Download A Few Things You Might Want To Know

A Few Things
You Might
Want To Know
For The
2013
AP
CHEMISTRY
EXAM
—Florian König
2012-2013 AP Chemistry
Hastings High School
2
"For a man to attain to an eminent degree in learning costs him time, watching,
hunger, nakedness, dizziness in the head, weakness in the stomach, and other
inconveniences."
– Miguel de Cervantes
“It’s not having been in the dark house, it’s having left it, that counts.”
– Theodore Roosevelt
3
TABLE OF CONTENTS
Basic Concepts
Atomic Theory and the Periodic Table
Electronic Structure of the Atom
Quantum Mechanics
Ionic Bonding
Covalent Bonding
VSEPR Model
Hybridization
MO Theory
Reaction Types and Stoichiometry
Gases
Liquids
Solids
Phase Diagrams
Solutions and Solubility Rules
Acids and Bases in Water
Redox Reactions
Kinetics
Equilibrium
Aqueous Equilibria I (Acids and Bases)
Aqueous Equilibria II (Solutions)
Thermodynamics
Electrochemistry
Organic Chemistry
Nuclear Chemistry
Before the Test
During the Test
4
5
6
7
10
11
13
15
16
18
19
21
23
25
26
29
30
31
33
34
40
41
45
48
50
52
52
Appendix A:
Appendix B:
Appendix C:
Appendix D:
Appendix E:
Appendix F:
53
100
102
104
105
109
Equations in the AP Examinations
Solubility Rules For Ionic Compounds in Water
Nomenclature of Inorganic Compounds
Nomenclature of Organic Compounds
Miscellaneous
Helpful Websites
4
BASIC CONCEPTS
1.
Physical changes do not affect the nature of a compound; chemical
changes do. Physical and chemical properties go along with according
changes.
2.
Mixtures can be heterogeneous or homogeneous (= solutions).
They consist of substances that can be separated by physical changes
(distillation, crystallization, chromatography). Substances can be either
elements or compounds. Compounds can be separated into elements
by chemical changes (redox reactions).
3.
Law of Definite Proportions (Law of Constant Composition)
The composition of a compound is constant. The relative masses of
elements in a compound form simple, whole-number ratios.
4.
Law of Multiple Proportions
If two elements combine with each other in more than one ratio, the
relative masses of the varying element form simple, whole-number ratios.
5.
Intensive and Extensive Properties
Intensive properties do not depend on the amount of matter (solubility,
density, color, reactivity, conductivity, etc.); extensive properties do (mass,
volume, etc.).
6.
Significant Figures

Significant figures = all certain digits + 1 uncertain digit

Determination of significant figures: Atlantic-Pacific Rule
Calculation
Significant figures are the …
Addition/subtraction
least number of decimal places
Multiplication/division
least number of significant figures
pH
only decimal places are significant
Notes:
5
ATOMIC THEORY AND THE PERIODIC TABLE
7.
Atomic Theory
Contributions by LAVOISIER and DALTON ( 3. Law of Definite
Proportions)
8.
THOMSON: “cathode rays”

Proved existence of electrons

Determination of charge/mass ratio of e
Leads to “plum pudding model” of atoms
9.
MILLIKAN: “oil-drop experiment”

Determination of e- charge (see Appendix E)

Determination of e- mass (see Appendix E)
10.
RUTHERFORD: “gold foil experiment”

Atoms consist of small, heavy, positive nuclei.

Electrons are negative, and circle the nucleus.

The nucleus-electron distance far exceeds the nucleus size.
(“Atoms are mostly empty space.”)
11.
Nuclides are nuclei with

Specific proton and neutron number (same #p + = isotopes)
12.
The Periodic Table
(see: DÖBEREINER, LOWLAND, MENDELEEV, MOSELEY)

Periodic arrangement by atomic number and properties

Group numbers and names, representative elements

Distinction between metals, semimetals and nonmetals

Trends in metallic/nonmetallic character, ionization energy, electron
affinity, electronegativity, activity series of metals
Notes:
6
ELECTRONIC STRUCTURE OF THE ATOM
13.
Electronic Structure of the Atom

Speed of light c = (frequency ) x (wavelength )

Visible spectrum: 400 nm (blue)  700 nm (red)

Energy is found in “quanta” with E = h (h = PLANCK’s constant).

Line spectrum: distinct wavelengths, i.e., electron transitions
between distinct energy levels

Continuous spectrum: no distinct wavelengths
14.
BOHR’s Model of the Hydrogen Atom

Absorption or release of energy = jump of e- between discrete
energy levels

All e- in lowest possible state: “ground state”

Flaw: electrons in ground-state atoms do not lose energy as
predicted by classical mechanics
15.
Matter-Wave Dualism (deBROGLIE)

Very small particles (photons, electrons) can exhibit wave behavior
h
 = ——
mv

16.
m = mass (kg); v = velocity (m/s)
“Allowed orbit” = orbit length is a multiple of the electron
wavelength
HEISENBERG’s Uncertainty Principle

It is impossible to know position and momentum of an electron
exactly.

Uncertainty of momentum (p = mv):
p

Uncertainty of position:
q
h
p q  ——
2
Notes:
7
QUANTUM MECHANICS
17.
Quantum-mechanical Description of the Atom:

SCHRÖDINGER: Electrons are described by wave-functions 
(exact only in one-electron systems, e.g., H, He+).

2 = probability of finding e- = deBROGLIE “orbital “ (replaces
BOHR’s “orbits”)

Quantum numbers:
 principal quantum number ( energy): n = 1, 2, 3, 4, …
 azimuthal quantum number ( orbital shape):
l = 0, 1, …, (n-1) 
corresponds to s, p, d, f
Orbital type
s
p
d


Orbital shape
spherical
“dumbbells”
“double dumbbells”
magnetic quantum number ( orbital orientation):
ml = - l, …, 0, … , + l
spin quantum number ( orientation of e- in magnetic field)
ms = + ½ , - ½
18.
PAULI Exclusion Principle

No two electrons in an atom have the same four quantum numbers.
19.
The Aufbau Principle

Arrange orbitals in energetic sequence

Start filling sublevels with lowest energy first
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
7p
8


20.

Electron arrangement:
 all electrons are in lowest possible state  ground state
 one or more electrons occupy higher levels  excited state
s- and p-electrons of the outermost principal energy level (“valence
shell”) are called “valence electrons.” (determine reactivity)
HUND’s Rule
Orbitals with the same energy (= same sublevel) are called “degenerate.”
They are filled with one electron each before pairing up occurs.
9
Notes:
10
IONIC BONDING
21.
“Octet Rule”

Atoms tend to attain noble-gas configuration, i.e., a filled valence
shell with 8 (or 2) valence electrons.
22.
Transition Metals

s-electrons ionize first; half-filled or filled sublevels (especially dsublevels) are relatively stable.
23.
The Ionic Bond

Octet Rule is followed by gaining or losing valence electrons.

Metal (low EN) + nonmetal (high EN)

Determining factor for stability of compound: ionization energy +
electron affinity + lattice energy

Cations are smaller than metal atoms, anions bigger than nonmetal
atoms.

Ionic compounds: hard, brittle, high m.p./ b.p., conductors when
liquid or dissolved in polar solvents (= electrolytes)
Notes:
11
COVALENT BONDING
24.
The Covalent Bond

Nonmetal + nonmetal

Octet Rule is followed by sharing electrons (1-3 per atom).
25.
Multiple Bonds

Up to six electrons (= triple bond) are shared between two atoms.

Multiple bonds are shorter and stronger than single bonds.
26.
Lewis Structures

Add up available valence e-, add/subtract extra charges.

Derive central atom (often C, N, P, S).

Draw single bonds first.

Put excess valence e- at central atom.

Electron deficiencies? (if yes: form multiple bonds)

Structures with low formal charges are favored.

Negative charges on electronegative atoms are favored.

Formal charges: bonds are broken homolytically (= evenly)
(compare # of valence e- to “normal” # in a neutral atom).
27.
Resonance

Degenerate “resonance formulas” add to stability.
28.
Exceptions to Octet Rule

Odd number of valence e- (e.g., NO, NO2)

Less than 8 valence e- (Per. 2 elements: e.g., BeH2, BF3)

More than 8 valence e- require available d-orbitals (Per. 3+
elements: e.g., ICl4-, SF6)
29.
Bond and Molecule Polarity and Electronegativity

EN < 0.4
nonpolar

0.4 < EN < 1.7
polar

EN > 1.7
ionic

Polar molecules only if polar bonds and unsymmetrical geometry

Bond and molecule polarities are not discrete properties, i.e., they
change gradually. The EN differences serve as rough guidelines
only for the determination of bind polarity.
12
30.
Oxidation Numbers

Bonds are formally broken heterolytically, i.e., both e - are counted
towards the more electronegative atom

Elements: 0

Monoatomic ions: oxidation number = charge (e.g., Na +, N3-)

Polyatomic compounds: sum of oxidation numbers = total ionic
charge
Notes:
13
VSEPR MODEL
31.
Maximum Repulsion ( angle) between electron pairs

VSEPR = Valence Shell Electron Pair Repulsion

Repulsion: lone pair > bonding pair > unpaired single electron

Distinction between molecular geometry and electron-pair
geometry

“Electron-pair geometry” is also called “electron-domain (ED)
geometry.”

Multiple bonds are treated as one ED, repel stronger (higher edensity).
32.
Common Electron-Domain and Molecule Geometries
# of ED
2
3
4
5
6
Example
CO2
BeH2
BCl3
NO2CH4
NH3
H2O
PF5
SF4
ClF3
XeF2
SF6
BrF5
XeF4
ED geometry
linear
linear
trigonal planar
trigonal planar
tetrahedral
tetrahedral
tetrahedral
trigonal bipyramidal
trigonal bipyramidal
trigonal bipyramidal
trigonal bipyramidal
octahedral
octahedral
octahedral
Molecule Geometry
linear
linear
trigonal planar
bent
tetrahedral
pyramidal
bent
trigonal bipyramidal
seesaw
T-shaped
linear
octahedral
square pyramidal
square planar
The ideal bond angles and electron-domain geometries are:
 180o
(linear)
o
 120
(trigonal planar)
o
 109.5
(tetrahedral)
 90o/120o (trigonal bipyramidal)
 90o
(octahedral)
Actual bond angles can vary due to the repulsive effect of
1) lone pair < single bond and 2) single bond < double bond < triple bond
14
Notes:
15
HYBRIDIZATION
33.
Symmetrical Overlap

-bond (between nuclei)
34.
Unsymmetrical Overlap

-bond (above nucleus-nucleus axis)
35.
Orbitals

s-orbitals: only -bonds

p-orbitals: - and -bonds possible
36.
Hybridization

At least 2 orbitals from degenerate hybrid orbitals

Number of hybridizing orbitals = Number of hybrid orbitals
37.
Common Hybridizations
Hybridization
sp
sp2
sp3
sp2d
sp3d
sp3d2
Notes:
Example
BeH2, HgCl2, C2H2
C2H4, BF3
CH4, NH3, H2O
PdBr42PF5, SF4
SF6, SbCl6-, ICl4-
Electron-domain geometry
linear
trigonal planar
tetrahedral
square planar
trigonal bipyramidal
octahedral
16
MO THEORY
38.
“O2 problem”

Valence-bond and hybridization approach provide no explanation
for the paramagnetism (= unpaired e-) of O2
39.
Atomic Orbitals (AO) and Molecular Orbitals (MO)

AOs combine to form bonding (lower energy) and antibonding
(higher energy) MOs  see textbook for diagram

Bonding: , ; antibonding: * , * (subscript = constituting AOs)

e.g.: 1s — *1s — 2s — *2s — 2p — 2p (2) — 2p* (2) — 2p*
40.
Bond Order (B.O.)

B.O. = ½ [(# of bonding e-) – (# of antibonding e-)]

B.O. > 0
stable bond

B.O.  0
no stable bond
17
41.
Rules

# of AO = # of MO

AOs combine effectively if they have good overlap and similar
energies.

MOs belong to the whole molecule.

HUND’s Rule and the PAULI Exclusion Principle apply to MOs.
42.
Delocalization

p-orbitals that don’t participate in hybridization form a -MO that
stretches across the molecule.

e- are evenly distributed (example: benzene C6H6, nitrate NO3-).

Stabilization of molecule leads to lower reactivity.
Notes:
18
REACTION TYPES AND STOICHIOMETRY
43.
Reaction Types:

see Appendix A
44.
Gram Atomic Mass (gam) / gram molecular mass (gmm) / gram formula
mass (gfm) = mass of 1 mole of substance
45.
Conversions

(# of moles) x NA = # of particles (atoms, molecules, ions)

(# of moles) x gfm = mass

(# of moles) x 22.4 l mol-1 = volume (for ideal gases at STP)
46.
Empirical Formula from Quantitative Analysis

Assume 100g-sample.

Change % into masses.

Change masses into moles.

Write formula.

Divide by smallest subscript.

Multiply by integer, if necessary, to remove simple fractions such as
0.5 or 0.33.
47.
Limiting Reactant

Check stoichiometry of reaction for “perfect” ratio.

Check reactant that is “used up” first.
48.
Concentrations

moles solute (n)
Molarity M = ————————
liters solution (V)

Dilution: n1 = M1V1 = M2V2 = n2
Notes:
19
GASES
49.
Properties

take shape of container

can be easily compressed
50.
Pressure

Manometer (standard pressure: 760 mm Hg = 760 Torr = 1 atm
= 101.3 kPa)
51.
Gas Laws

BOYLE’s Law: pV = constant (T = const.)



V
CHARLES’ Law : —— = constant (p = const.)
T
AVOGADRO’s Hypothesis:
Equal volumes of gas at equal temperature and equal pressure
contain equal numbers of molecules.
AVOGADRO’s Law:
The volume of a gas is proportional to the quantity. (V  n)
52.
The Ideal Gas Equation

pV = nRT
(p = pressure, V = volume, n = number of moles,
R = Universal Gas Constant, T = absolute temperature [Kelvin])
53.
DALTON’s Law of Partial Pressures

ptotal = p1 + p2 + p3 + … + pn (total pressure = sum of partial
pressures)
54.
Molecular Weight and Gas Densities

55.
Molecular mass MM =
dRT
——
p
Gas Pressure (of gas collected over water)

external pressure = ptotal = pgas + pwater
d = density (gram/liter)
20
56.
The Kinetic-Molecular Theory of Ideal Gases

Gases consist of large numbers of molecules that are in constant,
random, straight-lined motion.

The volume of gas molecules is negligible compared to the total
volume.

There are only negligible intermolecular forces between molecules.

Collisions are elastic, i.e., no thermal energy is lost.

Average kinetic energy  temperature
57.
GRAHAM’s Law of Diffusion and Effusion

At equal temperatures heavier molecules move slower.
58.
 m2
——— (v1, v2, m1, m2 = molecular speeds & masses)
 m1

v1
——— =
v2


Effusion: rate of movement through small hole
Diffusion: rate of movement through space
Nonideal Gases

Deviations at high pressure and/or low temperature
an2

(p + ——) (V- nb) = RT
V2

a = corrective term for intermoleclar attractions (reduces pressure)

b = corrective term for finite molecular volume (increases overall
volume)

Manifestations of nonideal behavior:
 condensation at high pressure and/or low temperatures
 condensate takes up space
Notes:
21
LIQUIDS
59.
Solids

Regular arrangement (crystalline lattice)

Particles close together

Little vibration, no free motion
60.
Liquids

Partial disorder (clusters and particles)

Still relatively close together
61.
Gases

Maximum disorder

Free motion

Big intermolecular distance
62.
Equilibria

Between phases: solid
liquid
gas
63.
Phase Changes

For each phase change there is a molar enthalpy, e.g., Hfusion,
Hvapor., etc.)

No temperature change, i.e., enthalpy change = change in potential
energy
64.
Properties of Liquids

Strong intermolecular forces  high enthalpy of vaporization 
 low volatility  low vapor pressure  high boiling point
65.
Critical Temperature Tc / Critical Pressure pc

There is a maximum temperature at which a gas can be liquefied.
 T > Tc
no condensation possible
 T = Tc
condensation at pc
 T < Tc
gas will condensate at p < pc
66.
Viscosity

How easily molecules move around each other

Depends on attractive forces and structure (long hydrocarbons!)

Decreases with increasing temperature
67.
Surface Tension

Minimizes surface, makes liquids “bead up”
22
68.
Cohesive Forces

Between molecules of a sample

Causes surface tension and viscosity

Affects vapor pressure, boiling point, Hvap.
69.
Adhesive forces

Between different substances, e.g., liquid and container

Intermolecular forces:
 ion-dipole (esp. in aqueous solutions)
 dipole-dipole (between polar molecules)
 London dispersion forces (temporary, weak dipoles)
 hydrogen bonding (between molecules with H+), require
electronegative atoms, usually N, O, F
Notes:
23
SOLIDS
70.
Crystalline Solids

Highly regular arrangement

Crystalline planes

Show cleavage
71.
Amorphous Solids

Only short-range order, long-range disorder

Example: glass, soot
72.
Atomic Solids

Atoms

London dispersion forces

Soft, low m.p., poor conductors

Example: noble gases
73.
Molecular Solids

Molecules

London dispersion forces, dipole-dipole, hydrogen bond

Soft, low to moderate m.p., poor conductor

Example: CH4, carbohydrates, CO2, H2O
74.
Ionic Solids

Ions (mono- or polyatomic)

Electrostatic attraction (ionic bond)

Hard and brittle, high m.p., poor conductors as solids

Electrolyte strength depends on solubility.

Example: NaCl, FeSO4, Ag2Cr2O7
75.
Network Solids

Atoms

Network of covalent bonds

Hard, high m.p., poor conductor

Example: diamond C, quartz SiO2, β-boron nitride (BN)∞
76.
Metallic Solids

Metallic bond (ions and delocalized valence e-)

Soft to very hard, low to high m.p., good conductors

Example: Cu, Fe, Na, Ag, Au, alloys
24
Notes:
25
PHASE DIAGRAMS
77.
See Diagrams in Textbook, Notes and Review Books
78.
State of matter is a function of

temperature AND

pressure
79.
Parts of Phase Diagram

Three curves connecting four points A, B, C, D

solid
liquid
gas

AB: vapor pressure curve

AD: melting curve
 positive slope = solid is denser than liquid
 negative slope = solid is less dense than liquid (e.g., H2O)

AC: vapor pressure curve of solid (describes sublimation)

A: triple point (only point where all three phases coexist)

B: critical point (above B liquid and gas are indistinguishable)
Notes:
26
SOLUTIONS AND SOLUBILITY RULES
80.
Solutions = Homogeneous Mixtures

solute (smaller part, often undergoes phase change)

solvent (bigger part)
81.
Concentrations

parts per million
(ppm)
=
mass solute
—————— x 106
mass solution

parts per billion
(ppb)
=
mass solute
—————— x 109
mass solution
=
mass solute
—————— x 100%
mass solution

82.
Weight percent

Volume percent
=
volume solute
——————— x 100%
volume solution

Mole fraction X
=
moles solute
———————
total # of moles

Molarity M
=
moles solute
——————— (titrations)
liters of solution

Molality M
=
moles solute
——————— (colligative properties)
kg of solvent
Solution Process

Solute-solute interaction (e.g., ionic bond) H1

Solvent-solvent interaction (e.g., hydrogen bond) H2

Solute-solvent interaction (e.g., ion-dipole forces) H3

Hsolv. = H1 + H2 + H3
 Hsolv. > 0 (exothermic, e.g., NaOH)
 Hsolv. < 0 (endothermic, e.g., NH4NO3)
27
83.
Saturated Solution

Dynamic equilibrium of solution with undissolved solvent
84.
Factors Affecting Solubility

Molecular structure (“likes dissolves like”)

Pressure (gases only; Henry’s Law: cg = kpg)

Temperature
85.
Electrolytes

Electrolytes are substances that conduct electricity in aqueous
solutions or when molten.

Electrolytes in water dissociate into hydrated ions.

Strong electrolytes:
 strong acids/ bases (e.g., HNO3, H2SO4, NaOH, KOH)
 soluble salts (e.g., NaNO3, FeBr3)
 osmosis (osmotic pressure  = MRT)

Weak electrolytes:
 weak acids/ bases (e.g., CH3COOH, HNO2, HF, Ca(OH)2,
Ba(OH)2)
 poorly soluble salts (e.g., CaCO3, SrSO4, Hg2Cl2)

Nonelectrolytes:
 organic compounds (e.g., carbohydrates, proteins, alcohols)
86.
Vapor Pressure of Solutions

Volatile solvents (e.g., alcohol) increase vapor pressure, reduce
b.p.

Nonvolatile solvents (e.g., ionic compounds) decrease vapor
pressure, increases b.p.
87.
RAOULT’s Law

pA = XApAo
pA = partial vapor pressure of A above the solution
XA = mole fraction of solvent
pAo = vapor pressure of pure solvent
88.
Colligative Properties

Depend only on number of solvent molecules

Concentration is given in molality m (except for osmosis)

Examples:
 boiling point elevation (Tb = kb i m)
 freezing point depression(Tf = kf i m)
 osmosis (osmotic pressure  = MRT)

i: dissociation factor (“van’t HOFF factor”)

kb: boiling-point elevation constant (solvent-specific)

kf: freezing-point depression constant (solvent-specific)
28
89.
Related Systems

Solution
 homogeneous
 transparent
 particle size  10-9 m
 particles do not settle
 no Tyndall effect (scattering of light)
 no filtration possible

Suspension
 heterogeneous
 particle size > 10-7 m
 particles settle
 can be filtrated
 Tyndall effect

Colloid
 particle size = 10-9  10-7 m
 no filtration possible
 no settling
 Tyndall effect

Emulsion: colloid of liquid in liquid, e.g. oil in water, milk
90.
Solubility Rules:

see Appendix B
91.
Factors Driving a Chemical Reaction

Formation of a precipitate (e.g., Ba2+ + SO42-  BaSO4)

Neutralization (formation of a nonelectrolyte (e.g., H+ + OH-  H2O)

Formation of a weak electrolyte (e.g., Na2S + H+  H2S + Na+)

Redox reactions (e- transfer) (e.g., CrO42-  Cr3+)

Rupture or formation of covalent bonds
(e.g., CH3-CHO  CH3-COOH)
Notes:
29
ACIDS AND BASES IN WATER
92.
Acid Dissociation

forms H3O+ and bigger adducts such as H5O2+
93.
Oxides (Anhydrides) in Water (oxidation state remains constant)

Nonmetal oxides in water yield acids, e.g., P2O5 + H2O  H3PO4

Metal oxides in water yield bases, e.g., CaO + H2O  Ca(OH)2
94.
BRØNSTED-LOWRY Theory

Acids donate H+.

Bases accept H+.
95.
Conjugate Acids and Bases

Strong acids (big Ka) yield weak conjugate bases (small Kb) and
vice versa

Ka (acid) x Kb (conjugate base) = Kw = 1.0 x 10-14
Notes:
30
REDOX REACTIONS
96.
Balancing Redox Reactions

Write reactants/products, omit spectator ions if present.

Write oxidation number, determine # of transferred electrons.

Write separate oxidation/reduction reactions.

Balance charges with H+ (acidic solutions) or OH- (basic solutions).

Balance mass for each half-reaction with water.

Multiply half reactions  equal # of e- for reduction and oxidation

Add reduction and oxidation equations.

Simplify and add spectator ions, if necessary.
97.
Redox Reactants

Common oxidizing and reducing agents: see Appendix A
Notes:
31
KINETICS
98.
Reaction Rate

Reaction rate = (change of measurable property) / time

Property: mostly concentration ([…]), also pH, color, temperature

Rate decreases with time: rate (t = 0) > rate (t)

Rate has to be referred to one reactant or product.

Rates of appearance or disappearance of substances relate like
stoichiometric coefficients.
e.g.: 4 NH3 + 5 O2  4 NO + 6 H2O
rate (NH3) = - 2/3 rate (H2O)
99.
Collision Theory

Proper atoms collide

Proper orientation

Sufficient energy ( EA)
100.
Effect of Concentration
reaction
order
rate law
integrated rate law
half-life
[A]t =  kt + [A]0
[A]0
t1/2 = ———
2k
rate = k [A]
ln [A]t =  kt + ln [A]0
ln 2
t1/2 = ——
k
rate = k[A]2
1
1
—— = —— + kt
[A]t
[A]0
1
t1/2 = ———
k[A]0
0
rate = k
1
2
32
101.
Effect of Temperature

102.
103.
EA
Arrhenius equation: k = A x exp ( ——)
RT
EA

Plot lnT vs. 1/T yields straight line with slope =  ——
R
Reaction Mechanisms

Series of small, elementary steps

Intermediates = formed and consumed during reaction

Intermediates do not show up in overall equation.

ONLY the rate laws of elementary reactions can be deduced
from stoichiometry.

Multi-step reaction:
 stoichiometry of elementary reactions adds up
 rate law of slowest step = rate law of overall reaction
Catalyst

Reduces EA for forward and reverse reaction

Decreases time needed to reach equilibrium; does not affect
position of equilibrium

Takes part in reaction, but is reformed at the end of the reaction

Can be homo- or heterogeneous

Examples: enzymes, NH3 synthesis (Fe)
Notes:
33
EQUILIBRIUM
104.
Dynamic Equilibrium

Rate (forward reaction) = rate (reverse reaction)

Examples: phase, solution, chemical equilibrium

At equilibrium: all concentrations remain constant
105.
Mass-Action Expression

For aA + bB —> cC + dD Q =
[C]c [D]d

[A]a [B]b


If concentrations are equilibrium concentrations: Q = K eq
Only concentrations included: solutions and gases
106.
Equilibrium Constants Kc and Kp

Relationship between Kc (concentrations) and Kp (partial pressures)
Kp = Kc x (RT)n
n = (moles product) – (moles reactant)
107.
Direction of Reaction

Q > Keq
reaction proceeds towards reactants

Q = Keq
reaction is at equilibrium

Q < Keq
reaction proceeds towards products
108.
LeChatêlier’s Principle
A system at equilibrium will counteract – if possible – any stress, i.e.,
change in temperature, pressure or concentrations.







Notes:
Increased pressure  favors side with fewer gas molecules
Increased temperature  favors endothermic reaction
Increased reactant concentration  stronger formation of products
Add inert gas at const. V  partial pressures remain constant
(= no change of equilibrium)
Add inert gas at const. p  V increases
(= favors side with more gas molecules)
Add catalyst  no change
Keq will only be affected by a temperature change
34
AQUEOUS EQUILIBRIA I (ACIDS AND BASES)
(That’s the chapter of assumptions and negligences.)
109.
Ion Product of Water

Autodissociation of water: H2O
H+ + OH
Kw (25 oC) = [H+][OH-] = 1.0 x 10-14

pH = - log[H+]
pOH = -log[OH-]

pH + pOH =14
110.
Indicators

Weak acids HInd
H+ + Ind- (pH-dependent equilibrium)

Color of HInd and Ind- are different
Indicator
methyl orange
methyl red
litmus
bromothymol blue
phenolphthalein
Color change
red  yellow
red  yellow
red  blue
yellow  blue
colorless  pink
at pH
3.1 – 4.4
4.4 – 6.2
5.0 – 8.0
6.0 – 7.6
8.2 – 10.0
111.
Assumptions

Weak acids (less than 5% dissociation): [HA] = constant

Polyprotic acids, e.g., H3PO4, H2SO4:
If Ka(1) > 1000 x Ka (2), the pH is only determined by the first
deprotonation step.
112.
Acid Strength of Polyprotic Acids

pKa(1) < pKa(2) < pKa (3) (acid strength decreases)
113.
Hydrolysis
Salt + water  acid + base
Strong acid + strong bases 
Weak acid + strong base 
Strong acid + weak base 
Weak acid + weak base 
neutral
basic
acidic
(pH depends on relative strengths)
35
114.
Prediction of Strength

Binary acids/bases HX:
 metal hydrides (bases):
base strength increases with period #, e.g., LiH < NaH < KH
 nonmetal hydrides (acids):
acid strength increases with period #, e.g., HF < HCl < HBr < HI
(Reason: increasing EN of element X, increasing bond length X—H)

Oxyacids HnXOm
 same structure, different EN of central atom X:
higher EN  stronger acid (H2SeO3 < H2SO3)
 same central atom X, different number m of oxygen atoms:
increasing number m  stronger acid
HClO < HClO2 < HClO3 < HClO4
HNO2 < HNO3
H2SO3 < H2SO4
H3PO3 < H3PO4
115.
Lewis Theory of Acids and Bases

Lewis acids: electron-pair acceptors (esp. small, highly charged
metal ions, e.g., Al3+, Fe3+, Cr3+, also BCl3 and analogous
compounds)

Lewis bases: electron-pair donors (e.g., NH3, H2O, CN-, SCN- etc.)
116.
Common-Ion Effect

Addition of an already present ion to an equilibrium shifts it
according to LeChatêlier’s Principle.

Change of solubility
117.
Buffer Solutions

Aqueous solutions of weak acid and its conjugate base

Resists pH change upon addition of little H+/OH
Henderson-Hasselbalch Equation:
[X-]
pH = pKA + log ———
[HX]
or (for base B and conjugated acid BH+)
[BH+]
pOH = pKB + log ———
[B]
36


118.
Biological function: keep pH in blood and cells constant
Examples: HCO3-/CO32-, H2PO4-/HPO42-
Titration Curves

Plot of pH as a function of added titrant

Strong acid + strong base (e.g., HCl + NaOH):
(indicator range around pH = 7)
Titration curve strong acid/strong base
(50.00 ml 0.100 M HCl with 0.100 M NaOH)
14.00
12.00
245.000
10.00
195.000
8.00
pH
145.000
6.00
95.000
4.00
45.000
2.00
0.00
-5.000
0.00 10.00 20.00 30.00 40.00 50.00 60.00 70.00 80.00 90.00 100.0
0
0.100 M NaOH (ml)
37

Weak acid + strong base or weak base/strong acid: Theindicator
range has to be around pH = pKB / pKA of conjugate base/acid
Titration curve weak acid/strong base
(50.00 ml 0.100 M HC2H3O2 with 0.100 M NaOH)
14.00
45.000
pH
12.00
10.00
35.000
8.00
25.000
6.00
15.000
4.00
5.000
2.00
0.00
-5.000
0.00 10.00 20.00 30.00 40.00 50.00 60.00 70.00 80.00 90.00 100.0
0
0.100 M NaOH (ml)
Titration curve weak base/strong acid
(50.00 ml 0.100 M NH3 with 0.100 M HCl)
14.00
98.000
12.00
10.00
78.000
pH
8.00
58.000
6.00
38.000
4.00
18.000
2.00
0.00
0.00
-2.000
10.00 20.00 30.00 40.00 50.00 60.00 70.00 80.00 90.00 100.0
0
0.100 M HCl (ml)
38
Titration and pH at equivalence point:
“Equivalence point” means equal moles of acid and base have
been added. It does not necessarily mean a pH of 7.
Acid
strong
strong
weak
weak
Base
strong
weak
strong
weak
pH at equivalence point
7
<7
>7
 7 (depends on relative strength)

Deviations from pH = 7 at equivalence point are due to hydrolysis of salt

Polyprotic acid, e.g. H3PO4 (loss of all H+  different equiv. points)
H3PO4  H2PO4-  HPO42-  PO43pH (1), pH (2), and pH (3) correspond to the eq. points #1, 2, 3.
Titration curve multiprotic acid/strong base
(50.00 ml 0.100 M H3PO4 with 0.100 M NaOH)
14.00
4.000
12.00
3.500
3.000
10.00
2.500
pH
8.00
2.000
6.00
1.500
4.00
1.000
2.00
0.00
0.00
0.500
25.00
50.00
75.00
100.00
125.00
0.100 M NaOH (ml)
150.00
175.00
0.000
200.00
39
At the halfway equivalence points:
[H3PO4] = [H2PO4-]  pH = pKA (H3PO4)
[H2PO4-] = [HPO42-]  pH = pKA (H2PO4-)
[HPO42-] = [PO43-]  pH = pKA (HPO42-)
Note that the equivalence point “jumps” become smaller. This can
be explained by the decreasing acid strength H3PO4 > H2PO4- >
HPO42-. The third equivalence point becomes undetectable due to
the interference by the autodissociation of water.
Notes:
40
AQUEOUS EQUILIBRIA II (SOLUTIONS)
(We’re still assuming and neglecting.)
119.
Equilibria

Acid-base: homogeneous equilibrium

Solution-precipitate: heterogeneous equilibrium
120.
Dissociation of Ionic Compounds and Solubility Product

Ionic compound AaBb
a A+ + b B
Ksp = [A+]a [B-]b

Calculation of (molecular) solubility based on relative
concentrations of A+ vs. B
If a common ion is added, the contribution via dissociation is
negligible.

Ion product Q = (product of actual ion concentrations)
Solubility product Ksp = (product of equilibrium ion concentrations)

Q > Ksp
precipitation occurs until Q = Ksp
Q = Ksp
saturated solution
Q < Ksp
unsaturated solution, extra solute dissolves
121.
Solubility and pH

Solubility of metal sulfides (CuS, ZnS, FeS, MnS) depends on pH.

Notes:
7.4 x 10-22
In a saturated H2S solution ( 0.1 M): [S2-] = ——————
[H+]2
41
THERMODYNAMICS
122.
Energy

Ability to do work and/or

Transfer heat
123.
Forms of Energy

Kinetic energy Ek = ½ mv2

Potential energy Ep = mgh (gravitational potential)

q1q2
Potential Energy Ep = k —— (electrostatic potential)
d2
(q1, q2 = charges; d = distance between charges)
124.
125.
Unit


Unit: 1 J(oule) = 1 kg m2 s-2
1 J = 4.184 cal
The First Law of Thermodynamics (Law of Conservation of Energy)
The sum of all kinetic and potential energies is constant.

E = Efinal - Einitial
E < 0
system loses energy
E > 0
system gains energy

Energy change E = q + w (q = heat; w = work)
126.
State Functions

The value of a state function depends only on the present state,
not on the initial conditions.

Examples: internal energy, volume, pressure, temperature, density,
mass, refractive index, amount of matter
127.
Path Functions

Depend on how the present state was reached

Examples: work, heat
128.
p-V work and enthalpy

Volume work against constant outside pressure
w = – pV

constant volume
qV = E (heat change = change of internal
energy)

constant pressure qP = H (heat change = change of enthalpy)

reaction enthalpy H = Hproducts – Hreactants
(extensive quantity!)
42
129.
Hess’ Law
If a reaction is carried out in a series of steps, the individual enthalpies
add up to the total enthalpy.

For n steps: Htotal = H1 + H2 + … + Hn

Hess’ Law allows the indirect determination of reaction enthalpies
130.
Heat of Formation

Standard heat of formation is the enthalpy change during the
formation of one mole of substance from its elements in their
standard state.

Hf0 (elements) = 0 (by definition)

Reaction enthalpy = (sum of heats of formation of products) – (sum
of heats of formation of reactants)

H0 =  n Hf0 (products) –  m Hf0 (reactants)
(n, m: stoichiometric coefficients of products and reactants)
131.
Calorimetry

Transferred heat = heat capacity x temperature change
q = C x T (C = heat capacity)

Heat capacity = (mass) x (specific heat)
C=mxs
(m = mass; s = specific heat)
132.
The Second Law of Thermodynamics

The entropy of the universe increases constantly.

An entropy decrease is always accompanied by a bigger entropy
increase somewhere else.

Heat never flows spontaneously from a cold object to a warmer
one.
133.
Entropy S

Synonyms: randomness, chaos, disorder, state of higher
probability, degree of dispersion of heat

Unit: J(oule) / K(elvin)

S = Sfinal – Sinitial

Suniverse = Ssystem + Ssurroundings > 0

Ssolid < Sliquid < Ssolution < Sgas

Entropy increases with temperature, phase changes and number of
particles.
134.
Driving Forces of Reactions

Enthalpy towards minimum

Entropy towards maximum
43
135.
The Third Law of Thermodynamics

Absolute zero cannot be reached in a finite number of steps.

The entropy S of a pure, crystalline substance at T = 0 K is 0.
136.
Standard Entropies S0

Are tabulated and allow calculation of Standard Entropy of
Reaction S0

Standard Entropy of Reaction S0 =  n S0 (p) –  m S0 (r)
n,m: stoichiometric coefficients of products and reactants)
137.
The Free-Energy Function (Gibbs Function)

G = H – TS
H = enthalpy
T = absolute temperature
S = entropy




138.
G = H – TS
G < 0
spontaneous reaction
G > 0
nonspontaneous reaction
At low temperatures: –TS < H
Enthalpy (H) determines spontaneity.
At high temperatures: -TS > H
Entropy (S) determines spontaneity.
Gf0 =  n Gf0 (products) -  m Gf0 (reactants)
(n,m: stoichiometric coefficients of products and reactants)
Spontaneity and Free Energy
“Some things happen spontaneously, some things don't.”
— P.W. Atkins, Physical Chemistry

139.
Sign of G determines direction in which system moves
 G < 0 (forward) reaction is spontaneous
 G = 0 reaction has reached equilibrium
 G > 0 reaction is not spontaneous (reverse reaction is!)
Definition of Standard State
Standard state
pure solid
pure liquid
gas at p = 1 atm
c=1M
at 25 0C, Gf0 = 0
State of matter
solid
liquid
gas
solutions
elements
44
140.
Free Energy, Temperature and Concentrations

G = Go + RT lnQ = Go + 2.303 RT logQ

Equilibrium: Q = Keq; G = 0 (reaction stops), so Go = -RTlnKeq

Sign of G0 allows semiquantitative statement about Keq
 Go is negative 
Keq > 1
o
 G is positive 
Keq < 1
 Go is zero

Keq = 1
Notes:
45
ELECTROCHEMISTRY
141.
Electrochemical Cell

Converts chemical energy into electrical energy

Two half-cells connected with
 wire (= electron flow) and
 salt bridge/porous cup (= ion flow to ensure charge balance)

Reduction occurs at the cathode, oxidation at the anode.
(anode oxidation, reduction cathode)
142.
Cell EMF (electromotive force)






1 J(oule)
Unit: 1 V(olt) = ——————
1 C(oulomb)
Standard EMF E0 at standard cond.: p = 1 atm, c = 1 M, T = 298 K
Listed: standard reduction potentials E0red
Connect two half-reactions:
 bigger E0red will be the reduction
 smaller E0red will be the oxidation (reverse sign!!)
 multiply (if necessary) to eliminate e do not multiply the reduction potentials
 add equations and (unmultiplied) potentials
E0cell = E0red + E0ox
Reference half-cell: S(tandard) H(ydrogen) E(lectrode)
Hydrogen bubbling over a Pt-electrode in acidic solution
p (H2) = 1 atm; [H+] = 1 M  pH = 0
143.
Reducing and Oxidizing Agents

Oxidizing agents are reduced.

Reducing agents are oxidized.

Strong oxidizing agents have big E0red (F2, Au3+, Pb4+, Cl2).

Strong reducing agents have small E0red (Li, K, Na, Ca).
144.
EMF and Spontaneity

E0 > 0
spontaneous (G < 0)
0

E < 0
not spontaneous (G > 0)
46
145.
EMF and Free Energy
 G = nFE
(n = number of electrons transferred; F = Faraday
constant; E = cell potential)
0

G =  n F E0
146.
EMF and Equilibrium Constant/Mass-Action Expression
0.0591
– ———— log Q
n

NERNST Equation:

If electrochemical cell has used up one or more reactants:
G = 0
E=
 E = 0 and
E0
E=
E0
0.0591
– ———— log Keq
n
147.
Commercial Voltaic Cells

Lead battery: Pb + PbO2 + 4 H+ + 2 SO42-  2 PbSO4 + 2 H2O

Dry cell: Zn + NH4Cl + MnO2  ZnCl2 + Mn2O3 + 2 NH3 + H2O
148.
Electrolysis

Reversal of electrochemical cell

Converts electrical energy into chemical energy

Same process (an ox, red cat)

Electrolysis of aqueous solutions: often water is reduced
2 H2O + 2 e-  H2 + 2 OH-
149.
Electrodes

Active electrodes = electrode material participates in electrolysis
Application: electroplating, electroraffination
Examples: Cu, Zn, Co, Ni

Inert electrodes = electrode material does not participate
Examples: C (graphite), Pt
150.
Quantitative Analysis

Amount of charge of 1 mol e- = 1 F = 96,500 C

1 C(oulomb) = 1 A(mpère) x 1 s(econd)
47
151.
Electrical Work

wmax =  n F E

Work done by system = voltaic cell

Work done to system = electrolytical cell

Electrical work: 1 W(att) = 1 J/s 1 kWh = 3.6 x 106 J
152.
Corrosion

Common form of oxidation of metals

Often slowed down by formation of inert oxides, e.g. Al, Cr, Mg

Sped up by salts (electrolytes!)

pH dependent, e.g., Fe does not corrode if pH > 9

Corrosion protection
 coating with metal that is inert to H+ (Sn)
 coating with metal that corrodes first (Zn)
 “sacrificial anode” (Mg)
Notes:
48
ORGANIC CHEMISTRY
153.
Central Feature

Carbon-carbon bond

Second common element: hydrogen
154.
Nomenclature

See Appendix D
155.
Homologous Series

Compounds that differ by number of methylene units –CH2–

Examples: alkanes
156.
Isomerism

Same molecular formula

Different structure

Structural isomerism

CH3

CH3–C–CH3

CH3
CH3CH2CH2CH2CH3
vs.
n-pentane
2,2-dimethylpropane
Geometrical isomerism
CH3CH2CH2
\
CH2CH2CH3
CH3CH2CH2 H
/
\
/
C==C
vs.
C==C
/
\
/
\
H
H
H
CH2CH2CH3
cis
Substituents on
trans
same side of double bond:
different side of double bond:
cis
trans
49
157.
Aromatic Hydrocarbons

Mostly 6-membered rings with delocalized -electrons

Stabilization due to resonance

Examples:
 benzene C6H6 look up formula  notice sextet of electrons
 toluene C6H5CH3
C6H5 = phenyl
158.
Reactions of Hydrocarbons

See Appendix A
159.
Derivatives

Functional groups (often polar) affect properties, e.g. solubility, b.p.
(R = alkyl group or other organic fragment)

ROH
alcohols
CH3CH3 + <O>  CH3CH2OH
1
2

R OR
ethers
R1OH + R2OH  R1OR2 + H2O

O
//
RC
\
H
aldehydes
R—CHO
(R = H formaldehyde
CH3 acetaldehyde)
O

R1CR2
ketones
R1COR2
(R1 = R2 = CH3 acetone
(R1 = R2 = C2H5
diethylketone)

RCOOH
carboxylic acids

R1COOR2
ester
R1COOH + R2OH 
R1COOR2 + H2O

RNH2
amines
(organic ammonia derivatives)

Notes:
(R =
H
formic
CH3 acetic
C2H5 propanoic)
50
NUCLEAR CHEMISTRY
160.
Radioactivity

Spontaneous decay of nuclei (radioisotopes)
161.
Decay modes

Alpha () decay = loss of 42He (charge: +2)
222 Rn
218 Po + 4 He

86
84
2

Beta () decay = loss of electrons 0-1e (does not come from shell,
but from 10n  11p + 0-1e)
131 I
131 Xe + 0 e

53
54
-1

Gamma () decay = high-energy photon (very short wavelength)
(is usually not shown)

Positron emission = 11p  10n + 01e (= antiparticle to electron 0-1e)
11 C
11 B + 0 e

6
5
1

Electron capture from inner shell = conversion of
81 Rb + 0 e 
81 Kr
37
-1
36
1 p
1
+ 0-1e  10n
162.
Stability of Nuclei

All nuclei with more than 83 protons are unstable.

Magic numbers (2, 8, 20, 50, 82, 126) of protons or neutrons are
more stable than adjacent nuclei.

Nuclei with even numbers of protons or neutrons are more stable
than those with uneven numbers.

Nuclei with equal numbers of protons and neutrons are more stable
than those with different numbers.
163.
Radioactive Series

Certain nuclei keep disintegrating until a proton number of less
than 84 is reached.

Example: 23892U  (…)  (…)  20882Pb
164.
Transmutation

Formation of new nuclei

Target nuclei are bombarded with highly accelerated p + or n
51
165.
Half-life

Time needed for a radioisotope to decay to ½ of its mass
(first-order process)

Application: radiocarbon dating (= determination of ages of
minerals and organic matter, e.g., wood)

Useful formulas:
ln [A]t = ln [A]0 – kt
1
[A]0
k = — · ln ——
t
[A]t
ln 2
t1/2 = ——
k
t1/2
[A]0
t = —— · ln ——
ln 2
[A]t
166.
Mass-energy Conversions

Masses of nuclei are lower than the masses of individual nucleons.

Difference: mass converted to energy E = mc2 (= “mass defect”)

Mass defect is converted into binding energy.
Notes:
52
BEFORE THE TEST






Form study groups, but with no more than three people at a time.
Get a review book. Then use it. And use it. And use it.
Do as many sample AP exam questions as you possibly can.
Bother your teacher with questions. After all, it’s his job.
Stop studying the day before. Marathon runners don’t run the day
before a race. You won’t forget anything if you don’t study during the
last 24 hours. Work out. Go for a long walk or see a really bad movie.
Have a nice dinner. Distract yourself.
Get plenty of sleep and have breakfast. If you are tired or hungry,
you won’t do well.
DURING THE TEST







Read the questions carefully. Select essential from superfluous
information.
Scan the questions before you answer them. Rate their difficulty
from easy to medium to hard. Do the easy ones first.
Elimininate wrong answers in the Multiple Choice section. If you
don’t know the right answer, guess as soon as you have eliminated
two answers as wrong.
Do not ramble. The graders don’t want to know what else you know –
keep your answers brief. Stick to basic concepts. Don’t waste your
time.
If a question doesn’t ask you to explain something, don’t.
Keep an eye on the time.
Do not expect to be able to answer every question completely.
Don’t fret over the really hard ones you couldn’t answer.
53
APPENDIX A:
EQUATIONS IN THE AP CHEMISTRY EXAMINATIONS
As of May 2007, the equations in this section need to be balanced.
Balancing can be done either by trial-and-error as in the past, or by a more
sophisticated method that you will need to apply to redox equations.
Omit spectator ions, i.e., ions that stay in solution during the reaction and form
no precipitates (see worksheet “Solubility Rules”). Strong acids that are mostly
dissociated (HCl, HNO3, H2SO4) should be written in their ionized form, e.g., HCl
= H+ + Cl-. Weak acids that are mostly undissociated (HC2H3O2 and similar
organic acids) must be written as undissociated molecules.
Each question consists of two parts:


One equation that must be correctly balanced.
A question about that reaction. This question may involve numbers of
electrons transferred, oxidation numbers, observable changes during
reactions, formulas/names of precipitates, etc.
Since this change is a very recent one, my suggestions as how to approach this
new format can be only tentative.
I will therefore include my previous strategies.
54
1) ACID-BASE REACTIONS:
There are no changes of oxidation states in acid-base reactions.
a) acid and base solutions:
H2SO4 (aq) + Ca(OH)2 (aq)  CaSO4 (s) + 2 H2O should be written as:
H+ + SO42- + Ca2+ + 2 OH-  CaSO4 + 2 H2O
Some acids are unstable:
HCO3- / CO32- + H+  (H2CO3)*  H2O + CO2 (Don’t include H2CO3!)
HSO3- / SO32- + H+  (H2SO3)*  H2O + SO2
(Don’t include H2SO3!)
A star superscript after a compound in parentheses (…)* is often used to indicate
instable compounds or intermediates.
b) solution and anhydride:
N2O5 + 2 OH-  2 NO3- + H2O
or
6 H+ + Cr2O3  2 Cr3+ + 3 H2O
c) anhydrides:
P2O3 + 3 MgO  Mg3(PO3)2
d) amphoteric hydroxides:
Some hydroxides (Sn4+, Zn2+, Al3+, Cr3+)
dissolve in acids as well as excess base.
Al(OH)3 + 3 H+  Al3+ + 3 H2O
Al(OH)3 + 3 OH- (xs.)  Al(OH)63- (also: AlO33- + 3 H2O
e) Lewis acid-base reactions:
Compounds with lone electron pairs react with
electron-deficient compounds.
BF3 + NH3  BF3 • NH3 (F3B—NH3)
f) hydrolysis of nonmetal halides or oxyhalides:
PCl5 + 4 H2O  8 H+ + 5 Cl- + PO43-
Yields two acids.
SO2Cl2 + 2 H2O  4 H+ + SO42- + 2 Cl-
g) hydrolysis of certain transition metal halides:
UF6 + 2 H2O  UO2F2 + 4 H+ + 4 F-
Yields oxyhalides and
acid.
55
2) DOUBLE REPLACEMENT:
compound+ compound  compound + compound
AB
+ CD
 AD
+ BC
2 AgNO3(aq) + BaCl2(aq)  2 AgCl(s) + Ba(NO3)2(aq)
(Ag+ + Cl-  AgCl)
Driving force: Formation of precipitate.
3) SINGLE REPLACEMENT:
element + compound  compound + element
a) cationic single replacement: A + BC  AC + B
(A, B = metals or
hydrogen)
H2 + CuO  H2O + Cu
2 Al + Cr2O3  Al2O3 + 2 Cr
Mg + 2 H+  Mg2+ + H2
b) anionic single replacement: A + BC  AB + C
(A, C = nonmetals)
F2 + 2 I-  2 F- + I2
Cl2 + 2 Br-  2 Cl- + Br2
Driving force: Check Table of Standard Reduction Potentials, often hydrogen is
formed.
56
4) REDOX REACTIONS:
a) oxidizing agents:
oxygen:
halogens:
Elements in (unusually) high oxidation states.
O2 (0), peroxides O22- (-1)  O2-
X2 (0)
e.g. Cl2
XO- (+1)
e.g. ClOXO2- (+3)
e.g. ClO2XO3 (+5)
e.g. ClO3XO4- (+7)
e.g. ClO4(oxidation strength increases with higher oxidation number)
chromates and dichromates:
CrO42-, Cr2O72-  Cr3+
manganese dioxide and (per)manganates:
MnO2  Mn2+
MnO42-, MnO4-  MnO2 (basic)
MnO42-, MnO4-  Mn2+ (acidic)
oxyacids (esp. when concentrated), common: HNO3, H2SO4 and HClO4
HNO3  NO
H2SO4  SO2
HClO4  Cl-
b) reducing agents:
alkali and alkaline earth metals
Na, K, Ca  Na+, K+, Ca2+
aluminum
Al  Al3+
transition metals in low oxidation states
Fe2+ Fe3+
hydrogen
H2  H2O
carbon monoxide
CO  CO2
sulfite and thiosulfate ion
S2O32- / SO32-  SO42-
Cu+  Cu2+
57
c) miscellaneous:
Halogens often oxidize and reduce themselves
in water (= disproportionation),
e.g. Cl2 + H2O
 H+ + Cl- + ClOHalogens in low AND high oxidation states
reach one oxidation state
(= comproportionation),
e.g. I- + IO3- + H+  I2 + H2O
5) (THERMAL) DECOMPOSITIONS:
a) Carbonates (esp. alkaline earth):
CaCO3  CaO + CO2
b) Hydrogencarbonates:
2 NaHCO3  Na2O + H2O + CO2
c) Azides (esp. Na, Ag):
decompose explosively and form
nitrides and nitrogen, e.g.
3 NaN3  Na3N + 4 N2
Driving force: Entropy gain (gas!) and enthalpy gain by formation of NN bond
6) LIGAND EXCHANGE REACTIONS:
Molecules or anions with lone pairs can have a higher affinity to hydrated metal
ions than water and subsequently replace one or all water molecules. (examples:
NH3, CN-, CO)
Cu(H2O)42+ + 4 NH3  Cu(NH3)4+ + 4 H2O
Fe(H2O)63+ + 6 CN-  Fe(CN)63- + 6 H2O
AgCl + 2 CN-  Ag(CN)2- + ClAmmine ligands (NH3) can be removed by acids:
Ag(NH3)2+ + 2 H+  Ag+ + 2 NH4+
58
7) ORGANIC REACTIONS:
a) Addition to multiple bonds:
Alkynes and alkenes add halogens, acids and water.
HCCH + Cl2  HClC=CClH
H2C=CH2 + HBr  H3C—CH2Br
H2C=CH—CH3 + H2O  H3C—CHOH—CH3
b) Substitution of single bonds:
Alkanes and aromatic hydrocarbons substitute halogens for one or more
hydrogen.
CH4 + 4 Cl2  CCl4 + 4 HCl
C6H6 + Br2  C6H5Br + HBr
c) Ether formation: alcohol + alcohol  ether + water (catalysis: H+)
C2H5OH + C2H5OH  C2H5—O—C2H5 + H2O
d) Ester formation: alcohol + acid  ester + water (catalysis: H+)
CH3CH2COOH + CH3CH2CH2CH2OH  CH3CH2—COO—CH2CH2CH2CH3 +
H2O
59
EXAMPLES FROM AP EXAMS:
2007 (example):
(i) A strip of magnesium is added to a solution of silver (I) nitrate.
Mg + 2 Ag+ Mg2+ + 2 Ag
(ii) Which substance is oxidized in the reaction? Mg is oxidized.
2007 (a):
(i) A solution of sodium hydroxide is added to a solution of lead (II) nitrate.
Pb2+ + 2 OH-  Pb(OH)2
(ii) If 1.0 L volumes of 1.0 M solutions of sodium hydroxide and lead (II) nhitrate
are mixed together, how many moles of product(s) will be produced. Assume the
reaction goes to completion.
0.5 moles of lead (II) hydroxide Pb(OH)2 will be formed.
(Since OH- is used up at twice the rate, it is the limiting reactant.)
2007 (b):
(i) Excess nitric acid is added to solid sodium carbonate.
2 H+ + CaCO3  Ca2+ + CO2 + H2O
(ii) Briefly explain why statues made of marble (calcium carbonate) displayed
outdoors in urban areas are deteriorating.
The air in urban areas contains nonmetal oxides such as SO2, SO3, and
NOx. These react as acid anhydrides and create acid rain which has a
deteriorating effect on calcium carbonate.
2007 (c)
(i) A solution containing silver (I) ion (an oxidizing agent) is mixed with a solution
containing iron (II) ion (a reducing agent).
Ag+ + Fe2+  Ag + Fe3+
(ii) If the contents of the reaction mixture described above are filtered, what
substance(s), if any, would remain on the filter paper.
Silver metal (Ag).
60
1974/2
A sample of pure 2-butene is treated with hydrogen bromide gas.
Answer:
2-butene is an alkene and undergoes addition to the double bond.
CH3—CH=CH—CH3 + HBr  CH3—CH2—CHBr—CH3
1988/1
A solution of potassium iodide is added to an acidified solution of potassium
dichromate.
Answer:
a) potassium K+ is a spectator ion  omit!
b) Cr2O72- is an oxidizing agent  iodide is oxidized, dichromate
reduced
I- + H+ + Cr2O72-  I2 + Cr3+ + H2O
1989/2
Solutions of silver nitrate and lithium bromide are mixed.
Answer:
a) lithium Li+ and nitrate NO3- are spectator ions  omit!
b) silver bromide is insoluble.
Ag+ + Br-  AgBr
1990/5
Hydrogen sulfide gas is bubbled through a solution of potassium hydroxide.
Answer:
a) potassium K+ is a spectator ion  omit!
b) H2S is an acid, OH- a base
H2S + OH-  S2- + H2O
1993/4
Excess chlorine gas is passed over hot iron filings.
Answer:
a) chlorine acts as an oxidizing agent, iron as a reducing agent
b) synthesis reaction: metal + nonmetal  salt
Fe + Cl2  FeCl3
61
Please use
 “Equations in the AP Examinations” and
 Appendices D and E
in your textbook.
For each of the following reactions,
(i) write a balanced equation for the reaction
(ii) answer the question about the reaction.
In part (i), coefficients should be in terms of lowest whole numbers.
Assume that solutions are aqueous unless otherwise indicated.
Represent substances in solutions as ions if they are extensively ionized.
Omit formulas for ions or molecules that are unchanged by the reaction.
1)
2)
3)
4)
5)
6)
7)
8)
9)
10)
11)
12)
13)
14)
15)
—I—
(i) Solid lithium hydride is added to water.
(ii) How does the pH of the reaction mixture change?
(i) Excess ammonia is added to a solution of zinc sulfate.
(ii) What particles act as LEWIS acids and bases?
(i) Sodium acetate is acidified with dilute hydrochloric acid.
(ii) How does the pH of the reaction mixture change?
(i) A solution of potassium dichromate is added to an acidified solution of sodium iodide.
(ii) What is total number of electrons transferred?
(i) Solid barium carbonate is added to dilute nitric acid.
(ii) What is a simple analytic test for one of the products?
— II —
(i) Solid aluminum sulfide is added to water.
(ii) Why is the gaseous product toxic?
(i) Silver metal is added to 6M nitric acid.
(ii) What acts as an oxidizing agent?
(i) A solution of tin(II)nitrate is dropped onto a piece of zinc metal.
(ii) What observation can be made about the piece of zinc?
(i) Solid sodium dichromate is added to an acidified solution of iron(II) sulfate.
(ii) What is the total number of electrons transferred?
(i) Liquid titanium tetrachloride is given into excess water.
(ii) How does the pH of the reaction mixture change?
— III —
(i) Solid ammonium chloride is added to a concentrated solution of sodium hydroxide.
(ii) What ion(s), if any, act as spectator ions?
(i) Solid sodium hydrogencarbonate is added to concentrated sodium hydroxide.
(ii) Sodium hydrogencarbonate reacts with hydrochloric acid as well. Name this behavior.
(i) Solid calcium sulfite is acidified.
(ii) Name one unstable reaction intermediate.
(i) Liquid acetic acid and methanol are combined and warmed.
(ii) What specific organic reaction type occurs here?
(i) Aqueous solutions of sodium iodate and sodium iodide are combined, then acidified.
(ii) What specific type of redox reaction occurs here?
62
16)
17)
18)
19)
20)
21)
22)
23)
24)
25)
26)
27)
28)
29)
30)
31)
32)
33)
34)
35)
— IV —
(i) Solid calcium phosphate is given to excess hydrobromic acid.
(ii) What phosphorus-containing product would be formed if hydrobromic acid and calcium
phosphate reacted in a mole ratio of 1:1?
(i) A piece of magnesium is dropped into hydrochloric acid.
(ii) How many electrons are transferred in this reaction?
(i) A stream of carbon dioxide gas is bubbled through a suspension of barium carbonate.
(ii) What observation can be made regarding the appearance of the reaction mixture?
(i) Potassium permanganate solution is added to concentrated hydrochloric acid.
(ii) How does the color of the reaction mixture change?
(i) Ammonia gas is bubbled through a suspension of zinc hydroxide.
(ii) What observation can be made regarding the appearance of the reaction mixture?
—V—
(i) Hydrogen sulfide gas reacts with moist lead(II)chloride.
(ii) What color change could be observed?
(i) Moist sodium hydroxide reacts with carbon dioxide gas.
(ii) How and why does the pH of the reaction mixture change?
(i) An aqueous solution of ammonia is added to solid silver iodide.
(ii) Name this type of reaction.
(i) Rubidium permanganate is added to hydroiodic acid.
(ii) How many electrons are transferred in this reaction?
(i) Liquid phosporus trichloride is poured into an excess of sodium hydroxide.
(iii) What are the oxidation states of phosphorus in reactant(s) and product(s)?
— VI —
(i) 30% hydrogen peroxide is heated.
(ii) What are the oxidation states of oxygen in reactant(s) and product(s)?
(i) Liquid mercury is added to concentrated nitric acid.
(ii) What substance acts as a reducing agent?
(i) Solid iron(III)sulfate is added to a sodium iodide solution.
(ii) What substance could be used to detect one of the products in small concentrations?
(i) Hydrogen sulfide is bubbled through a solution of silver nitrate.
(ii) Name this specific type of reaction.
(i) Concentrated (15 M) ammonia solution is added in excess to a solution of
copper(II)nitrate.
(ii) Describe any observable color changes.
— VII —
(i) Magnesium metal is added to dilute nitric acid, giving as one of the products a
compound in which the oxidation number of nitrogen is -3.
(ii) Name said product.
(i) Excess water is added to solid calcium hydride.
(ii) Why must calcium hydride be stored in a well-sealed container?
(i) Silver acetate solution is dropped into a solution of sodium phosphate.
(ii) If the contents of the reaction mixture would be filtered, what substance(s) would
remain on the filter paper?
(i) Solid sodium cyanide is added to water.
(ii) Why is sodium cyanide toxic?
(i) Solid potassium hydride is added to anhydrous ethyl alcohol.
(ii) Name the solid reaction product.
63
36)
37)
38)
39)
40)
41)
42)
43)
44)
45)
46)
47)
48)
49)
50)
51)
52)
53)
54)
55)
— VIII —
(i) Lithium metal is burned in air.
(ii) Given sufficiently high temperatures, what byproduct could be formed?
(i) Aluminum metal is added to a solution of copper(II)chloride.
(ii) What, if any, color change would be observable?
(i) Manganese(II)nitrate is mixed with sodium hydroxide solution.
(ii) Name any spectator ions.
(i) Equal volumes of dilute equimolar solutions of sodium carbonate and hydrochloric acid
are mixed.
(ii) What ratio of sodium carbonate to hydrochloric acid would lead to the formation of a
gas?
(i) Solid sodium carbide is added to an excess of water.
(ii) How does the pH of the reaction mixture change?
— IX —
(i) An excess of sodium hydroxide solution is added to a solution of aluminum chloride.
(ii) Describe any visible changes during the reaction.
(i) Solid ammonium carbonate is heated.
(ii) Which of the products has the lowest boiling point?
(i) Phosphorus pentachloride is added to an excess of potassium hydroxide solution.
(ii) Name the type of reaction that the phosphorus pentachloride undergoes.
(i) Potassium chlorate is heated in the presence of manganese dioxide.
(ii) What function does the manganese dioxide have?
(i) Chlorine gas is bubbled into a solution of sodium hydroxide.
(ii) What oxidation state(s) change(s)?
—X—
(i) Magnesium carbonate is heated strongly.
(ii) Which of the product is an acid anhydride?
(i) Methanol and propanoic acid are mixed.
(ii) What general organic reaction takes place here?
(i) Barium acetate is added to dilute sulfuric acid.
(ii) What, if any, spectator ions occur here?
(i) Ammonia gas is bubbled through a suspension of silver chloride.
(ii) Describe any visible changes of the reaction mixture.
(i) Hot iron(III)oxide is reacted with carbon monoxide.
(ii) How many electrons are transferred during the reaction?
— XI —
(i) Silane (SiH4) is burned in air.
(ii) What type of solid is the product containing silicon?
(i) Cesium metal is heated with sulfur.
(ii) Name this type of reaction.
(i) Sodium phosphate is dissolved in water.
(ii) How does the pH of the reaction mixture change?
(i) Ethyl ethanoate is boiled with sodium hydroxide.
(ii) What type of organic compound is ethyl ethanoate?
(i) Solutions of sodium phosphate and calcium chloride are combined.
(ii) Name any spectator ion(s).
64
56)
57)
58)
59)
60)
61)
62)
63)
64)
65)
66)
67)
68)
69)
70)
71)
72)
73)
74)
75)
— XII —
(i) A suspension of aluminum hydroxide is added to hydroiodic acid.
(ii) How would the reaction be different if hydrochloric acid were used?
(i) Hydrogen peroxide is added to an acidified solution of sodium iodide.
(ii) How many electrons are transferred during this reaction?
(i) Hydrogen peroxide is added to an acidified solution of potassium dichromate.
(ii) What are the oxidation states of chromium in this reaction?
(i) Calcium is added to dilute hydrochloric acid.
(ii) What specific inorganic reaction type occurs here?
(i) Aluminum nitrate is dissolved in water.
(ii) Which substance acts as LEWIS acid here?
— XIII —
(i) Phosphorus tribromide is added to water.
(ii) Name the product containing phosporus.
(i) Solutions of sulfuric acid and lithium hydroxide are mixed.
(ii) What specific analytical method would use this type of reaction?
(i) Nitric acid is added to a solution of tetraamminecadmium(II).
(ii) What spectator ion(s), if any, occur in this reaction?
(i) 2-butene reacts with hydrobromic acid.
(ii) What functional group of 2-butene is active in this reaction?
(i) Hydrogensulfide is bubbled through a solution of lead(II)nitrate.
(ii) What color change could be observed?
— XIV —
(i) Solutions of iron(III)nitrate and sodium thiocyanate are combined.
(ii) What color change could be observed?
(i) Ammonia gas is bubbled through a suspension of zinc hydroxide.
(ii) How many electrons are transferred during this reaction?
(i) Sulfur dioxide gas is bubbled through an acidified solution of potassium permanganate.
(ii) How many electrons are transferred during this reaction?
(i) Aluminum is added to excess sodium hydroxide solution.
(ii) Name the ionic product of this reaction.
(i) A solution of tin(II)nitrate is added to an acidified solution of potassium dichromate.
(ii) How many electrons are transferred during this reaction?
— XV —
(i) Equal volumes of equimolar sodium hydroxide and sodium dihydrogenphosphate are
mixed.
(ii) What phosphorus-containing product would be formed if the ratio of hydroxide to
dihydrogenphosphate were 2:1?
(i) Excess sodium hydroxide is added to a solution of calcium hydrogencarbonate.
(ii) Name any spectator ions.
(i) Acetylene (ethyne) reacts with an an excess of chlorine.
(ii) What product would be formed if the reactants were used in a ratio of 1:1?
(i) Solid silver nitrate is added to a solution of sodium chromate.
(ii) Name any spectator ions.
(i) An excess of solid potassium hydroxide is added to a suspension of zinc hydroxide.
(ii) What changes in the reaction mixture can be observed?
65
76)
77)
78)
79)
80)
81)
82)
83)
84)
85)
86)
87)
88)
89)
90)
91)
92)
93)
94)
95)
— XVI —
(i) Benzene reacts with bromine.
(ii) Why does benzene not react as if it had double bonds?
(i) A direct electric current is passed through a dilute solution of sulfuric acid.
(ii) What gas is formed at the cathode?
(i) Ammonia and oxygen are heated in the presence of a catalyst.
(ii) How does the oxidation number for nitrogen change?
(i) Liquid bromine is added to a solution of sodium iodide.
(ii) What specific type of inorganic reaction occurs here?
(i) Dilute hydrochloric acid is added to solid calcium oxide.
(ii) Name the spectator ion.
— XVII —
(i) Sodium hydride is added to water.
(ii) Compare the relative reactivity of potassium hydride with sodium hydride.
(i) Ethene reacts with water in the presence of a catalyst.
(ii) Name the product.
(i) Solutions of barium hydroxide and iron(III)sulfate are combined.
(ii) If equal moles of reactants were combined, which would be the limiting reactant?
(i) Ammonia gas is bubbled through a solution of copper(II)nitrate.
(ii) What color change could be observed?
(i) Lead is added to a dilute solution of sulfuric acid.
(ii) How many electrons are transferred during this reaction?
— XVIII —
(i) Lead is added to a hot solution of concentrated sulfuric acid.
(ii) What acts as the oxidizing agent?
(i) Calcium oxide is exposed to an atmosphere of carbon dioxide.
(ii) Which reactant is an acid anhydride?
(i) Nitrogen(V)oxide is bubbled through water.
(ii) How does the pH of this reaction change?
(i) Solid sodium hydrogencarbonate is added to water.
(ii) How does the pH of this reaction change?
(i) Solid sodium hydroxide and solid ammonium chloride are mixed and heated.
(ii) Name the gaseous product.
— XIX —
(i) Methyl iodide (iodomethane) is heated with a solution of sodium hydroxide.
(ii) Compare the reactivity of methyl iodide with the reactivity of methyl chloride.
(i) Solid barium hydroxide and solid ammonium sulfate are mixed and heated.
(ii) What product has the highest melting point?
(i) A solution of sodium oxalate is added to an acidified solution of potassium
permanganate.
(ii) How many electrons are transferred during this reaction?
(i) Equal volumes of equimolar hydrochloric acid are added to a solution of sodium
hydrogenphosphate.
(iii) What would be the stoichiometric ratio of the reactants?
(i) Ethanol burns in air.
(ii) Name a reagent that would detect the heavier of the products.
66
96)
97)
98)
99)
100)
101)
102)
103)
104)
105)
106)
107)
108)
109)
110)
111)
112)
113)
114)
115)
— XX —
(i) A strip of magnesium is put in a solution of iron(III)nitrate.
(ii) What is reduced in this reaction?
(i) Hydrogen peroxide is heated.
(ii) What changes of oxidation state(s) occur here?
(i) Iron filings are sprinkled into a solution of iron(III)chloride.
(ii) What substance is oxidized?
(i) Chlorine is bubbled through a solution of sodium bromide.
(ii) What reaction would occur if a solution of sodium fluoride were used?
(i) Solid lithium oxide is added to water.
(ii) Name another metal whose oxide would react similarly.
— XXI —
(i) Methane reacts with an excess of chlorine gas.
(ii) What type of organic reaction occurs here?
(i) Hydrogen sulfide gas is bubbled into excess sodium hydroxide solution.
(ii) Name any spectator ion(s).
(i) Solutions of ammonia and carbon dioxide are mixed.
(ii) Name the acid that is formed as an intermediate.
(i) Solid copper(II)sulfide is added to a dilute solution of nitric acid.
(ii) How would the appearance of the reaction mixture change?
(i) Zinc is added to a solution of copper(II)sulfate.
(ii) Describe any visible changes to the zind metal.
— XXII —
(i) Solutions of potassium hydroxide and ammonium sulfate are mixed.
(ii) Name any spectator (ion)s.
(i) Ethanol reacts with methanoic acid.
(ii) Name this specific organic reaction.
(i) Dilute sulfuric acid is added to solid sodium fluoride.
(ii) Name the nonionic product.
(i) An electric current is passed through a solution of copper(II)sulfate.
(ii) What reaction occurs at the anode?
(i) Solid sodium acetate is mixed with dilute hydrochloric acid.
(ii) Name any spectator ion(s).
— XXIII —
(i) Formic acid is reacted with acidified potassium dichromate.
(ii) How would the color of the reaction mixture change?
(i) Sulfur dioxide gas is passed over solid calcium oxide.
(ii) Name the product.
(i) Ethene reacts with liquid bromine.
(ii) Name the product.
(i) Hydrochloric acid is given to a solution of dimercury(I)nitrate.
(ii) Name any spectator ion(s).
(i) Concentrated hydrobromic acid is heated with manganese dioxide.
(ii) How many electrons are transferred in this reaction?
67
116)
117)
118)
119)
120)
121)
122)
123)
124)
125)
126)
127)
128)
129)
130)
131)
132)
133)
134)
135)
— XXIV —
(i) Bromine is added to a dilute solution of sodium hydroxide.
(ii) How does the pH change during this reaction?
(i) Acetic acid is added to solid sodium hydrogencarbonate.
(ii) What gaseous product is formed?
(i) Solutions of sodium sulfide and zinc nitrate are combined.
(ii) Name this type of inorganic reaction.
(i) A basic solution of potassium permanganate is added to a solution of sodium sulfite.
(ii) How does the oxidation state of manganese change?
(i) A stoichiometric amount of sulfuric acid is added to a solution of lithium carbonate.
(ii) What product would be formed if the mole ratio of reactant were 1:1?
— XXV —
(i) Hydrogen sulfide gas is bubbled through a solution of nickel(II)nitrate.
(ii) How would the pH of the solution change?
(i) Magnesium is added to dilute nitric acid.
(ii) What acts as an oxidizing agent?
(i) Lead is added to a solution of silver nitrate.
(ii) What observable change would occur to the piece of lead?
(i) Borontrifluoride and ammonia gas are mixed.
(ii) What are the formal charges of B and N in the product?
(i) Propanol combusts in air.
(ii) If 25 g of propanol were used, what volume of products at STP would be formed,
assuming that all products are gaseous?
— XXVI —
(i) A solution of barium chloride is added to solid silver nitrate.
(ii) Name any spectator ion(s).
(i) Hydrogen gas is passed over hot iron(II)oxide.
(ii) What acts as an oxidizing agent?
(i) Hydrogen peroxide is mixed with an acidified solution of sodium bromide.
(ii) How many electrons are transferred during this reaction?
(i) Phosphorusoxytrichloride is added to an excess of potassium hydroxide solution.
(ii) Name this type of reaction.
(i) Rubidium is added to water.
(ii) How would the pH of the reaction change?
— XXVII —
(i) Solutions of iron(III)sulfate and tin(II)chloride are mixed.
(ii) How many electrons are transferred during this reaction?
(i) An excess of lauric acid CH3(CH2)10COOH reacts with glycerol.
(iii) Name the type of organic molecule formed.
(i) Solid sodium sulfite is added to an acidified solution of sodium permanganate.
(ii) What acts as a reducing agent?
(i) Sodium sulfite solution is added to hydrochloric acid.
(ii) Name the unstable intermediate in this reaction.
(i) Solid barium methanoate is added to dilute nitric acid.
(ii) Name any spectator ion(s).
68
136)
137)
138)
139)
140)
141)
142)
143)
144)
145)
146)
147)
148)
149)
150)
— XXVIII —
(i) Lithium reacts with nitrogen.
(ii) How many electrons are transferred during this reaction?
(i) Iron filings are boiled with a solution of iron(III)nitrate.
(ii) What substance is reduced?
(i) Water is added to solid sodium oxide.
(ii) What color would added litmus show?
(i) Water is added to solid sodium peroxide.
(ii) How would the pH of the reaction mixture change?
(i) Chlorine gas is bubbled through a solution of sodium iodide.
(ii) What different product would be formed if sodium bromide were used?
— XXIX —
(i) Ammonia is bubbled through dilute acetic acid.
(ii) Name any spectator ion(s).
(i) 1-Pentene reacts with hydrogen in the presence of a catalyst.
(ii) Name the product.
(i) Magnesium oxide is exposed to sulfur trioxide gas.
(ii) What species acts a base anhydride?
(i) Calcium is put in water.
(ii) Name Group 2 element that would show higher reactivity.
(i) Silver is dropped into a concentrated solution of nitric acid.
(ii) What acts as the oxidizing agent?
— XXX —
(i) Equal volumes of equimolar solutions of potassium hydroxide and sodium
dihydrogenphospate are mixed.
(ii) Name the product if stoichiometric amounts would be used.
(i) Excess dilute hydrochloric acid is added to a solution of sodium phosphate.
(ii) Name any spectator ion(s).
(i) Solutions of silver nitrate and lithium bromide are combined.
(ii) Name this type of inorganic reaction.
(i) Carbon disulfide is burned in an excess of fluorine.
(ii) What structure would you predict for the sulfur-containing product?
(i) Butyl ethanoate is boiled with potassium hydroxide.
(ii) Name the alcohol that is formed.
Notes:
69
HHS, AP Chemistry, 2012/2013
AP Equations (Answer Key)
Please use “Equations in the AP Examinations” and Appendices D and E in your textbook.
The answer key might not be free of errors.
Symbol
Reaction Type
SY
synthesis
DC
decomposition
DR
double replacement/precipitation
AS
anionic single replacement
CS
cationic single replacement
CO
combustion
RO
redox (other)
AB
acid/base
LE
ligand exchange
HY
hydrolysis
DH
dehydration
OA
OS
organic (addition to double or
triple bonds)
organic (substitution to alkanes
or aromatic compounds)
ET
ether reactions
ES
ester reactions
Examples
Li + O2  Li2O
Si + F2  SiF4
NaHCO3  Na2O + CO2 + H2O
CaSO4  CaO + SO3
OH- + Fe3+  Fe(OH)3
Ca2+ + PO43-  Ca3(PO4)2
Cl2 + Br-  Cl- + Br2
F2 + Cl-  F- + Cl2
Ca + Au3+  Ca2+ + Au
Al + Cu+  Al3+ + Cu
CH3OH + O2  CO2 + H2O
C8H18 + O2  CO2 + H2O
CrO42- + H2O2 + H+  Cr3+ + O2 + H2O
Sn4+ + Fe2+  Sn2+ + Fe3+
MgO + H+  Mg2+ + H2O
SO3 + OH-  SO42- + H2O
Cu(OH)2 + NH3  Cu(NH3)42+ + OHAgCl + CN-  Ag(CN)2- + ClNa3N + H2O  NH3 + Na+ + OHPCl5 + H2O  H+ + PO43- + ClFe(OH)3  Fe2O3 + H2O
H2SO3  H2O + SO2
C2H4 + Cl2  C2H4Cl2
C2H2 + Cl2  C2H2Cl2
CH4 + Br2  CH3Br + HBr
C6H6 + Br2  C6H5 Br + HBr
CH3OCH2CH3 + H2O
CH3OH + CH3CH2OH
CH3COOCH2CH3 + H2O
CH3COOH + CH3CH2OH
70
—I—
1)
(i) Solid lithium hydride is added to water.
(ii) How does the pH of the reaction mixture change?
LiH + H2O  Li+ + OH- + H2
AB
The pH would increase.
2)
(i) Excess ammonia is added to a solution of zinc sulfate.
(ii) What particles act as LEWIS acids and bases?
4 NH3 + Zn2+  Zn(NH3)42+
LE
LEWIS acid: Zn2+; LEWIS base: NH3
3)
(i) Sodium acetate is acidified with dilute hydrochloric acid.
(ii) How does the pH of the reaction mixture change?
CH3COO- + H+  CH3COOH
AB
The pH would increase.
4)
(i) A solution of potassium dichromate is added to an acidified solution of
sodium iodide.
(ii) What is total number of electrons transferred?
Cr2O72- + 14 H+ + 6 I-  2 Cr3+ + 3 I2 + 7 H2O
RO
Six electrons are transferred.
5)
(i) Solid barium carbonate is added to dilute nitric acid.
(ii) What is a simple analytic test for one of the products?
BaCO3 + 2 H+  Ba2+ + H2O + CO2
AB, DC
CO2 gas in calcium hydroxide solution precipitates CaCO3.
71
— II —
6)
(i) Solid aluminum sulfide is added to water.
(ii) Why is the gaseous product toxic?
Al2S3 + 6 H2O  2 Al(OH)3 + 3 H2S
AB/DR
7)
H2S is an acid and dissociates, forming sulfide ions S2- that can
precipitate Fe2+ in hemoglobin, thus inactivating it.
(i) Silver metal is added to 6M nitric acid.
(ii) What acts as an oxidizing agent?
3 Ag + 4 H+ + NO3-  3 Ag+ + NO + 2 H2O
RO
The nitrate ion NO3- acts an oxidizing agent.
8)
(i) A solution of tin(II)nitrate is dropped onto a piece of zinc metal.
(ii) What observation can be made about the piece of zinc?
Sn2+ + Zn  Sn4+ + Zn2+
RO
The zinc metal corrodes.
9)
(i) Solid sodium dichromate is added to an acidified solution of iron(II)
sulfate.
(ii) What is the total number of electrons transferred?
Na2Cr2O7 + 14 H+ + 6 Fe2+  2 Na+ + 2 Cr3+ + 6 Fe3+ + 7 H2O
RO
Six electrons are transferred.
10)
(i) Liquid titanium tetrachloride is given into excess water.
(ii) How does the pH of the reaction mixture change?
TiCl4 + 2 H2O  TiO2 + 4 H+ + 4 ClHY
The pH decreases.
72
— III —
11)
(i) Solid ammonium chloride is added to a concentrated solution of sodium
hydroxide.
(ii) What ion(s), if any, act as spectator ions?
NH4Cl + OH-  NH3 + H2O + ClAB
12)
Sodium ion Na+ is a spectator ion.
(i) Solid sodium hydrogencarbonate is added to concentrated sodium
hydroxide.
(ii) Sodium hydrogencarbonate reacts with hydrochloric acid as well.
Name this behavior.
NaHCO3 + OH-  Na+ + CO32- + H2O
AB
Amphoterism
13)
(i) Solid calcium sulfite is acidified.
(ii) Name one unstable reaction intermediate.
CaSO3 + 2 H+  Ca2+ + H2O + SO2
AB, DC
Sulfurous acid H2SO3
14)
(i) Liquid acetic acid and methanol are combined and warmed.
(ii) What specific organic reaction type occurs here?
CH3COOH + CH3OH  CH3COOCH3 + H2O
ES, DH
Esterification
15)
(i) Aqueous solutions of sodium iodate and sodium iodide are combined,
then acidified.
(ii) What specific type of redox reaction occurs here?
IO3- + 5 I- + 6 H+  3 I2 + 3 H2O
RO
Comproportionation
73
— IV —
16)
(i) Solid calcium phosphate is given to excess hydrobromic acid.
(ii) What phosphorus-containing product would be formed if hydrobromic
acid and calcium phosphate reacted in a mole ratio of 1:1?
Ca3(PO4)2 + 6 H+  3 Ca2+ + 2 H3PO4
AB
17)
Hydrogenphosphate ion HPO42-
(i) A piece of magnesium is dropped into hydrochloric acid.
(ii) How many electrons are transferred in this reaction?
Mg + 2 H+  Mg2+ + H2
CS
Two electrons are transferred.
18)
(i) A stream of carbon dioxide gas is bubbled through a suspension of
barium carbonate.
(ii) What observation can be made regarding the appearance of the
reaction mixture?
CO2 + H2O + BaCO3  Ba2+ + 2 HCO3AB
The precipitate dissolves.
19)
(i) Potassium permanganate solution is added to concentrated
hydrochloric acid.
(ii) How does the color of the reaction mixture change?
2 MnO4- + 16 H+ + 10 Cl-  2 Mn2+ + 5 Cl2 + 8 H2O
RO
The solution color changes from violet to colorless (faint pink).
20)
(i) Ammonia gas is bubbled through a suspension of zinc hydroxide.
(ii) What observation can be made regarding the appearance of the
reaction mixture?
4 NH3 + Zn(OH)2  Zn(NH3)42+ + 2 OHLE
The precipitate dissolves.
74
—V—
21)
(i) Hydrogen sulfide gas reacts with moist lead(II)chloride.
(ii) What color change could be observed?
H2S + Pb2+  PbS + 2 H+
AB, DR
The reaction mixture turns black.
22)
(i) Moist sodium hydroxide reacts with carbon dioxide gas.
(ii) How and why does the pH of the reaction mixture change?
2 OH- + CO2  H2O + CO32- (or: OH- + CO2  HCO3-)
AB
Carbon dioxide is an acid anhydride, therefore the pH decreases.
23)
(i) An aqueous solution of ammonia is added to solid silver chloride.
(ii) Name this type of reaction.
2 NH3 + AgCl  Ag(NH3)2+ + ClLE
Ligand-exchange reaction
24)
(i) Rubidium permanganate is added to hydroiodic acid.
(ii) How many electrons are transferred in this reaction?
2 MnO4- + 16 H+ + 10 I-  2 Mn2+ + 5 I2 + 8 H2O
RO
Ten electrons are transferred.
25)
(i) Liquid phosporus trichloride is poured into an excess of sodium
hydroxide.
(iii) What are the oxidation states of phosphorus in reactant(s) and
product(s)?
PCl3 + 6 OH-  PO33- + 3 Cl- + 3 H2O
HY, AB
The oxidation state is +3 in both reactant and product.
75
— VI —
26)
(i) 30% hydrogen peroxide is heated.
(ii) What are the oxidation states of oxygen in reactant(s) and product(s)?
2 H2O2  2 H2O + O2
RO/DC
Oxidation states of oxygen: H2O2: -1; H2O: -2; O2: 0
27)
(i) Liquid mercury is added to concentrated nitric acid.
(ii) What substance acts as a reducing agent?
3 Hg + 8 H+ + 2 NO3-  3 Hg2+ + 2 NO + 4 H2O
RO
Mercury Hg acts as reducing agent.
28)
(i) Solid iron(III)sulfate is added to a sodium iodide solution.
(ii) What substance could be used to detect one of the products in small
concentrations?
Fe2(SO4)3 + 2 I-  2 Fe2+ + I2 + 3 SO42RO
29)
Iodine I2 can be detected by starch; it forms a dark violet
solution.
(i) Hydrogen sulfide is bubbled through a solution of silver nitrate.
(ii) Name this specific type of reaction.
H2S + Ag+  Ag2S + 2 H+
AB, DR
Double replacement
30)
(i) Concentrated (15 M) ammonia solution is added in excess to a solution
of copper(II)nitrate.
(ii) Describe any observable color changes.
4 NH3 + Cu2+  Cu(NH3)42+
LE
The blue color intensifies during the reaction.
76
— VII —
31)
(i) Magnesium metal is added to dilute nitric acid, giving as one of the
products a compound in which the oxidation number of nitrogen is -3.
(ii) Name said product.
4 Mg + 10 H+ + NO3-  4 Mg2+ + NH4+ + 3 H2O
RO
32)
Ammonium ion NH4+
(i) Excess water is added to solid calcium hydride.
(ii) Why must calcium hydride be stored in a well-sealed container?
2 H2O + CaH2  Ca2+ + 2 OH- + H2
AB
33)
It could react with moisture in the air; the hydrogen could burst
the container.
(i) Silver acetate solution is dropped into a solution of sodium phosphate.
(ii) If the contents of the reaction mixture would be filtered, what
substance(s) would remain on the filter paper?
3 Ag+ + PO43-  Ag3PO4
DR
Silver phosphate
34)
(i) Solid sodium cyanide is added to water.
(ii) Why is sodium cyanide toxic?
NaCN + H2O  HCN + Na+ + OHAB
35)
Cyanide ions remove Fe2+ from hemoglobin as Fe(CN)64-,
therefore inactivating it.
(i) Solid potassium hydride is added to anhydrous ethyl alcohol.
(ii) Name the solid reaction product.
KH + CH3CH2OH  CH3CH2OK + H2
AB
Potassium ethanoate
77
— VIII —
36)
(i) Lithium metal is burned in air.
(ii) Given sufficiently high temperatures, what byproduct could be formed?
4 Li + O2  2 Li2O
SY
Lithium nitride Li3N
37)
(i) Aluminum metal is added to a solution of copper(II)chloride.
(ii) What, if any, color change would be observable?
2 Al + 3 Cu2+  2 Al3+ + 3 Cu
CS
The blue color of Cu2+ would fade over time; Al3+ is colorless.
38)
(i) Manganese(II)nitrate is mixed with sodium hydroxide solution.
(ii) Name any spectator ions.
Mn2+ + 2 OH-  Mn(OH)2
DR
Nitrate NO3- and sodium ion Na+
39)
(i) Equal volumes of dilute equimolar solutions of sodium carbonate and
hydrochloric acid are mixed.
(ii) What ratio of sodium carbonate to hydrochloric acid would lead to the
formation of a gas?
CO32- + H+  HCO3-
AB, DC
(sodium carbonate) : (hydrochloric acid) = 2:1
40)
(i) Solid sodium carbide is added to an excess of water.
(ii) How does the pH of the reaction mixture change?
Na2C2 + 2 H2O  2 Na+ + 2 OH- + C2H2
AB
The pH increases.
78
— IX —
41)
(i) An excess of sodium hydroxide solution is added to a solution of
aluminum chloride.
(ii) Describe any visible changes during the reaction.
6 OH- + Al3+  Al(OH)63AB
42)
An initial precipitate of Al(OH)3 would dissolve when further OHis added.
(i) Solid ammonium carbonate is heated.
(ii) Which of the products has the lowest boiling point?
(NH4)2CO3  2 NH3 + H2O + CO2
DC
Carbon dioxide CO2
43)
(i) Phosphorus pentachloride is added to an excess of potassium
hydroxide solution.
(ii) Name the type of reaction that the phosphorus pentachloride
undergoes.
PCl5 + 8 OH-  PO43- + 5 Cl- + 4 H2O
HY, AB
Hydrolysis
44)
(i) Potassium chlorate is heated in the presence of manganese dioxide.
(ii) What function does the manganese dioxide have?
2 KClO3  2 KCl + 3 O2
DC
Manganese dioxide acts as a catalyst.
45)
(i) Chlorine gas is bubbled into a solution of sodium hydroxide.
(ii) What oxidation state(s) change(s)?
Cl2 + 2 OH-  Cl- + OCl- + H2O
RO, AB
Chlorine: 0  -1 and +1
79
—X—
46)
(i) Magnesium carbonate is heated strongly.
(ii) Which of the product is an acid anhydride?
MgCO3  MgO + CO2
DC
Carbon dioxide CO2 is an acid anhydride.
47)
(i) Methanol and propanoic acid are mixed.
(ii) What general organic reaction takes place here?
CH3OH + CH3CH2COOH  CH3CH2COOCH3 + H2O
ES, DH
Dehydration synthesis (condensation)
48)
(i) Barium acetate is added to dilute sulfuric acid.
(ii) What, if any, spectator ions occur here?
Ba2+ + 2 CH3COO- + 2 H+ + SO42-  2 CH3COOH + BaSO4
AB, DR
There are no spectator ions.
49)
(i) Ammonia gas is bubbled through a suspension of silver chloride.
(ii) Describe any visible changes of the reaction mixture.
2 NH3 + AgCl  Ag(NH3)2+ + ClLE
The precipitate of AgCl would dissolve.
50)
(i) Hot iron(III)oxide is reacted with carbon monoxide.
(ii) How many electrons are transferred during the reaction?
Fe2O3 + 3 CO  2 Fe + 3 CO2
RO
Six electrons are transferred.
80
— XI —
51)
(i) Silane (SiH4) is burned in air.
(ii) What type of solid is the product containing silicon?
SiH4 + 2 O2  SiO2 + 2 H2O
RO, CO
Silicon dioxide SiO2 is a network solid.
52)
(i) Cesium metal is heated with sulfur.
(ii) Name this type of reaction.
16 Cs + S8  8 Cs2S
SY
Synthesis
53)
(i) Sodium phosphate is dissolved in water.
(ii) How does the pH of the reaction mixture change?
Na3PO4 + H2O  HPO42- + 3 Na+ + OHAB
The pH increases.
54)
(i) Ethyl ethanoate is boiled with sodium hydroxide.
(ii) What type of organic compound is ethyl ethanoate?
CH3COOCH2CH3 + OH-  CH3COO- + CH3CH2OH
ES, HY
Ethyl ethanoate is an ester.
55)
(i) Solutions of sodium phosphate and calcium chloride are combined.
(ii) Name any spectator ion(s).
2 PO43- + 3 Ca2+  Ca3(PO4)2
DR
Sodium Na+ and chloride ion Cl-
81
— XII —
56)
(i) A suspension of aluminum hydroxide is added to hydroiodic acid.
(ii) How would the reaction be different if hydrochloric acid were used?
Al(OH)3 + 3 H+  Al3+ + 3 H2O
AB
The reaction would be the same.
57)
(i) Hydrogen peroxide is added to an acidified solution of sodium iodide.
(ii) How many electrons are transferred during this reaction?
H2O2 + 2 H+ + 2 I-  2 H2O + I2
RO
Two electrons are transferred.
58)
(i) Hydrogen peroxide is added to an acidified solution of potassium
dichromate.
(ii) What are the oxidation states of chromium in this reaction?
3 H2O2 + 8 H+ + Cr2O72-  3 O2 + 7 H2O + 2 Cr3+
RO
The oxidation states of chromium are +6 (Cr2O72-) and +3 (Cr3+).
59)
(i) Calcium is added to dilute hydrochloric acid.
(ii) What specific inorganic reaction type occurs here?
Ca + 2 H+  Ca2+ + H2
CS
Cationic single replacement
60)
(i) Aluminum nitrate is dissolved in water.
(ii) Which substance acts as LEWIS acid here?
Al(H2O)63+  [Al(H2O)5(OH)]- + H+
AB
Aluminum ion Al3+
82
— XIII —
61)
(i) Phosphorus tribromide is added to water.
(ii) Name the product containing phosporus.
PBr3 + 3 H2O  PO33- + 3 Br- + 6 H+
HY
Phosphite ion PO33-
62)
(i) Solutions of sulfuric acid and lithium hydroxide are mixed.
(ii) What specific analytical method would use this type of reaction?
H+ + OH-  H2O
AB
Acid-base titration
63)
(i) Nitric acid is added to a solution of tetraamminecadmium(II).
(ii) What spectator ion(s), if any, occur in this reaction?
4 H+ + Cd(NH3)42+  Cd2+ + 4 NH4+
AB, LE
Nitrate ion NO3-
64)
(i) 2-butene reacts with hydrobromic acid.
(ii) What functional group of 2-butene is active in this reaction?
CH3CH=CHCH3 + H+ + Br-  CH3CHBr—CH2CH3
OA
The functional group is the C=C double bond.
65)
(i) Hydrogensulfide is bubbled through a solution of lead(II)nitrate.
(ii) What color change could be observed?
H2S + Pb2+  2 H+ + PbS
AB, DR
The solution would turn black.
83
— XIV —
66)
(i) Solutions of iron(III)nitrate and sodium thiocyanate are combined.
(ii) What color change could be observed?
Fe(H2O)63+ + SCN-  [Fe(SCN)(H2O)5]2+ + H2O
LE/DR
The solution would turn red.
67)
(i) Ammonia gas is bubbled through a suspension of zinc hydroxide.
(ii) How many electrons are transferred during this reaction?
4 NH3 + Zn(OH)2  Zn(NH3)42+ + 2 OHLE
No electrons would be transferred.
68)
(i) Sulfur dioxide gas is bubbled through an acidified solution of potassium
permanganate.
(ii) How many electrons are transferred during this reaction?
5 SO2 + 2 H2O + 2 MnO4-  5 SO42- + 4 H+ + 2 Mn2+
RO
Ten electrons are transferred.
69)
(i) Aluminum is added to excess sodium hydroxide solution.
(ii) Name the ionic product of this reaction.
2 Al + 6 OH- + 6 H2O  2 Al(OH)63- + 3 H2
RO, LE
Hexahydroxoaluminate ion
70)
(i) A solution of tin(II)nitrate is added to an acidified solution of potassium
dichromate.
(ii) How many electrons are transferred during this reaction?
3 Sn2+ + 14 H+ + Cr2O72-  3 Sn4+ + 2 Cr3+ + 7 H2O
RO
Six electrons are transferred.
84
— XV —
71)
(i) Equal volumes of equimolar sodium hydroxide and sodium
dihydrogenphosphate are mixed.
(ii) What phosphorus-containing product would be formed if the ratio of
hydroxide to dihydrogenphosphate were 2:1?
OH- + H2PO4-  HPO42- + H2O
AB
72)
The phosphate ion PO43-
(i) Excess sodium hydroxide is added to a solution of calcium
hydrogencarbonate.
(ii) Name any spectator ions.
OH- + HCO3- + Ca2+  H2O + CaCO3
AB, DR
73)
Sodium ion Na+
(i) Acetylene (ethyne) reacts with an excess of chlorine.
(ii) What product would be formed if the reactants were used in a ratio of
1:1?
HCCH + Cl2  CHCl2—CHCl2
OA
74)
1,2-dichloroethene ClCH2—CH2Cl
(i) Solid silver nitrate is added to a solution of sodium chromate.
(ii) Name any spectator ions.
DR
2 AgNO3 + CrO42-  Ag2CrO4 + 2 NO3Sodium ion Na+
75)
(i) An excess of solid potassium hydroxide is added to a suspension of
zinc hydroxide.
(ii) What changes in the reaction mixture can be observed?
2 KOH + Zn(OH)2  2 K+ + Zn(OH)42LE
The precipitate of zinc hydroxide would dissolve.
85
— XVI —
76)
(i) Benzene reacts with bromine.
(ii) Why does benzene not react as if it had double bonds?
C6H6 + Br2  C6H5Br + HBr
OS
77)
Benzene does not react as if it had double bond due to
resonance stabilization.
(i) A direct electric current is passed through a dilute solution of sulfuric
acid.
(ii) What gas is formed at the cathode?
2 H2O  2 H2 + O2
RO
Hydrogen gas is formed at the cathode.
78)
(i) Ammonia and oxygen are heated in the presence of a catalyst.
(ii) How does the oxidation number for nitrogen change?
4 NH3 + 5 O2  4 NO + 6 H2O
RO
The oxidation states of nitrogen are -3 (NH3) and +2 (NO).
79)
(i) Liquid bromine is added to a solution of sodium iodide.
(ii) What specific type of inorganic reaction occurs here?
Br2 + 2 I-  2 Br- + I2
AS
Anionic single replacement
80)
(i) Dilute hydrochloric acid is added to solid calcium oxide.
(ii) Name the spectator ion.
2 H+ + CaO  Ca2+ + H2O
AB
Chloride Cl-
86
— XVII —
81)
(i) Sodium hydride is added to water.
(ii) Compare the relative reactivity of potassium hydride with sodium
hydride.
NaH + H2O  Na+ + OH- + H2
AB
Potassium hydride is more reactive than sodium hydride.
82)
(i) Ethene reacts with water in the presence of a catalyst.
(ii) Name the product.
CH2=CH2 + H2O  CH3—CH2OH
OA
Ethanol
83)
(i) Solutions of barium hydroxide and iron(III)sulfate are combined.
(ii) If equal moles of reactants were combined, which would be the limiting
reactant?
3 Ba2+ + 6 OH- + 2 Fe3+ + 3 SO42-  3 BaSO4 + 2 Fe(OH)3
DR
The limiting reactant would be barium hydroxide Ba(OH)2.
84)
(i) Ammonia gas is bubbled through a solution of copper(II)nitrate.
(ii) What color change could be observed?
4 NH3 + Cu2+  Cu(NH3)42+
LE
The blue color intensifies during the reaction.
85)
(i) Lead is added to a dilute solution of sulfuric acid.
(ii) How many electrons are transferred during this reaction?
Pb + 2 H+ + SO42-  PbSO4 + H2
CS, DR
Two electrons are transferred.
87
— XVIII —
86)
(i) Lead is added to a hot solution of concentrated sulfuric acid.
(ii) What acts as the oxidizing agent?
Pb + 4 H+ + 2 SO42-  PbSO4 + SO2 + 2 H2O
RO, DR
The oxidizing agent is the sulfate ion SO42-.
87)
(i) Calcium oxide is exposed to an atmosphere of carbon dioxide.
(ii) Which reactant is an acid anhydride?
CaO + CO2  CaCO3
SY
The acid anhydride is carbon dioxide CO2.
88)
(i) Nitrogen(V)oxide is bubbled through water.
(ii) How does the pH of this reaction change?
N2O5 + H2O  2 H+ + 2 NO3AB
The pH decreases.
89)
(i) Solid sodium hydrogencarbonate is added to water.
(ii) How does the pH of this reaction change?
NaHCO3  Na+ + CO2 + OHAB
The pH increases.
90)
(i) Solid sodium hydroxide and solid ammonium chloride are mixed and
heated.
(ii) Name the gaseous product.
NaOH + NH4Cl  NaCl + NH3 + H2O
AB
Ammonia NH3
88
— XIX —
91)
(i) Iodomethane is heated with a solution of sodium hydroxide.
(ii) Compare the reactivity of iodomethane with chloromethane.
CH3I + OH-  CH3OH +IOS
Chloromethane is less reactive than iodomethane.
92)
(i) Solid barium hydroxide and solid ammonium sulfate are mixed and
heated.
(ii) What product has the highest melting point?
Ba2+ + 2 OH- + 2 NH4+ + SO42-  BaSO4 + 2 NH3 + 2 H2O
DR, AB
Barium sulfate BaSO4 has the highest melting point.
93)
(i) A solution of sodium oxalate is added to an acidified solution of
potassium permanganate.
(ii) How many electrons are transferred during this reaction?
5 -OOC—COO- + 16 H+ + 2 MnO4-  10 CO2 + 2 Mn2+ + 8 H2O
RO
Ten electrons are transferred.
94)
(i) Equal volumes of equimolar hydrochloric acid are added to a solution of
sodium hydrogenphosphate.
(iii) What would be the stoichiometric ratio of the reactants?
H+ + HPO42-  H2PO4AB
95)
The stoichiometric ratio would be
(hydrochloric acid) : (hydrogenphosphate) = 2:1
(i) Ethanol burns in air.
(ii) Name a reagent that would detect the heavier of the products.
CH3CH2OH + O2  CO2 + H2O
CO
Carbon dioxide would precipitate calcium carbonate out of a
solution of calcium hydroxide Ca(OH)2
— XX —
89
96)
(i) A strip of magnesium is put in a solution of iron(III)nitrate.
(ii) What is reduced in this reaction?
Mg + 2 Fe3+  Mg2+ + 2 Fe2+
CS/RO
The iron (III) ion would be reduced.
97)
(i) Hydrogen peroxide is heated.
(ii) What changes of oxidation state(s) occur here?
2 H2O2  2 H2O + O2
DC
Oxidation states of oxygen: -1 (H2O2) to -2 (H2O) and 0 (O2)
98)
(i) Iron filings are sprinkled into a solution of iron(III)chloride.
(ii) What substance is oxidized?
Fe + 2 Fe3+  3 Fe2+
RO
The iron Fe is oxidized.
99)
(i) Chlorine is bubbled through a solution of sodium bromide.
(ii) What reaction would occur if a solution of sodium fluoride were used?
Cl2 + 2 Br-  2 Cl- + Br2
AS
No reaction would occur.
100) (i) Solid lithium oxide is added to water.
(ii) Name another metal whose oxide would react similarly.
Li2O + H2O  2 Li+ + 2 OHAB
Any other Group 1 metal (Na, K, Rb, Cs)
90
— XXI —
101) (i) Methane reacts with an excess of chlorine gas.
(ii) What type of organic reaction occurs here?
CH4 + 4 Cl2  CCl4 + 4 HCl
OS
(Aliphatic) Substitution
102) (i) Hydrogen sulfide gas is bubbled into excess sodium hydroxide solution.
(ii) Name any spectator ion(s).
H2S + 2 OH-  S2- + 2 H2O
AB
Sodium ion Na+
103) (i) Solutions of ammonia and carbon dioxide are mixed.
(ii) Name the acid that is formed as an intermediate.
2 NH3 + CO2 + H2O  2 NH4+ + CO32AB/SY
Carbonic acid H2CO3
104) (i) Solid copper(II)sulfide is added to a dilute solution of nitric acid.
(ii) How would the appearance of the reaction mixture change?
CuS + 2 H+  Cu2+ + H2S
AB
The black CuS precipitate would disappear and the solution
would turn blue.
105) (i) Zinc is added to a solution of copper(II)sulfate.
(ii) Describe any visible changes to the zinc metal.
Zn + Cu2+  Zn2+ + Cu
CS
The zinc would corrode.
91
— XXII —
106) (i) Solutions of potassium hydroxide and ammonium sulfate are mixed.
(ii) Name any spectator (ion)s.
OH- + NH4+  NH3 + H2O
AB
Potassium ion K+ and sulfate SO42-
107) (i) Ethanol reacts with methanoic acid.
(ii) Name this specific organic reaction.
CH3CH2OH + HCOOH  HCOOCH2CH3 + H2O
ES, DH
Esterification
108) (i) Dilute sulfuric acid is added to solid sodium fluoride.
(ii) Name the nonionic product.
2 H+ + SO42- + CaF2  2 HF + CaSO4
AB, DR
Hydrogen fluoride
109) (i) An electric current is passed through a solution of copper(II)sulfate.
(ii) What reaction occurs at the anode?
2 Cu2+ + 2 OH-  2 Cu + O2 + 2 H+
RO
The oxidation at the anode is 2OH-  O2 + 4 e- + 2 H+
110) (i) Solid sodium acetate is mixed with dilute hydrochloric acid.
(ii) Name any spectator ion(s).
NaCH3COO + H+  Na+ + CH3COOH
AB
Chloride Cl-
92
— XXIII —
111) (i) Formic acid is reacted with acidified potassium dichromate.
(ii) How would the color of the reaction mixture change?
3 HCOOH + 8 H+ + Cr2O72-  3 CO2 + 2 Cr3+ + 7 H2O
RO
The color would change from orange to green.
112) (i) Sulfur dioxide gas is passed over solid calcium oxide.
(ii) Name the product.
SO2 + CaO  CaSO3
SY
Calcium sulfite
113) (i) Ethene reacts with liquid bromine.
(ii) Name the product.
CH2=CH2 + Br2  CH2Br—CH2Br
OA
1,2-dibromoethane
114) (i) Hydrochloric acid is given to a solution of dimercury(I)nitrate.
(ii) Name any spectator ion(s).
Cl- + Hg22+  Hg2Cl2
DR
Hydrogen ion H+ and nitrate NO3-
115) (i) Concentrated hydrobromic acid is heated with manganese dioxide.
(ii) How many electrons are transferred in this reaction?
4 H+ + 2 Br- + MnO2  Br2 + Mn2+ + 2 H2O
RO
Two electrons are transferred.
93
— XXIV —
116) (i) Bromine is added to a dilute solution of sodium hydroxide.
(ii) How does the pH change during this reaction?
Br2 + 2 OH-  Br- + OBr- + H2O
RO, AB
The pH decreases.
117) (i) Acetic acid is added to solid sodium hydrogencarbonate.
(ii) What gaseous product is formed?
CH3COOH + NaHCO3  Na+ + CH3COO- + CO2 + H2O
AB, DC
Carbon dioxide CO2
118) (i) Solutions of sodium sulfide and zinc nitrate are combined.
(ii) Name this type of inorganic reaction.
S2- + Zn2+  ZnS
DR
Double replacement
119) (i) A basic solution of potassium permanganate is added to a solution of
sodium sulfite.
(ii) How does the oxidation state of manganese change?
H2O + 2 MnO4- + 3 SO32-  2 MnO2 + 2 OH- + 3 SO42RO
Oxidation states of manganese: +7 (MnO4-) and +4 (MnO2)
120) (i) A stoichiometric amount of sulfuric acid is added to a solution of lithium
carbonate.
(ii) What product would be formed if the mole ratio of reactant were 1:1?
2 H+ + CO3-  H2O + CO2
AB, DC
Hydrogencarbonate HCO3-
94
— XXV —
121) (i) Hydrogen sulfide gas is bubbled through a solution of nickel(II)nitrate.
(ii) How would the pH of the solution change?
H2S + Ni2+  NiS + 2 H+
AB, DR
The pH would decrease.
122) (i) Magnesium is added to dilute nitric acid.
(ii) What acts as an oxidizing agent?
Mg + 2 H+  Mg2+ + H2
RO
Hydrogen ion H+
123) (i) Lead is added to a solution of silver nitrate.
(ii) What observable change would occur to the piece of lead?
Pb + 2 Ag+  Pb2+ + 2 Ag
CS
The lead would corrode.
124) (i) Borontrifluoride and ammonia gas are mixed.
(ii) What are the formal charges of B and N in the product?
BF3 + NH3  F3B—NH3
(addition)
The formal charges would be +1 (N) and -1 (B).
125) (i) Propanol combusts in air.
(ii) If 25.0 g of propanol were used, what volume of products at STP would
be formed, assuming that all products are gaseous?
2 CH3CH2CH2OH + 9 O2  6 CO2 + 8 H2O
CO
The total product volume would be 65.3 liters.
95
— XXVI —
126) (i) A solution of barium chloride is added to solid silver nitrate.
(ii) Name any spectator ion(s).
Cl- + Ag+  AgCl
DR
Barium ion Ba2+ and nitrate NO3-
127) (i) Hydrogen gas is passed over hot iron(II)oxide.
(ii) What acts as an oxidizing agent?
H2 + FeO  H2O + Fe
CS
The oxidizing agent is Fe2+ in FeO.
128) (i) Hydrogen peroxide is mixed with an acidified solution of sodium
bromide.
(ii) How many electrons are transferred during this reaction?
H2O2 + 2 H+ + 2 Br-  2 H2O + Br2
RO
Two electrons are transferred.
129) (i) Phosphorusoxytrichloride is added to an excess of potassium hydroxide
solution.
(ii) Name this type of reaction.
POCl3 + 6 OH-  PO43- + 3 Cl- + 3 H2O
HY, AB
Hydrolysis and subsequent deprotonation
130) (i) Rubidium is added to water.
(ii) How would the pH of the reaction change?
2 Rb + 2 H2O  2 Rb+ + 2 OH- + H2
RO
The pH would increase.
96
— XXVII —
131) (i) Solutions of iron(III)sulfate and tin(II)chloride are mixed.
(ii) How many electrons are transferred during this reaction?
2 Fe3+ + Sn2+  2 Fe2+ + Sn4+
RO
Two electrons are transferred.
132) (i) An excess of lauric acid CH3(CH2)10COOH reacts with glycerol.
(iii) Name the type of organic molecule formed.
3 CH3(CH2)10COOH + C3H5(OH)3  C3H5[OOC(CH2)10CH3]3 + 3 H2O
ES, DH
Triglyceride
133) (i) Solid sodium sulfite is added to an acidified solution of sodium
permanganate.
(ii) What acts as a reducing agent?
5 Na2SO3 + 6 H+ + 2 MnO4-  10 Na+ + 5 SO42- + 2 Mn2+ + 3 H2O
RO
The sulfite ion SO32- is the reducing agent.
134) (i) Sodium sulfite solution is added to hydrochloric acid.
(ii) Name the unstable intermediate in this reaction.
SO32- + 2 H+  H2O + SO2
AB, DC
Sulfurous acid H2SO3
135) (i) Solid barium methanoate is added to dilute nitric acid.
(ii) Name any spectator ion(s).
Ba(CH3O)2 + 2 H+  Ba2+ + CH3OH
AB
Nitrate NO3-
97
— XXVIII —
136) (i) Lithium reacts with nitrogen.
(ii) How many electrons are transferred during this reaction?
6 Li + N2  2 Li3N
SY
Six electrons are transferred.
137) (i) Iron filings are boiled with a solution of iron(III)nitrate.
(ii) What substance is reduced?
Fe + 2 Fe3+  3 Fe2+
RO
The iron (III) ion Fe3+ is reduced.
138) (i) Water is added to solid sodium oxide.
(ii) What color would added litmus show?
H2O + Na2O  2 Na+ + 2 OHAB
Litmus would turn blue.
139) (i) Water is added to solid sodium peroxide.
(ii) How would the pH of the reaction mixture change?
2 H2O + Na2O2  2 Na+ + 2 OH- + H2O2
AB
The pH would increase.
140) (i) Chlorine gas is bubbled through a solution of sodium iodide.
(ii) What different product would be formed if sodium bromide were used?
Cl2 + 2 I-  2 Cl- + I2
AS
Bromine Br2
98
— XXIX —
141) (i) Ammonia is bubbled through dilute acetic acid.
(ii) Name any spectator ion(s).
CH3COOH + NH3  CH3COO- + NH4+
AB
There are no spectator ions.
142) (i) 1-Pentene reacts with hydrogen in the presence of a catalyst.
(ii) Name the product.
CH2=CHCH2CH2CH3 + H2  CH3CH2CH2CH2CH3
OA
Pentane
143) (i) Magnesium oxide is exposed to sulfur trioxide gas.
(ii) What species acts a base anhydride?
MgO + SO3  MgSO4
AB
The base anhydride is magnesium oxide MgO.
144) (i) Calcium is put in water.
(ii) Name a Group 2 element that would show higher reactivity.
Ca + 2 H2O  Ca2+ + 2 OH- + H2
RO
Strontium Sr, barium Ba, or radium Ra
145) (i) Silver is dropped into a concentrated solution of nitric acid.
(ii) What acts as the oxidizing agent?
3 Ag + 4 H+ + NO3-  3 Ag+ + NO + 2 H2O
RO
The oxidizing agent is the nitrate ion NO3-.
99
— XXX —
146) (i) Equal volumes of equimolar solutions of potassium hydroxide and
sodium dihydrogenphospate are mixed.
(ii) Name the product if stoichiometric amounts would be used.
OH- + H2PO4-  HPO42- + H2O
AB
Phosphoric acid H3PO4
147) (i) Excess dilute hydrochloric acid is added to a solution of sodium
phosphate.
(ii) Name any spectator ion(s).
H+ + PO43-  H3PO4
AB
Chloride Cl- and sodium ion Na+
148) (i) Solutions of silver nitrate and lithium bromide are combined.
(ii) Name this type of inorganic reaction.
Ag+ + Br-  AgBr
DR
Double replacement
149) (i) Carbon disulfide is burned in an excess of fluorine.
(ii) What structure would you predict for the sulfur-containing product?
CS2 + 8 F2  CF4 + 2 SF6
RO
Sulfur hexafluoride SF6 has an octahedral structure.
150) (i) Butyl ethanoate is boiled with potassium hydroxide.
(ii) Name the alcohol that is formed.
CH3COOCH2CH2CH2CH3 + OH-  CH3COO- + CH3CH2CH2CH2OH
ES, HY
1-butanol
100
APPENDIX B:
SOLUBILITY RULES FOR IONIC COMPOUNDS
The following cations are considered:
1)
2)
3)
4)
5)
6)
7)
8)
alkali metals (Li+, Na+, etc.) and NH4+ (similar reactivity)
alkaline earth metals (Be2+, Mg2+, Ca2+, etc.)
Al3+
Sn2+, Pb2+
Cr3+
Mn2+
Fe2+, Fe3+, Co2+, Ni2+
Cu2+, Ag+, Zn2+, Cd2+, Hg22+ (= "Hg+"), Hg2+
There is no completely insoluble compound; they are often classified as "mostly
soluble" or "mostly (nearly) insoluble." For brevity I will use “soluble” and “insoluble.”
Soluble salts with the following anions:
NO3-, C2H3O2-, ClO3-, ClO4-, FCl-
(except: Ag+, Hg22+, Pb2+)
Br-
(except: Ag+, Hg22+, Hg2+, Pb2+)
I-
(except: Ag+, Hg22+, Hg2+, Pb2+)
SO42- (except: Ca2+, Sr2+, Ba2+, Hg22+, Hg2+, Pb2+)
Insoluble salts with the following anions
S2-
(except: alkali and alkaline earth , NH4+)
CO32- (except: alkali, NH4+)
SO32- (except: alkali, NH4+)
PO43- (except: alkali, NH4+)
OH-
(except: alkali, Ba2+, Sr2+, Ca2+)
NH4+ decomposes: NH4+ + OH-  NH3 + H2O
O2-
(except: alkali, Ba2+, Sr2+, Ca2+)
NH4+ decomposes: NH4+ + O2-  NH3 + OHReaction with water: O2- + H2O  2 OH-
101
Notes:
102
APPENDIX C:
NOMENCLATURE OF INORGANIC COMPOUNDS
1)
Binary Acids:
Binary acids contain only one hydrogen and one more element. The
names of acids and corresponding ions are derived from the element
name.
Formula
Acid: “hydro” + (element) + “ic”
Anion: (element) + “ide”
HCl
hydrochloric acid
Cl-
chloride
HBr
hydrobromic acid
Br-
bromide
H2S
hydrosulfuric acid
S2(HS-
sulfide
hydrogensulfide)
analogous names: F- fluoride, N3- nitride, P3- phosphide, O2- oxide, etc.
2)
Ternary Acids:
Ternary acids contain hydrogen and two more elements, one of them
usually being oxygen ( oxyacids). The amount of oxygen present (i.e.,
the oxidation state of the third element) is denoted by various pre- and
suffixes.
Group 17
Formula
Name
Formula
Name
HClO4
HClO3
HClO2
HClO
perchloric
acid
chloric
acid
chlorous acid
hypochlorous acid
ClO4ClO3ClO2ClO-
perchlorate
chlorate
chlorite
hypochlorite
(analogous: Br, I)
Fluorine forms only hypofluorous acid HOF.
103
Group 16
Formula
Name
Formula
Name
H2SO4
sulfuric acid
SO42(HSO4-
sulfate
hydrogensulfate)
H2SO3
sulfurous acid
SO32(HSO3-
sulfite
hydrogensulfite)
analogous: Se, Te
Group 15
Formula
Name
Formula
Name
H3PO4
H3PO3
phosphoric acid
phosphorous acid
PO43PO33-
phosphate
phosphite
H3AsO5
H3AsO4
perarsenic acid
arsenic acid
AsO53AsO53-
perarsenate
arsenate
CO32BO33SiO44-
carbonate
borate
silicate
Some elements form only one type of oxyacid:
H2CO3
H3BO3
H4SiO4
carbonic acid
boric acid
silicic acid
In general, salts of elements in the same group follow the same pattern (formula
and name). This is not true for carbon (no H4CO4) and nitrogen (no H3NO4).
And some ions do not follow these “rules” at all:
MnO4RuO43)
permanganate
perruthenate
but
but
MnO42RuO42-
manganate
ruthenate
Molecular or ionic compounds
Using prefixes (Appendix E) or oxidation states, the nomenclature is
similar for both.
The cationic (less electronegative) particle/ion is written first.
Notes:
104
APPENDIX D:
NOMENCLATURE OF ORGANIC COMPOUNDS
Alkanes
C—C
CnH2n+2
CH4 methane
C2H6 ethane
C3H8 propane
C4H10 butane
C5H12 pentane
(cont. with Greek prefixes)
Alkenes
C==C
CnH2n
C2H4 ethene
C3H6 propene etc.
Alkynes
CC
CnH2n-2
C2H2 ethyne (acetylene)
C3H4 propyne etc.
Nomenclature:
1) find longest continuous chain  name for skeleton
2) branches = alkyl groups (methyl, ethyl, propyl, etc.) or other functional groups
3) position alkyl groups to minimize numbers indicating their position
4) more than one substituent of a kind: di, tri, tetra, penta, etc.
The nature of functional groups is given by using specific pre- and suffixes:
Compound
halides
alcohols
aldehydes
ketones
ethers
acid
esters
amines
amides
Notes:
Functional group
-F, -Cl, -Br, -I
R-OH
R-CHO
R1-CO-R2
R1-O-R2
R-COOH
R1-COO-R2
R-NH2
R-CO-NH2
Pre-/suffix
“fluoro-“, “chloro-“, “bromo-“, “iodo-“
“-ol”
“-al”
“-one”
“(Rest 1)(Rest 2) ether”
(carbon chain) + “-oic acid”
(Rest 2) (Rest 1) + “oate”
“amino-“
“-amide”
105
APPENDIX E:
MISCELLANEOUS
SI units: Basic units are by definition and use an arbitrary standard.
Basic quantity
Name
Symbol/Unit
Mass
kilogram
kg
Length
meter
m
Time
second
s
Electric current
Ampere
A
Luminous intensity
Candela
Cd
Amount of matter
Mole
mol
Derived quantity
Name
Symbol/Unit
Force
Newton
1 N = 1 kg m s-2
work, energy
Joule
1 J = 1 N m = 1 kg m2 s-2
Pressure
Pascal
1 Pa = 1 N m-2 = 1 kg m-1 s-2
Density
––
g cm-3
106
Constants and other assorted data
Constant
Atomic mass unit
Symbol
amu
Numerical Value
1.66055 x
10-24 g
Application
Stoichiometry
Avogadro’s Number
NA
Boltzmann’s constant
k
1.381 x 10-23 J mol-1 K-1
Thermodynamics,
entropy
Gas Constant
R
8.314 J mol-1 K-1
0.082 l atm mol-1 K-1
Thermodynamics,
gas problems
Faraday constant
F
96,000 C mol-1
Electrochemistry
6.0221367 x
10-23
10-31
mol-1
electron mass
me
9.110 x
proton mass
mp
1.673 x 10-27 kg
neutron mass
mn
1.675 x 10-27 kg
Planck’s constant
h
6.626 x 10-34 J s-1
108
Stoichiometry
kg
m
s-1
Quantum mechanics
speed of light
c
2.9979 x
Quantum mechanics
Standard Pressure

101.3 kPa = 1 atm
Thermodynamics,
gas problems
Standard Temperature

273 K = 0 oC
Thermodynamics,
gas problems
107
Prefixes
Fractions and multiples of 10
Giga
G
Mega
M
Kilo
k
Hecto
h
deci
d
milli
m
micro

nano
n
pico
p
Numerical prefixes from 1 to 10
Mono1
Di2
Tri3
Tetra4
Penta5
Hexa6
Hepta7
Octa8
Nona9
Deca10
Logarithms
log 0
log 1
log 2
log 3
log 4
log 5
log 6
log 7
log 8
log 9
=
=
=
=
=
=
=
=
=
=
(not defined)
0
0.3
0.5
0.6
0.7
0.8
0.85
0.9
0.95
Logarithmic rules
 log (ab) = log a + log b

log (a/b) = log a – log b

log (ab) = b log a
109
106
103
102
10-1
10-3
10-6
10-9
10-12
1,000,000,000
1,000,000
1,000
100
0.1
0.001
0.000 001
0.000 000 001
0.000 000 000 001
108
Notes:
109
APPENDIX F:
HELPFUL WEBSITES

Collegeboard — Advanced Placement Courses (general information)
http://apcentral.collegeboard.com/

Common misconceptions in chemistry
http://www.princeton.edu/~lehmann/BadChemistry.html

Various chemistry links (might be too much)
http://www.liv.ac.uk/Chemistry/Links/links.html

History of chemistry (biographies of chemists)
http://www.woodrow.org/teachers/ci/1992/

Nomenclature of oxyacids
http://chem01.usca.sc.edu/chemistry/genchem/nomen.htm

Results that any search engine gives you for the search term “AP
Chemistry”
Notes: