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East Meck Chemistry
STUDENT NAME:___________________
TEACHER/Period:________________
Table of Contents
Pages
Unit 1- Matter and Change ................................................................................................... 1 - 9
Unit 2- Atomic Theory and Structure ................................................................................... 10 - 23
Unit 3- Electrons and Periodicity ......................................................................................... 24 - 38
Unit 4- Bonding .................................................................................................................... 39 - 50
Unit 5- Nomenclature ........................................................................................................... 51 - 65
Unit 6- Chemical Reactions .................................................................................................. 66 - 76
Unit 7- Stoichiometry ........................................................................................................... 77 - 84
Unit 8- Gas Laws .................................................................................................................. 85 - 97
Unit 9- Solids, Liquids and Phase Changes .......................................................................... 98 - 106
Unit 10- Solutions and Solubility ......................................................................................... 107 - 115
Unit 11- Acids and Bases ..................................................................................................... 116 - 130
Unit 12- Kinetics and Thermochemistry .............................................................................. 131 - 145
Unit 13- Oxidation and Reduction ........................................................................................ 146 - 152
Reference Materials .............................................................................................................. 153 - 160
Unit 1- Matter and Change
CHEMISTRY: THE STUDY OF MATTER
Chemistry is the science that investigates and explains the structure and properties of ________________. This
includes its composition, properties and the changes it undergoes.
SCIENTIFIC METHOD
A scientific method is a systematic approach to answer a ______________________ or study a situation. It starts
with _____________________ - noting and recording facts. A _________________________ is a possible
explanation for what has been observed. It is an educated _________________ as to the cause of the problem or
answer to the question. An experiment is a set of controlled observations that _____________ a hypothesis. The
variable that is changed in an experiment is called the ________________________ variable. The variable that you
watch to see how it _________________ as a result of your changes to the independent variable is called the
dependent variable. The cycle (hypothesis followed by experimentation) repeats many times, and the hypothesis
gets more and more certain. The hypothesis becomes a _______________________, which is a thoroughly tested
model that explains why things behave a certain way. Theories can never be ____________________; they are
always subject to additional research. Another outcome is that certain behavior is repeated many times. A scientific
____________ describes a relationship in nature that is supported by many experiments and for which no exception
has been found.
Identify the dependent variable and the independent variable in the following experiments.
a)
A student tests the ability of a given chemical to dissolve in water at three different temperatures.
independent variable: _________________________________
dependent variable: _________________________________
b) A farmer compares how his crops grow with and without phosphorous fertilizers.
independent variable: _________________________________
dependent variable: _________________________________
MATTER
Matter is anything that takes up __________________ and has mass. ______________ is the measure of the
amount of matter that an object contains. Virtually all of the matter around us consists of mixtures. A mixture can
be defined as something that has _____________________ composition. Soda is a mixture (carbon dioxide is
dissolved in it), and ____________________ is a mixture (it can be strong, weak or bitter). If matter is not uniform
throughout, then it is a _______________________ mixture. If matter is uniform throughout, it is homogeneous.
Homogeneous mixtures are called ___________________. A heterogeneous mixture contains regions that have
____________________ properties from those of other regions. When we pour sand into water, the resulting
mixture contains two distinct regions. ___________________ pavement, which has small rocks mixed with tarry
goo, is a simple example of a heterogeneous mixture. Oil-and-vinegar salad dressing, which has a layer of oil
floating on a layer of vinegar, is another example. Homogeneous mixtures (also known as solutions) are mixtures in
1
Unit 1- Matter and Change
which the composition is _______________________, there are no chunks or layers. Salt water,
___________________ ___________________ and dust free air (mixture of nitrogen, oxygen, argon, carbon
dioxide, water vapor and other gases) are examples of homogeneous mixtures. Brass (solid mixture of copper and
______________) is also a homogeneous mixture. Brass is a(n) _________________, which is a mixture of metals.
Since heterogeneous mixtures contain chunks or layers, they are often easier to separate than homogeneous
mixtures. A mixture of solid particles in a liquid can be separated by pouring the mixture through a
___________________ that traps the solid particles while the liquid passes through in a process called filtering.
Some simple methods also exist for separating homogeneous mixtures. A solid dissolved in a liquid solution can be
separated by letting it dry out in the process of ___________________. Mixtures are separated into pure
_____________________. A pure substance always has the same composition. Pure substances are either elements
or _________________________. Elements are substances that cannot be broken down into other substances
chemically or _______________________. Examples include sodium, carbon and aluminum. Compounds are
substances made of two or more ______________________ combined chemically. Compounds have properties
___________________________ from those of the original elements. Examples of compounds include water
(hydrogen and oxygen) and table salt (sodium and chlorine).
Classify each of the following as a pure substance, a homogeneous mixture or a heterogeneous mixture.
A. gasoline ____________________
B. copper metal ____________________
C. a stream with gravel at the bottom ____________________
D. maple syrup _________________
E. chunky peanut butter _____________
F. common salt ____________________
G. margarine ____________________
H. a Spanish omelet ________________
I. a multivitamin tablet ________________
J. oxygen gas ____________________
K. carbon dioxide gas _________________
PROPERTIES
The properties of matter describe the characteristics and behavior of matter, including the changes that matter
undergoes. _____________________ properties are characteristics that a sample of matter exhibits without any
change in its identity. This property can be observed and measured without _____________________ the
substance.
Examples of the physical properties of a chunk of matter include its:
1. __________________________________
2. _________________________________
3. __________________________________
4. _________________________________
5. __________________________________
6. _________________________________
7. __________________________________
Chemical properties are those that can be observed only when there is a change in the
___________________________ of the substance. Rusting is a chemical reaction in which iron combines with
__________________ to form a new substance, iron (III) oxide.
2
Unit 1- Matter and Change
Classify each of the following as a chemical or physical property.
density ___________________________
reactivity ___________________________
color _____________________________
melting point ________________________
Using the Chemistry Reference Tables, which substance has a
A. density = 19.31 g/cm3
_____________________
B. melting point = -119°C
_____________________
C. boiling point = 65°C
_____________________
D. melting point = -73°C
_____________________
Using the Chemistry Reference Tables, are the following substances soluble or insoluble in water?
A. zinc nitrate
_____________________
B. sodium sulfate
_____________________
C. calcium carbonate
_____________________
D. potassium oxide
_____________________
E. lead (II) fluoride
_____________________
F. barium hydroxide
_____________________
G. copper (II) sulfide
_____________________
H. silver chloride
_____________________
CHANGES
A physical change is a change in matter that does not involve a change in the chemical identity of individual
substances. The matter only changes in appearance. Examples: ______________, _________________,
__________________, _________________, ___________________, and _____________________. A chemical
property always relates to a chemical change, the change of one or more substances _____________ other
substances. Another term for chemical change is chemical ___________________. Indications of a chemical
reaction: __________________ absorbed or released, _________________ change, formation of a precipitate ______________ that separates from solution, and formation of a ___________. All matter is made of atoms, and
any chemical change involves only a rearrangement of the atoms. Atoms do not just appear. Atoms do not just
disappear. This is an example of the law of conservation of mass (or matter), which says that in a chemical change,
matter is neither ________________ nor destroyed. All chemical changes also involve some sort of energy change.
Energy is either taken in or __________________ ____________ as the chemical change takes place. Energy is the
capacity to do _________________. Work is done whenever something is moved. Chemical reactions that give off
heat energy are called ____________________ reactions. Chemical reactions that _________________ heat energy
are called endothermic reactions. Freezing, condensation and ___________________ are exothermic. Melting,
_______________________ and sublimation are endothermic.
3
Unit 1- Matter and Change
State whether each of the following is an endothermic or exothermic process.
1. melting of ice __________________________
2. combustion of gasoline __________________________
3. Natural gas is burned in a furnace. __________________________
4. When solid potassium bromide is dissolved in water, the solution gets colder.
_____________________
5. When concentrated sulfuric acid is added to water, the solution gets very hot.
_____________________
In an endothermic graph, the ____________________ have greater energy than the reactants. The change in energy
is a ______________________ value. In a(n) ____________________ graph, the reactants have greater energy
than the products. The change in ________________ is a negative value.
Sketch endothermic and exothermic graphs below. Label the reactants, products and activation energy.
Endothermic
Exothermic
E
n
e
r
g
y
E
n
e
r
g
y
Reaction Progress
Reaction Progress
Conservation of Energy: Energy can be neither created nor destroyed in ordinary changes (not nuclear); it can only
change _______________.
DENSITY
Density is the amount of matter (mass) contained in a unit of ___________________. Styrofoam has a low density
or small mass per unit of volume.
denisty 
mass
volume
D
m
V
Solve the following density problems.
1. The density of sugar is 1.59 g/cm3. Calculate the mass of sugar in 15.0 ml. (1 mL = 1 cm3).
4
Unit 1- Matter and Change
2. The density of helium is 0.178 g/L. Calculate the volume of helium that has a mass of 23.5 g.
3. A 14.95 g sample of gold has a volume of 0.774 cm3. Calculate the density of gold.
4. Balsa wood has a density of 0.12 g/cm3. What is the mass of a sample of balsa wood if its volume is 134 cm 3?
5. The density of ice at 0°C is 0.917 g/cm3. Calculate the volume of 145 g of ice.
6. The density of a sample of lead is found by the process of water displacement. A graduated cylinder is filled
with water to the 30.0 mL mark. The cylinder with the water is placed on an electronic balance and weighs
106.82 g. A piece of lead is added to the cylinder. The cylinder is reweighed with the water and the lead and
the scale reads 155.83 g. The volume of all the material in the cylinder is 34.5 mL. Calculate the density of the
lead.
7. The density of an unknown solid was found by the process of water displacement. The object was massed on an
electronic balance. The balance reads 125 g. 50.0 cm3 of water was poured into a 100.0 mL graduated
cylinder. The unknown sample was then gently placed into the graduated cylinder. The volume in the cylinder
rose to 60.7 cm3. Calculate the density of the unknown solid.
Practice / Homework
Reference Packet Study
Density: Identify the substance based on the density value given D = m / V
1.
D = 0.66g/cm3
3.
m = 20 g, V = 4.44 cm3
2.
D = 2.702g/cm3
4.
m = 3 g, V = 2.1 L
Melting and Boiling points: Identify the substance based on the given temperature value.
5.
Melting point = 801oC
7.
Melting point = 1455oC
6.
Boiling point = 79oC
8.
Boiling point = 1413oC
Solubility: Identify if the substance is soluble or insoluble.
9.
Lithium sulfate
11. Lead (IV) bromide
10. Strontium oxide
12. Ammonium carbonate
Identify each of the following as an element, a compound, a homogeneous mixture or a heterogeneous mixture.
13. Water
16. Silver
14. Cheerios in milk
17. Salsa
15. Apple juice
18. A bag of nuts and bolts
Identify each of the following as a chemical or physical property
19. Combustible
21. Volume
20. Mass
22. Ability to rust
5
Unit 1- Matter and Change
Identify each of the following as a chemical or physical change
23. Melts
26. Rips
24. Burns
27. Tarnishes
25. Dissolves
28. Shatters
Density Practice: Solve each problem below, writing the equation and showing the substitution. Provide a unit for
each answer.
1.
A block of aluminum occupies a volume of 15.0 mL and weighs 40.5 g. What is its density?
2.
Find the mass of gold that occupies 965 cm3 of space.
3.
Mercury metal is poured into a graduated cylinder that holds exactly 22.5 mL. The mercury used to fill the
cylinder weighs 306.0 g. From this information, calculate the density of mercury.
4.
Find the volume occupied by 250.0 g of O2.
5.
A cube of metal has a side length of 1.55 cm. If the sample is found to have a mass of 26.7 g, find the density
and identity of the metal.
6.
An irregularly-shaped sample of aluminum (Al) is put on a balance and found to have a mass of 43.6 g. The
student decides to use the water-displacement method to find the volume. The initial volume reading is
25.5 mL and, after the Al sample is added, the water level has risen to 41.7 mL. Find the density of the Al
sample in g/cm3. (Remember: 1 mL = 1 cm3.)
7.
A gas has a mass of 7914 g and takes up enough space to fill a room that is 2.00 m X 2.00 m X 2.50 m.
Determine the density of the gas in g/m3.
8.
A flask that weighs 345.8 g is filled with 225 mL of carbon tetrachloride. The weight of the flask and carbon
tetrachloride is found to be 703.55 g. From this information, calculate the density of carbon tetrachloride.
The Study of Matter Practice Test
Directions: Define and/or describe the following terms relating to the scientific method.
1.
Dependent Variable _________________________________________________
2.
Hypothesis __________________________________________________________
Directions: Indicate if the process listed is a physical or chemical change.
3.
Food digests _________________________________________________________
4.
Bending a piece of copper wire. _______________________________________
5.
Two clear liquids react to form a yellow clumps ___________________________
Directions: Solve the following problems. Show all work! Be sure to include the correct unit with your final
answer.
6. What is the density of a substance that has a volume of 2.8 cm3 and mass of 25 grams?
6
Unit 1- Matter and Change
7.
What is the density of a solid that has a volume of 4 cm3 and a mass of 6 grams?
Directions: For each sample of matter below, correctly classify it as a pure substance or a mixture.
8. Trail Mix
____________________
9.
Helium
____________________
Directions: For each, correctly classify as homogeneous or heterogeneous mixture.
10. Vegetable soup
______________________
11. Gatorade
______________________
12. Orange juice, with pulp
______________________
Multiple Choice Practice
13. The amount of mass per unit volume refers to the
a. Density
b. Specific weight
c.
d.
Volume
Weight
14. A substance that can be further simplified may be either
a. An element or a compound
b. An element or a mixture
c.
d.
A mixture or a compound
A mixture or an atom
15. A substance composed of two or more elements chemically united is called
a. An isotope
c. An element
b. A compound
d. A mixture
16. An example of a chemical change is the
a. Breaking of a glass bottle
b. Sawing of a piece of wood
c.
d.
Rusting of iron
Melting of an ice cube
17. A substance that cannot be further decomposed by ordinary chemical means is
a. Water
c. Sugar
b. Air
d. Silver
18. An example of a physical change is
a. The fermenting of sugar to alcohol
b. The rusting of iron
c.
d.
The burning of paper
A solution of sugar in water
19. What Kelvin temperature is equal to 25°C?
a. 248 K
b. 298 K
c.
d.
100 K
200 K
20. Which substance cannot be decomposed into simpler substances?
a. ammonia
b. aluminum
c. methane
d. methanol
7
Unit 1- Matter and Change
21. A compound differs from a mixture in that a compound always has a
a. homogeneous composition
b. maximum of two components
c.
d.
minimum of three components
heterogeneous composition
22. Which statement describes a chemical property?
a. Its crystals are a metallic gray.
b. It dissolves in alcohol.
c.
d.
It is a violet-colored gas.
It reacts with hydrogen.
23. To determine the density of an irregularly shaped object, a student immersed the object in 21.2 milliliters of
H2O in a graduated cylinder, causing the level of the H2O to rise to 27.8 milliliters. If the object had a mass of
22.4 grams, what was the density of the object.
a. 27.8 g / mL
c. 3.0 g / mL
b. 6.6 g / mL
d. 3.4 g/ mL
24. An experiment for a new asthma medication was set up into two groups. Group one was given the new drug for
asthma, while group 2 was given a sugar pill. The sugar pill serves as
a. Control
c. experimental variable
b. Constant
d. dependent variable
25. A scientist plants two rows of corn for experimentation. She puts fertilizer on row 1 but does not put fertilizer
on row 2. Both rows receive the same amount of water and light intensity. She checks the growth of the corn
over the course of 5 months. What is a constant in this experiment?
a. Plant height
c. Corn with fertilizer
b. Corn without fertilizer
d. Amount of water
26. Which sentence best states the importance of using control groups?
a. Control groups eliminate the need for large sample sizes, reducing the number of measurements
needed.
b. Control groups eliminate the need for statistical tests and simplify calculations.
c. Control groups provide a method by which statistical variability can be reduced.
d. Control groups allow comparison between subjects receiving a treatment and those receiving no
treatment
27. The measurable factor in an experiment is known as the:
a. Control
b. independent variable
c.
d.
constant
dependent variable
28. A student decides to set up an experiment to see if detergent affects the growth of seeds. He sets up 10 seed
pots. 5 of the seed pots will receive a small amount of detergent in the soil and will be placed in the sun. The
other 5 seed pots will not receive detergent and will be placed in the shade. All 10 seed pots will receive the
same amount of water, the same number of seeds, and the same type of seeds. He grows the seeds for two
months and charts the growth every 2 days. What is wrong with his experiment?
a. More than one variable is being tested.
b. The student should have a larger number of pots.
c. There is no way of measuring the outcome.
d. There is no control set-up.
8
Unit 1- Matter and Change
29. A scientific study showed that the depth at which algae were found in a lake varied from day to day. On clear
days, the algae were found as much as 6 meters below the surface of the water but were only 1 meter below the
surface on cloudy days. Which hypothesis best explains these observations?
a. Nitrogen concentration affects the growth of algae.
b. Precipitation affects the growth of algae.
c. Light intensity affects the growth of algae.
d. Wind currents affect the growth of algae.
9
Unit 2- Atomic Theory and Structure
ATOMIC THEORY
HISTORY OF THE ATOM
The original idea (400 B.C.) came from ______________________, a Greek philosopher. He expressed the belief
that all matter is composed of very small, indivisible particles, which he named atomos. John Dalton (1766-1844),
an English school teacher and chemist, proposed his atomic theory of matter in 1803. Dalton’s Atomic Theory
states that:
1. All matter is made of tiny __________________________ particles called atoms.
2. Atoms of the ____________ element are identical; those of different atoms are different.
3. Atoms of different elements combine in whole number ________________ to form compounds
4. Chemical reactions involve the rearrangement of atoms. No _______ atoms are created or destroyed.
PARTS OF THE ATOM
Because of Dalton’s atomic theory, most scientists in the 1800s believed that the atom was like a tiny solid ball that
could not be broken up into parts. In 1897, a British physicist, J.J. Thomson, discovered that this solid-ball model
was not accurate. Thomson’s experiments used a __________________ ray tube. It is a vacuum tube - all the air
has been pumped out. Because these rays originate at the ____________________, they are called cathode rays.
Thomson concluded that cathode rays are made up of invisible, _________________________ charged particles
referred to as electrons. From Thomson’s experiments, scientists had to conclude that atoms were not just neutral
_________________, but somehow were composed of electrically charged particles. Matter is not negatively
charged, so atoms can’t be negatively charged either. If atoms contained extremely light, negatively charged
particles, then they must also contain positively charged particles — probably with a much greater _____________
than electrons. J.J. Thomson said the atom was like ______________ pudding, a popular English dessert.
In 1909, a team of scientists led by Ernest Rutherford in England carried out the first of several important
experiments that revealed an arrangement far different from the plum pudding model of the atom. The
experimenters set up a lead-shielded box containing radioactive polonium, which emitted a beam of positively
charged subatomic particles through a small hole. The sheet of ________________ foil was surrounded by a screen
coated with zinc sulfide, which glows when struck by the positively charged particles of the beam. The
________________ particles were expected to pass through without changing direction very much because
Rutherford thought the mass was evenly distributed in the atom. Because most of the particles passed through the
foil, they concluded that the atom is nearly all _______________ ______________. Because so few particles were
deflected, they proposed that the atom has a small, dense, positively charged central core, called a
____________________. Alpha particles are deflected by it if they get close enough to the nucleus.
10
Unit 2- Atomic Theory and Structure
R.A. Millikan found the charge of an electron to be -1.60 x 10-19 Coulombs in his famous oil drop
experiment. In 1910, J.J. Thomson discovered that neon consisted of atoms of two different masses. Atoms of an
element that are chemically alike but differ in mass are called ______________________ of the element. Because
of the discovery of isotopes, scientists hypothesized that atoms contained still a third type of particle that explained
these differences in mass. Calculations showed that such a particle should have a mass ____________________ to
that of a proton but no electrical _______________. The existence of this neutral particle, called a neutron, was
confirmed in the early 1930s. James _________________ is given credit for discovering the neutron.
NAME
SYMBOL
CHARGE
RELATIVE MASS
1/2000
proton
no
MODERN VIEW OF THE ATOM
The atom has two regions and is ___-dimensional. The nucleus is at the ___________________ and contains the
protons and _____________________. The electron cloud is the region where you might find an electron and most
of the volume of an atom. The atomic _________________ of an element is the number of protons in the nucleus of
an atom of that element. The number of protons determines ____________________ of an element, as well as
many of its chemical and physical properties. Because atoms have no overall electrical charge, an atom must have
as many ____________________ as there are protons in its nucleus. Therefore, the atomic number of an element
also tells the number of electrons in a neutral atom of that element. The mass of a neutron is almost the same as the
mass of a ________________. The sum of the protons and neutrons in the nucleus is the ________________
number of that particular atom. _____________________ of an element have different mass numbers because they
have different numbers of _______________, but they all have the same atomic number.
11
Unit 2- Atomic Theory and Structure
AVERAGE ATOMIC MASS
The atomic mass is the weighted average mass of all the naturally occurring isotopes of that element.
To determine the average atomic mass, first calculate the contribution of each isotope to the average atomic mass,
being sure to convert each ___ to a fractional abundance. The average atomic mass of the element is the sum of the
mass contributions of each isotope.
Elements can be represented by using the symbol of the element, the mass number and the atomic number. The
mass number is the __________________ mass rounded to a whole number.
1.
Determine the following for the fluorine-19 atom.
a) number of protons
b) number of neutrons
d) atomic number
e) mass number
c) number of electrons
2. Repeat #1 for bromine-80.
3. If an element has an atomic number of 34 and a mass number of 78, what is the
a) number of protons
b) number of neutrons
c) number of electrons
d) complete symbol
4. If an element has 91 protons and 140 neutrons, what is the
a) atomic number
b) mass number
c) number of electrons
d) complete symbol
5. If an element has 78 electrons and 117 neutrons what is the
a) atomic number
b) mass number
c) number of protons
d) complete symbol
Practice/Homework
Isotopes and Subatomic Particles
Complete the chart
Isotope symbol
207
82
Atomic #
Mass
Protons
8
Electrons
77
54
12
13
Pb
38
238
92
U
75
33
Neutrons
9
As
32
16
S
12
50
Unit 2- Atomic Theory and Structure
Isotope symbol
65
29
Atomic #
Mass
Protons
11
Neutrons
12
7
14
Electrons
11
Cu
126
202
80
85
Hg
261
104
47
61
40
51
24
28
Rf
MOLES
We measure ________________ in grams. We measure volume in __________________. We count pieces in
_________________. The number of moles is defined as the number of __________________ atoms in exactly
____ grams of carbon-12. ____ mole is 6.022 x 1023 particles. 6.022 x 1023 is called __________________ number.
Representative particles are the smallest pieces of a substance. For a molecular compound it is a(n)
______________________. For an ionic compound it is a ______________________ ______________. For an
element it is a(n) ________________.
How many oxygen atoms are in the following?
a) CaCO3
b) Al2(SO4)3
How many total ions are in the following?
a) CaCl2
b) NaF
c) Al2S3
MOLE CONVERSIONS
1.
How many atoms of carbon are there in 1.23 moles of carbon?
2.
How many molecules of CO2 are in 4.56 moles of CO2?
3.
How many atoms of iron are in 0.600 moles of iron?
4.
How many moles are in 7.78 x 1024 formula units of MgCl2?
5.
How many moles of water are 5.87 x 1022 molecules of water?
6.
How many moles of aluminum are 1.2 x 1024 atoms of aluminum?
13
Unit 2- Atomic Theory and Structure
Calculate the number of particles (atoms, ions or molecules) in each of the following.
a) 3.4 moles Na2S
c) 1.77 x 10
-11
b) 0.0020 moles Zn
moles C
d) 92.35 moles O2
Calculate the number of moles in each of the following.
a) 3.4 x 1024 molecules HCl
b) 8.7 x 1021 atoms Zn
c) 1.77 x 1018 ions Al+3
d) 2.66 x 1026 atoms Cu
Representative
particles
Volume
Use
22.4
L
Use
6.022x1023
moles
Remember DIMO:
Divide
In
Multiply
Out
Use
molar mass
mass
MOLAR MASS
Molar mass is the generic term for the mass of one _____________. It may also be referred to as gram molecular
mass, gram formula mass, and gram atomic mass. The unit is ______________. To determine the molar mass of an
element, find the element’s symbol on the periodic table and round the mass so there is __________ digit beyond
the decimal.
Determine the molar mass of the each of the following elements.
a) sulfur (S)
b) chromium (Cr)
14
c) bromine (Br)
Unit 2- Atomic Theory and Structure
To determine the molar mass of a compound, find the mass of all elements in the compound.
If necessary, ___________________ an element’s mass by the subscript appearing beside that element in the
compound’s formula (or ________________ of the subscripts).
Calculate the molar mass of each of the following compounds.
a) Na2S
b) N2O4
c) C6H12O6
d) Ca(NO3)2
MASS-PARTICLE/MOLE CONVERSIONS
1.
How many atoms of lithium are in 1.00 g of Li?
2.
How many molecules of sodium oxide are in 42.0 g of Na 2O?
3.
How much would 3.45 x 1022 atoms of uranium (U) weigh?
4.
How many moles of magnesium are in 56.3 g of Mg?
5.
How many moles is 5.69 g of NaOH?
6.
How many grams of sodium chloride are in 3.45 moles of NaCl?
7.
How many moles is 4.8 g of CO2?
8.
How many grams is 9.87 moles of H2O?
9.
How many molecules are in 6.8 g of CH4?
10. What is the mass of 49.0 molecules of C6H12O6?
GASES
Many of the chemicals we deal with are gases. They are difficult to weigh, and we need to know how many moles
of gas we have. Two things affect the volume of a gas: temperature and pressure. Standard temperature is
______ ºC, and standard pressure is ______ atm. Standard temperature and pressure is abbreviated STP. At STP 1
mole of gas occupies ______ L. 22.4 L is called the _____________ volume. Avogadro’s Hypothesis - At the same
temperature and pressure equal volumes of gas have the same number of _______________________.
GAS CONVERSIONS
1.
What is the volume of 4.59 mole of CO2 gas at STP?
2.
How many moles is 5.67 L of O2 at STP?
3.
What is the volume of 8.8 g of CH4 gas at STP?
4.
How many grams is 16.2 L of O2 at STP?
Calculate the number of liters in each of the following.
a) 3.10 x 1024 molecules Cl2
b) 8.7 moles Ne
c) 2.77 x 1018 atoms He
d) 266 grams SO2
15
Unit 2- Atomic Theory and Structure
Homework/Practice
Part 1--Convert between particles and moles
1.
24 atoms of sodium = _____ moles of sodium atoms
2.
5 molecules of chlorine gas = _____ moles of chlorine molecules
3.
900 atoms of silver = _____ moles of silver atoms
4.
2.89 x 1023 molecules of ammonia = _____ moles of ammonia molecules
5.
15 moles of arsenic atoms = ______ atoms of arsenic
6.
4.00 x 103 moles of barium atoms = __________ atoms of barium
Part 2--Convert between mass and moles
7.
Calculate the mass of 1.000 mole of CaCl2
11. Calculate moles in 510.0 g of Al2S3
8.
Calculate grams in 3.0000 moles of CO 2
12. How many moles are in 27.00 g of H2O
9.
Calculate number of moles in 32.0 g of CH4
13. What is the mass of 2.55 moles Cu2CrO4
10. Calculate moles in 168.0 g of HgS
Part 3- Multiple steps
14. Arrange the following in order of increasing weight.
a.
10.4 g of sulfur
c.
6.33 x 1025 atoms of hydrogen
b.
0.179 moles of iron
d.
0.77 moles of N2
15. How many grams would 8.1  1021 molecules of sucrose (C12H22O11) weigh?
16. How many atoms are in a 2.0 kg ingot of gold? (Note mass units.)
17. What is the mass of 2.3 x 1024 molecules of KCl?
18. Calculate the number of molecules in 50.0 grams of H2SO4
19. Calculate the number of molecules in 100. grams of KClO4
20. Calculate the number of molecules in 8.76 grams of NaOH
21. Calculate the mass of 1.2 x 1022 molecules of Fe3(PO4)2
22. Calculate mass of 7.2 x 1024 molecules of Na2CO3
NUCLEAR CHEMISTRY
Nuclear chemistry is the study of the structure of _________________ nuclei and the changes they undergo. Marie
Curie named the process by which materials such as uranium give off rays radioactivity; the rays and particles
emitted by a radioactive source are called __________________. As you may recall, isotopes are atoms of the same
16
Unit 2- Atomic Theory and Structure
element that have different numbers of _________________. Isotopes of atoms with unstable nuclei are called
______________________. These unstable nuclei emit radiation to attain more stable atomic configurations in a
process called radioactive ________________. During radioactive decay, unstable atoms lose _________________
by emitting one of several types of radiation.
TYPES OF RADIATION
The three most common types of radiation are alpha (α), ____________ (β), and gamma (γ). An alpha particle (α)
has the same composition as a __________________ nucleus - two protons and ________ neutrons - and is
therefore given the symbol _________. The charge of an alpha particle is 2+ due to the presence of the two
___________________. Because of their mass and charge, alpha particles are relatively slow-moving compared
with other types of radiation. Thus, alpha particles are not very ________________________ - a single sheet of
paper stops alpha particles. A beta particle is a very-fast moving ______________________ that has been emitted
from a neutron of an unstable nucleus. Beta particles are represented by the symbol _________. The zero
superscript indicates the insignificant mass of an electron in comparison with the mass of a
____________________. The –1 subscript denotes the _____________________ charge of the particle. Beta
radiation consists of a stream of fast-moving electrons. Because beta particles are both lightweight and fast moving,
they have _____________________ penetrating power than alpha particles. A thin metal foil is required to stop
beta particles. Gamma rays are high-energy (_________________ wavelength) electromagnetic radiation. They are
denoted by the symbol __________. As you can see from the symbol, both the subscript and superscript are zero.
Thus, the emission of gamma rays does not change the __________________ number or mass number of a nucleus.
Gamma rays almost always accompany alpha and beta radiation, as they account for most of the energy loss that
occurs as a nucleus decays.
NAME
SYMBOL
FORMULA
4
2
Alpha
MASS
CHARGE
DESCRIPTION
He
β
-1
0
High energy
radiation
NUCLEAR STABILITY and DECAY
Radioactive nuclei undergo decay in order to gain _____________________. All elements with atomic numbers
greater than 83 are radioactive. Nuclear equations are used to show nuclear transformations. Balanced nuclear
equations require that both the ____________________ number and the mass number must be balanced.

When beryllium-9 is bombarded with alpha particles (helium nuclei), a neutron is produced. The balanced
nuclear reaction is given as: ________________________________________________
The atomic number (the number on the bottom) determines the identity of the element.
14
Unit 2- Atomic Theory and Structure

When nitrogen-14 is bombarded with a neutron, a proton is produced. The balanced nuclear equation can be
written as: _________________________________________________________

Polonium-230 undergoes alpha decay: ________________________________________________

Uranium-234 undergoes alpha decay: ________________________________________________

Cobalt-50 undergoes beta decay: ____________________________________________________
Provide symbols for each of the following: neutron ___________, proton ___________ or ___________, and the
positron ___________.

What element is formed when iron-60 undergoes beta decay? Give the atomic number and mass number of the
element. ____________

Write a balanced nuclear equation for the alpha decay of the following radioisotope, uranium-235.
____________________________________________________________

Nitrogen-12 decays into a positron and another element. Write the balanced nuclear equation.
____________________________________________________________

Uranium-238 is bombarded with a neutron. One product forms along with gamma radiation. Write the
balanced nuclear equation.
____________________________________________________________

Nitrogen-14 is bombarded with deuterium (hydrogen-2). One product forms along with an alpha particle.
Write the balanced nuclear equation.
____________________________________________________________
RADIOACTIVE DECAY RATES
Radioactive decay rates are measured in half-lives. A half-life is the time required for one-half of a radioisotope’s
nuclei to ________________ into its products. For example, the half-life of the radioisotope strontium-90 is 29
years. If you had 10.0 g of strontium-90 today, 29 years from now you would have 5.0 g left. The decay continues
until negligible strontium-90 remains.

Iron-59 is used in medicine to diagnose blood circulation disorders. The half-life of iron-59 is 44.5 days. How
much of a 2.000-mg sample will remain after 133.5 days?

Cobalt-60 has a half-life of 5.27 years. How much of a 10.0 g sample will remain after 21.08 years?

Carbon-14 has a half-life of 5730 years. How much of a 250. g sample will remain after 5730 years?
FISSION and FUSION
Heavy atoms (mass number > 60) tend to break into smaller atoms, thereby increasing their
________________________. Using a neutron to split a nucleus into fragments is called nuclear
15
Unit 2- Atomic Theory and Structure
_______________________. Nuclear fission releases a large amount of energy and several neutrons. Since
neutrons are products, one fission reaction can lead to more fission reactions, a process called a ________________
reaction. A chain reaction can occur only if the starting material has enough mass to sustain a chain reaction; this
amount is called __________________ mass. The _____________________ of atomic nuclei is called nuclear
fusion. For example, nuclear fusion occurs within the Sun, where hydrogen atoms fuse to form
__________________ atoms. Fusion reactions can release very large amounts of energy but require extremely high
temperatures. For this reason, they are also called _____________________________ reactions.
APPLICATIONS OF NUCLEAR REACTIONS
Geiger counters, scintillation counters, and film badges are devices used to detect and measure radiation. Geiger
counters use ________________________ radiation, which produces an electric current in the counter, to rate the
strength of the radiation on a scale. Film badges are often used to monitor the approximate radiation
______________________ of people working with radioactive materials. Scintillation counters measure ionizing
radiation. With proper safety procedures, radiation can be useful in industry, in scientific experiments, and in
medical procedures. Nuclear power plants use the process of nuclear fission to produce heat in nuclear reactors.
The heat is used to generate steam, which is then used to drive __________________ that produce electricity. A
radiotracer is a radioisotope that emits non-ionizing radiation and is used to signal the presence of an element or of a
specific substance. Radiotracers are used to detect ______________________ and to analyze complex chemical
reactions. Ionizing radiation has many uses. A(n) _______________ is ionizing radiation, and ionizing radiation
can be used in medicine to kill cancerous cells. Most medical devices require sterilization after they are packaged,
and another trend has been the move to sterilization by __________________ radiation as opposed to other methods
such as ethylene oxide gas. Advantages of gamma irradiation include ______________, cost-effectiveness, and the
elimination of the need for special packaging. Chemical reaction rates are greatly affected by changes in
temperature, _____________________, and concentration, and by the presence of a catalyst. In contrast, nuclear
reaction rates remain ____________________ regardless of such changes. In fact, the half-life of any particular
radioisotope is constant. Because of this, radioisotopes, especially carbon-14, can be used to determine the
____________ of an object. The process of determining the age of an object by measuring the amount of a certain
radioisotope remaining in that object is called radiochemical dating.
EFFECTS OF NUCLEAR REACTIONS
Any exposure to radiation can damage living ____________. Gamma rays are very dangerous because they
penetrate ______________________ and produce unstable and reactive molecules, which can then disrupt the
normal functioning of cells. The amount of radiation the body absorbs (a dose) is measured in units called rads and
____________. Everyone is exposed to radiation, on average 100–300 millirems per year. A dose exceeding
____________ rem can be fatal.
16
Unit 2- Atomic Theory and Structure
Atomic Theory, The Mole, and Nuclear Chemistry- Practice Test
1.
Given the work of Dalton, please check the box for the postulate(s) that have since been proven to be incorrect.
Explain what we now know to be the true case.
[] All atoms of a specific element are identical.
[] Compounds consist of atoms of different elements combined together.
[] Atoms of different elements have different masses.
Directions: For the scientist listed below, explain what was done in the experiment, what knowledge was
developed as a result.
2.
Rutherford
3.
Thomson
Directions: Fill in the table for the following isotopes.
4.
5.
6.
7.
8.
Isotope
H-1
Cu-65
Atomic #
Mass #
Protons
18
40
19
9
Neutrons
Electrons
What is the charge of a beta particle?
Directions: Solve the following problems be sure to include the correct unit with your final answer.
9.
Given the equation:
X 
4
2
He +
220
84
Po The nucleus represented by X is
10. How many moles of sodium are 6.02 x 10 23 atoms of sodium?
11. What is the mass of 6 moles of Carbon?
12. How many atoms are in 45 g of Neon?
Multiple Choice Practice
13. What is the approximate formula mass of Ca(NO3)2
a. 70
c. 102
b. 82
d. 150
e.
164
14. How many molecules are in 1 mole of water?
a. 3
c.
b. 54
d.
e.
3 (6.02 x 1023)
6.02 x 1023
2 (6.02 x 1023)
15. How many atoms are represented in the formula Ca3(PO4)2
a. 5
c. 9
b. 8
d. 12
16. Four grams of hydrogen gas at STP contain
a. 6.02 x 1023 atoms
b. 12.04 x 1023atoms
c. 12.04 x 1046atoms
e.
d.
e.
17
13
1.2 x 1022molecules
12.04 x 1023molecules
Unit 2- Atomic Theory and Structure
17. What is the mass in grams of 1 mole of KAl(SO4)2•12H2O
a. 132
c. 394
b. 180
d. 474
e.
516
18. Compared to the charge and mass of a proton, an electron has
a. the same charge and a smaller mass
b. the same charge and the same mass
c. an opposite charge and a smaller mass
d. an opposite charge and the same mass
19. When alpha particles are used to bombard gold foil, most of the alpha particles pass through undeflected.
This result indicates that most of the volume of a gold atom consists of ____.
a. deuterons
c. protons
b. neutrons
d. unoccupied space
20. A proton has approximately the same mass as
a. a neutron
b. an alpha particle
c.
d.
a beta particle
an electron
21. A neutron has approximately the same mass as a
a. an alpha particle
b. a beta particle
c.
d.
an electron
a proton
22. Which symbols represent atoms that are isotopes?
a. C-14 and N-14
b. O-16 and O-18
c.
d.
I-131 and I-131
Rn-222 and Ra-222
23. Which atom contains exactly 15 protons?
a. P-32
b. S-32
c.
d.
O-15
N-15
24. An ion with 5 protons, 6 neutrons, and a charge of 3+ has an atomic number of
a. 5
b. 6
c. 8
d.
11
25. What is the mass number of an atom which contains 28 protons, 28 electrons, and 34 neutrons?
a. 56
b. 62
c. 90
d.
28
26. What is the total number of atoms represented in the formula CuSO 4 . 5H2O?
a. 8
b. 13
c. 21
d.
27
27. What is the gram formula mass of K2CO3?
a. 138 g
b. 106 g
c.
d.
28. What is the total number of atoms contained in 2.00 moles of nickel?
a. 58.9
c.
b. 118
d.
18
99 g
67 g
6.02 x 1023
1.2 x 1024
Unit 2- Atomic Theory and Structure
29. What is the total number of moles of hydrogen gas contained in 9.03 x 10 23 molecules
a. 1.5 moles
c. 6.02 moles
b. 2.00 moles
d. 9.03 moles
30. What is the mass in grams of 3.0 x 1023 molecules of CO2?
a. 22 g
b. 44 g
c.
d.
66 g
88 g
31. The amount of substance having 6.022 x 1023 of any kind of chemical unit is called a(n):
a. formula
c. mole
b. mass number
d. atomic weight
32. The total number of atoms in the formula of aluminum dichromate, Al2(Cr2O7)3 is:
a. 5
b. 29
c. 17
d.
11
33. The total number of OXYGEN atoms in the formula of aluminum dichromate, Al 2(Cr2O7)3 is:
a. 21
b. 10
c. 29
d.
7
34. The formula mass of calcium hydroxide, Ca(OH) 2 is:
a. 57.05 grams
b. 74.10 grams
c.
d.
128 grams
97.07 grams
35. The formula mass of ammonium chlorate, NH4ClO3 is:
a. 101.5 g
b. 78.06 g
c.
d.
211.43 g
172.40 g
36. What is the molar mass of the gas butane, C4H10?
a. 13.02 grams
b. 485.2 grams
c.
d.
68 24 grams
58.14 grams
37. The molar mass of sodium chloride, NaCl is:
a. 69.71 grams
b. 2 grams
c.
d.
6.022 x 1023 grams
58.44 grams
38. The formula mass of magnesium hydroxide, Mg(OH) 2 is:
a. 42.33 grams
b. 58.33 grams
c.
d.
41.32 grams
5 grams
39. What is the mass in grams of 3 moles of water molecules, H 2O?
a. 54.06 grams
b. 21.02 grams
c.
d.
0.166 grams
6.01 grams
19
Unit 2- Atomic Theory and Structure
40. What is the mass of 4 moles of hydrogen molecules (H2)?
a. 4.04 grams
b. 8.08 grams
c.
d.
3.96 grams
1.98 grams
41. What is the mass in grams of 10 moles of ammonia, NH3?
a. 170.4 grams
b. 0.587 grams
c.
d.
1.704 grams
27.04 grams
42. How many moles of water molecules, H2O, are present in a 42 gram sample of water?
a. 23.98 moles
c. 2.33 moles
b. 0.429 moles
d. 757 moles
43. How many moles of methane molecules, CH4, are in 80 grams of methane?
a. 0.201 moles
c. 6.022 x 1080 moles
b. 4.98 moles
d. 1284 moles
44. How many moles of calcium hydroxide, Ca(OH)2 are in 150 grams of the compound?
a. 2.02 moles
c. 0.494 moles
b. 224.1 moles
d. 11115 moles
45. About how many atoms of helium would be found in 2 grams of helium?
a. 4.00260
c. 2
b. 6.02 x 1023
d. 3.01 x 1023
46. By knowing the number of electrons in a neutral atom, you should also be able to determine
a. the number of neutrons in the neutral atom
b. the number of protons in the neutral atom
c. the atomic number of the neutral atom
d. the mass of the neutral atom
e. two of these
47. How many oxygen atoms are there in one formula unit of Al2(SO4)3?
a. 3
b. 4
c.
7
d.
12
48. The average mass of a magnesium atom is 24.31 amu. If you were able to select and measure a single atom
of magnesium, the chance that you would select an atom of mass 24.31 is about:
a. 0%
d. greater than 50%
b. 0.31%
e. 100%
c. 24.31%
20
Unit 2- Atomic Theory and Structure
49. Which of the following arrangements represent different isotopes of the same element?
i.
ii.
iii.
iv.
v.
a.
b.
c.
12 protons, 11 neutrons, 12 electrons
11 protons, 12 neutrons, 11 electrons
10 protons, 12 neutrons, 12 electrons
11 protons, 12 neutrons, 10 electrons
12 protons, 12 neutrons, 12 electrons
1 and 5
2 and 4
2, 3, 4 and 5
d.
e.
all of these qualify
None of these qualify
50. If the abundance of 6Li (6.015121 amu) is 7.500% and the abundance of 7Li (7.016003 amu) is 92.500%,
what is the average atomic mass?
a. 6.0750 amu
c. 6.9250 amu
b. 6.0902 amu
d. 6.9409 amu
51. Which of following is not true of the carbon-14 atom?
a. It has six protons
b. It has an average mass of 12.011 amu.
c. It has six electrons
d. It has eight neutrons
e. It is the less common than carbon-12
52. Which of these is the correct number of particles in this nuclide?
a.
b.
c.
d.
e.
79
34
Se 2
34 protons, 79 neutrons, 2 electrons
34 protons, 45 neutrons, 32 electrons
34 protons, 45 neutrons, 2 electrons
34 protons, 45 neutrons, 36 electrons
34 protons, 113 neutrons, 36 electrons
53. Which of the following have equal numbers of neutrons?
a.
b.
c.
I, II and III
II and III
I and V
d.
e.
I and IV
II, III and IV
54. The element hafnium (Hf) has five stable isotopes. The correct number of nuclear particles in an atom of
hafnium-178 is:
a. 72 protons, 178 neutrons
d. 72 protons, 106 neutrons
b. 72 protons, 72 electrons
e. 72 protons, 106 neutrons,
c. 106 protons, 72 neutrons
72 electrons
21
Unit 2- Atomic Theory and Structure
55. J.J. Thomson's model of the atom can be summarized with the visual image of:
a. planets orbiting the sun
d. a small central nucleus and an
b. plum pudding
electron cloud
c. bees around a hive
e. none of the above
56. The number of errors contained in the table below is:
Nuclide
protons
neutrons
electrons
mass
Sodium-23
11
12
11
23
Cobalt-59
27
22
27
59
Tungsten-184
74
110
184
184
Fluorine-19
10
9
10
19
a.
b.
57.
10
5
one
two
c.
d.
three
four
e.
five
e.
4
2
B + _____  13 H + 2 24 He In the equation shown here, the missing particle is:
a.
b.
1
1
1
0
H
n
c.
d.
o
1
1
1
e
p
58. Identify the missing particle in the following nuclear reaction:
a.
37
18
Ar
b.
38
18
37
19
Ar
K → _____ +
36
c. 18 Ar
0
1
He
e
d.
59. For the most common types of radioactive decay, the order of least penetrating to human tissue, to most
penetrating to human tissue is:
a. gamma, beta, alpha
c. beta, gamma, alpha
b. alpha, beta, gamma
d. gamma, alpha, beta
60. Very large nuclei tend to be unstable because of the:
a. repulsive forces between protons
b. attraction of protons for neutrons
c. repulsive forces between neutrons
d. attraction of electrons for the positively charged nucleus
e. repulsive forces between electrons
61. An alpha () particle is essentially a ____________________ nucleus.
a. plutonium
c. hydrogen
b. helium
d. uranium
22
e.
carbon-12
37
20
Ca
Unit 2- Atomic Theory and Structure
62. Phosphorus-15 has a half-life of 14 days. What proportion of the original phosphorus-15 remains after
8 weeks?
a. 1/2
c. 1/4
e. 1/8
b. 1/16
d. 1/32
63. The nuclide radium-226 is the daughter nuclide resulting from the  decay of what parent nuclide?
a. radon-222
d. thorium-228
b. polonium-214
e. radium-225
c. thorium-230
64. An electron emitted from the nucleus during some kinds of radioactive decay is known as:
a. A gamma ray
c. A beta () particle
b. A positron
d. An alpha () particle
65. A process in which a very heavy nucleus splits into more-stable nuclei of intermediate mass is called:
a. nuclear fission
d. nuclear fusion
b. radiocarbon dating
e. radioactive decay
c. a chain reaction
23
Unit 3- Electrons and Periodicity
ELECTRONS IN ATOMS
LIGHT
Light is a kind of electromagnetic _____________________. All forms of electromagnetic radiation move at
3.00 x 108 m/s. The ______________ is the baseline of a wave. The crest is the high point on a wave, and the
trough is low point on a wave. The amplitude of a wave is the wave’s _____________ from the origin to a crest, or
from the origin to a trough. Wavelength (represented by λ, the Greek letter lambda) is the ___________________
distance between equivalent points on a continuous wave. Wavelength is the distance from crest to crest or trough
to trough and is usually expressed in meters (m). _____________________ (represented by f ) is the number of
“waves” that pass a given point per second, and the units are cycles/sec or hertz (Hz)
c = fλ
c = the speed of light
Frequency and wavelength are __________________ related, which means that as one goes up the other goes
_____________. Different frequencies of light correspond to different colors of light. In 1900, the German
physicist Max Planck began searching for an explanation as he studied the light emitted from
___________________ objects. Matter can gain or lose energy only in small, specific amounts called
_______________. That is, a quantum is the minimum amount of energy that can be gained or lost by a(n)
____________. That is, while a beam of light has many wavelike characteristics, it also can be thought of as a
stream of tiny particles, or bundles of energy, called ________________. Thus, a photon is a particle of
electromagnetic radiation with no _____________ that carries a quantum of energy. Planck went further and
demonstrated mathematically that the energy of a quantum is ___________________ related to the frequency of the
emitted radiation.
E=hf
E = energy of the photon (J – Joules); f = frequency (Hz); h = Planck’s constant (J.s)
The energy of radiation increases as the radiation’s frequency, f, __________________.
Scientists knew that the wave model of light could not explain a phenomenon called the ____________________
effect. In the photoelectric effect, electrons, called __________________________, are emitted from a metal’s
surface when light of a certain _______________________ shines on the surface. Einstein proposed that for the
photoelectric effect to occur, a photon must possess, at a minimum, the energy required to _______________ an
electron from an atom of the metal.
THE BOHR MODEL OF THE ATOM
Niels Bohr, a young Danish physicist working in Rutherford’s laboratory in 1913, suggested that the single electron
in a ___________________ atom moves around the nucleus in only certain allowed circular orbits. The atom
looked like a miniature _________________ system. The nucleus is represented by the sun, and the electrons act
like the planets. The orbits are circular and are at different levels. Amounts of ___________________ separate one
level from another. (Modern View: The atom has two regions and is 3-dimensional. The nucleus is at the
24
Unit 3- Electrons and Periodicity
_________________ and contains the protons and neutrons. The electron _________________ is the region where
you might find an electron and most of the volume of an atom.) Bohr proposed that electrons must have enough
energy to keep them in constant motion around the ___________________. Electrons have energy of motion that
enables them to overcome the attraction of the _________________ nucleus. Further away from the nucleus means
more energy. Electrons reside in ________________ levels.
THE QUANTUM MECHANICAL MODEL
Building on Planck’s and Einstein’s concepts of ____________________ energy (quantized means that only certain
values are allowed), Bohr proposed that the hydrogen atom has only certain allowable energy ______________.
The lowest allowable energy state of an atom is called its _______________ state. When an atom gains energy, it is
said to be in a(n) __________________ state. When the atom is in an excited state, the electron can drop from the
higher-energy orbit to a _______________-energy orbit. As a result of this transition, the atom emits a
____________________ corresponding to the difference between the energy levels associated with the two orbits.
ATOMIC AND EMISSION SPECTRA
By heating a gas of a given element with electricity, we can get it to give off _______________. Each element
gives off its own characteristic colors. The spectrum can be used to __________________ the atom. These are
called line _______________. Each is unique to an element. The spectrum of light released from excited atoms of
an element is called the _________________ spectrum of that element. As the electrons fall from the excited state,
they __________________ energy in the form of light. The further they fall, the ________________ the energy.
This results in a higher frequency.
Use the Chemistry Reference Tables to answer the following:
(a) An electron falls from energy level 5 to 3. What is the wavelength of the light emitted?
(b) An electron falls from energy level 6 to 2. What is the wavelength of the light emitted?
(c) An electron falls from energy level 3 to 1. What type of electromagnetic radiation is emitted (infrared,
visible or ultraviolet)?
(d) An electron falls from energy level 4 to 2. What type of electromagnetic radiation is emitted (infrared,
visible or ultraviolet)?
(e) An electron falls from energy level 5 to 2. What color of visible light is emitted?
(f) An electron falls from energy level 3 to 2. What color of visible light is emitted?
MORE QUANTUM MECHANICAL MODEL
Like Bohr’s model, the quantum mechanical model limits an electron’s energy to certain values. The space around
the nucleus of an atom where the atom’s electrons are found is called the electron ________________. A threedimensional region around the nucleus called an atomic __________________ describes the electron’s probable
location. In general, electrons reside in principal ________________ levels. As the energy level number increases,
25
Unit 3- Electrons and Periodicity
the orbital becomes _______________, the electron spends more time ___________________ from the nucleus, and
the atom’s energy level increases. Principal energy levels contain energy ___________________. Principal energy
level 1 consists of a single sublevel, principal energy level 2 consists of __________ sublevels, principal energy
level 3 consists of three sublevels, and so on. Sublevels are labeled s, p, d, or f. The s sublevel can hold 2 electrons,
the p sublevel can hold _____ electrons, the d sublevel can hold 10 electrons, and the f sublevel can hold 14
electrons. Sublevels contain __________________. Each orbital may contain at most ________ electrons. There is
one s orbital for every energy level, and the s orbital is ____________________ shaped. They are called the 1s, 2s,
3s, etc… orbitals. The p orbitals start at the second energy level, reside along ______ different directions and have 3
different ________________ shapes. The d orbitals start at the ________________ energy level and have ____
different shapes. The f orbitals start at the fourth energy level and have ______ different shapes.
ELECTRON CONFIGURATIONS
Electron configurations represent the way electrons are arranged in atoms. The Aufbau principle states that
electrons enter the __________________ energy first. This causes difficulties because of the ________________ of
orbitals of different energies. At most there can be only 2 electrons per orbital, and they must have
__________________ “spins.” Hund’s rule states that when electrons occupy orbitals of equal energy, they don’t
_________ up with an electron of opposite spin until they have to.
Let’s determine the electron configuration for phosphorus. ______________________________
Let’s determine the electron configuration for chromium. _______________________________
 Write the electron configuration for aluminum (Al). ________________________________
 Write the electron configuration for neon (Ne). ____________________________________
 Write the electron configuration for calcium (Ca). __________________________________
 Write the electron configuration for iron (Fe). _____________________________________
 Write the electron configuration for bromine (Br). _________________________________
To identify an element with a given electron configuration, add the _________________ numbers together and find
the element with that atomic number.
Directions: Identify the element with the following electron configuration:
a.
1s2 2s2 2p6 3s2 3p4 _________________________________
b.
1s2 2s2 2p6 3s2 3p6 4s2 3d9 _________________________________
c.
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2 _________________________________
Electron Configuration Using a Noble Gas Abbreviation - In order to write this type of configuration, find the
_______________ gas (from Group 8A) that comes before the element in question. Put the symbol for the noble
gas in _____________________ and then write the part of the configuration that follows to reach the desired
element.
26
Unit 3- Electrons and Periodicity
Write the electron configuration using a noble gas abbreviation for:
 magnesium (Mg) _________________________
• nickel (Ni) ___________________
 fluorine (F) _________________________
• silicon (Si) ___________________
 zirconium (Zr) _________________________
VALENCE ELECTRONS
The electrons in the ______________________ energy level are called valence electrons. You can also use the
periodic table as a tool to predict the number of valence electrons in any atom in Groups 1, 2, 13, 14, 15, 16, 17, and
18. All atoms in Group 1, like hydrogen, have __________ valence electron. All atoms in Group 2 have two, in
Group 13 have _______, in Group 14 have four, in Group 15 have five, in Group 16 have six, and in Group 17 have
________ valence electrons. All atoms in Group 18 have eight valence electrons, except helium which only has
two. All atoms in sublevels d and f have _________ valence electrons.
How many valence electrons does each of the following elements have?
a) carbon (C)
c) iron (Fe)
b) bromine (Br)
d) potassium (K)
e) aluminum (Al)
LEWIS DOT DIAGRAMS
Because valence electrons are so important to the behavior of an atom, it is useful to represent them with symbols.
A Lewis dot diagram illustrates ___________________ electrons as dots (or other small symbols) around the
chemical symbol of an element. Each dot represents _____________ valence electron. In the dot diagram, the
element’s symbol represents the core of the atom - the nucleus plus all the _______________ electrons.
Write a Lewis dot diagram for
a) chlorine
b) calcium
c) potassium
PERIODIC TABLE
HISTORY
The Russian chemist, Dmitri ______________________ was studying the properties of the elements and realized
that the chemical and physical properties of the elements repeated in an orderly way when he organized the elements
according to increasing atomic ___________. Mendeleev later developed an improved version of his table with the
elements arranged in horizontal ___________. This arrangement was the forerunner of today’s periodic table.
Patterns of changing properties repeated for the elements across the horizontal rows. Elements in vertical
___________________ showed similar properties. Mendeleev grouped elements in columns by similar properties
in order of increasing atomic mass. He found some inconsistencies and felt that the properties were more important
than the mass, so he switched order. Mendeleev left some _____________ in his periodic table, deciding there must
be undiscovered elements. He predicted their properties before they were found. Mendeleev is considered to be the
27
Unit 3- Electrons and Periodicity
_________________ of the Periodic table. This repeated pattern (when Mendeleev grouped elements in columns by
similar properties) is an example of __________________ in the properties of elements. Periodicity is the tendency
to recur at regular intervals. By 1860, scientists had already discovered _________ elements and determined their
atomic masses.
THE MODERN PERIODIC TABLE
Fifty years after Mendeleev, the British scientist Henry ________________ discovered that the number of protons in
the nucleus of a particular type of atom was always the same. When atoms were arranged according to increasing
atomic ___________________, the few problems with Mendeleev's periodic table disappeared. Because of
Moseley's work, the modern periodic table is based on the atomic numbers of the elements. The statement that the
physical and chemical properties of the elements repeat in a regular pattern when they are arranged in order of
increasing atomic number is known as the periodic _____________. On the periodic table a _________________,
sometimes also called a series, consists of the elements in a horizontal row. A __________________, sometimes
also called a family, consists of the elements in a vertical column. Elements are placed in columns by similar
properties.
The elements in the A groups are called the __________________ elements. The B groups are called the
____________________ elements. The two rows at the bottom of the table are called the inner transition elements.
Group 1A elements are the _________________ metals. Group 1A elements have ______ valence electron and form
_______ ions after losing the one valence electron. Group 2A elements are the alkaline earth metals. Group 2A
elements have ______ valence electrons and form 2+ ions after losing the two __________________ electrons.
Group 3A is called the _________________ group. Group 3A elements have ________ valence electrons and form
3+ ions after losing the three valence electrons. Group _______ is called the carbon group. Group 4A elements
have four valence electrons and form 4+ ions after ___________________ the four valence electrons or 4- ions after
___________________ four additional electrons. Group 5A is called the _____________________ group. Group
5A elements have five valence electrons and form ________ ions after gaining three more electrons. Group 6A is
called the oxygen group. Group 6A elements have _______ valence electrons and form 2- ions after
____________________ two more electrons. Group 7A is called the ____________________. Group 7A elements
have seven valence electrons and form 1- ions after gaining one more electron. The word halogen is from the Greek
words for “______________ former” so named because the compounds that halogens form with metals are salt-like.
Group 8A elements are the ________________ gases. Group 8A elements have eight valence electrons except for
helium which only has ________. The noble gases, with a full complement of valence electrons, are generally
unreactive. All transition elements have _______ valence electrons.
 How many valence electrons are in an atom of each of the following elements?
a) Magnesium (Mg) ______
b) Selenium (Se) ______
28
c) Tin (Sn) _____
Unit 3- Electrons and Periodicity
METALS, NONMETALS AND METALLOIDS
Metals are elements that have ________________, conduct ____________ and electricity, and usually bend without
breaking. Most metals have one, two, or three valence electrons. All metals except _________________ are solids
at room temperature; in fact, most have extremely _____________ melting points. A metal’s
___________________ is its ability to react with another substance.
 Consult the “Activity Series of Metals” in the Chemistry Reference Tables to determine the more active metal.
a) cobalt (Co) or manganese (Mn) ________
b) barium (Ba) or sodium (Na) ________
Although the majority of the elements in the periodic table are _________________, many nonmetals are abundant
in nature. Most nonmetals don’t conduct electricity, are much poorer conductors of heat than metals, and are
__________________ when solid. Many are ______________ at room temperature; those that are solids lack the
luster of metals. Their _____________________ points tend to be lower than those of metals. With the exception
of carbon, nonmetals have five, six, seven, or eight valence electrons. A nonmetal’s reactivity is its ability to react
with another substance.
 Consult the “Activity Series of Metals” in the Chemistry Reference Tables to determine the less active
nonmetal.
a) fluorine (F2) or chlorine (Cl2) _________
b) chlorine (Cl2) or iodine (I2) _________
__________________ have some chemical and physical properties of metals and other properties of nonmetals. In
the periodic table, the metalloids lie along the border between metals and nonmetals. Some metalloids such as
silicon, germanium (Ge), and arsenic (As) are _____________________. A semiconductor is an element that does
not conduct electricity as well as a ________________, but does conduct slightly better than a nonmetal.
 Match each element in Column A with the best matching description in Column B. Each
Column A element
may match more than one description from Column B.
Column A
Column B
1. strontium
a. halogen
2. chromium
b. alkaline earth metal
3. iodine
c. representative element
d. transition element
PERIODIC TRENDS
Because the periodic table relates group and period numbers to valence electrons, it’s useful in predicting atomic
structure and, therefore, ______________________ properties.
29
Unit 3- Electrons and Periodicity
Atomic Radius
Atomic radius is half the distance between two __________________ of a diatomic molecule. Atomic size is
influenced by two factors: (1) energy level – A _________________ energy level is further away. (2) charge on
nucleus - More charge (_________________) pulls electrons in closer. As you go down a ___________________,
each atom has another energy level so the atoms get bigger. As you go across a period, the radius gets
____________________. Atoms are in the same energy level, but as you move across the chart, atoms have a
greater ___________________ charge (more protons). Therefore, the outermost electrons are closer.
 Choose the element from the pair with the larger atomic radius.
a) lithium (Li) or beryllium (Be) _________
b) silicon (Si) or tin (Sn) _________
 Choose the element from the pair with the smaller atomic radius.
a) silver (Ag) or gold (Au) _________
b) cesium (Cs) or barium (Ba) _______
Ionic Radius
When an atom gains or loses one or more electrons, it becomes a(n) ______________. Because an electron has a
negative charge, gaining electrons produces a _______________________ charged ion, an anion, whereas losing
electrons produces a positively charged ion, a ________________. As you might expect, the loss of electrons
produces a positive ion with a radius that is ___________________ than that of the parent atom. Conversely, when
an atom gains electrons, the resulting negative ion is larger than the parent atom. Practically all of the elements to
the _____________ of group 4A of the periodic table commonly form positive ions. As with neutral atoms,
___________________ ions become smaller moving across a period and become larger moving down through a
group. As you go down a group, you are adding a(n) _________________ level. Ions get bigger as you go down.
Most elements to the right of group 4A (with the exception of the noble gases in group 8A) form negative ions.
These ions, although considerably larger than the positive ions to the left, also decrease in ______________ moving
across a period. Like the positive ions, the negative ions increase in size moving down through a group. Across the
period, nuclear charge __________________ so they get smaller. Energy level changes between anions and cations.
 Choose the element from the pair with the smaller radius.
a) silver (Ag) or the silver ion (Ag1+) _________
b) oxygen (O) or the oxygen ion (O2-) _________
 For each of the following pairs, predict which atom is larger.
a) Mg, Sr _________
b) Sr, Sn _________
d) Ge, Br _________
e) Cr, W _________
c) Ge, Sn _________
 For each of the following pairs, predict which atom or ion is larger.
a) Mg, Mg2+ _________
b) S, S2– _________
d) Cl–, I– _________
e) Na+, Al3+ _________
30
c) Ca2+, Ba2+ ______
Unit 3- Electrons and Periodicity
Ionization Energy
Ionization energy (IE) is the amount of energy required to completely _____________________ an electron from a
gaseous atom. Removing one electron makes a ________ ion. The energy required to do this is called the first
ionization energy. The _____________________ the nuclear charge (# of protons), the greater IE. The distance
from the ____________________ increases IE. As you go down a group, first IE decreases because the electron is
further away, thus there is more shielding by the _______________ electrons from the pull of the positive nucleus.
All the atoms in the same period have the same energy level. They have the same shielding, but as you move across
the chart there is a(n) _____________________ nuclear charge. Therefore, IE generally increases from left to right.
 Choose the element from the pair with the greater ionization energy.
a) silver (Ag) or iodine (I) _________
b) oxygen (O) or selenium (Se) ________
 Choose the element from the pair with the smaller ionization energy.
a) chromium (Cr) or tungsten (W) ______
b) sodium (Na) or magnesium (Mg) _______
Electronegativity
Electronegativity is the tendency for an atom to ___________________ electrons to itself when it is chemically
combined with another element. Large electronegativity means it _______________ the electron toward it. The
further you go down a group, the farther the electron is away from the nucleus and the _____________ electrons an
atom has. It is harder to attract extra electrons if the available energy level is far from the nucleus, so the
electronegativity _____________________. As you go across a row, electronegativity increases as the
________________________ character of the elements decreases.
 Choose the element from the pair with the greater electronegativity.
a) sodium (Na) or rubidium (Rb) _______
b) selenium (Se) or bromine (Br) _______
 Choose the element from the pair with the smaller electronegativity.
a) magnesium (Mg) or calcium (Ca) _______
b) nitrogen (N) or oxygen (O) _______
Homework / Practice
Write the configuration notation for each of the following elements:
1) sodium
2) iron
3) bromine
4) barium
Write the noble gas notation for each of the following elements:
5) cobalt
6) silver
7) tellurium
8) radium
Determine what elements are denoted by the following electron configurations:
9) 1s22s22p63s23p4
11) [Kr] 5s24d105p3
10) 1s22s22p63s23p64s23d104p65s1
12) [Rn] 7s25f11
31
Unit 3- Electrons and Periodicity
Write the orbital notation for the following:
13) carbon
14) neon
15) sulfur
16) P
17) B
18) Na
Write configuration notation for atoms containing the following number of electrons:
19) 3
20) 6
21) 8
22) 13
Draw the Lewis Dot Notation for the following elements
23) Sodium
25) Silver
27) Antimony
24) Sulfur
26) Aluminum
28) Argon
Reference Labeling- Label the following on the blank periodic table on the next page-29)
30)
31)
32)
33)
34)
35)
36)
37)
38)
39)
40)
41)
42)
43)
44)
45)
Metals
Nonmetals
Metalloids
Transition metals
Actinides
Lanthanides
Alkali metals
Alkaline earth metals
Halogens
Noble gases (inert gases)
Group numbers: 1A, 2A, 3A, 4A, 5A, 6A, 7A, 8A
Label the most likely charge of an ion from groups 1A, 2A, 5A, 6A, 7A, and Noble gases
Show which way are periods (left to right or up and down)
Show which way are groups/families (left to right or up and down)
Using color, identify the s, p, d and f blocks on the table
Label the rows with energy level numbers
Using arrows, indicate the trends of increasing atomic radius, ionization energy, and electronegativity
(from left to right and top to bottom)
46) Put a star in the box for each diatomic element
47) Show the charge of the ions that form for silver (+1), zinc (+2), and aluminum (+3)
32
Unit 3- Electrons and Periodicity
33
Unit 3- Electrons and Periodicity
Electrons and Periodicity Practice Test
Directions: For questions 1-4, match each of the following terms with a number or chemical symbol from the
periodic table below.
1. Alkaline earth metals:
3. Noble gases
2.
4. The transition metals
Halogens:
5. Draw the orbital notation for sodium.
6. Given the electron configuration, identify the element 1s2 2s2 2p6 3s2 2p6 3s2 3p6 4s2 3d7
7. Write the complete configuration notation for silver.
8. Write the shorthand method (Noble Gas notation) for antimony.
9. Give the energy level for the valence electrons in helium.
10. Determine the color of light emitted when an electron jumps from the following quantum levels n=4 to n=2.
11. How many valence electrons does carbon have?
12. Draw the Lewis Dot notation for sodium.
13. Describe why the atomic radius of elements increases as you go down a group.
Multiple Choice Practice
14. The two main parts of an atom are the
a.
b.
c.
d.
Principle energy levels and energy sublevels
Nucleus and kernel
Nucleus and energy levels
Planetary electrons and energy levels
15. The sublevel that has only one orbital is identified by the letter
a.
s
b.
p
c.
34
d
d.
f
Unit 3- Electrons and Periodicity
16. The sublevel that can be occupied by a maximum of ten electrons is identified by the letter
a.
s
b.
p
c.
d
c.
d.
3 electrons
0 electrons
d.
f
17. An orbital may never be occupied by
a.
b.
1 electron
2 electrons
18. An atom of beryllium consists of 4 protons, 5 neutrons, 4 electrons. The mass number of this atom is
a.
13
b.
9
c.
8
d.
5
19. Which of the following is the correct electron configuration for the bromide ion, Br1- ?
a.
b.
c.
[Ar] 4s24p5
[Ar] 4s23d104p5
[Ar] 4s23d104p6
d.
e.
[Ar] 4s23d104p65s1
[Ar] 4s23d103p6
20. Which is the first element to have 4d electrons in its electron configuration?
a.
b.
Ca
Sc
c.
d.
Rb
Y
e.
La
21. When electrons in an atom in an excited state fall to lower energy levels, energy is
a.
b.
absorbed, only
released, only
c.
d.
neither released nor absorbed
both released and absorbed
c.
S
22. Which of the following elements has the greatest electronegativity?
a.
Mg
b.
K
d.
F
23. Which of the following elements would react with chlorine in a one to one ratio
a.
Mg
b.
K
c.
S
d.
F
c.
S
d.
F
c.
S
d.
F
c.
S
d.
F
c.
d.
sodium
sulfur
24. Which of the following elements would have the smallest radius
a.
Mg
b.
K
25. Which of following elements has the lowest first ionization energy
a.
Mg
b.
K
26. Which of the following elements is an alkali metal?
a.
Mg
b.
K
27. Which element's ionic radius is smaller than its atomic radius?
a.
b.
neon
nitrogen
28. Which three groups of the Periodic Table contain the most elements classified as metalloids (semimetals)?
a.
b.
1, 2, and 13
2, 13, and 14
c.
d.
35
14, 15, and 16
16, 17, and 18
Unit 3- Electrons and Periodicity
29. Which element has the highest first ionization energy?
a.
b.
sodium
aluminum
c.
d.
calcium
phosphorus
c.
d.
calcium
potassium
c.
d.
Ba, Ag, Sn, Xe
Fr, F, O, Rn
30. Which of the following elements has the smallest atomic radius?
a.
b.
nickel
cobalt
31. Which set of elements contains a metalloid?
a.
b.
K, Mn, As, Ar
Li, Mg, Ca, Kr
32. Atoms of elements in a group on the Periodic Table have similar chemical properties. This similarity is
most closely related to the atoms'
a. number of principal energy levels
b. number of valence electrons
c. atomic numbers
d. atomic masses
33. As atoms of elements in Group 16 are considered in order from top to bottom, the electronegativity of each
successive element
a. decreases
b. increases
c.
remains the same
34. An atom of which of the following elements has the greatest ability to attract electrons?
a.
b.
silicon
sulfur
c.
d.
nitrogen
chlorine
c.
d.
sulfur
silver
35. At STP, which substance is the best conductor of electricity?
a.
b.
nitrogen
neon
36. A strontium atom differs from a strontium ion in that the atom has a greater
a.
b.
number of electrons
number of protons
c.
d.
atomic number
mass number
c.
d.
neon
nitrogen
c.
8
c.
d.
the number of its electrons
its atomic mass.
37. Which gas is monatomic at STP?
a.
b.
chlorine
fluorine
38. How many valence electrons does an oxygen atom have?
a.
2
b.
6
d.
39. The identity of an element is determined by...
a.
b.
the number of its protons.
the number of its neutrons.
40. Which of the following electron configurations represents the electron configuration for a magnesium
cation (Mg2+)?
a. 1s22s22p63s2
b. 1s22s22p63s23p2
c.
d.
36
1s22s22p6
1s22s22p4
16
Unit 3- Electrons and Periodicity
41. Which of the following atoms has the largest diameter?
a.
F
b.
Cl
c.
Br
d.
I
c.
N
d.
O
42. Which of the following elements has the greatest electronegativity?
a.
Si
b.
P
43. Which scientist noted a definite pattern in valence numbers and arranged an early periodic table in order of
the elements atomic mass?
a. Enrico Fermi
b. Dmitri Mendeleev
c.
d.
Albert Einstein
Madame Curie
44. The periodic law states that the properties of elements are periodic functions of their
a.
b.
Mass
Symbol
c.
d.
atomic number
valence
c.
d.
Chlorine
Neon
45. Which of the following is a noble gas?
a.
b.
Sodium
Gold
46. A gas is called "noble" because
a.
b.
c.
d.
it is normally unreactive
it is normally inert
it has a complete outer energy level of electrons
all of the above
47. Of the following elements, the one that forms cations with varying positive charges is:
a.
b.
c.
iron, Fe
sodium, Na
aluminum, Al
d.
e.
strontium, Sr
nitrogen, N
48. Which of the following statements is incorrect?
a.
b.
c.
d.
e.
metals generally form cations
nonmetals are generally poor conductors of electricity
metals are malleable
nonmetals are generally brittle
metalloids are metals with some nonmetallic characteristics
49. Which of the following statements are true?
a.
b.
c.
d.
e.
It is easier to remove an electron from Na than from Na +.
As the atomic number increases within a group of the representative elements, the tendency is for
first ionization energy to increase.
All particles with the electron configuration [Ar]4s2 have the same ionization energy.
The first ionization energy of fluorine is greater than the first ionization energy of oxygen.
All are false
37
Unit 3- Electrons and Periodicity
50. An element having the configuration [Xe]6s1 belongs to the Group:
a.
b.
c.
alkali metals
halogens
alkaline earth metals
d.
e.
None of these
noble gases
51. How many unpaired electrons are there in an atom of tin in its ground state?
a.
b.
4
0
c.
d.
3
2
e.
1
52. Of the following elements, which one is most likely to form an ion through the loss of two electrons?
a.
b.
strontium
chlorine
c.
d.
aluminum
sodium
e.
sulfur
e.
P
53. Which of the following particles has the greatest atomic radius?
a.
b.
Al
Si
c.
d.
S
Al3+
54. Which of the following forms of electromagnetic radiation has the shortest wavelength?
a.
b.
ultraviolet
radio waves
c.
d.
infrared
visible light
e.
microwaves
55. For which of the following transitions does the light emitted have the shortest wavelength?
a.
b.
c.
n = 4 to n = 2
n = 2 to n = 1
n = 5 to n = 3
d.
e.
n = 4 to n = 3
n = 3 to n = 2
56. Researchers at Lawrence Berkeley National Lab have recently formed a new synthetic element with atomic
number 118 and mass number 293. Which of the following elements would have chemical properties most
similar to this new element?
a. Ir
c. Ta
e. S
b. Xe
d. Pb
38
Unit 4 – Types of Bonding
BONDING
As atoms bond with each other, they _____________________ their potential energy, thus creating more stable
arrangements of matter. The force that holds two ________________ together is called a chemical bond. There are
3 types of bonding: ionic, ___________________, and metallic. The number of valence electrons are easily found
by looking up the group number on the periodic table.
Group 1A (Li, Na, K, etc.): 1 valence electron
Group 2A (Be, Mg, Ca, etc.): ______ valence electrons
Group 3A (B, Al, Ga, etc.): 3 valence electrons
Group 4A (C, Si, Ge, etc.): 4 valence electrons
Group 5A (N, P, As, etc.): _______ valence electrons
Group 6A (O, S, Se, etc.): 6 valence electrons
Group 7A (F, Cl, Br, etc.): 7 valence electrons
Group 8A (He, Ne, Ar, etc.): 8 valence electrons (except He has _______ valence electrons)
Recall the rules for drawing Lewis dot diagrams: Write the _____________________ of the element. Put one dot
for each valence electron. Don’t _____________ electrons up until you have to.
Electron Configurations and Electron Dot Diagrams for Cations
Metals lose electrons to attain noble gas configuration. They make positive ions, ____________.
If we look at an electron configuration, it makes sense. Example: Sodium (Na), 1s22s22p63s1, has _________
valence electron(s). The electron that is removed comes from the ____________ energy level. As a result of the
loss of the electron, the sodium ion (Na+) has the following electron configuration: 1s22s22p6
Calcium has 2 valence electrons. These will come off, forming a positive ion.
Electron Configurations and Electron Dot Diagrams for Anions
Nonmetals gain electrons to attain noble gas configuration. This means they want a(n) ________________ of
electrons, 8 electrons. They make negative ions, ___________________.
If we look at an electron configuration, it makes sense. Example: Sulfur (S), 1s22s22p63s23p4, has _______ valence
electrons and needs to gain 2 more to have an octet.
The sulfur ion (S-2) has the same electron configuration as a noble gas: 1s22s22p63s23p6 Phosphorous has 5 valence
electrons. It will gain _________ electrons to fill the outer shell.
Stable Electron Configurations
All atoms react to achieve __________________ gas configuration. Noble gases, except He, have 2 s electrons and
6 p electrons, totaling 8 valence electrons. They obey the ____________________ rule.
39
Unit 4 – Types of Bonding
IONIC BONDING
Anions and cations are involved in ionic bonding and are held together by __________________ charges,
electrostatic attraction. The bond is formed through the ______________________ of electrons. Electrons are
transferred to achieve noble gas configuration. Ionic bonds occur between _________________ and nonmetals. All
the electrons must be accounted for! A compound that is composed of _______________ is called an ionic
compound. Note that only the arrangement of electrons has changed. Nothing about the atom’s nucleus has
changed. Ionic compounds have a _______________________ structure, a regular repeating arrangement of ions in
the solid. Even though the ions are ___________________ bonded to one another, ionic compounds are
__________________. Strong repulsion breaks crystal apart. The structure is rigid. They have _______________
melting points because of strong forces between ions. They also conduct electricity in the _________________ and
dissolved states. Any compound that conducts electricity when melted or dissolved in water is a(n)
___________________________.

How many valence electrons must an atom have in its outer energy level in order to be considered stable?
The energy required to separate one mole of the ions of an ionic compound is called ____________________
energy, which is expressed as a negative quantity. The greater (that is, the more negative) the lattice energy is, the
______________________ the force of attraction between the ions. Lattice energy tends to be
__________________________ for more-highly-charged ions (those atoms that have more electrons to give or
those atoms that can take more electrons). Lattice energy also tends to be greater for __________________ ions.

Between the following ionic compounds, which would be expected to have the higher (more negative)
lattice energy?

LiF or KBr
Between the following ionic compounds, which would be expected to have the higher (more negative)
lattice energy?
NaCl or MgS
The electronegativity difference for two elements in an ionic compound is greater than or equal to
_______________.
COVALENT BONDING
A _______________________ is an uncharged group of two or more atoms held together by covalent bonds.
Covalent compounds occur between two ___________________ or a nonmetal and hydrogen. The attraction of two
atoms for a shared _______________ of electrons is called a covalent bond. In a covalent bond, atoms share
electrons and neither atom has an ionic ______________________. Covalent bonds occur between 2
___________________________ because nonmetals hold onto their valence electrons. They can’t give away
electrons to bond, yet, they still want _______________ gas configuration. They get it by sharing valence electrons
with each other. By sharing both atoms get to count the electrons toward noble gas configuration. A
____________________ bond is formed from the sharing of two valence electrons. The electronegativity difference
for two elements in a covalent compound is between _________ and 1.7.

Do atoms that share a covalent bond have an ionic charge?
40
Unit 4 – Types of Bonding
Sometimes atoms share more than one pair of valence electrons. A ____________________ bond is when atoms
share two pair of electrons, 4 electrons. A triple bond is when atoms share three pair of electrons, _____ electrons.
Triple bonds are ________________________ and shorter than double bonds. Double bonds are stronger and
shorter than ______________________ bonds.
THE WETTER WAY
You can easily determine the number of bonds in a compound by performing the “Wetter Way.”
# bonds = (Σe- after – Σe- before ) / 2
The number of electrons before bonding is equal to the __________________________ number. To get the number
of electrons after bonding, double the number of electrons before bonding, BUT DO NOT EXCEED ___________!
Example: CO2
C is in column 4A and therefore has 4 valence electrons before bonding. O is in column 6A and
therefore has 6 valence electrons before bonding.
Σ electrons before bonding
Carbon = 4
Two oxygens = 6 x 2 = 12
(The formula CO 2 implies there are 2 oxygen atoms.)
Σ electrons before bonding = 4 + 12 = 16
Σ electrons after bonding
C has 4 valence e- before bonding, so it has (4x2) = 8 electrons after bonding
O has 6 valence e- before bonding, so it has (6x2) = 12 electrons after bonding
Remember 8 is the maximum electrons after bonding, so oxygen can only have 8 electrons even though we
calculated 12.
Carbon = 8
Two oxygens = 8 x 2 = 16
Σ electrons after bonding = 8 + 16 = 24
# bonds = (Σe- after – Σe- before) / 2 = (24 - 16) / 2 = 4 bonds
The element you have only 1 of goes in the center. The other elements surround it. Connect the elements with a
single line (a single bond). You have only used 2 of your calculated 4 bonds, so you need to double up. A line
represents 2 electrons. Count your lines for each element to determine if extra electrons need to be added. Carbon
has 4 lines attached which represent 8 electrons. No extra electrons are needed around carbon. Each oxygen has
2 lines attached which represent 4 electrons. Oxygen needs 8 electrons after bonding, so each oxygen needs
4 electrons.

Determine the number of bonds and draw the dot-dash diagram for HBr.

Determine the number of bonds and draw the dot-dash diagram for N2.
41
Unit 4 – Types of Bonding
METALLIC BONDING
The bonding in metals is explained by the _______________________ ____________ model, which proposes that
the atoms in a metallic solid contribute their valence electrons to form a “sea” of electrons that surrounds metallic
__________________. These delocalized electrons are not held by any specific atom and can ________________
easily throughout the solid. Metal atoms release their valence electrons into a sea of electrons shared by all of the
metal atoms. The bond that results from this _____________________ pool of valence electrons is called a metallic
bond. Metals hold onto their valence electrons very _______________________. Think of them as positive ions
floating in a sea of electrons. Because electrons are free to move through the solid, metals conduct
_______________________. Metals generally have extremely ______________ melting points because it is
difficult to pull metal atoms completely away from the group of cations and attracting electrons. Metals are
________________________ (able to be hammered into sheets). Metals are also ________________________
(able to be drawn into wire) because of the mobility of the particles. Electrons allow atoms to slide by. A mixture
of elements that has metallic properties is called a(n) _____________________.
Homework / Practice
Complete the table by identifying the charge of each of the elements listed and then indicating the formula for ionic
compounds formed between the two substances
Charge
Charge
Na
O
N
P
S
Cl
F
21+
Na2O
Mg
Ca
Al
Li
Zn
What type of bonding is present in the following compounds:
1) SbBr3
2) Ag
3) MgBr2
4) ClO2
5) KCl
6) Fe
7) PbO
8) FeCl2
15) N2O3
16) Mg3P2
17) CCl4
18) H2O
What type of bonding is present in the following compounds:
9) NI3
10) CO2
11) Ni
12) Au
13) LiF
14) Al2O3
POLARITY and VSEPR
How each atom fares in a tug-of-war for shared electrons is determined by comparing the
_________________________________ of the two bonded atoms. Recall that electronegativity is the measure of
42
Unit 4 – Types of Bonding
the ability of an atom in a bond to ________________________ electrons. Atoms with large electronegativity
values, such as fluorine, attract shared valence electrons more __________________ than atoms such as sodium that
have small electronegativities. Electronegativity is a periodic property. With only a few exceptions,
electronegativity values_____________________ as you move from left to right in any period of the periodic table.
Within any group, electronegativity values decrease as you go ___________________ the group. Fluorine has the
highest value of ____________. The greater the difference between the electronegativities of the bonding atoms,
the more _____________________________ the electrons are shared and the more polar the bond.
If the electronegativity difference between the two elements in question is:
between 0.0 – 1.7, the bond is ______________________
greater than 1.7, the bond is _____________________
When the electronegativity difference in a bond is 1.7 or greater, the sharing of electrons is so unequal that you can
assume that the electron on the less electronegative atom is ________________________ to the more
electronegative atom. For example, ∆EN for cesium and fluorine is 4.0 − 0.7 = 3.3. Therefore the bond is
_________________.
COVALENT BONDS AND POLARITY
When the atoms in a bond are the same, the electrons are shared ________________________. This results in a
_______________________ covalent bond. _________________________ elements (H 2, O2, N2, Cl2, Br2, I2, and
F2) have pure nonpolar covalent bonds. All other covalent bonds are polar. The electron sharing is not equal, but it
is not so unequal that a complete _____________________ of electrons takes place.
Consider hydrogen and chlorine. Hydrogen has an electronegativity of 2.20, and chlorine has an electronegativity of
3.16. The ________ pulls harder on the electrons because its electronegativity is greater. The electrons spend more
time near the Cl. These symbols, __________________ plus (δ+) and delta minus (δ-), represent a partial positive
charge and a partial negative charge.
Polar molecules are molecules with a positive and a negative ______________. This requires two things to be true:
The molecule must contain _______________ bonds. (This can be determined from differences in
electronegativity.) Symmetry cannot ______________________ out the effects of the polar bonds. (Must
determine geometry first.)

In the following compounds, determine whether the molecule is polar or nonpolar
a.
hydrogen fluoride (HF)
d.
ammonia (NH3)
b.
water (H2O)
e.
carbon dioxide (CO2)
c.
carbon tetrachloride (CCl4)
VSEPR
VSEPR stands for Valence Shell ______________________ ________________ Repulsion. It predicts threedimensional geometry of molecules. The valence shell includes the ______________________ electrons. The
electron pairs try to get as far away as possible to _______________________ repulsion. You can determine the
43
Unit 4 – Types of Bonding
angles of the bonds. VSEPR is based on the number of pairs of valence electrons, both bonded and unbonded. An
unbonded pair of electrons is referred to as a _______________pair. Use the Wetter Way to calculate the number of
bonds and then draw the dot-dash diagram. The shape of the molecule and bond angle can be determined from this
diagram.
LINEAR
Consider the simplest molecule that exists—hydrogen, H2. Determine the number of bonds using the Wetter Way.
H is in column 1A and therefore has 1 valence electron before bonding.
Σ electrons before bonding
Two hydrogens = 1 x 2 = 2
H – 1 valence e- before bonding, so it has (1x2) = 2 electrons after bonding
Σ electrons after bonding
Two hydrogens = 2 x 2 = 4
(4 – 2) / 2 = 1, so H2 has 1 bond!
Connect the elements with a single bond. You have used all of the calculated bonds.
Each hydrogen has 1 line attached which represents _______________ electrons. No extra electrons are needed
around hydrogen to have the 2 electrons needed after bonding.
A hydrogen molecule is linear. The electrons attempt to maximize their distance from one another by having bond
angle of ____________. Linear compounds are NOT ____________________.
TETRAHEDRAL
Consider CH4. The Wetter Way shows that CH4 has __________ bonds! The element you have only one of goes in
the _______________________. The other elements surround it. Connect the elements with a single
_________________ (a single bond). You have used all 4 of your calculated bonds. Remember a line represents
___________ electrons. Count your lines for each element to determine if extra electrons need to be added. Carbon
has 4 lines attached which represent _________ electrons. No extra electrons are needed around carbon. Each
hydrogen atom has one line attached which represents 2 electrons. No extra electrons are needed around hydrogen.
Single bonds fill all atoms. There are _________ bond pairs of electrons pushing away. The electrons can
_________________ their distance from one another by forming a 3-D shape. The furthest they can get away is
___________. This basic shape is a tetrahedral, a pyramid with a triangular base. The tetrahedral is the shape for
everything with 4 bond pairs and ____________ lone pairs around the central atom.
TRIGONAL PYRAMIDAL
Perform the Wetter Way for phosphorous trichloride (PCl3). How many bonds are in this molecule? ________ .
Sketch the dot-dash diagram for phosphorous trichloride. Please include all electrons. Only the electrons around the
______________________ atom affect the shape. The shape is a basic _______________________________ but
44
Unit 4 – Types of Bonding
you can’t see the lone pair. The shape is called trigonal pyramidal. The bond angle is ____________ between the
chlorines because the electron pair forces the chlorines closer to each other.
BENT
Perform the Wetter Way for water (H2O). How many bonds are in this molecule? _____ Sketch the dot-dash
diagram for water. Please include all electrons. Only the electrons around the central atom affect the shape. The
shape is still basic tetrahedral, but you can’t see the _________ lone pairs. The shape is called bent. The bond angle
between hydrogens is ____________.
TRIGONAL PLANAR
Perform the Wetter Way for H2CO. How many bonds are in this molecule? _____. Sketch the dot-dash diagram
for H2CO. Please include all electrons. (Carbon is the central atom.) The farthest you can get the elements apart is
__________. The shape is flat and called trigonal planar.

Determine the number of bonds, draw the dot-dash diagram, state the VSEPR shape and provide the bond
angle for the following compounds
a.
CO2
b.
BCl3
c.
SCl2
d.
SiF4
HYBRID ORBITALS
Hybrid orbitals combine bonding with geometry. ______ hybridization has a linear shape. sp 2 hybridization has a
trigonal ____________________ shape. ____________ hybridization has tetrahedral, trigonal pyramidal and bent
shapes.
Homework/Practice
Draw the Lewis structure for each of the following compounds, identify the shape of the molecule, and identify the
polarity of the molecule.
1.
CCl4
4.
SiO2
2.
BF3
5.
H2S
3.
NF3
INTERMOLECULAR FORCES
Intermolecular forces are forces of _______________________. They are what make solid and liquid molecular
compounds possible. The three intermolecular forces are _________________ bonds, dipole–dipole forces and
London ____________________________ forces.
45
Unit 4 – Types of Bonding
Hydrogen Bonding
A hydrogen bond is a _________________________________________ attraction that occurs between molecules
containing a hydrogen atom bonded to a small, highly electronegative atom with at least ____________ lone
electron pair. For a hydrogen bond to form, hydrogen must be bonded to a fluorine,
__________________________, or nitrogen atom. F, O, and N are very electronegative so it is a very
_______________________ dipole. Hydrogen bonding is the _________________________ of the intermolecular
forces. Examples include H2O, NH3, and HF.
Dipole-dipole Forces
Polar molecules contain ___________________________ dipoles; that is, some regions of a polar molecule are
always ___________________________ negative and some regions of the molecule are always partially positive.
Attractions between _____________________________ charged regions of polar molecules are called dipole–
dipole forces. Neighboring polar molecules orient themselves so that oppositely charged regions _______________
up. Opposites attract but are not completely hooked as in ionic solids. Dipole-dipole forces depend on the number
of _______________________. Bigger molecules result in more electrons, and more electrons mean
________________________ forces. Dipole–dipole forces are stronger than dispersion forces as long as the
molecules being compared have approximately the same mass. Examples of compounds that exhibit dipole-dipole
forces include CO, HCl, and PH3.
London Dispersion Forces
Dispersion forces are ____________________ forces that result from temporary shifts in the
______________________ of electrons in electron clouds. Remember that the electrons in an electron cloud are in
constant _____________________. When two nonpolar molecules are in close contact, especially when they
collide, the electron cloud of one molecule _______________________ the electron cloud of the other molecule.
The electron density around each nucleus is, for a moment, greater in one region of each cloud. Each molecule
forms a __________________________ dipole. When temporary dipoles are close together, a weak dispersion
force exists between oppositely charged regions of the dipoles. Due to the temporary nature of the dipoles,
dispersion forces are the __________________________ intermolecular force. Dispersion forces exist between
____________ gases and compounds that are nonpolar. Examples include Ar, Cl 2, Br2, CH4, and CO2. Dispersion
forces ______________________ as the mass of the molecule increases. C2H6 (MW = 30.0 g/mol) has stronger
dispersion forces than CH4 (MW = 16.0 g/mol). This difference in dispersion forces explains why fluorine and
chlorine are gases, bromine is a __________________________, and iodine is a solid at room temperature. The
molecular mass of iodine is greater than that of bromine, and bromine has a greater mass than chlorine.
Intermolecular Forces
To determine what type of intermolecular force a compound has, ask yourself the following questions.

Does the compound contain hydrogen attached to N, O, or F?
46
Unit 4 – Types of Bonding
o
If yes, the force is hydrogen bonding.
Determine the number of bonds from the Wetter Way and draw the dash-dot diagram.

Does the central element of the compound contain any lone pairs of electrons?
o

If yes, the force is dipole-dipole.
Does the central element of the compound contain ZERO lone pairs of electrons?
o
If yes, the force is dispersion.
Determine the type of intermolecular force in each of the following compounds
1) BCl3 _____________________________
2) Xe _____________________________
3) NH3 _____________________________
4) CH4 _____________________________
5) SO2 _____________________________
6) H2 _____________________________
7) SO3 _____________________________
8) CH3Cl ___________________________
9) HF _____________________________
10) HBr ____________________________
Types of Bonding Practice Test
1.
In a complete sentence, compare and contrast metallic bonds and ionic bonds.
Directions- For each of the following pairs of elements, write the formula for the ionic compound that would form
between them
2. K and Cl
5. Calcium and Chlorine
3. Na and N
6. Zinc and Sulfur
4. Al and O
7. Lithium and Phosphorous
Directions- Draw the Lewis structure , Identify the shape of the molecule, Identify the polarity of the bonds,
Identify the polarity of the molecule, Identify the IMF that would be exhibited
8.
9.
CCl4
SF2
10. SiO2
11. BI3
12. PCl3
13. N2
14. What does IMF stand for? Which of the three IMF’s is the weakest?
Multiple Choice Practice
15. What type of bond exists between atoms of potassium and chloride in a crystal of potassium chloride?
a. Hydrogen bond
d. Nonpolar covalent bond
b. Ionic bond
e. Metallic bond
c. Polar covalent bond
16. What type of bond exists between atoms in a nitrogen molecule?
a. Hydrogen bond
b. Ionic bond
c. Polar covalent bond
d.
e.
Nonpolar covalent bond
Metallic bond
17. What type of bond exists between atoms of calcium in a crystal of calcium?
a. Hydrogen bond
d. Nonpolar covalent bond
b. Ionic bond
e. Metallic bond
c. Polar covalent bond
47
Unit 4 – Types of Bonding
18. All of the following have covalent bonds except
a. HCl
b. CCl4
c.
d.
H2O
CsF
e.
19. Which of the following atoms normally forms monatomic molecules?
a. Cl
c. O
b. H
d. N
CO2
e.
He
20. The complete loss of an electron of one atom to another atom with the consequent formation of electrostatic
charges is said to be
a. Covalent bonding
c. Ionic bonding
b. Polar covalent bonding
d. Coordinate covalent bonding
21. When a metal atom combines with a nonmetal atom, the nonmetal atom will
a. lose electrons and decrease in size
b. lose electrons and increase in size
c. gain electrons and decrease in size
d. gain electrons and increase in size
22. Which formula represents a molecular substance?
a. CaO
b. CO
c.
d.
Li2O
Al2O3
23. Which combination of atoms can form a polar covalent bond?
a. H and H
b. H and Br
c.
d.
N and N
Na and Br
24. Fluorine atoms tend to.______.when they form chemical compounds with metals.
a. lose electrons
b. gain electrons
c. neither lose nor gain electrons...they usually share electrons equally with metals.
d. Fluorine atoms do not form compounds with other atoms...fluorine is an inert gas.
25. What is a compound composed of?
a. two or more different elements that are physically combined in a fixed proportion
b. two or more different mixtures that are physically combined in a fixed proportion
c. two or more different elements that are chemically combined in a fixed proportion
d. two or more different elements that are chemically combined in a variable proportion
26. Which of the following compounds is most likely to be ionic?
a. CO2
c. FeCl3
b. CCl4
d. MgCl2
e.
HBr
27. How many unshared electron pairs must be included in the Lewis structure for water, H 2O?
a. 3
c. 1
b. 2
d. 4
e.
0
48
Unit 4 – Types of Bonding
28. How can a chemical compound be broken?
a. can be broken down by physical means
b. can be broken down by chemical means
c. cannot be broken down
d. can be broken down by physical or chemical means
29. Nitrogen triiodide, NI3, is an unstable molecule that is used as a contact explosive. Its molecular structure is:
a. none of these
d. tetrahedral
b. octahedral
e. pyramidal
c. square planar
30. In which of the following compounds does the bond between the central atom and chlorine have the greatest
ionic character?
a. BCl3
c. CCl4
e. CaCl2
b. FeCl2
d. HCl
31. Which of the following molecules must contain at least one double bond
a. H2O
b. CCl4
c. H2O2
d.
e.
CH3I
CH3COOH
32. The Lewis structure for hydrogen cyanide is:
c.
a.
d.
b.
e.
33. In the Lewis structure for CH2Cl2, the number of unshared electron pairs is:
a. 10
c. 2
b. 8
d. 4
34. The only intermolecular forces existing between oxygen molecules are:
a. ion-ion attractive forces
b. hydrogen bonding forces
c. permanent dipole forces
e.
d.
e.
6
nuclear forces
London dispersion forces
35. Reactions between alkali metals and phosphorous result in compounds with the formula:
a. M3P
d. M2P3
b. None of these
e. MP3
c. M2P
36. A particle X contains 10 electrons, seven neutrons and has a net charge of 3-. The particle is:
a. a nitride ion
d. a neon ion
b. obviously polyatomic
e. none of these are correct
c. an oxide ion
49
Unit 4 – Types of Bonding
37. Which of the following arrangements represent ions?
i.
ii.
iii.
iv.
v.
a.
b.
c.
12 protons, 12 neutrons, 11 electrons
12 protons, 11 neutrons, 12 electrons
11 protons, 12 neutrons, 12 electrons
11 protons, 12 neutrons, 11 electrons
12 protons, 12 neutrons, 12 electrons
1 and 2
1, 3, and 4
1 and 3
d.
e.
2 and 3
all of these are ions
d.
e.
1 and 3 are correct
2 and 4 are correct
38. Ions are formed in chemical reactions by:
i. Gaining electrons
ii. Losing electrons
iii. Gaining protons
iv. Losing protons
v. All of these
a.
b.
c.
1 and 2 are correct
3 and 4 are correct
5 is correct
50
Unit 5 – Nomenclature
NAMING COMPOUNDS AND WRITING FORMULAS
Recall that the periodic table is more than a list of elements. Elements are put in columns because of
_____________________ properties. Each column is called a group. A compound is made of two or more
______________________. The name should tell us how many and what type of atoms. There are two types of
compounds: ___________________ compounds and molecular compounds. The simplest ratio of the ions
represented in an ionic compound is called a ______________________ unit. The overall charge of any formula
unit is ________________. In order to write a correct formula unit, one must know the charge of each ion. Atoms
are electrically _____________________. They have the same number of protons and electrons.
________________ are atoms, or groups of atoms, with a charge. Ions have a different numbers of electrons. An
anion is a _____________________ ion. An anion has gained electrons. Nonmetals can ________________
electrons. The charge is written as a superscript on the right. F 1- has gained _________ electron. O2- has gained
__________ electrons. A ___________________ is a positive ion. It is formed by __________________ electrons.
There are more _____________________ than electrons. ______________________ form cations. K 1+ has lost
one electron. Ca2+ has lost __________ electrons. The charges of monatomic ions, or ions containing only one
atom, can often be determined by referring to the periodic table or table of common ions based on group number.
The charge of a monatomic ion is equal to its _________________________ number. For most of the Group
________ elements, the Periodic Table can tell what kind of ion they will form from their location. Elements in the
same group have similar properties, including the charge when they are ions.
NAMING CATIONS
We will use the systematic way. For cations, if the charge is always the same (Group A) just write the
_________________ of the metal. Transition metals (as well as tin and lead) can have more than one type of
charge. The charge is indicated with ___________________ numerals in parenthesis. Zinc (Zn2+) and silver (Ag1+),
although transition metals, only have __________ possible charge. Roman numerals ARE NOT used for zinc and
silver. Li1+ is called the ____________________ ion. __________ is called the Strontium ion. Fe2+ is called the
iron (II) ion. Iron is a transition metal, so the charge is not always the same. The name of the metal is written, and
the charge is denoted with Roman numerals in parenthesis. Pb 2+ is called the lead __________ ion.
Name the following cations.
a) Ca2+ _________________________
b) Al3+ ___________________________
c) Sn4+ _________________________
d) Na+ _________________________
e) Fe3+ _________________________
f) Cu+ _________________________
WRITING FORMULAS FOR CATIONS
Write the formula for the metal. If a Roman numeral is in parenthesis use that number for the
_____________________. Indicate the charge with a superscript. If no Roman numeral is given, find the Group A
51
Unit 5 – Nomenclature
metal on the periodic table and determine the charge from the _____________________ number. The formula for
the nickel (II) ion is Ni2+. The formula for the gallium ion is ____________.
Write the formulas for the following cations.
a) magnesium ion ________________
b) copper (II) ion ___________________
c) potassium ion ________________
d) silver ion _________________
e) chromium (VI) ion ________________
f) mercury (II) ion ________________
NAMING ANIONS
Naming monatomic anions is always the same. Change the element ending to – ___________. F is the symbol for
fluorine, F1- is fluoride. Cl1- is called the chloride ion. _______ is called the oxide ion.
Name the following anions.
a) S2- _________________________
b) Br1- ___________________________
c) N3- _________________________
d) As3- _________________________
e) Te2- _________________________
WRITING FORMULAS FOR ANIONS
Write the formula for the nonmetal. Find the Group A nonmetal on the periodic table and determine the charge from
the column number.
Write the formulas for the following anions.
a) iodide ion ________________
b) phosphide ion ___________________
c) selenide ion ________________
d) carbide ion _________________
IONIC COMPOUNDS
Oxidation numbers can be used to determine the chemical formulas for ionic compounds. If the oxidation number
of each ion is _________________________ by the number of that ion present in a formula unit, and then the
results are added, the sum must be _______________. In the formula for an ionic compound, the symbol of the
_________________ is written before that of the anion. Subscripts, or small numbers written to the lower
______________________ of the chemical symbols, show the numbers of ions of each type present in a formula
unit.
BINARY IONIC COMPOUNDS
Binary ionic compounds are composed of a metal bonded with a ________________________. Name the metal ion
using a Roman numeral in parenthesis if necessary. Follow this name with the name of the nonmetal ion.
Name the following binary ionic compounds.
a) NaCl __________________
b) Ca3P2 __________________
c) CuO __________________
d) SnBr2 __________________
e) Fe2S3 __________________
f) AlF3 __________________
52
Unit 5 – Nomenclature
g) KCl __________________
h) Na3N __________________
i) CrN __________________
j) PbO2 __________________
Write the symbol for the metal. Determine the oxidation number from either the column number or the Roman
numeral and write it as a superscript to the right of the metal’s symbol. To the right of the metal’s symbol, write the
symbol for the nonmetal. Determine the oxidation number from the column number and write it as a superscript to
the right of the nonmetal’s symbol.

Example: potassium fluoride - K1+ F1- If the two oxidation numbers add together to get zero, the formula
is a one-to-one ratio of the elements. Answer = KF

Example: aluminum sulfide - Al3+ S2- If the two oxidation numbers DO NOT add together to get zero, you
will need to “criss-cross” the superscripts. These numbers now become subscripts. Omit all positive and
negative signs and omit all 1’s. Answer = Al2S3
Write the formulas for the following binary ionic compounds.
a) lithium selenide __________________
b) tin (II) oxide __________________
c) tin (IV) oxide __________________
d) magnesium fluoride ________________
e) copper (II) sulfide __________________
f) iron (II) phosphide _________________
g) gallium nitride __________________
h) iron (III) sulfide __________________
TERNARY IONIC COMPOUNDS
Ternary ionic compounds are composed of at least _________ elements. Name the metal ion, using a Roman
numeral in parenthesis if necessary. Follow this name with the name of the polyatomic ion. Polyatomic ions are
groups of atoms that stay together and have a __________________. Examples are provided on page 7 of the
NCDPI Reference Tables for Chemistry. There is one polyatomic ion with a positive oxidation number (NH 4+) that
may come first in a compound. Name the ion. Follow this name with the name of the anion or second polyatomic
ion. Certain polyatomic ions, called ________________________, contain oxygen and another element.
Name the following ternary ionic compounds.
a) LiCN __________________
b) Fe(OH)3 ___________________
c) (NH4)2CO3 __________________
d) NiPO4 __________________
e) NaNO3 __________________
f) CaSO4 __________________
g) (NH4)2O __________________
h) CuSO3 __________________
Write the symbol for the metal or ammonium ion. Write the oxidation number as a superscript to the right of the
metal’s/ammonium ion’s symbol. To the right of the metal’s symbol, write the symbol for the nonmetal or
polyatomic ion. Write the oxidation number as a superscript to the right of the nonmetal’s/polyatomic ion’s symbol.

Example: potassium nitrate - K1+ NO31- If the two oxidation numbers add together to get zero, the formula is a
one-to-one ratio of the elements. Answer = KNO3
53
Unit 5 – Nomenclature

Example: aluminum hydrogen sulfate – Al3+ HSO4 1- If the two oxidation numbers DO NOT add together to
get zero, you will need to “criss-cross” the superscripts. These numbers now become subscripts. Parentheses
are to be placed around polyatomic ions before criss-crossing. Omit all positive and negative signs and omit all
1’s. Answer = Al(HSO4)3
Write the formulas for the following ternary ionic compounds.
a) ammonium chloride __________________
b) ammonium sulfide _________________
c) barium nitrate __________________
d) zinc iodate __________________
e) sodium hypochlorite __________________
f) chromium (III) acetate ______________
g) iron (II) dichromate __________________
h) mercury (I) bromate ________________
MOLECULAR COMPOUNDS
Molecular compounds are made of molecules. They are made by joining _______________________ atoms
together into molecules. A molecular compound’s name tells you the number of atoms through the use of
____________________________.
1 mono-
4
tetra-
7
hepta-
2 di-
5
penta-
8
octa-
3 tri-
6
hexa-
9
nona-
The name will consist of two words.
Prefix name
prefix name –ide
10
deca-
One exception is we don’t write mono- if
there is only one of the first element. The following double vowels cannot be used when writing names: (oa) and
(oo).

Example: NO2
There is one nitrogen. Mononitrogen But, you cannot use mono- on the first element, so drop
the prefix. There are two oxygens. dioxygen You need the suffix –ide. dioxide (Answer: nitrogen dioxide).

Example: N2O
There are two nitrogens. Dinitrogen There is one oxygen. monooxygen You cannot run (oo)
together, so monoxygen. You need the suffix –ide. monoxide (Answer: dinitrogen monoxide).
Name the following molecular compounds.
a) Cl2O7 ____________________________
b) CBr4 ____________________________
c) CO2 ________________________
d) BCl3 ___________________________
When writing a formula of a molecular compound from the name, you will not need to criss-cross oxidation
numbers. Molecular compounds name tells you the number of atoms through the use of prefixes.

Example: diphosphorus pentoxide
The name implies there are 2 phosphorous atoms and 5 oxygens.
Answer: P2O5

Example: sulfur hexafluoride The name implies there is 1 sulfur atom and 6 fluorines. Answer: SF 6
54
Unit 5 – Nomenclature
Write the formulas for the following molecules.
a) tetraiodide nonoxide __________________
b) nitrogen trioxide __________________
c) carbon tetrahydride __________________
d) phosphorus trifluoride ______________
IONIC
MOLECULAR
Smallest Piece
Molecule
Types of Elements
metal and nonmetal
State of Matter
solid
Melting Point
Low <300°C
ACIDS
Acids are compounds that give off hydrogen ions (H +) when dissolved in water. Acids will always contain one or
more hydrogen ions next to an anion. The anion determines the name of the acid.
Binary Acids
Binary acids contain hydrogen and an anion whose name ends in –ide. When naming the acid, put the prefix hydroand change -ide to -ic acid.

Example: HCl
The acid contains the hydrogen ion and chloride ion. Begin with the prefix hydro-, name
the nonmetallic ion and change -ide to -ic acid. Answer: hydrochloric acid

Example: H2S
The next step is change -ide to -ic acid, but for sulfur the “ur” is added before -ic.
Answer: hydrosulfuric acid
Name the following binary acids.
a) HF ____________________________
b) H3P ____________________________
The prefix hydro- lets you know the acid is binary. Determine whether you need to criss-cross the oxidation
numbers of hydrogen and the nonmetal.

Example: hydrobromic acid
The acid contains the hydrogen ion and the bromide ion. H1+ Br1- The two
oxidation numbers add together to get zero. Answer: HBr

Example: hydrotelluric acid
The acid contains the hydrogen ion and the telluride ion. H1+ Te2- The two
oxidation numbers do NOT add together to get zero, so you must criss-cross. Answer: H2Te
Write the formulas for the following binary acids.
a) hydrocyanic acid __________________
b) hydroselenic acid __________________
Ternary Acids
The acid is a ternary acid if the anion has oxygen in it. The anion ends in -ate or -ite. Change the suffix -ate to -ic
acid. Change the suffix -ite to -ous acid. The hydro- prefix is NOT used!
55
Unit 5 – Nomenclature

Example: HNO2 The acid contains the hydrogen ion and nitrite ion. Name the polyatomic ion and change
-ite to -ous acid. Answer: nitrous acid

Example: H3PO4
The acid contains the hydrogen ion and phosphate ion. Name the polyatomic ion and
change -ate to -ic acid. Answer: phosphoric acid
Name the following ternary acids.
a) H2CO3 ____________________________
b) H2SO4 __________________________
c) H2CrO4 ________________________
d) HClO2 __________________________
The lack of the prefix hydro- from the name implies the acid is ternary, made of the hydrogen ion and a polyatomic
ion. Determine whether you need to criss-cross the oxidation numbers of hydrogen and the polyatomic ion.

Example: acetic acid
The polyatomic ion must end in –ate since the acid ends in -ic. The acid is made of
+
H and the acetate ion.

Example: sulfurous acid
H1+ C2H3O21- The two charges when added equal zero. Answer: HC2H3O2
Again the lack of the prefix hydro- implies the acid is ternary, made of the
hydrogen ion and a polyatomic ion. The polyatomic ion must end in –ite since the acid ends in -ous. The
acid is made of H+ and the sulfite ion.
H1+ SO32-
The two charges when added do not equal zero, so
you must crisscross the oxidation numbers. Ignore the negative sign and ones are understood. Answer:
H2SO3
Write the formulas for the following binary acids.
a) perchloric acid __________________
b) iodic acid __________________
c) dichromic acid __________________
d) hypochlorous acid ________________
Homework / Practice
Name each of the following compounds
1.
CuS
15. Mg(OH)2
29. NI3
43. SnSe2
2.
FeI2
16. Na2CrO4
30. CO2
44. GaAs
3.
Cu2SO4
17. Ba(CN)2
31. Ni
45. Pb(SO4)2
4.
CuCl2
18. K2SO4
32. Au
46. Be(HCO3)2
5.
Ni(C2H3O2)2
19. NH4NO3
33. LiF
47. Mn2(SO3)3
6.
MnO
20. FePO4
34. Al2O3
48. Al(CN)3
7.
CoO
21. SbBr3
35. N2O3
49. P4O6
8.
Mn2O3
22. Ag
36. Mg3P2
50. N2O3
9.
Co2S3
23.
37. CCl4
51. N2O
MgBr2
10. AuF
24. ClO2
38. H2O
52. BrO3
11. CrBr2
25. KCl
39. NH4Cl
53. SiF4
12. AlPO4
26. Fe
40. Fe(NO3)3
54. P4S10
13. KNO2
27. PbO
41. TiBr3
55. Cl2O3
14. CaCO3
28. FeCl2
42. Cu3P
56. PCl3
56
Unit 5 – Nomenclature
57. PF5
60. CCl4
63. N2O4
58. P2O5
61. N2O5
64. NI3
59. Cl2O
62. As4O6
65. ClF3
66. I2O5
Write the formula for each of the following compounds
67. iron (II) chloride
83. cobalt (III) nitrite
99. silver dichromate
68. copper (I) sulfide
84. ammonium sulfite
100. barium Iodide
69. mercury (I) bromide
85. dihydrogen oxide
101. carbon tetrafluoride
70. chromium (III) oxide
86. lithium Iodide
102. sodium sulfide
71. gold (I) iodide
87. silver chloride
103. chromium (VI) phosphate
72. manganese (II) nitride
88. carbon monoxide
104. vanadium (IV) carbonate
73. cobalt (III) phosphide
89. selenium difluoride
105. tin (II) nitrite
74. iron (III) chloride
90. calcium sulfide
106. cobalt (III) oxide
75. copper (II) sulfide
91. silicon diselenide
107. titanium (II) acetate
76. manganese (III) chloride
92. diphosporus trioxide
108. vanadium (V) sulfide
77. potassium perchlorate
93. iron (III) chloride
109. chromium (III) hydroxide
78. aluminum sulfate
94. magnesium fluoride
110. lithium iodide
79. iron (II) carbonate
95. zinc oxide
111. lead (II) nitride
80. barium iodate
96. arsenic tribromide
112. silver bromide
81. magnesium phosphate
97. carbon dioxide
82. silver phosphate
98. gold (III) chlorite
Write the formula for the compound formed between the two ions listed
carbonate
chromate
chlorate
sulfide
chloride
sulfite
sodium
Na2CO3
Na2CrO4
·
·
·
·
silver
·
·
·
·
·
·
ammonium
·
·
·
·
·
·
Tin (II)
·
·
·
·
·
·
magnesium
·
·
·
·
·
·
lead (II)
·
·
·
·
·
·
aluminum
·
·
·
·
·
·
manganese (III)
·
·
·
·
·
·
potassium
·
·
·
·
·
·
barium
·
·
·
·
·
·
57
Unit 5 – Nomenclature
APPLICATIONS OF THE MOLE
Molar Mass
The molar mass of a compound is the mass of a mole of the _________________________________ particles of the
compound. Because each representative particle is composed of two or more atoms, the molar mass of the
compound is found by adding the molar masses of all of the ________________ in the representative particle. To
determine the molar mass of an element, find the element’s symbol on the periodic table and round the mass so there
is __________ digit beyond the decimal. For example, the molar mass of carbon (C) is ____________ g/mol, of
chlorine (Cl) is ___________ g/mol and of iron (Fe) is _____________ g/mol. In the case of NH 3, the molar mass
equals the mass of one mole of nitrogen atoms plus the mass of ___________ moles of hydrogen atoms.
Molar mass of NH3 = molar mass of N + 3 (molar mass of H)
Molar mass of NH3 = 14.0 + 3 (1.0) = 17.0 g/mol
Mole Conversions
You can use the molar mass of a compound to convert between mass and moles, just as you used the molar mass of
elements to make these conversions.

How many moles of magnesium in 56.3 g of Mg?

How many moles are in 146 grams of NH 3?

How many moles are in 295 grams of Cr(OH)3?

How many moles are in 22.5 grams of HCl?

How many grams of sodium chloride in 3.45 moles of NaCl?

How many grams are in 0.120 moles of AlF3?

How many grams are in 13.0 moles of H2SO4?

How many grams are in 1.6 moles of K2CrO4?
Percent Composition
Recall that every chemical compound has a definite composition - a composition that is always the same wherever
that compound is found. The composition of a compound is usually stated as the percent by mass of each element in
the compound. The percent of an element (X) in a compound can be found in the following way.
%X 

molarmassX # X ' s 
MolarMassC ompound
Determine the percent composition of chlorine in calcium chloride (CaCl 2). First, analyze the information
available from the formula. A mole of calcium chloride consists of one mole of calcium ions and
___________ moles of chloride ions. Next, gather molar mass information from the atomic masses on the
periodic table. To the mass of one mole of CaCl2, a mole of calcium ions contributes ______________ g,
and two moles of chloride ions contribute 2 x 35.5 g = 71.0 g for a total molar mass of 111.1 g/mol for
58
Unit 5 – Nomenclature
CaCl2. Finally, use the data to set up a calculation to determine the percent by mass of an element in the
compound.

Determine the percent composition of carbon in sodium acetate (NaC 2H3O2).

Calculate the percent composition aluminum of aluminum oxide (Al2O3).

Determine the percent composition of oxygen in magnesium nitrate, which has the formula Mg(NO 3)2.

Determine the percent composition of sulfur in aluminum sulfate, which has the formula Al 2(SO4)3.

Determine the percent composition of oxygen in zinc nitrite, which has the formula Zn(NO 2)2.
Percent Water in a Hydrate
Hydrates are compounds that incorporate ________________________ molecules into their fundamental solid
structure. In a hydrate (which usually has a specific crystalline form), a defined number of water molecules are
associated with each formula unit of the primary material. Gypsum is a hydrate with __________ water molecules
present for every formula unit of CaSO4. The chemical formula for gypsum is CaSO4 • 2 H2O and the chemical
name is calcium sulfate __________________. Note that the dot in the formula (or multiplication sign) indicates
that the waters are there.
Other examples of hydrates are: lithium perchlorate trihydrate - LiClO4 • 3 H2O;
magnesium carbonate pentahydrate - MgCO3 • 5 H2O;
and copper (II) sulfate pentahydrate - CuSO4 • 5 H2O.
The water in the hydrate (referred to as "water of hydration") can be removed by heating the hydrate. When all
hydrating water is removed, the material is said to be __________________________ and is referred to as a(n)
___________________________.
Experimentally measuring the ________________________ water in a hydrate involves first heating a known mass
of the hydrate to remove the waters of hydration and then measuring the mass of the anhydrate remaining. The
difference between the two masses is the mass of water _______________. Dividing the mass of the water lost by
the original mass of hydrate used is equal to the fraction of water in the compound. Multiplying this fraction by
___________ gives the percent water.

Determine the percent water in CuSO4 • 5 H2O (s).

Determine the percent water in MgCO3 •5 H2O (s).

Determine the percent water in LiClO4 • 3 H2O (s).
Empirical Formula
You can use percent composition data to help identify an unknown compound by determining its empirical formula.
The empirical formula is the ________________________ whole-number ratio of atoms of elements in the
compound. In many cases, the empirical formula is the actual formula for the compound. For example, the simplest
ratio of atoms of sodium to atoms of chlorine in sodium chloride is 1 atom Na : 1 atom Cl. So, the empirical
formula of sodium chloride is Na1Cl1, or NaCl, which is the true formula for the compound. The formula for
glucose is C6H12O6. The coefficients in glucose are all divisible by 6. The empirical formula of glucose is CH 2O.
59
Unit 5 – Nomenclature

Determine the empirical formula for Tl2C4H4O6.

Determine the empirical formula for N2O4.
The percent composition of an unknown compound is found to be 38.43% Mn, 16.80% C, and 44.77% O.
Determine the compound’s empirical formula. Because percent means “parts per hundred parts,” assume that you
have ___________ g of the compound. Then calculate the number of moles of each element in the 100 g of
compound. To obtain the simplest whole-number ratio of moles, _________________ each number of moles by the
smallest number of moles. Find the whole number mole ratio for the compound. These numbers become
the____________________________ in the empirical formula.

Determine the empirical formula of the following compound: 31.9 g Mg, 27.1 g P

The composition of an unknown acid is 40.00% carbon, 6.71% hydrogen, and 53.29% oxygen. Calculate
the empirical formula for the acid.

The composition of an unknown ionic compound is 60.7% nickel and 39.3% fluorine. Calculate the
empirical formula for the ionic compound.

The composition of a compound is 6.27 g calcium and 1.46 g nitrogen. Calculate the empirical formula for
the compound.

Find the empirical formula for a compound consisting of 63.0% Mn and 37.0% O.
Molecular Formula
For many compounds, the empirical formula is not the true formula. A molecular formula tells the
___________________ number of atoms of each element in a molecule or formula unit of a compound. The
molecular formula for a compound is either the same as the empirical formula or a whole-number
_______________________ of the empirical formula. In order to determine the molecular formula for an unknown
compound, you must know the molar mass of the compound in addition to its empirical formula. Then you can
compare the molar mass of the compound with the molar mass represented by the empirical formula.

The molecular mass of benzene is 78 g/mol and its empirical formula is CH. What is the molecular
formula for benzene? HINT: Calculate the molar mass represented by the formula CH. Calculate the
whole number multiple, n, and apply it to its empirical formula.

The simplest formula for butane is C2H5 and its molecular mass is about 60.0 g/mol. What is the molecular
formula of butane?

What is its molecular formula of cyanuric chloride, if the empirical formula is CClN and the molecular
mass is 184.5 g/mol?

The simplest formula for vitamin C is C3H4O3. Experimental data indicates that the molecular mass of
vitamin C is about 180. What is the molecular formula of vitamin C?

Maleic acid is a compound that is widely used in the plastics and textiles industries. The composition of
maleic acid is 41.39% carbon, 3.47% hydrogen, and 55.14% oxygen. HINT: Start by determining the
empirical formula for the compound.
60
Unit 5 – Nomenclature

The composition of silver oxalate is 71.02% silver, 7.91% carbon, and 21.07% oxygen. If the molar mass
of silver oxalate is 303.8 g/mol, what is its molecular formula?

The composition of silver oxalate is 71.02% silver, 7.91% carbon, and 21.07% oxygen. If the molar mass
of silver oxalate is 303.8 g/mol, what is its molecular formula?
Homework / Practice
1) A compound is found to have (by mass) 48.38% carbon, 8.12% hydrogen and the rest oxygen. What is its
empirical formula?
2) A compound is found to have 46.67% nitrogen, 6.70% hydrogen, 19.98% carbon and 26.65% oxygen. What is
its empirical formula?
3) A compound is known to have an empirical formula of CH and a molar mass of 78.11 g/mol. What is its
molecular formula?
4) Another compound, also with an empirical formula if CH is found to have a molar mass of 26.04 g/mol. What is
its molecular formula?
5) A compound is found to have 1.121 g nitrogen, 0.161 g hydrogen, 0.480 g carbon and
0.640 g oxygen. What
is its empirical formula? (Note that masses are given, NOT percentages.)
6) A compound containing only carbon, hydrogen and oxygen is found to be 48.38% carbon and 8.12% hydrogen
by mass. What is its empirical formula?
7) Phentyfloroform contains 57.54% C, 3.45% H, and 39.01% F. What is its empirical formula?
8) Calculate the empirical formula of the compound that contains 1.0 g S for each 1.5 g O.
9) Calculate the empirical formula of the compound containing 75.0% C and 25.0% H.
10) Calculate the empirical formula of the compound containing 81.8% C and 18.2% H.
11) The active ingredient in chocolate is theobromine; a sample was analyzed and determined to be composed of:
147.0 g C 14.0 g H 56.0 g O 98.0 g N
a.
Determine the % composition for each element.
b.
Determine the empirical formula for theobromine.
c.
The molecular weight of theobromine is known to be 180.0 g/mole. What is the molecular
formula?
12) Determine the empirical formula of the compound containing 37.5% C, 12.5% H, and 50.0% O by weight.
13) Determine the empirical formula of the compound containing 26.1% C, 4.3% H, and 69.6% O by weight.
14) Determine the empirical formula of the compound containing 38.7% C, 16.1% H, and 45.2% N by weight.
15) What is the empirical formula of a compound if a 50.0 g sample of it contains 9.1 g Na, 20.6 g Cr, and
22.2 g O?
16) A compound was analyzed and found to contain 13.5 g Ca, 10.8 g O, and 0.675 g H. What is the empirical
formula of the compound?
17) NutraSweet is 57.14% C, 6.16% H, 9.52% N, and 27.18% O. Calculate the empirical formula of NutraSweet
and find the molecular formula. (The molar mass of NutraSweet is 294.30 g/mol)
61
Unit 5 – Nomenclature
18) A compound consists of 85% silver and 15% florine by mass. What is the empirical formula?
19) A compound consists of 40% calcium, 12% carbon, and 48% oxygen by mass. What is the empirical formula by
mass?
20) A compound consists of 75% Magnesium, and 25% oxygen by mass. What is the empirical formula by mass?
21) A compound contains 50% Magnesium, 24% carbon, 16% oxygen, and 10% hydrogen. What is the empirical
formula?
22) Phentyfloroform contains 57.54% C, 3.45% H, and 39.01% F. What is its empirical formula?
23) Benzoic acid contains 68.8% Carbon, 4.95% Hydrogen and 26.2% Oxygen. Find the empirical formula?
24) Freons are gaseous compounds used in refrigeration. A particular freon contains 9.93% carbon, 56% chlorine,
and 31.4% fluorine. What is the empirical formula?
Naming and Formula Math Practice Test
Directions: Write the formula for the compound
1.
2.
3.
sodium phosphide
iron (II) perchlorate
vanadium (V) nitrite
4.
5.
6.
nickel (I) oxide
magnesium hydroxide
cesium nitride
7.
8.
9.
nitrogen trichloride
hydroxic acid
carbon tetrahydride
Directions: Name the compound
10. KCl
11. FeSO4
12. Li2O
13. Cr2S3
14. Ca3N2
15. Fe2S3
16. CuI2
17. PBr3
18. CO2
19. HNO3
20. What is the percent nitrogen in potassium nitrate?
21. What is the ratio of barium ions to Nitrogen ions in a formula unit of barium nitrate?
22. A compound is found to have 46.67% nitrogen, 6.70% hydrogen, 19.98% carbon and 26.65% oxygen. What is
its empirical formula?
23. A compound is found to have (by mass) 48.38% carbon, 8.12% hydrogen and the rest oxygen. What is its
empirical formula?
24. Phentyfloroform contains 57.54% C, 3.45% H, and 39.01% F. What is its empirical formula?
25. Determine the empirical formula of N3O6
26. The empirical formula of a compound is CH2. Its molecular mass is 70g/mol. What is its molecular formula?
27. Determine the empirical formula of a compound that is 82% carbon, 18% hydrogen.
28. Determine the molecular formula of a compound that is composed of 40.0% carbon, 6.7% hydrogen and
53.5% oxygen. The molecular mass is 120.0g/mol.
62
Unit 5 – Nomenclature
Multiple Choice Practice
29. Which of the following is a binary compound?
a. hydrogen sulfide
b. hydrogen sulfate
c.
d.
ammonium sulfide
ammonium sulfate
30. What is the formula for sodium oxalate?
a. NaClO
b. Na2ClO
c.
d.
Na2C2O4
NaC2H3O2
31. Which of the following is a binary compound?
a. potassium chloride
b. ammonium chloride
c.
d.
potassium chlorate
ammonium chlorate
32. Which is the correct formula for nitrogen monoxide?
a. NO
b. N2O
c.
d.
NO2
N2O3
33. Which of the following represents the correct formula for aluminum oxide?
a. AlO
c. AlO2
b. Al2O3
d. Al2O
34. Which of the following is the correct name for NaHCO3?
a. sodium hydrogen carbonate
b. sodium acetate
c. nitrogen hydrogen carbonate
d. sodium hydrogen carbon trioxide
35. In which of the following compounds does ionic bonding predominate?
a. NH4Cl
c.
b. CO2
d.
CH4
LiBr
36. Which of the following molecules contains only one non-bonding pair of valence electrons?
a. NH4+
c. C2H4
b. HCN
d. N2
37. What is the name of CaCl2?
a. calcium dichloride
b. calcium (II) chloride
c.
d.
monocalcium dichloride
calcium chloride
38. What is the name of Mg(NO3)2
a. Magnesium nitrate
b. Magnesium (II) nitrate
c.
d.
Magnesium dinitrate
Magnesium nitrogen oxide
39. What is the name of P2O5?
a. phosphorus oxide
b. phosphorus pentaoxide
c.
d.
diphosphorus pentaoxide
phosphorus (III) oxide
40. What is the formula for sulfur hexachloride?
a. S5Cl
b. SHCl
c.
d.
SCl5
SCl6
41. What is the name of the formula Fe(NO3)2?
a. iron nitrate
b. iron (II) nitrate
c.
d.
iron dinitrate
iron (III) nitrate
63
Unit 5 – Nomenclature
42. What is the formula for the compound nitrogen (II) oxide.
a. N2O3
b. NO
c.
d.
NO2
N2O
43. Which of the following is not a type of chemical formula?
a. Empirical
b. Molecular
c.
d.
Structural
Parabola
44. What is the approximate percentage oxygen in the formula mass of Ca(NO 3)2?
a. 28
c. 58
b. 42
d. 96
e.
164
45. Which formulas could represent the empirical formula and the molecular formula of a given compound?
a. CH2O and C4H6O4
d. CH2 and C3H6
b. CHO and C6H12O6
e. CO and CO2
c. CH4 and C5H12
46. When combining with nonmetallic atoms, metallic atoms generally will
a. lose electrons and form negative ions
b. lose electrons and form positive ions
c. gain electrons and from negative ions
d. gain electrons and form positive ions
47. What is the empirical formula of the compound whose molecular formula is P 4O10?
a. PO
c. P2O5
b. PO2
d. P8O20
48. What is the percent by mass of oxygen in magnesium oxide, MgO?
a. 20%
b. 40%
c.
d.
50%
60%
49. A compound is 86% carbon and 14% hydrogen by mass. What is the empirical formula for this compound?
a. CH
c. CH3
b. CH2
d. CH4
50. What is the percent by mass of water in the hydrate Na2CO3 • 10H2O (formula mass = 286)?
a. 6.89%
c. 26.1%
b. 14.5%
d. 62.9%
51. What is the gram formula mass of (NH4)3PO4?
a. 113 g
b. 121 g
c.
d.
149 g
404 g
52. What is the empirical formula of a compound that contains 85% Ag and 15% F by mass?
a. AgF
c. AgF2
b. Ag2F
d. Ag2F2
53. What is the percent water in hydrated calcium chloride CaCl 2•2H2O?
a. 66.67%
b. 32.47%
64
c.
d.
24.51%
12.26%
Unit 5 – Nomenclature
54. What is the empirical formula for a compound that contains 17.34% hydrogen and
a. C5H
c. CH3
b. C2H5
d. CH2
82.66% carbon?
55. What is the molecular formula for a compound that is 46.16% carbon, 5.16% hydrogen and 48.68% fluorine if
the molar mass of this compound is 156.12 g?
a. C3H4F2
c. C6H8F4
b. C5H10F5
d. C6H6F3
56. One mole of (NH4)2HPO4 contains _?_ moles of hydrogen atoms.
a. 1
b. 5
c.
6
d.
57. Which of the following is named incorrectly?
a. H2CO2 : carbonous acid
b. HClO2 : chlorous acid
c. H2SO4 : sulfuric acid
d.
e.
HClO : hydrochlorous acid
H3PO3 : phosphorous acid
58. The name for the compound NaHSO4 is:
a. sodium hydrogen sulfate
b. sodium persulfate
c. sodium bisulfate
d.
e.
two of these are correct
none of these is correct
9
59. A sample of an alcohol is tested and found to contain 52% carbon, 35% oxygen, and 13% hydrogen by mass.
Tests indicate that the molecular weight of the molecule is between 30 and 80. What is the molecular formula of
the alcohol?
a.
b.
c.
C2H5OH
C3H7OH
C5H11OH
d.
e.
C4H9OH
CH3OH
60. A 5.15 gram sample of a hydrocarbon is burned in oxygen, producing 15.6 grams of carbon dioxide and
8.45 grams of water. Assuming an excess of oxygen, what is the empirical formula of the hydrocarbon?
a.
b.
c.
CH4
C7H16
C9H20
d.
e.
C5H12
C3H8
61. For which of the following compounds does 0.400 mol have a mass of 12.8 grams?
a.
b.
c.
CH3OH
CH4
CCl4
d.
e.
CO2
C4H10
62. What is the molar mass of glucose, C6H12O6?
a.
b.
18 g
220.17 g
c.
d.
65
12.01 g
180.18 g
e.
160.00 g
Unit 6- Chemical Reactions
CHEMICAL REACTIONS
All chemical reactions have two parts: (1) A substance that undergoes a reaction is called a
__________________________. In other words, reactants are the substances you start with. (2) When reactants
undergo a chemical change, each new substance formed is called a ___________________________. In other
words, the products are the substances you end up with. The reactants turn into the products.
Reactants → Products. In a chemical reaction, the way atoms are joined is changed. Atoms aren’t
__________________________ or destroyed.
Chemical reactions can be described several ways.

In a sentence: Copper reacts with chlorine to form copper (II) chloride.

In a word equation: Copper + chlorine → copper (II) chloride
The arrow separates the reactants from the products. The arrow reads “reacts to ________________.” The plus
sign reads “_____________.” (s) after the formula implies the substance is a ___________________. (g) after the
formula implies the substance is a gas. (l) after the formula implies the substance is a ______________________.
(aq) after the formula implies the substance is aqueous, a solid dissolved in _____________________. __________
used after a product indicates a gas, same as (g). ↓ used after a product indicates a ________________, same as (s).
_____________ indicates a reversible reaction. ________________ or ________________ shows that heat is
supplied to the reaction. ___________________ is used to indicate a catalyst used supplied, in this case, platinum.
A catalyst is a substance that ____________________ ____________ a reaction without being changed by the
reaction. Enzymes are biological or ______________________ catalysts.
There are seven elements that never want to be alone. They form ________________________ molecules. H2 , N2,
O2 , F2 , __________ , Br2 , I2. (1 + 7 pattern on the periodic table)
The following are indications that a chemical reaction has occurred: formation of a
____________________________, evolution of a gas, _____________________ change, and absorption or release
of ________________.
A ________________________ formula uses formulas and symbols to describe a reaction. All chemical equations
are sentences that describe reactions.

Convert the following sentences to chemical equations.
a)
Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form solid iron (II) chloride and
hydrogen sulfide gas.
b) Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon
dioxide gas and sodium nitrate dissolved in water.

Convert the following chemical equations to sentences.
a)
Fe (s) + O2 (g) → Fe2O3 (s)
b) Cu (s) + AgNO3 (aq) → Ag (s) + Cu(NO3)2 (aq)
66
Unit 6- Chemical Reactions
Balancing Equations
Atoms can’t be ______________________ or destroyed. All the atoms we start with we must end up with. A
balanced equation has the same number of each element on both _________________ of the equation. Example: C
+ O2 → CO
This equation is NOT balanced. There is one carbon atom on the left and ________ on the right.
There are two oxygen atoms on the left and only one on the right. We need one more oxygen atom in the products.
We can’t change the formula, because it describes what it is. In order to have two oxygen atoms, another CO must
be produced. But where did the other carbon come from? We must have started with two carbon atoms. The
balanced chemical equation is 2 C + O2 → 2 CO
Rules for Balancing

Write the correct formulas for all the reactants and products.

Count the number of atoms of each type appearing on both sides.

Balance the elements one at a time by adding coefficients (the numbers in front).

Check to make sure it is balanced.
Never change a _________________________ to balance an equation. If you change the formula you are
describing a different reaction. Never put a coefficient in the middle of a formula. 2 NaCl is okay; Na2Cl is not.

Balance the following reaction. H2 + O2 → H2O
Balance elements in the following order: (1) metals; (2) nonmetals; (3) hydrogen; and (4) oxygen
If an atom appears more than once on a side, balance it last. If you fix everything except one element, and it is even
on one side and odd on the other, double the first number, then move on from there.

Balance the following equations.
1) _____ CH4 + _____ O2 → _____ CO2 + _____ H2O
2) _____ AgNO3 + _____ Cu → _____ Cu(NO3)2 + ______ Ag
3) _____ Mg + _____ N2 → _____ Mg3N2
4) _____ P + _____ O2 → ______ P4O10
5) _____ Na + _____ H2O → _____ H2 + _____ NaOH
6) _____ Pb(NO3)2 + _____ K2CrO4  ______ PbCrO4 + ______ KNO3
7) _____ MnO2 + _____ HCl  _____ MnCl2 + ______ H2O + _____ Cl2
8) _____ Ba(CN)2 + _____ H2SO4  _____ BaSO4 + _____ HCN
9) _____ Zn(OH)2 + _____ H3PO4  _____ Zn3(PO4)2 + _____ H2O
TYPES OF REACTIONS
Reactions fall into 5 categories. We will recognize the type by the reactants. We will be able to predict the
products. For some we will be able to predict whether they will happen at all.
67
Unit 6- Chemical Reactions
Synthesis Reactions
Synthesize means to put together. Whenever two or more substances combine to form one single product, the
reaction is called a synthesis reaction.
Examples: Ca + O2 → CaO
and
P2O5 + 3 H2O → 2 H3PO4
We can predict the products if they are two elements. All you need to do is combine the elements, metals first, and
criss-cross oxidation numbers if necessary. After predicting the product, the reaction must be balanced.

Mg + N2 →

CaO + H2O →
On page 6 of the Chemistry Reference Packet, this reaction is an example of “Metal oxide - water reactions.” The
product listed in the packet is “base.” A base is a metallic hydroxide.

SO2 + H2O →
On page 6 of the Chemistry Reference Packet, this reaction is an example of “Nonmetal oxide - water reactions.”
The product listed in the packet is “________________.” The acid is a ternary acid. Ternary acids start with
_________ and end in O. The other element goes in the center. This is the only compound for which you can add
the number of elements and use these numbers as subscripts.

Write and balance the following synthesis reactions.
a) Ca + Cl2 →
d) Al + O2 →
b) Fe + O2 →
e) SO3 + H2O →
f) N2O5 + H2O →
HINT: Use iron (II).
c) K2O + H2O →
Decompositions Reactions
The word decompose implies the compound will “fall apart.” In a decomposition reaction, one compound breaks
down into _____________ or more simple substances.
NaCl → Na + Cl2
CaCO3 → CaO + CO2
We can easily predict the products if it is a binary compound. A binary compound is made up of only two elements.
The compound merely falls apart into its elements.

H 2O →

HgO →
If the compound has more than two elements, you must consult the Reference Tables, page 6.

NiCO3 →
NiCO3 is called nickel (II) carbonate. The packet states that a metallic carbonate decomposes to form a MO
(metallic oxide) and CO2. The metallic oxide is nickel (II) oxide.

Use the Chemistry Reference Tables to write and balance the following decomposition reactions.
a) KClO3 →
b) CaBr2 →
c) Li2CO3 →
68
Unit 6- Chemical Reactions
d) Cr(OH)2 →
e) NaHCO3 →
f) HNO3 →
(Dinitrogen pentoxide is one of the products.)
Single Replacement
In a single-displacement reaction, one element takes the place of another in a compound. One reactant must be an
element, and the one reactant must be a _______________________. The products will be a different element and a
different compound.
F2 + LiCl → LiF + Cl2
Remember zinc, Zn, always forms a ___________ ion doesn’t need parenthesis. ZnCl 2 is zinc chloride. In addition,
silver, Ag, always forms a ___________ ion. AgCl is silver chloride.
Some single replacement reactions do not occur because some elements are not as ________________ as others. A
more active element _________________________ a less active element. There is a list referred to as the Activity
Series on page 7 of your Chemistry Reference Packet. A higher element on the list replaces lower element. If the
element by itself is lower on the list, the reaction will ___________ occur.
Metals replace metals (and hydrogen)

K + NaCl →
Potassium wants to replace ________________________. You must check the activity series on page 7 of your
Chemistry Reference Packet to see if this is possible. Because K is higher, potassium can replace sodium. The
potassium will bond with the _____________________ and the sodium will be alone. You must always check to
see if the compound formed needs criss-crossing. Check for balancing.

Sn + FeCl3 →
Because Sn is NOT higher, tin cannot replace iron. No reaction occurs.

Write and balance the following single replacement reaction.
a) Rb + AlN →
c) Ag + CoBr2 →
b) Zn + HCl →
Metals replace hydrogen

Na + H2O (cold) →
Think of water as HOH. Metals high enough on the activity series replace the first ______ and combine with the
OH1- (hydroxide) according to page 6 of the Reference Tables. Is sodium above hydrogen and higher than the line
marked “Replace hydrogen from cold water” on the activity series? Since the answer is yes, sodium replaces the
first H, bonding with hydroxide.

Mg + HCl →
Metals higher on the activity series replace the H and combine with the nonmetal according to page 6 of the
Reference Tables. Hydrogen gas is a second product. Is magnesium above hydrogen on the activity series?
69
Unit 6- Chemical Reactions

Write and balance the following single replacement reactions.
a) Ag + H2O (steam) →
c) Cr + H3PO4 → (HINT: Use Cr3+ )
b) Cu + H2SO4 →
d) Ca + H2O (steam) →
Nonmetals can replace other ________________________. This is limited to F 2 , Cl2 , Br2 and I2 The order of
activity is listed in the Chemistry Reference Packet, page 7. Higher replaces _____________.

F2 + HCl →
Is fluorine above chlorine in the activity series of halogens? Since the answer is yes, fluorine replaces the chlorine,
bonding with hydrogen.

Write and balance the following single replacement reactions.
a) Br2 + KCl →
b) Cl2 + KI →
Double Replacement
In double-displacement reactions, the positive portions of two ___________________ compounds are interchanged.
The reactants must be two ionic compounds or ______________. Double replacement reactions usually take place
in ________________________ solution.

NaOH + FeCl3 →
The positive ions change place. You must check to see if you need to criss-cross the products. Now balance. A
double replacement reaction will only happen if one of the products: (1) doesn’t dissolve in water and forms a
__________________, (2) is a _____________ that bubbles out, or (3) is a _________________________
compound usually water.
3NaOH + FeCl3 → Fe(OH)3 + 3NaCl
None of the products are familiar gases. Both products are ionic (not covalent) because they start with metals. We
must consult the Solubility Rules on page 6 of the Chemistry Reference Tables to determine if a solid (a
________________________) is formed. The “Soluble” side of the Solubility Rules states that Group 1 (IA) salts
are soluble; therefore, NaCl is soluble and is NOT the precipitate. The “Insoluble” side of the Solubility Rules states
that all hydroxides except Group 1, Sr, Ba and NH41+ are INSOLUBLE. Therefore, Fe(OH)3 is the precipitate
(solid). In molecular equations, the formulas of the compounds are written as though all species existed as
molecules or whole units. An ionic equation shows dissolved ionic compounds in terms of their free ions. Ions that
are not involved in the overall reaction are called spectator ions. The net ionic equation indicates only the species
that actually take part in the reaction. The following steps are useful for writing ionic and net ionic equations:
1) Write a balanced molecular equation for the reaction.
2) Rewrite the equation to indicate which substances are in ionic form in solution. Remember that all soluble
salts (and other strong electrolytes), are completely dissociated into cations and anions. This procedure
gives us the ionic equation.
3) Lastly, identify and cancel spectator ions on both sides of the equation to arrive at the net ionic equation.
Example: sodium hydroxide + iron (III) chloride yields iron (III) hydroxide + sodium chloride
70
Unit 6- Chemical Reactions
Balanced Molecular Equation: 3 NaOH + FeCl3  Fe(OH)3 + 3 NaCl
Complete Ionic Equation:
3Na1+ + 3OH1- + Fe3+ + 3Cl1-  Fe(OH)3 + 3Na1+ + 3Cl1Net Ionic Equation: 3OH1- + Fe3+  Fe(OH)3

Write and balance the following double replacement reaction. Assume the reaction takes place. In addition,
identify the precipitate and write the net ionic equation.
a) CaCl2 + NaOH →
c) KOH + Fe(NO3)3 →
b) CuCl2 + K2S →
d) (NH4)2SO4 + BaF2 →
Combustion
A combustion reaction is one in which a substance rapidly combines with ____________________ to form one or
more oxides. Combustion reactions involve a compound composed of only C and H (and maybe O) that is reacted
with oxygen gas. If the combustion is complete, the products will be CO 2 and __________________. Combustion
reactions produce heat, and are therefore considered exothermic reactions.

Complete and balance the following combustion reactions.
a) C4H10 + O2 →
c) C8H8 + O2 →
b) C6H12O6 + O2 →
d) C3H8O3 + O2 →
To determine which type a reaction is, look at the reactants. (E = element and C = compound)
E+E
Synthesis
C
Decomposition
E+C
Single replacement
C+C
Double replacement
CH cpd + O2
Combustion
Note: Two other common synthesis reactions include: nonmetallic oxide + water and metallic oxide + water.

Identify whether the reaction is synthesis, decomposition, single replacement, double replacement or
combustion.
a) H2 + O2 →
e) KBr + Cl2 →
b) H2O →
f) Zn + H2SO4 →
c) Mg(OH)2 + H2SO3 →
g) AgNO3 + NaCl →
d) HgO →
h) C6H6 + O2 →
Homework / Practice
Directions: Balance the following equations.
SYNTHESIS
1.
S + O2  SO2
3.
P + O2  P2O3
2.
S + O2  SO3
4.
Mg + N2  Mg3N2
71
Unit 6- Chemical Reactions
5.
N2 + O2  NO2
8.
Al + N2  AlN
6.
Na + O2  Na2O
9.
Hg + I2  HgI2
7.
Cu + S  Cu2S
10. Fe + O2  Fe2O3
DECOMPOSITION
11. HgO  Hg + O2
12. MgSO4•7H2O  MgSO4 + H2O
16. BaO2  BaO + O2
13. KClO3  KCl + O2
18. NO2  N2 + O2
14. NH4NO3  N2O + H2O
19. CaCO3  CaO + CO2
15. NaNO3  NaNO2 + O2
20. H2O  H2 + O2
17. H2O2  H2O + O2
SINGLE REPLACEMENT (SINGLE DISPLACEMENT)
21. AlI3 + Cl2  AlCl3 + I2
26. ZnS + O2  ZnO + SO2
22. CH4 + Cl2  CHCl3 + HCl
27. Na + H2O  NaOH + H2
23. Al + CuSO4  Al2(SO4)3 + Cu
28. Al + H2SO4  Al2(SO4)3 + H2
24. Fe2O3 + Al  Al2O3 + Fe
29. Zn + NaOH  Na2ZnO2 + H2
25. Zn + HCl  ZnCl2 + H2
30. AgNO3 + Zn  Zn(NO3)2 + Ag
DOUBLE REPLACEMENT
31. Fe(OH)3 + H2SO4  Fe2(SO4)3 + H2O
36. AgNO3 + H2S  Ag2S + HNO3
32. AgNO3 + K2CrO4  Ag2CrO4 + KNO3
37. CaCO3 + HCl  CaCl2 + H2CO3
33. AgNO3 + CuCl2  AgCl + Cu(NO3)2
38. Hg2(NO3)2 + NaCl  Hg2Cl2 + NaNO3
34. Pb(NO3)2 + HCl  PbCl2 + HNO3
39. BaCl2 + (NH4)2CO3  BaCO3 + NH4Cl
35. MgCl2 + NaOH  Mg(OH)2 + NaCl
40. Al(OH)3 + NaOH  NaAlO2 + H2O
COMBUSTION
41. CH4 + O2  CO2 + H2O
44. C5H8 + O2  CO2 + H2O
42. C4H10 + O2  CO2 + H2O
45. CH3OH + O2  CO2 + H2O
43. C3H6 + O2  CO2 + H2O
46. C6H12O6 + O2  CO2 + H2O
Predict the products of the following reactions
47. MgCl2 + O2
51. BaO + CO2 
55. NaCl + O2 
48. Na + O2 
52. BeO + CO2 
56. Ra + Cl2 
49. P2O3 + H2O 
53. Al2O3 + H2O 
57. Ni(ClO3)2 
50. K2O + H2O 
54. N2O5 + H2O 
58. Ag2O 
72
Unit 6- Chemical Reactions
59. HNO2 
70. Al + Pb(NO3)2 --->
81. K2CO3 + BaCl2 
60. Fe(OH)3 
71. Cl2 + NaI 
82. Cd3(PO4)2 + (NH4)2S 
61. ZnCO3 
72. Fe + AgC2H3O2
83. Co(OH)3 + HNO3 
62. Cs2CO3 
73. Al + CuCl2 
84. AgNO3 + KCl 
63. Al(OH)3 
74. Br2 + CaI2 
85. Na2CO3 + H2SO4 
64. H2SO4 
75. Al + HCl 
86. Al(OH)3 + HC2H3O2 
65. RbClO3 
76. Mg + HCl 
87. Al2(SO4)3 + Ca3(PO4)2 
66. RaCl2 
77. Zn + H2SO4 
88. Cr2(SO3)3 + H2SO4 
67. ZnS + O2 
78. Fe + CuSO4 
89. AgC2H3O2 + K2CrO4 
68. K + H2O 
79. Cl2 + MgI2 
69. Fe + HCl 
80. Ca(OH)2 + H3PO4 
Write the complete ionic equation, identify the spectator ions, write the net ionic equation
90. K2CO3 + BaCl2  BaCO3 + KCl
91. Cd3(PO4)2 + (NH4)2S  CdS + (NH4)3PO4
92. AgNO3 + KCl  AgCl2 + K2NO3
93. Na2CO3 + H2SO4  H2CO3 + Na2SO4
94. Al(OH)3 + HC2H3O2  Al(C2H3O2)3 + HOH
95. Al2(SO4)3 + Ca3(PO4)2  AlPO4 + CaSO4
96. Cr2(SO3)3 + H2SO4  Cr2(SO4)3 + H2SO3
97. AgC2H3O2 + K2CrO4  Ag2CrO4 + KC2H3O2
Write the balanced equation (including states) and identify the type of reaction:
98. Aqueous solutions of ammonium chloride and lead (II) nitrate produce lead (II) chloride precipitate and
aqueous ammonium nitrate.
99.
Solid carbon disulfide burns in oxygen to yield carbon dioxide and sulfur dioxide gases.
100. Iron metal reacts with aqueous silver nitrate to produce aqueous iron (III) nitrate and silver metal.
101. Solid potassium nitrate yields solid potassium nitrite and oxygen gas.
102. Calcium metal reacts with chlorine gas to produce solid calcium chloride.
Chemical Reactions and Balancing Practice Test
1.
Write the balanced equation of the reaction that occurs when calcium carbonate decomposes to form
calcium oxide and carbon dioxide.
Directions: Balance and identify the type of reaction
2. ____K3PO4 + ____Al(NO3)3  ____KNO3 + ____AlPO4
3. ____Fe2O3 + ____Al  ____Fe + ____Al2O3
73
Unit 6- Chemical Reactions
4.
5.
6.
____NaOH  ____Na2O + ____H2O
____HCl + ____Mg  ____MgCl2 + ____H2
____C2H4 + ____O2  ____CO2 + ____H2O
7.
Write the balanced equation of the synthesis reaction that occurs when iron metal and oxygen gas react to
form iron (III) oxide.
8.
Write the balanced equation of the combustion reaction that occurs when ethane (C 2H6) reacts with oxygen
to form carbon dioxide and water.
9.
Write the balanced equation of the reaction that occurs when calcium carbonate decomposes to form
calcium oxide and carbon dioxide.
Directions: Predict the products and balance the reaction
10. _____Na + ______O2 
11. _____Al + _____Pb(NO3)2 
12. _____NaF + _____Br2 
Multiple Choice Practice
13. In the following unbalanced reaction, what is the coefficient of HOH once the reaction is balanced?
H2CO3 + KOH --> HOH + K2CO3
a. 1
b. 2
c. 3
d. 4
14. What are the different types of chemical reactions?
a. synthesis, fusion, combustion, fission, and decomposition
b. single replacement, combustion, and double replacement
c. synthesis, fission, single replacement, combustion, and fusion
d. synthesis, decomposition, combustion, single and double replacement
15. What type of reaction is represented by 2 or more elements forming a compound?
a. Decomposition
c. Combustion
b. Synthesis
d. Single replacement
16. Decomposition is the burning of hydrocarbons in the presence of oxygen.
a. True
b. False
17. Which equation represents a double replacement reaction?
a. CaCO3  CaO + CO2
b. CH4 + 2O2  CO2 + 2H2O
c.
d.
LiOH + HCl  LiCl + H2O
C3H8 + 5O2  3CO2 + 8H2O
18. MgSO4 + BaCl2  MgCl2 + BaSO4, is an example of what type of chemical reaction?
a. Single replacement
c. Combustion
b. Synthesis
d. Double replacement
19. Zn + 2 AgNO3  2 Ag + Zn(NO3)2, is an example of what type of chemical reaction?
a. Synthesis
c. Decomposition
b. Single replacement
d. Double replacement
20. The cation of one compound replaces the cation in another compound in a double replacement reaction.
a. True
b. False
21. Which statement best describes the conservation of atoms in all balanced chemical equations?
a. There is a conservation of mass, number of protons, and charge.
b. There is a conservation of mass, electronegativity, and charge.
74
Unit 6- Chemical Reactions
c.
d.
There is a conservation of only energy, and charge.
There is a conservation of mass, energy, and charge.
22. Given the unbalanced equation CuS + O2  CuO + SO2 When it is balanced, what is the sum of the
coefficients?
a. 8
b. 9
c. 10
d. 11
23. Which one of these chemical reactions is balanced?
a. Na + Cl2  NaCl
b. H2 + O2  H2O
c.
d.
CuCO3  CuO + CO2
KClO3  KCl + O2
24. Which is the correct way of setting up a word equation for this balanced chemical equation,
2Na + Cl2  2NaCl?
a. Sodium react with chlorine gas to produce sodium chloride.
b. 2 moles of sodium react with 1 mole of chlorine to yield 1 mole of sodium chloride.
c. 2 moles of sodium react with 1 mole of chlorine gas to yield 2 mole of sodium chloride.
d. 2 moles of sodium added with 1 mole of chlorine gas to yield 1 mole of sodium chloride.
25. 2 moles of copper react with 1 mole of oxygen gas to yield 2 moles of copper (ll) oxide. How would you
express this word equation into a balanced chemical equation?
a. Cu + O  CuO
c. 2Cu + O  CuO
b. Cu + O2  CuO
d. 2Cu + O2  2CuO
26. When the following reaction is balanced, the sum of all of the coefficients in the equation is:
NaCl + H2SO4 + MnO2 → Na2SO4 + MnSO4 + H2O + Cl2
a.
b.
11
10
c.
d.
6
16
e.
14
27. Which equation represents combustion?
a. 4 Fe + 3 O2  2 Fe2O3
b. 2H2O - 2H2 + O2
c. CH4 + 2O2  CO2 + 2 H2O
d. Cu + 2 AgNO3  Cu(NO3)2 + 2 Ag
28. Given the unbalanced equation: Al + O2 Al2O3 When this equation is completely balanced using the
smallest whole numbers, what is the sum of the coefficients?
a. 9
b. 7
c. 5
d. 4
29. Given the reaction CH4 + 2O2 --> CO2 + 2H2O,
What amount of oxygen is needed to completely react with 1 mole of CH 4?
a. 2 moles
c. 2 grams
b. 2 atoms
d. 2 molecules
30. The balanced equation for the complete combustion of benzene, C6H6, is
a.
b.
c.
d.
C6H6 + 12 H2O 
2 + 15 H2
2 C6H6 + 9 O2  12 CO + 6 H2O
C6H6 + O2  CO2 + H2O
2 C6H6 + 15 O2  12 CO2 + 6 H2O
75
Unit 6- Chemical Reactions
31. Which equation shows conservation of atoms?
a. H2 + O2  H2O
b. H2 + O2  2H2O
c.
d.
2H2 + O2  2H2O
2H2 + 2O2  2H2O
32. The reaction of magnesium with elemental iodine, I2, yields magnesium iodide. Write a balanced chemical
equation for this reaction. Which equation shows conservation of mass and charge?
a.
b.
c.
d.
NH4Br  NH3 + Br2
2Mg + Fe3+  Mg2+ + 3Fe
H2SO4 + LiOH  Li2SO4 + H2O
2+
Cu + 2Ag

+ 2Ag
33. When ethanol undergoes complete combustion, the products are carbon dioxide and water.
__ C2H5OH + __ O2  __ CO2+ __ H2O What are the respective coefficients when the equation is
balanced with the smallest whole numbers?
a.
b.
c.
2, 7, 4, 6
1, 3, 2, 3
2, 2, 1, 4
34. What is the net ionic equation?
a.
b.
c.
d.
e.
d.
e.
1, 2, 3, 2
2, 4, 6, 4
AgNO3(aq) + KBr(aq)  AgBr(s) + KNO3(aq)
K+(aq) + NO3-(aq)  KNO3(s)
AgNO3(aq) + KBr(aq)  AgBr(s)
K+(aq) + NO3-(aq)  KNO3(aq)
AgNO3(aq) + KBr(aq)  AgBr(s) + KNO3(aq)
Ag+(aq) + Br-(aq) AgBr(s)
76
Unit 7- Stoichiometry
STOICHIOMETRY
The word stoichiometry is Greek for “________________________ elements.” The calculations of quantities in
chemical reactions are based on a ______________________ equation. We can interpret balanced chemical
equations several ways. Using the methods of stoichiometry, we can measure the amounts of substances involved in
chemical reactions and relate them to one another. The group or unit of measure used to count numbers of atoms,
molecules, or formula units of substances is the ______________ (abbreviated mol).
Moles in Chemical Reactions
The coefficients tell us how many moles of each kind of element or compound we have.
2 Al2O3 → 4 Al + 3 O2
2 moles of aluminum oxide form 4 moles of aluminum and 3 moles of oxygen gas.
2 H2 + O2 → 2 H2O
___ mole(s) of hydrogen gas and ___ mole of oxygen form ___ mole(s) of water.
2 Na + 2 H2O → 2 NaOH + H2
___ moles of sodium and ___ moles of water form ___ moles of sodium hydroxide and ___ mole of hydrogen gas.
2 Al2O3 → 4 Al + 3 O2
Every time we use 2 moles of Al2O3 we make 3 moles of O2. Every time we use 2 moles of Al2O3 we make 4 moles
of Al.

Using the balanced equation above, how many moles of O2 are produced when 3.34 moles of Al 2O3
decompose?

2 C2H2 + 5 O2 → 4 CO2 + 2 H2O
a) If 3.84 moles of C2H2 are burned, how many moles of O2 are needed?
b) How many moles of C2H2 are needed to produce 8.95 moles of H2O?
c) If 2.47 moles of C2H2 are burned, how many moles of CO2 are formed?

SiCl4 + 2 Mg → 2 MgCl2 + Si
3.74 mol of Mg would make how many moles of Si?
Mass in Chemical Reactions
2 Al2O3 → 4 Al + 3 O2
2 x (102.0) grams of aluminum oxide form 4 x (27.0) grams of aluminum and 3 x (32.0) grams of oxygen. The mass
of Al2O3 was found by adding the masses of 2 aluminums & 3 oxygens.
(2 x 27.0 + 3 x 16.0 = 102.0)
2 H2 + O2 → 2 H2O
2 x (_________) grams of hydrogen and ___ x (16.0) of oxygen form ___ x (_______) grams of water.
77
Unit 7- Stoichiometry
2 Na + 2 H2O → 2 NaOH + H2
___ x (23.0) grams of sodium and 2 x (________) grams of water form ___ x (_________) grams of sodium
hydroxide and ___ x (________) grams of hydrogen gas.
The law of conservation of _____________ applies in chemical reactions. The mass of the reactants equals the mass
of the ________________________.

Show that the following equation follows the Law of Conservation of Mass.
2 Al2O3 → 4 Al + 3 O2
Mass – Mole Stoichiometry
The mass of 1 mole of a pure substance is called its _________________ mass. To convert the mass of an element
or compound to the number of moles, use the mass of 1 mol as a conversion factor. We can convert
___________________ to moles using the periodic table. Then we must apply the mole to mole conversion to
change chemicals using the balanced equation. Finally we will turn the moles back to grams using the periodic
table.

2 C2H2 + 5 O2 → 4 CO2 + 2 H2O
a) How many moles of C2H2 are needed to produce 8.95 g of H2O?
b) If 2.47 moles of C2H2 are burned, how many grams of CO2 are formed?

SiCl4 + 2 Mg → 2 MgCl2 + Si
How many moles of Mg are needed to make 9.3 g of Si?

3 Al + 3 NH4ClO4 → Al2O3 + AlCl3 + 3 NO + 6 H2O
How many moles of water are produced when 32 grams of aluminum are used?

CO2 + 2 LiOH → Li2CO3 + H2O
What mass of water can be produced from 3.66 moles of lithium hydroxide (LiOH)?

2 Al + 3 I2 → 2 AlI3
Calculate the mass of AlI3 (Aluminum Iodide) that can be produced from 3.00 mol of Al.
Mass – Mass Stoichiometry

2 Fe + 3 CuSO4 → Fe2(SO4)3 + 3 Cu
If 10.1 g of Fe are added to a solution of copper (II) sulfate, how much solid copper would form?

2 Al + 3 I2 → 2 AlI3
Calculate the mass of I2 needed just to react with 35.0 g of Al.

SiCl4 + 2 Mg → 2 MgCl2 + Si
How many grams of MgCl2 are produced along with 9.3 g of silicon?

3 Al + 3 NH4ClO4 → Al2O3 + AlCl3 + 3 NO + 6 H2O
a) How many grams of Al must be used to react with 652 g of NH4ClO4?
b) How many grams of NO are produced if 150.0 grams of AlCl 3 are also produced?
78
Unit 7- Stoichiometry
Particles in Chemical Reactions
The number of things in one mole is 6.022 x 10 23. This big number has a short name: the Avogadro constant.
Atom - ______________________
Molecule - Molecular compound (non-metals) or ______________________ (O2 etc.)
Formula unit - _________________ Compounds (Metal and non-metal or metal and a polyatomic ion)
2 Al2O3 → 4 Al + 3 O2
2 x (6.022 x 1023) formula units of aluminum oxide form 4 x (6.022 x 10 23) atoms of aluminum and
3 x (6.022 x 1023) molecules of oxygen.
2 H2 + O2 → 2 H2O
2 x (___________) molecules of hydrogen and ___ x (6.022 x 10 23) molecules of oxygen form ___ x
(___________) molecules of water.
2 Na + 2 H2O → 2 NaOH + H2
23
___ x (6.022 x 10 ) atoms of sodium and ___ x (___________________) molecules of water form ___ x
(__________________) formula units of sodium hydroxide and ___ x (6.022 x 1023) molecules of hydrogen gas.

SiCl4 + 2 Mg → 2 MgCl2 + Si
How many moles of MgCl2 are produced along with 8.76 x 1024 atoms of silicon?
Gases and Reactions
In gas conversions, liters of a gas are converted to moles and vice-versa. ____________ stands for standard
temperature and pressure. 0ºC is standard _________________________, and
1 atmosphere is standard
pressure. At STP, ____________ L of a gas = 1 mole

2 H2O → 2 H2 + O2
If 6.45 grams of water are decomposed, how many liters of oxygen will be produced at STP?

CH4 + 2 O2 → CO2 + 2 H2O
How many liters of CH4 at STP are required to completely react with 17.5 L of O 2?

2 C8H18 + 25 O2 → 16 CO2 + 18 H2O
Octane, C8H18, is one of the hydrocarbons in gasoline. How many liters of oxygen are required, at STP, to
burn 1.00 g of octane?

2 C4H10 + 13 O2 → 8 CO2 + 10 H2O
How many liters of CO2 at STP will be produced from the complete combustion of 23.2 g C 4H10?

2 NiS + 3 O2 → 2 NiO + 2 SO2
What volume of sulfur dioxide is produced from 123 grams of nickel (II) sulfide at STP?
According to Avogadro, equal volumes of gas, at the _____________ temperature and pressure, contain the same
number of particles. _______________ are numbers of particles. We can also change between particles and liters at
STP.
79
Unit 7- Stoichiometry

2 C4H10 + 13 O2 → 8 CO2 + 10 H2O
a) How many molecules of CO2 at STP will be produced from the complete combustion of 18.2 L C 4H10 ?
b) How many molecules of O2 at STP are needed to produce 18.2 L of steam?
c) How many liters of CO2 at STP are produced from 3.2 x 1024 molecules of butane, C4H10?

4 NH3 + 6 NO → 5 N2 + 6 H2O
Nitrogen monoxide is a pollutant found in smokestack emissions. How many liters of ammonia, NH3, at
STP are needed to produce 1.4 x 1023 molecules of H2O?
Homework / Practice
Solve the following problems. The reactions may not be balanced.
1.
If 20.0 g of magnesium react with excess hydrochloric acid, how many grams of magnesium chloride are
produced? Mg + HCl  MgCl2 + H2
2.
How many grams of chlorine gas must be reacted with excess sodium iodide if 10.0 g of sodium chloride
are needed? NaI + Cl2  NaCl + I2
3.
How many moles of oxygen gas are produced in the decomposition of 5.00 g of potassium chlorate?
KClO3  KCl + O2
4.
What mass of copper is required to replace silver from 4.00 g of silver nitrate dissolved in water?
Cu + AgNO3  Cu(NO3)2 + Ag
5.
If excess ammonium sulfate reacts with 20.0 g of calcium hydroxide, how many grams of ammonia (NH 3)
are produced? (NH4)2SO4 + Ca(OH)2  CaSO4 + NH3 + H2O
6.
If excess sulfuric acid reacts with 0.2564 moles of sodium chloride, how many grams of hydrochloric acid
are produced? H2SO4 + NaCl  HCl + Na2SO4
7.
How many grams of silver phosphate are produced if 10.0 g of silver acetate react with excess sodium
phosphate? AgC2H3O2 + Na3PO4  Ag3PO4 + NaC2H3O2
8.
How many moles of sodium hydroxide are needed to completely neutralize 25.0 g of sulfuric acid?
NaOH + H2SO4  Na2SO4 + H2O
9.
When calcium carbonate is heated strongly, carbon dioxide gas is released. What volume of carbon
dioxide, measured at STP, is produced if 15.2 g of calcium carbonate is heated? CaCO3  CaO + CO2
10. What volume of oxygen gas at STP is needed for complete combustion of 5.63 g of propane?
C3H8 + O2  CO2 + H2O
11. What volume of chlorine gas, measured at STP, is needed to produce 10.0 g of potassium permanganate
(KMnO4)? K2MnO4 + Cl2  KMnO4 + KCl
80
Unit 7- Stoichiometry
12. Suppose that you could decompose 0.250 mol of Ag2S into its elements.
a. How many moles of silver would form?
b. How many moles of sulfur would form from 38.8 g of silver sulfide?
13. Ammonia (NH3) is made industrially by reacting nitrogen gas and hydrogen gas under pressure, at high
temperature and in the presence of a catalyst. If 4.0 mol of hydrogen react, how many moles of ammonia
will be produced?
14. How many grams of sodium hydroxide can be produced from 500. g of calcium hydroxide according to the
equation: Ca(OH)2 + Na2CO3  2 NaOH + CaCO3?
15. How many liters of Cl2 can be produced from 5.60 mole HCl at STP? 4 HCl + O2  2 Cl2 + 2 H2O
16. Given the equation Al4C3 + 12 H2O  4 Al(OH)3 + 3 CH4
to react with 100 g Al4C3?
How many moles of water are needed
17. How many grams of zinc phosphate are formed when 10.0 g of Zn are reacted with the phosphoric acid?
The other product is hydrogen gas.
18. Given the equation 4 FeS2 + 11 O2  2 Fe2O3 + 8 SO2
with 4.50 mol of FeS2 at STP?
How many liters of O2 are required to react
19. Given the equation C2H4 + 3 O2  2 CO2 + 2 H2O
a. If 6.0 mol of CO2 are produced, how many moles of O2 were reacted?
b. How many liters of O2 are required for the complete reaction of 45 g of C2H4 at STP?
c. If 18.0 g of CO2 are produced, how many grams of H2O are produced?
Balance the following equations to use with questions 24 – 31:
20. ____ Al + ____ O2  ____ Al2O3
21. ____ Cu + ____ AgNO3  ____ Ag + ____ Cu(NO3)2
22. ____ Zn + ____ HCl  ____ ZnCl2 + ____ H2
23. ____ Fe + ____ Cl2  ____ FeCl3
Perform the following calculations:
24. Zinc reacts with hydrochloric acid to produce zinc chloride and hydrogen. How many moles of HCl are
required to produce 7.50 moles of ZnCl2?
25. Copper metal reacts with silver nitrate to form silver and copper (II) nitrate. How many grams of copper
are required to form 250 g of silver?
26. When aluminum is burned in excess oxygen, aluminum oxide is produced. How many grams of oxygen
are required to produce 0.75 moles of Al2O3?
27. How many grams of iron (III) chloride are produced when 15.3 g of iron react with excess chlorine gas?
81
Unit 7- Stoichiometry
28. Copper metal reacts with silver nitrate to form silver and copper (II) nitrate. How many moles of silver
will be produced from 3.65 moles of silver nitrate?
29. When 9.34 g of zinc react with excess hydrochloric acid how many grams of zinc chloride will be
produced?
30. How many grams of iron are needed to react with 31.0 L of chlorine gas at STP to produce iron (III)
chloride?
31. How many liters of oxygen gas at STP are required to react with 65.3 g of aluminum in the production of
aluminum oxide?
32. How many liters of ammonia gas are formed when 25 L of hydrogen at STP? N2 + H2  NH3
33. Given the following equation: Na2O + H2O ---> 2 NaOH How many grams of Na2O are required to
produce 1.60 x 102 grams of NaOH?
34. Given the following equation:
a.
b.
8 Fe + S8  8 FeS
What mass of iron is needed to react with 16.0 grams of sulfur?
How many grams of FeS are produced?
35. Given the following equation: 2 NaClO3  2 NaCl + 3 O2
a. 12.00 moles of NaClO3 will produce how many grams of O2?
b. How many grams of NaCl are produced when 80.0 grams of O2 are produced?
36. Given the following equation: Cu + 2 AgNO3 ---> Cu(NO3)2 + 2 Ag
a. How many moles of Cu are needed to react with 3.50 moles of AgNO 3?
b. If 89.5 grams of Ag were produced, how many grams of Cu reacted?
37. Molten iron and carbon monoxide are produced in a blast furnace by the reaction of iron (III) oxide and
coke (pure carbon). If 25.0 kilograms of pure Fe 2O3 is used, how many kilograms of iron can be produced?
The reaction is: Fe2O3 + 3 C ---> 2 Fe + 3 CO
38. The average human requires 120.0 grams of glucose (C6H12O6) per day. How many grams of CO2 (in the
photosynthesis reaction) are required for this amount of glucose? The photosynthetic reaction is:
6 CO2 + 6 H2O ---> C6H12O6 + 6 O2
39. Given the reaction: 4 NH3 (g) + 5 O2 (g) ---> 4 NO (g) + 6 H2O (l) When 1.20 mole of ammonia reacts, the
total number of moles of products formed is:
Stoichiometry Practice Test
Directions: Solve the following problems, showing all work.
1.
Balance the equation ______NaOH  ______Na2O + _____H2O
2.
How many moles of water are produced from 4 moles of sodium hydroxide?
3.
Balance the equation ____KCl + _____O2  _____KClO3
82
Unit 7- Stoichiometry
4.
How many moles of potassium chlorate are produced from 9 moles of oxygen?
5.
Fe2O3 + 2Al  2Fe + Al2O3 What mass of aluminum oxide is produced when 4 moles of aluminum
react?
6.
P4O10 + 6H2O  4H3PO4 How many grams of phosphoric acid are produced by the reaction of 12.5 g of
water?
7.
Using the reaction in question 6, how many grams of P 4H10 must react if the reaction produces 25 moles
H3PO4?
8.
Balance the equation _____HNO3 + ______Cu  _______Cu(NO3)2 + ______H2
9.
How many moles of nitric acid must react in order to form 83 g of copper nitrate?
10. Balance the equation ______NH3 + ________O2  _______NO + ______H2O
11. What mass of nitrogen monoxide will be formed when 7.2 g nitrogen trihydride react?
12. C2H4 + 3O2  2CO2 + 2H2O Determine the mass of water produced if 50 g C2H4 and 50 g O2 react.
Multiple Choice Practice
13. Consider the balanced equation Zn + 2HCl  ZnCl2 + H2 How many moles of ZnCl2 will be produced if 7
moles of HCl are used?
a. 2 moles
c. 3.5 moles
b. 2.5 moles
d. 4 moles
14. Given : C2H2(g) + 5O2(g)  4CO2(g) + 2H2O(g) Is this chemical equation balanced?
a. True
b. False
15. In the reaction below, how many moles of oxygen gas is produced by the decomposition of 4 moles of
mercury (II) oxide? 2HgO  2Hg + O2
a. 1 mole
c. 3 moles
b. 2 moles
d. 4 moles
16. True or False, 6 moles of H2 is needed to completely react with 2 moles of N2 in the balanced chemical
reaction N2 + 3H2  2NH3
a. True
b. False
17. If 18.0 grams of carbon are burned in 55.0 grams of oxygen, how many grams of carbon dioxide are
formed?
a. 44.01 grams CO2
c. 151 grams CO2
b. 75.6 grams CO2
d. 66.0 grams CO2
18. How many moles of Al2O3 are formed when a mixture of 0.36 moles Al and 0.36 moles O2 is ignited?
a. 0.12
c. 0.28
e. 0.72
b. 0.18
d. 0.46
83
Unit 7- Stoichiometry
19. A mass of 21.5 grams of calcium hydroxide reacts with an excess of phosphoric acid. What mass of
calcium phosphate could be recovered from solution?
a. 284 grams
c. 94.7 grams
e. 326 grams
b. 186 grams
d. 31.6 grams
20. If 3.64 g of calcium hydroxide react with excess sodium sulfate in aqueous solution to produce solid sulfate
and aqueous sodium hydroxide, how many moles of calcium atoms are reacting here?
Ca(OH)2(aq) + Na2SO4(aq)  CaSO4(s) + 2NaOH(aq)
a. 0.00982 mol
c. 0.0266 mol
e. 0.0909 mol
b. 0.0246 mol
d. 0.0491 mol
21. If one mole of the rocket fuel ammonium perchlorate, NH 4ClO4 (s) is allowed to react with excess Al so all
of the NH4ClO4 is consumed, how many molecules of water will be produced?
3NH4ClO4 (s) + 3Al (s)  Al2O3 (s) + AlCl3 (s) + 3NO (g) + 6H2O (g)
a. 3.61 x 1023
c. 6.02 x 1023
e. 3.01 x 1024
23
24
b. 1.0 x 10
d. 1.20 x 10
22. How many grams of potassium cyanide, PCl3, is produced from 93.0 grams of P 4 (s) and 213 g of Cl2 (g),
assuming the reaction goes to completion? The balanced equation for the reaction is:
P4 (s) + 6Cl2 (g)  4PCl3 (g)
a. 277 g
c. 213 g
e. 69.3 g
b. 416 g
d. 104 g
23. In the oxidation of ethane: 2 C2H6 + 7 O2 → 4 CO2 + 6 H2O how many moles of O2 are required to react
with 1 mole of ethane?
a. 7 moles
b. 2 moles
c. 3.5 moles
24. In the reaction: 2 C2H6 + 7 O2 → 4 CO2 + 6 H2O how many moles of CO2 are formed when 1 mole of O2
is consumed?
a. 7 moles
b. 1.75 moles
c. 0.57 moles
25. In the reaction: 2 C2H6 + 7 O2 → 4 CO2 + 6 H2O how many moles of CO2 are formed when 5moles of
ethane are consumed?
a. 10 moles
b. 4 moles
c. 2 moles
26. How many mL of 0.1 M HCl are required to react with 0.01 mole of Na 2CO3?
HCl + Na2CO3  NaCl + H2CO3
a. 100 mL
b. 200 mL
27. How many liters of H2 at STP are required to react with 2.3 g of Fe3O4?
a. 0.22 L
b. 0.44 L
c.
50 mL
H2 + Fe3O4  FeO + H2O
c. 0.56 L
28. When 0.05 mole H2 is mixed with 0.05 mole CO, what is the maximum number of moles of methanol
(CH3OH) that can be obtained? H2 + CO  CH3OH
a. 0.10 mole
b. 0.05 mole
c. 0.025 mole
84
Unit 8- Gas Laws
THE GAS LAWS
The gas laws describe how gases behave. They can be predicted by theory and the amount of change can be
calculated with mathematical equations. One ____________________________ is equal to 760 mm Hg, 760 torr,
or _______________ kPa (kilopascals).

Perform the following pressure conversions.
a) 144 kPa = ______________ atm
b) 795 mm Hg = ______________ atm
c) 669 torr = ______________ kPa
d) 1.05 atm = ______________ mm Hg
Air pressure at higher altitudes, such as on a mountaintop, is slightly ______________________ than air pressure at
sea level. Air pressure is measured using a ________________________. More molecules mean more
____________________ between the gas molecules themselves and more collisions between the gas molecules and
the walls of the container. Number of molecules is ________________________ proportional to pressure.
Doubling the number of gas particles in a basketball _____________________ the pressure. Gases naturally move
from areas of high pressure to ____________ pressure because there is empty space to move in. If you double the
number of molecules, you _____________________ the pressure. As you remove molecules from a container, the
pressure ________________________ until the pressure inside equals the pressure outside. In a smaller container,
molecules have less room to move. The molecules hit the sides of the container _________________ often, striking
a smaller area with the same force. As volume decreases, pressure increases. Volume and pressure are
______________________ proportional. As the pressure on a gas increases, the volume decreases. Raising the
temperature of a gas increases the _______________________ if the volume is held constant. At higher
temperatures, the particles in a gas have greater ________________________ energy. They move faster and collide
with the walls of the container more often and with greater ___________________, so the pressure rises. If you
start with 1 liter of gas at 1 atm pressure and 300 K and heat it to 600 K, one of 2 things happens. Either the
volume will increase to
2 liters at ______ atm, or the pressure will increase to ______ atm while the volume
remains constant.
Ideal Gases and the Kinetic Molecular Theory
In this unit we will assume the gases behave ideally. _____________________ gases do not really exist, but this
makes the math easier and is a close approximation. Gas particles are much smaller than the spaces between them.
The particles have negligible _____________________. There are no attractive or repulsive
___________________ between gas molecules. Gas particles are in constant, _________________________
motion. Until they bump into something (another particle or the side of a container), particles move in a straight line.
No kinetic energy is ____________________ when gas particles collide with each other or with the walls of their
container. All gases have the same ______________________ energy at a given temperature. Temperature is a
measure of the average kinetic energy of the particles in a sample of matter. There are no gases for which this is
true. Real gases behave more ideally at ________________ temperature and _________________ pressure. At low
temperature, the gas molecules move more _______________________, so attractive forces are no longer
85
Unit 8- Gas Laws
negligible. As the pressure on a gas increases, the molecules are forced closer together and
_________________________ forces are no longer negligible. Therefore, real gases behave more ideally at high
temperature and low pressure.
Avogadro’s Law
Avogadro’s law states that equal volumes of different gases (at the same temperature and pressure) contain equal
numbers of ________________ or molecules. 2 liters of helium has the same number of particles as ______ liters
of oxygen. The molar volume for a gas is the volume that one mole occupies at 0.00ºC and 1.00 atm.
1 mole = 22.4 L at STP (standard temperature and pressure). As a result, the volume of gaseous reactants and
products can be expressed as small whole numbers in reactions.

How many moles are in 45.0 L of a gas at STP?

How many liters are in 0.636 moles of a gas at STP?
The volume of a gas is directly proportional to the number of moles.
V1 V2

n1 n2

Consider two samples of nitrogen gas. Sample 1 contains 1.5 mol and has a volume of 36.7 L. Sample 2
has a volume of 16.5 L at the same temperature and pressure. Calculate the number of moles of nitrogen in
sample 2.

If 0.214 mol of argon gas occupies a volume of 652 mL at a particular temperature and pressure, what
volume would 0.375 mol of argon occupy under the same conditions?

If 46.2 g of oxygen gas occupies a volume of 100. L at a particular temperature and pressure, what volume
would 5.00 g of oxygen gas occupy under the same conditions?
Boyle’s Law
At Boyle’s law states that the pressure and volume of a gas at constant temperature are inversely
proportional. Inversely proportional means as one goes up the other goes ________________.
P 1 V1 = P 2 V2

Sketch the PV graph that represents Boyle’s law.

The P-V graph for Boyle’s law results in a _____________________________ because pressure and
volume are inversely proportional.

A balloon is filled with 25 L of air at 1.0 atm pressure. If the pressure is changed to 1.5 atm, what is the
new volume? (Make sure the pressure and volume units in the question match.)

A balloon is filled with 73 L of air at 1.3 atm pressure. What pressure is needed to change the volume to
43 L?

A gas is collected in a 242 cm3 container. The pressure of the gas in the container is measured and
determined to be 87.6 kPa. What is the volume of this gas at standard pressure?
86
Unit 8- Gas Laws

A gas is collected in a 24.2 L container. The pressure of the gas in the container is determined to be 756
mm Hg. What is the pressure of this gas if the volume increases to 30.0 L?
Charles’ Law
The volume of a gas is directly proportional to the Kelvin temperature if the pressure is held constant.
V1 V2

T1 T2
K = °C + 273

Sketch the PV graph that represents Charles’ law.
The V-T graph for Charles’ law results in a _____________________________ _________________ because
pressure and volume are directly proportional.

What is the temperature of a gas that is expanded from 2.5 L at 25 ºC to 4.1 L at constant pressure? (Make
sure the volume units in the question match and make sure to convert degrees Celsius to Kelvin.)

What is the final volume of a gas that starts at 8.3 L and 17 ºC and is heated to 96 ºC?

A 225 cm3 volume of gas is collected at 57 ºC. What volume would this sample of gas occupy at standard
temperature?

A 225 cm3 volume of gas is collected at 42 ºC. If the volume is decreased to 115 cm3, what is the new
temperature?
Gay-Lussac’s Law
The temperature and the pressure of a gas are directly related at constant volume.
P1 P2

T1 T2

Sketch the PT graph that represents Gay-Lussac’s law.

What is the pressure inside a 0.250 L can of deodorant that starts at 25 ºC and 1.2 atm if the temperature is
raised to 100 ºC? Volume remains constant. (Make sure the pressure units in the question match and make
sure to convert degrees Celsius to Kelvin.)

A can of deodorant starts at 43 ºC and 1.2 atm. If the volume remains constant, at what temperature will
the can have a pressure of 2.2 atm?

A can of shaving cream starts at 25 ºC and 1.30 atm. If the temperature increases to 37 ºC and the volume
stays constant, what is the pressure of the can?

A 12 ounce can of a soft drink starts at STP. If the volume remains constant, at what temperature will the
can have a pressure of 2.20 atm?
87
Unit 8- Gas Laws
The Combined Gas Law
The gas laws may be combined into a single law, called the combined gas law, which relates two sets of conditions
of pressure, volume, and temperature by the following equation.
P1V1 P2V2

T1
T2

A 15 L cylinder of gas at 4.8 atm pressure at 25 ºC is heated to 75 ºC and compressed to 17 atm. What is
the new volume?

If 6.2 L of gas at 723 mm Hg at 21 ºC is compressed to 2.2 L at 4117 mm Hg, what is the temperature of
the gas?

A sample of nitrogen monoxide has a volume of 72.6 mL at a temperature of 16 °C and a pressure of
104.1 kPa. What volume will the sample occupy at 24 °C and 99.3 kPa?

A hot air balloon rises to an altitude of 7000 m. At that height the atmospheric pressure drops to
300 mm Hg and the temperature cools to -33 °C. Suppose on the hot air balloon there was a small balloon
filled to 1.00 L at sea level and a temperature of 27 °C. What would its volume ultimately be when it
reached the height of 7000 m?
Dalton’s Law of Partial Pressures
Dalton’s law of partial pressures states that the _________________ pressure of a mixture of gases is equal to the
sum of the pressures of all the gases in the mixture, as shown below.
Pt = P1 + P 2 + P3 + …
Pt = total pressure
The partial pressure is the contribution by that gas.

What is the total pressure in a balloon filled with air if the pressure of the oxygen is 170 mm Hg and the
pressure of nitrogen is 620 mm Hg?

In a second balloon the total pressure is 1.30 atm. What is the pressure of oxygen (in mm Hg) if the
pressure of nitrogen is 720 mm Hg?

A container has a total pressure of 846 torr and contains carbon dioxide gas and nitrogen gas. What is the
pressure of carbon dioxide (in kPa) if the pressure of nitrogen is 50 kPa?

When a container is filled with 3 moles of H2, 2 moles of O2 and 4 moles of N2, the pressure in the
container is 8.7 atm. The partial pressure of H2 is _____.
It is common to synthesize gases and collect them by displacing a volume of ________________.

Hydrogen was collected over water at 21°C on a day when the atmospheric pressure is 748 torr. The
volume of the gas sample collected was 300 mL. The vapor pressure of water at 21°C is 18.65 torr.
Determine the partial pressure of the dry gas.

A sample of oxygen gas is saturated with water vapor at 27ºC. The total pressure of the mixture is
772 mm Hg and the vapor pressure of water is 26.7 mm Hg at 27ºC. What is the partial pressure of the
oxygen gas?
88
Unit 8- Gas Laws
The Ideal Gas Law
Remember ideal gases do not exist. Molecules do take up ______________________. There are
_________________________ forces; otherwise, there would be no liquids.
PV = nRT
Pressure times volume equals the number of ___________________ (n) times the ideal gas constant (R) times the
temperature in Kelvin.

R = 0.0821 (L atm)/(mol K)
or R = 8.314 (L kPa)/(mol K)
or R = 62.4 (L mm Hg)/(mol K)
The one you choose depends on the unit for pressure!

How many moles of air are there in a 2.0 L bottle at 19 ºC and 747 mm Hg?

What is the pressure in atm exerted by 1.8 g of H 2 gas exerted in a 4.3 L balloon at 27 ºC?

Sulfur hexafluoride (SF6) is a colorless, odorless and very unreactive gas. Calculate the pressure (in atm)
exerted by 1.82 moles of the gas in a steel vessel of volume 5.43 L at 69.5 ºC.

Calculate the volume (in liters) occupied by 7.40 g of CO 2 at STP.

A sample of nitrogen gas kept in a container of volume 2.30 L and at a temperature of 32 ºC exerts a
pressure of 476 kPa. Calculate the number of moles of gas present.

A 1.30 L sample of a gas has a mass of 1.82 g at STP. What is the molar mass of the gas?

Calculate the mass of nitrogen gas that can occupy 1.00 L at STP.
Homework / Practice
1.
Identify whether the descriptions below describe an ideal gas or a real gas.
a) Gas particles move in straight lines until they collide with other particles or the walls of their
container.
b) Individual gas particles have a measurable volume.
c) The gas will not condense even when compressed or cooled.
d) Collisions between molecules are perfectly elastic.
e) Gas particles passing close to one another exert an attraction on each other.
2.
The formula for kinetic energy is KE = 1/2mv2.
a) What happens to KE if the mass is tripled (at constant speed)?
b) What happens to KE if the speed is halved (at constant mass)?
c) If two gases at the same temperature share the same KE, it follows that the molecules of greater
mass have the _____ speed. (faster or slower)
3.
Explain the following using the kinetic-molecular theory:
a) As a gas is heated, its rate of effusion through a small hole increases if all other factors remain
constant.
b) A strong-smelling gas released from a container in the middle of a room is soon detected in all
areas of the room.
4.
Pressure = force / area. For a constant force, when the surface area is quadrupled, the pressure
5.
Rank the following in increasing order.
a) 50 kPa
6.
b) 2 atm
Does atmospheric pressure increase or decrease as altitude above sea level increases?
89
c)
76 torr
Unit 8- Gas Laws
7.
Convert the following:
a. 0.200 atm = _____ mm Hg
b. 790 mm Hg = _____ Pa
c.
d.
123 kPa = _____ atm
0.935 atm = ______ torr
8.
The mercury level in an open manometer is 25 mm Hg higher in the arm connected to the atmosphere. If
atmospheric pressure is 765 mm Hg, calculate the pressure of the enclosed gas.
9.
The mercury level in an open manometer is 32 mm Hg lower in the arm connected to the atmosphere. If
atmospheric pressure is 765 mm Hg, calculate the pressure of the enclosed gas.
10. A 24 L sample of a gas (at fixed mass and constant temperature) exerts a pressure of 3.0 atm. What
pressure will the gas exert if the volume is changed to 16 L?
11. An air bubble with a volume of 2.0 mL forms at the bottom of a lake, where the pressure is 3.0 atm. As it
rises, the pressure on the bubble decreases. Assume the temperature remains constant.
a. Will the bubble expand or contract as it rises?
b. Predict the volume of the bubble just as it reaches the surface, where the pressure is 1.0 atm.
12. A common laboratory system to study Boyle’s law uses a gas trapped in a syringe. The pressure in the
system is changed by adding or removing identical weights on the plunger. The original gas volume is
50.0 mL when two weights are present. Predict the new gas volume when 4 more weights are added.
13. Helium gas in a balloon occupies 2.40 L at 400. K. What volume will it occupy at 300 K?
14. A bicycle tire is inflated to 55 lb/in2 at 15 °C. Assume that the volume of the tires does not change
appreciably once it is inflated.
a. The tire and the air inside it are heated to 30 °C by road wear, does the pressure in the tire increase
or decrease?
b. Because the temperature has doubled, does the pressure double to 110 psi? Why or why not?
15. If 0.864 mol of neon gas occupies a volume of 452 mL at a particular temperature and pressure, what
volume would 0.575 mol of neon occupy under the same conditions?
16. If 26.5 g of oxygen gas occupies a volume of 100. L at a particular temperature and pressure, how many
moles of oxygen gas will there be in 350. L under the same conditions?
17. At one point in the cycle of a piston in an automobile engine, the volume of the trapped fuel mixture is
400 cm3 at a pressure of 1.0 atm and a temperature of 27 °C. In the compression of the piston, the
temperature reaches 77 °C and the volume decreases to 50.0 cm3. What is the new pressure?
18. On a cold winter morning when the temperature is - 13 °C, the air pressure in an automobile tire is 1.5 atm.
If the volume does not change, what will the pressure be after the tire has warmed to 13 °C?
19. A gas storage tank has a volume of 3.5 x10 5 m3 when the temperature is 27 °C and the pressure is 1.0 atm.
What is the new volume of the tank if the temperature drops to - 10.0°C and the pressure drops to
0.95 atm?
20. Explain how to correct for the partial pressure of water vapor when calculating the partial pressure of a dry
gas that is collected over water.
21. A sample of chlorine gas is collected by water displacement at 23 °C. If the atmospheric pressure is
751 torr, what is the partial pressure due to the chlorine? The vapor pressure of water at 23 °C is
21.1 mm Hg.
90
Unit 8- Gas Laws
22. When an explosive like TNT is detonated, a mixture of gases at high temperature is created. Suppose that
gas X has a pressure of 50 atm, gas Y has a pressure of 20 atm, and gas Z has a pressure of 10 atm.
a. What is the total pressure of the system?
b. Once the gas mixture combines with air, Ptotal soon drops to 2 atm. By what factor does the
volume of the mixture increases? (Assume mass and temperature are constant.)
23. A gas occupies a volume of 180 mL at 35.0 °C and 740 mm Hg. What is the volume of the gas at STP?
24. Perform the following calculations
a. How many moles of methane, CH4, are present in 5.6 L of the gas at STP?
b. How many moles of gas are present in 5.6 L of any ideal gas at STP?
c. What is the mass of the 5.6 L sample of methane gas?
25. What is the pressure exerted by 32 g of oxygen gas in a 20. L container at 30.0 °C?
26. How many grams of nitrogen gas are in a flask with a volume of 250 mL at a pressure of 3.0 atm and a
temperature of 300. K?
27. A container holds three gases: oxygen, carbon dioxide, and helium. The partial pressures of the three gases
are 2.00 atm, 3.00 atm, and 4.00 atm, respectively. What is the total pressure inside the container?
28. A gas occupies 12.3 liters at a pressure of 40.0 mm Hg. What is the volume when the pressure is increased
to 60.0 mm Hg?
29. If a gas at 25.0 °C occupies 3.60 liters at a pressure of 1.00 atm, what will be its volume at a pressure of
2.50 atm?
30. A gas occupies 1.56 L at 1.00 atm. What will be the volume of this gas if the pressure becomes 3.00 atm?
31. A gas occupies 11.2 liters at 0.860 atm. What is the pressure if the volume becomes 15.0 L?
32. How much will the volume of 75.0 mL of neon change if the pressure is lowered from 50.0 torr to
8.00 torr?
33. Calculate the decrease in temperature when 2.00 L at 20.0 °C is compressed to 1.00 L.
34. 600.0 mL of air is at 20.0 °C. What is the volume at 60.0 °C?
35. A gas occupies 900.0 mL at a temperature of 27.0 °C. What is the volume at 132.0 °C?
36. What change in volume results if 60.0 mL of gas is cooled from 33.0 °C to 5.00 °C?
37. A gas occupies 1.00 L at standard temperature. What is the volume at 333.0 °C?
38. Determine the pressure change when a constant volume of gas at 1.00 atm is heated from 20.0°C to
30.0 °C.
39. A gas has a pressure of 0.370 atm at 50.0 °C. What is the pressure at standard temperature?
40. A gas has a pressure of 699.0 mm Hg at 40.0 °C. What is the temperature at standard pressure?
41. If a gas is cooled from 323.0 K to 273.15 K and the volume is kept constant what final pressure would
result if the original pressure was 750.0 mm Hg?
91
Unit 8- Gas Laws
42. If a gas in a closed container is pressurized from 15.0 atmospheres to 16.0 atmospheres and its original
temperature was 25.0 °C, what would the final temperature of the gas be?
43. A gas has a volume of 800.0 mL at negative 23.00 °C and 300.0 torr. What would the volume of the gas be
at 227.0 °C and 600.0 torr of pressure?
44. 500.0 liters of a gas are prepared at 700.0 mm Hg and 200.0 °C. The gas is placed into a tank under high
pressure. When the tank cools to 20.0 °C, the pressure of the gas is 30.0 atm. What is the volume of the
gas?
45. What is the volume of gas at 2.00 atm and 200.0 K if its original volume was 300.0 L at 0.250 atm and
400.0 K.
46. At conditions of 785.0 torr of pressure and 15.0 °C temperature, a gas occupies a volume of 45.5 mL.
What will be the volume of the same gas at 745.0 torr and 30.0 °C?
47. A gas occupies a volume of 34.2 mL at a temperature of 15.0 °C and a pressure of 800.0 torr. What will be
the volume of this gas at standard conditions?
48. The volume of a gas originally at standard temperature and pressure was recorded as 488.8 mL. What
volume would the same gas occupy when subjected to a pressure of 100.0 atm and temperature of minus
245.0 °C?
49. At a pressure of 780.0 mm Hg and 24.2 °C, a certain gas has a volume of 350.0 mL. What will be the
volume of this gas under STP?
50. A gas sample occupies 3.25 liters at 24.5 °C and 1825 mm Hg. Determine the temperature at which the gas
will occupy 4250 mL at 1.50 atm.
51. If 10.0 liters of oxygen at STP are heated to 512 °C, what will be the new volume of gas if the pressure is
also increased to 1520.0 mm Hg?
52. If 2.00 liters of hydrogen, originally at 25.0 °C and 750.0 mmHg, are heated until a volume of 20.0 liters
and a pressure of 3.50 atm is reached. What is the new temperature?
53. How many moles of gas are contained in 890.0 mL at 21.0 °C and 750.0 mm Hg pressure?
54. Calculate the volume 3.00 moles of a gas will occupy at 24.0 °C and 762.4 mm Hg.
55. What volume will 20.0 g of argon occupy at STP?
56. How many moles of gas would be present in a gas trapped within a 100.0 mL vessel at 25.0 °C at a
pressure of 2.50 atmospheres?
57. How many moles of a gas would be present in a gas trapped within a 37.0 liter vessel at 80.00°C at a
pressure of 2.50 atm?
58. At what pressure would 0.150 mole of nitrogen gas at 23.0 °C occupy 8.90 L?
59. What volume would 32.0 g of NO2 gas occupy at 3.12 atm and 18.0 °C?
60. Find the volume of 2.40 mol of gas whose temperature is 50.0 °C and whose pressure is 2.00 atm.
61. How many moles of gas are contained in a 50.0 L cylinder at a pressure of 100.0 atm and a temperature of
35.0 °C?
92
Unit 8- Gas Laws
62. Determine the number of moles of Krypton contained in a 3.25 liter gas tank at 5.80 atm and 25.5 °C. If
the gas is Oxygen instead of Krypton, will the answer be the same? Why or why not?
63. A 500.0 mL sample of a gas is collected at 745.0 mm Hg. What will the volume be at standard pressure?
64. Convert 350.0 mL at 740.0 mm of Hg to its new volume at standard pressure.
65. The temperature of a sample of gas in a steel container at 30.0 kPa is increased from -100.0 °C to
1.00 x 103 °C. What is the final pressure inside the tank?
Gas Laws Practice Test
1.
Explain how the temperature is related to the kinetic energy and motion of gas particles.
2.
If the volume of a gas contained within balloon were to be tripled, what would be the impact upon the pressure
if Kelvin temperature is maintained as constant?
Directions: Solve the following problems. Show all your work, including units
3.
In a mixture of carbon dioxide, oxygen gas, sulfur dioxide and carbon monoxide, the pressure of the carbon
dioxide is 0.3 atm, oxygen gas is 0.5 atm, sulfur dioxide is 0.6 atm, and the pressure of the carbon monoxide is
0.1 atm. What is the total pressure in the container?
4.
A high-altitude balloon contains 4 Liters of helium gas at 1.35 atm. What is the volume when the balloon rises
to an altitude where the pressure is only 1.20 atm? (Assume that the temperature remains constant.)
5.
If a sample of gas occupies 27 Liters at 12 Celsius, what will be its volume at 112 Celsius if the pressure does
not change?
6.
A gas has a pressure of 122 kPa at -6 Celsius (negative 6). What will be the pressure at 85 Celsius if the
volume does not change?
7.
A gas at 10 kPa and 45 Celsius occupies a container with an initial volume of 4 Liters. By changing the
volume, the pressure of the gas increases to 25 kPa as the temperature is raised to 190 Celsius. What is the new
volume?
8.
You fill a rigid steel cylinder that has a volume of 840 milliliters with oxygen gas to a final pressure of
1.1 atmospheres at 145 Celsius. How many moles of nitrogen gas does the cylinder contain?
9.
What is the temperature when 4 moles of carbon dioxide occupies a 2 L container and exerts a pressure of
745 torr?
10. What pressure, in atm, will be exerted by 1.25 moles of a gas at 39 Kelvin if it is contained in a 5 Liter vessel?
11. What volume will 29 grams of nitrogen gas occupy at 10 Celsius and a pressure of 620 torr?
12. A 35mL sample of hydrogen gas is collected over water at a temperature of 24 oC, the vapor pressure of the
water at that temperature is 2.99 kPa, and the atmospheric pressure is 765.5 torr. What is the pressure of the dry
hydrogen gas?
93
Unit 8- Gas Laws
Multiple Choice Practice
13. As the pressure of a gas at 2 atm is changed to 1 atm at constant temperature, the volume of the gas
a. decreases
b. increases
c. remains the same
14. According to the kinetic molecular theory, molecules increase in kinetic energy when they
a. Are mixed with other molecules at lower temperature
b. Are frozen into a solid
c. Are condensed into a liquid
d. Are heated to a higher temperature
15. Collide with each other in a container at lower temperature At STP, 32.0 liters of O 2 contain the same number
of molecules as
a. 22.4 L Ar
c. 32. 0 L of H2
b. 28.0 L of N2
d. 44.8 L of He
16. What is the total pressure, in atmospheres, of a 10.0 L container that contains 10 moles of nitrogen gas and 10
moles of oxygen gas at 300 K?
a. 24.6 L
c.
2460 L
b. 49.3 L
d.
4930 L
17. An 8.25 L sample of oxygen is collected at 25°C and 1.022 atm pressure. What volume will the gas occupy
0.940 atm and -15°C?
a. 1.78 L
c. 10.4 L
e. 7.77 L
b. 5.00 L
d. 8.76 L
18. A motorist fills his car tires to 32 lb/in2 pressure at a temperature of 30°C. Assuming no change in volume,
what will be the pressure in the tires when the motorist drives across Death Valley, with a pavement
temperature of 78°C?
a. 12 lb/in2
c. 37 lb/in2
e. 83 lb/in2
b. 28 lb/in2
d. 4.8 lb/in2
19. The mass of 2.37 liters of a gas is 8.91 grams. What is the density of the gas?
a. 3.76 g/L
c. None of these
b. 6.54 g/L
d. 0.266 g/L
e.
21.1 g/L
20. If temperature is constant, the relationship between pressure and volume is
a. Direct
b. inverse
21. A 268 cm3 sample of an ideal gas at 18°C and 748 torr pressure is placed in an evacuated container of volume
648cm3. To what centigrade temperature must the assembly be heated so that the gas will fill the whole
chamber at 748 torr?
a. 431°C
c. 704°C
e. 324°C
b. 120°C
d. 597°C
94
Unit 8- Gas Laws
22. How big a volume of dry oxygen gas at STP would you need to take to get the same number of oxygen
molecules as there are hydrogen molecules in 25.0 liters at 0.850 atm and 35°C
a. 18.8 L
c. 0.656 L
e. 32.3 L
b. 0.068 L
d. 4.2 L
23. Nitrogen has a molar mass of 28.02 g/mol. What is the density of nitrogen at 1.05 atm and 37°C?
a. None of these
c. 0.89 g/L
e. 4.72 g/cm3
b. 2.82 g/L
d. 1.25 g/L
24. How many moles of gas would it take to fill an average man's lungs, total capacity of which is about 4.5 liters?
Assume 1.00 atm pressure and 37.0°C.
a. 37.0 mol
c. 0.75 mol
e. 11.2 mol
b. 1.24 mol
d. 0.18 mole
25. Which flask contains the greatest number of molecules?
a.
b.
c.
Flask 3 (O2)
Flask 1 (NH3)
Flask 2 (CH4)
d.
e.
Flasks 2 and 3
All are the same
26. You have two samples of the same gas in the same size container, at the same pressure. The temperature in the
first container is -23°C and the temperature in the second container is 227°C. What is the ratio of the number of
moles of gas in the first container compared to that in the second container?
a. 2:1
c. 1:2
e. 4:1
b. 1:4
d. 1:1
27. If pressure is constant, the relationship between temperature and volume is
a. Direct
b. Inverse
28. If pressure of a gas is increased and its volume remains constant, what will happen to its temperature?
a. Increase
b. Decrease
c. Stay the same
29. One way to increase pressure on a gas is to
a. decrease temperature
b. increase volume
c.
d.
increase the number of gas particles
lower the kinetic energy of the gas molecules
30. How do gas particles respond to an increase in volume?
a. increase in kinetic energy and decrease in temperature
b. decrease in kinetic energy and decrease in pressure
c. increase in temperature and increase in pressure
d. increase in kinetic energy and increase in temperature
95
Unit 8- Gas Laws
31. If a gases volume is decreased and pressure is constant, its temperature will
a. Increase
b. Decrease
c.
Stay the same
32. If the temperature of a gas remains constant but pressure is decreased, the volume will
a. Increase
b. Decrease
c.
Stay the same
33. Convert 2.3 atm into mmHg
a. 2300 mmHg
b. 1750 mmHg
c.
d.
2.3 mmHg
0.0030 mmHg
34. The pressure of a gas is 750.0 torr when its volume is 400.0 mL. Calculate the pressure (in atm) if the gas is
allowed to expand to 600.0 mL at constant temperature.
a. 0.660 atm
c. 500.0 atm
b. 1.48 atm
d. 1125 atm
35. The volume of a gas is increased from 150.0 mL to 350.0 mL by heating it. If the original temperature of the
gas was 25.0 °C, what will its final temperature be (in °C)?
a. - 146°C
c. 58.3°C
e. 695°C
b. 10.7°C
d. 422°C
36. Standard temperature and pressure (STP) refers to which conditions?
a. 0 oC and 1 kPa
d.
b. 0 oC and 1 mm Hg
e.
c. 0 K and 1 kPa
0 K and 1 atm
273 K and 1 atm
37. If 4 moles of a gas are added to a container that already holds 1 mole of gas, how will the pressure change
within the container? (Assume volume and temperature are constant.)
a. The pressure will be 5 times as great.
b. The pressure will be 2 times as great.
c. The pressure will be 4 times as great.
d. The pressure will not change.
e. None of the above are correct.
38. A 4.0 L sample of hydrogen gas at 700 mm Hg would occupy what volume at 250 mm Hg? (Assume
temperature and number of particles stays constant.)
a. 1.4 x 10 -7 L
c. 11.2 L
e. 7.0 x 10 5 L
b. 1.4 L
d. 2.4 L
39. A 25 L tank of oxygen under a pressure of 80. atm would require what pressure to decrease the volume to
1.0 L? (Assume temperature and number of particles stays constant.)
a. 0.31 atm
b. 3.2 atm
c. 2000 atm
d. There is not enough information to answer the question.
e. None of these is correct.
40. A balloon containing 2.50 L of gas at 1 atm would be what volume at a pressure of 300 kPa? (Assume
temperature and number of particles stays constant.)
a. 6.33 L
c. 0.844 L
e. 000833 L
b. 8.11 L
d. 120. L
96
Unit 8- Gas Laws
41. A syringe containing 75.0 mL of air is at 298 K. What will the volume of the syringe be if it is placed in a
boiling water bath (373 K). Assume pressure and the number of particles are held constant.
a. 59.9 mL
b. 188 mL
c. 300. mL
d. 8.34 x 106 mL
e. None of the above are correct.
42. A gas occupies 40.0 mL at 127 oC. What volume will it occupy at -73 oC? (Assume pressure and number of
particles is constant.)
a. 182 mL
b. 8.80 mL
c. 80.0 mL
d. 20.0 mL
e. None of these is correct
43. If 88.0 grams of solid carbon dioxide evaporates, how many liters of CO 2 gas will be formed at a temperature of
300 K and 2.00 atmospheres of pressure?
a. 98.5 liters
c. 24.6 liters
b. 2170 liters
d. 1080 liters
44. Which of the following equations correctly combines Boyle's and Charles' Laws?
a. P1V1  P2V2
P1T1 P2T2
d.

b. T1V1  T2V2
V1
V2
c.
P1V1 P2V2

T1
T2
e.
T1V1 T2V2

P1
P2
45. A 50.0 mL sample of a gas is at 3.00 atm of pressure and a temperature of 298 K. What volume would the gas
occupy at STP?
a. 0.00728 mL
c. 18.2 mL
e. None of these is
b. 15.3 mL
d. 137 mL
correct.
46. A syringe contains 60.0 mL of air at 740 mm Hg pressure and 20 oC. What would be the temperature at which
the syringe would contain 30.0 mL at a pressure of 370 mm Hg? (Assume no gas could leak in or out of the
syringe.)
a. -200 oC
c. 5 oC
e. None of these is
o
b. 0.0137 C
d. 73.3 oC
correct
47. A sealed container contains 1.0 mol of hydrogen and 2.0 moles of nitrogen gas. If the total pressure in the
container is 1.5 atm, what is the amount of pressure exerted by each gas?
a. H2 = 1.0 atm and N2 = 0.50 atm
c. H2 = 1.0 atm and N2 = 2.0 atm
b. H2 = 0.50 atm and N2 = 1.0 atm
d. H2 = 2.0 atm and N2 = 1.0 atm
e. There is not enough information given to answer the question.
48. A sample of gas is collected by water displacement. The atmospheric pressure in the room is 757 mm Hg and
the vapor pressure of water is 17 mm Hg. What is the partial pressure of hydrogen under these conditions?
a. 17 mm Hg
c. 757 mm Hg
b. 740 mm Hg
d. 774 mm Hg
e. You cannot answer this question because you do not know the temperature.
97
Unit 9- Solids, Liquids and Phase Changes
SOLIDS AND LIQUIDS
States of Matter
There are ______ states of matter. A solid is a form of matter that has its own definite _____________ and volume.
A solid cannot _________________. The particles can vibrate but cannot move around. The particles of matter in a
solid are very tightly ____________________; when heated, a solid expands, but only slightly. A liquid is a form of
matter that flows, has ____________________ (definite) volume, and takes the _________________ of its
container. The particles in a liquid are not rigidly held in place and are _______________ closely packed than are
the particles in a solid; liquid particles are able to move past each other. A liquid is not very
__________________________. Like solids, liquids tend to expand when heated. A gas is a form of matter that
flows to conform to the ____________________ of its container and fills the entire _______________________ of
its container. Compared to solids and liquids, the particles of gases are very far apart. Because of the significant
amount of space between particles, gases are easily compressed. _____________________ is composed of
electrons and positive ions at temperatures greater than ____________ °C. The sun and other stars are examples of
plasma.

Identify the following as a property of a solid, liquid or gas. The answer may include more that one state of
matter.
1. flows and takes the shape of a container
2. compressible
3. made of particles held in a specific arrangement
4. has definite volume
5. always occupies the entire space of its container
6. has a definite volume but flows
The word_____________________ refers to the gaseous state of a substance that is a solid or a liquid at room
temperature. For example, steam is a vapor because at room temperature water exists as a liquid. Some substances
are described as _______________________, which means that they change to a gas easily at room temperature.
Alcohol and gasoline are ______________ volatile than water. Kinetic-molecular theory predicts the constant
motion of the liquid particles. Individual liquid molecules do not have fixed positions in the liquid. However,
forces of ________________________ between liquid particles limit their range of motion so that the particles
remain closely packed in a fixed volume. These attractive forces are called ___________________________
forces. Inter = between. Molecular = molecules. A liquid diffuses more _______________________ than a gas at
the same temperature, however, because intermolecular attractions interfere with the flow.
__________________________ is a measure of the resistance of a liquid to flow. Viscosity decreases with
________________________ temperature. Particles in the middle of the liquid can be attracted to particles above
them, below them, and to either side. For particles at the surface of the liquid, there are no attractions from above to
balance the attractions from _______________. Thus, there is a net attractive force pulling down on particles at the
surface. _____________________ ____________________ is a measure of the inward pull by particles in the
98
Unit 9- Solids, Liquids and Phase Changes
interior. Soaps and detergents decrease the surface tension of water by disrupting the _______________________
bonds between water molecules. For a substance to be a solid rather than a liquid at a given temperature, there must
be strong attractive forces acting between particles in the solid. These forces limit the motion of the particles to
__________________________ around fixed locations in the solid. Thus, there is more order in a solid than in a
liquid. The particles can only vibrate and revolve in place. Because of this order, solids are much less
_________________ than liquids and gases. In fact, solids are not classified as fluids. Most solids are more
_________________ than most liquids. A crystalline solid is a solid whose atoms, ions, or molecules are arranged
in an orderly, geometric, three-dimensional structure. Most solids are _____________________. Amorphous solids
lack an orderly internal structure. Think of them as __________________________ liquids. Examples of
amorphous solids include ____________________ and glass.
Phase Changes
If a substance is usually a liquid at room temperature (as water is), the gas phase is called a _________________.
Vaporization is the process by which a liquid changes into a gas or vapor. Vaporization is an endothermic process it requires _______________. When vaporization occurs only at the _____________________ of an uncontained
liquid (no lid on the container), the process is called evaporation. Molecules at the surface break away and become
gas. Only those with enough _____________________ energy (KE) escape. Evaporation is a
_______________________ process. It requires heat, which is endothermic. __________________ pressure is the
pressure exerted by a vapor over a liquid. As temperature increases, water molecules gain kinetic energy and vapor
pressure ______________________. When the vapor pressure of a liquid equals atmospheric pressure, the liquid
has reached its boiling point, which is 100°C for water at sea level. Recall that standard atmospheric pressure equals
______ atm. At this point, molecules throughout the liquid have the energy to enter the gas or vapor phase. The
temperature of a liquid can never ______________ above its boiling point. Boiling is an
__________________________ process. It requires the addition of heat. As you go up into the mountains (increase
in elevation), atmospheric pressure ______________. Lower external pressure requires ______________________
vapor pressure. Lower vapor pressure means lower ______________________ point. As a result, spaghetti cooks
slower in the mountains than at the beach. When you use a pressure cooker to can vegetables, the external pressure
around the mason jars rises. This raises the vapor pressure needed in order for water to boil. In turn, the boiling
point is raised so the food cooks ______________________.
Some phase changes release energy into their surroundings. For example, when a vapor loses energy, it may change
into a __________________. Condensation is the process by which a gas or vapor becomes a liquid. It is the
___________________ of vaporization. In a closed system, the rate of vaporization can equal the rate of
condensation. When first sealed, the molecules gradually _________________ the surface of the liquid. As the
molecules build up above the liquid, some condense back to a liquid. Equilibrium is reached when the rate of
vaporization __________________ the rate of condensation. Molecules are constantly changing phase. The total
amount of liquid and vapor remains _______________________.
99
Unit 9- Solids, Liquids and Phase Changes
The melting point of a solid is the temperature at which the ____________________ holding the particles together
are broken and the solid becomes a liquid. When heated the particles vibrate more _____________________ until
they shake themselves free of each other. The freezing point is the temperature at which a liquid becomes a
_________________________ solid. The freezing point is the _______________ as the melting point. The process
by which a solid changes directly into a gas without first becoming a liquid is called _______________________.
Solid air fresheners and dry ice are examples of solids that sublime. When a substance changes from a gas or vapor
directly into a solid without first becoming a liquid, the process is called _________________________. Deposition
is the reverse of sublimation. _______________ is an example of water deposition.

Classify the following phase changes.
1. dry ice (solid carbon dioxide) to carbon dioxide gas ____________________________
2. ice to liquid water ________________________________
3. liquid water to ice ________________________________
4. water vapor to liquid water ________________________________
Phase Diagrams
Temperature and _____________________ control the phase of a substance. A phase diagram is a graph of
pressure versus temperature that shows in which phase a substance exists under different conditions of temperature
and pressure. A phase diagram typically has ______ regions, each representing a different phase and three curves
that ________________________ each phase.
0.0098
Temperature (°C)
The points on the curves (lines) indicate conditions under which two phases coexist. The critical point indicates the
critical pressure and the critical temperature above which a substance cannot exist as a ____________________.
The triple point is the point on a phase diagram that represents the temperature and pressure at which three phases of
a substance can __________________________. The __________________________ slope of the solid-liquid line
in the phase diagram for water indicates that the solid floats on its liquid.

What happens to solid CO2 at -100 ºC and 1 atm pressure as it is heated to room temperature?
100
Unit 9- Solids, Liquids and Phase Changes

What happens to water at 1 atm as the temperature rises from -15°C to 60°C?

What state of matter is water at 50°C and 20 atm?

At what temperature does the triple point occur for water?

At what temperature does the critical point occur for carbon dioxide?

At standard pressure and -78°C, what phase change occurs for carbon dioxide?

What state of matter is carbon dioxide at -80°C and 2 atm?
Solids and Liquids Practice Test
Directions: Identify the proper sections by indicating the interval between 2 letters.
1.
2.
Heating Curve
Which section
represents the gas
being warmed?
_____ to _______
Which section
represents a phase
change from solid to
liquid?
_____ to _______
H
A
F
200º C
G
60º C
B
D
E
3.
Define viscosity-
4.
The temperature at
which the vapor
Energy
pressure of a liquid
equals the external or atmospheric pressure is known as the________________________
C
Directions: Using the phase diagram below,
answer questions 5-7:
5.
What does letter T represent?
Pressure
(atm)
C
_______________
6.
At 0.75 atm, what is the melting point
0.75
and boiling point of the substance?
T
________and _________
7.
0.25
At 0.25 atm, what is the freezing point?
________
125
101
175
250
Temperature (oC)
Unit 9- Solids, Liquids and Phase Changes
Multiple Choice
8.
Under the same conditions of temperature and pressure, a liquid differs from a gas because the particles of
the liquid
a. are in constant straight-line motion
b. take the shape of the container they occupy
c. have no regular arrangement
d. have stronger forces of attraction between them
9.
The phase change represented by the equation I2 (s)  I2 (g) is called
a. sublimation
c.
b. condensation
d.
melting
boiling
10. Which of the following terms represents the temperature and pressure at which three states of a compound
can coexist
a. Law of definite composition
d. Triple point
b. Van der Waals forces
e. Critical point
c. Graham’s Law of Diffusion
11. What is the smallest portion of a crystal lattice that reveals the 3-dimensional pattern?
a. unit cell
c. coordinate system
b. crystal structure
d. crystalline symmetry
12. What forces hold nonpolar particles together?
a. magic
b. hydrogen bonding
c.
d.
London dispersion
dipole-dipole
13. Compared with the particles in a solid, the particles in a liquid usually are
a. higher in energy
c. more massive
b. closer together
d. less fluid
14. What is the process of a substance changing from a vapor to a solid without passing through the liquid
phase?
a. condensation
c. sublimation
b. deposition
d. evaporation
15. A liquid forms when the average energy of a solid substance's particles
a. increases
c. creates an orderly arrangement
b. changes form
d. decreases
16. Which of the following is an NOT an amorphous solid?
a. silly putty
b. play dough
c.
d.
ice
glass
17. Which term best describes the process by which particles escape from both the surface of a liquid and from
within the liquid itself and enter the gas phase?
a. boiling
c. aeration
b. evaporation
d. surface tension
102
Unit 9- Solids, Liquids and Phase Changes
18. The attractive forces in a solid are
a. too weak to prevent the particles from changing positions
b. strong enough to hold the particles in fixed positions
c. less effective than those in a liquid
d. weaker than those of a liquid particles
19. When electrons in a covalent bond spend more time around on nucleus of the compound than the other, the
molecule is considered
a. weak
c. ionic
b. polar
d. nonpolar
20. Which of the following phase changes results in an overall increase in randomness of particles over the
course of the change?
a. deposition
c. melting
b. condensation
d. freezing
21. What type of crystals are like giant molecules?
a. covalent network
b. covalent molecular
c.
d.
metallic
ionic
22. The difference between crystalline and amorphous solids is determined by
a. temperature changes
b. pressure when the substances are formed
c. strength of molecular forces
d. the particle arrangement
23. Which of the following statements is false?
a. Condensed states have much higher densities than gases.
b. Molecules are very far apart in gases and closer together in liquids and solids.
c. Gases completely fill any container they occupy and are easily compressed.
d. Vapor refers to a gas formed by evaporation of a liquid or sublimation of a solid.
e. Solid water (ice), unlike most substances, is denser than its liquid form (water).
24. Which physical state/ property is incorrectly matched?
a. liquids and solids - rigid shape
b. gases - easily compressed
c. gases and liquids – flow
d.
e.
solids - higher density than gases
liquids – incompressible
25. Which one of the following statements does not describe the general properties of liquids accurately?
a. Liquids have characteristic volumes that do not change greatly with changes in temperature.
(Assuming that the liquid is not vaporized.)
b. Liquids diffuse only very slowly when compared to solids.
c. The liquid state is highly disordered compared to the solid state.
d. Liquids have high densities compared to gases.
103
Unit 9- Solids, Liquids and Phase Changes
26. For which of the following would permanent dipole-dipole interactions play an important role in
determining physical properties in the liquid state?
a. BF3
c. BeCl2
e. CCl4
b. ClF
d. F2
27. Identify which property liquids do not have in common with solids.
a. rigid shape
b. volumes do not change significantly with pressure
c. hydrogen bonding forces can be significant
d. practically incompressible
e. volumes do not change significantly with temperature
28. Which one of the following statements does not describe the general properties of solids accurately?
a. Solids have characteristic volumes that do not change greatly with changes in temperature.
b. Solids have characteristic volumes that do not change greatly with changes in pressure.
c. Solids diffuse only very slowly when compared to liquids and gases.
d. Solids are not fluid.
e. Most solids have high vapor pressures at room temperature.
29. Which of the following interactions are the strongest?
a. hydrogen bonding force
b. ion-ion interactions
c.
d.
dipole- dipole force
London-dispersion force
30. For which of the following would dispersion forces be the most important factor in determining physical
properties in the liquid state?
a. H2O
c. F2
e. NH4Cl
b. NaCl
d. HF
31. For which of the following would hydrogen bonding not be an important factor in determining physical
properties in the liquid state?
a. HI
c. HF
e. H2O2
b. H2O
d. NH3
32. Which technique listed below separates a mixture of liquids on the basis of their boiling points?
a. Distillation
c. Filtration
e. None of the
b. Extraction
d. Reflux
above
33. The melting point of a solid is the same as the ____ of its liquid.
a. Boiling point
b. Freezing point
c. Sublimation point
104
d.
e.
Condensation point
Critical point
Unit 9- Solids, Liquids and Phase Changes
34. Which one of the following statements does not describe the general properties of liquids accurately?
a. In the liquid state the close spacing of molecules leads to large intermolecular forces that are
strongly dependent on the nature of the molecules involved.
b. Liquids are practically incompressible.
c. As the temperature of a liquid is increased, the vapor pressure of the liquid decreases.
d. The normal boiling point of a liquid is the temperature at which the vapor pressure of the liquid
becomes equal to exactly 760 torr.
e. Vapor pressures of liquids at a given temperature differ greatly, and these differences in vapor
pressure are due to the nature of the molecules in different liquids.
35. Some solids can be converted directly to the vapor phase by heating. The process is called ____.
a. Fusion
d. Condensation
b. Sublimation
e. Distillation
c. Vaporization
36.
Which of the images shown here depicts a phase that has definite volume but not definite shape?
a. The one on
b. The one in the
c. The one on
the left
middle
the right
37.
Which phase depicted here has both a definite shape and a definite volume?
a. The one in the middle
d. The one in the middle and the one
b. The one on the right
on the right
c. The one on the left
38.
Which of the phases depicted here can be easily compressed?
a. The one in the middle
b. The one on the right
c. The one on the left
105
d.
The one in the middle and the one
on the right
Unit 9- Solids, Liquids and Phase Changes
39.
Which phase of matter is depicted here?
a.
b.
Liquid
Gas
c.
d.
Plasma
Solid
40. Ice floats in water because:
a. Water is denser than ice
b. Ice is colder than water
c. Water has a substantial surface tension
d. Ice is denser than water
41.
Which phase(s) depicted here have the ability to flow?
a. The one on the right
b. The one on the left
c. The ones on the right and the left
d. The one in the middle and the one on the right
e. The one in the middle
42. During the phase change from liquid to solid:
a. energy must be removed
b. energy must be absorbed
c. there is no change in energy
43. Definite shape, definite volume, and a low rate of diffusion are characteristics of:
a. Fluids
c. Gases
b. Liquids
d. Solids
44.
Which phase of matter is depicted here?
a.
b.
Solid
Gas
c.
d.
106
Liquid
Plasma
Unit 10- Solutions and Solubility
SOLUTIONS
A solution is made up of a solute and a _______________________________. The solvent does the
________________________________. The solute is the substance that is dissolved. If a solution is made of two
liquids, the one in ______________________ quantity is the solute. _________________________ is the universal
solvent. Water is a versatile solvent because of its attraction to other molecules and its
___________________________. Most of the water on the Earth is not pure, but rather is present in solutions.
Table salt (NaCl), like a great many ionic compounds, is _________________________ in water. The salt solution
is also an excellent ___________________________ of electricity. This high level of electrical conductivity is
always observed when ionic compounds dissolve to a significant extent in water. The process by which the charged
particles in an ionic solid separate from one another is called _____________________________. You can represent
the process of dissolving and dissociation in shorthand fashion by the following equation.
________________________________________ Water is not only good at dissolving ionic substances. It also is a
good solvent for many _________________________________ compounds. Consider the covalent substance
sucrose, commonly known as table sugar, as an example. Although water dissolves an enormous variety of
substances, both ionic and covalent, it does not dissolve everything. The phrase that scientists often use when
predicting solubility is “________________ dissolves like.” The expression means that dissolving occurs when
similarities exist between the solvent and the solute. A salt dissolves faster if it is _________________________ or
shaken, if the particles are made ________________________ and if the temperature is ______________________.
In order to dissolve the solvent molecules must come in ______________________________ with the solute.
Stirring moves fresh _____________________________ next to the solute. The solvent touches the surface of the
solute. __________________________________ pieces increase the amount of surface of the solute. For solids in
liquids, as the temperature goes up the solubility goes ______________________. A higher temperature makes the
molecules of the solvent move around ______________________________ and contact the solute harder and
________________________ often. It speeds up dissolving. Higher temperature usually increases the
_______________________________ that will dissolve.

How many grams of sodium hydroxide (NaOH) will dissolve in 100 g of water at 15ºC?

How many grams of sodium hydroxide will dissolve in 100 g of water at 40ºC?

At what temperature will 90 grams of Pb(NO3)2 dissolve in 100 g of water?

At what temperature will 30 grams of KNO3 dissolve in 100 g of water?
For gases in a liquid, as the temperature goes up the solubility goes _______________________. For gases in a
liquid, as the pressure goes up the solubility goes ______________________.
Solubility is the ________________________________ amount of substance that will dissolve at that temperature
(usually measured in grams/liter). If the amount of solute dissolved is less than the maximum that could be
dissolved, the solution is called a(n) ___________________________ solution. A solution which holds the
maximum amount of solute per amount of the solution under the given conditions is called a(n)
107
Unit 10- Solutions and Solubility
_____________________________ solution. A(n) _________________________________ solution contains more
solute than the usual maximum amount and are unstable. They cannot permanently hold the excess solute in solution
and may release it suddenly. A(n) __________________ crystal will make the extra come out. Generally, a
supersaturated solution is formed by dissolving a solute in the solution at an elevated temperature, at which
solubility is _______________________ than at room temperature, and then slowly cooling the solution.
Figure 1
Figure 2
108
Unit 10- Solutions and Solubility

If 45 g of KCl is dissolved in 100 g of water at 60ºC, is the solution unsaturated, saturated or
supersaturated?

If 90 g of Pb(NO3)2 is dissolved in 100 g of water at 40ºC, is the solution unsaturated, saturated or
supersaturated?

If 30 g of KNO3 is dissolved in 100 g of water at 20ºC, is the solution unsaturated, saturated or
supersaturated?

If 10 g of KClO3 is dissolved in 100 g of water at 50ºC, is the solution unsaturated, saturated or
supersaturated?
___________________________ means that two liquids can dissolve in each other.
________________________________ means they cannot. Oil and ______________________ are immiscible.
Measuring Solutions
Chemists never apply the terms strong and weak to solution concentrations. Instead, use the terms concentrated and
_________________________. Concentration is a measure of the amount of solute dissolved in a certain amount of
solvent. A concentrated solution has a _________________________ amount of solute. A dilute solution has a
__________________________ amount of solute. For chemistry applications, the concentration term molarity is
generally the most useful. Molarity is the number of moles of _______________________ in 1 Liter of the
solution.
Molarity 
Moles
Liters
Note that the volume is the total solution volume that results, not the volume of solvent alone. Suppose you need
1.0 Liter of a 1 M copper (II) sulfate solution.
STEP 1: Measure a mole of copper (II) sulfate.
STEP 2: Place the CuSO4 in a volumetric flask.
STEP 3: Add some water to dissolve the CuSO4 and then add enough additional water to bring the total volume
of the solution to 1.0 L.

What is the molarity of a solution with 2.0 moles of NaCl in 4.0 Liters of solution?

What is the molarity of a solution with 3.0 moles dissolved in 250 mL of solution?

How many moles of NaCl are needed to make 6.0 L of a 0.75 M NaCl solution?

0.200 moles of NaOH are dissolved in a small amount of water then diluted to 500. mL. What is the
concentration?

1.25 moles of NaCl are dissolved in a small amount of water then diluted to 625 mL. What is the
concentration?

How many moles are in 2.00 L of a 3.00 M solution of sulfuric acid (H 2SO4)?

How many moles are in 1500 mL of a 3.2 M solution of nitric acid (HNO 3)?

10.3 g of NaCl are dissolved in a small amount of water then diluted to 250 mL. What is the concentration?
109
Unit 10- Solutions and Solubility

20.3 g of NaOH are dissolved in a small amount of water then diluted to 500. mL. What is the
concentration?

80.6 g of KCl are dissolved in a small amount of water then diluted to 500. mL. What is the concentration?

125 g of NaC2H3O2 are dissolved in a small amount of water then diluted to 750. mL. What is the
concentration?

How many grams of CaCl2 are needed to make 625 mL of a 2.00 M solution?

How many grams of sugar are needed to make 125 mL of a 0.500 M C 6H12O6 solution?

How many grams of sodium hydroxide are needed to make 500. mL of a 0.750 M NaOH solution?

How many grams of aluminum nitrate are needed to make 600. mL of a 0.500 M Al(NO 3)2 solution?
Refer to Figure 1 to answer the following questions:

What is the molarity of a KNO3 solution at 10ºC? (100 g of water = 100 mL of water)

What is the molarity of a Pb(NO3)2 solution at 61ºC?

What is the molarity of a KNO3 solution at 71ºC?
Dilution
The number of moles of solute doesn’t change if you add more solvent.
M 1 x V1 = M 2 x V2
M1 and V1 represent the starting concentration and volume. M2 and V2 represent the ______________ concentration
and volume.

2.0 L of a 0.88 M solution are diluted to 3.8 L. What is the new molarity?

6.0 L of a 0.55 M solution are diluted to 8.8 L. What is the new molarity?

You have 150 mL of 6.0 M HCl. What volume of 1.3 M HCl can you make?

6.0 liters of a 0.55 M solution are diluted to a 0.35 M solution. What is the final volume?

You need 450 mL of 0.15 M NaOH. All you have available is a 2.0 M stock solution of NaOH. How do
you make the required solution?
Compounds in Aqueous Solution and Double Replacement Reactions
The _________________________________ of ions when an ionic compound dissolves in water is called
dissociation. Although no compound is completely insoluble, compounds of very low solubility can be considered
insoluble.

Using the solubility rules printed on page 6 of the NCDPI Reference Tables for Chemistry, determine
whether the following salts are soluble in water.
a) sodium chloride
d) nickel carbonate
b) mercury (I) acetate
e) barium sulfate
c) potassium nitrate
f) ammonium bromide
110
In a double-replacement reaction, two compounds exchange partners with each other to produce two different
compounds. The general form of the equation is
AB + CD  AD + CB
Signs that a double-replacement reaction has taken place include a color change, the release or absorption of energy,
evolution of a gas, and formation of a _______________________________.

Write the net ionic equation for each of the following precipitation reactions.
a) barium chloride + silver nitrate ___________________________________________
b) lead (II) nitrate + potassium iodide _______________________________________
c) ammonium sulfate + barium nitrate _______________________________________
d) potassium sulfide + zinc acetate __________________________________________
e) strontium phosphate + aluminum hydroxide __________________________________
Homework / Practice
1.
Suppose you had 2.00 moles of solute dissolved into 1.00 L of solution. What's the molarity?
2.
Calculate the molarity of 25.0 grams of KBr dissolved in 750.0 mL.
3.
80.0 grams of glucose (C6H12O6, mol. wt = 180. g/mol) is dissolved in enough water to make 1.00 L of
solution. What is its molarity?
4.
What is the molarity when 0.75 mol is dissolved in 2.50 L of solution
5.
What is the molarity of 245.0 g of H2SO4 dissolved in 1.00 L of solution?
6.
What is the molarity of 5.00 g of NaOH in 750.0 mL of solution?
7.
How many moles of Na2CO3 are there in 10.0 L of 2.0 M soluton?
8.
How many moles of Na2CO3 are in 10.0 mL of a 2.0 M solution?
9.
How many grams of Ca(OH)2 are needed to make 100.0 mL of 0.250 M solution?
10. What is the molarity of a solution made by dissolving 20.0 g of H 3PO4 in 50.0 mL of solution?
11. What weight (in grams) of KCl is there in 2.50 liters of 0.50 M KCl solution?
12. What is the molarity of a solution containing 12.0 g of NaOH in 250.0 mL of solution?
13. Determine the number of moles of solute to prepare these solutions:
a.
2.35 liters of a 2.00 M Cu(NO3)2 solution.
b.
16.00 mL of a 0.415-molar Pb(NO3)2 solution.
c.
3.00 L of a 0.500 M MgCO3 solution.
d.
6.20 L of a 3.76-molar Na2O solution.
14. Determine the final volume of these solutions:
a.
4.67 moles of Li2SO3 dissolved to make a 3.89 M solution.
b.
4.907 moles of Al2O3 to make a 0.500 M solution.
c.
0.783 grams of Na2CO3 to make a 0.348 M solution.
d.
8.97 grams of (NH4)2CO3 to make a 0.250-molar solution.
e.
48.00 grams of PbCl2 to form a 5.0-molar solution.
Solutions Practice Test
Directions: For credit, show all steps in your
calculations and include units.
1. What is the molarity of a solution of
NaOH if 12 liters of the solution contains
3 moles of NaOH?
2.
You have a 3.5 L solution that contains
20 grams of NaCl. What is the molarity
of the solution?
Directions: Using the solubility curve below,
answer the following questions.
3.
Which is most soluble at 40ºC? _______
4.
How many grams of KClO3 can be
dissolved in 100g H2O at 90ºC? _____
5.
At 40ºC, how much HCl can be dissolved
in 300 g. H2O? _________
For these common substances, identify what is the solute and solvent.
6. KoolAid (sugar, water)
7. Vinegar (acetic acid, water)
Solute
Solvent
Determine whether, according to the solubility rules, the mixing of these
substances will make a solution.
8. Water and Mg(OH)2
9. Water and Na2CO3
Yes
No
Faster
Slower
Determine how the following conditions can affect the rate of dissolving KCl
in water.
10. Decrease the temperature of the water
11. Agitate the mixture
Multiple Choice
12. Which of these compounds are soluble in water?
a. CaBr2
b. PbCl2
c.
d.
SrS
CaCO3
13. Which of these compounds is insoluble in water?
a. CuI
b. CaCl2
c.
d.
MgS
NaBr
14. Iron (III) sulfide is soluble in water.
a. True
b.
False
15. LiBr is
a. Soluble
b. Insoluble
c.
d.
can't tell the solubility
a covalent compound
16. NH4OH is insoluble.
a. True
b.
False
17. Which of these compounds is soluble?
a. Pb(OH)4
b. NaHCO3
c.
d.
BaCrO4
Mg3(PO4)2
18. Powdered NaCl will dissolve slower then NaCl crystals because there is less surface area for the reaction to
take place.
a. True
b. False
19. Which term indicates that there is a large quantity of solute, compared to the amount of solvent in a
solution
a. Dilute
c. Unsaturated
b. Concentrated
d. Saturated
20. Ten grams of sodium hydroxide is dissolved in enough water to make 1L of solution. What is the molarity
of the solution?
a. 0.25 M
c. 1 M
b. 0.5 M
d. 1.5 M
21. Which solution is the most concentrated?
a. 1 mole of solute dissolved in 1 liter of solution?
b. 2 moles of solute dissolved in 3 liters of solution?
c. 6 moles of solute dissolved in 4 liters of solution?
d. 4 moles of solute dissolved in 8 liters of solution?
22. What is the total number of moles of H2SO4 needed to prepare 5.0 liters of a 2.0 M solution of H 2SO4?
a. 2.5
c. 10
b. 5.0
d. 20
23. What is the molarity of a KF (aq) solution containing 116 grams of KF in 1.00 liter of solution?
a. 1.00 M
c. 3.00 M
b. 2.00 M
d. 4.00 M
24. The solubility of a gas will ___ when a solution containing the gas is heated and the solubility of a gas in a
solution will ___ when the pressure over the solution is decreased.
a. decrease...decrease
c. increase...decrease
b. decrease...increase
d. increase...increase
25. How many grams of potassium nitrate are required to prepare 3.00 x 10 2 mL of 0.750 M solution?
a. 2.28 x 104 g
c. 22.8 g
e. 0.00223 g
b. 84.5 g
d. 2.4 g
26. How many grams of sodium chloride are dissolved in 50.0 mL of 1.50 M solution?
a. 0.00324 g
c. 4.38 x 103 g
b. 117 g
d. 23.4 g
e.
4.38 g
27. A 500 mL sample of a 0.350 M solution is left open on a lab counter for two weeks, after which the
concentration of the solution is 0.955 M. What is the new volume of the solution?
a. 183 L
c. 0.183 L
e. 0.605 L
b. 223 mL
d. 1.83 mL
28. A chemist makes a stock solution of potassium chromate solution by dissolving 97.1 grams of the
compound in 1.00 liter of solution. What volume of the solution must be diluted with water in order to
prepare 200.0 mL of 0.200 M solution?
a. 80.0 mL
c. 750. mL
e. 120. mL
b. 0.150 L
d. 0.0800 mL
29. A 25.0-g sample of sodium hydroxide is dissolved in 400. mL of water. What is the concentration of the
solution?
a. 0.10 M
c. 1.56 M
e. 1.56 x 10-3 M
b. 62.5 M
d. 100. M
30. How many milliliters of 6.0 M HNO3 are needed to prepare 500 mL of 0.50 M HNO3?
a. 0.25 mL
c. 15.76 mL
b. 300 mL
d. 40 mL
e. None of these are correct
31. How many grams of calcium chloride are needed to prepare 300 mL of a 0.250 M solution?
a. 832 g
c. 566 g
e.
b. 5.66 g
d. 8.32 g
32. In a solution of sugar and water, the solvent is the:
a. sugar
b.
water
33. In a solution of sugar and water, the solute is the:
a. sugar
b.
water
112 g
34. Gases dissolve best in liquids when:
a. the pressure is high and the temperature is low
b. the pressure is low and the temperature is low
c. the pressure is low and the temperature is high
d. the pressure is high and the temperature is high
35. The solubility of potassium nitrate in water at 35 °C is about 60 grams KNO 3 per 100 grams of water. How
many grams of KNO3 should dissolve in 300 grams of water at 35 °C?
a. 180 grams
b. 335 grams
c. 20 grams
36. Breaking up a solid speeds dissolving in a liquid by:
a. decreasing the pressure
b. slowing hydration
c.
d.
raising the temperature
increasing surface area
37. Most salts become more soluble in water as the:
a. temperature is decreased
b. pressure is decreased
c.
d.
pressure is increased
temperature is increased
38. Calculate the molarity of a solution made when 80 grams of NaOH is dissolved in 2 L of solution
a.
b.
40 M NaOH
82 M NaOH
39. Which procedure will increases the solubility of KCl in water?
a. stirring the solute and solvent mixture
b. increasing the surface area of the solute
c. raising the temperature of the solvent
d. increasing the pressure on the surface of the solvent
40. Under which conditions are gases most soluble in water?
a. high pressure and high temperature
b. high pressure and low temperature
c. low pressure and high temperature
d. low pressure and low temperature
c.
d.
1 M NaOH
160 M NaOH
Unit 11- Acids and Bases
Acids and Bases
Properties of Acids and Bases
Acids taste _________________. Lemon juice and _____________________________, for example, are both
aqueous solutions of acids. Acids conduct electricity; they are ________________________. Some are strong
electrolytes, while others are _________________ electrolytes. An acetic acid solution, which is a weak electrolyte,
contains only a few ions and does not conduct as much current as a strong electrolyte. The bulb is only
_____________________ lit. Acids cause certain colored dyes (_________________________) to change color.
(Litmus paper turns _______________.) Acids react with metals to form ______________________________ gas.
This property explains why acids corrode most metals. Acids react with hydroxides (bases) to form water and a
___________________. Bases taste _________________________ and feel ______________________________.
Bases can be strong or weak electrolytes.
Naming Acids
Acids are compounds that give off _________________________ ions (H +) when dissolved in water. Acids will
always contain one or more hydrogen ions next to an __________________________. The anion determines the
name of the acid.
Naming Binary Acids
Binary acids contain hydrogen and an anion whose name ends in –ide. When naming the acid, put the prefix
______________________- and change -ide to -ic acid.
Example: HCl
The acid contains the hydrogen ion and chloride ion. Begin with the prefix hydro-, name the
nonmetallic ion and change -ide to -ic acid.
Example: H2S
The acid contains the hydrogen ion and sulfide ion. Begin with the prefix hydro- and name the
nonmetallic ion. The next step is change -ide to -ic acid, but for sulfur the “ur” is added before -ic.

Name the following binary acids.
a) HF ___________________________________________
b) H3P __________________________________________
Writing the Formulas for Binary Acids
The prefix hydro- lets you know the acid is binary. Determine whether you need to criss-cross the oxidation
numbers of hydrogen and the nonmetal.
Example: Hydrobromic acid
The acid contains the hydrogen ion and the bromide ion. The two oxidation
numbers add together to get zero. The prefix hydro- lets you know the acid is binary.
Example: Hydrotelluric acid
The acid contains the hydrogen ion and the telluride ion. The two oxidation
numbers do NOT add together to get zero, so you must criss-cross.

Write the formulas for the following binary acids.
a) Hydrocyanic acid _______________
b) Hydroselenic acid _______________
116
Unit 11- Acids and Bases
Naming Ternary Acids
The acid is a ternary acid if the anion has oxygen in it. The anion ends in -ate or -ite. Change the suffix -ate to _______ acid. Change the suffix -ite to -ous acid The hydro- prefix is NOT used!
Example: HNO3
The acid contains the hydrogen ion and nitrate ion. Name the polyatomic ion and change -ate
to -ic acid.
Example: HNO2
The acid contains the hydrogen ion and nitrite ion. Name the polyatomic ion and change -ite to
-ous acid.
Example: H3PO4
The acid contains the hydrogen ion and phosphate ion. Name the polyatomic ion and change -
ate to -ic acid.

Name the following ternary acids.
a) H2CO3 ___________________________________________________
b) H2SO4 ___________________________________________________
c) H2CrO4 ___________________________________________________
d) HClO2 ___________________________________________________
Writing the Formulas for Ternary Acids
The lack of the prefix hydro- from the name implies the acid is ternary, made of the hydrogen ion and a polyatomic
ion. Determine whether you need to criss-cross the oxidation numbers of hydrogen and the polyatomic ion.
Example: Acetic acid
The polyatomic ion must end in –ate since the acid ends in -ic. The acid is made of H+ and
the acetate ion. The two charges when added equal zero.
Example: Sulfurous acid
Again the lack of the prefix hydro- implies the acid is ternary, made of the hydrogen ion
and a polyatomic ion. The polyatomic ion must end in –ite since the acid ends in -ous. The acid is made of H+ and
the sulfite ion. The two charges when added do not equal zero, so you must crisscross the oxidation numbers.

Write the formulas for the following ternary acids.
a) perchloric acid ______________________
b) iodic acid _____________________
c) nitrous acid _______________________
d) bromic acid ___________________
Types of Acids and Bases
Arrhenius Definitions - The simplest definition is that an acid is a substance that produces _____________________
ions when it dissolves in water. A hydronium ion, H3O+, consists of a hydrogen ion attached to a
__________________ molecule. A hydronium ion, H3O+, is equivalent to H+. HCl and H3PO4 are acids according
to Arrhenius. A base is a substance that produces ________________________ ions, OH –, when it dissolves in
water. Ca(OH)2 and NaOH are Arrhenius bases. NH3, ammonia, could not be an Arrhenius
___________________.
117
Unit 11- Acids and Bases
Monoprotic acids have only ____________ ionizable hydrogen. Some acids have more than one ionizable hydrogen
and are called ______________________________ acids.
Bronsted-Lowry Definitions - An Bronsted-Lowry acid is a ________________________ (H+) donor. HBr and
H2SO4 are Bronsted-Lowry acids. When a Bronsted-Lowry acid dissolves in water it gives its proton to water. HCl
(g) + H2O (l) ↔ H3O+ + Cl- A Bronsted-Lowry base is a proton acceptor. B + H2O ↔ BH+ + OH- A
Brønsted-Lowry base does not need to contain OH-.
Consider
HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq)
HCl donates a proton to water. Therefore, HCl is an
_______________. H2O accepts a proton from HCl. Therefore, H2O is a ______________.

Identify the acid and base in the following reactions.
a) H2SO3 + H2O ↔ HSO3- + H3O+
Acid _____________________________
b) NH3 + H2SO4 ↔ NH4 +
+
base _________________________
HSO4-
Acid _____________________________
base _________________________
Molarity and Dilution
The concentration of a solution is the amount of solute present in a given quantity of solution.
_________________________ is the number of moles of solute in 1 liter of solution.
moles solute
Molarity = -------------------------liters of solution
The procedure for preparing a less concentrated solution from a more concentrated one is called a
___________________________.
M 1 V1 = M 2 V2
PRACTICE:
 What is the molarity of an acetic acid (HC2H3O2) solution with 4.0 moles dissolved in 250 mL of solution?

How many moles of hydrochloric acid (HCl) are needed to make 3.0 L of a 0.55 M HCl solution?

0.600 moles of the base sodium hydroxide (NaOH) are dissolved in a small amount of water then diluted to
500 mL. What is the concentration?

3.25 moles of the base potassium hydroxide (KOH) are dissolved in a small amount of water then diluted to
725 mL. What is the concentration?

How many moles are in 2.00 L of a 6.00 M solution of sulfuric acid (H 2SO4)?

How many moles are in 1250 mL of a 3.60 M solution of nitric acid (HNO 3)?

6.0 L of a 1.55 M LiOH solution are diluted to 8.8 L. What is the new molarity of the lithium hydroxide
solution?

You have 250 mL of 6.0 M HCl. How many milliliters of 1.2 M HCl can you make?
118
Unit 11- Acids and Bases

4.0 liters of a 0.75 M solution of sulfuric acid (H 2SO4) are diluted to a 0.30 M solution. What is the final
volume?

You need 350 mL of 0.25 M NaOH. All you have available is a 2.0 M stock solution of NaOH. How do you
make the required solution?
Strength of Acids and Bases
The strength of a base is based on the percent of units ___________________________________, not the number
of OH– ions produced. The strength of a base does NOT depend on the _____________________________. 1A
and _______ hydroxides, excluding __________, are strong bases. Some bases, such as Mg(OH) 2, are not very
soluble in water, and they don’t produce a large number of OH – ions. However, they are still considered to be
strong bases because all of the base that does dissolve completely dissociates. The strength of an acid is based on
the percent of units dissociated, not the number of ____________ ions produced. The strength of an acid does NOT
depend on the _______________________________. There are 6 strong acids: HCl, HBr, HI, HClO 4, HNO3, and
H2SO4. Strong acids and bases are strong __________________________________ because they dissociate
completely. Electrolytes conduct ______________________________. Weak acids and bases don’t completely
ionize, so they are weak electrolytes. Although the terms weak and strong are used to compare the
_____________________________ of acids and bases, dilute and concentrated are terms used to describe the
_____________________________ of solutions.
pH Scale
Water ionizes; it falls apart into _________________. H2O  H+ + OH- The preceding reaction is called the
_____________________________________ of water. [H+ ] = [OH -] = 1 x 10-7 M When [H+ ] = [OH-], the
solution is _________________________. At 25°C, Kw = [H+] [OH-] = 1 x 10-14 Kw is called the ion-product
constant. If [H+] > 10-7 then [OH-] < 10-7. The solution is ______________________ when [H +] > [OH-]. If [H+]
< 10-7 then [OH-] > 10-7. The solution is __________________________ when [OH -] > [H+]. In most applications,
the observed range of possible hydronium or hydroxide ion concentrations spans 10 –14 M to ______M. To make
this range of possible concentrations easier to work with, the pH scale was developed. pH is a mathematical scale in
which the concentration of hydronium ions in a solution is expressed as a number from _________ to __________.
pH meters are instruments that measure the exact pH of a solution. Indicators register different colors at different
pH’s. In neutral solution, pH = 7. In an acidic solution, pH < 7. In a basic solution, pH > 7. As the pH drops from
7, the solution becomes more acidic. As pH increases from 7, the solution becomes more basic.
The pH of a solution equals the negative logarithm of the hydrogen ion concentration.
pH = - log [H+]
Chemists have also defined a pOH scale to express the basicity of a solution.
pOH = - log [OH-]
If either pH or pOH is known, the other may be determined by using the following relationship.
pH + pOH = 14.00
119
Unit 11- Acids and Bases

Find the pH of the following solutions.
a) The hydronium ion concentration equals: 10 –2 M. pH = _________________
b) The hydronium ion concentration equals: 10 –11 M. pH = _________________
c) The hydronium ion concentration equals: 1 x 10 –6 M. pH = _________________
d) The hydroxide ion concentration equals: 10 –8 M. pH = _________________
e) The hydroxide ion concentration equals: 10 –5 M. pH = _________________
f) The hydroxide ion concentration equals: 10–3 M. pH = _________________

If a certain carbonated soft drink has a hydrogen ion concentration of 1.0 x 10 –4 M, what are the pH and
pOH of the soft drink?
Calculating Ion Concentrations From pH
If either pH or pOH is known, the hydrogen ion or hydroxide ion can be found.
[H+] =10-pH
[OH-] =10-pOH
On the calculator, hit
 2nd
 log
 (-)
 and then the number.

Find the [H+] of a solution that has a pH equal to 6.

Find the [H+] of a solution that has a pH equal to 12.

Find the [H+] of a solution that has a pH equal to 5.

Find the [H+] of a solution that has a pOH equal to 6.

Find the [H+] of a solution that has a pOH equal to 6.

Find the [H+] of a solution that has a pOH equal to 2.

Find the [H+] of a solution that has a pOH equal to 4.

Find the [OH-] of a solution that has a pH equal to 10.
Calculating Ion Concentration From Ion Concentration
If either [H+] or [OH-] is known, the hydrogen ion or hydroxide ion can be found.
[H+] [OH-] = 1 x 10-14
 Find the hydrogen ion concentration if the hydroxide ion concentration equals: 1 x 10 –8 M.

Find the hydrogen ion concentration if the hydroxide ion concentration equals: 1 x 10 –2 M.

Find the hydroxide ion concentration if the hydrogen ion concentration equals: 1 x 10 –4 M.

Find the hydroxide ion concentration if the hydrogen ion concentration equals: 1 x 10 –9 M.
Indicators
120
Unit 11- Acids and Bases
Chemical _____________________ whose colors are affected by acidic and basic solutions are called indicators.
Many indicators do not have a sharp color change as a function of ____________. Most indicators tend to be
__________________ in more acidic solutions.

Which indicator is best to show an equivalence point pH of 4?

Which indicator is best to show an equivalence point pH of 11?

Which indicator is best to show an equivalence point pH of 2?
Neutralization Reactions
The reaction of an acid and a base is called a neutralization reaction. Acid + base  salt + water
A salt is an
___________________ compound.

Predict the products of and balance the following neutralization reactions. (Remember to check the
oxidation numbers of the ions in the salt produced.)
a) HNO3 + KOH 
The salt is composed of the ________________ ion and the _________________ ion.
b) HCl + Mg(OH)2 
c) H2SO4 + NaOH 
Neutralization
 How many moles of HNO3 are needed to neutralize 0.86 moles of KOH?

How many moles of HCl are needed to neutralize 3.5 moles of Mg(OH) 2?

How many moles of H3PO4 are needed to neutralize 3.5 moles of Mg(OH)2?
121
Unit 11- Acids and Bases

How many moles of HC2H3O2 are needed to neutralize 3.5 moles of Cr(OH)3?

If it takes 87 mL of an HCl solution to neutralize 0.67 moles of Mg(OH) 2 what is the concentration of the
HCl solution?

If it takes 58 mL of an H2SO4 solution to neutralize 0.34 moles of NaOH what is the concentration of the
H2SO4 solution?

If it takes 85 mL of an HNO3 solution to neutralize 0.54 moles of Mg(OH)2 what is the concentration of the
HNO3 solution?

If it takes 150. mL of an Ca(OH)2 solution to neutralize 0.800 moles of HCl what is the concentration of the
Ca(OH)2 solution?
Acid Rain
Acid Rain is any rain with a pH less than __________. Pure rain is naturally acidic because of dissolved
_______________. It is caused by the man-made oxides of _______________________ and nitrogen.
SO3 + H2O  H2SO4 Research shows acid rain is associated with parts of a country where heavy
______________________________ are situated and also down-wind from such sites. Analysis of acid rain
indicates that especially sulfur oxides, SOx and nitrogen oxides, NOx are mostly responsible from rain acidity. Snow
_________________, sleet, hail and drizzle all become contaminated with acids when SO x and NOx are present as
pollutants.
Titration
The known reactant molarity is used to find the unknown _________________________ of the other solution.
Solutions of known molarity that are used in this fashion are called _________________________ solutions. In a
titration, the molarity of one of the reactants, acid or base, is known, but the other is unknown.

A 15.0-mL sample of a solution of H2SO4 with an unknown molarity is titrated with 32.4 mL of 0.145M
NaOH to the bromothymol blue endpoint. Based upon this titration, what is the molarity of the sulfuric acid
solution?
First find the number of moles of the solution for which you know the molarity and volume. Next, use the molemole ratio to determine the moles of the unknown. Finally, determine the molarity of the unknown solution.

If it takes 45 mL of a 1.0 M NaOH solution to neutralize 57 mL of HCl, what is the concentration of the
HCl ?

If it takes 67.0 mL of 0.500 M H2SO4 to neutralize 15.0 mL of Al(OH)3 what was the concentration of the
Al(OH)3 ?

How many moles of 0.275 M HCl will be needed to neutralize 25.0 mL of 0.154 M NaOH?
Titration Curves
A plot of ___________ versus volume of acid (or base) added is called a titration curve.
122
Unit 11- Acids and Bases
Strong Base-Strong Acid Titration Curve
Consider adding a strong base (e.g. NaOH) to a solution of a strong acid (e.g. HCl). Before any base is added, the
pH is given by the strong _________________ solution. Therefore, pH ____ 7. When base is added, before the
equivalence point, the pH is given by the amount of strong acid in _________________________. Therefore, pH <
7. At ________________________________ point, the amount of base added is stoichiometrically equivalent to
the amount of acid originally present. Therefore, pH =_________. To detect the equivalent point, we use an
indicator that changes ____________________somewhere near 7.00. Past the equivalence point all acid has been
consumed. Thus one need only account for excess __________________. Therefore, pH ______ 7.
Homework / Practice
1.
2.
Name the following compounds as acids.
a.
H2SO4
c.
H2 S
b.
H2SO3
d.
HClO4
c.
phosphoric acid
Write formulas for the following acids.
a.
nitrous acid
b.
hydrobromic acid
3.
Use an activity series to identify two metals that will not generate hydrogen gas when treated with an acid.
4.
Write the equation that represents the following reaction: the ionization of HClO3 in water
5.
Explain how strong acid solutions conduct an electric current.
123
Unit 11- Acids and Bases
6.
7.
Write balanced molecular equations for the reactions of acids and bases.
a.
aluminum metal with dilute nitric acid
b.
calcium hydroxide solution with acetic acid
CaCO3 (s) + HCl (aq)  CaCl2 (aq) + H2O (l) + CO2 (g)
a.
Balance the above equation.
b.
How many liters of CO2 form at STP if 5.0 g of calcium carbonate are treated with excess
hydrochloric acid?
8.
9.
Consider the following reaction: NH4+ (aq) + CO3-2 (aq) ↔ NH3 (aq) + HCO3-1 (aq)
a.
What reactant serves as the base?
b.
What reactant serves as the acid?
Given the following reaction: HCO3-1 (aq) + OH-1 (aq) ↔ CO3-2 (aq) + H2O (l)
a.
What reactant serves as the base?
b.
What reactant serves as the acid?
10. Write the formula for the salt formed in each of the following neutralization reactions.
a.
hydrobromic acid combined with barium hydroxide
b.
lithium hydroxide combined with sulfuric acid
11. H2SO4 (aq) + NaOH (aq)  Na2SO4 (aq) + H2O (l)
a.
Balance the above neutralization equation
b.
In order to completely consume all reactants, what should be the mole ratio of acid to base?
12. Consider the reaction represented by the following incomplete equation:
Ba(OH)2 (aq) + H2SO4 (aq) 
a. Predict the products of this reaction, and write the balanced equation.
b.
Use the solubility rules to determine the solubility of the salt produced in the reaction.
c.
If 0.030 mol of Ba(OH)2 is consumed, how many grams of water are produced?
13. Perform the following calculations.
a.
If the hydronium concentration is 1 x 10-6 M for a solution, calculate the hydroxide concentration.
b.
If the hydroxide concentration is 1 x 10-12 M for a solution, calculate the hydronium concentration.
c.
If the pOH = 4.00 for a solution, calculate the pH. Is the solution acidic or basic?
124
Unit 11- Acids and Bases
d.
If the hydronium concentration is 1.00 x 10 -3 M, calculate the pOH.
e.
If the pOH = 5.0 for a solution, calculate the hydroxide concentration.
f.
If the pH = 12.0 for a solution, calculate the hydronium concentration.
g.
If the pH = 3.00 for a solution, calculate the hydroxide concentration.
h.
If the hydronium concentration = 1.0 x 10 -8 M for a solution, calculate the hydroxide
concentration.
14. Summarize the two main acid-base theories in the table below.
ACID
BASE
Arrhenius
Brønsted-Lowry
15. Label the acid (A), base (B) in each of the following reactions.
a.
H2SO4 + NH3  HSO4
b.
HC2H3O2 + H2O  H3O+ + C2H3O2
c.
NaHCO3 + HCl  NaCl + H2CO3
+ NH4
16. Give the products of the dissociation of each of the following Brønsted-Lowry acids.
a.
HI
c.
H2CO3
b.
NH4+
d.
HNO3
17. Give the product formed when each of the following Brønsted-Lowry bases accepts a proton.
a.
CN–
c.
CH3COO–
b.
O2–
d.
NH3
18. Find [OH ] for 1.0 × 10-12 M HClO4.
19. What is the pH of 1.0 × 10-4 M HCl?
20. What is the pH of 1.5 × 10-3 M NaOH?
21. A solution of HNO3 has a pH of 4.0. What is the molarity of HNO 3?
22. What is the molarity of KOH in a solution that has a pH of 10.0?
125
Unit 11- Acids and Bases
Acids and Bases Practice Test
Directions: Give the names or formulas for the following acids, bases, and salts:
1. KOH________________________
3. Sulfuric Acid__________________
2.
HNO3 _______________________
4.
Magnesium hydroxide ___________
Directions:.
5. In complete sentences, define an acid according to the Arrhenius theory.
Directions: Label (according to Bronsted-Lowry) the Bronsted-Lowry acid, Bronsted-Lowry base, conjugate acid,
and conjugate base in each of the equations below:
6.
H₂O + HC2H3O2 ⇆ H₃O⁺ + C2H3O2⁻
7.
CN-- + H₃O⁺ ⇆ H₂O + HCN
Directions Identify the following as an acid or a base, strong or weak.
a.Acid or base b.Strong or weak
8. 2 M KOH
__________
_____________
9.
7 M H₂SO₄
10. 0.12 M H2S
__________
_____________
_________
_____________
Directions: Complete and balance the following neutralization reactions.
11.
NaOH +
HCl → _______________ + __________________
12.
H₂SO₄ +
KOH → _____________
+ __________________
13. Determine the pOH for a solution of HNO3 that has a concentration of 0.01 M.
14. Determine the pH for a solution of CuOH that has an [OH-] of 0.000001 M.
Directions: Complete the following chart.
pH
[H₃O⁺]or [H+]
15.
16.
[OH⁻]
pOH
1 x 10⁻6
2
17.
4
Multiple Choice Practice
18. According to the Arrhenius theory, a base yields
a. H+ as the only positive ion in an aqueous solution
b. OH+ as the only positive ion in an aqueous solution
c. OH- as the only negative ion in an aqueous solution
d. H- as the only negative ion in an aqueous solution
126
Acidic or basic
Unit 11- Acids and Bases
19. In the reaction H2SO4(aq) --> 2H+(aq) + SO4-2(aq) H2SO4 is a(n)
a. Arrhenius acid
b. Arrhenius base
c.
Salt
21. Which of the following is an Arrhenius base?
a. HCl
b. H2SO4
c.
d.
NaCl
NH3
22. Which of the following is a salt?
a. HOH
b. NH4NO3
c.
d.
HCl
H2CO3
20. Arrhenius acids yield
a. OH- as the only negative ion in an aqueous solution
b. H- as the only negative ion in an aqueous solution
c. H3O+ as the only positive ion in an aqueous solution
d. OH+ as the only positive ion in an aqueous solution
23. A substance that conducts an electrical current when dissolved in water is called
a. an acid
c. an ionic compound
b. an electrolyte
d. a nonelectrolyte
24. Which of the following can conduct an electric current?
a. Mg(OH)2 (s)
b. H2O (s)
c.
d.
NaOH (aq)
NH4Cl (s)
25. Which electrolyte is best at conducting electricity when dissolved in an aqueous solution?
a. KCl (s)
c. CaCl2 (s)
b. Na2SO4 (s)
d. H3PO4 (s)
26. In the process of neutralization a salt and a base react to yield water and an acid.
a. True
b. False
27. A student dissolved NaCl (s) in water, and tested with a battery, wire, and a light blub to see if it conducted
an electric current. The solution conducted an electric current. This is because NaCl (s) is
a. a salt and an electrolyte
b. a salt and a nonelectrolyte
c. a Arrhenius acid and an electrolyte
d. a Arrhenius base and an electrolyte
28. In the process of neutralization, an Arrhenius acid and an Arrhenius base react to form which of the
following?
a. Water only
c. Water and Carbon dioxide
b. Salt and Carbon dioxide
d. Water and Salt
29. What type of chemical reaction is neutralization?
a. single replacement
b. double replacement
c.
d.
synthesis
decomposition
30. Which of the following can be a product of neutralization?
a. LiOH(s)
b. HCl(l)
c.
d.
MgI2(l)
NaCl(aq)
31. A titration reaction is a complete neutralization reaction where the moles of H+ equal the moles of OH-.
a. True
b. False
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Unit 11- Acids and Bases
32. Which of the following reactants will represent a neutralization?
a. BaCl2 + CaSO4
b. HCl + F
c.
d.
Ca(OH)2 + H2SO4
NaCl + H2O
33. When NaOH and HCl react, what will be on the product side?
a. only NaCl
b. only HOH
c.
d.
NaCl and HOH
NaCl and Cl2
34. Titration is a process in which
a. which a volume of solution of unknown concentration is used to determine the concentration of
another solution.
b. which a volume of solution of known concentration is used to determine the volume of another
solution.
c. which a volume of solution of known concentration is used to determine the concentration of
another solution.
d. which a volume of solution of known concentration is used to determine the curve of another
solution.
35. In MaVa=MbVb, what is Mb?
a. molarity of OHb. molarity of H+
c.
d.
molarity of Hmolarity of O2
36. What is the molarity of HCl (aq) if 25 mL of 8.0 M NaOH (aq) neutralizes exactly 20.0 mL of HCl (aq)?
a. 5M
c. 15M
b. 10M
d. 20M
37. At the end point of titration, what is the relationship between moles of H + and OH-?
a. the moles of H+ are greater than OHb. the moles of OH- are greater than H+
c. the moles of H+ are equal to moles of OHd. there is no relationship between moles of H+ and OH38. The molarity of HBr (aq) is 2M when 10 milliliters of 8 M NaOH (aq) neutralizes exactly 20 milliliters of
HBr.
a. True
b. False
39. What is the molarity of NaOH if 5 milliliters of 4M HCl (aq) neutralizes exactly 10 mL of NaOH (aq)?
a. .5M
c. 1.5M
b. 1M
d. 2M
40. If 10.milliliters of a 0.40 M HBr solution is required to neutralize exactly 0.2 M of NaOH, what is the
volume of the base?
a. 10ml
c. 30ml
b. 20ml
d. 40ml
41. The molarity of an acid can be calculated if a base of known concentration (standard base) is added, drop
by drop, to a specific volume of the acid until the indicator changes color.
a. True
b. False
42. One acid-base theory states that an acid is an H +
a. Acceptor
b. Eliminator
c.
d.
Dissolver
Donor
43. According to the Bronsted-Lowry acid-base theory, a base is a substance that can
a. donate an electron
c. donate a proton
b. accept a proton
d. accept a electron
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Unit 11- Acids and Bases
44. In the following reaction, NH3 + HCl --> NH4+ + Cl- NH3 acts as a(n)
a. base in the reverse reaction.
c. base in the forward reaction.
b. acid in the forward reaction.
d. acid in the reverse reaction.
45. The acidity or alkalinity of a solution can be measured by its pH value.
a. True
b. False
46. The relative level of acidity or alkalinity of a solution can be shown by using their pH values.
a. True
b. False
47. Which of these 1 M solutions will have the highest pH?
a. H3PO4
b. HCl
c.
d.
NaCl
NaOH
48. Which pH indicates an acidic solution?
a. 1
b. 7
c.
9
d.
12
49. Which of these pH numbers indicates the lowest level of acidity?
a. 1
b. 3
c.
8
d.
12
50. Which formula represents a salt?
a. KOH
b. KCl
c.
d.
CH3OH
CH3COOH
c.
d.
LiOH
KOH
52. Which solution will change red litmus to blue?
a. HCl (aq)
b. NaCl (aq)
c.
d.
CH3OH (aq)
NaOH (aq)
53. An acidic solution could have a pH of
a. 7
b.
c.
3
d.
14
c.
5
51.
Which substance can be classified as an Arrhenius acid?
a. HCl
b. NaCl
10
54. What is the pH of a 0.00001 molar HCl solution?
a. 1
b. 9
55. What is the pH of a solution with a hydronium ion concentration of 0.01 moles per liter?
a. 1
b. 2
c. 10
56. Which 0.1 M solution will turn phenolphthalein pink?
a. HBr (aq)
b. CO2 (aq)
c.
d.
LiOH (aq)
CH3OH (aq)
57. Given the equation: H+ + OH- ↔ H2O
Which type of reaction does the equation represent?
a. esterification
b. decomposition
c.
d.
hydrolysis
neutralization
129
d.
d.
4
14
Unit 11- Acids and Bases
58. As the hydrogen ion concentration of an aqueous solution increases, the hydroxide ion concentration of this
solution will
a. decrease
c. remain the same
b. increase
59. A student wishes to prepare approximately 100 milliliters of an aqueous solution of 6 M HCl using
12 M HCl. Which procedure is correct?
a.
b.
c.
d.
adding 50 mL of 12 M HCl to 50 mL of water while stirring the mixture steadily.
adding 50 mL of 12 M HCl to 50 mL of water and then stirring the mixture steadily.
adding 50 mL of water to 50 mL of 12 M HCl while stirring the mixture steadily.
adding 50 mL of water to 50 mL of 12 M HCl and then stirring the mixture steadily.
60. The reaction of an acid like HCl and a base like NaOH always
a. Forms a precipitate
b. Forms a volatile product
c.
d.
Forms a salt and water
Forms a sulfate and water
61. What is the pH of an acetic acid solution if the [H3O+] = 1 x 10-4 mol/L?
a. 1
c. 3
b. 2
d. 4
130
e.
5
Unit 12- Kinetics and Thermochemistry
REACTION KINETICS
Energy Diagrams
Reactants always start a reaction so they are on the _________________ side of the diagram. Products are on the
right. The exothermic reaction gives off ___________________ because the products are at a lower energy level
than the reactants. In an exothermic graph, the reactants have _____________________ energy than the products.
The change in energy is a _________________________ value.
The endothermic reaction absorbs heat because the products are at a _________________________ energy level
than the reactants. In an endothermic graph, the products have _____________________ energy than the reactants.
The change in energy is a _______________________ value.
Scientists have observed that the energy released in the formation of a compound from its elements is always
identical to the energy required to ______________________ that compound into its elements.
_____________________________ energy is the minimum amount of energy that reacting particles must have to
form the activated complex. The activated complex is a short-lived, _________________ arrangement of atoms that
may break apart and re-form the reactants or may form products. To calculate the activation energy,
______________________ the energy of the reactants from the energy at the top of the peak. The enthalpy or heat
of reaction (ΔH) is the amount of ___________________ released or absorbed in the reaction. To determine ΔH,
take the energy of the products and _______________________ the energy of the reactants.
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Unit 12- Kinetics and Thermochemistry

The heat content of the reactants of the forward reaction is about ________ kilojoules.

The heat content of the products of the forward reaction is about ________ kilojoules.

The heat content of the activated complex of the forward reaction is about _______ kilojoules.

The activation energy of the forward reaction is about _______ kilojoules.

The heat of reaction (ΔH) of the forward reaction is about _______ kilojoules.

The forward reaction is (endothermic or exothermic).

The heat content of the reactants of the reverse reaction is about ________ kilojoules.

The heat content of the products of the reverse reaction is about ________ kilojoules.

The heat content of the activated complex of the reverse reaction is about _______ kilojoules.

The activation energy of the reverse reaction is about _______ kilojoules.

The heat of reaction (ΔH) of the reverse reaction is about _______ kilojoules.

The reverse reaction is (endothermic or exothermic).
The activation energy can be lowered by adding a __________________________. The catalyst
_________________ the activation energy by providing an alternate pathway for the reaction to occur.
Expressing Reaction Rates
We generally define the average _____________________ of an action or process to be the change in a given
quantity during a specific period of time. Reaction rates cannot be calculated from balanced equations as
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Unit 12- Kinetics and Thermochemistry
stoichiometric amounts can. Reaction rates are determined experimentally by measuring the
_____________________________ of reactants and/or products in an actual chemical reaction.
Collision Theory
According to the collision theory, atoms, ions, and molecules must collide with each other in order to react. The
following three statements summarize the collision theory.
1. Particles must ________________________ in order to react.
2. The particles must collide with the correct _______________________________.
3. The particles must collide with enough ________________________ to form an unstable activated
complex, also called a _____________________________ state, which is an intermediate particle made
up of the joined reactants.
The _____________________________ amount of energy that colliding particles must have in order to form an
activated complex is called the activation energy of the reaction. Particles that collide with energy less than the
activation energy ____________________________ form an activated complex. In an exothermic reaction,
molecules collide with enough energy to overcome the activation energy barrier, form an activated complex, then
__________________________ energy and form products at a lower energy level. In the reverse endothermic
reaction, the reactant molecules lying at a _______________ energy level must absorb energy to overcome the
activation energy barrier and form high-energy products.
Factors Affecting Reaction Rates
The reaction rate for almost any chemical reaction can be modified by varying the conditions of the reaction.
1) An important factor that affects the rate of a chemical reaction is the reactive nature of the reactants. As
you know, some substances react more readily than others. The more reactive a substance is, the
_____________________________ the reaction rate.
2) Another important factor that affects the rate of a chemical reaction is the concentration of the reactants.
Reactions ______________________ _________________ when the concentrations of reacting particles
are increased. Increasing the number of reactants increases probability of collisions. The rate of gaseous
reactions can be ___________ by pumping more gas into the reaction container.
3) Surface area of the reactants affects the rate of a chemical reaction. ______________________ the surface
area of reactants provides more opportunity for collisions with other reactants, thereby increasing the
reaction rate.
4) Temperature affects the rate of a chemical reaction. Generally, increasing the temperature at which a
reaction occurs _________________________ the reaction rate. Raising the temperature raises both the
collision frequency and the collision energy.
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Unit 12- Kinetics and Thermochemistry
5) Adding a catalyst affects the rate of a chemical reaction. A catalyst is a substance that increases the rate of
a chemical reaction without itself being consumed in the reaction. In fact, catalysts are not included in the
chemical equation.
6) Compressing gases affects the rate of a chemical reaction. When two gases react, compressing the gases
generally ___________ the rate of reaction.
REACTION ENERGY
Energy is the ability to do___________________ or produce heat. It exists in two basic forms, potential energy and
______________________ energy. Potential energy is energy due to the ______________________ or position of
an object. Kinetic energy is energy of ________________. The potential energy of the dammed water is converted
to kinetic energy as the dam gates are opened and the water flows out. Chemical systems contain
_________________ kinetic energy and potential energy. As temperature increases, the motion of submicroscopic
particles ______________________, so its kinetic energy __________________________. The potential energy of
a substance depends upon its composition: the type of atoms in the substance, the number and type of chemical
bonds joining the atoms, and the particular way the atoms are arranged.
Law of Conservation of Energy and Heat
The law of conservation of energy states that in any chemical reaction or physical process, energy can be converted
from one form to another, but it is neither created nor ________________________.
Heat, which is represented by the symbol ____, is energy that is in the process of flowing from a
_____________________ object to a cooler object. The SI unit of heat and energy is the joule (J). Heat involves a
transfer of energy between 2 objects due to a ________________________ difference. When the warmer object
loses heat, its temperature decreases and q is _________________________. When the cooler object absorbs heat,
its temperature ________________ and q is positive.
The specific heat of any substance is the amount of heat required to raise the temperature of ____ gram of that
substance by one degree Celsius. Because different substances have different compositions, each substance has its
own specific heat.
q = m Cp ∆T
q = heat (J); m = mass (g); Cp = specific heat (J/(g.°C); ∆T = change in temperature = T f – Ti (°C)
Exothermic: Heat flows _________ of the system (to the surroundings). The value of ‘q’ is negative. Endothermic:
Heat flows _________ the system (from the surroundings). The value of ‘q’ is positive.

The temperature of a sample of iron with a mass of 10.0 g changed from 50.4°C to 25.0°C with the release
of 114 J heat. What is the specific heat of iron?

A piece of metal absorbs 256 J of heat when its temperature increases by 182°C. If the specific heat of the
metal is 0.301 J/g.°C, determine the mass of the metal.
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Unit 12- Kinetics and Thermochemistry

If 335 g water at 65.5°C loses 9750 J of heat, what is the final temperature of the water? The specific heat
of water is 4.18 J/g.°C.

As 335 g of aluminum at 65.5°C gains heat, its final temperature is 300.°C. The specific heat of aluminum
is 0.897 J/g.°C. Determine the energy gained by the aluminum.
Heat changes that occur during chemical and physical processes can be measured accurately and precisely using a
___________________________. A calorimeter is an insulated device used for measuring the amount of heat
absorbed or released during a chemical or physical process. A coffee-cup calorimeter made of ________ Styrofoam
cups.

Suppose you put 125 g of water into a foam-cup calorimeter and find that its initial temperature is 25.6°C.
Then, you heat a 50.0 g sample of the unknown metal to a temperature of 115.0°C and put the metal sample
into the water. Both water and metal have attained a final temperature of 29.3°C. Heat flows from the hot
metal to the cooler water and the temperature of the water rises. The flow of heat stops only when the
temperature of the metal and the water is equal. Assuming no heat is lost to the surroundings, the heat
gained by the water is equal to the heat lost by the metal. Determine the specific heat of the metal.

You put 352 g of water into a foam-cup calorimeter and find that its initial temperature is 22.0°C. What
mass of 134°C lead, Clead = 0.129 J/g°C, can be placed in the water so that the equilibrium temperature is
26.5°C?

You put water into a foam-cup calorimeter and find that its initial temperature is 25.0°C. What is the mass
of the water if 14.0 grams of 125°C nickel, CNi = 0.444 J/g°C, can be placed in the water so that the
equilibrium temperature is 27.5°C?
Phase Changes Review
Solid → liquid ________________________
Liquid → solid ___________________
Liquid → gas ________________________
Gas → liquid _____________________
Solid → gas ________________________
Gas → solid ______________________
Energy and Phase Changes
q = m Hf
q = m Hv
Hf = latent heat of fusion (J/g) ; Hv = latent heat of vaporization (J/g)
135
Unit 12- Kinetics and Thermochemistry
Heating Curve for Water
120
The heating curve has _____ distinct regions.
Steam
Water and
Steam
100
The ______________________ lines are where
80
phase changes occur. Temperature is
60
___________________ during a phase change!
Water
40
20
0
Water
and Ice
Ice
-20
0
40
120
220
760
800
Heat of vaporization (Hv) is the energy required to change one gram of a substance from ________________ to gas.
Heat of fusion (Hf) is the energy required to change one gram of a substance from __________________ to liquid.

How much heat does it take to melt 12.0 g of ice at 0 °C? H f for water is 334 J/g.

How much heat must be removed to condense 5.00 g of steam at 100 °C? H v = 2260 J/g.
Three equations can be used in calculating energy.
q = m Cp ΔT
q = m Hf
q = m Hv
Solving Problems
The total heat equals the sum of all the heats you have to use.
Go in the following order when energy is being added to the system.
1) Heat ice
q = m Cice ∆T
2) Melt ice
q = m Hf
3) Heat water
q = m Cwater ∆T
4) Boil water
q = m Hv
5) Heat steam
q = m Csteam ∆T
Numbers Needed For Energy Problems Involving Water (look up in reference tables)
For ice, specific heat = __________
For water, specific heat = _________
For steam, specific heat = _____________
Heat of vaporization = ___________
Heat of fusion = _____________

How much heat does it take to heat 12 g of ice at – 6 °C to 25 °C water? Round to a whole number.

How much heat does it take to heat 35 g of ice at 0 °C to steam at 150 °C? Round to a whole number.

How much heat does it take to convert 16.0 g of ice to water at 0 °C?

How much heat does it take to heat 21.0 g of water at 12.0°C to water at 75.0°C?

How much heat does it take to heat 14.0 g of water at 12.0°C to steam at 122.0°C?
Entropy
Entropy (S) is a measure of the ______________________ or randomness of the particles that make up a system.
Spontaneous processes always result in a(n) _____________________ in the entropy of the universe. Entropy of a
solid < Entropy of a liquid << Entropy of a gas
A solid has an orderly arrangement. A liquid has the molecules
next to each other. A gas has molecules moving all over the place.
136
Unit 12- Kinetics and Thermochemistry
Several factors affect the change in entropy of a system.
1.
Changes of state. Entropy ___________________ when a solid changes to a liquid and when a liquid
changes to a gas because these changes of state result in freer movement of the particles.
2.
Dissolving of a gas in a solvent. When a gas is dissolved in a liquid or solid solvent, the motion and
randomness of the particles are limited and the entropy of the gas _____________.
3.
Change in the number of gaseous particles. When the number of gaseous particles increases, the entropy
of the system usually ___________________ because more random arrangements are possible.
4.
Dissolving of a solid or liquid to form a solution. When solute particles become dispersed in a solvent, the
disorder of the particles and the entropy of the system usually ________________.
5.
Change in temperature. A temperature increase results in increased disorder of the particles and a(n)
____________________ in entropy.

Predict the sign of ∆Ssystem for: O2 (g)  O2 (aq)

Predict the sign of ∆Ssystem for: C6H6 (s)  C6H6 (l)

Predict the sign of ∆Ssystem for: C (s) + CO2 (g)  2 CO (g)
Homework / Practice
1.
If 200. g of water at 20.0 °C absorbs 41840 J of heat, what will its final temperature be?
2.
Aluminum has a specific heat of 0.900 J/g.°C. How much energy is needed to raise the temperature of a
625 g block of aluminum from 30.7 °C to 82.1 °C?
3.
The products in a reaction have a total heat content of 458 kJ and the reactants have a total heat content of
656 kJ. What is the value of ∆H?
4.
If a reaction is exothermic, are the products or the reactants more stable?
5.
The heat of combustion of ethene gas, C2H2, is - 1301.1 kJ/mol of ethene.
(a) If 0.250 mol of ethene react, how much energy is released?
(b) How many grams of ethene are needed to react to release 3900. kJ of heat?
(c) Write a balanced reaction for the combustion of ethene.
6. List 4 factors that can speed up a chemical reaction.
137
Unit 12- Kinetics and Thermochemistry
7. For each of the following examples, state whether the change in entropy is positive, negative or remains the
same.
(a) HCl (l) → HCl (g)
(c) 2 NH3 (g) → N2 (g) + 3 H2 (g)
(b) C6H12O6 (aq) → C6H12O6 (s)
(d) 3 C2H4 (g) → C6H12 (l)
8. Consider the following equilibrium equation: H2O (g) + C (s) → H2 (g) + CO (g) + heat energy. Will
the reaction rate increase or decrease when each of the following occurs?
(a) extra CO gas is introduced
(b) a catalyst is introduced
(c) the temperature of the system is lowered
(d) the pressure of the system is increased
9. Convert from one unit to the other:
(a) 1.69 Joules to calories
(i) 556 kilocalories to cal
(b) 0.3587 J to cal
(j) 6.78 kilocalories to kilojoules
(c) 820.1 J to kilocalories
(k) 59.6 calories to kcal
(d) 68 calories to kilocalories
(l) 449.6 joules to kilojoules
(e) 423 calories to kilocalories
(m) 9.806 kJ to J
(f) 20.0 calories to Joules
(n) 5.567 cal to J
(g) 252 cal to J
(o) 5467.9 kcal to J
(h) 2.45 kilocalories to calories
10. How much energy must be absorbed by 20.0 g of water to increase its temperature from 283.0 °C to 303.0
°C?
11. When 15.0 g of steam drops in temperature from 275.0 °C to 250.0 °C, how much heat energy is released?
12. How much energy is required to heat 120.0 g of water from 2.0 °C to 24.0 °C?
13. If 720.0 g of steam at 400.0 °C absorbs 800.0 kJ of heat energy, what will be its increase in temperature?
14. How much heat (in kJ) is given out when 85.0 g of lead cools from 200.0 °C to 10.0 °C?
15. If it takes 41.72 joules to heat a piece of gold weighing 18.69 g from 10.0 °C to 27.0 °C, what is the
specific heat of the gold?
16. A certain mass of water was heated with 41,840 Joules, raising its temperature from 22.0 °C to 28.5 °C.
Find the mass of water.
17. Determine the energy required (in Joules) when the temperature of 3.21 grams of liquid water increases by
4.0 °C.
138
Unit 12- Kinetics and Thermochemistry
18. Determine the temperature change when:
(a) 20.0 g of water is heated from 16.8 °C to 39.2 °C.
(b) 35.0 g of water is cooled from 56.5 °C to 5.9 °C.
(c) 50.0 g liquid water is heated from 0.0 °C to 100.0 °C.
(d) 25.0 g of ice is warmed from -25.0 °C to 0.0 °C, but does not melt.
(e) 30.0 g of steam heats from 373.2 K to 405.0 K.
19. Determine the energy needed (in Joules) when 55.6 grams of water at 43.2 °C is heated to 78.1°C.
20. Determine the energy required (in kilojoules) when cooling 456.2 grams of water at 89.2 °C to a final
temperature of 5.9 °C.
21. Determine the specific heat of a 150.0 gram object that requires 62.0 cal of energy to raise its temperature
12.0 °C.
22. Determine the energy required to raise the temperature of 46.2 grams of aluminum from 35.8 °C to 78.1 °C.
23. Determine the energy required to:
(a) melt 5.62 moles of ice at 0 °C.
(b) melt 74.5 grams of ice at 0 °C.
(c) boil 0.345 moles of water at 100.0 °C.
(d) boil 43.89 grams of water at 100.0 °C.
(e) Convert 16.2 grams of ice to liquid water.
(f) Convert 5.8 grams of water to steam
(g) Convert 98.2 grams of water to ice.
(h) Convert 52.6 grams of steam to water
(i) Convert 34.0 grams of water at 20.0 °C to steam at 100.0 °C.
(j) Convert 125.0 grams of ice at 0.0 °C to steam at 100.0 °C.
(k) Convert 25.9 grams of steam at 100.0 °C.to ice at 0.0 °C.
24. How many degrees of temperature rise will occur when a 25.0-g block of aluminum absorbs 10.0 kJ of
heat?
Kinetics and Thermochemistry Practice Test
Directions: Match the terms below with their correct definitions. (1-4)
a.
calorimeter
c.
enthalpy
b.
thermochemistry
d.
system
1.
The study of heat changes that accompany chemical reactions and phase changes
2.
The specific part of the universe that contains the reaction or process you wish to study
3.
The heat content of a system at constant pressure
4.
An insulated device used to measure the amount of heat absorbed or released during a chemical or physical
process
139
Unit 12- Kinetics and Thermochemistry
5.
Predict the change in entropy (ΔS) for the following reaction (will it increase, decrease, or stay the
same).CH4(g) + 2O2(g)  2H2O(l) + CO2(g)
6.
Which of the following species has the highest entropy at 25°C? Explain your answer.
a)
CH3OH(l)
d) H2O(l)
b) CO(g)
c)
7.
e)
Ni(s)
MgCO3(s)
If the temperature of a 25-g sample of liquid water is raised 40°C, how much heat is absorbed by the
water?
8.
Copper metal has a specific heat of 0.385 J/g·°C , calculate the amount of heat required to raise the
temperature of 38 g of copper from 20.0°C to 575°C.
9.
When a 50.0-g nugget of pure gold is heated from 35.0°C to 50.0°C, it absorbed 5200.0 J of energy.
Find the specific heat of gold.
10.
When 80.0 grams of a certain metal at 90.0 °C was mixed with 100.0 grams of water at 30.0 °C, the
final equilibrium temperature of the mixture was 36.0 °C. What is the specific heat of the metal?
Directions: Use the energy diagram for the rearrangement
reaction of methyl isonitrile to acetonitrile to answer the
following questions. (11-13)
11.
What kind of reaction is represented by this diagram,
endothermic or exothermic?
12.
What does the symbol E represent?
13.
How does a catalyst speed a reaction?
Multiple Choice Practice
14.
As ice cools from 273 K to 263 K, the average kinetic energy of its molecules will
a) decrease
c) remain the same
b) increase
15.
The heat of fusion is defined as the energy required at constant temperature to change 1 unit mass of a
a) gas to a liquid
c) solid to a gas
b) gas to a solid
d) solid to a liquid
140
Unit 12- Kinetics and Thermochemistry
16.
What is the total number of joules of heat energy absorbed by 15 grams of water when it is heated from
30°C to 40°C?
a) 10
c) 150
b) 63
d) 630
17.
How many joules of heat are absorbed when 70.0 grams of water is completely vaporized at its boiling
point?
a) 23, 352
c) 15, 813
b) 7, 000
d) 158, 130
18.
What occurs as potassium nitrate is dissolved in a beaker of water, indicating that the process is
endothermic?
a) The temperature of the solution decreases.
b) The temperature of the solution increases.
c) The solution changes color.
d) The solution gives off a gas.
19.
Given the change of phase: CO2(g) changes to CO2(s), the entropy of the system
a) decreases
c) remains the same
b) increases
20.
A solid is dissolved in a beaker of water. Which observation suggests that the process is endothermic?
a) The solution gives off a gas.
b) The solution changes color.
c) The temperature of the solution decreases.
d) The temperature of the solution increases.
21.
When a catalyst is added to a system at equilibrium, a decrease occurs in the
a) activation energy
b) heat of reaction
c) potential energy of the reactants
d) potential energy of the products
22.
Which statement explains why the speed of some chemical reactions is increased when the surface area
of the reactant is increased?
a) This change increases the density of the reactant particles.
b) This change increases the concentration of the reactant.
c) This change exposes more reactant particles to a possible collision.
d) This change alters the electrical conductivity of the reactant particles.
141
Unit 12- Kinetics and Thermochemistry
23.
Which conditions will increase the rate of chemical reaction?
a) decreased temperature and decreased concentration of reactants?
b) decreased temperature and increased concentration of reactants?
c) increased temperature and decreased concentration of reactants?
d) increased temperature and increased concentration of reactants?
24.
In a chemical reaction, a catalyst changes the
a) potential energy of the products
b) potential energy of the reactants
c) heat of reaction
d) activation energy
25.
The following reaction coordinate diagram represents...
a) an endothermic reaction
b) an exothermic reaction
c) a reaction that is neither
endothermic nor exothermic
d) a reaction in which a
catalyst is used
26.
Catalysts increase the rate of reaction by being consumed.
a) True
b) False
27.
If 1.45 J of heat are added to a 2.00 g sample of aluminum metal and the temperature of the metal
increases by 0.798 oC, what is the specific heat of aluminum?
a) 0.579 J/g deg
c) 1.68 J/g deg
b) 0.909 J/g deg
d) 3.63 J/g deg
28.
Water has a specific heat of 4.184 J/g deg while glass (Pyrex) has a specific heat of 0.780 J/g deg. If
10.0 J of heat is added to 1.00 g of each of these, which will experience the larger increase of
temperature?
a) glass
b) water
c) They both will experience the same change in temperature since only the amount of a
substance relates to the increase in temperature.
29.
How many PopTarts® are needed to convert 1,000.0 g of water at 20.0 oC to steam at 100.0 oC? One
PopTart is equal to 800,000 J of energy!
a) 0.419 PopTart
c) 3.25 PopTarts
b) 2.83 PopTarts
30.
The collision theory states that a reaction is most likely to occur if reactant particles collide with the
proper
a) energy and concentration
b) energy and orientation
c) concentration and orientation
d) pressure and orientation
142
Unit 12- Kinetics and Thermochemistry
31.
What will happen to the rate of reaction when temperature increases?
a) Increase
b) Decrease
c) remains the same
d) increase then decrease
32.
Given the following reaction H+(aq) + OH-(aq) --> H2O(l) if the concentration of the reactants is increased,
the rate of reaction will
a) Increase
c) remains the same
b) Decrease
d) increase then decrease
33.
What does a catalyst decrease when introduced to a reaction?
a) the rate of reaction
b) the energy released
during the reaction
34.
Entropy is a measure of the randomness or disorder of a system.
a) True
c) the activation energy
d) the kinetic energy
b) False
35.
Systems in nature tend to undergo changes toward what kind of energy and entropy?
a) lower energy and lower entropy
b) lower energy and higher entropy
c) higher energy and higher entropy
d) higher energy and lower entropy
36.
Which process is accompanied by a increase in entropy?
a) melting of ice
b) freezing of water
37.
38.
39.
What phase change represents a decrease in entropy?
a) solid to liquid
b) liquid to gas
c)
condensing of water
vapor
c) gas to liquid
d) solid to gas
H2O(g)  H2O(l) The entropy in this equation
a) Increases
b) Decreases
c)
remains the same
Which sample has the lowest entropy?
a) 1 mole of KNO3(l)
b) 1 mole of KNO3(g)
c)
1 mole of H2O(g)
1 mole of KNO3(s)
d)
40.
In a chemical reaction, if the products have more entropy than the reactants, the change in entropy is
negative.
a) True
b) False
41.
What is the change in entropy in the following reaction C + O 2 -> CO2
a) Increases
b) Decreases
c) remains the same
d) not enough information to tell
42.
What is the change in entropy in the following reaction 4Al(s) + 3O2(g) -> 2Al2O3(s)
a) Increase
b) Decrease
c) remains the same
d) not enough information to tell
143
Unit 12- Kinetics and Thermochemistry
43.
What is the change in entropy in the following reaction N2(g) + O2(g) -> 2NO(g)
a) Increase
b) Decrease
c) remains the same
d) not enough information to tell
44.
A 47.5 gram sample of a metal at a temperature of 425°C is placed in 1.00 liters (1000 g) of water
which had an initial temperature of 18°C. What is the specific heat capacity of the metal if the final
temperature of the metal and water at equilibrium is 21°C? (The specific heat capacity of water is
4.18 J/°C·g.
a) 0.03 J/°C·g
d) 0.65 J/°C·g
b) 1.47 J/°C·g
e) 12.54 J/°C·g
c) -0.75 J/°C·g
45.
A bomb calorimeter has a heat capacity of 3.18 kJ/K. When 0.0038 mol of a gas is burned in the
calorimeter, the temperature increased from 25.0°C to 27.3°C. Calculate the energy released by the
combustion of one mole of the gas.
a) 2.8 x 10-2 kJ
d) -2.8 x 10-2 kJ
b) 7.3 kJ
e) -3.6 x 102 kJ
3
c) -1.9 x 10 kJ
46.
A sample of wood has a heat of combustion of 3.29 kJ/g. What quantity of the wood must be burned to
heat 250. g of water from 18°C to 85°C? Once again, the specific heat capacity of water is 4.18 J/°C·g.
a) 85.1 g
d) 2.13 x 104 g
b) 0.45 g
e) 21.3 g
c) 12.4 g
47.
20.0 mL of pure water at 285 K is mixed with 48 mL of water at 315 K. What is the final temperature
of the mixture in kelvins?
a) 306 K
d) None of these
b) 290 K
e) 275 K
c) 318 K
48.
As a result of an exothermic reaction,
a) the energy of the system is increased and the energy of the surroundings are decreased.
b) the energy of the system and the energy of the surroundings are decreased.
c) the energy of the system is decreased and the energy of the surroundings are increased.
d) the energy of the system and the energy of the surroundings are increased.
e) None of these are accurate
49.
How much energy is needed to convert 50.0 g of ice at -5.00°C to water at 25°C?
a) 22.4 kJ
d) 18.0 J
b) 37.3 kJ
e) 175 J
c) 21.9 kJ
144
Unit 12- Kinetics and Thermochemistry
50.
The specific heat of liquid water is 4.18 J/°C·g and the specific heat of carbon is 0.71 J/°C·g. A
10 gram sample of water and a 10 gram sample of carbon are each subjected to 155 J of heat. If both
samples started at 25°C, which substance will have the higher final temperature, and by what
magnitude?
a) Neither. They will have the same final temperature because the started at the same
temperature and were exposed to the same quantity of heat.
b) carbon, by 18.1°C
c) water, by 3.47°C
d) carbon, by 215°C
e) water, by 28.7°C
51.
The specific heat of iron is 0.450 J/(g·°C). How much heat is required to raise the temperature of a
5.00 gram sample of iron from 22°C to 53°C?
a) -43 J
d) 18 J
b) 155 J
e) 69.8 J
c) 344 J
52.
Which of the following is not an endothermic process?
a) combustion
b) melting
c) crystallization
145
d) vaporization
e) sublimation
Unit 13- Oxidation and Reduction
OXIDATION-REDUCTION
One of the defining characteristics of single-replacement and combustion reactions is that they always involve the
transfer of ____________________________________ from one atom to another. So do many, but not all,
synthesis and decomposition reactions. The ________________________________ number of an atom in an ionic
compound is the number of electrons lost or gained by the atom when it forms ions. Oxidation numbers are tools
that scientists use in written chemical equations to help them keep track of the movement of electrons in a
________________________ reaction. Like some of the other tools you have learned about in chemistry, oxidation
numbers have a specific notation. Oxidation numbers are written with the positive or negative sign
__________________________ the number (+3, +2), whereas ionic charge is written with the sign after the number
(3+, 2+). Oxidation number: +3. Ionic charge: 3+. In order to understand all kinds of redox reactions, you must
have a way to determine the oxidation number of the atoms involved in the reaction. Chemists use a set of rules to
make this determination easier.
1.
The oxidation number of an uncombined atom is _______________________. This is true for elements
that exist as polyatomic molecules such as O2, Cl2, H2, N2, S8.
2.
The oxidation number of a _____________________ ion is equal to the charge on the ion. For example,
the oxidation number of a Ca2+ ion is +2, and the oxidation number of a Br– ion is –1.
3.
The oxidation number of the _______________________ electronegative atom in a molecule or a complex
ion is the same as the charge it would have if it were an ion.
•
In ammonia (NH3), for example, nitrogen is more electronegative than hydrogen, meaning that it
attracts electrons more strongly than does hydrogen. So nitrogen is assigned an oxidation number
of –3, as if it had gained three electrons to complete an octet. In the compound silicon
tetrachloride (SiCl4), chlorine is more electronegative than silicon, so each chlorine has an
oxidation number of –1 as if it had taken an electron from silicon. The silicon atom is given an
oxidation number of +4 as if it had lost electrons to the four chlorine atoms.
4.
The most electronegative element, _________________________________, always has an oxidation
number of –1 when it is bonded to another element.
5.
The oxidation number of oxygen in compounds is always ___________, except in peroxides, such as
hydrogen peroxide (H2O2), where it is –1. When it is bonded to fluorine, the only element more
electronegative than oxygen, the oxidation number of oxygen is +2.
6.
The oxidation number of hydrogen in most of its compounds is _____________. The exception to this rule
occurs when hydrogen is bonded to less electronegative metals to form _______________________ such
as LiH, NaH, CaH2, and AlH3. In these compounds, hydrogen’s oxidation number is –1 because it attracts
electrons more strongly than does the metal atom.
146
Unit 13- Oxidation and Reduction
7.
The metals of groups 1A and 2A and aluminum in group 3A form compounds in which the metal atom
always has a ___________________________ oxidation number equal to the number of its valence
electrons (+1, +2, and +3, respectively).
8.
The sum of the oxidation numbers in a neutral compound is ______________________.
9.
The sum of the oxidation numbers of the atoms in a ________________________________ ion is equal to
the charge on the ion.
Determining Oxidation Numbers
Determine the oxidation number of each element in the following compounds or ions. Assign the known oxidation
numbers to their elements, set the sum of all oxidation numbers to zero or to the ion charge, and solve for the
unknown oxidation number. (Let nelement = oxidation number of the element in question.)
•
KClO3 (potassium chlorate) Potassium chlorate is a _________________________ salt, so oxidation
numbers must add up to zero. K’s oxid. # + Cl’s oxid. # + 3 (O’s oxid. #) = 0 Rule _________ states
that Group 1 metals have a +1 oxidation number in compounds. According to rule ___________, the
oxidation number of oxygen in compounds is –2. (+1) + (nCl) + 3 (-2) = 0. Chlorine’s oxidation number
must equal _______.
•
SO32– (sulfite ion) Sulfite ion has a charge of 2–, so oxidation numbers must add up to ______. S’s oxid. #
+ 3 (O’s oxid. #) = –2. According to rule 5, the oxidation number of oxygen in compounds is –2. (nsulfur)
+ 3 (-2) = –2. Sulfur’s oxidation number must equal _______.
•
Determine the oxidation number of each element in the following compound or ion.
a)
Na2CO3 ___________________________________________
b) H2SO4 ___________________________________________
c)
Mg3P2 ___________________________________________
d) SrH2 ___________________________________________
e)
PO4-3 ___________________________________________
f)
ClO4-1 ___________________________________________
g) C4H10 ___________________________________________
Electron Transfer and Redox Reactions
Complete Chemical Equation: 2 Na (s) + Cl2 (g)  2 NaCl (s)
Net Ionic Equation: 2 Na (s) + Cl2 (g)  2 Na+ + 2 Cl•
Solid sodium, on the reactant side, has an oxidation number of _______. The sodium ion product has an
oxidation number of _______.
Oxidation is defined as the ____________________ of electrons from atoms of a substance. When an element’s
oxidation number ________________________, the element is said to be oxidized.
•
Sodium’s oxidation number changes from 0 to +1. Sodium is oxidized because it loses ______ electron.
To state this reaction more clearly, Na → Na+ + e-. For oxidation to take place, the electrons lost by the
147
Unit 13- Oxidation and Reduction
substance that is oxidized must be accepted by atoms or ions of another substance. In other words, there
must be an accompanying process that involves the ________________ of electrons.
_____________________________ is defined as the gain of electrons by atoms of a substance.
Net Ionic Equation: 2 Na (s) + Cl2 (g)  2 Na+ + 2 Cl•
Chlorine gas, on the reactant side, has an oxidation number of _______. The chlorine ion product has an
oxidation number of _______. Chlorine’s oxidation number changes from 0 to -1. Chlorine is reduced
because it _________________ one electron. Following our sodium chloride example further, the
reduction reaction that accompanies the oxidation of sodium is the reduction of chlorine.
Cl2 + 2 e- → 2 Cl-. In this reaction, an electron from each of two sodium atoms is transferred to the Cl2
molecule to form two Cl– ions.
Complete Chemical Equation: 2 Mg (s) + O2 (g)  2 MgO (s)
Net Ionic Equation: 2 Mg (s) + O2 (g)  2 Mg2+ + 2 O2•
Solid magnesium, on the reactant side, has an oxidation number of ____. The magnesium ion product has
an oxidation number of ____. Magnesium’s oxidation number changes from 0 to +2. Magnesium is
oxidized because it loses _________ electrons. To state this reaction more clearly,
_______________________________________________. Oxygen gas, on the reactant side, has an
oxidation number of _______. The oxygen ion product has an oxidation number of _______. Oxygen’s
oxidation number changes from 0 to -2. Oxygen is reduced because it gains _________ electrons.
Following our magnesium oxide example further, the reduction reaction that accompanies the oxidation of
magnesium is the reduction of oxygen.
______________________________________________________ When magnesium reacts with oxygen,
each magnesium atom transfers two electrons to each oxygen atom. The two magnesium atoms become
Mg2+ ions and the two oxygen atoms become O2– ions (oxide ions).
A reaction in which electrons are ______________________________ from one atom to another is called an
oxidation–reduction reaction. For simplicity, chemists often refer to oxidation–reduction reactions as ____________________ reactions.
Now consider the single-replacement reaction in which chlorine in an aqueous solution replaces bromine from an
aqueous solution of potassium bromide.
Complete Chemical Equation: 2 KBr (aq) + Cl2 (aq)  2 KCl (aq) + Br2 (aq)
Net Ionic Equation: 2 Br- (aq) + Cl2 (aq)  2 Cl- (aq) + Br2 (aq)
•
Note that chlorine “steals” electrons from bromide ions to become _______________________ ions.
When the bromide ions lose their extra electrons, the two bromine atoms form a covalent bond with each
other to produce Br2 molecules. The result of this reaction, the characteristic color of elemental bromine in
148
Unit 13- Oxidation and Reduction
solution. The formation of the covalent bond by ________________________ of electrons also is an
oxidation–reduction reaction.
•
Note that there is no change in the oxidation number of potassium. The potassium ion takes no part in the
reaction and as a result DOES NOT appear in the net ionic equation. It is called a
_________________________________ ion.
Can oxidation occur without reduction? By our definitions, oxidation and reduction are complementary processes;
oxidation cannot occur unless reduction also occurs.
•
Determine which element is oxidized and which element is reduced in the following reaction. (Provide the
oxidation numbers.)
a)
3Cu + 2NO3- + 8H+  3Cu2+ + 2NO + H2O
b) 2KNO3  2KNO2 + O2
c)
H2 + CuO  Cu + H2O
d) H2S + NO3-  SO42- + NO2
How DO Oxidation and Reduction Differ?
LEO the lion says GER or, for short, LEO GER
This phrase will help you remember that Loss of Electrons is Oxidation, and Gain of Electrons is Reduction.
•
Determine how many electrons are lost and how many electrons are gained in the following reaction.
(Provide the oxidation numbers.)
a)
Fe2+ + MnO4-  Fe3+ + Mn2+
b) Na2SnO2 + Bi(OH)3  Bi + Na2SnO3 + H2O
c)
2Fe + O2 + 2H2O  2Fe(OH)2
Redox and Electronegativity
The chemistry of oxidation–reduction reactions is not limited to atoms of an element changing to ions or the reverse.
Example: N2 (g) + 3H2 (g) → 2NH3 (g)
•
Nitrogen’s oxidation number changes from _______ to _______; therefore nitrogen is
____________________________. For the purpose of studying oxidation–reduction reactions, the more
electronegative atom (nitrogen) is treated as if it had been reduced by gaining electrons from the other
atom. Hydrogen’s oxidation number changes from _______ to _______; therefore hydrogen is
_______________________. The less electronegative atom (hydrogen) is treated as if it had been
oxidized by losing electrons.
149
Unit 13- Oxidation and Reduction
Half-Reactions
The oxidation process and the reduction process of a redox reaction can each be expressed as a half-reaction.
•
For example, consider the unbalanced equation for the formation of aluminum bromide.
Al + Br2 → AlBr3
•
The oxidation half-reaction shows the loss of electrons by aluminum. Al → Al 3+ + 3ehalf-reaction shows the gain of electrons by bromine. Br 2 + 2e- →2Br

The reduction
-
Write balanced half-reactions for each of the following molecular equations.
a)
K + S → K2S
b) Ca + Cl2 → CaCl2
c)
AlBr3 → Al + Br2
d) Zn3P2 → Zn + P
Practical Applications of Redox
A common application of redox chemistry is to remove ____________________________ from metal objects, such
as a silver cup. When you add chlorine ________________________ to your laundry to whiten clothes, you are
using an aqueous solution of sodium hypochlorite (NaClO). Chlorine is ________________________ in the
process. Hydrogen peroxide (H2O2) can be used as an antiseptic because it _____________________________
some of the vital biomolecules of germs, or to lighten hair because it oxidizes the dark pigment of the hair. When
you put a battery into a flashlight, radio, or CD player, you complete the electrical circuit of a galvanic cell(s),
providing a path for the electrons to flow through as they move from the site of ____________________________
to the site of the reduction. Electroplating is important in protecting objects from
______________________________ and uses the process of redox. Reduction of silver ions onto cheaper metals
forms _______________________________.
Homework / Practice
What is the oxidation number of .
1) N in NO3¯
2
5) Fe in Fe2O3
2) C in CO3 ¯
6) Pb in PbOH
3) Cr in CrO42¯
7) V in VO2+
4) Cr in Cr2O72¯
8) V in VO2+
9) Mn in MnO4¯
+
10) Mn in MnO42¯
Balance each half-reaction for atoms and charge:
11) Cl2  Cl¯
13) Fe2+  Fe3+
12) Sn  Sn2+
14) I3¯  I¯
150
Unit 13- Oxidation and Reduction
Separate each of these redox reactions into their two half-reactions (but do not balance):
15) Sn + NO3¯  SnO2 + NO2
16) HClO + Co  Cl2 + Co2+
17) NO2  NO3¯ + NO
In each of the following, identify what is being oxidized and what is being reduced.
18) Cu+2 + Fe  Fe+2 + Cu
19) Co + Sn+2  Co+2 + Sn
20) 2 Cr + 3 Sn+2  2 Cr+3 + 3 Sn
21) CH4 + 2 O2  CO2 + 2H2O
22) Mg + CO2  2MgO + C
23) 4H+1 + NO3-1 + 3 Fe+2  2 H2O + NO + 3 Fe+3
24) S-2 + 2 NO3-1 + 4H+1  SO2 + NO2 + 2H2O
25) 2 NH4+1 + 2 NO3-1  2 N2 + 4 H2O + O2
Oxidation and Reduction Practice Test
1.
Explain what takes place in an oxidation-reduction reaction.
Directions: Identify the oxidation number for the listed element within each compound or ion
2.
3.
4.
Cr in Cr2O72¯ __________
Pb in PbOH
V in VO
+
2+
__________
__________
5.
Mn in Mn2O7 __________
6.
S in H2SO4
__________
7.
BrO3-
__________
Br in
Directions: Label the oxidation states on all of the elements and trace the paths of oxidation and reduction.
8.
HNO3 + 3Cu2O  6Cu(NO3)2 + 2NO + 7H2O
9.
H2SO4 + HI  H2S + I2 + H2O
Directions: Balance each half-reaction for atoms and charge:
12. Ag  Ag1+
10. Sn ---> Sn2+
11. Cr5+  Cr3+
Directions: Separate each of these redox reactions into their two half-reactions:
15. 2Ca + O2  2CaO
13. Sn + NO3¯ ---> SnO2 + NO2
14. C2H4 + 3O2  2CO2 + 2H2O
151
Unit 13- Oxidation and Reduction
Multiple Choice Practice
16. The oxidation number of sulfur in H2SO4 is
a. +2
c.
b. +3
d.
+4
+6
e.
+8
17. In the following reaction, the oxidation numbers on sulfur are:
NaCl + H2SO4 + MnO2 → Na2SO4 + MnSO4 + H2O + Cl2
a.
b.
c.
+6, +6 and +4 respectively
+6, +6 and +6 respectively
+4, +6 and +4 respectively
d.
e.
+4, +4 and +4 respectively
+4, +6 and +6 respectively
18. NaCl + H2SO4 + MnO2 → Na2SO4 + MnSO4 + H2O + Cl2
Which of the reactants contains an element that is oxidized?
a. H+
c. Na+
4+
b. Mn
d. Cl-
e.
O2-
19. In the balanced molecular equation for the neutralization reaction between phosphoric acid and potassium
hydroxide, the products are:
a. KPO4 + H3OH
d. K+(aq) + PO43-(aq) + 3H+ + 3OHb. 3H2O
e. K3PO4 + 3H2O
c. None of these
20. In the reaction 4Al + 3O2 → 2Al2O3 the substance oxidized is:
a. O2b. O2
e. None of these, because this is not a redox reaction
c.
d.
Al3+
Al
21. In which of the following compounds does sulfur have an oxidation state of +4?
a. H2SO4
d. MgSO4
b. H2S
e. SO3
c. H2SO3
152
Reference Materials
Name
Avogadro’s number
Value
6.022 x1023 particles/mole
Gas constant (R)
0.0821 L atm/ mole K
62.4 L mmHg/ mole K
8.314 L kPa/ mole K
Standard pressure
1.00 atm = 101.3 kPa =760.0 mm Hg = 760.0 torr
Standard temperature
0°C or 273K
Volume of 1 mole of any gas at STP
22.4 L
Thermodynamic Constants
Symbol
Value
Heat of fusion of water
Hf (water)
334 J/g
Heat of vaporization of water
Hv (water)
2,260 J/g
Specific heat of water
Cp (water)
2.05 J/g°C for ice,
2.02 J/g°C for steam,
4.18 J/g°C for liquid
Metal
Specific Heat J/ g°C
Density (g/cm3)
Melting Point (°C)
Aluminum
0.897
2.702
660
Copper
0.385
8.92
1083
Gold
0.129
19.31
1064
Iron
0.449
7.86
1535
Lead
0.129
11.3437
328
Magnesium
1.023
1.74
649
Mercury
0.140
13.5939
-39
Nickel
0.444
8.90
1455
Titanium
0.523
4.5
1660
Zinc
0.388
7.14
420
153
Reference Materials
Organic Substances
Name
Density
Melting Point (°C)
Boiling Point (°C)
Ethanol (CH3CH2OH)
0.7893 g/cm3
−119
79
Glucose (C6H12O6)
1.54 g/cm3
86
Decompose
Hexane (C6H14)
0.6603 g/cm3
−95
69
Methane (CH4)
0.716 g/L
−182
−164
Methanol (CH3OH)
0.7914 g/cm3
−94
65
Sucrose (C12H22O11)
1.27 g/cm3
86
Decompose
Inorganic Substances
Name
Density @ STP
Melting Point (°C)
Boiling Point (°C)
Chlorine
3.21 g/L
−101
−35
Hydrogen
0.0899 g/L
−259
−253
Hydrogen chloride
1.640 g/L
−115
−85
Hydrogen sulfide
1.54 g/L
−85
−61
Nitrogen
1.25 g/L
−210
−196
Nitrogen monoxide
1.34 g/L
−164
−152
Oxygen
1.43 g/L
−218
−183
Sodium carbonate
2.532 g/cm3
851
Decomposes
Sodium chloride
2.165 g/cm3
801
1413
Sulfur dioxide
2.92 g/L
−73
−10
Water
1.00 g/cm3
0
100
154
Reference Materials
Formulas
D
m
V
D = density
K= °C +273
m = mass
P1V1 P2V2

T1
T2
V = volume
Pt = P1 + P2 + P3 + …
K = Kelvin
M1V1=M2V2
P = pressure
PV = nRT
R = gas constant
M 
moles.solute
LitersSolu tion
T = temperature
q = m Cp Δ T
M = molarity
q = mHv
n = number of moles
q = mHf
q = quantity of heat energy
pH + pOH=14
Cp = specific heat
pH= −log[H+]
Hv = heat of vaporization
pOH = −log[OH−]
Hf = heat of fusion
Kw = [H+] [ OH-] = 1 x 10-14
Kw = equilibrium constant for the ionization of water
[H+] = 10−pH
[OH−] = 10−pOH
155
Reference Materials
SOLUBILITY RULES
Soluble:
 All Nitrates, Acetates, Ammonium, and
Group 1 (IA) salts
 All Chlorides, Bromides, and Iodides,
except Silver, Lead, and Mercury(I)
 All Fluorides except Group 2 (IIA),
Lead(II), and Iron(III)
 All Sulfates except Calcium, Strontium,
Barium, Mercury, Lead(II), and Silver
Insoluble (0.10 M or greater):
 All Carbonates and Phosphates except
Group 1 (IA) and Ammonium
 All Hydroxides except Group 1 (IA),
Strontium, Barium, and Ammonium
 All Sulfides except Group 1 (IA),
2 (IIA), and Ammonium
 All Oxides except Group 1 (IA)
Guidelines for Predicting the Products of Selected Types of Chemical Reaction
Key: M = Metal
NM = Nonmetal
1. SYNTHESIS:
a. Formation of binary compound: A + B AB
b. Metal oxide-water reactions: MO + H2O base
c. Nonmetal oxide-water reactions: (NM)O + H2O acid
2. DECOMPOSITION:
a. Binary compounds: AB A + B
b. Metallic carbonates: MCO3 MO + CO2
c. Metallic hydrogen carbonates: MHCO3 MO+ H2O(l) + CO2(g)
d. Metallic hydroxides: MOH MO + H2O
e. Metallic chlorates: MClO3 MCl + O2
f. Oxyacids decompose to nonmetal oxides and water: acid (NM)O + H2O
3. SINGLE REPLACEMENT:
a. Metal-metal replacement: A + BC AC + B
b. Active metal replaces H from water: M + H2O MOH + H2
c. Active metal replaces H from acid: M + HX MX + H2
d. Halide-Halide replacement: D + BC BD + C
4. DOUBLE REPLACEMENT: AB + CD AD + CB
a. Formation of a precipitate from solution
b. Acid-Base neutralization reaction
5. COMBUSTION REACTION
Hydrocarbon + oxygen carbon dioxide + water
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Reference Materials
ACTIVITY SERIES of Halogens:
F2
Cl2
Br2
I2
ACTIVITY SERIES of Metals
Li
NH4
+
Polyatomic Ions
Ammonium
BrO3-
Bromate
CN-
Cyanide
C2H3O2-
Acetate
Rb
(CH3COO-)
K
ClO4-
Perchlorate
Ba
ClO3-
Chlorate
Sr
ClO2-
Chlorite
ClO-
Hypochlorite
IO3-
Iodate
Mg
MnO4-
Permanganate
Al
NO3-
Nitrate
Mn
NO2-
Nitrite
Zn
OH-
Hydroxide
Cr
HCO3-
Hydrogen carbonate
HSO4-
Hydrogen sulfate
Cd
SCN-
Thiocyanate
Co
CO32-
Carbonate
Ni
Cr2O72-
Dichromate
Sn
CrO42-
Chromate
SO42-
Sulfate
[H2 ]
SO32-
Sulfite
Sb
PO43-
Phosphate
Ca
Na
Fe
Pb
Replaces hydrogen
from cold water
Replaces hydrogen from steam
Replaces hydrogen from acids
Bi
Cu
Reacts with oxygen to form oxides
Hg
Ag
Pt
Au
157
Reference Materials
158
Reference Materials
159
Reference Materials
160