AP
CHEMISTRY
AP Problem Set:
Summer Assignment (2010)
1
The first part of your assignment is to read the review material in this packet.
You will need
to
know the names and chemical element symbols on page 3, the metric prefixes on page 3, one of the
conversion factors from each column on page 4, the tables on pages 7-8, prefixes for naming covalent
compounds on page 9, and the solubility rules on page 10.
The textbook we will be using is, Chemistry, 6th edition, by Zumdahl. I recommend picking up a copy
from ebay or amazon so that you have you own personal copy in which you can write and take notes.
However, the school will also be providing you with a copy on the first day of class.
The final part of your assignment is to complete the “Reading, ‘Riting, and ‘Rithmetic” assignments at
the end of this packet.
If you have any questions, please email me at mgangluff@whrhs.org.
The packet (in additional to the 1988 National Chemistry Olympiad Local Exam, see page 11) its
entirety is due the first day of class. All work must be shown to receive full credit.
2
1
Matter – Its Properties
and Measurement
Section Objectives
1. Write the names and chemical symbols of the more common elements.
This is a memorization task: to know the symbol that goes with the name and vice versa. “Common
elements” means different things to different people. A reasonable goal would be the main group
elements along with those in the first transition series (Sc through Zn) plus Ag, Au, Cd, and Hg. These
are elements with atomic numbers 1-38, 47-56, and 79-88. The atomic number is the whole number in
each box in the periodic table, often given the symbol Z. The best way to learn names is by groups
(vertical columns) in the periodic table. By learning them together, you are also learning chemistry,
not just memorizing symbols.
2. Know common units in the English system, and the relationships between them.
Common English units, their abbreviations, and the relationships between them follow. They should
be memorized if you do not know them already. Most scientists, and practically all engineers, in the
U.S. do.
Volume measure: Gallon (gal) = 4 quarts (qt) Pint (pt) = 16 fluid ounces (fl oz) = 2 cups (c)
Quart (qt) = 2 pt = 32 fl oz
Tablespoon (T) = 2 teaspoons (tsp) = 1/2 fl oz
Linear measure:
Yard (yd) = 3 feet (ft)
Mile (mi) = 1760 yd
Foot (ft) = 12 inches (in)
Mass measure:
Pound (lb) = 16 ounces (oz) Ton (t) = short ton = 2000 lb
3. For the metric system, state the basic units of mass, length, and volume, and the common
prefixes.
At present there are three units to be learned: gram (g) for mass, meter (m) for length, and liter (L) for
volume. The following prefixes should be memorized.
Mega (M)
or
1,000,000
One million of
106
(k)
103
or
1000
One thousand of
Deci
(d)
10
-1
or
1/10
One tenth of
Centi
(c)
10-2
or
1/100
One hundredth of
Milli
(m)
10
-3
or
1/1000
One thousandth of
Micro
(μ)
10-6
or
1/1,000,000
One millionth of
Nano
(n)
10-9
or
1/1,000,000,000
One billionth of
Kilo
3
4. State the relationships between English and metric units.
There are many relationships, but they can be reduced to only three if you know the interrelationships
within each system. These three can be one from each column that follows.
Mass
453.6 g = 1 lb
1 kg = 2.205 lb
28.35 g = 1 oz
Length
2.5400 cm = 1 in
1 m 39.37 in
30.48 cm = 1 ft
Volume
0.9464 L = 1 qt
1 L = 1.057 qt
29.57 mL = 1 fl oz
5. Determine the number of significant digits in a numerical calculation.
6. Express the result of a calculation with the appropriate number of significant digits.
Here are the rules that govern significant digits.
a) Significant digits include all non-zero digits, zeros located between significant digits (“captive
zeros”), and zeros located after non-zero digits to the right of the decimal point (“trailing
zeros”).
b) Zeros that precede non-zero digits and zeros that end a number with no decimal point are not
significant (“placeholders”).
c) The result of a multiplication or division should have as many significant digits as the factor
with the fewest number of significant digits.
d) The result of an addition or a subtraction is rounded to the same number of decimal places as
the term with the fewest number of decimal places.
EXAMPLE 1-1
Express the result of the following calculation with the appropriate number of significant digits.
[(725.6 - 19.1)/760. 00]75 x 10 -3
 ? The result is 2.9 x 10-3
(0.082057)( 293.2)
7. Express numbers in scientific notation.
8. Write a conversion factor from a relationship between two quantities, and use conversion
factors to solve problems.
Probably the most powerful problem-solving method you can learn is the conversion factor method.
The method is valuable no only in chemistry but in any numerical problem-solving course.
EXAMPLE 1-2
How many teaspoons are in 5.00 gallons?
4 qt 32 fl oz
3 tsp
5.00 gal x
x
x
 3.84 x 103 tsp
1 gal
1 qt
0.5 fl oz
EXAMPLE 1-3
How many gallons are in 1.00 cubic foot?
3
3
1L
1 qt
1 gal
 12 in   2.5400 cm 
3
1.00 ft x 
x
x
 7.48 gal
 x
 x
1 in
 1 ft  

1000 cm3 0.946 L 4 qt
9. Express and use density in the form of conversion factors.
Density is both a physical property of a substance and the means of interconverting mass and volume
m
of that substance. The defining equation ( d  ) has three variables: density, mass, and volume.
v
4
EXAMPLE 1-4
An empty container weighs 206 g. Filled with 242 mL of liquid it weighs 938 g. What is the density
of the liquid?
(938 g  206 g )
d
 3.02 g/mL
242 mL
EXAMPLE 1-5
A 27.4 mL gold (19.3 g/mL) object has what mass?
19.3 g
27.4 mL x
 529 g Au
1 mL
EXAMPLE 1-6
A 75.2 g piece of zirconium (6.42 g/mL) has what volume?
1 mL
75.2 g x
 11.6 mL Zr
6.42 g
10. Express and use percent composition in terms of conversion factors.
Percent means part per hundred. Thus, 40.0% C in acetic acid means 40.0 g C in 100.0 g acetic acid.
EXAMPLE 1-7
What mass of acetic acid contains 247 g C?
100.0 g acetic acid
247 g C x
 618 g acetic acid
40.0 g C
11. Solve algebraic equations that arise in the course of working chemistry problems.
Solving an algebraic equation generally means obtaining a new equation, with the symbol for one
variable isolated on one side and the remainder of the equation of the of the equation on the other side.
EXAMPLE 1-8
Solve the following equation for P.
2

 P  an V - nb   nRT


V2 

P
nRT
an 2
V - nb  V 2
5
2
Atoms and the
Atomic Theory
Section Objectives
1. List the numbers of protons, neutrons, and electrons present in atoms and ions, using the
symbolism A
ZE.
The complete symbol for an atom or ion consists of the elemental symbol surrounded by subscripts and
superscripts.
a) The leading superscript (upper left) is the mass number. This is also the number of nucleons; a
nucleon is a proton or a neutron.
b) The leading subscript (lower left) is the atomic number or proton number.
c) The trailing superscript (upper right) is the charge or the number of protons (atomic number)
minus the number of electrons. The sign (+ or -) always must be included. The number is zero
for a neutral atom, but the zero is written only for emphasis.
Mass
Charge
number
235 U3
92
Atomic number
EXAMPLE 2-1
What is the atomic number, mass number, and charge of 19 F  ?
19 F  has an atomic number of 9; a mass number of 19; and a charge of –1.
2. Use the periodic table to predict the charges of ions of main group elements.
Elements in the same column have similar properties. Each column is referred to as a periodic family
or group. The horizontal rows are called periods. Elements on the right side of the periodic table are
nonmetals; they form anions, or negatively charged ions. Elements on the left side of the periodic
table are metals; they form cations, or positively charged ions. Elements within the same group will
form ions with the same charge.
6
3
Chemical Compounds
Section Objectives
1. Know and apply the conventions used in determining oxidation states.
Because oxidation state is a formal rather than an experimental concept, it is possible to devise a rigid
set of rules that work in all but the most unusual circumstances. One set of rules is given below.
1.
2.
3.
a.
b.
c.
d.
e.
f.
g.
h.
i.
j.
Method of applying the rules
Apply the rules from the top to the bottom of the list.
Search the list to find a rule that fits. Apply it.
Then start again at the top of the list to find the next rule that fits.
Oxidation State rules
The OS (oxidation state) of all uncombined elements = 0.
The sum of the OS in compound = 0.
The sum of the OS in an ion = ionic charge.
Alkali metals (group 1A) have OS = +1.
Alkaline earth metals (2A) have OS = +2.
F has OS = -1 and H has OS = +1.
O has OS = -2.
Cl, Br, I (in order) have OS = -1.
S, Se, Te (in order) have OS = -2.
N, P, As (in order) have OS = -3.
2. Know the names, formulas, and charges of ions in the following tables and be able to write
formulas and names of the compounds formed from these ions.
Al
H
3+
-
Aluminum
Hydride
Names, Formulas, and Charges of Some Common Ions
Iron (III) or ferric
Au3+ Gold (III) or auric
Fe3+
2+
2+
Tin (II) or stannous
Cobalt (II) or cobaltous
Sn
Co
Mn2+ Manganese (II)
Ni2+ Nickel (II)
2+
Sn4+
Tin (IV) or stannic
Co3+
Cobalt (III) or cobaltic
Pb2+
Lead (II) or plumbous
Cu+
Copper (I) or cuprous
Zinc
Pb
Cd2+
Cadmium
Ag+
Au+
Zn
4+
2+
Lead (IV) or plumbic
Cu
Copper (II) or cupric
Cr2+
Chromium (II) or chromous
Silver
Cr3+
Chromium (III) or chromic
Hg22+ Mercury (I) or mercurous
Hg2+ Mercury (II) or mercuric
Gold (I) or aurous
Fe2+
Iron (II) or ferrous
7
NH4
Names, Formulas, and Charges of Some Common Polyatomic Ions
Ammonium
Sulfate
Hypofluorite
SO42FO-
C2H3O2-
Acetate
HSO4-
Hydrogen sulfate
ClO-
CO32-
Carbonate
SO32-
Sulfite
ClO2-
Chlorite
-
+
HCO3C2O42-
-
Hydrogen carbonate
HSO3
Oxalate
S2O32-
CN
Cyanide
HS
-
OCN-
Cyanate
SCN-
Thiocyanate
Hydrogen sulfite
ClO3
Chlorate
Thiosulfate
ClO4-
Perchlorate
-
Hydrogen sulfide
BrO
OH-
Hydroxide
BrO3-
Bromate
O22-
Peroxide
BrO4-
Perbromate
-
Hypoiodite
NO2
-
Nitrite
CrO4
Chromate
IO
NO3-
Nitrate
Cr2O72-
Dichromate
IO3-
3-
2-
-
Phosphate
MnO4
Permanganate
HPO42-
Hydrogen phosphate
MnO42-
Manganate
H2PO4-
Dihydrogen phosphate
PO4
Hypochlorite
IO4
-
Hypobromite
Iodate
Periodate
The lists presented above may seem rather extensive, but they contain practically all the ions you are
likely to encounter in AP Chemistry. (You may get a few more in September, but this will give you a
good start!)
Naming an ionic compound is simple. Write down the name of the cation (positive ion), followed with
a space, and then the name of the anion (negative ion).
Writing formulas from names is not quite so simple. The formula contains more than just the symbols
for the cation and anion. The cation and anion symbols are multiplied so that the total charge from the
cations just balances the total charge of the anions. The total cation charge plus the total anion charge
equals zero.
Cations
M+
[NH4+]
2+
[Ca ]
M3+
[Al3+]
M4+
[Ce4+]
M
2+
-
X
MX
-
[F ]
2-
X
Anions
[SO42-]
[NH4F] M2X [(NH4)2SO4]
MX2 [CaF2] MX
[CaSO4]
MX3 [AlF3] M2X3 [Al2(SO4)3]
MX4 [CeF4] MX2 [Ce(SO4)3]
X3M 3X
[PO43-]
[(NH4)3PO4]
M3X2 [Ca3(PO4)2]
MX
[AlPO4]
M3X4 [Ce3(PO4)4]
8
3. Be able to write formulas and names of simple binary covalent compounds and of binary
acids.
Covalent compounds are formed between nonmetallic elements. The names of binary covalent
compounds are obtained from the names of the two elements. The elements are named in the same
order as they appear in the formula. The first element name is unchanged; the ending of the second
becomes “-ide.” The element names have prefixes depending on the subscript of that element in the
formula, except that the prefix mono- (meaning one of) is rarely used for the first element in a formula.
Other prefixes are: di = 2, tri = 3, tetra = 4, penta = 5, hexa = 6, hepta = 7, octa = 8, nona = 9, and deca
= 10.
Binary acids consist of hydrogen and a nonmetal. HCl is a binary acid. The name of a binary acid has
the prefix “hydro-” and the suffix “-ic” surrounding the root name of the element. HCl is hydrochloric
acid. The binary acid names are used when the compound is dissolved in water, that is, in aqueous
solution. When the compound is not an aqueous solution the name is the same as any ionic compound.
4. Use oxidation states to name oxoacids and oxoanions.
Salts and acids of chlorine oxoanions
Ox.
State
+1
Salt
Example
Acid
Example
Hypo- -ite
NaClO
Sodium hypochlorite
NaClO2
Sodium chlorite
NaClO3
Sodium chlorate
NaClO4
Sodium perchlorate
Hypo- -ous
HClO
Hypochlorous acid
HClO2
Chlorous acid
HClO3
Chloric acid
HClO4
Perchloric acid
+3
-ite
+5
-ate
+7
Per-
-ite
-ous
-ic
Per-
-ic
All oxoanions of the same family with the same oxidation state have similar names. Another
generality is that the –ate anion and the –ic acid endings are used when the oxidation state of the
central atom equals the periodic table family number. The only exceptions to this occur in the
halogens, where the –ate and the –ic endings correspond to a +5 oxidation state and the noble gases
where they correspond to +6.
5. Use solubility rules to predict products of reactions.
The attached table of rules is one form of the solubility rules. You are responsible for learning
these rules in some format. Each textbook gives a slightly different approach.
9
SOLUBILITY RULES
LEARN!!
The solubility of a solute is the amount that can be dissolved in a given quantity of solvent at a given
temperature. For example, the solubility of lead (II) nitrate is
56 g/100 mL at 20oC. The solubilities of ionic solids in water vary over a wide range of values. For
convenience, we divide compounds into three categories called soluble, slightly soluble and insoluble.
Insoluble is a relative term and does not mean that no solute dissolves! Compounds are classified as
insoluble if their solubility is less than
0.1 g/100 mL of water. On the other hand, soluble compounds are those whose
solubilities are greater than 1.0 g/100 mL of water. The following “solubility rules” summarize the
solubilities of various compounds in water at 25oC.
1. All Group IA salts are soluble (aq).
2. All ammonium salts are soluble (aq).
3. All salts containing nitrate, acetate, chlorate and perchlorate are soluble (aq).
4. All salts containing halides (chlorides, bromides, iodides, and fluorides) are
soluble (aq) EXCEPT silver, mercury(I) and lead (s). (Lead halides are soluble in
hot water.)
5. All sulfate salts are soluble (aq) EXCEPT barium, calcium, strontium, silver, mercury(I) and lead
(s).
6. All salts containing carbonates, phosphates, and chromates are insoluble (s) EXCEPT
for rules #1 and 2 (aq).
7. All sulfide salts are insoluble (s) EXCEPT for rules #1 and 2 and calcium, strontium, and barium
(aq).
8. All hydroxide salts are insoluble (s) EXCEPT for rules #1 and 2 and barium and strontium (aq).
(Calcium hydroxide is very slightly soluble.)
Note: Rule #8 is the one that varies from text book to text book and causes the most trouble for
people writing net ionic equations. Are the Group IIA hydroxides soluble or not? At best they are
only moderately soluble – barium and strontium are a little more soluble than calcium and usually are
called soluble. Calcium hydroxide is usually called insoluble.
10
AP Chemistry Summer Reading and ‘Riting and ‘Rithmetic – 2010
I know – it is the end of a long school year and after final exams you do not want to even think of
opening a school book again for a very long time. Sorry! A little refresher is in order before we start
off in September. You have some review chapters to read and some problems to work out. Since
these problems will be graded, it is important that you take them seriously! In addition, a few class
periods at the start of the school year will be used for questions on these assigned problems and then
you will be tested on the material. This material is definitely not busy work to make your summer
miserable!!
All assignments and the Flinn Safety Contract must be completed and turned in on the first day of
class. No credit will be given for late assignments. You will be having a safety quiz on the first day.
Remember: to receive credit for any math problems, you must show all work! Otherwise credit
will not be given – even if a numerical answer is “correct”!
It wouldn’t hurt to pick up an AP Chem Review Book early so that you can use it all year
Also, you will need a scientific calculator (not your fancy graphing calculator!) for taking your
tests and quizzes. If you don’t have one from your chem or physics class, pick one up and start
learning how to use it! I recommend the Texas Instruments TI-36X Solar. Bring that with you at the
start of school in September also.
In addition to the problems in this packet, you are to go to the following website and complete the
1988 National Chemistry Olympiad Local Exam.
http://www.chemteam.info/NChO/NChO-88-Local.html
You must show all work on a separate sheet of paper. Please be aware that there will be some
questions which you do not how to do, but this is meant to keep you fresh with the topics you learned
the last time you had chemistry.
Turn in that work with the work from this packet.
11
Summer Assignment #1
Please place all answers on a separate sheet of paper. Do not try to cram the
answers onto these pages!
I. Chemical Formulas
1. Write formulas for the following substances:
a.
b.
c.
d.
e.
f.
g.
h.
i.
j.
k.
l.
m.
n.
o.
p.
q.
r.
s.
t.
u.
v.
w.
x.
y.
z.
Barium sulfate
Ammonium chloride
Chlorine monoxide
Silicone tetrachloride
Magnesium fluoride
Sodium oxide
Sodium peroxide
Copper (I) iodide
Zinc sulfide
Potassium carbonate
Hydrobromic acid
Perbromic acid
Lead (II) acetate
Sodium permanganate
Lithium oxalate
Potassium cyanide
Iron (III) hydroxide
Silicone dioxide
Nitrogen trifluoride
Chromium (III) oxide
Calcium chlorate
Sodium thiocyanate
Cobalt (III) nitrate
Nitrous acid
Ammonium phosphate
Potassium chromate
12
2. Name each of the following compounds (Give acid names where appropriate)
a.
b.
c.
d.
e.
f.
g.
h.
i.
j.
k.
l.
m.
n.
o.
p.
q.
r.
s.
t.
u.
v.
w.
x.
y.
z.
CuSO4
PCl3
Li3N
BaSO3
N2F4
KClO4
NaH
(NH4)2Cr2O7
HNO2 (aq)
Sr3P2
Mg(OH)2
Al2S3
AgBr
P4O10
HC2H3O2 (aq)
CaI2
MnO2
Li2O
FeI3
Cu3PO4
PCl3
NaCN
Cs3N
Zn (NO3)2
N2O
HF (aq)
13
II.
Chemical Equations
Tell the type of reaction, predict the products and write a balanced chemical
equation for each of the following, as shown in the example:
Ex:
Solutions of silver nitrate and magnesium iodide are combined.
This is a double replacement reaction.
2AgNO3 + MgI2 2AgI (s) + Mg(NO3)2
1. Ammonium sulfate reacts with barium nitrate,
2. Zinc metal is added to a solution of copper (II) chloride.
3. Propane gas (C3H8) is burned in excess oxygen.
4. Dinitrogen tetroxide gas is added to distilled water.
5. Solid calcium chlorate is heated strongly.
6. Sodium hydroxide solution is added to a solution of iron (III) bromide.
7. Chlorine gas is bubbled through a solution of sodium bromide.
8. Solutions of lead nitrate and calcium iodide are combined.
9. Sulfuric acid is combined with solid magnesium hydroxide.
10. Solid barium oxide is added to distilled water.
11. Isopropyl alcohol (C3H7OH) is burned in air.
12. Iron metal shavings are added to hydrochloric acid.
13. Solid sodium carbonate is heated in a crucible.
14. Solid aluminum hydroxide is added to perchloric acid.
15. Sodium metal is added to distilled water.
14
III.
Metric Conversions, Dimensional Analysis, Atomic Structure
1) The English unit, the rod, is equal to 16.5 ft. What is this length expressed in meters?
2) A certain brand of coffee is offered for sale at $7.26 for a 3-lb. can or $5.42 for an l-kg can, which
is the better buy?
3) A sprinter runs the 100-yd dash in 9.3 s. What would be his time for a 100-m run if he ran at the
same rate?
4) The unit of length, the furlong, is used in horse racing. The units of length, the chain and the link,
are used in surveying. There are 8 furlongs in 1 mi., 10 chains in 1 furlong, and 100 links in 1
chain. To three significant figures, what is the length of 1 link in inches?
5) An English unit of mass used in pharmaceutical work is the grain (gr). 15 gr = 1.0 g. An aspirin
tablet contains 5.0 gr of aspirin. A 145-lb person takes two aspirin tablets.
a) What is the quantity of aspirin taken, expressed in milligrams?
b) What is the dosage rate of the aspirin, expressed in milligrams of aspirin per kilogram of body
weight?
6) A block of ice measures 24 in. x 36 in. x 18 in.
a) What is the volume of this block in cubic meters?
b) What is the total surface area of the block in square centimeters?
7) Many times errors are expressed in terms of percentage. The percentage error is the absolute value
of the difference of the accepted value and the experimental value, divided by the accepted value,
and multiplied by 100.
accepted value - experiment al value
Percent Error 
x 100
accepted value
Calculate the percent error for the following measurements.
a) The density of an aluminum block determined in an experiment was 2.64 g/mL. (Accepted
value = 2.70 g/mL)
b) The experimental determination of iron in iron ore was 16.48%. (Accepted value = 16.12%)
c) A balance measured the mass of 1.0000 g standard as 0.9981 g.
8) How many protons, neutrons and electrons are in each of the following?
a) 227 Ac
f) 64 Cu
k) 127 I 53
3
b) 70 Ga
g) 56
l) 127 I 7 
26 Fe
53
c) 11 B
h) 40 Ca 2 
20
d) 251 Cf
i) 19 F 9
e) 239 Pu
j) 31 P 315
15
9) The density of water is 1.0 g/cm3. Express this value in units of kg/m3 and lb/ft3.
10) Diamonds are measured in carats, and 1 carat = .200 g. The density of diamond is
3.51 g/cm3. What is the volume of a 5.0 carat diamond?
11) In the opening scenes of the movie Raiders of the Lost Ark, Indiana Jones tries to
remove a gold idol from a booby-trapped pedestal. He replaces the idol with a bag of
sand of approximately equal volume. (Density of gold = 19.32 g/mL; density of sand
≈ 2 g/mL.)
1. Did he have a reasonable chance of not activating the masssensitive booby trap?
2. In a later scene he and an unscrupulous guide play catch with the
idol. Assume that the volume of the idol is about 1.0 L. If it
were solid gold, what mass would the idol have? Is playing
catch with it plausible? Why or why not?
12) An experiment was performed in which an empty 100 mL graduated cylinder was
weighed. It was weighed once again after it had been filled to the 10.0 mL mark with
dry sand. A 10 mL pipet was used to transfer 10.00 mL of methanol to the cylinder.
The sand-methanol mixture was stirred until bubbles no longer emerged from the
mixture and the sand looked uniformly wet. The cylinder was then weighed again.
Use the data obtained from this experiment (and displayed below) to find the density
of dry sand, the density of methanol, and the density of sand particles. Does the
bubbling that occurs when the methanol is added to the dry sand indicate that the sand
and methanol are reacting?
Mass of cylinder plus wet sand
Mass of cylinder plus dry sand
Mass of empty cylinder
Volume of dry sand
Volume of sand + methanol
Volume of methanol
45.2613 g
37.3488 g
22.8317 g
10.0 mL
17.6 mL
10.00 mL
13) The German chemist Fritz Haber proposed paying off the reparations imposed against
Germany after World War I by extracting gold from seawater. Given that (a) the
amount of the reparations was 28.8 billion dollars, (b) the value of gold at the time
was about $21.25 per troy ounce (12 troy ounces = 1 lb), and (c) gold occurs in
seawater to the extent of 4.67 x 1017 atoms per ton of seawater (1 ton = 2000 lb), how
many cubic kilometers of seawater would have had to be processed to obtain the
required amount of gold? Assume that the density of seawater is 1.03 g/cm3.
Summer Assignment #2
Stoichiometry
These problems must be worked out on a separate page and all work
shown for credit.
1. Benzene contains only carbon and hydrogen and has a molar mass of 78.1
g/mol. Analysis shows the compound to be 7.74% H by mass. Find the
empirical and molecular formulas of benzene.
2. Find the mass percent of nitrogen in each of the following compounds:
a. NO
b. NO2
c. N2O4
d. N2O
3. Calcium carbonate decomposes upon heating, producing calcium oxide and
carbon dioxide gas.
a. Write a balanced chemical equation for this reaction.
b. How many grams of calcium oxide will remain after 12.25 g of calcium
carbonate is completely decomposed?
c. What volume of carbon dioxide gas is produced from this amount of
calcium carbonate? The gas is measured at 0.95 atm and 10C.
4. Hydrogen gas and bromine gas react to form hydrogen bromide gas.
a. Write a balanced chemical equation for this reaction.
b. How many grams of hydrogen bromide gas can be produced from 3.2 g
of hydrogen gas and 9.5 g of bromine gas?
c. How many grams of which reactant is left unreacted?
d. What volume of HBr, measured at STP, is produced in b)?
5. When ammonia gas (NH3), oxygen gas (O2) and methane gas (CH4) are
combined, the products are hydrogen cyanide gas (HCN) and water.
a. Write a balance chemical equation for this reaction.
b. Calculate the mass of each product produced when 225 g of oxygen gas
is reacted with an excess of the other two reactants.
c. If the actual yield of the experiment in b) is 105 g of HCN, calculate
the percent yield.
6. A 2.29 g sample of an unknown acid is dissolved in 1.0 liter of water. A
titration required 25.0 ml of 0.500 M NaOH to completely react with all
the acid present. What is the molar mass of the acid?
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7. What mass of aluminum hydroxide is produced when 50.0 ml of 0.200 M
Al(NO3)3 reacts with 200.0 ml of 0.100 M KOH?
8. Cinnamic acid contains only carbon, hydrogen and oxygen, and is found by
analysis to be 73.0% C and 5.4% hydrogen. In a titration, 18.02 ml of
0.135M NaOH is found to neutralize 0.3602 g of cinnamic acid.
a. Find the empirical formula of this compound.
b. Find the molar mass of this compound.
c. Write the molecular formula for this compound.
9. Potassium nitrate decomposes when heated, forming potassium nitrite and oxygen
gas.
a. Write a balanced chemical equation for this reaction.
b. What mass of KNO3 would be needed to produce 18.4 liters of oxygen
gas, measured at 775 mmHg and 15C?
c. What mass of KNO2 would also be produced?
10. A compound contains only carbon, hydrogen and oxygen. Combustion of
10.68 mg of the compound yields 16.01mg CO2 and 4.37 mg H2O. The
molar mass of the compound is 176.1 g/mole. What are the empirical and
molecular formulas of the compound?
18
AP Chemistry Summer Review #3
Writing Net Ionic Equations
Your first review assignment reminded you of how to write formulas and balance
equations. For the AP Chemistry exam, all equations must be written in net ionic form
and you must be able to recognize types of reactions and then predict products.
We will be writing net ionic equations all year and our beginning lab work will involve
recognizing types of reactions and predicting products. The pages that follow give
tutorials for recognizing types of reactions. The first section of this assignment will have
you predicting products for a given type of reaction and then writing balanced equations.
The final sections of this assignment will be a “mixed review” where you will have to
recognize the type of reaction before being able to predict the products and write the
balanced equations! You will also be asked to write net ionic equations, rather than just
molecular equations.
Write the balanced molecular equations for exercises 1, 2, 3, 4, and 5. As you do these
exercises you will be learning your solubility rules, strong and weak electrolytes, and
decomposition products.
Write the balanced molecular equations for exercise 6, the write the balanced overall
ionic equations for exercise 6 and then write the balanced net ionic equations for
exercise 6.
Now go back to your answers to exercises 1, 2, 3, 4, and 5 and write balanced net ionic
equations for each problem that goes to completion. (Remember, sometimes the overall
ionic and net ionic equations are the same; sometimes the molecular equation is all there
is and there are no ions at all!
Do the equations in order and on separate paper not the pages of the tutorial packet.
Bring the tutorial pages to class on the second class in September and we will discuss the
work.
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Exercise 1. Synthesis and Decompostion Reactions
Synthesis reactions occur when two or more reactants combine to form a single product.
There are several common types of synthesis reactions.
A metal combines with a nonmetal to form a binary salt.
Example: A piece of lithium metal is dropped into a container of nitrogen gas.
6Li + N2 → 2Li3N
Metallic oxides and water form bases (metallic hydroxides).
Example: Solid sodium oxide is added to water.
Na2O + HOH  2NaOH
Example: Solid magnesium oxide is added to water.
MgO + HOH → Mg(OH)2
Nonmetallic oxides and water form acids. The nonmetal retains its oxidation number.
Example: Carbon dioxide is bubbled into water.
CO2 + H2O → H2CO3 (Oxidation number of C = +4)
Example: Dinitrogen pentoxide is bubbled into water.
N2O5 + H2O → 2HNO3
Metallic oxides and nonmetallic oxides form salts.
Example: Solid sodium oxide is added to carbon dioxide.
Na2O + CO2 → Na2CO3
Example: Solid calcium oxide is added to sulfur trioxide.
CaO + SO3 → CaSO4
Decomposition reactions occur when a single reactant is broken down into two or more
products.
Metallic carbonates decompose into metallic oxides and carbon dioxide.
Example: A sample of magnesium carbonate is heated.
MgCO3 → MgO and CO2
Metallic chlorates decompose into metallic chlorides and and oxygen.
Example: A sample of magnesium chlorate is heated.
Mg(ClO3)2 → MgCl2 + 3O2
Ammonium carbonate decomposes into ammonia, water and carbon dioxide.
Example: A sample of ammonium carbonate is heated.
(NH4)2CO3 → 2NH3 + H20 + CO2
20
Sulfurous acid decomposes into sulfur dioxide and water.
Example: A sample of sulfurous acid is heated.
H2SO3 → H2O + SO2
Carbonic acid decomposes into carbon dioxide and water.
Example: A sample of carbonic acid is heated.
H2CO3 → H2O + CO2
Ammonium hydroxide decomposes into ammonia and water.
Example: A sample of ammonium hydroxide is heated.
NH4OH  NH3 + H2O
Hydrogen peroxide decomposes into water and oxygen.
Example: A sample of hydrogen peroxide is heated.
2H2O2  2H2O + O2
A binary compound may break down into two elements.
Example: Molten sodium chloride is electrolyzed.
2NaCl  2Na + Cl2
Your Turn! Predict and balance the following synthesis and decomposition
reactions. Use abbreviations to indicate the phase (aq, s, l, or g) of reactants and
products where possible. Write answers on a separate page!
1. A sample of calcium carbonate is heated.
2. Sulfur dioxide gas is bubbled through water.
3. Solid potassium oxide is added to a container of carbon dioxide gas.
4. Liquid hydrogen peroxide is warmed.
5. Solid lithium oxide is added to water.
6. Molten aluminum chloride is electrolyzed.
7. A pea-sized piece of sodium is added to a container of iodine vapor.
8. A sample of carbonic acid is heated.
9. A sample of potassium chlorate is heated.
10. Solid magnesium oxide is added to sulfur trioxide gas.
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Exercise 2. Single Replacement Reactions
Single replacement reactions are reactions that involve an element replacing one part of
a compound. The products include the displaced element and a new compound. An
element can only replace another element that is less active than itself.
General activity series for metals:
(Most active) Li Ca Na Mg Al Zn Fe Pb [H2] Cu Ag Pt (Least active)
General activity series for nonmetals:
(Most active) F2 Cl2 Br2 I2 (least active)
Here are some common types of single replacement reactions.
Active metals replace less active metals from their compounds in aqueous solution.
Example: Magnesium turnings are added to a solution of iron (III) chloride.
3Mg + 2FeCl3  2 Fe + 3MgCl2
Active metals replace hydrogen in water.
Example: Sodium is added to water.
2Na + 2HOH  H2 + 2NaOH
Active metals replace hydrogen in acids.
Example: Lithium is added to hydrochloric acid.
2Li + 2HCl  H2 + 2LiCl
Active nonmetals replace less active nonmetals from their compounds in aqueous
solution.
Example: Chlorine gas is bubbled into a solution of potassium iodide.
Cl2 + 2KI  I2 + 2KCl
If a less reactive metal is combined with a more reactive element in compound form,
there will be no resulting reaction.
Example: Chlorine gas is bubbled into a solution of potassium fluoride.
Cl2 + KF  no reaction
Example: Zinc is added to a solution of sodium chloride.
Zn + NaCl  no reaction
NOTE: On the AP reaction prediction section, all reactions “work”; in other words,
there will be no “no reaction” problems on the AP exam.
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Your turn! Using the activity series, predict and balance the following single
replacement reactions. Use abbreviations to indicate the appropriate phase of
reactants and products where possible. Write answers on a separate page! Note: Not
all of the reactions will occur. For those that do not, write “no reaction”.
1. A piece of copper is dropped into a container of water.
2. Liquid bromine is added to a container of sodium iodide crystals.
3. An aluminum strip is immersed in a solution of silver nitrate.
4. Zinc pellets are added to a sulfuric acid solution.
5. Fluorine gas is bubbled into a solution of silver nitrate.
6. Magnesium turnings are added to a solution of lead (II) acetate.
7. Iodine crystals are added to a solution of sodium chloride.
8. Calcium metal is added to a solution of nitrous acid.
9. A pea-sized piece of lithium is added to water.
10. A solution of iron (III) chloride is poured over a piece of platinum wire.
Exercises 3, 4 and 5. Double Replacement (Metathesis) Reactions
In many reactions between two compounds in aqueous solution, the cations and anios
appear to switch partners according to the following equation:
AX + BY → AY + BX
The two compounds react to form two new compounds. No changes in oxidation
numbers occur. Reactions of this type are known as double replacement or metathesis
reactions. An example of such a reaction would be the mixing of aqueous solutions of
potassium bromide and silver nitrate forming insoluble silver bromide (precipitate) and
aqueous potassium nitrate;
KBr (aq) + AgNO3 (aq) → AgBr (s) + KNO3 (aq)
Note that each cation pairs up with the anion in the other compound, thus switching
partners. Anions do not pair up with anions and cations do not pair up with cations.
Likes repel; opposites attract!
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All double replacement reaction must have a “driving force” or a reason why the reaction
will occur or “go to oompletion”. The “driving force” in metathesis reactions is the
removal of a least one pair of ions from solution. This removal can occur in one of three
ways:
1. Formation of a precipitate: A precipitate is an insoluble substance (solid) formed by
the reaction of two aqueous substances. It is the result of ions binding together so
strongly that the solvent (water) cannot pull them apart. The insoluble solid (or solids if a
double precipitate occurs!) will settle out (precipitate) from the solution and this results in
the removal of ions from solution.
2. Formation of a gas: Gases may form directly in a double replacement reaction or
from the decomposition of one of the products. The gases will bubble off or evolve from
the solution.
3. Formation of primarily molecular species: The formation of primarily unionized
molecules in solution removes ions from the solution and the reaction “works” or is said
to go to completion. Unionized or partially ionized molecules give solutions that are
known as nonelectrolytes or weak electrolytes. The best know nonelectrolyte is water
formed in acid-base neutralization reactions. Acetic acid is an example of an acid that is
primarily molecular (weak electrolyte) when placed in water.
Reversible Reactions
If a double replacement reaction does not go to completion (no precipitate, gas or
molecular species is formed), then the reaction is reversible (no ions have been removed).
Reversible reactions are at equilibrium and have both forward and reverse reactions
taking place. In a reversible reaction, evaporation of the water solvent will result in solid
residues of both reactants and products. The reaction is not driven to completion
(products) because no ions have been removed. A double arrow is used to designate a
reversible reaction at equilibrium.
BaCl2 (aq) + 2 NaNO3 (aq)  Ba(NO3)2 (aq) + 2 NaCl (aq)
Solubility Rules
The solubility classification of ionic substances according to their solubility in water is
difficult. Nothing is completely “insoluble” in water. The degree of solubility varies
from one “soluble” substance to another. Nevertheless, a solubility classification scheme
is useful even though it must be regarded as an approximate guideline. Learn the
solubility rules!
Formation of a Gas
Common gases formed in metathesis reactions are:
H2S: Any sulfide (salt of S2-) plus any acid for H2S (g) and a salt.
CO2: Any carbonate (salt of CO32-) plus any acid for CO2 (g) , HOH and a salt.
SO2: Any sulfite (salt of SO32-) plus any acid form SO2 (g), HOH and a salt.
24
NH3: Any ammonium salt (salt of NH4+) plus any soluble strong hydroxide react upon
heating to form NH3 (g), HOH and a salt.
Reactions that produce three of the gases (CO2, SO2, and NH3) involve the initial
formation of a substance that breaks down to give the gas ad HOH.
Example: The reaction of Na2SO3 and HCl produces H2SO3.
Na2SO3 (aq) + 2HCl (aq) → H2SO3 (aq) + 2 NaCl (aq)
Bubbling is observed in this reaction because the H2SO3 (sulfurous acid) is
unstable and immediately decomposes to give HOH and SO2 gas:
H2SO3 (aq) → HOH (l) and SO2 (g)
The molecular equation for the complete reaction, therefore, is
Na2SO3 (aq) + 2HCl (aq) → HOH (l) + SO2 (g) + 2 NaCl (aq)
Example: A typical reaction of a carbonate and an acid is:
K2CO3 (aq) and 2HNO3 (aq) → HOH (l) + CO2 (g) + 2KNO3 (aq)
Bubbling is also observed in this reaction. Theoretically, H2CO3,
carbonic acid, is formed, but the acid is unstable and immediately
decomposes to form carbon dioxide gas and water according to the
following equation:
H2CO3 → HOH (l) and CO2 (g)
Example: Ammonium salts and soluble bases react as follows (particularly when the
solution is warmed):
NH4Cl (aq) + NaOH (aq) → NH3 (g) +HOH (l) + NaCl (aq)
The odor of ammonia gas is noted and moist red litmus paper held near the
mouth of the container will turn blue. Theoretically, NH4OH, ammonium
hydroxide, is produced (also known as ammonia water). The compound is
unstable and decomposes into ammonia gas and water:
NH4OH (aq) → NH3 (g) + HOH (l)
Example: The odor of rotten eggs and bubbling are noted when an acid is added to a
sulfide. A typical reaction producing hydrogen sulfide gas is:
FeS (s) + 2HCl (aq) → FeCl2 (aq) + H2S (g)
NOTE: Be aware of reactions involving the formation of carbon dioxide, sulfur dioxide,
ammonia and hydrogen sulfide gases on the AP exam. Over the years these reactions
have appeared many, many times. Know these four gases and how they are produced!
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Formation of a Molecular Species (Weak or Nonelectrolytes)
Metathesis reactions that produce primarily molecules in the form of partially dissociated
or ionized molecules (weak electrolytes) or molecules that do not ionized of dissociated
at all (nonelectrolytes) serve as the driving force in some aqueous reactions. Forming
molecular products in double replacement reactions results in the removal of ions from
solution. Such reactions tend to go to completion (shift to the right) and form primarily
products.
General rules:
1. The six common strong acids, and thus strong electrolytes, are HClO4, HCl, HBr, HI,
HNO3 and H2SO4. (Memorize these!) All other common acids are weak acids and thus
weak electrolytes. (HC2H3O2 or CH3COOH, H3PO4, HF and HNO2 are examples of
weak acids. Note: All organic acids (R-COOH) are weak electrolytes.) All strong acids
in their pure form (as opposed to dilute aqueous form!) are nonelectrolytes (molecular).
When water is added, the action of the solvent water with a strong acid produces a
hydrated proton (hydronium ion) and a negatively charged anion. The process of making
ions from molecular species is known as ionization. Strong acids ionize 100% in water.
An example of a strong electrolyte undergoing ionization is as follows:
HCl (aq) + H2O (l) → H3O+ (aq) + Cl- (aq)
This reaction may be abbreviated as:
HCl (aq) → H+ (aq) + Cl- (aq)
2. The common strong bases are the soluble hydroxides (those of Group IA elements and
Ba2+) and the slightly soluble hydroxides (those of Ca2+ and Sr2+). Strong bases, like
strong acids, are strong electrolyte. (Memorize these!) NH4OH is a soluble weak
electrolyte which normally decomposes into NH3 (g) and HOH (l). Technically speaking,
the pure compound ammonium hydroxide has never been isolated and the substance is
more correctly known as aqueous ammonia. Most other hydroxides are insoluble. Pure
liquid hydroxides are strong hydroxides because they already contain ions. The action of
the solvent water releasing the ions of a base into solution is known as dissociation.
Acids ionize in water; bases dissociate.
3. Most common (soluble) salts are strong electrolytes and thus dissociate inot ions when
placed into water.
4. Water is a weak electrolyte which is typically produced in acid-base neutralization
reactions.
Some examples of weak electrolytes forming as products (shown in bold):
Ca(CH3COO)2 (aq) + 2HCl (aq) → CaCl2 (aq) + 2CH3COOH (aq)
2Na3PO4 (aq) + 3H3PO4 (aq) → 3Na2SO4 (aq) + 2H3PO4 (aq)
HCl (aq) + NaOH (aq) → NaCl (aq) + HOH (l)
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Acid-Base Neutralization Reactions
Acids react with bases to produce salts and water. One mole of hydrogen ions will react
with one mole of hydroxide ions to produce one mole of water. Learn which acids are
strong (written in ionic form!) and which acids are weak (written in molecular form)!
Check the solubility rules for the solubility of the salt produced! If it is soluble, it is
written in ionic form; if it is insoluble it is written in molecular form.
Acid + Base → Salt + Water
(A salt consists of a cation from a base and an anion from an acid. For example, the salt
sodium sulfate contains sodium ions form sodium hydroxide and sulfate ions from
sulfuric acid.)
Example: Hydrogen sulfide gas is bubbled through excess potassium hydroxide solution.
H2S (g) + KOH (aq) → K2S (aq) + 2HOH (l)
Polyprotic acids can be tricky when it comes to predicting neutralization reactions.
Sulfuric acid and phosphoric acid are classic examples frequently encountered on AP
exams. If the base is in excess, all hydrogen ions will react with the strong base to
produce water.
Example: Dilute sulfuric acid is reacted with excess sodium hydroxide solution.
H2SO4 (aq) + 2NaOH (aq) → Na2SO4 (aq) + 2HOH (l)
If, however, the reaction above stated that equal numbers of moles of sulfureic acid and
sodium hydroxide react, then the coefficients for both reactants must be one and the salt
that forms is sodium hydrogen sulfate.
Example: Equal number of moles of sulfuric acid and sodium hydroxide solution mix.
H2SO4 (aq) + NaOH (aq) → NaHSO4 (aq) + HOH (l)
Take into account information dealing with the quantity of each reactant!
Example: Equal volumes of 0.1M phosphoric acid and 0.2M sodium hydroxide are
mixed together.
H3PO4 (aq) + 2NaOH (aq) → Na2PO4 (aq) + 2HOH (l)
Watch for substances that react with water before reacting with an acid or a base. The
acid and basic anhydrides behave in such a manner. These are really two-step reactions.
Example: Excess sulfur dioxide gas is bubbled into a saturated solution of calcium
hydroxide.
SO2 (g) + Ca(OH)2 (aq) → CaSO3 (s) + HOH (l)
(Remember, SO2 (g) is an acid anhydride!)
27
If an acid + base yields a salt + water, then an acid anhydride + basic anhydride will
yield a salt.
Example: Sulfur dioxide gas and solid calcium oxide are reacted together.
SO2 (g) + CaO (s) → CaSO3 (s)
(SO2 (g) is the acid anhydride for sulfurous acid; CaO (s) is the basic anhydride
for calcium hydroxide.)
Exercise 3. Formation of a Precipitate
Predict and balance the following metathesis reactions based on the solubility of the
products. Use the abbreviations (aq) and (s) for the reactants and the products. All
reactants are aqueous. Write answers on a separate page! Note: Some of these
reactions do not go to completion! Correct molecular formulas must be written for
both reactants and products before an equation may be balanced.
1. silver nitrate + postasium chromate
2. ammonium chloride + cobalt (II) sulfate
3. lithium hydroxide + sodium chromate
4. zinc acetate + cesium hydroxide
5. ammonium sulfide + lead (II) nitrate
6. iron (III) sulfate + barium iodide
7. chromium (III) bromide + sodium nitrate
8. rubidium phosphate + titanium (IV) nitrate
9. ammonium carbonate + nickel (II) chloride
10. tin (IV) nitrate + potassium sulfite
Exercise 4. Formation of a Gas
Predict and balance the following metathesis reactions based on the solubility of the
products. Use the abbreviations (aq), (s), (l) and (g) for the reactants and the
products. All reactants are aqueous unless otherwise stated. Write answers on a
separate page!
1. ammonium sulfate and potassium hydroxide are mixed together
2. ammonium sulfide is reacted with hydrochloric acid
3. cobalt (II) chloride is combined with silver nitrate
28
4. solid calcium carbonate is reacted with sulfuric acid
5. potassium sulfite is reacted with hydrobromic acid
6. potassium sulfide is reacted with nitric acid
7. ammonium iodide + magnesium sulfate
8. solid titanium (IV) carbonate + hydrochloric acid
9. solid calcium sulfite + acetic acid
10. strontium hydroxide + ammonium sulfide
Exercise 5.
Predict and balance the following metathesis reactions based on the solubility of the
products. Use the abbreviations (aq), (s), (l) and (g) for the reactants and the
products. All reactants are aqueous unless otherwise stated. Write answers on a
separate page!
1. Carbon dioxide gas is bubbled through a solution of lithium hydroxide
2. Sodium nitrite is reacted with hydrochloric acid
3. Ammonium bromide + sodium hydroxide
4. Carbon dioxide gas is reacted with solid potassium oxide
5. Solid magnesium oxide is reacted with solid potassium oxide
6. Equal numbers of moles of potassium hydroxide and phosphoric acid react
7. sodium chloride reacts with silver nitrate
8. ammonium carbonate + potassium bromide
9. oxalic aid (0.l M) reacts with an equal volume of cesium hydroxide (0.1M)
10. silver nitrate + sodium chromate
Aqueous Solutions and Ionic Equations
It is now time to concentrate on writing chemical equations in the form required for the
AP Chem exam! All equations in the previous sections were written as if the reactants
and products were molecular. In overall (total) ionic equations, formulas of the
reactants and products are written to show the predominant form of each substance as it
29
exists in aqueous solution. Soluble salts, strong acids and strong bases are written as
separated ions. Insoluble salts, suspensions, solids, weak acids, weak gases, gases, water
and organic compounds are always written as individual molecules. Consider the
following molecular equation:
Cd(NO3)2 (aq) + Na2S (aq) → CdS (s) + 2NaNO3 (aq)
The overall ionic equation for this reaction is
Cd2+ (aq) + 2NO3- (aq) + 2Na+ (aq) + S2- (aq) → CdS (s) + 2Na+(aq) + 2NO3- (aq)
Note that this equation illustrates that cadmium nitrate dissociates into three ions (one
cadmium ion and two nitrate ions); sodium sulfide dissociates into two sodium ions and
one sulfide ion; the soluble sodium nitrate formed remains dissociated as two sodium
and two nitrate ions. The precipitated (insoluble) cadmium is undissociated. Also note
that the parentheses that appear in a molecular formula are not used when representing
the ionic form in solution. For example, dissociated cadmium nitrate contains no
parentheses.
Note: The only common substances that should be written as ions in ionic equations
are soluble salts, strong acids and strong bases.
Net ionic equations are written to show only the species that react or undergo change in
aqueous solution. The net ionic equation is obtained by eliminating the spectator ions
form an overall ionic equation. All that is left are the ions that have changed chemically.
Spectators at a sporting event watch the action unfolding in front of them rather than
participating; spectator ions likewise do not participate in the reaction. The elimination
of spectator ions allows us to concentrate only on the reacting species.
Molecular equations provide complete chemical formulas which are of necessity when it
comes to doing stoichiometric calculations. Overall ionic equations, the intermediated
between molecular and net ionic equations, show what is happening to all species in the
solution. Such equations are very helpful when dealing with hydrolysis, electrical
conductivity, and colligative properties. Net ionic equations are the simplest form of
equations and show only the reacting species.
Molecular:
Cd(NO3)2 (aq) + Na2S (aq) → CdS (s) + 2NaNO3 (aq)
Overall Ionic
Cd2+ (aq) + 2NO3- (aq) + 2Na+ (aq) + S2- (aq) → CdS (s) + 2Na+(aq) + 2NO3- (aq)
(The spectator ions have been highlighted in bold!)
Net Ionic
Cd2+ (aq) + S2- (aq) → CdS (s)
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Example: Aqueous solutions of sulfuric acid and excess sodium hydroxide are reacted.
H2SO4 (aq) + 2NaOH (aq) → 2 HOH (l) + Na2SO4 (aq)
H+ + HSO4- + 2Na+ +2OH- → 2HOH (l) + 2Na+ (aq) + SO42- (aq)
H+ + HSO4- + 2OH- → 2HOH (l) + SO42- (aq)
Special note: The first proton in sulfuric acid is ionized completely; the second
proton is only partially ionized. Sulfuric acid is the only polyprotic acid that
exhibits this property. All other polyprotic acids are weak and are written in their
molecular forms!
Example: Aqueous solutions of sulfuric acid and excess sodium hydroxide are
combined.
H2SO4 (aq) + Ba(OH)2 (aq) → 2 HOH (l) + BaSO4 (s)
H+ + HSO4- + Ba2+ +2OH- → 2HOH + BaSO4
H+ + HSO4- + Ba2+ +2OH- → 2HOH + BaSO4
Example: Sodium acetate undergoes hydrolysis when placed in water.
NaCH3COO (aq) + HOH (l) → NaOH (aq) + CH3COOH (aq)
Na+ + CH3COO- + HOH → Na+ + OH- + CH3COOH
CH3COO- + HOH → OH- + CH3COOH
Example: Equal volumes of 0.2 M potassium hydroxide and 0.2 M phosphoric acid are
reacted.
KOH (aq) + H3PO4 (aq) → HOH (l) + KH2PO4 (aq)
K+ + OH- + H3PO4 → HOH (l) + K+ + H2PO4OH- + H3PO4 → HOH (l) + H2PO4Example: Aqueous solutions of potassium chromate and silver nitrate are reacted.
K2CrO4 (aq) + 2AgNO3 (aq) → 2KNO3 (aq) + Ag2CrO4 (s)
2K+ + CrO42- + 2Ag+ + 2NO3- → 2K+ + 2NO3- + Ag2CrO4 (s)
CrO42- + 2Ag+ → Ag2CrO4 (s)
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Example: Aqueous ammonium sulfide reacts with excess aqueous lithium hydroxide.
(NH4)2S (aq) + 2LiOH (aq) → Li2S (aq) + 2NH3 (g) + 2HOH (l)
2NH4- + S- + 2Li+ + OH- → 2Li+ +S2- + 2NH3 (g) + 2HOH (l)
2NH4- + OH- → 2NH3 (g) + 2HOH (l)
Example: Excess acetic acid reacts with aqueous ammonium sulfite.
2CH3COOH (aq) + Na2SO3 (aq) → 2NaCH3COO (aq) + HOH (l) + SO2 (g)
2CH3COOH + 2Na+ + SO3- → 2Na+ + 2CH3COO- + HOH + SO2
2CH3COOH + SO3- → 2CH3COO- + HOH + SO2
Example: Calcium carbonate in aqueous suspension reacts with dilute hydrochloric acid.
CaCO3 (s) + 2HCl (aq) → CaCl2 (aq) + HOH (l) + CO2 (g)
CaCO3 + 2H+ + 2Cl- → Ca2+ +2Cl- + HOH + CO2
CaCO3 + 2H+ → Ca2+ + HOH + CO2
Exercise 6.
Write balanced molecular, overall ionic and ionic equations for the following
reactions. Write answers on a separate page!
1. aqueous nickel (II) nitrate + aqueous cesium hydroxide
2. equal volumes of equal molar concentrations of sulfuric acid and sodium hydroxide
3. solid potassium chlorate is strongly heated
4. potassium tartrate solution + water
5. solid lithium metal is added to water
6. aqueous solutions of magnesium nitrate and sodium bromide are mixed together
7. aqueous solutions of oxalic acid and excess potassium hydroxide
8. solid cobalt (II) hydroxide + hydroiodic acid
9. aqueous solutions of manganese (II) sulfate undergoing hydrolysis
10. aqueous sodium carbonate + chlorous acid
11. aqueous solutions of potassium phosphate and excess hydrobromic acid
(Modified from “Chemical Equations Handbook”, Flinn Scientific)
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In addition please answer the following questions on a separate sheet of paper.
Chapter 1
1. Explain the difference between a law and a theory. What is the law of
conservation of mass?
2. Explain the difference between Qualitative and Quantitative data. Is an
experiment designed to determine the identity of a compound by its color
quantitative or qualitative?
3. Explain the difference between an extensive property and an intensive property.
Give an example of each
4. Classify the following mixtures as heterogeneous or homogenous
a. Salt water
b. Lemonade
c. Italian salad dressing
d. Potassium Hydroxide in water
e. Soil
f. Hershey bar with almonds
5. State whether a Physical or Chemical property is being described
a. Ice melts at temperatures above 0oC
b. Ethane gas combusts in air
c. Liquid Toluene is more dense than liquid ethanol
d. Helium balloons rise
e. Gold does not dissolve in Nitric Acid
6. Classify the following as physical or chemical changes
a. Water boiling on a stove
b. A piece of Iron rusting in the presence of water
c. Sodium metal reacting with water to produce Hydrogen gas
d. Cesium metal melts in the palm of your hand
7. Define Precision and Accuracy. Which one implies the other for a set of data?
8. If a liquid is being measured in a graduated cylinder and the meniscus is at the
55mL mark and it reaches the 56mL mark on the sides. What volume of the liquid
is in the graduated cylinder?
9. A 14.3g piece of silver is dropped in a graduated cylinder with water and it
displaced 1.36mL of water. What is the density of Silver?
33
10. Express the following numbers in scientific notation
a. .00098
b. 341281.4
c. 810
d. .0975
e. 3.4
f. 3/40
11. Express the following numbers in decimal notation.
a. 3.75 x 10-4
b. 8.1 x 10-3
c. 4.9 x 101
d. 8.51 x 10-7
e. 1.05 x 10-1
12. Classify the following numbers as exact or inexact.
a. There are 100cm in one meter.
b. There are 2.54cm one inch
c. The mass of the earth is approximately 6 x 1024kg
d. There are 433 students in the WHRHS 2006 graduating class.
e. A student measures out 850mL of water in a 1 liter beaker.
13. If a volume of a liquid is measured out in a 100mL graduated cylinder with each
milliliter marked, which digit is uncertain?
14. Which of the following measuring devices is best for measuring out 17mL of a
liquid?
a. A 25mL volumetric pipette
b. A 50mL graduated burette
c. A 100mL volumetric flask
d. A 50mL beaker
15. How many significant figures are in each of the following numbers
a. .0121
b. 4570
c. 9.901
d. 2.8 x 104
e. 5
f. .1
g. 2401
h. 0020
i. 951.1
16. How many significant figures are in each of the following numbers?
a. 8.01 x 10-4
b. 700
34
c. .0805
d. 6.626 x 10-34
e. 200.
f. 2.9979 x 108
g. 2.178 x 10-18
h. 100.78
i. 6.022 x 10-23
j. 46.485
k. 9.11 x 10-31
17. Express the number 96500 in
a. 3 sig figs
b. 2 sig figs
c. 1 sig fig
d. 4 sig figs
e. 5 sig figs
f. 6 sig figs
18. Express the answers of the following calculations in the correct number of
significant figures
a. 3.5 x 4.0
b. 10 x 21
c. 12.4/4.00
d. 3.94 x 102 x 8.01
e. 6.626 x 10-34 x 9.10 x 10-15
f. 2.998 x 10 8/450
g. 1.602 x 10-19 x 6.022 x 1023
h. 298.15 x .08206 x 1.0/5.4
19. Express the answer to the following calculations in the correct number of
significant figures
a. 37.1 + 4.8
b. 8.01 – 5.8
c. 32 + 51.7
d. 89.257 – 87
e. 197 – 4.00
f. 492.015- 301.4
20. Express the answers to the following calculations in the correct number of
significant figures.
a. 3.31/(29.5 – 26.2)
b. (39.2 x 8.48) + (8.97 x 9.1)
c. (21.1 + 8.1)(22.8 – 14.7)
d. 8 x 90 – 41.4
e. ((42.97-41.01)/42.97) x 100 (100 is exact)
35
21. In an experiment to find the mass of a solid compound in a mixture, the result
from five trials are 3.47g, 3.56g, 3.44g, 3.49g, and 3.50g. What is the average of
the data?
22. The relationship between frequency and wavelength of light is given by the
following equation
vλ = c
If c = 3.00 x 108 and v = 9.3 x 1014 what is λ.
23. Express the answer to the following calculations in the correct # of significant
figures.
a. 9.0 x 1012 x 8.11 x 10-9
b. 4.0 x 10-7 x 1.12-4
c. (9.0 x 109)(1.602 x 10-19)2/(5.29 x 10-11) x ½ (1/2 is exact).
d. (2.178 x 10-18)(1/4 – 1/9)
e. (9.81)(9.0) + 8.2 x 101
f. ((9.11 x 10-31)(9.81) + (1.67 x 10-27)(2.02 x 102)) / 2.18
24. Carry out the following conversions
a. 90.1mL to liters
b. 5.0km to meters
c. 5.01kg to grams
d. 9.2 Megagrams to micrograms
e. 456 nanometers to meters
f. 5 teragrams to kilograms
25. Carry out the following conversions
a. 8.47 x 104cm3 to m3
b. 96ft2 to in2
c. 14.0yd2 to ft2
d. 29dm3 to mL
e. 4.0L to m3
26. If you see a sign that says your are 97km from Warren, and you are driving 65
miles per hour, how long, in seconds, will it take you to get to Warren.
27. In Norway gasoline costs 1.01 euros per Liter. Convert this to dollars per 9/10
gallon. (1.28 dollars = 1 euro)
28. If the density of aluminum is 2.70 g/mL, what is its density in pounds per cubic
yard
29. Does 2.08cm3 of platinum or 4.84cm3of iron weigh more?
36
30. At 250C, water has a density of 1.00g/cm3 and another liquid which does not mix
with water has a density of 6.02kg/ft3. Which one layers on top of the other when
they are poured into a beaker at the same time?
31. A box has dimensions 4.2cm x 3.9cm x 6.5cm has a mass of 138g. What are the
volume and the density of the box?
32. Kobe Bryant bought his wife an 8-carat, $4 million ring. If the density of diamond
is 3.51g/cm3, how big is the diamond? How many dollars per cubic cm did Kobe
pay for the ring? (1 carat = .200g). If someone spent that much per cubic
centimeter on you, would you forgive him?
33. A spherical star has a mass of 3.0 x 1031kg and a radius of 2.3 x 108m. What is the
star’s density in g/cm3? (Vsphere = 4/3πr3).
34. The mass of the Earth is 5.98 x 1024kg and its radius is 6.37 x 106m. What is the
density of Earth in Megagrams per cubic foot?
35. Osmium, the densest metal, has a density of 22.57g/cm3. If a 205g sheet of
Osmium foil has a surface area of 3.85ft2, What is the thickness of the foil in
inches?
36. Convert the following temperatures to Fahrenheit
a. 00C
b. 1000C
c. 58.00C
d. -273.150C
e. 37.00C
37. Convert the following temperatures to Fahrenheit
a. 300.K
b. 298.15K
c. 373.15K
d. .0001K
e. 210K
38. Convert the following to Celsius
a. 420F
b. 1210F
c. 51.90F
d. 82.00F
e. 91.30F
39. Convert the following to Celsius
a. 402K
b. 301K
c. 288.15K
d. 550K
37
40. Convert the following temperatures to Kelvin
a. 590F
b. 84.40F
c. 98.60F
d. 2120F
41. How many seconds are there in one solar year (1 year = 365.24 days)
42. Every four calendar years, there are 365 x 4 days + 1 for leap year. Assuming
there are 365.24 days in a solar year, and that we’ve been counting leap day for
2000. years, how far, in days are we ahead or behind from the “real date?”
43. If a star is 1225 light years away from Earth, how far is it in miles (speed of light
= 3.00 x 108m/s).
44. The average distance from the Earth to the Sun is 1.50 x 108km. If the sun were to
explode, how long would it take before we could see it happen? (Hint: use the
speed of light).
45. Mrs. Gershman drives her convertible to school at a speed of 60.m/s. How fast
does she drive in mils per hour? If the speed limit is 1.48km/min, by how many
miles per hour is she speeding?
46. Mr. Amendola is on Pluto where the units of measurement are different. If his
Lamborghini gets 12.7 miles per gallon on Earth and gets 27.3 peeps per liter,
how many peeps are equal to 1 mile? If 1 quart = 3.01 quacks, how many peeps
per quack does his supercar get?
47. Sarah O’Meara runs a 400meter in 45 seconds. What is her average speed in yards
per minute? If the previous state record-holder ran it in 55 seconds, how many
miles per hour faster did she run?
48. A new temperature scale, measured in &, is created that is based on the freezing
and boiling point of anti-freeze solution. If 0& is equal to the freezing point (530C) and 100& is its boiling point (1280C), derive an equation the converts
temperatures in Celsius to & and one that converts from Fahrenheit to &.
49. In another galaxy, there are different kinds of subatomic particles. A new particle
called a Wangium is discovered and it is smaller than an electron and spherical in
shape. If the radius of a Wangium .000493 picometers and it has a mass of 4.03 x
10-21lbs, what is the density of a Wangium g/cm3. How much mass would
Wangium have if it were as big as a baseball (radius = 1.40in).
38
Chapter 2
1. What role did Democritus play in the development of atomic theory?
2. What does the word atomos mean in Greek and how does the word atom come
from it.
3. What was Antoine Lavoisier’s main contribution to atomic theory?
4. What was Joseph Proust’s major contribution to atomic theory? What is “Proust’s
law” or the Law of definite proportions?
5. Is the converse of the Law of definite proportions true? Give an example to
support your answer.
6. Which postulate of Dalton’s atom theory is false? Give two examples that make it
false.
7. What was the major flaw in Dalton’s relative mass theory? How did Gay-Lussac
improve upon this theory?
8. If equal volumes of H2 and Cl2 gas react, what volume of HCl is produced?
9. What determines the identity of an element? How does the number of neutrons of
an atom of an element affect its identity?
10. What is an isotope? Which subatomic particle must be constant in number for
each isotope of a given element?
11. The mass ratio of a proton to an electron is about 1840. What is their charge ratio?
What is the ratio of their charge/mass ratios?
12. Who discovered the electron and what was his experiment called? What property
of the electron did he discover?
13. What did Milikan determine with his oil drop experiment? How did he use
Thomson’s findings to determine the mass of an electron?
14. Describe the “plum pudding” model of the atom.
39
15. Describe Rutherford’s gold foil experiment. What happened to the alpha particles
when they were shot at the foil?
16. What did Rutherford conclude from this experiment? How did His new model of
the atom compare to the previously accepted plum pudding model?
17. Who discovered the Neutron and when? Why did it take longer to discover the
neutron than it did to discover the electron and proton?
18. Which is the most massive of the subatomic particles?
19. What are the mass number and atomic number of an atom? What letter represents
each?
20. Write the symbol for the Carbon atom with 6 protons and 6 neutrons using the
mass number – atomic number notation
21. Why is the atomic number often omitted from this notation?
22. If you were to choose a point at random within an atom, what are you most likely
going to find there?
23. What is true of the relative number of protons and electrons for cations? Anions?
24. How are elements arranged on the periodic table?
25. What the horizontal rows and the verticals columns called?
26. What is the name of
a. Group 1
b. Group 2
c. Group 7A
d. Group 8A
27. Where does the Lanthanide series start? Actinide series?
28. Why are the Lanthanides and Actinides placed at the bottom of periodic table?
29. What are the properties of metals, nonmetals and metalloids? Where are the
relative positions on the periodic table?
30. What are allotropes? Give an example.
31. What do parentheses around an atomic mass indicate about the element?
32. Give the symbols to the following metals
a. Titanium
40
b.
c.
d.
e.
f.
Scandium
Cobalt
Sodium
Potassium
Zinc
33. Give the symbols to the following metals
a. Lead
b. Iron
c. Tin
d. Silver
e. Gold
f. Mercury
34. Give the symbols of the following nonmetals
a. Chlorine
b. Fluorine
c. Nitrogen
d. Carbon
e. Silicon
f. Gallium
g. Xenon
35. Give the names that correspond to the following symbols
a. K
b. Fr
c. P
d. At
e. Be
f. Se
36. Five the names that correspond to the following symbols
a. La
b. Hg
c. Pt
d. Y
e. Pu
f. Cr
37. Give the number of protons and neutrons in a nucleus of the following.
a. 238U
b. 56Co
c. 38K
d. 31P
e. 1H
f. 52Cr
41
38. Write the following in the AZX format or mass-atomic number notation
a. Oxygen with 8 neutrons
b. Sulfur with 17 neutrons
c. Barium with 123 neutrons
d. Plutonium with 150 neutrons
e. Lead with 125 neutrons
f. Chlorine with 18 neutrons
39. Give the charge of the most common ion of the following elements.
a. Na
b. F
c. H
d. Ba
e. Br
f. S
g. N
40. How many protons and electrons are in the following
a. Ca2+
b. Ic. Rb+
d. Ni2+
e. O2f. Se241. How many protons and electrons are in the following
a. Mn2+
b. Mn7+
c. Cr3+
d. Cr6+
e. Sn2+
f. Sn4+
42. Name the following ions
a. Clb. Fc. O2d. N3e. S2f. I-
42
43. Name the ions in question 41
44. Name the following ions
a. Mo3
b. Zn2+
c. Fe2+
d. Cu+
e. Hg2+
f. Hg22+
45. Name the following polyatomic ions.
a. SO42b. PO43c. NO3d. ClO3e. BrO346. Name the following polyatomic ions
a. SO32b. SO22c. SO52d. HSO447. Name the following polyatomic ions
a. PO33b. PO23c. PO53d. HPO42e. H2PO448. Name the following acids
a. HNO2
b. H3PO3
c. H3PO2
d. H2SO3
e. H2SO2
f. HClO4
g. HClO2
49. Write the formula of the following polyatomic ions
a. Permanganate
b. Chromate
c. Dichromate
d. Acetate
e. Bicarbonate
f. Hydroxide
43
50. Write the formula of the following polyatomic ions
a. Nitrite
b. Cyanide
c. Thiosulfate
d. Oxalate
e. Ammonium
f. Mercury(I)
g. Thiocyanate
h. Peroxide
i. Superoxide
51. What is the empirical formula of a compound?
52. Name the following ionic compounds
a. NaCl
b. KF
c. LiOH
d. CaCl2
e. Mg(OH)2 (milk of magnesia)
f. MgS
g. SrO
53. Name the following ionic compounds
a. NaNO3
b. K2SO4
c. SrCl2
d. Ca(NO2)2
e. MnCO3
f. Na2O
54. Name the following ionic compounds
a. Hg2Cl2
b. CrPO4
c. Na2O2
d. BaO2
e. Fe(SCN)2
f. Mg(ClO)2
55. Write the formula of the following compounds
a. Carbon Dioxide
b. Nitrogen Dioxide
c. Hydrogen Chloride
d. Sulfur Trioxide
44
e. Phosphorus Pentachloride
f. Sulfur Hexafluoride
56. Write the formula of the following ionic compounds
a. Iron(III) Chloride
b. Zinc Oxide
c. Cerium(IV) Nitride
d. Potassium Selenate
e. Silver Peroxide
f. Copper(II) Phosphide
57. Write the formula of the following compounds
a. Nitrous Oxide (laughing gas)
b. Nitric Oxide
c. Stannous Chloride
d. Cupric Oxide
e. Ferrous Chloride
58. Write the formula of the following hydrates
a. Copper Sulfate Pentahydrate
b. Magnesium Sulfate Heptasulfate (Epsom)
c. Strontium Nitrate tetrahydrate
d. Sodium carbonate decahydrate (Washing soda)
45
Chapter 3
1. What is Atomic mass?
2. Do the elements in the periodic table increase in atomic mass. Provide an
example of an element with a higher atomic number than another element, but
a lower atomic mass. How could the lower atomic number element be lighter
than the other?
3. In what units is the atomic mass given in on the periodic table?
4. What does the atomic mass on the periodic table represent considering that
there is more than one isotope of each element in nature?
5. What is the atomic mass unit scale based on?
6. What device is used to determine the relative mass of two atoms?
7. What is a mole? Who introduced the mole?
8. What is the convenience in using Avogadro’s number?
9. How many grams are in one mole of Sulfur atoms?
10. The atomic mass of Nitrogen is 14.01amu. Is there a Nitrogen atom that exists
with a mass of 14.01amu? Explain.
11. What is the molecular mass of water? Molar mass?
12. How many grams are there in 1 mole of Cl2
13. How many grams are there in 2.01mol of C
14. What is the molar mass of Sulfuric acid?
15. What is the molar mass of ethane (C2H6)
16. Assuming a piece of paper is .00200in thick, how high would a stack of a
mole of papers be in light years?
17. How many moles are there in 75.0g of Fluorine gas?
46
18. If the atomic mass unit were based on a system in which Carbon-12 is set to
be 20amu, what would Avogadro’s number be?
19. How many Phosphorus atoms are there in 8.92g of P4O10?
20. How many grams are in one Sodium atom?
21. What is the molar mass of Urea? ([(NH2)2CO])
22. What is the mass of hydrogen atoms in grams in one mole of Acetic acid?
(CH3COOH)
23. The relative abundances of the two Chlorine isotopes are 35Cl (75.53 percent)
and 37Cl (24.47 percent). If the atomic masses of the two isotopes are
34.968amu and 36.956amu respectively, what is the atomic mass of Chlorine?
24. What is the Percent composition (percent by mass) of each atom in Ammonia?
25. What is the percent composition of each atom in Guanidine? (HNC(NH2)2)
26. What is the percent composition of each atom in Calcium Silicate? (Silicates
are similar to Carbonates). How many grams of Silicon does 3.40mol of
Calcium Silicate contain?
27. A compound that contains Carbon, Hydrogen, and Oxygen has a percent
composition of C: 53.30%, H: 11.20%, and O: 35.50%. What is the empirical
formula of this compound?
28. Iron(III) Oxide, or rust, is represented by the formula Fe2O3. How many
grams of Iron are there in 45.02g of rust?
29. What is the percent composition of Acrylonitrile? (C3H3N)
30. What is the molar mass of Chloral Hydrate? (C2H3Cl3O2) How many Chlorine
atoms are in 435g of Chloral hydrate? What is the percent composition of
Oxygen in Chloral hydrate?
31. A Chromium Oxide has a percent composition of Chromium of 52.00%. What
is the charge of the Chromium ion in the compound?
32. A compound that contains Carbon, Nitrogen, and Hydrogen has a percent
composition of C: 38.65%, N: 45.09%, and H: 16.25%. What is the empirical
formula of the compound? If 1.23mol of this compound is 38.22g, what is the
molecular formula of the compound?
47
33. Balance the following equations
a. Mg(s) + H2O(l) → Mg(OH)2(s) + H2(g)
b. N2(g) + Ca(s) → Ca3N2(s)
c. F2 + Au → AuF3
34. Balance the following equations
a. HCl(aq) + MnO2(s) → MnCl2(aq) + Cl2(g) + H2O(l)
b. Ag(s) + NH3(l) → [Ag(NH3)2]+(aq)
c. HCl(aq) + Na2CO3(aq) → H2O(l) + CO2(g) + NaCl(aq)
35. Balance the following equations
a. Cu(s) + HNO3(aq) → Cu(NO3)2(aq) + NO(g) + H2O(l)
b. Br2(g) + KClO3(s) + H2SO4(aq) + H2O(l) → HBrO3 + K2SO4 + HCl
c. Ca3P2(s) + H2O(l) → CaO(s) + PH3(aq)
36. Balance the equation of the reactions described
a. Iron filings are sprinkled over powdered Sulfur
b. Calcium Hydroxide is reacted with Hydroiodic acid
c. Hydrochloric acid is added Calcium Sulfite (SO2 is formed)
37. Phosphorus naturally occurs Fluorapatite, CaF2 * 3Ca3(PO4)2. When
Fluorapatite is reacted with Sulfuric Acid, Phosphoric Acid, Hydrogen
Fluoride, and gypsum (CaSO4 * 2H2O) are produced. Write a balanced
chemical equation for this reaction.
38. Hydrogen Peroxide combusts according to the following equation.
.
H2O2 → 2H2O + O2
If the reaction goes to, what mass of O2 gas is produced if 43.2g of
Hydrogen Peroxide is completely combusted?
.
39a Sodium reacts with water according to the following equations
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
If 12.2g of Na reacts with excess water, how many moles of H2 gas are
produced?
39. When 53.1g of Magnesium metal is reacted with excess Nitrogen gas, the
reaction yield is 92.0%. How much Magnesium Nitride is produced?
40. 40.8g CH4 is reacted with excess Oxygen gas to produce Carbon Dioxide and
water. How many grams of Carbon dioxide are produced?
41. Coke is an impure form of Carbon. If a sample of Coke is 94% percent
Carbon by mass, how much coke is needed to completely react with 2520kg
of Copper(II) Oxide to produce Copper metal and Carbon Dioxide?
48
42. A sample of a mixture containing Sodium Chloride and Potassium Chloride
has a mass of 5.100g. When it’s reacted with excess Silver Nitrate, 10.41g of
Silver Chloride (a white solid) is formed. What is the percent composition of
Potassium Chloride and Sodium Chloride in the original compound?
43. For the following unbalanced equation, what is the maximum mass of NH3
that can be produced when 16.5 g of Mg3N2 and 28.5 g H2O react?
Mg3N2(s) + H2O(l)
Mg(OH)2(s) + NH3(aq)
44. For the following unbalanced equation, a 0.620 g sample of impure Al2(SO4)3
reacts with excess BaCl2. If the sample produces 0.700 g of BaSO4, what is
the mass percent of Al2(SO4)3 in the sample?
Al2(SO4)3(aq) + BaCl2(aq)
BaSO4(s) + AlCl3(aq)
45. If 11.5g of Calcium and 15.0g of Phosphorus react and 14.02g of Calcium
Phosphide is produced, what is the percent yield of the reaction?
46.
For the unbalanced equation, what is the maximum mass of N2O5 that can be
produced when 35.0 g of N2 and 65.5 g O2 react?
N2(g) + O2(g)
N2O5(g)
47. Assume the reaction in #46 goes to completion. Calculate the mass of the
excess reactant.
48. A sample of Copper Sulfate Pentahydrate is heated in a crucible until the
water is vaporizes and leaves the crucible. The change in mass of the crucible
from before to after the heating is 5.67g. How many moles of Copper Sulfate
Pentahydrate were present in the sample?
49. A hydrocarbon is combusted in excess oxygen gas to produce CO2 and H2O.
If 17.7g of CO2 and 9.04g of H2O are formed, what is the empirical formula of
the hydrocarbon? If 208g of Oxygen gas is needed to completely combust
1.00 mole of the compound, what is its molecular formula?
49
Chapter 4
1. What is the solute and solvent in a solution?
2. What is an aqueous solution?
3. What is a polar molecule? Is water polar or nonpolar?
4. What is the process of dissolving an ionic compound in water called?
5. Which aqueous solution would be a better electrical conductor, NaCl or
Sucrose?
6. What makes a compound a weak electrolyte or a strong electrolyte?
7. CH3COOH(aq) ↔ CH3COO-(aq) + H+(aq). Why is this reaction written
with a double arrow? What does this say about the electrical conductivity
of an aqueous solution of acetic acid?
8. How does solubility differ with electrolytes?
9. Classify the following as soluble or insoluble in water
a. Sodium Hydroxide
b. Calcium Carbonate
c. Magnesium Sulfate
d. Silver Chloride
e. Ammonium Sulfide
10. Classify the following compounds as soluble or insoluble in water
a. Barium Hydroxide
b. Iron(II) Nitrate
c. Rubidium Chromate
d. Mercury(II) Chloride
e. Strontium Chlorate
11. Write the complete and net ionic equations for the following reaction.
What is the color of the precipitate formed?
Pb(NO3)2 + 2NaI → 2NaNO3 + PbI2
50
12. Write the complete and net ionic equations for the following reaction.
What is the color of the precipitate formed?
AgNO3 + KCl → AgCl + KNO3
13. Write the complete and net ionic equations for the following reactions.
What is the color of the precipitate formed?
Al(NO3)3 + 3LiOH → 3LiNO3 + Al(OH)3
14. Write the complete and net ionic equations for the following reaction.
What is the color of the precipitate?
NiCl2 + 2KOH → Ni(OH)2 + 2KCl
15. What is the definition of an Arrhenius acid and Arrhenius base?
16. What is the definition of a Brønsted acid and base
17. Which acid and base definition is more encompassing?
18. Why is H+ often expressed as H3O+ in acid dissociation reactions? Which
one is more accurate in terms of what exists in aqueous solution?
19. Classify the following as an acid or a base
a. NaOH
b. HNO3
c. HCl
d. Ba(OH)2
20. What is a monoprotic acid. Give two examples
21. What is a diprotic acid. Give an example
22. What is another name for Ammonium Hydroxide?
23. What is always produced in a Brønsted acid-base reaction?
24. What is an oxidation-reduction reaction?
25. What is the oxidation state of Nitrogen in Nitric Acid? Ammonia?
26. What is the oxidation state of Silver in Silver Oxide?
27. What is the oxidation state of Chromium in the following compounds
a. Chromate
b. Potassium Dichromate
c. Chromium(II) Oxide
28. Identify the oxidizing agent and reducing agent in the reactions in question
35 of chapter 3, a and b.
Colors of various compounds are very useful to know for the AP test.
51
29. What is the color of the following solids?
a. NaCl
b. Manganese(IV) Oxide
c. Lead Sulfate
d. Cadmium Sulfide
30. What is the color of Potassium Dichromate?
31. What is the color of Barium Chromate?
32. What color is Potassium Permanganate?
33. What is the color of the following transition metal ions in aqueous solution
a. CrCl2
b. CrCl3
c. CoCl2
d. MnCl2
e. FeCl3
f. NiCl2
g. CuCl2
h. TiCl3
34. What is the color of the following gases
a. Cl2
b. Br2
c. I2
d. NO2
e. CO2
35. Will Lithium react with Hydrochloric Acid? (use activity series)
36. Will Platinum react with Hydrochloric Acid?
37. Which metal reacts more vigorously with a strong acid, Zinc or
Magnesium? Why?
38. Define molarity.
39. How many moles are in 233mL of a 1.43M NaCl solution?
40. How many Chloride Ions are in 1.00L of 3.20M CaCl2 solution?
41. What volume of a .400M KI solution contains 2.41moles of KI?
42. How much distilled water needs to be added to 205.0mL of a 4.00 molar
BaBr2 solution so that the concentration of Br- becomes 3.50M?
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43. A 230.mL sample of a 0.275M CaCl2 solution is heated. The concentration
of the solution afterwards is1.10 M. What volume of water evaporated
from the 0.275M CaCl2 solution?
44. A safe amount of Lead in tap water is below 15 parts per billion (ppb). If
the concentration of Lead(II) ions a sample of tap water is 9.65 x 10-8M, is
the water safe to drink? (Hint: Assume the density of water is 1.00g/mL.)
45. The following reaction occurs in developing black and white film. What
mass of AgBr can be dissolved by 1.27 L of 0.242 M Na2S2O3?
AgBr(s) + 2S2O32-(aq)
Ag(S2O3)23-(aq) + Br -(aq)
46.
For the unbalanced reaction, what is the maximum mass of Mg3(PO4)2 that
can be produced when 16.5 mL of 0.310 M MgCl2 and 10.0 mL of 0.240 M
Na3PO4 react?
MgCl2(aq) + Na3PO4(aq)
Mg3(PO4)2(aq) + NaCl(aq)
47. What mass of KBr is required to prepare 8.70L of 0.100M solution?
48. What is the molarity of each ion in 0.300M Na3PO4?
49. What is the molarity of a solution of sodium carbonate that contains 6.30 g
of Na2CO3 in 575 mL of solution?
50. What is the main piece of equipment used in a titration? What is the
advantage in using it?
51. Why are Potassium Permanganate and Potassium Dichromate useful as
oxidizing agents in redox titrations that occur in an acidic medium?
52. If 34.2mL of .120M NaOH is required to titrate 23.5mL of an HCl
solution, what is the concentration of the HCl solution?
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53. The pH scale is used to measure the acidity of a solution. pH = -log(H+
concentration in mol/L). If the pH of an HCl solution is 2.43, what is the
concentration of Cl- in the solution? (HCl is a strong acid and strong
electrolyte)
54. 34.2ml of a .405M Ba(OH)2 solution is mixed with 50.0 ml of a 2.55M
KOH solution. What is the concentration of OH- in the resulting solution?
55. 3.530g of an unknown metal is added to excess Hydrochloric Acid. The
mass of the beaker and hydrochloric Acid is 88.322g and after the metal is
added and the reaction is completed, the mass of the beaker and its
contents is 91.674g. After the reaction, the contents of the beaker are clear
and a gas is released during the reaction. The possibilities of the identity of
the metal are Ag, Ca, Cr, and Al. What is the metal? Use qualitative and
quantitative reasoning.
56. 54.0mL of a 2.33M CaCl2 is mixed with 12.3mL of a 3.60M AlCl3
solution. 35.5mL of a solution of copper(I) chloride is added and the final
concentration of Cl- is 4.89M. What was the concentration of the copper(I)
chloride solution?
57. An alloy with a mass of 46.8g that contains Silver, Zinc, and Copper is
reacted with excess concentrated Hydrochloric Acid to produce .522g of a
gas and then the remaining solid is reacted with 210.0mL of a 2.00M
Sodium Thiosulfate solution. The remaining solution has a Thiosulfate ion
concentration of 1.24M. Only Silver reacts with Thiosulfate and according
to the following equation
Ag(s) + 2S2O32-(aq) → Ag(S2O3)23What is the percent composition of the alloy? (All reactions go to completion)
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