Chemistry
Final Exam Review
Chapter 2 Study Questions
1. Define the following:
a) element
b) compound
c) pure substance
2. Classify each of the following as a pure substance or a mixture. For each pure substance,
indicate whether it is an element or a compound. Which of the mixtures are solutions?
a) air
b) titanium
c) oak
d) baking soda
e) oxygen
f) 7-Up
g) wine
h) carbon monoxide
3. Label each of the following drawings as element, compound, or mixture: (Assume each type
of circle represents a different type of atom.)
a)
b)
c)
d)
____________
____________
____________
____________
Which of the boxes above contain molecules?
4. List four physical states of matter. For each physical state, indicate whether the particles are
in motion and whether they are close or far apart.
5. List one chemical and one physical property of the element chlorine. (You may use your
textbook.)
6. Describe three observations that frequently accompany chemical reactions and explain why
they might indicate that a chemical reaction is occuring.
7. Classify each of the following processes as physical or chemical changes.
a) combustion of natural gas
b) evaporation of alcohol
c) condensation of water vapor
d) photosynthesis
e) splitting of carbon dioxide into carbon and oxygen
f) formation of sodium chloride (NaCl) from it elements
g) distillation of a sodium chloride solution to collect pure water
8. Record one a) qualitative and one b) quantitative observation about this page. State a theory
about this page.
9. Give the name of the following types of containers:
a)
b)
c)
d)
Summary of Chapter 2: Matter and Energy
matter
physical states: solid, liquid, gas
physical & chemical properties
physical & chemical changes
elements
compounds
atoms
molecules
pure substances
homogenous and heterogenous mixtures
solutions
separation of mixtures: distillation, filtration, chromatography
Chemistry
Chapter 5 Study Questions
1. Express the following numbers or answers in scientific notation:
a) 650 (2 sig fig)
b) 0.0005 (1 sig fig)
3
2
d) (5.0 x 10 ) x (2.0 x 10 )
e) (3.0 x 102) ÷ (6.0 x 103)
c) 207,000 (3 sig fig)
2. For each of the following, indicate the metric unit and a device used to measure it.
a) volume
b) mass
c) length
3. Indicate the number of significant figures in the following numbers:
a) 2,348
b) 7.0001
c) 0.0023
d) 24,500
e) 0.1060
4. Perform the following operations and express the answers in significant figures:
a) 1.24 x 8.2 =
b) 6.78 - 3.3 =
c) 9.999 + 0.22 =
d) (5.67 x 103) x (2.1 x 102)
5. Bozo determined the density of a sample of aluminum. For his sample, he found the volume
was 0.350 cm3 and the mass was 0.822 g.
a) Calculate the density of aluminum from Bozo’s data.
b) The actual density of aluminum is 2.70 g/cm3. Calculate Bozo’s percent accuracy error.
NOTE: Use conversion factors to answer the problems below. Show all work. Answers must
be in significant figures and include units. Use the table on the inside back cover of the
textbook as needed. (Or you may use the Table handed out in class.)
6. Calculate the mass in milligrams of a person with a mass of 50.0 kg.
7. Find the mass in pounds (lbs) of a 275-gram sample of sugar.
8. Find the number of cm in 0.286 miles.
9. Find the volume in microliters of 11.8 kg of iron. The density of iron is 7.87 g/cm3.
10. Tungsten is a very dense metal, with a density of 19.3 g/cm3. Convert the density of
tungsten to pounds/quart.
11. The volume of a sample is recorded from three different containers as indicated below. The
“true” value for the volume is exactly 61.2 mL.
Measurement
Container A
Container B
Container C
1
63.40 mL
61 mL
59 mL
2
63.48 mL
60 mL
59 mL
3
63.42 mL
62 mL
59 mL
a) Which of the three containers is the most precise?
b) Which container is the most accurate?
c) Which container(s) show a systematic error?
12. (OPTIONAL) Assuming each ant is 5.0 mm long, how many ants would it take to make a
line, single file, from one end to the other of a 100-yard football field? (2 sig fig)
Summary of Chapter 5: Measurements and Calculations
qualitative vs. quantitative observations
scientific notation
metric units: gram, liter, meter
metric prefixes: nano, micro, milli, centi, kilo
measuring devices: balance, graduated cylinder
significant figures
recording, counting and in arithmetic
exact quantities
accuracy & precision
| true value  measured value |
x 100%
percent accuracy error: % accuracy error 
true value
problem solving using conversion factors (dimensional analysis)
density
Chemistry
Chapters 3 & 4 Study Questions
1. What are two conclusions supported by Rutherford’s experiment?
2. Fill in the following table:
Nuclear
Symbol
Atomic
Number
Mass
Number
Number of
Protons
Number of
Electrons
Number of
Neutrons
Charge
40
18 Ar
______
______
________
_________
_________
______
_______
______
39
19
18
_________
______
_______
16
______
_________
20
-2
_________
3. Write the nuclear symbols for the isotopes of neon which contain 10 neutrons and 12
neutrons.
4. For each of the following elements, indicate whether it is a main group element (MG),
transition metal (TM), or inner transition metal (ITM). If the element is a main group
element, indicate the group number and whether it is a metal, a nonmetal or a metalloid.
Also indicate the Period of each element.
a) Sr (atomic # 38)
b) Br (atomic # 35)
c) Mo (atomic # 42)
d) P (atomic # 15)
e) B (atomic # 5)
f) U (atomic # 92)
g) Sn (atomic # 50)
h) Hg (atomic # 80)
5. Provide the common names of Groups 1, 2, 17 and 18.
6. Give an example of
a) an element made of molecules
c) a compound made of ions
b) a compound made of molecules
7. For each of the following atoms, indicate whether it forms a positive or a negative ion, and
include the ion charge.
a) Na
b) Ba
c) Cl
d) S
e) Ag
8. Which of the following are ionic compounds? Which are covalent compounds? Name each
compound.
a) N2O
b) K2O
c) PCl3
d) AlPO4
e) HCl
f) NH4F
g) Pb(NO2)2
h) H2SO3
9. Name the following ionic compounds:
a) CaCO3
b) ZnS
c) CuOH
d) Mg(ClO4)2
10. Give the formulas for the following ionic compounds:
a) potassium phosphate
b) ammonium sulfate
c) cobalt(II) hyroxide
d) iron(III)nitride
11. Provide the formulas for the following covalent compounds:
a) phosphorus triiodide
b) dinitrogen pentoxide
c) chloric acid
Summary of Chapter 3: Elements, Atoms, and Ions
element symbols
atomic theory
law of constant composition
elements
atoms
compounds
chemical formulas
Rutherford’s experiment
structure of atom
nucleus
protons, electrons, neutrons
atomic number, mass number
isotopes
nuclear symbol
periods & groups
regions of the Periodic Table: main groups, transition metals, inner transition metals
noble gases, halogens, alkali metals, alkaline earth metals
metals, nonmetals, metalloids
ionic & covalent compounds
molecules
diatomic molecules
allotropes
ions
formulas for ionic compounds
Summary of Chapter 4: Nomenclature
naming binary ionic compounds
Type I (no roman numeral)
Type II (roman numeral)
naming binary covalent compounds
polyatomic ions
naming ionic compounds involving polyatomic ions
naming acids
writing formulas from names
Chemistry
Chapters 6 & 7 Study Questions – note some questions numbers have intentially been
deleted since these types of questions will not appear on the final.
1. Glycerol (C3H8O3) is sold in drug stores as glycerine and is commonly found in soaps and
shampoos.
a) What is the molar mass of glycerol?
b) What is the mass in grams of 1.00 mole of glycerol?
c) How many molecules are in one mole of glycerol?
d) How many grams are in 0.217 moles of glycerol?
e) How many moles are in 783 grams of glycerol?
2. Ammonia (NH3) is the active ingredient in many kitchen cleansers. How many atoms are in
a) one molecule of ammonia?
b) one mole of ammonia?
c) 3.40 grams of ammonia?
3. Sodium nitrite is a controversial food preservative added to processed meat and thought to
form cancer-causing compounds when heated. What are the mass percentages of each
element in sodium nitrite?
6. Describe three observations that frequently accompany chemical reactions and explain why
they might indicate that a chemical reaction is occuring.
7. Balance the following equations:
a) the reaction between iron and oxygen to form iron(III) oxide,
Fe(s) + O2(g)  Fe2O3(s)
b) the combustion of the rocket fuel diborane,
B2H6(l) + O2(g)  B2O3(s) + H2O(l)
c) the combustion of the poisonous gas, PH3,
PH3(g) + O2(g)  H2O(l) + P4O10(s)
8. In the balanced equation for 7c above
a) What are the reactants? b) What are the products? c) What is the coefficient for water?
9. Write a balanced equation for each of the following reactions:
a) the reaction of lithium with nitrogen gas to form lithium nitride.
b) the reaction of propane (C3H8) gas with oxygen to form carbon dioxide and water.
Summary of Chapter 6: Chemical Composition
atomic mass
mole
Avogadro’s number
molar mass (molecular mass, formula weight)
calculations: # particles  moles  mass
percent composition (mass percent)
empirical formula
molecular formula
percent composition  empirical formula
finding molecular formula from empirical formula and molar mass
empirical formula from experimental data
Summary of Chapter 7: Introduction to Chemical Reactions
chemical reactions
chemical equations
reactants, products
coefficients
writing and balancing chemical equations
Chemistry
Chapters 8 & 9 Study Questions
1. Sodium carbonate and iron(III) chloride react to form a precipitate.
a) Write a balanced molecular equation for this reaction.
.
2. Using a solubility table, decide whether a precipitate will form when the following solutions
are mixed. If a precipitate forms, write a net ionic equation for the reaction.
a) iron(III) nitrate and potassium hydroxide
b) ammonium chloride and lithium carbonate
c) sodium sulfide and nickel(II) sulfate
3. For each of the following equations
i. indicate whether it is a combustion (C), synthesis (S), decomposition (D), single
replacement (SR), or double displacement (DD)
ii. indicate which reactions are oxidation-reductions (OR), precipitations (P) or acid-base
(AB) reactions.
iii. predict the products and record their formulas. If no reaction occurs, write “NR.”
(Use Tables as needed to help you decide if a reaction occurs.)
iv. balance the equation.
a)
c)
e)
g)
Li(s) + Cl2(g) 
C3H6(g) + O2(g) 
Fe(s) + MgSO4(aq) 
Al(s) + HCl(aq) 
b)
d)
f)
h)
Sr(NO3)2(aq) + K2SO4(aq) 
CaCl2(aq) + NaNO3(aq) 
KI(l) 
HNO3(aq) + KOH(aq) 
4. Write a chemical equation for the ionization of iron(III) nitrate when it dissolves in water.
5. Chromium reacts with hydrochloric acid in a single replacement reaction. The balanced
equation is:
2 Cr(s) + 6 HCl(aq)  2 CrCl3(aq) + 3 H2(g)
a) How many moles of HCl are needed to produce 1.60 moles of CrCl3?
b) How many grams of chromium are required to react with 0.450 moles of HCl?
c) How many atoms of chromium are required to produce 12 moles of H2?
d) How many grams of Cr are needed to produce 3.20 g H2?
e) In an experiment, 10.2 grams of CrCl3 are produced starting from 8.30 grams of HCl.
What was the theoretical yield and the percent yield in this experiment?
f) When 6.0 moles of Cr are combined with 12.0 moles of HCl, which reactant is limiting?
How many moles of excess reactant are left over?
g) How many grams of CrCl3 are produced starting from 13.0 g of Cr and 43.8 g of HCl?
Summary of Chapter 8: Reactions in Aqueous Solutions
predicting whether a reaction will occur
precipitation reactions
strong electrolytes
using a solubility table
predicting whether a precipitate occurs
writing equations for precipitation reactions
molecular equations
complete ionic equations
net ionic equations
acids, bases
acid-base reactions
common strong acids
common strong bases
double displacement reactions
precipitation reactions
acid-base reactions
oxidation-reduction reactions
synthesis (combination)
decomposition
combustion reactions
single replacement reactions
Summary of Chapter 9: Chemical Quantities
interpreting balanced chemical equations
stoichiometric calculations:
mole relationships between reactants and products
mass relationships between reactants and products
limiting reactant
theoretical yield
experimental yield
calculating percent yield
Chemistry
Chapter 11 Study Questions
1. How are wavelength and frequency related? How are energy and frequency related?
2. What was the revolutionary new idea in Bohr's model of the hydrogen atom? What was the
most significant difference between the quantum mechanical atom and the Bohr hydrogen
atom? Briefly explain the relationship between electronic transitions and atomic spectra.
3. Explain, in terms of their electron configurations, why the most reactive metals are in Group
1, the most reactive nonmetals are in Group 17, and the noble gases are chemically inert.
4. What is the electron capacity of
a) any Principal energy level?
b) each sublevel?
c) each orbital?
5. Which of the following sublevels do not exist? List the ones that do exist in order of
increasing energy.
a) 1s
b) 2s
c) 2d
d) 3d
e) 4p
f) 4f
6. Which sublevel is in the process of being filled in the following regions of the periodic table?
a) Groups 1 and 2
b) Transition metals
c) Group 15
d) Inner transitional metals
7. What is the outer electron configuration of the following groups?
a) alkali metals
b) halogens
c) noble gases
8. Classify each of the following electron configurations as ground state, excited state or
impossible:
a) 1s22s22p1
b) 1s21p62s2
c) 1s22s22p43s1
d) 1s22s22p63d1
9. Give the complete ground state electron configuration of
a) sulfur
b) the element with atomic number 29
10. Give the abbreviated ground state electron configuration of
a) strontium
b) lead (Z = 82)
11. Give the symbol of the element which (in the ground state)
a) has the outer electron configuration 6s2
b) is in Group 18 but has no p electrons
c) has three unpaired 4p electrons
d) has four valence electrons in the Second Principal Energy level.
e) is in Period 3 and has the same outer electron configuration as F.
f) has only five 3d electrons.
12. Sketch the shape of s and p orbitals. How do orbitals change as n increases?
13. Draw a complete orbital diagram for
a) oxygen
b) titanium (Z = 22)
14. Which is a better predictor of chemical properties: Period number or Group number?
15. a) Which element has a greater ionization energy? Cl or Ar? Na or K?
b) Which element has a larger atomic radius? Mg or Ca? S or Cl?
Summary of Chapter 11: Modern Atomic Theory
wavelength (), frequency ()
atomic spectra
Bohr model of the hydrogen atom
ground state, excited states
quantum mechanics
electron clouds
orbitals
principle energy levels (n)
sublevels (s, p, d, f): electron capacity and relative energies
ground state electron configuration of atoms
abbreviated electron configurations
outer electron configuration
valence electrons
orbital diagrams
Hund’s rule
Pauli exclusion principle
ionization energy
atomic radius
electron configuration & the Periodic Table
Chemistry
Chapter 12 Study Questions
1. What is a chemical bond? Why do atoms form chemical bonds? How are covalent bonds and
ionic bonds different? How are they the same?
2. How is the valence of an atom related to the number of bonds it usually forms?
3. What types of substances contain covalent bonds?
4. List the atoms in each of the following sets in order of increasing electronegativity:
a) N, As, P
b) O, Li, C
c) Mg, K, B
5. Choose the bond from each pair which is most polar.
a) Cl  Cl, H  Cl
b) O  C, F  C
6. Choose the atom or ion in each set with the smallest atomic radius.
a) Li, Li+, H
b) Na+, Cl, K+
c) F, O2, F
7. For each of the following ions, give its electron configuration, another ion with the same
configuration and a noble gas with the same configuration:
a) O2
b) Sc3+
8. Write a balanced chemical equation (always include physical states) for the reaction between
calcium and iodine.
9. Draw Lewis structures for the following atoms:
a) Be
b) C
c) F
10. Draw Lewis structures for the following compounds:
a) H2S
b) Br2
c) NH2F
d) CH2I2
f) diphosphorus dichloride
g) dinitrogen tetroxide
h) C3H4Cl2
i) C3H4
e) CO32-
13. Add hydrogen atoms and electrons in order to complete Lewis structures of the following
following compound: C3H6O (acetone; nail polish remover)
CCC
|
O
Chemistry
Summary of Chapter 12: Chemical Bonding
chemical bonds
ionic bonds
covalent bonds
polar and nonpolar covalent bonds
electronegativity
bond polarity
dipole moment
electron configurations of ions
ion size
valence
Lewis structures of atoms
Lewis structures of molecules
Octet rule
lone pairs
resonance
VSEPR Model
Chemistry
Chapter 13 Study Questions
NOTE: Vapor Pressure of Water Chart is at the bottom of this page section.
1. A flask contains air at 722 mm Hg and 22°C. What would the temperature of the gas be if the
pressure is increased to 1.07 atm?
2. A sample of air collected at STP contains 0.039 moles of N2, 0.010 moles of O2, and 0.001
moles of Ar. (Assume no other gases are present.)
a) Find the partial pressure of O2.
b) What is the volume of the container?
3. A sample of hydrogen gas (H2) is collected over water at 19C.
a) What are the partial pressures of H2 and water vapor if the total pressure is 756 mm Hg?
b) What is the partial pressure of hydrogen gas in atmospheres?
4. If 600 cm3 of H2 at 25C and 750 mm Hg is compressed to a volume of 480 cm3 at 41C,
what does the pressure become?
5. Find the density of helium gas at STP.
6. a) Write a balanced chemical equation for the reaction of butane gas with oxygen gas to form
carbon dioxide and water vapor.
b) How many liters of oxygen are required to produce 2.0 liters of CO2?
c) How many liters of CO2 are produced from 11.6 g of butane at STP?
d) How many molecules of water vapor are produced from 5.6 liters of butane gas at STP?
7. A flask contains 0.25 moles of argon and 0.75 moles of helium. If the partial pressure of the
helium is 0.60 atm, what is the total pressure in the flask?
8. Calculate the density of carbon dioxide at 546 K and 4.00 atmospheres pressure.
9. How many grams of methane are contained in a 28.0 liter flask at 68C and 2.00 atmospheres
pressure?
10. What volume of O2 at 710 mm Hg pressure and 36C is required to react with 6.52 g of
CuS?
CuS(s) + 2 O2(g)  CuSO4(s)
11. What is the molar mass of a gas if 7.00 grams occupy 6.20 liters at 29C and 760 mm Hg
pressure?
12. At a particular temperature and pressure, 15.0 g of CO2 occupy 7.16 liters. What is the
volume of 12.0 g of CH4 at the same temperature and pressure?
Summary of Chapter 13: Gases
Kinetic-molecular theory
pressure
barometer, manometer
1 atm = 760 mmHg = 760 torr
temperature
absolute zero temperature
T(K) = T(°C) + 273
relationship between pressure, volume, temperature
Boyle’s Law
Charles’ Law
Avogadro's Law
Ideal Gas Law
R = 0.08206 L atm/mol K
partial pressure
molar volume
STP
molar volume @ STP = 22.4 L
gas stoichiometry
molar mass and density of a gas
formulas:
Ptotal = Px + Py + . . .
P1V1 P2 V2

T1
T2
PV = nRT; R = 0.08206 L atm/mol K
mm
d=
mV
n 
P1   1  PT
 nT 
Temp (°C)
PH2O(mm Hg)
Chemistry
15
13
16
14
Vapor Pressure of Water
17
18
19
20
21
15
15
16
18
19
22
20
23
21
24
22
25
24
Chapter 14 Study Questions
1. Discuss the differences between the states of matter (gas, liquid, solid) with respect to: a)
distance between the particles, b) mobility of the particles, c) shape of the substance, and d)
attractive forces between the particles.
2. List the types of intermolecular attractive forces in order of increasing strength.
3. Is the heat of vaporization endothermic or exothermic? Why?
4. Define boiling point in terms of vapor pressure.
5. For each of the following sets, indicate the substance with the highest boiling point:
a) C2H6, C8H18, or C4H10
b) HF, HCl, or HBr
c) CH3OCH3 or CH3CH2OH
d) H2O, C3H8, or MgO
e) CH3CH2CH3 or CH3OCH3
f) Cs or CH3OH
6. For each of the following types of solids, describe its structure and the nature of the forces
holding it together:
a) ionic
b) covalent (molecular)
c) metallic
d) network covalent
7. List the substance types in (7) in order of increasing melting point.
8. Which of the types of substances in (7) conduct electricity as solids? as liquids?
9. Classify the following substances according to the types in (7):
a) NH3
b) SiO2
c) sodium oxide
d) magnesium
e) O2
f) Rb
g) KNO3
h) carbon disulfide
10. List the following substances in order of increasing vapor pressure: CO2, SiO2, or TiO2.
11. The heat of fusion of ethanol is 26.1 cal/g. Calculate the number of moles of ethanol needed
to produce 1.31 kilojoules when it freezes.
Summary of Chapter 14: Liquids and Solids
Differences between gas, liquid, solid
heating/cooling curve
heat of fusion and heat of vaporization
calculating the heat needed to melt or boil a given amount of a substance
sublimation
intramolecular forces and intermolecular forces
Intermolecular forces:
London dispersion forces, dipole forces, hydrogen bonds
Relationship between interparticle forces and melting pt, boiling pt, vapor pressure
vapor pressure
boiling point
Properties of the following types of solids (nature of particles, electrical conductivity, melting
points, solubility, examples): molecular, network covalent, ionic, metallic
Chapter 15 Study Questions
Mass percent
1. What is the mass percentage of KMnO4 in a solution containing 1.00 mole of KMnO4 and
158 g of water?
2. How many moles of KMnO4 are needed to prepare 335 g of a 22.0% solution?
Molarity
3. How many moles of NaCl are in 275 mL of 0.500 M NaCl?
4. What mass of NaCl is needed to prepare 250. mL of a 2.00 M NaCl solution?
5. What volume of a 2.00 M NaCl solution is needed to make 125 mL of a 0.350 M NaCl
solution?
6. What is the molarity of a solution made by dissolving 90.0 grams of glucose (C6H12O6; molar
mass = 180. g/mole) in enough water to yield 200. mL of solution?
7. (Optional) What is the molarity of a 24.0% sucrose solution? The density of this solution is
1.10 g/cm3. (The molar mass of sucrose is 342 g/mole.)
Summary of Chapter 15: Solutions
solution
solute
solvent
molecular and ionic solutes
like dissolves like
nonpolar vs. polar solutes and solvents
Saturated, supersaturated and unsaturated solutions
solubility and temperature
solution composition
mass percent
molarity
dilution: V1 x M1 = V2 x M2