AP CHEMISTRY SUMMER ASSIGNMENT
Over the summer, you will be responsible for completing this summer assignment packet.
To complete the assignment, refer to your Chem I notes or the online Chem I textbook. It is
not necessary to purchase an AP Chem book to complete these assignments. But if you
would like to do so, our current AP Chem book is “Chemistry” fifth edition by Zumdahl, ISBN
# 0-395-98583-8. The book can be purchased on Amazon.com or ebay.com.
The completed summer assignment packet (in your own handwriting) will be due on the
first day of school – Mon. August 26, 2013. A quiz over the “Common and Polyatomic Ions to
be Memorized” list will be on Tues. August 27, 2013. You must know everything on that list –
including the common names! A test over all the summer assignments will be on Friday.
August 30, 2013.
 To access the online Chemistry I textbook: Go to http://my.hrw.com/
The login is rangers2122 and the password is chemistry.
SIGNIFICANT FIGURES AND CALCULATIONS
Rules for Counting Significant Figures
1. Nonzero integers. Nonzero integers are always significant.
2. Zeros. There are three classes of zeros:
a. Leading zeros are zeros that precede all the nonzero digits. These do not count as
significant figures.
Example: 0.0025 m has only 1 significant figures.
b. Captive zeros are zeros between nonzero digits. These always count as significant
figures.
Example: 1.008 m has 4 significant figures.
c. Trailing zeros are zeros at the right end of a number. They are significant only if the
measurement contains a decimal point.
Examples: 100 m has only 1 significant figure.
100. m has 3 significant figures
1.00 x 103 has 3 significant figures.
3. Exact Numbers Many times calculations involve numbers that were not obtained using
measuring devices but were determined by counting. For example, 10 experiments, 3
apples, 8 books. Such numbers are called exact numbers. They
can be assumed to have an infinite number of significant figures.
Rules for Significant Figures in Mathematical Operations
1. For multiplication or division, the number of significant figures in the result is the same as the
number in the least precise measurement used in the calculation.
Example: 4.56 m x 1.4 m = 6.4 m2
2. For addition or subtraction, the result has the same number of decimal places as the least
precise measurement used in the calculation.
Example: 12.11 m + 18.0 m + 1.013 m + 31.123 m = 62.2 m
MORE PRACTICE!
1. How many significant figures are in the following measurements:
(a) 1406.20 g
(d) 8 x 108 nm
(b) 1600.0 cm
(e) 0.0261140 dm
(c) 0.0007 m
(f) 200 km
2. Use exponential notation to express the number 37,100,000 with
(a) 1 sig fig
(c) 3 sig figs
(b) 2 sig figs
(d) 6 sig figs
3. Perform the indicated calculations on the following measured values, giving the final answer with the
correct number of significant figures.
(a) 16.81 + 3.2257
(c) [(3.901-3.887)/3.901] x 100.0
(b) 324.6 + 815.991
(d) 3.14159 x 68 / (5.18 x 10-11 – 6 x 10-4)
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SI PREFIXES
Name
Multiples
Symbol
deka- hecto- kilo- mega- giga- tera- peta- exa- zetta-
da
Factor 100 101
Name
Subdivisions Symbol
H
k
M
G
T
P
E
Z
102
103
106
109
1012
1015
1018
1021
deci- centi- milli- micro- nano- pico- femto- atto- zepto-
d
Factor 100 10−1
C
m
µ
10−2
10−3
10−6
n
p
f
a
z
10−9 10−12 10−15 10−18 10−21
DIMENSIONAL ANALYSIS
This method involves using conversion factors to cancel units until you have the proper unit in the
proper place.
EXAMPLES: (Show all work!!)
1. How many micrometers are in 2.60 kilometers?
2. An average adult has 5.2 L of blood. What is the volume of blood in m3?
3. In the vacuum of space, light travels at a speed of 186,000 miles per second. How many
centimeters can light travel in 1 year?
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ATOMS, MOLECULES, AND IONS
Know:
The Law of Conservation of Mass
The Law of Definite Proportion
The Law of Multiple Proportions
Dalton’s Atomic Theory
Isotopes - atoms of the same elements that have the same number of protons, but different number of neutrons.
Examples:
1) Identify each of the following elements:
(a) 9140X
(b) 10847X
(c) 3316X
2) How many protons and neutrons are in each of the following elements?
(a) 89Y
(b) 73Ge
(c) 24Mg2+
3) How many protons, neutrons, and electrons are in each of the following ions?
(a) 56Fe3+
(b) 40Ca2+
(c) 19F-
Calculation of Average Atomic Mass from Isotopic Data (Show all work!)
1)A noble gas consists of three isotopes of masses of 19.99 amu, 20.99 amu, and 21.99 amu. The relative
abundances of these isotopes are 90.92%, 0.257%, and 8.82% respectively. What is the average atomic
mass of this noble gas? What noble gas is this?
2) The element Europium exists in nature as 2 isotopes: 151Eu has a mass of 150.9196 amu, and 153Eu has a mass
of 152.9209 amu. The average atomic mass of Europium is 151.96 amu. Calculate the relative abundances of
the two Europium isotopes.
Questions:
1) What discoveries were made by J.J. Thompson, Robert Millikan, and Lord Rutherford? How did Dalton’s
atomic theory have to be modified to account for these discoveries?
2) Distinguish between the terms group (family) and period in connection to the periodic table. On the periodic
table, know the location of the : alkali metals, alkaline earth metals, halogens, noble gases.
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NAMING SIMPLE COMPOUNDS
NOMENCLATURE
A. Naming Compounds
1. Binary Salt - a salt that contains only one kind of cation and one kind of anion.
a. Cations- Ions with a positive charge.
Two classes: (i)Cations with only one possible charge/oxidation state(ex. Group I, II and III elements
form ions with only one charge/oxidation state).
(ii) Cations with more than one possible charge/oxidation state.
(ex. Most transition metals form ions with more than one charge/oxidation state)
b. Anions – Ions with a negative charge. All monoatomic anion names end with –ide.
Example: Sulfur will form the anion with a charge of 2- called the sulfide ion.
c. Naming binary salts – combine the names of the cation and anion. The cation always comes
first in a name (include roman numeral if necessary) and the total charge of the compound must
equal zero.
Examples Name the following compounds
1) KCl
3) CaF2
2) SnCl4
4) Fe2O3
2. Salts with Polyatomic Ions (Ternary Salts)
The same rules are followed when naming compounds containing polyatomic ions.
Examples: Name the following compounds
1) KMnO4
3) Fe(OH)3
2) (NH4)2Cr2O7
4) Cr(NO3)2
3. Binary Covalent Compounds
These compounds are formed between 2 nonmetals.
Know the prefixes from Chem I or found in Table 7.3 in the Chem I online textbook.
The prefix mono- is never used with the first element.
Examples: Name the following compounds
1) N2O3
3) CO
2) P4O10
4) BrF3
4. Acids
For now, we will consider an acid to be any compound in which hydrogen is the only cation. We will
learn much more about acids (and different definitions) of acids in later chapters.
a. Binary acids (nonoxyacids) – contain an anion that ends in –ide. Change the anion ending to –ic
and add the prefix hydro- along with the word acid.
Examples: Name the following acids
1) HF
2) HCl
One exception to the rule is CN-. Even though the anion is polyatomic, the anion still ends in –ide
and the acid is named hydrocyanic acid.
b. Oxyacids – contains a polyatomic ion with oxygen and either ends in –ite or –ate.
If the anion ends in –ite, change the ending to –ous and add the word acid.
If the anion ends in –ate, change the ending to –ic and add the word acid.
Examples: Name the following acids
1) HNO3
2) H2SO3
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3) HClO4
4) HClO
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B. Writing Formulas from Names
1. Ionic Compounds (Salts)– the compound is neutral; therefore, the charge of the cation plus the charge
of the anion must equal zero.
Examples: Write formulas for the following ionic compounds.
1) potassium chloride
3) iron (III) nitrate
2) calcium phosphate
4) magnesium hydroxide
2. Binary Covalent (Molecular) Compounds –use prefixes to indicate number of each atom in the molecule.
Examples: Write formulas for the following covalent compounds.
1) dinitrogen tetroxide
2) phosphorus pentachloride
3. Acids- use suffixes and in some case the prefixes associated with the acid to determine the anion that is
part of the acid. Remember H+ is the cation and the charge of the cation plus the charge of the anion
must equal zero (acids, like salts, are neutral).
Examples: Write formulas for the following acids.
1) nitrous acid
3) bromic acid
2) sulfuric acid
4) phosphorus acid
MORE PRACTICE:
A. Name the following compounds:
1) MgSO4
9) Fe2O3
2) N2O3
10) NH3
3) KMnO4
11) Cr(OH)3
4) NiO
12) AgNO3
5) HCN
13) HBr
6) NaHCO3
14) HIO3
7) P4O6
15) BF3
8) LiH
B. Write formulas for the following compounds
16) sulfur trioxide
24)plumbous nitrate
17) potassium chlorate
25) ammonium phosphate
18) hydrosulfuric acid
26) diphosphorus pentoxide
19) iron (III) chloride
27) zinc carbonate
20) acetic acid
28) ferrous chlorite
21) sodium hydrogen sulfate
29) carbonic acid
22) uranium hexafluoride
30) ammonium oxalate
23) barium chloride
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NOMENCLATURE OF ALKANES AND ALCOHOLS
Hydrocarbon - a compound composed of hydrogen and carbon.
Alkane – the simplest hydrocarbon. General formula: CnH2n+2
The following prefixes are used to indicate the number of carbon in the alkane:
Prefix
methethpropbutpenthexheptoctnondec-
Number of carbon atoms
1
2
3
4
5
6
7
8
9
10
To name the alkane, use the prefix that corresponds to the number of carbon atoms and add the
ending –ane.
Examples:
A. Name the following alkanes
1) CH4
3) C2H6
2) C4H10
4) C7H16
B. Write the correct formulas for the following:
1) octane
3) ethane
2) propane
4) decane
Alcohols are an example of a substituted hydrocarbon. They have an –OH group (hydroxyl group) substituted
in place of a hydrogen on a carbon.
General Formula: CnH2n+1OH
When naming alcohols, use the same prefixes as above and add the ending –anol.
Examples:
A. Name the following alcohols
1) CH3OH
2) C5H11OH
B. Write the correct formula for the following alcohols:
1) ethanol
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2) hexanol
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STOICHIOMETRY
I. Composition Stoichiometry (Show all work!!)
A. Percent Composition of a Compound
Calculate the mass percent of each element in barium sulfite.
B. Determining the Formula of a Compound
The empirical formula is represented by the simplest whole number ratio of atoms in a compound.
Examples:
The molecular formula is represented by the actual ratio of atoms in a compound.
Examples:
Can the empirical and molecular formula be the same? Yes !!
1. Determining the Empirical Formula of a Compound:
Vanillin is often used in the production of vanilla extract, Vanillin contains the following percentages
by mass:
C = 63.18 %, H = 5.26 %, O = 31.56%
Determine the empirical formula of vanillin.
2. Determining the Molecular Formula of a Compound
A compound is found, by mass spectral analysis, to contain the following percentages of elements by
mass:
C = 49.97 %, Cl = 48.92%, H = 1.39 %
The molar mass of the compound is 289.9 g/mole. Determine the empirical and molecular formulas of
the compound.
C. Conversion between Atoms, Moles, and Mass
1. How many moles are in a sample of 300. atoms of Nitrogen? How many grams?
2. How many atoms of gold does it take to make 1.00 g of gold?
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D. Conversions between Ions, Formula Units, Atoms, Molecules, Moles, and Mass
1. How many atoms of nitrogen are present in 1.0 g of ammonia?
2. Calculate the number of phosphate ions in 25.0g calcium phosphate. Number of oxygen atoms?
II. Reaction Stoichiometry (Show all work!!)
A. Reaction Review
1. Balance the following equations, then classify each reaction as either synthesis, decomposition, single replacement,
double replacement, or combustion:
(a) N2 + H2  NH3
(b) Li + H2O  LiOH + H2
(c) KClO3  KCl + O2
(d) C6H14 + O2  CO2 + H2O
(e) H2SO4 + NaOH  Na2SO4 + H2O
2. Complete and balance the following reactions, also classify each reaction
(a) calcium + nitrogen 
(b) octane + oxygen 
(c) bromine + potassium iodide 
(d) cupric chloride + sodium sulfide 
(e) sodium chlorate 
(f) barium oxide + water 
(g) magnesium carbonate 
(h) zinc + hydrochloric acid 
B. Reaction Stoichiometry Calculations
1. How many moles of water vapor can be generated from the combustion of 18.74 g of ethanol?
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2. Sodium hydroxide reacts with lead (II) nitrate. How many grams of sodium hydroxide are required to form
51.63 g of lead (II) hydroxide.
C. Calculations Involving a Limiting Reactant
1. A reaction combines 64.81 g of silver nitrate with 92.67 g of potassium bromide. How many grams of silver
bromide are formed?
2. Sodium hydroxide reacts with phosphoric acid. If 17.80 g of sodium hydroxide is mixed with 15.40 g of
phosphoric acid
(a) How many grams of water can be formed?
(b) How many grams of the excess reactant remain unreacted?
(c) If the actual yield of water is 5.00 g , what is the percent yield of water?
ADDITIONAL PROBLEMS
1. Mercury and bromine react in a synthesis reaction.
(a) What mass in g of the product can be produced from the reaction of 10.0 g of Hg and 9.00 g Br2?
(b) What mass (g) of excess reactant is left unreacted?
(c) If 16.5 g of the product in produced when this reaction is carried out in a lab, what is the percent yield?
2. If a reaction of 5.0 g of hydrogen with 5.0 g of carbon monoxide produced 4.5 g of methanol, what is the
percent yield?
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ADDITIONAL STOICHIOMETRY PROBLEMS
Show all work!
1. A compound containing only sulfur and oxygen is 50.0% sulfur by mass. What is the empirical formula
of the compound?
2. How many grams of potassium iodide are necessary to completely react with 20.61 g of mercury (II)
chloride?
3. Calculate the mass percent of each element in ferric acetate.
4. How many grams of H2O will be formed when 32.0 g H2 is mixed with 32.0 g O2 and allowed to react
and form water? How many grams of excess reactant remain? If 34.5 g of water was produced in a lab
situation, what is the percent yield of water?
5. Adipic acid is an organic compound composed of 49.31% C, 43.79% O, and the rest hydrogen. If the
molar mass of adipic acid is 146.1 g/mol, what are the empirical and molecular formulas for adipic
acid?
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