Download 1 • Matter and Measurement - Chemistry with Mr. Gansle

South Pasadena  AP Chemistry
[Keep for Reference]
1  Matter and Measurement
BLUFFER’S
1. Matter
Normally exists in 3 physical states:
LiquidFixed volume, Fluid; Takes on the shape of
lower part of container; well-defined surface
Solid
Rigid Shape; very little volume change as
temperature and pressure change
Gas
Volume expands to fill the container; volume
varies according to temperature and pressure
Kinetic Molecular Theory
The idea that matter consists of molecules or
atoms that are in constant, random motion.
Kinetic Energy = Energy of motion; higher
temperature = more motion
Macroscopic – seen with the eyes.
Microscopic – seen with a microscope
Particulate or Submicroscopic – Structures
at the atomic level (what we think about)
Mixtures
Heterogeneous Mixture – A mixture where
the properties of the mixture vary
throughout. (Like an Oatmeal cookie, the
different components are visible)
Homogeneous Mixture – Also called a
solution, where the components mix at a
molecular level; different properties of the
mixtures are unnoticeable.
Purification – The separation of a mixture
into its components. (techniques: distillation,
filtration, & chromatography)
2. Elements
GUIDE
3. Physical Properties
Properties of a lone sample (ex. mass, volume,
boiling temp, melting temp, conductivity, etc.)
Density is the physical property that relates the
mass of an object to its volume
Mass
Density =
/Volume
Extensive Property – Properties, like mass and
volume, that depend on the amount of substance
Intensive Properties – Properties like color and
density; independent of the amount of substance
Temperature -- how hot a substance is;
physical properties (like density) vary with temp
Celsius 0C for freezing point of water and
100C for melting point of water.
Kelvin – same scale as Celsius; 0C = -273 K;
0 K = no motion; Celsius o + 273 = Kelvin
4. Chemical Properties
How substance interacts with other substances.
Ex. forms gas with acid; burns in air, etc.
5. Physical and Chemical Change
Physical Change – where the identity of all the
substances remain unchanged (melting, boiling,
grinding, pounding into sheets, etc.)
Chemical Change (Reaction) – atoms
rearrange to convert one substance into another
Chemical Equation – A representation of the
chemical reaction taking place
Example: P4 + 6Cl2  4PCl3
6. Measurements/Calculations
Accuracy – how close to a “true value”;
quantified by percent error.
A substance that cannot be decomposed
further by chemical means
Precision – how close measurements are to
each other. Measured by significant figures or
± notation. [I assume you know metric system.]
Names given by symbols: Example: Helium
= He, Gold = Au, Aluminum = Al
Dimensional Analysis – use of a conversion
factor to change units (ex: metric conversions,
mass & volume, time units, etc.)
South Pasadena • AP Chemistry
Name ________________________________
Period ___ Date ___/___/___
2 • Atoms and Elements
STUDY
LIST
The Development of the Atomic Theory:
Molar Mass Calculations:
I can:
 State the four “signpost scientists”, their
experiments, what they added to the atomic
theory, and the name of their model.
 Define the three theories that Dalton explained
in terms of atoms:
o Law of Conservation of Matter
o Law of Definite/Constant Proportions
o Law of Multiple Proportions
 Give examples and solve calculation problems
related to each of the three theories.
 Sketch a cathode ray tube as demonstrated in
class and state how J.J. Thomson’s
experiments led to the idea that atoms have
positive and negative parts, the negative parts
are all the same, and the negative parts (called
electrons) have a certain charge/mass ratio.
 Define cathode rays.
 State the factors that determine how much a
moving charged particle will be deflected by
an electric or magnetic field.
 Explain Millikan’s oil drop experiment & how
it added to the atomic theory.
 Sketch the set-up used by Ernest Rutherford
(the gold-foil experiment), show what he
observed, and explain how these observations
led to the idea that most of the mass of the
atom is concentrated into a tiny, amazingly
massive, positively-charged nucleus.
 Calculate the isotopic mass of an atom given
the resting mass of protons and neutrons.
 Explain that a mole of any element is actually
made up of various isotopes in a constant
percentage abundance.
 Calculate the average atomic mass of an
element using the percent abundance and mass
of each isotope.
 Calculate the percent abundance of isotopes
given the average atomic mass and isotopic
masses of an element.
Parts of the Atom:
 State the three particles that make up an atom,
their symbol, their charge, their mass, and
their location.
 State the number of protons, neutrons, and
electrons in any atom or ion.
 Explain that isotopes are two atoms with the
same atomic number (number of protons) but
different mass numbers (number of
nucleons—protons + neturons).
 Represent the nucleus with isotopic notation,
220
such as: 86 Rn
 Recognize when two nuclei are isotopes of
each other.
The Families of the Periodic Table:
 List the common families of the periodic table
and recognize to which family any element
belongs.
 Recognize metals, non-metals, and metalloids
(semi-metals) on the periodic table.
 State and define the terms conductivity,
malleability, ductility, and sectility.
 State some element facts such as which
elements are too radioactive to exist, which is
the largest non-radioactive element, which
element has the greatest density, and which
element has the highest melting point.
 Explain how Dmitri Mendeleev put together
the periodic table and why we give him credit
for the table even though others were working
along the same lines.
 List the three elements that Mendeleev
predicted and where they are located on the
periodic table.
A Little Nuclear Chemistry:
 State that Henri Becquerel discovered
radioactivity and Marie Curie studied it.
 List the three “Becquerel rays” (alpha, beta,
and gamma) and state why alpha particles
were the perfect tool for Ernest Rutherford to
study the structure of atoms.
 State that the alpha particle is the same as a
helium nucleus, a beta particle is a high-speed
electron, and a gamma ray is a high-energy
form of light.
South Pasadena  AP Chemistry
Name __________________________________
Period ___ Date ___/___/___
3  Chemical Formulas
STUDY
I can:
Formulas
 Look at a formula and state how many
elements and atoms are in that compound.
 Calculate the molecular mass or molar
mass of any compound.
 State that the mass of a molecule is
measured in amu’s and the mass of a mole
is measured in grams.
 Give examples of empirical formulas,
molecular formulas, and structural
formulas.
 Identify a formula as empirical,
molecular, or structural.
Ionic Compounds
I can:
 List the names and formulas of 60 ions.
 State whether a compound is an ionic
compound or a nonmetal compound.
 Write the formula of an ionic compound
given the two ions or its name. Know
when to use parentheses.
 Name an ionic compound given the
formula.
 Determine the charge on an ion from
information in an ionic formula.
Nonmetal Compounds
aka Molecular Compound
 Write the formula of a binary nonmetal
compound (molecular compound) given
its name.
 Name a binary nonmetal compound
(molecular compound) given its formula.
LIST
Percent Composition
 Calculate the percent composition (by
mass) for any compound.
 Calculate the empirical formula from
percent composition data.
 Determine the molecular formula of a
compound given its empirical formula
and molar mass.
Hydrates
 Give examples of hydrates and anhydrous
compounds.
 Calculate the formula of a hydrate from
dehydration data.
The Mole
 State the significance of the mole.
 State the three mole facts for any
substance (molar volume, molar mass,
Avogadro’s number)
1 mole = 22.4 Liters @ STP (gases only)
1 mole = 6.02 x 1023 particles
(particles = molecules or atoms)
1 mole = gram molecular mass of chemical
 Use dimensional analysis to convert
between moles, mass, volume, and number
of particles for a chemical.
 Use density as a conversion factor in mole
problems.
 Use gas density to calculate molar mass.
South Pasadena  AP Chemistry
Name __________________________________
Period ___ Date ___/___/___
4  Chemical Equations and Stoichiometry
STUDY
I can:
Chemical Equations
Limiting Reactant Problems



Give examples of products and reactants in a
chemical equation.
State that Antoine Lavoisier introduced the
law of conservation of matter.
Combustion



State that combustion is another name for
burning.
Write an equation for a combustion reaction
given only the fuel that is burned.
Correctly label substances in an equation as
solid (s) , gas (g), liquid (l), or aqueous (aq)
Balancing Equations



Balance equations by adding coefficients.
Recognize when an equation is balanced.
State that the formulas of reactants and
products should not be changed in order to
balance equations.
Stoichiometry Problems



Use the stoichiometric factor ( of the
problem) to convert from moles of one
substance to moles of a different substance.
(i.e. In the equation: N2 + 3H2  2NH3,
3 mol H2  2 mol NH3)
Convert between the quantities of mass,
volume, molecules and moles using
dimensional analysis
(i.e. use 1 mol = 22.4 L, 1 mol = 6.02 x 1023
molecules, and 1 mol = gram molecular mass)
Show the units of molar mass as grams/mol or
g·mol-1.



LIST
Recognize that a problem with two “given
values” is a limiting reactant problem.
Determine the limiting reactant and excess
reactant in a problem.
Solve problems involving Limiting Reactants
Calculate how much excess chemical is left
over after a reaction.
Percent Yield Problems



Use stoichiometry to calculate the theoretical
yield (mass of a product) in a problem.
State that actual yields are usually given in a
problem.
Use the theoretical yield and actual yield to
calculate the percent yield.
Chemical Analysis Problems



Calculate the mass of each element in a given
compound given data such as the masses of
CO2 and H2O formed in a combustion
reaction.
Use mass and mole information to calculate
the empirical formula of an unknown
substance.
Use percent composition to equalize mass
and mole information derived from different
samples.
Continuous Variation Data

Use “continuous variation” laboratory data to
determine the correct mole ratio of an
equation. (Micro Mole Rockets Lab)
South Pasadena  AP Chemistry
5  Reactions in Aqueous Solution
Properties of Aqueous Solutions
 Define solute, solvent, and solution. Give
examples.
 Define electrolytes. Give operational and
theoretical definitions of electrolytes.
 Know that soluble ionic compounds and
strong acids are strong electrolytes. Ionic
compounds of low solubility [e.g.
Mg(OH)2] and weak acids/bases are weak
electrolytes.
 Know that molecular compounds (except
acids) are non-electrolytes.
 Know that alcohols (e.g. CH3OH )are not
ionic hydroxides. Bases are usually
metallic hydroxides.
 Know the solubility rules. State whether
an ionic compound is soluble in water.
Precipitation Reactions
 Know that ppt reactions are double
replacement reactions that produce an
insoluble product.
 Given two ionic compounds in solution,
correctly determine the products. (Know
your ions).
 Determine which products are precipitates.
Use (aq) and (s) symbols correctly.
 Correctly write the ions in a soluble ionic
compound. [e.g. CaCl2(aq)  Ca2+ + 2Cl]
 Identify spectator ions.
 Write molecular, detailed ionic, and net
ionic equations for a ppt reaction.
Acids and Bases
 Give operational (cabbage juice) and
theoretical (ions) definitions of acids and
bases.
 Know that acids increase the H+ ion
concentration in an aqueous solution.
(Theoretical definition)
 Memorize the 8 strong acids.
 Know that acids are molecular
compounds that form ions when in
aqueous solution.
 Be able to name acids according to their
anion.



S T U D Y L I S T
[ide  hydro__ic acid; ate  __ic acid; ite
 __ous acid; S: add “ur”; P: add “or”]
Know that bases increase the OH ion
concentration in an aqueous solution.
(Theoretical definition)
Memorize the soluble hydroxides (except
NH4OH) that are the strong bases.
Understand that ammonia(aq),
H2O




NH3 +
NH4+ + OH
forms a weak basic
solution.
Know that metal oxides form bases [CaO
+ H2O  Ca(OH)2] while nonmetal oxides
form acids [CO2 + H2O  H2CO3]
Know that acids react with bases to form
H2O and a salt. (Neutralization)
Write equations for acid-base reactions
including NH3 (example on page 199) as
the base.
Know that strong acids and strong bases
are written as ions in the ionic equations.
Gas Forming Reactions
 Recognize the six products that turn into
gases. Memorize the gases formed.
Organizing Reactions in Aqueous Solution
 Double Replacement reactions (text calls
them exchange reactions) (FredWilma/Barney-Betty reactions) also have
the old fashioned name: metathesis
reactions.
 Know the three examples of double
replacement reactions and the “driving
force” for each.
Precipitate reactions form an insoluble
product. Acid-Base reactions form water (a
very weak electrolyte therefore, a very
stable product). Gas-forming reactions
form a gas.
 Know that a driving force is something
that keeps the new combinations of ions
from reforming the old combinations of
ions.
 Oxidation-Reduction is a fourth type of
reaction driven by the transfer of
electrons.
Oxidation-Reduction Reactions
Fe2O3 + 3 CO  2 Fe(s) + 3 CO2
 Know that an important type of reaction
gets its name from atoms that combine
with oxygen. During the refining of iron,
carbon monoxide combines with oxygen
(from the iron ore), CO  CO2 and is
oxidized. Large masses of iron ore (Fe2O3)
are reduced to a smaller amount of iron
metal.
 Understand that since CO helps the iron ore
to be reduced, CO is called the reducing
agent. Since Fe2O3 causes the C to be
oxidized, iron ore is called the oxidizing
agent. What ever is oxidized acts as the
reducing agent. What ever is reduced
acts as the oxidizing agent.
 Know that oxidation-reduction (redox) is
driven by the transfer of electrons.
Mnemonics to help: GROL (Gain=Reduce
/ Oxidize=Lose); LeO the lion says GeR
(Losing e’s = Oxidation / Gaining e’s =
Reduction); OIL RIG (Oxidation is Losing
e’s / Reduction is Gaining e’s)
 A redox reaction can be divided into two
half-reactions. The oxidation halfreaction has electrons as a product. The
reduction half-reaction has electrons as a
reactant.
 Be able to assign oxidation numbers to
any atom in any substance. Learn the rules
on page 207.
 Recognize redox reactions because
oxidation numbers change. (#  =
oxidation / #  = reduction), electrons are
gained or lost, or oxygen atoms are gained
or lost.
 Know several common oxidizing agents
and reducing agents and what they turn
into.
Measuring Concentrations of Cmpds. in Solution
 Know the definition of molarity, M, as one
way to communicate concentration of
solute.
 Know that the symbol [X] means the
concentration of X in moles/Liter.




Be able to determine the concentration of
ions in an ionic compound.
For example, in 0.25 M AlCl3
[AlCl3] = 0.25 M [Al3+] = 0.25 M
[Cl] = 0.75 M
Use the molarity formula to calculate
moles, mass, volume, or molarity of a
solution.
Know that Volume x Molarity = moles of
solute. Dilution problems use ViMi =
VfMf.
Describe how to make a solution correctly.
Know what a volumetric flask is.
Stoichiometry of Reactions in Aqueous Solution
 Use molarity as another conversion factor
to solve stoichiometry problems.
 Know that titration is a technique called
quantitative chemical analysis because
you are measuring. It is also called
volumetric analysis (because you are
measuring volumes).
[Note: qualitative analysis involves no
measurements such as using solubility rules
to determine the identity of an unknown
ionic compound.]
 Understand the terms indicator,
equivalence point, standardization, and
primary standard.
[Note: you saw a titration being done in
the Measurement video early in the
summer. Chloride ion from the
Chesapeake Bay was being titrated against
silver nitrate to determine the salinity
(saltiness) of the water. Yellow K2CrO4
was used as an indicator because it formed
the reddish-brown ppt, Ag2CrO4 (which
looked pink) when all the chloride ion was
used up.]
 Know common indicators such as
phenolphthalein for titrations with strong
bases.
 Understand that a titration can be done with
an acid-base reaction or a redox reaction.
In each case, some sort of indicator must be
used to tell when equivalent amounts of
reactants have been mixed.
South Pasadena • AP Chemistry
[Keep for Reference]
6 • Energy and Chemical Reactions
Driving Forces
I can…
 state that product-favored (spontaneous) reactions
tend toward maximum entropy, S, and
minimum enthalpy, H.
 state the sign of H based on observation of
warming or cooling of the surroundings.
 correctly apply the terms exothermic and
endothermic to situations where the surroundings
are warming or cooling.
 draw a PE curve (uphill or downhill) based on
information about warming or cooling of the
surroundings.
Measuring Heat
 state the units of heat capacity, specific heat, and
molar heat capacity as well as the significance of
each.
 convert between the heat units of calories and
Joules. (4.184 J = 1 calorie)
 use calorimetry (q=mCT) to calculate heat
changes during temperature changes.
 calculate the heat transferred when two objects, at
different temperatures, come into contact.
Energy = Heat and Work
 state the difference between work and heat
energy.
 state the difference between system and
surroundings.
 recognize the system and the surroundings in a
chemical or physical system.
 calculate the change in internal energy based on
changes in heat absorbed by the system and work
done by the system.
 state that H is a more general (and useful)
measure of energy than E and that H = q when
a reaction occurs at constant pressure.
S T U D Y L I S T
Chemical Work = Expanding Gases
 relate physical work (w=F·d) and chemical work
(w=P·V).
 calculate PV work done by an expanding gas.
 state that no work is done in a constant volume
situation such as a bomb calorimeter.
Calculating H -- Hess’s Law
 state the definition of a state function.
 list examples of properties that are and are not
state functions.
 write the equation for the heat of formation of a
substance.
 state that the heat of formation of an element
under standard conditions has a value of zero.
 use Hess’s Law to calculate the energy of a
chemical or physical change.
Calculating Heat During Phase
Changes – Heats of Fusion and
Vaporization
 use heats of vaporization or heats of fusion to
calculate heat changes during phase changes.
 write an equation showing the heat of fusion or
heat of vaporization.
South Pasadena • AP Chemistry
[Keep for Reference]
12 • Gases and Their Properties
£
STUDY LIST
Know the pressure of the atmosphere at sea level measured in atm, kPa, mmHg, torr, psi
Convert one pressure unit into another
Understand how to measure pressure using a U-tube manometer, open-end manometer, and a
barometer
Convert gauge pressure into total pressure
£
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Sketch a P vs. V graph
Manipulate P V data so a straight- line graph is obtained
Form a mathematical law from straight- line graph
State Boyle’s Law
Recognize situations of Boyle’s Law
Do Boyle’s Law problems
£
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Sketch a V vs. T graph
Graphically determine a value for absolute zero
State Charles’s Law
Explain why temperatures must be in K
Do Charles’s Law problems
Recognize situations of Charles’s Law
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Know the General Gas Law (P,V&T)
Know Avogadro’s Law (V&n)
Know the Ideal Gas Law
Given the molar volume of a gas (22.414 L at STP) determine values of R, the ideal gas constant
Do Ideal Gas Law problems
£
£
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£ Derive the gas density equation from the Ideal Gas Law
£ Do gas density problems
£ Calculate molar mass from P, V, and T data
£ Do Gas Laws and Stoichiometry problems
£ Know Dalton’s Law of Partial Pressures
£ Do Partial Pressure problems
£ Apply this to gases collected over water
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Know the principal features of the Kinetic Molecular Theory of gases
Be able to explain why each of the gas laws works in terms of the Kinetic Molecular Theory
Understand the significance of the Maxwell-Boltzmann distribution curve s on pages 566-567
Derive Graham’s Law of Effusion from rms or KE of two gases
Do Graham’s Law problems
£ Compare van der Waal’s equations for Real gases with the Ideal Gas Law
£ Know the correction factors that appear in the Real Gas Law
South Pasadena • AP Chemistry
[Keep for Reference]
15 • Chemical Kinetics: Rates of Reaction
• How to talk about Rate
rate = ∆[chemical]/∆time
rate of disappearance of reactant or
rate of appearance of product
use coefficients to change one rate to another
watch your signs (∆[React.] = -∆[Prod.])
instantaneous rate is slope of [R] vs. time
graph. Initial rate is often used.
• How to Speed Up a Reaction
[Use Collision Theory, Kinetic Molecular Theory]
increase the concentration of reactants
- increase molarity of solutions
- increase partial pressure of gases
[collision model: more collisions]
more surface area between unlike phases
[collision model: more collisions]
increase the temperature
[collision model: more & harder collisions]
add a catalyst
- homogeneous catalyst (used & reformed)
- heterogeneous catalyst (surface catalyst)
[collision model: alternate mechanism that
requires lower energy collision or
ensures that correct particles collide]
• Two Important Diagrams
PE energy profile of a reaction
+50
c
+35
PE
a
0
d
b
–35
–50
e
reaction coordinate
? H of the reaction relates reactant and product
PE’s / exo- or endothermic/ downhill, -∆H, or
uphill, +∆H
activation energy (Ea) = energy barrier
• activated complex (at the peak)
• whether a reaction is fast or slow depends on
the activation energy in the PE profile
• PE profile does not change with change in
temperature of the reactants?
• adding a catalyst lowers the Ea
A BLUFFER’S GUIDE
The KE distribution of a substance
threshold energy
KE
- temperature is the average KE
-increasing temperature spreads out curve to the right,
increases average KE
threshold energy
KE
- adding a catalyst moves the threshold energy to the
left.
threshold energy
KE
How do these two picture relate to each other (turn
the KE on its side... the particles use their KE to
provide the needed PE to react)
• Reaction mechanisms
- step-by-step...two particles at a time
- example
overall:
4 HBr + O2 → 2 Br2 + 2 H2O
mechanism:
HBr + O2 → HOOBr
HOOBr + HBr → 2 HOBr
HOBr + HBr → Br2 + H2O
HOBr + HBr → Br2 + H2O
[note: HOOBr and HOBr are not in the overall
reaction because they are neither reactants nor
products, they are “reactive intermediates”]
- overall reaction is sum of steps
- slowest step is rate-determining step
Rate Laws
- what they mean
- how to determine them
- how they relate to the rate determining step
- how they help you choose a mechanism
General Form:
Equation: A + B → C
Rate = k [A]x[B]y
k is the “specific rate constant”
Use experimental data to determine x, y, and k.
The Rate Law CANNOT be determined from the
overall reaction. It MUST be determined
experimentally because the rate law reflects only the
“rate determining step.”
Rate law can be determined from the initial rates.
See Example 15.3 and Exercise 15.3
Rate Law matches the Mechanism
Examples for: 2A + 3B → C (fill in from lecture)
Rate Determining Step
Rate Law
in the mechanism
Rate = k [A][B]
Rate = k [A]2
2
Rate = k [A] [B]
Rate = k
order of rxn
- first and second order reactions
- what these look like graphically
- how you can graphically tell the order of a
reaction
order straight-line plot
Slope
0
[R]t vs. t
-k
1
ln[R]t vs. t
-k
2
1/[R]t vs. t
k
- how this relates to the rate law
half-life
- relationship to radioactivity
(a first order reaction)
- the equation
[ A]o
ln
= kt
[ A]t
- the special case of half-life
ln(2) = 0.693 = kt½
chain reactions (fill in from lecture & video)
- initiation steps
- propagation steps
A + B → X (slow)
- termination steps
A + A → X (slow)
examples:
- H2 + Cl2 → 2 HCl
- polymerization reactions (addition)
- ozone depletion
A + A D X (fast)
B + X → Y (slow)
Each step is usually bimolecular.
A third order overall reaction
often comes from a fast
equilibrium before a slow step.
This could be a mechanism that
depends on a catalyst
only. The concentrations would
not matter.
ozone layer
- specifics on why CFC’s are dangerous to the
ozone layer and are economically desirable here
on the surface
Determining Ea from calculations using the Arrhenius
Equation
South Pasadena • AP Chemistry
[Keep for Reference]
16 • Chemical Equilibria
BLUFFER’S GUIDE
1. aA +bB + . . .
rR +sS + . . .
r
s
[R] [S] ⋅ ⋅ ⋅
Kc =
[A] a [B] b ⋅ ⋅ ⋅
and for gases:
(PR ) r (PS ) s
Kp =
(PA ) a (PB ) b
2. K > 1 products favored
K < 1 reactants favored
3. Excluded: solids; pure liquids; water (in
aqueous solutions) because their [ ]’s do
not change.
4. Convert from Kc to Kp
Kp = Kc(RT)∆n
where ∆n = moles of gaseous product –
moles of gaseous reactant.
5. Typical question: Given Kc and the
starting concentrations of reactants, find
concentrations of products at equilibrium.
10. If out of equilibrium: Calculate the
reaction quotient (Q) similar to the
way an equilibrium constant would be
found. If:
Q < K forward reaction occurs to
reach equilibrium
Q > K reverse reaction occurs to
reach equilibrium
11. Problem solving:
• Set up problems using the “magic box”
(or ICE box) C = “change” or ∆.
Example: A
B+C
A
B
C
initial
5.0 M
0M
0M
∆
equilibrium
“∆” row only follows the stoichiometry of
the equation.
•
Example: Kc for acetic acid = 1.8 x 10-5.
What is the equilibrium concentration of
[H+] in a 0.100 M solution of the acid?
6. Equilibrium constant for a reverse
reaction = 1 K the value of the forward
reaction.
Learn when to make an approximation
(needed for multiple choice questions!)
5% rule usually works when value of K is
103 smaller than value of known
concentrations.
Example: A
B+C
K = 3.0 x 10-6
if [A] = 5.0M initially; find [C] at
equilibrium.
•
7. Equilibrium constant for a doubled
reaction = K2 .
If greater than 5% use the quadratic
equation: (not usual on the AP exam)
ax2 + bx + c = 0
x=
− b ± b 2 − 4ac
2a
8. When using Hess’s Law:
Koverall = K1 x K 2
•
9. Le Châtelier’s Principle: effect of changes
in concentration, pressure, & temperature.
Equilibrium always “shifts” away from
what you add. “Stress” means too much
or too little: chemical, heat, or room.
Based on a handout by William Bond, Snohomish HS
Another easy to solve situation is the
perfect squares situation.
Example: H2 + I2 2HI K = 3.5 x 102
Calculate [HI] when [H2 ] = [I2 ] = 0.10 M
South Pasadena • AP Chemistry
[Keep for Reference]
17 • Acid-Base Equilibria
BLUFFER’S GUIDE
1. H2 O H+ + OH− Kw = [H+][OH−] = 10−14
pH = -log[H+ ] pH+pOH = 14 [H+] =10−pH
Convert between pH, pOH, [H +], & [OH−]
2. Acid Ionization Constant (K a):
HA + H2O
H3O+ + AKa = [A- ][H 3O+]/[HA]
H3O+ + F-
Example: HF + H2O
Ka = [F- ][H3O+]/[HF]
3. Typical question: Given Ka and the starting
concentrations of acid, find concentrations
(or pH) of [H+] at equilibrium.
Example: Ka for acetic acid = 1.8 x 10-5.
Find the pH of 0.100M acetic acid.
4. Polyprotic Acids: H3PO4, H2SO4, H2C 2O4,
etc. The 1st dissociation is strong for H2 SO4 .
When using Hess’s Law with a polyprotic
acid: Koverall = Ka1 x Ka2
Calculating pH, use Ka1
5. Bronsted-Lowry Definitions.
Acids = H+ donors; Bases = H+ acceptors
Conjugate acid-base pairs.
6. Base Ionization Constant (K b):
B + H2O
BH+ + OHKb = [BH+ ][OH- ]/[B]
Example: F- + H2O
HF + OHKb = [HF][OH- ]/[F- ]
7. Salt solns can have pH’s ≠ 7 (hydrolysis)
ions from weak acids → basic solutions
C2 H3O2 − + H2O HC2 H3 O2 + OH−
ions from weak bases → acidic solutions
NH4 + + H2 O NH4 OH + H+
8. Ka x Kb = Kw = 10-14
only applies for conjugate acids & bases!
Example: Ka HC2 H3 O2 = 1.8 x 10-5
Kb C2 H3O2- = 10-14 / 1.8 x 10-5
9. Percent ionization =
[H+]eqilibrium /[HA]initial x 100
10. Acid Strength-know the 6 strong acids: HCl,
HBr, HI, HNO3, HClO 4, and H2SO4 (removal
of the first H+ only)
(a) binary acids - acid strength increases with
increasing size and electronegativity of the
“other element”. ( NOTE: Size predominates
over electronegativity in determining acid
strength.)
Examples: H2 Te > H2O & HF > NH3
(b) Oxoacids - Acid strength increases with
increasing:
(1) electronegativity
(2) number of bonded oxygen atoms
(3) oxidation state of the “central atom”.
Example: HClO 4 or [O 3Cl(OH)]
is very acidic
NaOH is very basic
Acid strength also increases with decreasing
radii of the “central atom”.
Example:
HOCl (bond between Cl and OH is
covalent--making HOCl acidic)
HOI (bond between I and OH is ionic-making HOI basic)
11. Lewis Acids and Bases:
(This applies to coordinate covalent bonds.)
Lewis Acid--electron pair acceptor
Lewis Base--electron pair donor
“Have Pair…Will Share” – Lewis Base
In complex ion formation, metal ions are Lewis
acids, and ligands are Lewis bases.
Example: Cu2+ + 4NH3
Cu(NH3) 42+
Cu2+ acts as an acid; NH3 acts as a base.
12. Strong Bases: amide ion, NH2 −
hydride ion, H−, methoxide ion, CH3 O−
Based on a handout by William Bond, Snohomish HS
South Pasadena  AP Chemistry
[Keep for Reference]
18  Acid-Base Reactions
BLUFFER’S
1. Buffers:
A solution that resists a change in pH when
small amounts of acid or base are added.
GUIDE
Similarly, for bases with conjugate acids:
B + H2O
HB+ + OH-
[OH  ][HB  ]
[B]
[B]
[OH-] = Kb
[HB  ]
[B]
pOH = pKb - log
[HB  ]
Kb =
Buffers are a mixture of
a weak acid & its conjugate base or
a weak base & its conjugate acid.
Examples:
HC2H3O2 & C2H3O2or NH3 & NH4+
Your blood is a buffer.
The equilibrium is the acid equilibrium:
H3O+ + Ain which both the acid and its conjugate base
are available to counteract the stress of
adding acid or base (Le Châtelier’s). The
equilibrium shifts, but is almost completely
counteracted by the proton donor or acceptor.
HA + H2O
Similarly, for bases with conjugate acids:
B + H2O
HB+ + OH-
The best buffer contains large, equal amounts
of the proton donor and the proton acceptor.
[HA] = [A-] (Note: they cancel out of Ka)
Buffers can also be formed by changing a weak
acid into its conjugate base by neutralizing
some of the acid..
HA + OH-  H2O + AThe same can be done with a weak base:
B + H+  HB+
So, a weak acid and some strong base can form
a buffer. A weak base and some strong acid can
also form a buffer.
2. Titration:
A carefully measured neutralization.
Acid + Base  H2O + Salt
Since volumes are measured, this is a
“volumetric analysis.”
Consider:
[H  ][A - ]
[HA]
+
The best buffer, Ka = [H ]; pH = pKa.
The pH of a buffer can be adjusted by
changing the ratio of acid and base.
HA
H+ + A-
Ka =
[HA]
[H+] = Ka
[A - ]
[HA]
pH = pKa - log
[A - ]
==============================
HA + OH-  H2O + AThe salt is the conjugate base of the weak acid
or the conjugate acid of the weak base. A “halftitration” (neutralizing half the weak acid or
weak base) forms a buffer. In that case the pH
= pKa of the acid. (pOH = pKb)
3. Equivalence Point:
The point in a titration when stoichiometric
amounts of acid and base have reacted.
Note that the salt solution that is formed may
have a pH >, <, or = 7. (Remember hydrolysis.)
Indicators for a titration are selected based on
the pH at the equivalence point.
acid
STRONG
STRONG
weak
weak
base
STRONG
weak
STRONG
weak
pH at eq. pt.
pH = 7
pH < 7
pH > 7
it depends
4. Titration Curves:
This graph shows how the pH changes as a
titration occurs.
(A) Strong acid/Strong Base
HCl + NaOH
H2O + NaCl
NOTE: Graph should have “pH” as the vertical
axis and “added base” as the horizontal axis.
The graph should be in an “S” shape. The
middle of the “S” is the equivalence point
(pH = 7). The top part of the “S” levels off at
the pH of the base solution.
(B) Weak acid/Strong Base
HA + OHA- + H2O
NOTE: The middle of the lower part of the “S”
indicates the point of maximum buffering
where [HA]/[A-] = 1. The middle of the “S”
is the equivalence point (above pH = 7) and
[HA] = 0. The top part of the “S” levels off at
the pH of the base solution.
(C) Weak base/Strong acid
B + H3O+
BH+ + H2O
NOTE: Graph should have “pH” as the vertical axis
and “added acid” as the horizontal axis. The
graph should be in a “backwards S” shape. The
middle of the upper part of the “backwards S”
indicates the point of maximum buffering where
[B]/[HB+] = 1. The middle of the “backwards
S” is the equivalence point (below pH = 7) and
[B] = 0. The bottom part of the “backwards S”
levels off at the pH of the acid solution.
(D) Weak diprotic acid/Strong base
H2A + OHHA- + H2O
HA- + OHA2- + H2O
NOTE: Graph should have “pH” as the vertical axis
and “added base” as the horizontal axis. The
graph should be in a “double S” shape. The
middle of the lower part of the “first S”
indicates the point of maximum buffering of the
first buffering zone where [H2A]/[HA-] = 1.
The middle of the “first S” is the first
equivalence point where [H2A] = 0. The top
part of the “first S” (i.e. the lower part of the
“second S”) indicates the point of maximum
buffering of the second buffering zone where
[HA-]/[A2-] = 1. The middle of the “second S”
is the second equivalence point where [HA-] =
0. The top part of the “second S” levels off at
the pH of the base solution.
From a handout by William Bond, Snohomish HS
South Pasadena  AP Chemistry
[Keep for Reference]
18  Acid-Base Reactions
STUDY
LIST
I can:












describe how a pH buffer behaves when small amounts of acid or base are added.
explain why a buffer works (buffering capacity) based on the presence of the weak acid (H+ donor) and
conjugate base (H+ acceptor). I can show mathematically that diluting the buffer does not change the pH
of the buffer; but it reduces its buffering capacity.
calculate the pH of the best buffer you can make from a given acid and its conjugate base given Ka’s of
weak acids (or Kb’s of weak bases)
choose the acid / conjugate base needed to get a buffer of specified pH. (Given Ka’s of acids.)
choose pairs of substances that will make a buffer:
--weak acid & its conjugate base
or
--weak base & some strong acid
--weak acid & some strong base
calculate the pH of a buffer using the ICE box or the Henderson-Hasselbach equation.
solve titration equivalence point problems using VH+ MH+ = VOH- MOHexplain that at the endpoint of a weak acid titration the solution only contains the conjugate base of the
acid. I can calculate the concentration of the conjugate base and the pH at the endpoint of a titration.
explain that weak acids and strong acids require the same amount of base to be neutralized because the
weak acids will dissociate during neutralization.
determine the equivalence point (end point) of the titration by looking at a titration curve.
determine the pKa of the weak acid being titrated by looking at a titration curve.
do the eight calculations that will allow me to sketch the pH curve for a weak acid or weak base.





--weak base & its conjugate acid
pH of the weak acid solution initially
amount of based needed for titration
concentration of conjugate base at endpoint
pH of the solution at the endpoint




pH halfway to the equivalence point (e.p.)
pH a little before halfway to the e.p.
pH a little after halfway to the e.p.
pH after all of the acid has been neutralized
translate all of my knowledge and skills from a weak acid titration to a weak base titration.
South Pasadena • AP Chemistry
[Keep for Reference]
19 • Precipitation Reactions
1. Solubility Rules
Review/memorize these rules. They can
be split into four groups:
ALWAYS SOLUBLE:
alkali metal ions (Na+, K+, Li+, Rb+, Cs+ ),
NH4 +, NO3-, C2 H3 O2 -, ClO 3-, ClO 4 USUALLY SOLUBLE:
chlorides, bromides, iodides (Cl- , Br-, I-)
except “AP/H” (Ag+, Pb2+, Hg2 2+)
sulfates (SO4 2-) except “CBS/PBS” (Ca2+,
Ba2+, Sr2+, Pb2+)
fluorides (F-) except “CBS/PM” (Ca2+,
Ba2+, Sr2+, Pb2+, Mg2+)
USUALLY INSOLUBLE:
oxides/hydroxides (O 2-, OH-) except
“CBS” ((Ca2+, Ba2+, Sr2+)
NEVER SOLUBLE:
CO32-, PO4 3-, S2-, SO 32-, CrO 42-, C2 O4 2except alkali metals & NH4 +
2.
Solubility Product (Ksp)
This type of equilibrium involves solids
of low solubility. A saturated solution is
a solution at equilibrium. The constant
has no denominator.
Example: Co(OH)2 (s)
Co2+ + 2OH2+
2
Ksp = [Co ][OH-] = 2.5 x 10-16
What is the pH of a saturated solution?
Let x = the amount (moles) of solid that
will just saturate 1 L of solution.
Co(OH)2 (s)
x
-x
0
Co2+ + 2OH0
0
+x
+2x
x
2x
(x) (2x) 2 = 4x3 = 2.5 x 10-16
x = 3.97 x 10-6 [OH-] = 2x = 7.94 x 10-6
pOH = 5.1 pH = 14- pOH = 8.9
BLUFFER’S GUIDE
3. Solubility vs. Ksp
“Molar solubility” is the concentration of
the saturated solution in moles/Liter.
(Solubility is sometimes reported in g/100
mL of water.)
As in the example, for a 1:2 compound,
Ksp = 4x3 (where x = solubility)
1:1
Ksp = x2
1:2
Ksp = 4x3
1:3
Ksp = 27x4
2:3
Ksp = 108x5
4. Will a Precipitate Form?
Ion Product (Q sp ) = “reaction quotient”.
Qsp < Ksp
more solid will dissolve
Qsp = Ksp
solution is saturated
Qsp > Ksp
ppt will form until Qsp = Ksp
Note: Be sure to calculate concentration of
DILUTED ions.
Example:
50. mL of 2.0 x 10-4 M Co(NO3 )2 is
mixed with 200 mL of 1.0 x 10-3 M
NaOH. Will a precipitate form?
[Note:K sp given in other example problem.]
50
= 4.0 x 10-5 M
250
200
[OH-] = 1.0 x 10-3 M x
= 8.0 x 10-4 M
250
Qsp = (4 x 10-5 ) (8 x 10-4 )2 = 2.56 x 10-11
Qsp > Ksp ; a precipitate will form!
[Co2+] = 2.0 x 10-4 M x
5. Solubility can be influenced by pH.
If the anion came from a weak acid, the
salt will be more soluble in a solution of
strong acid.
Example: CaCO3 (s)
Ca2+ + CO3 2-
In a strong acid, H+ combines with CO32to re- form the weak acid, H2 CO3 (which
may decompose into CO2 & H2 O). More
CaCO3 (s) will dissolve to reach
equilibrium.
South Pasadena  AP Chemistry
[Keep for Reference]
20  Entropy and Free Energy
BLUFFER’S
1. There are two driving forces for reactions.
Reactions tend toward:
minimum Enthalpy, H (heat energy)
H , H<0, downhill
maximum Entropy, S (randomness)
S +, S>0, uphill
2. Recognize whether S >0 or < 0.
Entropy increases, S +, S > 0:
 from solid to liquid to gas
 fewer moles (g) to more moles (g)
 simpler molecules to more complex
molecules
 smaller molecules to longer molecules
 ionic solids with strong attractions to
ionic solids with weaker attractions
 separate solute & solvent to solutions
 gas dissolved in water to escaped gas
3. Product or Reactant favored reactions
depend on H, S, and absolute Temp
Product-Favored…
H S
+


+
+

+

at higher temperatures
at lower temperatures
at all temperatures
never
(reactant-favored at all temps)
GUIDE
4. Many books use the term “spontaneous”
for “product-favored.”
A spontaneous reaction does not
necessarily mean a fast reaction.
The SPEED of a reaction is Kinetics
(Ch 15)… we are discussing whether a
reaction CAN OCCUR which is
Thermodynamics (Ch 6 and Ch 20).
5. Gibbs Free Energy, G, puts the effects
of H, S, and Temperature together.
G = H - TS
G<0, G , product-favored reaction
G>0, G +, reactant-favored reaction
G=0, reaction is at equilibrium
Important:
Note that H is usually in kJ/mol
S is usually in J/mol·K
6. Convert between K, G, and E
using equations given on the AP Exam.
South Pasadena • AP Chemistry
[Keep for Reference]
21 • Electrochemistry
BLUFFER’S GUIDE
1. Electrochemistry is all oxidationreduction chemistry.
Leo Ger
OIL RIG
Oxidation: loss of e− ; ox # increases
Reduction: gain of e−; ox # decreases
example: Fe2+ + 2e− → Fe(s) (reduction)
2. In a reaction, the
oxidizing agent gets reduced; the
reducing agent gets oxidized.
3. Balancing redox reactions:
oxidation number method
§ assign ox #’s to every atom
§ determine changes in ox #
§ balance changes
§ balance all atoms except H & O
§ balance O’s (add H2 O’s)
§ balance H’s (add H+’s)
§ adjust for basic solution if needed
half-reaction method.
§ determine oxidation & reduction
§ write two separate half-reactions
§ balance all atoms except H & O
§ balance O’s (add H2 O’s)
§ balance H’s (add H+’s)
§ add e− ‘s to more positive side
§ balance e-‘s between half-reactions
§ combine half- reactions
§ adjust for basic solution if needed
4. Electricity can either cause a reaction
(electrolysis, electrolytic cell) or can be
produced by the reaction (Galvanic cell,
electrochemical cell, Voltaic cell).
5. Electrolysis / Electroplating
coulomb (C) = an amount of charge
amp = current = charge per second
1 amp · 1 second = 1 Coulomb
1 C / amp·s
Faraday constant, F:
1 mole e- = 96,500 C
6. Electrolysis calculations begin with amp·s
Example:
How many moles of copper metal can be
plated using a 10 amp circuit for 30 s?
10amp x 30s x 1C x 1 mol e- x 1 mol Ag =
1 amp·s 96500C 1 mol e-3
= 3.1 x 10 mole Ag
7. Spontaneous redox reactions (unlike
electrolysis/electroplating) can simply
occur (as in the ornament lab) or can be
separated so the oxidation and reduction
occur in different containers (half- cells).
In this way, the electrons must move
through an outside wire (this is an
electrochemical cell—a battery).
8. Every atom has a different “potential” to
accept electrons… “reduction potential”
Ag+(aq) + e¯ → Ag(s)
E° = +0.80 v
2+
Cd (aq) + 2e¯ → Cd(s) E° = −0.40 v
These are measured by comparing every
chemical to the same “standard half-cell.”
The reduction with the more positive E°
value will occur as written; the other
reaction will reverse (oxidation).
Ex: 2Ag+ + Cd
2Ag + Cd2+
The difference in the E° values is the
voltage of a cell made using these two
reactions.
Ex: +0.80 v – (-0.40 v) = 1.20 volts
NOTE that you do not multiply the Cd
voltage by 2. Comparing every cell to the
same standard cell accounts for this.
9. Any change that drives the reaction
forward will increase the cell’s voltage.
10. In all electrochemical cells:
Oxidation occurs at the Anode
Reduction occurs at the Cathode
South Pasadena • AP Chemistry
Name __________________________________
Period ___ Date ___/___/___
21 • Electron Transfer Reactions
General Terms
I can…
o Determine the oxidation number of any element.
o State that oxidation number is the charge an
STUDY LIST
o Explain that during the electrolysis of an ionic
solution, either the + ion can be reduced or
water can be reduced. In the same way, either
the – ion can be oxidized or water can be
oxidized.
atom would have if all of the shared electrons
were assigned to the more electronegative atom.
o Use a reduction potential chart to determine
is gaining or losing electrons (LeO GeR).
o State that electrical current is measured in
o Identify for any element in a reaction whether it
o Explain that when oxidation occurs, reduction
must also occur (RedOx).
o Correctly apply the terms oxidizing agent and
reducing agent to a redox reaction.
o State that there are two big topics in
electrochemistry, (1) Electrolysis—in which
electricity (moving electrons) causes chemical
change, and (2) Electrochemical Cells—in
which chemical changes cause a flow of
electrons (electricity).
Electrolysis
which of two substances is more likely to be
reduced or oxidized.
Coulombs and 1 Coulomb = 1 amp·1 sec.
o State that 1 Faraday (F) = 1 mole of electrons =
96,500 Coulombs.
o Use the Faraday, amps and seconds to quantify
electrolysis problems.
Electrochemical Cells
(Voltaic Cells & Galvanic Cells)
I can…
o State that oxidization always occurs at the anode
and reduction always occurs at the cathode.
o Draw a simple electrochemical cell:
I can…
o State that during electrolysis, electricity applied
to a solution causes ions to migrate to the
electrodes.
o State that an electrode is the part of the
conductor that touches the solution.
o State that reduction always occurs at the cathode
(red cat).
o State that oxidation always occurs at the anode
(an ox).
o Write equations for the reactions that occur at
the electrodes when water undergoes
electrolysis (memorize how to derive these).
(−) cathode: 2 H2 O(l) + 2 e− → H2 (g) + 2 OH−
(+) anode: 2 H2 O(l) → O2 (g) + 4 H+ + 4 e−
o Use the reduction potential chart to determine
which chemical is the anode (smaller E°) and
which chemical is the cathode (larger E°).
o State that standard conditions are 25°C,
solutions are 1 M, and gases are 1 atm.
o Calculate the voltage of a standard cell as the
difference in the two E° values. (not like Hess)
o State that the anode is the (−) electrode because
−
the chemicals are being oxidized (losing e ’s).
o State that for non-standard cells, changes that
drive the reaction forward increase the voltage.
(The Nernst equation allows you to calculate
this voltage for a non-standard cell.)