution, these concepts will not
be examined.
TRA-2.C.1
Balanced chemical equations for redox
reactions can be constructed from
half-reactions.
TRA-3.A.1
The kinetics of a chemical reaction is defined
as the rate at which an amount of reactants is
converted to products per unit of time.
TRA-3.A.2
The rates of change of reactant and
product concentrations are determined
by the stoichiometry in the balanced
chemical equation.
TRA-3.A.3
The rate of a reaction is influenced by reactant
concentrations, temperature, surface area,
catalysts, and other environmental factors.
TRA-3.B.1
Experimental methods can be used to monitor
the amounts of reactants and/or products
of a reaction and to determine the rate of
the reaction.
TRA-3.B.2
The rate law expresses the rate of a reaction
as proportional to the concentration of each
reactant raised to a power.
TRA-3.B.3
The power of each reactant in the rate law
is the order of the reaction with respect to
that reactant. The sum of the powers of the
reactant concentrations in the rate law is the
overall order of the reaction.
TRA-3.B.4
The proportionality constant in the rate law
is called the rate constant. The value of this
constant is temperature dependent and the
units reflect the overall reaction order.
TRA-3.B.5
Comparing initial rates of a reaction is a
method to determine the order with respect to
each reactant.
TRA-3.C.1
The order of a reaction can be inferred from a
graph of concentration of reactant versus time.
TRA-3.C.2
If a reaction is first order with respect to a
reactant being monitored, a plot of the natural
log (ln) of the reactant concentration as a
function of time will be linear.
TRA-3.C.3
If a reaction is second order with respect
to a reactant being monitored, a plot of the
reciprocal of the concentration of that reactant
versus time will be linear.
TRA-3.C.4
The slopes of the concentration versus
time data for zeroth, first, and second order
reactions can be used to determine the rate
constant for the reaction.
Zeroth order:
EQN: [A]t − [A]0 = −kt
First order:
EQN: ln[A]t − ln[A]0 = −kt
Second order:
EQN: 1/[A]t − 1/[A]0 = kt
TRA-3.C.5
Half-life is a critical parameter for first order
reactions because the half-life is constant and
related to the rate constant for the reaction by
the equation:
EQN: t
1/2 = 0.693/k.
TRA-3.C.6
Radioactive decay processes provide an
important illustration of first order kinetics.
TRA-4.A.1
The rate law of an elementary reaction can
be inferred from the stoichiometry of the
molecules participating in a collision.
TRA-4.A.2
Elementary reactions involving the
simultaneous collision of three or more
particles are rare.
TRA-4.B.1
For an elementary reaction to successfully
produce products, reactants must successfully
collide to initiate bond-breaking and bondmaking events.
TRA-4.B.2
In most reactions, only a small fraction of
the collisions leads to a reaction. Successful
collisions have both sufficient energy to
overcome energy barriers and orientations
that allow the bonds to rearrange in the
required manner.
TRA-4.B.3
The Maxwell-Boltzmann distribution curve
describes the distribution of particle energies;
this distribution can be used to gain a
qualitative estimate of the fraction of collisions
with sufficient energy to lead to a reaction, and
also how that fraction depends on temperature.
TRA-4.C.1
Elementary reactions typically involve the
breaking of some bonds and the forming of
new ones.
TRA-4.C.2
The reaction coordinate is the axis along
which the complex set of motions involved
in rearranging reactants to form products can
be plotted.
TRA-4.C.3
The energy profile gives the energy along the
reaction coordinate, which typically proceeds
from reactants, through a transition state, to
products. The energy difference between
the reactants and the transition state is the
activation energy for the forward reaction.
TRA-4.C.4
The Arrhenius equation relates the
temperature dependence of the rate of an
elementary reaction to the activation energy
needed by molecular collisions to reach the
transition state.
X CALCULATIONS INVOLVING THE
ARRHENIUS EQUATION WILL NOT BE
ASSESSED ON THE AP EXAM.
TRA-5.A.1
A reaction mechanism consists of a series
of elementary reactions, or steps, that occur
in sequence. The components may include
reactants, intermediates, products, and catalysts.
TRA-5.A.2
The elementary steps when combined should
align with the overall balanced equation of a
chemical reaction.
TRA-5.A.3
A reaction intermediate is produced by some
elementary steps and consumed by others, such
that it is present only while a reaction is occurring.
TRA-5.A.4
Experimental detection of a reaction
intermediate is a common way to build
evidence in support of one reaction
mechanism over an alternative mechanism.
X COLLECTION OF DATA PERTAINING
TO DETECTION OF A REACTION
INTERMEDIATE WILL NOT BE ASSESSED
ON THE AP EXAM.
Rationale: Designing an experiment to
identify reaction intermediates often requires
knowledge that is beyond the scope of a general
chemistry course.
TRA-5.B.1
For reaction mechanisms in which each
elementary step is irreversible, or in which the
first step is rate limiting, the rate law of the
reaction is set by the molecularity of the slowest
elementary step (i.e., the rate-limiting step).
X COLLECTION OF DATA PERTAINING
TO DETECTION OF A REACTION
INTERMEDIATE WILL NOT BE ASSESSED
ON THE AP EXAM.
Rationale: Designing an experiment to identify
reaction intermediates often requires knowledge that
is beyond the scope of a general chemistry course.
TRA-5.C.1
If the first elementary reaction is not rate limiting,
approximations (such as steady state) must be
made to determine a rate law expression.
TRA-5.D.1
Knowledge of the energetics of each
elementary reaction in a mechanism allows
for the construction of an energy profile for a
multistep reaction.
ENE-1.A.1
In order for a catalyst to increase the rate of
a reaction, the addition of the catalyst must
increase the number of effective collisions and/
or provide a reaction path with a lower activation
energy relative to the original reaction coordinate.
ENE-1.A.2
In a reaction mechanism containing a catalyst,
the net concentration of the catalyst is constant.
However, the catalyst will frequently be
consumed in the rate-determining step of the
reaction, only to be regenerated in a subsequent
step in the mechanism.
ENE-1.A.3
Some catalysts accelerate a reaction by binding to
the reactant(s). The reactants are either oriented
more favorably or react with lower activation
energy. There is often a new reaction intermediate
in which the catalyst is bound to the reactant(s).
Many enzymes function in this manner.
ENE-1.A.4
Some catalysts involve covalent bonding
between the catalyst and the reactant(s). An
example is acid-base catalysis, in which a
reactant or intermediate either gains or loses
a proton. This introduces a new reaction
intermediate and new elementary reactions
involving that intermediate.
ENE-1.A.5
In surface catalysis, a reactant or intermediate
binds to, or forms a covalent bond with,
the surface. This introduces elementary
reactions involving these new bound
reaction intermediate(s).
ENE-2.A.1
Temperature changes in a system indicate
energy changes.
ENE-2.A.2
Energy changes in a system can be described
as endothermic and exothermic processes
such as the heating or cooling of a substance,
phase changes, or chemical transformations.
ENE-2.A.3
When a chemical reaction occurs, the energy
of the system either decreases (exothermic
reaction), increases (endothermic reaction), or
remains the same. For exothermic reactions,
the energy lost by the reacting species
(system) is gained by the surroundings, as
heat transfer from or work done by the system.
Likewise, for endothermic reactions, the
system gains energy from the surroundings by
heat transfer to or work done on the system.
ENE-2.A.4
The formation of a solution may be an
exothermic or endothermic process,
depending on the relative strengths of
intermolecular/interparticle interactions before
and after the dissolution process. 
ENE-2.B.1
A physical or chemical process can be
described with an energy diagram that shows
the endothermic or exothermic nature of
that process.
ENE-2.C.1
The particles in a warmer body have a greater
average kinetic energy than those in a
cooler body.
ENE-2.C.2
Collisions between particles in thermal
contact can result in the transfer of energy.
This process is called “heat transfer,” “heat
exchange,” or “transfer of energy as heat.”
ENE-2.C.3
Eventually, thermal equilibrium is reached as
the particles continue to collide. At thermal
equilibrium, the average kinetic energy of
both bodies is the same, and hence, their
temperatures are the same.
ENE-2.D.1
The heating of a cool body by a warmer body is
an important form of energy transfer between
two systems. The amount of heat transferred
between two bodies may be quantified by the
heat transfer equation:
EQN: q = mcΔT.
Calorimetry experiments are used to measure
the transfer of heat.
ENE-2.D.2
The first law of thermodynamics states
that energy is conserved in chemical and
physical processes.
ENE-2.D.3
The transfer of a given amount of thermal
energy will not produce the same temperature
change in equal masses of matter with differing
specific heat capacities.
ENE-2.D.4
Heating a system increases the energy of the
system, while cooling a system decreases the
energy of the system.
ENE-2.D.5
The specific heat capacity of a substance
and the molar heat capacity are both used in
energy calculations. 
ENE-2.D.6
Chemical systems change their energy through
three main processes: heating/cooling, phase
transitions, and chemical reactions.
ENE-2.E.1
Energy must be transferred to a system to
cause a substance to melt (or boil). The energy
of the system therefore increases as the
system undergoes a solid-to-liquid (or liquidto-gas) phase transition. Likewise, a system
releases energy when it freezes (or condenses).
The energy of the system decreases as the
system undergoes a liquid-to-solid (or gasto-liquid) phase transition. The temperature of
a pure substance remains constant during a
phase change.
ENE-2.E.2
The energy absorbed during a phase change
is equal to the energy released during a
complementary phase change in the opposite
direction. For example, the molar heat of
condensation of a substance is equal to the
negative of its molar heat of vaporization.
ENE-2.F.1
The enthalpy change of a reaction gives the
amount of heat energy released (for negative
values) or absorbed (for positive values) by a
chemical reaction at constant pressure. 
THE TECHNICAL DISTINCTIONS BETWEEN
ENTHALPY AND INTERNAL ENERGY WILL
NOT BE ASSESSED ON THE AP EXAM.
Rationale: These distinctions are beyond the
scope of the AP Chemistry course. Most reactions
studied at the AP level are carried out at constant
pressure. Under these conditions the enthalpy
change of the process is equal to the heat
(and by extension, the energy) of reaction. For
example, in the AP Chemistry course the terms
“bond energy” and “bond enthalpy” are often
used interchangeably. 
ENE-3.A.1
During a chemical reaction, bonds are broken
and/or formed, and these events change the
potential energy of the system.
ENE-3.A.2
The average energy required to break all of
the bonds in the reactant molecules can be
estimated by adding up the average bond
energies of all the bonds in the reactant
molecules. Likewise, the average energy
released in forming the bonds in the product
molecules can be estimated. If the energy
released is greater than the energy required,
the reaction is exothermic. If the energy
required is greater than the energy released,
the reaction is endothermic.
ENE-3.B.1
Tables of standard enthalpies of formation can
be used to calculate the standard enthalpies of
reactions.
EQN: ΔH°reaction = ΣΔHf
°products − ΣΔHf
°reactants
ENE-3.C.1
Although the concept of “state function” is
not required for the course, two principles of
Hess’s law should be understood. First, when a
reaction is reversed, the enthalpy change stays
constant in magnitude but becomes reversed
in mathematical sign. Second, when two (or
more) reactions are added to obtain an overall
reaction, the individual enthalpy changes of
each reaction are added to obtain the net
enthalpy of the overall reaction.
ENE-3.D.1
When the products of a reaction are at a
different temperature than their surroundings,
they exchange energy with the surroundings
to reach thermal equilibrium. Thermal energy
is transferred to the surroundings from the
products of an exothermic reaction. Thermal
energy is transferred from the surroundings to
the products of an endothermic reaction.
TRA-6.A.1
Many observable processes are reversible.
Examples include evaporation and
condensation of water, absorption and
desorption of a gas, or dissolution and
precipitation of a salt. Some important
reversible chemical processes include the
transfer of protons in acid-base reactions and
the transfer of electrons in redox reactions.
TRA-6.A.2
When equilibrium is reached, no observable
changes occur in the system. Reactants and
products are simultaneously present, and
the concentrations or partial pressures of all
species remain constant.
TRA-6.A.3
The equilibrium state is dynamic. The
forward and reverse processes continue
to occur at equal rates, resulting in no net
observable change.
TRA-6.A.4
Graphs of concentration, partial pressure, or
rate of reaction versus time for simple chemical
reactions can be used to understand the
establishment of chemical equilibrium.
TRA-6.B.1
If the rate of the forward reaction is greater
than the reverse reaction, then there is a net
conversion of reactants to products. If the
rate of the reverse reaction is greater than
that of the forward reaction, then there is a
net conversion of products to reactants. An
equilibrium state is reached when these rates
are equal.
TRA-7.A.1
The reaction quotient Q
c c
 describes the relative
oncentrations of reaction species at any time.
For gas phase reactions, the reaction quotient
may instead be written in terms of pressures
as Qp
. The reaction quotient tends toward the
equilibrium constant such that at equilibrium
Kc = Qc
 and K
p = Qp
. As examples, for the
reaction
aA + bB  cC + dD
the equilibrium expression for (Kc
, Qc
) is
EQN: [ ] [ ]
[ ] [ ]
K = C D
A B c
c d
a b
and that for (K
p
, Qp
) is
EQN: K = (P ) (P )
(P ) (P ) p
c d
a b
C D
A B
X CONVERSION BETWEEN Kc AND Kp WILL
NOT BE ASSESSED ON THE AP EXAM.
Rationale: Conversion between K and K is an c p
algorithm that does not deepen understanding of
equilibrium. However, students should be aware
of the conceptual differences and pay attention to
whether Kc or Kp is used in an exam question.
X EQUILIBRIUM CALCULATIONS ON systems where a dissolved species in equilibrium with that species in the gas phase will not be assesed on the ap exam.
Rationale: The need to account for the relative
volumes of the liquid and gas phase is complex in
these mixed equilibria.
TRA-7.A.2
The reaction quotient does not include
substances whose concentrations (or partial
pressures) are independent of the amount,
such as for solids and pure liquids.
TRA-7.B.1
Equilibrium constants can be determined
from experimental measurements of the
concentrations or partial pressures of the
reactants and products at equilibrium.
TRA-7.C.1
Some equilibrium reactions have very large K
values and proceed essentially to completion.
Others have very small K values and barely
proceed at all. 
TRA-7.D.1
When a reaction is reversed, K is inverted.
TRA-7.D.2
When the stoichiometric coefficients of a
reaction are multiplied by a factor c, K is raised
to the power c.
TRA-7.D.3
When reactions are added together, the K of
the resulting overall reaction is the product of
the K’s for the reactions that were summed.
TRA-7.D.4
Since the expressions for K and Q have
identical mathematical forms, all valid algebraic
manipulations of K also apply to Q.
TRA-7.E.1
The concentrations or partial pressures of
species at equilibrium can be predicted given
the balanced reaction, initial concentrations,
and the appropriate K.
TRA-7.F.1
Particulate representations can be used to
describe the relative numbers of reactant
and product particles present prior to
and at equilibrium, and the value of the
equilibrium constant.
TRA-8.A.1
Le Châtelier’s principle can be used to predict
the response of a system to stresses such
as addition or removal of a chemical species,
change in temperature, change in volume/
pressure of a gas-phase system, or dilution of a
reaction system.
TRA-8.A.2
Le Châtelier’s principle can be used to
predict the effect that a stress will have on
experimentally measurable properties such as
pH, temperature, and color of a solution.
TRA-8.B.1
A disturbance to a system at equilibrium
causes Q to differ from K, thereby taking
the system out of equilibrium. The system
responds by bringing Q back into agreement
with K, thereby establishing a new
equilibrium state.
TRA-8.B.2
Some stresses, such as changes in
concentration, cause a change in Q only.
A change in temperature causes a change in K.
In either case, the concentrations or partial
pressures of species redistribute to bring Q
and K back into equality.
SPQ-5.A.1
The dissolution of a salt is a reversible process
whose extent can be described by K
sp, the
solubility-product constant.
SPQ-5.A.2
The solubility of a substance can be calculated
from the K
sp for the dissolution process. This
relationship can also be used to predict the
relative solubility of different substances.
SPQ-5.A.3
The solubility rules (see TRA-2.A.5) can be
quantitatively related to K
sp, in which K
sp
values >1 correspond to soluble salts.
SPQ-5.B.1
The solubility of a salt is reduced when it is
dissolved into a solution that already contains
one of the ions present in the salt. The impact
of this “common-ion effect” on solubility can be
understood qualitatively using Le Châtelier’s
principle or calculated from the K
sp
for the
dissolution process.
SPQ-5.B.1
The solubility of a salt is reduced when it is
dissolved into a solution that already contains
one of the ions present in the salt. The impact
of this “common-ion effect” on solubility can be
understood qualitatively using Le Châtelier’s
principle or calculated from the K
sp
for the
dissolution process.
SPQ-5.D.1
The free energy change (ΔG°) for dissolution
of a substance reflects a number of factors:
the breaking of the intermolecular interactions
that hold the solid together, the reorganization
of the solvent around the dissolved species,
and the interaction of the dissolved species
with the solvent. It is possible to estimate the
sign and relative magnitude of the enthalpic
and entropic contributions to each of these
factors. However, making predictions for the
total change in free energy of dissolution can
be challenging due to the cancellations among
the free energies associated with the three
factors cited.
SAP-9.A.1
The concentrations of hydronium ion and
hydroxide ion are often reported as pH and
pOH, respectively.
EQN: pH = −log[H3
O+
]
EQN: pOH = −log[OH−
]
The terms “hydrogen ion” and “hydronium ion” and
the symbolsH+
(aq) and H3
O+
(aq) are often used
interchangeably for the aqueous ion of hydrogen.
Hydronium ion andH3
O+
(aq) are preferred, but
H+
(aq) is also accepted on the AP Exam.
SAP-9.A.2
Water autoionizes with an equilibrium constant Kw.
EQN: Kw = [H3
O+
][OH−
] = 1.0 × 10−14 at 25°C
SAP-9.A.3
In pure water, pH = pOH is called a neutral
solution. At 25°C, pKw = 14.0 and thus
pH = pOH = 7.0.
EQN: pKw = 14 = pH + pOH at 25°C
SAP-9.A.4
The value of Kw is temperature dependent,
so the pH of pure, neutral water will deviate
from 7.0 at temperatures other than 25°C.
SAP-9.B.1
Molecules of a strong acid (e.g., HCl, HBr,
HI, HClO4
, H2
SO4
, and HNO3
) will completely
ionize in aqueous solution to produce
hydronium ions. As such, the concentration
of H3
O+ in a strong acid solution is equal to
the initial concentration of the strong acid,
and thus the pH of the strong acid solution is
easily calculated.
SAP-9.B.2
When dissolved in solution, strong bases (e.g.,
group I and II hydroxides) completely dissociate
to produce hydroxide ions. As such, the
concentration of OH− in a strong base solution
is equal to the initial concentration of the
strong base, and thus the pOH (and pH) of the
strong base solution is easily calculated.
SAP-9.C.1
Weak acids react with water to produce
hydronium ions. However, molecules of a
weak acid will only partially ionize in this way.
In other words, only a small percentage of
the molecules of a weak acid are ionized in a
solution. Thus, the concentration of H3
O+ is
much less than the initial concentration of the
molecular acid, and the vast majority of the
acid molecules remain un-ionized.
SAP-9.C.2
A solution of a weak acid involves equilibrium
between an un-ionized acid and its conjugate
base. The equilibrium constant for this reaction
is Ka
, often reported as pKa
. The pH of a weak
acid solution can be determined from the initial
acid concentration and the pKa
.
EQN: K [H O ][A ]
[HA] a
3 =
+ −
EQN: pKa = -log Ka
SAP-9.C.3
Weak bases react with water to produce
hydroxide ions in solution. However, ordinarily
just a small percentage of the molecules of
a weak base in solution will ionize in this way.
Thus, the concentration of OH-
 in the solution
does not equal the initial concentration of
the base, and the vast majority of the base
molecules remain un-ionized.
SAP-9.C.4
A solution of a weak base involves equilibrium
between an un-ionized base and its conjugate
acid. The equilibrium constant for this reaction
is Kb
, often reported as pKb
. The pH of a weak
base solution can be determined from the initial
base concentration and the pKb
.
EQN: K [OH ][HB ]
[B] b =
− +
EQN: pKb = -log Kb
SAP-9.C.5
The percent ionization of a weak acid (or base)
can be calculated from its pKa
 (pKb
) and the
initial concentration of the acid (base).
SAP-9.D.1
When a strong acid and a strong base are
mixed, they react quantitatively in a reaction
represented by the equation:
H+
(aq) + OH-
(aq) → H2
O(l).
The pH of the resulting solution may be
determined from the concentration of
excess reagent.
SAP-9.D.2
When a weak acid and a strong base are
mixed, they react quantitatively in a reaction
represented by the equation:
HA(aq) + OH-
(aq)  A-
(aq) H2
O(l).
If the weak acid is in excess, then a buffer
solution is formed, and the pH can be
determined from the Henderson-Hasselbalch
(H-H) equation (see SAP-10.C.1). If the
strong base is in excess, then the pH can
be determined from the moles of excess
hydroxide ion and the total volume of solution.
If they are equimolar, then the (slightly basic)
pH can be determined from the equilibrium
represented by the equation:
A-
(aq) + H2
O(l)  HA(aq) + OH-
(aq).
SAP-9.D.3
When a weak base and a strong acid are mixed,
they will react quantitatively in a reaction
represented by the equation:
B(aq) + H3
O+
(aq)  HB+
(aq) + H2
O(l).
If the weak base is in excess, then a buffer
solution is formed, and the pH can be
determined from the H-H equation. If the
strong acid is in excess, then the pH can
be determined from the moles of excess
hydronium ion and the total volume of solution.
If they are equimolar, then the (slightly acidic)
pH can be determined from the equilibrium
represented by the equation:
HB+
(aq) + H2
O(l)  B(aq) + H3
O+
(aq).
SAP-9.D.4
When a weak acid and a weak base are mixed,
they will react to an equilibrium state whose
reaction may be represented by the equation:
HA(aq) + B(aq)  A-
(aq) + HB+
(aq).
SAP-9.E.1
An acid-base reaction can be carried out under
controlled conditions in a titration. A titration
curve, plotting pH against the volume of titrant
added, is useful for summarizing results from
a titration.
SAP-9.E.2
At the equivalence point, the number of moles
of titrant added is equal to the number of moles
of analyte originally present. This relationship
can be used to obtain the concentration of the
analyte. This is the case for titrations of strong
acids/bases and weak acids/bases.
SAP-9.E.3
For titrations of weak acids/bases, it is useful to
consider the point halfway to the equivalence
point, that is, the half-equivalence point. At
this point, there are equal concentrations of
each species in the conjugate acid-base pair,
for example, for a weak acid [HA] = [A−
].
Because pH = pK
and bas a
 when the conjugate acid
e have equal concentrations, the pKa
can be determined from the pH at the halfequivalence point in a titration.
SAP-9.E.4
For polyprotic acids, titration curves can
be used to determine the number of acidic
protons. In doing so, the major species present
at any point along the curve can be identified,
along with the pKa associated with each proton
in a weak polyprotic acid.
X Computation of the concetration of each species present in the titration curve for polyprotic acids will not be assesed on the ap exam.
Rationale: Such computations for titration of
monoprotic acids are within the scope of the
course, as is qualitative reasoning regarding
what species are present in large versus
small concentrations at any point in titration
of a polyprotic acid. However, additional
computations of the concentration of each species
present in the titration curve for polyprotic acids
may encourage algorithmic calculations rather
than deepen understanding.
SAP-9.F.1
The protons on a molecule that will participate
in acid-base reactions, and the relative
strength of these protons, can be inferred from
the molecular structure.
a. Strong acids (such as HCl, HBr, HI,
HClO4
, H2
SO4
, and HNO3
) have very
weak conjugate bases that are stabilized
by electronegativity, inductive effects,
resonance, or some combination thereof.
b. Carboxylic acids are one common class of
weak acid.
c. Strong bases (such as group I and II
hydroxides) have very weak conjugate acids.
d. Common weak bases include nitrogenous
bases such as ammonia as well as
carboxylate ions.
e. Electronegative elements tend to stabilize
the conjugate base relative to the conjugate
acid, and so increase acid strength.
SAP-10.A.1
The protonation state of an acid or base (i.e.,
the relative concentrations of HA and A−
) can
be predicted by comparing the pH of a solution
to the pKa of the acid in that solution. When
solution pH < acid pKa
, the acid form has a
higher concentration than the base form. When
solution pH > acid pKa
, the base form has a
higher concentration than the acid form.
SAP-10.A.2
Acid-base indicators are substances that
exhibit different properties (such as color) in
their protonated versus deprotonated state,
making that property respond to the pH of
a solution.
SAP-10.B.1
A buffer solution contains a large concentration
of both members in a conjugate acid-base pair.
The conjugate acid reacts with added base
and the conjugate base reacts with added acid.
These reactions are responsible for the ability
of a buffer to stabilize pH.
SAP-10.C.1
The pH of the buffer is related to the pKa of
the acid and the concentration ratio of the
conjugate acid-base pair. This relation is a
consequence of the equilibrium expression
associated with the dissociation of a weak
acid, and is described by the HendersonHasselbalch equation. Adding small amounts
of acid or base to a buffered solution does not
significantly change the ratio of [A−
]/[HA] and
thus does not significantly change the solution
pH. The change in pH on addition of acid or
base to a buffered solution is therefore much
less than it would have been in the absence of
the buffer.
EQN: pH = pKa +
[ ]
−
log [A ]
SAP-10.D.1
Increasing the concentration of the buffer
components (while keeping the ratio of these
concentrations constant) keeps the pH of the
buffer the same but increases the capacity of
the buffer to neutralize added acid or base.
SAP-10.D.2
When a buffer has more conjugate acid than
base, it has a greater buffer capacity for
addition of added base than acid. When a
buffer has more conjugate base than acid, it
has a greater buffer capacity for addition of
added acid than base.
ENE-4.A.1
Entropy increases when matter becomes more
dispersed. For example, the phase change
from solid to liquid or from liquid to gas results
in a dispersal of matter as the individual
particles become freer to move and generally
occupy a larger volume. Similarly, for a gas, the
entropy increases when there is an increase
in volume (at constant temperature), and the
gas molecules are able to move within a larger
space. For reactions involving gas-phase
reactants or products, the entropy generally
increases when the total number of moles of
gas-phase products is greater than the total
number of moles of gas-phase reactants.
ENE-4.A.2
Entropy increases when energy is dispersed.
According to kinetic molecular theory (KMT),
the distribution of kinetic energy among
the particles of a gas broadens as the
temperature increases. As a result, the entropy
of the system increases with an increase
in temperature.
ENE-4.B.1
The entropy change for a process can be
calculated from the absolute entropies of
the species involved before and after the
process occurs.
EQN: ΔSo
reaction = ΣSo
products - ΣSo
reactants
ENE-4.C.1
The Gibbs free energy change for a chemical
process in which all the reactants and
products are present in a standard state
(as pure substances, as solutions of 1.0 M
concentration, or as gases at a pressure of
1.0 atm (or 1.0 bar)) is given the symbol ∆Go
.
The standard Gibbs free energy change for a
chemical or physical process is a measure of
thermodynamic favorability. Historically, the
term “spontaneous” has been used to describe
processes for which ΔGo < 0. The phrase
“thermodynamically favored” is preferred
instead so that common misunderstandings
(equating “spontaneous” with “suddenly”
or “without cause”) can be avoided. When G<0 for the process, it is said to thermodynamically favored.
The standard Gibbs free energy change for
a physical or chemical process may also
be determined from the standard Gibbs
free energy of formation of the reactants
and products.
EQN: ΔG°
reaction = ΣΔGf
°products - ΣΔGf
°reactants
ENE-4.C.4
In some cases, it is necessary to consider both
enthalpy and entropy to determine if a process
will be thermodynamically favored. The freezing
of water and the dissolution of sodium nitrate
are examples of such phenomena.
ENE-4.C.5
Knowing the values of ΔHo
 and ΔSo
 for a
process at a given temperature allows ΔGo to
be calculated directly.
EQN: ΔG° = ΔH° - T ΔS°
ENE-4.C.6
In general, the temperature conditions for a
process to be thermodynamically favored
(ΔGo < 0) can be predicted from the signs of H and S.
In cases where H<0 and S>0, no 
calculation of ΔGo is necessary to determine
that the process is thermodynamically
favored (ΔGo < 0). In cases where ΔHo > 0
and ΔSo < 0, no calculation of ΔGo is
necessary to determine that the process is
thermodynamically unfavored (ΔGo > 0).
ENE-4.D.1
Many processes that are thermodynamically
favored do not occur to any measurable extent,
or they occur at extremely slow rates.
ENE-4.D.2
Processes that are thermodynamically favored,
but do not proceed at a measurable rate, are
under “kinetic control.” High activation energy
is a common reason for a process to be under
kinetic control. The fact that a process does
not proceed at a noticeable rate does not mean
that the chemical system is at equilibrium. If
a process is known to be thermodynamically
favored, and yet does not occur at a
measurable rate, it is reasonable to conclude
that the process is under kinetic control.
ENE-5.A.1
The phrase “thermodynamically favored”
(ΔGo < 0) means that the products are favored
at equilibrium (K > 1).
ENE-5.A.2
The equilibrium constant is related to free
energy by the equations
EQN: K = e-ΔG°/RT
and
EQN: ΔG° = -RT ln K.
ENE-5.A.3
Connections between K and ΔG° can be made
qualitatively through estimation. When ΔG° is
near zero, the equilibrium constant will be close
to 1. When ΔG° is much larger or much smaller
than RT, the value of K deviates strongly from 1.
ENE-5.A.4
Processes with ΔG° < 0 favor products
(i.e., K > 1) and those with ΔG° > 0 favor
reactants (i.e., K < 1).
ENE-5.B.1
An external source of energy can be used
to make a thermodynamically unfavorable
process occur. Examples include:
a. Electrical energy to drive an electrolytic cell
or charge a battery.
b. Light to drive the overall conversion of carbon
dioxide to glucose in photosynthesis.
ENE-5.B.2
A desired product can be formed by coupling
a thermodynamically unfavorable reaction
that produces that product to a favorable
reaction (e.g., the conversion of ATP to
ADP in biological systems). In the coupled
system, the individual reactions share one
or more common intermediates. The sum of
the individual reactions produces an overall
reaction that achieves the desired outcome and
has ΔG° < 0.
ENE-6.A.1
Each component of an electrochemical cell
(electrodes, solutions in the half-cells, salt
bridge, voltage/current measuring device) plays
a specific role in the overall functioning of the
cell. The operational characteristics of the cell
(galvanic vs. electrolytic, direction of electron flow,
reactions occurring in each half-cell, change in
electrode mass, evolution of a gas at an electrode,
ion flow through the salt bridge) can be described
at both the macroscopic and particulate levels.
ENE-6.A.2
Galvanic, sometimes called voltaic, cells involve
a thermodynamically favored reaction, whereas
electrolytic cells involve a thermodynamically
unfavored reaction. Visual representations
of galvanic and electrolytic cells are tools of
analysis to identify where half-reactions occur
and in what direction current flows.
ENE-6.A.3
For all electrochemical cells, oxidation
occurs at the anode and reduction occurs
at the cathode.
X LABELING AN ELECTRODE AS POSITIVE OR Negative will not be assessed on the ap exam.
Rationale: Sign conventions vary depending
on the type of electrochemical cell, even though
oxidation always occurs at the anode.
ENE-6.B.1
Electrochemistry encompasses the
study of redox reactions that occur within
electrochemical cells. The reactions are
either thermodynamically favored (resulting
in a positive voltage) or thermodynamically
unfavored (resulting in a negative voltage and
requiring an externally applied potential for the
reaction to proceed).
ENE-6.B.2
The standard cell potential of electrochemical
cells can be calculated by identifying the
oxidation and reduction half-reactions and their
respective standard reduction potentials.
ENE-6.B.3
Standard gibbs free energy change is proportional to the negative of the cell proportional to the negative of the cell potential for the redox reaction from which it is constructed. Thus a cell with a positive Eo involves a thermodynamically favored
reaction, and a cell with a negative Eo
 involves a
thermodynamically unfavored reaction.
EQN: ΔGo = -nFEo
ENE-6.C.1
In a real system under nonstandard conditions,
the cell potential will vary depending on the
concentrations of the active species. The cell
potential is a driving force toward equilibrium;
the farther the reaction is from equilibrium, the
greater the magnitude of the cell potential.
ENE-6.C.2
Equilibrium arguments such as Le Châtelier’s
principle do not apply to electrochemical
systems, because the systems are not in
equilibrium.
ENE-6.C.3
The standard cell potential Eo
 corresponds to
the standard conditions of Q = 1. As the system
approaches equilibrium, the magnitude (i.e.,
absolute value) of the cell potential decreases,
reaching zero at equilibrium (when Q = K).
Deviations from standard conditions that take
the cell further from equilibrium than Q = 1 will
increase the magnitude of the cell potential
relative to Eo
. Deviations from standard
conditions that take the cell closer to equilibrium
than Q = 1 will decrease the magnitude of the
cell potential relative to Eo
. In concentration
cells, the direction of spontaneous electron flow
can be determined by considering the direction
needed to reach equilibrium. 
ENE-6.C.4
Algorithmic calculations using the Nernst
equation are insufficient to demonstrate an
understanding of electrochemical cells under
nonstandard conditions. However, students
should qualitatively understand the effects
of concentration on cell potential and use
conceptual reasoning, including the qualitative
use of the Nernst equation:
EQN: E = Eo - (RT/nF) ln Q
to solve problems. 
ENE-6.D.1
Faraday’s laws can be used to determine the
stoichiometry of the redox reaction occurring
in an electrochemical cell with respect to the
following:
a. Number of electrons transferred
b. Mass of material deposited on or removed
from an electrode
c. Current
d. Time elapsed
e. Charge of ionic species
EQN: I = q/t