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AP CHEMISTRY Summer Assignment
AP Chemistry is a difficult course. It is not all about memorization; however,
having these items memorized is essential for success in learning the concepts covered in
the course. Make flashcards, have your friends and family quiz you, take the lists with
you on vacation, or do whatever it takes to get this information firmly planted in your
head. Do not wait until the night before school begins.
The first day test will cover six areas of memorization:
1. Polyatomic Ions (including name, symbol and charge)
2. Variable Valences (Charge) for Transition Metals
3. Rules for Naming Acids
4. Rules for Naming Ionic Compounds
5. The Solubility Rules
6. Determining Oxidation Numbers
7. Predicting Chemical Reactions
Rules for Determining Oxidation Number
Oxidation Number: A number assigned to an atom in a molecular compound or
molecular ion that indicates the general distribution of electrons among the bonded
atoms.
1. The oxidation number of any uncombined element is O.
2. The oxidation number of a monatomic ion equal the charge on the ion.
3. The more electronegative element in a binary compound is assigned the number
equal to the charge it would have if it were an ion.
4. The oxidation number of fluorine in a compound is always –1
5. Oxygen has an oxidation number of –2 unless it is combined with F, when it is +2,
or it is in a peroxide, when it is –1.
6. The oxidation state of hydrogen in most of its compounds is+1 unless it combined
with a metal, in which case it is –1.
7. In compounds, the elements of groups 1 and 2 as well as aluminum have oxidation
number of +1, +2, and +3, respectively
8. The sum of the oxidation numbers of all atoms in a neutral compound is O.
9. The sum of the oxidation number of all atoms in a polyatomic ion equals the charge
of the ion.
Rules for Naming Ionic Compounds
1. Balance Charges (charges should equal zero)
2. Cation is always written first ( in name and in formula)
3. Change the ending of the anion to -ide
Solubility Rules
1. All compounds containing alkali metal cations and the ammonium ion are soluble.
2. All compounds containing NO3-, ClO4-, ClO3-, and C2H3O2- anions are soluble.
3. All chlorides, bromides, and iodides are soluble except those containing Ag+, Pb2+,
or Hg2+.
4. All sulfates are soluble except those containing Hg2+, Pb2+, Sr2+, Ca2+, or Ba2+.
5. All hydroxides are insoluble except compounds of the alkali metals, Ca2+, Sr2+, and
Ba2+.
6. All compounds containing PO43-, S2-, CO32-, and SO32- ions are insoluble except
those that also contain alkali metals or NH4+.
Variable Valences for Transition Metals
Chromium +2, +3
Manganese +2, +3
Iron +2, +3
Cobalt +2, +3
Copper +1, +2
Lead +2, +4
Mercury (This one is weird) Hg2+2 , Hg+2
Tin +2, +4
Gold +1, +3
Rules for Naming an Acid
1. When the name of the anion ends in –ide, the acid name begins with the prefix
hydro-, the stem of the anion has the suffix –ic and it is followed by the word
acid. -ide becomes hydro _____ic Acid
Cl- is the Chloride ion so HCl = hydrochloric acid
2. When the anion name ends in –ite, the acid name is the stem of the anion with
the suffix –ous, followed by the word acid.
-ite becomes ______ous Acid
ClO2
- is the Chlorite ion so HClO2. = Chlorous acid.
3. When the anion name ends in –ate, the acid name is the stem of the anion with
the suffix –ic, followed by the word acid.
-ate becomes ______ic Acid
ClO3
 is the Chlorate ion so HClO3 = Chloric acid.
Review writing chemical formulas (also known as nomenclature).
a. Go to the following links to review writing formulas:
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http://members.aol.com/profchm/naming.html
http://www.chem.vt.edu/RVGS/ACT/notes/Nomenclature.html
http://dbhs.wvusd.k12.ca.us/webdocs/Nomenclature/Nomenclature.html
b. Go to the following links to practice writing formulas. Try at least 10 from
each page. Read the directions at the top of the page before trying.
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http://chemistry2.csudh.edu/newlechelp/namingcs.html (Formulas to names.)
http://chemistry2.csudh.edu/homework/hwnaming.html (Formulas to names.)
http://proton.csudh.edu/homeworkcs/hwstocknamingcsn7.html
(Don't forget to use Roman numerals when using the Stock System.)
http://proton.csudh.edu/homeworkcs/hwnamestoformulascsn7.html (Names to
formulas)
Chemical Reactions: Things to Memorize
Some of the common ions will produce a specific compound when reacting, and the
reaction can be reversed to re-form the ion.
SO42-
SO3
PO43-
P2O5
NO31-
N2O5
SO32-
SO2
CO32-
CO2
NO21-
N2O3
OH1-
H2O
ClO31-
NH41+
NH3
Cl1- + O2
Example: When Na2CO3(s) is heated, it will decompose to form CO2(g). Solid Na2O will
also be formed as the remainder of the original compound. The Na2O(s) will slowly reabsorb CO2(g) from the air and convert back into Na2CO3(s). Other carbonates follow
this same pattern, as seen in the reaction of carbonic acid;
H2CO3(aq)
CO2(g) + H2O.
Other generalizations include:

There are 7 diatomic elements (HONClBrIF) that exist as molecules in the elemental
state. Remember Hey Oh No, Clyde Brought In Firearms

There are 7 strong acids that ionize completely in water: HCl, HBr, HI, HClO4,
HClO3, HNO3 and H2SO4

All metal hydroxides are strong bases, but only the Group IA hydroxides are very
soluble.

Metal oxides react with water to form basic solutions. Nonmetal oxides react with
water to form acidic solutions.

The conjugate base of a weak acid reacts with water (hydrolysis) to reform the acid.
Likewise, the conjugate acid of a weak base reacts with water to reform the base.

During double replacement (metathesis) reactions, there are no changes in the
oxidation states of the elements involved; only a re-combination of the existing ions
occurs. One of the new compounds formed must be separable from the original
mixture, such as a precipitate or gas.

When working with water solutions, it is helpful to have a few rules concerning which
substances are soluble, and which will form precipitates. The more common solubility
rules are listed below:
1. All common salts of the Group IA(Li, Na, K, etc) elements and the ammonium ion
are soluble. – This one is really important to memorize
2. All common acetates, nitrates, chlorates and hydrogen carbonates are soluble. –
This one is also really good to memorize
3. All binary compounds of Group VIIA elements(other than F) with metals are
soluble, except those of silver, mercury (I), and lead.
4. All sulfates are soluble except those of barium, strontium, lead, calcium, silver, and
mercury (I).
5. Except for those in Rule #1, everything else is insoluble.

During complete combustion reactions of hydrocarbons, CxHy , with oxygen, the
products formed will be CO2 and H2O.

Single replacement reactions occur only if the free element is more reactive than the
element it would replace in the compound. (check the activity series or Standard
Reduction Potential, SRP, chart)

If an element has several possible positive oxidation states, assume complete
oxidation occurs during single replacement reactions and use the highest charge, but
during composition reactions use the most stable oxidation state.

Ionic compounds are formed when a metal and a nonmetal react. Covalent
(molecular) compounds are formed when two nonmetals react.

When two nonmetals react, the element with the higher electronegativity (closer to
the upper right corner of the periodic table) will have the negative oxidation number.

Transition metals typically form colored solutions and colored precipitates.

Transition metal ions act as Lewis acids by accepting an electron pair from anions
and/or polar molecules to form complex ions. The number of attached groups (ligands)
is usually twice the charge of the metallic ion.

The total loss of electrons by the oxidation reaction(s) must equal the total gain of
electrons by the reduction reaction(s). Adjust the coefficients of the reactants and
products that contain the oxidized and reduced elements first, then balance any other
elements by inspection.

When only net ionic equations are used, spectator ions or molecules not involved in
the reaction are excluded. These types of equations are balanced by the following
steps:
Step 1:
Write the equation in ionic form. Assign the oxidation numbers for the
elements.
Step 2:
Write the separate oxidation and reduction half-reactions.
Step 3:
Adjust the coefficients to get an equal number of atoms of all elements
EXCEPT oxygen and hydrogen.
Step 4:
Balance the hydrogen and oxygen atoms by adding H1+ and H2O.
Step 5:
Balance the electric charge on both sides of each half-reaction by adding
electrons.
Step 6:
Multiply each half-reaction by an appropriate number to make the electron
changes equal.
Step 7:
Combine the two half-reactions and simplify(subtract substances that
appear on both the reactant and product sides of the equation). Check to see
if the numbers of atoms of each element and total charges of reactants
equals the atoms and charges of the products.
Step 8:
If the solution is acidic, you are done. If the solution is basic, add one OH1for each H1+ present to form water, then add an equivalent number of OH1to the other side of the equation. Simplify the number of H2O present.
A.P. Chemistry Problem Set
Chapter 1 - Introduction: Matter and Measurement
1. Using your knowledge of metric units, English units, and the information in your text,
write down the conversion factors that you would need to convert the following:
a. mm to nm
b. mg to kg
c. km to ft
d. in.3 to cm3
2. A sample of ascorbic acid (vitamin C) is synthesized in the laboratory. It contains 1.50
grams of carbon and 2.00 grams of oxygen. Another sample of ascorbic acid isolated from
citrus fruit contains 6.35 grams of carbon. How many grams of oxygen does it contain?
What law are you assuming in answering this question?
3. The US quarter has a mass of 5.67 grams and is approximately 1.55 mm thick.
a. How many quarters would have to be stacked to reach 575 feet, the height of the
Washington monument?
b. What would be the mass of the stack?
c. How much money would the stack contain?
d. The US National Debt Clock showed the outstanding public debt to be
$11,687,233,914,811.11 on August 19, 2009. How many stacks like this one would
be necessary to pay off the debt?
4. The liquid substances mercury (density = 13.6 g/mL), water (D=1.00 g/mL), and
cyclohexane (D=0.778 g/mL) do not form a solution when mixed, but separate into distinct
layers. Sketch how the liquids would position themselves in a beaker?
A.P. Chemistry Problem Set
Chapter 2 – Atoms, Molecules, and Ions
1. The element oxygen has three naturally occurring isotopes, with 8, 9, and 10 neutrons in
their nucleus, respectively.
a. Write the full chemical symbols for the three isotopes
b. Describe the similarities and differences between the three kinds of atoms of
oxygen
2. It is common in mass spectrometry to assume that the mass of a cation is the same as that
of its parent atom,
a. Given that the mass of a proton is 1.0073 amu, a neutron is 1.0087 amu and an
electron is 5.486 x 10-4 amu, determine the number of significant figures that must
be reported before the difference between 1H and 1H+ is significant?
b. What percentage of the mass of an 1H atom does the electron represent?
3. Name each of the following ionic oxides. Identify the charge of the metal in each
compound:
a. NiO
b. MnO2
c. Cr2O3
d. MoO3
4. Many familiar substances have common, unsystematic names. For each of the following,
give the correct, systematic name:
a. Saltpeter, KNO3
b. Soda ash, Na2CO3
c. Lime, CaO
d. Muriatic acid, HCl
e. Epsom salts, MgSO4
f. Milk of magnesia, Mg(OH)2
A.P. Chemistry Problem Set
Chemical Formulas
Chapter 3 – Stoichiometry: Calculations with
1. Consider a sample of calcium carbonate in the form of a cube measuring 2.005 in. on each
edge. If the sample has a density of 2.71 g/cm3, how many oxygen atoms does it contain?
2. If an automobile travels 225 miles with a gas mileage of 20.5 miles/gallon, how many kg of
CO2 are produced, assuming the gasoline is composed of octane, whose density is 0.69
g/mL?
a. Repeat the calculation for a truck that has a gas mileage of 5 mi/gal.
3. A particular coal contains 2.5% sulfur by mass. When this coal is burned in a power plant,
the sulfur is converted into sulfur dioxide gas, which is a pollutant. To reduce sulfur
dioxide emissions, calcium oxide (lime) is used. The sulfur dioxide reacts with calcium
oxide to form solid calcium sulfite.
a. Write a balanced chemical equation for the reaction
b. If the coal is burned in a power plant that uses 2000 tons of coal per day, what is
the mass of calcium oxide that is required per day?
c. How many grams of calcium sulfite are produced daily by this power plant?
4. Balance and Classify the following reactions:
a. C3H6(g) +
O2(g) 
CO2(g) +
H2O(g)
b. NH4NO3(s) 
N2O(g) +
H2O(g)
c. C5H6O(l) +
O2(g) 
CO2(g) +
H2O(g)
d. N2(g) +
H2(g) 
NH3(g)
e. K2O(s) +
H2O(l) 
KOH(aq)
Chapter 4: Reactions in Aqueous Solution Open-Ended Response Questions
1. Answer the following questions about the solubility and reactions of the ionic
compounds M(OH)2 and MCO3 , where M represents an unidentified metal.
(a) Identify the charge of the M ion in the ionic compounds above.
(b) At 25°C, a saturated solution of M(OH)2 has a pH of 9.15. Calculate the molar
concentration of OH-(aq) in the saturated solution.
2. Zinc metal is added to a hydrobromic acid solution.
(a) Write the Balanced equation:
(b) Write the oxidation half-reaction for the reaction.
3. Solid lithium is added to distilled water.
(a) Write the Balanced equation:
(b) Indicate whether the pH of the resulting solution is less than 7, equal to 7, or greater
than 7. Explain.
4. A 100 mL sample of 1 M strontium chloride solution is mixed with a 100 mL
sample of 1 M sodium carbonate solution, resulting in the formation of a
precipitate.
(a) Write the Balanced equation:
(b) Describe what will occur if the precipitate is dried and a few drops of 1 M hydrochloric
acid are added. Explain.
5. A student is instructed to prepare 100.0 mL of 1.250 M NaOH from a stock solution
of 5.000 M NaOH. The student follows the proper safety guidelines.
(a) Calculate the volume of 5.000 M NaOH needed to accurately prepare 100.0 mL of
1.250 M NaOH solution.
(b) Describe the steps in a procedure to prepare 100.0 mL of 1.250 M NaOH solution
using 5.000 M NaOH and equipment selected from the list below.
Balance
50 mL buret
25 mL Erlenmeyer
100 mL graduated
100 mL volumetric
flask
cylinder
flask
100 mL Florence
25 mL pipet
100 mL beaker
Wash bottle of
Crucible
flask
Eyedropper
Drying oven
distilled H2O
The student is given 50.0 mL of a 1.00 M solution of a weak, monoprotic acid, HA. The
solution is titrated with the 1.250 M NaOH to the endpoint. (Assume that the endpoint is
at the equivalence point.)
(c) Explain why the solution is basic at the equivalence point of the titration. Include a
chemical equation as part of your explanation.
(d) Identify the indicator in the table below that would be best for the titration. Justify
your choice.
Indicator
pKa
Methyl red
5
Bromthymol blue
7
Phenolphthalein
9