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.
Steps of scientific method with description:
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
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
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).
List the examples of each from in class notes:
Homogeneous:
Heterogeneous
Pure Substance
_______________________
_______________________
_____________________
_______________________
_______________________
_____________________
_______________________
_______________________
_____________________
_______________________
_______________________
_____________________
_______________________
_______________________
_____________________
_______________________
_______________________
_____________________
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.
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
__________
E. lead (II) flouide
__________
B. sodium sulfate
__________
F. barium hydroxide
__________
C. calcium carbonate
__________
G. copper (II) sulfide
__________
D. potassium oxide
__________
H. silver chloride
__________
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).
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 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.
6. 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.
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.
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.
_____________________
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.
Unit 1 Practice / Homework
Density: Use the ref. packet to 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: Use the ref. packet to 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: Use the ref. packet to 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
Identify each of the following as a chemical or physical change
23. Melts
25. Dissolves
27. Tarnishis
24. Burns
26. Rips
28. Shatters
25. Determine the following for the fluorine atom.
a) number of protons
b) number of neutrons
d) atomic number
e) mass number
c) number of electrons
26. 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) element
27. If an element has 91 protons and 140 neutrons, what is the
a) atomic number
b) mass number
c) number of electrons
d) element
28. If an element has 78 electrons and 117 neutrons what is the
a) atomic number
b) mass number
c) number of protons
d) element
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 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?
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
______________________
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.
d.
methane
methanol
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. 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
24. 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
Unit 2- Atomic Theory and Structure
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. The measurable factor in an experiment is known as the:
a. Control
b. independent variable
c.
d.
constant
dependent variable
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 elements 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 & HISTORY
In 1897, a British physicist, J.J. Thomson, discovered that this solid-ball model was not accurate.
JJ Thomson
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
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.
Rutherford:____________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
NAME
SYMBOL
CHARGE
RELATIVE MASS
1/2000
proton
no
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.
Practice/Homework
Isotopes and Subatomic Particles
Complete the chart
Isotope symbol
207
82
Atomic #
Mass
Protons
8
Neutrons
9
Electrons
77
54
12
13
Pb
38
50
238
92
U
75
33
As
32
16
S
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
Unit 2- Atomic Theory and Structure
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?
Calculate the number of particles (atoms, ions or molecules) in each of the following.
a) 3.4 moles Na2S
b) 0.0020 moles Zn
c) 1.77 x 10-11 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
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)
c) bromine (Br)
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
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?
d) Ca(NO3)2
Unit 2- Atomic Theory and Structure
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 CO2at STP?
3. What is the volume of 8.8 g of CH4 gas at STP?
2. How many moles is 5.67 L of O2 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
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
10. Calculate moles in 168.0 g of HgS
8.
Calculate grams in 3.0000 moles of CO 2
11. Calculate moles in 510.0 g of Al2S3
9.
Calculate number of moles in 32.0 g of CH4
12. How many moles are in 27.00 g of H2O
Unit 2- Atomic Theory and Structure
13. What is the mass of 2.55 moles Cu2CrO4
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
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
Unit 2- Atomic Theory and Structure
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.

When nitrogen-14 is bombarded with a neutron, a proton is produced. The balanced nuclear equation can be
written as: _________________________________________________________

Thorium-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.
____________________________________________________________
Unit 2- Atomic Theory and Structure

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.

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
_______________________. 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.
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.
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
e.
13
16. What is the mass in grams of 1 mole of KAl(SO4)2•12H2O
a. 132
c. 394
b. 180
d. 474
e.
516
Unit 2- Atomic Theory and Structure
17. 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
18. 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
19. A proton has approximately the same mass as
a. a neutron
b. an alpha particle
c.
d.
a beta particle
an electron
20. A neutron has approximately the same mass as a
a. an alpha particle
b. a beta particle
c.
d.
an electron
a proton
21. 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
22. Which atom contains exactly 15 protons?
a. P-32
b. S-32
c.
d.
O-15
N-15
23. 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
24. 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
d.
11
25. What is the gram formula mass of K2CO3?
a. 138 g
b. 106 g
c.
d.
99 g
67 g
26. What is the total number of atoms contained in 2.00 moles of nickel?
a. 58.9
c.
b. 118
d.
6.02 x 1023
1.2 x 1024
27. What is the mass in grams of 3.0 x 1023 molecules of CO2?
a. 22 g
b. 44 g
66 g
88 g
c.
d.
28. 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
29. The total number of atoms in a formula unit of aluminum dichromate, Al2(Cr2O7)3 is:
a. 5
b. 29
c. 17
30. 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
Unit 2- Atomic Theory and Structure
31. 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
32. 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
33. 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
34. 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
35. 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
36. 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
37. 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
38. How many oxygen atoms are there in one formula unit of Al 2(SO4)3?
a. 3
b. 4
c.
7
d.
12
39. Which of the following arrangements represent different isotopes of the same element?
i. 12 protons, 11 neutrons, 12 electrons
ii. 11 protons, 12 neutrons, 11 electrons
iii. 10 protons, 12 neutrons, 12 electrons
iv. 11 protons, 12 neutrons, 10 electrons
v. 12 protons, 12 neutrons, 12 electrons
a. 1 and 5
d. all of these qualify
b. 2 and 4
e. None of these qualify
c. 2, 3, 4 and 5
40. 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
41. An alpha () particle is essentially a ____________________ nucleus.
a. plutonium
c. hydrogen
b. helium
d. uranium
e.
carbon-12
Unit 2- Atomic Theory and Structure
42. 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
43. 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
44. 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
45. Identify the missing particle in the following nuclear reaction:
a.
37
18
Ar
b.
38
18
Ar
37
19
K → _____ +
36
c. 18 Ar
0
1
e
d.
46. 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
47. 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
48. 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
49. An electron emitted from the nucleus during some kinds of radioactive decay is known as:
a. A gamma ray
d. An alpha () particle
b. A positron
c. A beta () particle
50. A process in which a very heavy nucleus splits into more-stable nuclei of intermediate mass is called:
a. nuclear fission
c. a chain reaction
b. radiocarbon dating
d. nuclear fusion
37
20
Ca
Unit 3- Electrons and Periodicity
ELECTRONS IN ATOMS
LIGHT
Light is a kind of electromagnetic _______________. All forms of electromagnetic radiation move at _______ m/s.
Origin:
Crest:
Trough:
Amplitude:____________________________________________________________________________________
Wavelength:___________________________________________________________________________________
Frequency:____________________________________________________________________________________
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.
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.
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.
Unit 3- Electrons and Periodicity
ATOMIC 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?
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
_________________ 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.
Unit 3- Electrons and Periodicity
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,
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.
Figure 28.5: The Aufbau principle.
Unit 3- Electrons and Periodicity
How do the s , p,d and f orbitals correlate to the periodic table?
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 _________________________________
Unit 3- Electrons and Periodicity
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.
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)
b) bromine (Br)
d) potassium (Al)
e) aluminum (Al)
c) iron (Fe)
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
d) sodium chloride
Unit 3- Electrons and Periodicity
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
_________________ 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
Unit 3- Electrons and Periodicity
____________________ 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) ______
c) Tin (Sn) _____
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.
Unit 3- Electrons and Periodicity
PERIODIC TRENDS
Because the periodic table relates group and period numbers to valence electrons, it’s useful in predicting atomic
structure and, therefore, ______________________ properties.
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+ _________
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) _______
Unit 3- Electrons and Periodicity
Homework / Practice
Write the configuration notation for each of the following elements:
1) sodium
3) bromine
2) iron
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
Write the orbital notation for the following:
13) C
14) Ne
15) S
16) P
17) B
Write configuration notation for atoms containing the following number of electrons:
19) 3
20) 6
21) 8
Draw the Lewis Dot Notation for the following elements
23) Sodium
25) Silver
24) Sulfur
26) Aluminum
18) Na
22) 13
27) Antimony
28) Argon
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 3p6 4s2 3d7
7. Write the complete configuration notation for silver.
Unit 3- Electrons and Periodicity
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.
14. The two main parts of an atom are the
a.
b.
Principle energy levels and energy
sublevels
Nucleus and kernel
c.
d.
Nucleus and energy levels
Planetary electrons and energy
levels
c.
d
15. The sublevel that has only one orbital is identified by the letter
a.
s
b.
p
d.
f
d.
f
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
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, Br 1- ?
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
d.
F
c.
S
d.
F
c.
S
d.
F
c.
S
d.
F
22. Which of the following elements has the greatest electronegativity?
a.
Mg
b.
K
23. Which of the following elements would have the smallest radius
a.
Mg
b.
K
24. Which of following elements has the lowest first ionization energy
a.
Mg
b.
K
25. Which of the following elements is an alkali metal?
a.
Mg
b.
K
Unit 3- Electrons and Periodicity
26. Which element's ionic radius is smaller than its atomic radius?
a.
b.
neon
nitrogen
c.
d.
sodium
sulfur
27. 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.
14, 15, and 16
16, 17, and 18
c.
d.
calcium
phosphorus
c.
d.
calcium
potassium
c.
d.
Ba, Ag, Sn, Xe
Fr, F, O, Rn
28. Which element has the highest first ionization energy?
a.
b.
sodium
aluminum
29. Which of the following elements has the smallest atomic radius?
a.
b.
nickel
cobalt
30. Which set of elements contains a metalloid?
a.
b.
K, Mn, As, Ar
Li, Mg, Ca, Kr
31. 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.
d.
atomic numbers
atomic masses
32. 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
33. 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
34. At STP, which substance is the best conductor of electricity?
a.
b.
nitrogen
neon
35. 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.
36. Which gas is monatomic at STP?
a.
b.
chlorine
fluorine
37. How many valence electrons does an oxygen atom have?
a.
2
b.
6
38. The identity of an element is determined by...
a.
b.
the number of its protons.
the number of its neutrons.
d.
16
Unit 3- Electrons and Periodicity
39. Which of the following atoms has the largest diameter?
a.
F
b.
Cl
c.
Br
d.
I
c.
N
d.
O
40. Which of the following elements has the greatest electronegativity?
a.
Si
b.
P
41. 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
c.
d.
Chlorine
Neon
c.
it has a complete outer energy level
of electrons
all of the above
42. Which of the following is a noble gas?
a.
b.
Sodium
Gold
43. A gas is called "noble" because
a.
b.
it is normally unreactive
it is normally inert
d.
44. Of the following elements, the one that forms cations with varying positive charges is:
a.
b.
Fe
Na
c.
d.
Al
Sr
e.
N
45. An element having the configuration [Xe]6s1 belongs to the Group:
a.
b.
alkali metals
halogens
c.
alkaline earth
metals
e.
d. None of these
noble gases
46. Using the Lewis Dot notation, how many unpaired electrons are there in an atom of tin in its ground state?
a. 4
b. 0
c. 3
d. 2
e.
47. Which of the following particles has the greatest atomic radius?
a.
b.
Al
Si
c.
d.
S
Al3+
e.
P
48. Which of the following forms of electromagnetic radiation has the shortest wavelength?
a.
b.
ultraviolet
radio waves
c.
d.
infrared
visible light
e.
microwaves
49. For which of the following transitions does the light emitted have the shortest wavelength?
a.
b.
n = 4 to n = 2
n = 2 to n = 1
c.
d.
n = 5 to n = 3
n = 4 to n = 3
e.
n = 3 to n = 2
50. 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
1
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 ___ valence electron
Group 5A ___ valence electron
Group 2A ___ valence electron
Group 6A ___ valence electron
Group 3A ___ valence electron
Group 7A ___ valence electron
Group 4A ___ valence electron
Group 8A ___ valence electron
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), 1s 22s22p63s1, 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), 1s 22s22p63s23p4, 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.
36
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
_______________.
37
Unit 4 – Types of Bonding
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?
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.
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. 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) _____________________.
38
Unit 4 – Types of Bonding
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
O
Charge
Na
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
9) NI3
10) CO2
11) Ni
12) Au
13) LiF
14) Al2O3
15) N2O3
16) Mg3P2
17) CCl4
18) H2O
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
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
_________________.
39
Unit 4 – Types of Bonding
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
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. 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
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 ____________________.
40
Unit 4 – Types of Bonding
TETRAHEDRAL
Consider CH4. which 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). 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
Consider 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
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
Consider 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 ____________.
41
Unit 4 – Types of Bonding
TRIGONAL PLANAR
Consider 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
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.
2.
3.
CCl4
4.
SiO2
5.
H2S
BF3
NF3
42
Unit 4 – Types of Bonding
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.
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. C 2H6 (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
43
Unit 4 – Types of Bonding
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?
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) H2 _____________________________
6) CH3Cl ___________________________
7) HF _____________________________
8) HBr ____________________________
44
Unit 4 – Types of Bonding
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?
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 iron in a sample of iron?
a. Hydrogen bond
b. Ionic bond
c. Polar covalent bond
d.
e.
Nonpolar covalent bond
Metallic bond
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
e.
CO2
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
c. gain electrons and decrease in size
b. lose electrons and increase 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
45
Unit 4 – Types of Bonding
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
b. CCl4
c.
d.
MgCl2
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
28. Which of the following molecules must contain at least one double bond
a. H2O
d.
b. CCl4
e.
c. H2O2
e.
0
CH3I
CH3COOH
29. 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
30. 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
31. 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
32. The Lewis structure for hydrogen cyanide is:
a.
d.
b.
c.
e.
33. In the Lewis structure for CH2Cl2, the number of unshared electron pairs is:
a. 10
c. 2
b. 8
d. 4
46
e.
6
Unit 4 – Types of Bonding
34. The only intermolecular forces existing between oxygen molecules are:
a. ion-ion attractive forces
d.
b. hydrogen bonding forces
e.
c. permanent dipole forces
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
b. obviously polyatomic
c. an oxide ion
d. a neon ion
e. none of these are correct
47
Unit 5 – Naming and Formula Writing
NAMING COMPOUNDS AND WRITING FORMULAS
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.
F1- has gained _________ electron. O2- has gained __________ electrons. A ___________________ is a positive
ion. It is formed by __________________ electrons. There are more _____________________ than electrons.
______________________ form cations. K1+ 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
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
Name the following cations.
a) Ca2+ _________________________
b) Al3+ ___________________________
c) Sn4+ _________________________
d) Na+ _________________________
e) Fe3+ _________________________
f) Cu+ _________________________
48
Unit 5 – Naming and Formula Writing
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
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 (III) 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
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
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.
49
Unit 5 – Naming and Formula Writing
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 _________________
d) SnBr2 ________________
g) KCl _________________
b) Ca3P2 ________________
e) Fe2S3 ________________
h) Na3N ________________
c) CuO _________________
f) AlF3 _________________
i) CrN _________________
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 (III) 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 __________________
50
Unit 5 – Naming and Formula Writing
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

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
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
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 ___________________________
51
Unit 5 – Naming and Formula Writing
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
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. H 1+ 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
52
Unit 5 – Naming and Formula Writing
Write the formulas for the following binary acids.
a) hydroiodic 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!

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.

The polyatomic ion must end in –ate since the acid ends in -ic. The acid is made of
Example: acetic acid
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
8.
Mg(OH)2
15. PbO
22. Pb(SO4)2
2.
CuCl2
9.
Ba(CN)2
16. FeCl2
23. P4O6
3.
Ni(C2H3O2)2
10. K2SO4
17. Al2O3
24. N2O3
4.
Co2S3
11. NH4NO3
18. Mg3P2
25. SiF4
5.
CrBr2
12. SbBr3
19. NH4Cl
26. P4S10
6.
AlPO4
13. Ag
20. Fe(NO3)3
27. Cl2O3
7.
CaCO3
14. KCl
21. TiBr3
28. PCl3
53
Unit 5 – Naming and Formula Writing
Write the formula for each of the following compounds
29. iron (II) chloride
37. silver phosphate
45. zinc oxide
30. mercury (I) bromide
38. cobalt (III) nitrite
46. arsenic tribromide
31. chromium (III) oxide
39. ammonium sulfite
47. carbon tetrafluoride
32. manganese (II) nitride
40. carbon monoxide
48. sodium sulfide
33. cobalt (III) phosphide
41. selenium difluoride
49. vanadium (IV) carbonate
34. copper (II) sulfide
42. calcium sulfide
50. tin (II) nitrite
35. potassium perchlorate
43. diphosporus trioxide
51. cobalt (III) oxide
36. aluminum sulfate
44. magnesium fluoride
52. titanium (II) acetate
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/Mass 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.
1.
How many moles are in 56.3 g of Mg?
3.
How many grams are in 3.45 moles of NaCl?
2.
How many moles are in 295 grams of Cr(OH)3?
4.
How many grams are in 1.6 moles of K2CrO4?
Percent Composition
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
54
Unit 5 – Naming and Formula Writing
%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 ___________ for CaCl 2. Finally, use the data to set up a
calculation to determine the percent by mass of an element in the compound.
1.
Determine the percent composition of carbon in sodium acetate (NaC 2H3O2).
2.
Calculate the percent composition aluminum of aluminum oxide (Al 2O3).
3.
Determine the percent composition of oxygen in magnesium nitrate, which has the formula Mg(NO3)2.
4.
Determine the percent composition of sulfur in aluminum sulfate, which has the formula Al 2(SO4)3.
5.
Determine the percent composition of oxygen in zinc nitrite, which has the formula Zn(NO 2)2.
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.
1.
Determine the empirical formula for Tl2C4H4O6.
2.
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.
55
Unit 5 – Naming and Formula Writing
1.
Determine the empirical formula of the following compound: 31.9 g Mg, 27.1 g P
2.
The composition of an unknown acid is 40.00% carbon, 6.71% hydrogen, and 53.29% oxygen. Calculate
the empirical formula for the acid.
3.
The composition of an unknown ionic compound is 60.7% nickel and 39.3% fluorine. Calculate the
empirical formula for the ionic compound.
4.
The composition of a compound is 6.27 g calcium and 1.46 g nitrogen. Calculate the empirical formula for
the compound.
5.
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.
1.
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.
2.
The simplest formula for butane is C2H5 and its molecular mass is about 60.0 g/mol. What is the molecular
formula of butane?
3.
What is its molecular formula of cyanuric chloride, if the empirical formula is CClN and the molecular
mass is 184.5 g/mol?
4.
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?
5.
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?
56
Unit 5 – Naming and Formula Writing
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 known to have an empirical formula of CH and a molar mass of 78.11 g/mol. What is its
molecular formula?
3) 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?
4) 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.)
5) Phentyfloroform contains 57.54% C, 3.45% H, and 39.01% F. What is its empirical formula?
6) Calculate the empirical formula of the compound containing 81.8% C and 18.2% H.
7) 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?
8) 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?
9) 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)
10) A compound consists of 85% silver and 15% florine by mass. What is the empirical formula?
11) A compound consists of 40% calcium, 12% carbon, and 48% oxygen by mass. What is the empirical formula by
mass?
12) Phentyfloroform contains 57.54% C, 3.45% H, and 39.01% F. What is its empirical formula?
13) 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
Directions: Name the compound
10. KCl
11. FeSO4
12. Li2O
4.
5.
6.
nickel (I) oxide
magnesium hydroxide
cesium nitride
13. Cr2S3
14. Ca3N2
15. Fe2S3
16. CuI2
17. PBr3
18. CO2
57
7.
8.
9.
nitrogen trichloride
hydroxic acid
carbon tetrahydride
19. HNO3
Unit 5 – Naming and Formula Writing
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. Determine the empirical formula of N3O6
24. The empirical formula of a compound is CH2. Its molecular mass is 70g/mol. What is its molecular formula?
25. 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.
26. Which of the following is a binary compound?
a. hydrogen sulfide
b. hydrogen sulfate
c. ammonium sulfide
d. ammonium sulfate
27. Which of the following is a binary compound?
a. potassium chloride
b. ammonium chloride
c.
d.
potassium chlorate
ammonium chlorate
28. Which is the correct formula for dinitrogen monoxide?
a. NO
b. N2O
c.
d.
NO2
N2O3
29. Which of the following represents the correct formula for aluminum oxide?
a. AlO
b. Al2O3
c. AlO2
d. Al2O
30. Which of the following is the correct name for NaHCO3?
a. sodium hydrogen carbonate
b. sodium acetate
c.
d.
nitrogen hydrogen carbonate
sodium hydrogen carbon trioxide
31. What is the name of CaCl2?
a. calcium dichloride
b. calcium (II) chloride
c.
d.
monocalcium dichloride
calcium chloride
32. What is the name of Mg(NO3)2
a. Magnesium nitrate
b. Magnesium (II) nitrate
c.
d.
Magnesium dinitrate
Magnesium nitrogen oxide
c.
d.
diphosphorus pentaoxide
phosphorus (III) oxide
33. What is the name of P2O5?
a. phosphorus oxide
b. phosphorus pentaoxide
34. What is the formula for sulfur hexachloride?
a. S5Cl
b. SHCl
c.
35. What is the name of the formula Fe(NO3)2?
a. iron nitrate
b. iron (II) nitrate
SCl5
c.
d.
36. Which of the following is not a type of chemical formula?
a. Empirical
b. Molecular
58
c.
d.
iron dinitrate
iron (III) nitrate
Structural
Parabola
d.
SCl6
Unit 5 – Naming and Formula Writing
37. 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
38. 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
39. When combining with nonmetallic atoms, metallic atoms generally will
a. lose electrons and form negative
c.
ions
b. lose electrons and form positive ions
d.
gain electrons and from negative
ions
gain electrons and form positive ions
40. What is the empirical formula of the compound whose molecular formula is P 4O10?
a. PO
b. PO2
c. P2O5
d.
P8O20
41. What is the percent by mass of oxygen in magnesium oxide, MgO?
a. 20%
b. 40%
c.
d.
60%
50%
42. A compound is 86% carbon and 14% hydrogen by mass. What is the empirical formula for this compound?
a. CH
b. CH2
c. CH3
d. CH4
43. What is the gram formula mass of (NH4)3PO4?
a. 113 g
b. 121 g
c.
d.
149 g
404 g
44. What is the empirical formula of a compound that contains 85% Ag and 15% F by mass?
a. AgF
b. Ag2F
c. AgF2
d. Ag2F2
45. 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
46. 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
47. 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. C2H5OH
c. C5H11OH
e. CH3OH
b. C3H7OH
d. C4H9OH
59
Unit 6 – Chemical Reactions
CHEMICAL REACTIONS
All chemical reactions have two parts:
(1) ___________________________________________________________
(2) ________________________________ ___________________________
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. H 2 , 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.
60
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.
C + O2 → CO
Example:
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
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________

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
61
Unit 6 – Chemical Reactions
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.
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.
62
Unit 6 – Chemical Reactions

Use the Chemistry Reference Tables to write and balance the following decomposition reactions.
a) KClO3 →
b) CaBr2 →
c) Li2CO3 →
d) Cr(OH)2 →
e) NaHCO3 →
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 →
63
Unit 6 – Chemical Reactions
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?

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.
64
Unit 6 – Chemical Reactions
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
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+C
Double replacement
C
Decomposition
CH cpd + O2
Combustion
E+C
Single replacement
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 →
65
Unit 6 – Chemical Reactions
Homework / Practice
Directions: Balance the following equations.
1.
S + O2  SO3
12. Na + H2O  NaOH + H2
2.
P + O2  P2O3
13. Al + H2SO4  Al2(SO4)3 + H2
3.
Na + O2  Na2O
14. Fe(OH)3 + H2SO4  Fe2(SO4)3 + H2O
4.
Al + N2  AlN
15. AgNO3 + K2CrO4  Ag2CrO4 + KNO3
5.
Fe + O2  Fe2O3
16. MgCl2 + NaOH  Mg(OH)2 + NaCl
6.
MgSO4•7H2O  MgSO4 + H2O
17. AgNO3 + H2S  Ag2S + HNO3
7.
NH4NO3  N2O + H2O
18. Al(OH)3 + NaOH  NaAlO2 + H2O
8.
NaNO3  NaNO2 + O2
19. C4H10 + O2  CO2 + H2O
9.
H2O2  H2O + O2
20. C3H6 + O2  CO2 + H2O
10. Al + CuSO4  Al2(SO4)3 + Cu
21. C5H8 + O2  CO2 + H2O
11. ZnS + O2  ZnO + SO2
22. C6H12O6 + O2  CO2 + H2O
Predict the products of the following reactions
23. Na + O2 
35. Al + Pb(NO3)2 --->
24. K2O + H2O 
36. Cl2 + NaI 
25. Al2O3 + H2O 
37. Al + CuCl2 
26. N2O5 + H2O 
38. Br2 + CaI2 
27. Ag2O 
39. Al + HCl 
28. HNO2 
40. Zn + H2SO4 
29. Fe(OH)3 
41. Fe + CuSO4 
30. ZnCO3 
42. Ca(OH)2 + H3PO4 
31. Cs2CO3 
43. AgNO3 + KCl 
32. RbClO3 
44. Na2CO3 + H2SO4 
33. ZnS + O2 
45. Al(OH)3 + HC2H3O2 
34. K + H2O 
46. Cr2(SO3)3 + H2SO4 
Write the balanced equation (including states) and identify the type of reaction:
47. Aqueous solutions of ammonium chloride and lead (II) nitrate produce lead (II) chloride precipitate and
aqueous ammonium nitrate.
48. Iron metal reacts with aqueous silver nitrate to produce aqueous iron (III) nitrate and silver metal.
49. Solid potassium nitrate yields solid potassium nitrite and oxygen gas.
50. Calcium metal reacts with chlorine gas to produce solid calcium chloride.
66
Unit 6 – Chemical Reactions
Chemical Reactions and Balancing Practice Test
1.
2.
3.
4.
5.
Directions: Balance and identify the type of reaction
____K3PO4 + ____Al(NO3)3  ____KNO3 + ____AlPO4
____Fe2O3 + ____Al  ____Fe + ____Al2O3
____NaOH  ____Na2O + ____H2O
____HCl + ____Mg  ____MgCl2 + ____H2
____C2H4 + ____O2  ____CO2 + ____H2O
6.
Write the balanced equation of the synthesis reaction that occurs when iron metal and oxygen gas react to form
iron (III) oxide.
7.
Write the balanced equation of the combustion reaction that occurs when ethane (C2H6) reacts with oxygen to
form carbon dioxide and water.
8.
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
9. _____Na + ______O2 
10. _____Al + _____Pb(NO3)2 
11. _____NaF + _____Br2 
12. In the following unbalanced reaction, what is the coefficient of HOH once the reaction is balanced?
+ KOH --> HOH + K2CO3
a. 1
b. 2
c. 3
d.
H 2CO3
4
13. 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
14. What type of reaction is represented by 2 or more elements forming a compound?
a. Decomposition
c. Combustion
b. Synthesis
d. Single replacement
15. Decomposition is the burning of hydrocarbons in the presence of oxygen.
a. True
b.
False
16. Which equation represents a double replacement reaction?
a. CaCO3  CaO + CO2
b. CH4 + 2O2  CO2 + 2H2O
LiOH + HCl  LiCl + H2O
C3H8 + 5O2  3CO2 + 8H2O
c.
d.
17. MgSO4 + BaCl2  MgCl2 + BaSO4, is an example of what type of chemical reaction?
a. Single replacement
c. Combustion
b. Synthesis
d. Double replacement
18. 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
19. The cation of one compound replaces the cation in another compound in a double replacement reaction.
a. True
b. False
67
Unit 6 – Chemical Reactions
20. 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
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.
c. There is a conservation of only energy, and charge.
d. There is a conservation of mass, energy, and charge.
22. Which one of these chemical reactions is balanced?
a. Na + Cl2  NaCl
b. H2 + O2  H2O
c.
d.
CuCO3  CuO + CO2
KClO3  KCl + O2
23. 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.
24. 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
25. 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. 11
c. 6
e. 14
b. 10
d. 16
26. Which equation represents combustion?
a. 4 Fe + 3 O2  2 Fe2O3
d. Cu + 2 AgNO3  Cu(NO3)2 + 2 Ag
b. 2H2O - 2H2 + O2
c. CH4 + 2O2  CO2 + 2 H2O
27. 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
28. The balanced equation for the complete combustion of benzene, C6H6, is
a. C6H6 + 12 H2O  6 CO2 + 15 H2
b. 2 C6H6 + 9 O2  12 CO + 6 H2O
c. C6H6 + O2  CO2 + H2O
d. 2 C6H6 + 15 O2  12 CO2 + 6 H2O
29. Which equation shows conservation of atoms?
a. H2 + O2  H2O
b. H2 + O2  2H2O
c.
d.
2H2 + O2  2H2O
2H2 + 2O2  2H2O
30. 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. 2, 7, 4, 6
d. 1, 2, 3, 2
b. 1, 3, 2, 3
c. 2, 2, 1, 4
68
Unit 7 – Stoichiometry
STOICHIOMETRY
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
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?
69
Unit 7 – Stoichiometry
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 (32.0) grams of oxygen form ___ x (_______) grams of water.
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 ________________________in a balanced chemical equation.

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)?
70
Unit 7 – Stoichiometry

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 NH 4ClO4?
b) How many grams of NO are produced if 150.0 grams of AlCl3 are also produced?
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
___ x (6.022 x 1023) atoms of sodium and ___ x (___________________) molecules of water form ___ x
(__________________) formula units of sodium hydroxide and ___ x (6.022 x 10 23) molecules of hydrogen gas.
71
Unit 7 – Stoichiometry
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 C4H10?

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.

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?
72
Unit 7 – Stoichiometry
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 moles of oxygen gas are produced in the decomposition of 5.00 g of potassium chlorate?
KClO3  KCl + O2
3.
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
4.
If excess sulfuric acid reacts with 0.2564 moles of sodium chloride, how many grams of hydrochloric acid
are produced? H2SO4 + NaCl  HCl + Na2SO4
5.
How many grams of silver phosphate are produced if 10.0 g of silver acetate react with excess sodium
phosphate? AgC2H3O2 + Na3PO4  Ag3PO4 + NaC2H3O2
6.
What volume of chlorine gas, measured at STP, is needed to produce 10.0 g of potassium permanganate
(KMnO4)? K2MnO4 + Cl2  KMnO4 + KCl
7.
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?
8.
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?
9.
How many liters of Cl2 can be produced from 5.60 mole HCl at STP? 4 HCl + O 2  2 Cl2 + 2 H2O
10. 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
11. 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.
12. 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 16 – 21:
13. ____ Al + ____ O2  ____ Al2O3
14. ____ Cu + ____ AgNO3  ____ Ag + ____ Cu(NO3)2
15. ____ Zn + ____ HCl  ____ ZnCl2 + ____ H2
73
Unit 7 – Stoichiometry
Perform the following calculations:
16. 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?
17. 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?
18. 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?
19. 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?
20. When 9.34 g of zinc react with excess hydrochloric acid how many grams of zinc chloride will be
produced?
21. How many liters of oxygen gas at STP are required to react with 65.3 g of aluminum in the production of
aluminum oxide?
22. Given the following equation: Na2O + H2O ---> 2 NaOH How many grams of Na2O are required to
produce 1.60 x 102 grams of NaOH?
23. 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?
24. 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?
25. 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
74
Unit 7 – Stoichiometry
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
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 P4O10 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.
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 O 2 is ignited?
a. 0.12
c. 0.28
e. 0.72
b. 0.18
d. 0.46
75
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 one mole of the rocket fuel ammonium perchlorate, NH4ClO4 (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
b. 1.0 x 1023
d. 1.20 x 1024
21. How many grams of phosphorous trichloide, 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
22. 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
23. 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
24. 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
25. 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
H2 + Fe3O4  FeO + H2O
c. 0.56 L
26. 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
76
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
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
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.
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
77
Unit 8 – Gas Laws
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 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
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________

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
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
P 1 V1 = P 2 V2

Sketch the PV graph that represents Boyle’s law.

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?
78
Unit 8 – Gas Laws

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?

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
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
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
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________

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?
79
Unit 8 – Gas Laws

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?
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.

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?
80
Unit 8 – Gas Laws

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?
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.
81
Unit 8 – Gas Laws
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.
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.
3.
Does atmospheric pressure increase or decrease as altitude above sea level increases?
4.
Convert the following:
a. 0.200 atm = _____ mm Hg
b. 790 mm Hg = _____ Pa
c.
d.
123 kPa = _____ atm
0.935 atm = ______ torr
5.
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?
6.
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.
7.
Helium gas in a balloon occupies 2.40 L at 400. K. What volume will it occupy at 300 K?
8.
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?
9.
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.
i. The tire and the air inside it are heated to 30 °C by road wear, does the pressure in the tire
increase or decrease?
ii. Because the temperature has doubled, does the pressure double to 110 psi? Why or why
not?
10. 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?
11. 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?
12. 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?
13. A gas occupies 11.2 liters at 0.860 atm. What is the pressure if the volume becomes 15.0 L?
82
Unit 8 – Gas Laws
14. 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?
15. Calculate the decrease in temperature when 2.00 L at 20.0 °C is compressed to 1.00 L.
16. What change in volume results if 60.0 mL of gas is cooled from 33.0 °C to 5.00 °C?
17. A gas occupies 1.00 L at standard temperature. What is the volume at 333.0 °C?
18. 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?
19. 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?
20. 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?
21. 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.
22. 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?
23. How many moles of gas are contained in 890.0 mL at 21.0 °C and 750.0 mm Hg pressure?
24. Calculate the volume 3.00 moles of a gas will occupy at 24.0 °C and 762.4 mm Hg.
25. 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?
26. Find the volume of 2.40 mol of gas whose temperature is 50.0 °C and whose pressure is 2.00 atm.
27. 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?
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?
83
Unit 8 – Gas Laws
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?
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. 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
17. 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
18. 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
19. If temperature is constant, the relationship between pressure and volume is
a. Direct
b. inverse
84
e.
21.1 g/L
Unit 8 – Gas Laws
20. 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
21. 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
22. 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
23. 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
24. 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
25. If pressure is constant, the relationship between temperature and volume is
a. Direct
b. Inverse
26. 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
27. 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
28. If a gases volume is decreased and pressure is constant, its temperature will
a. Increase
b. Decrease
c.
Stay the same
c.
Stay the same
29. 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
30. If the temperature of a gas remains constant but pressure is decreased, the volume will
a. Increase
b. Decrease
85
Unit 8 – Gas Laws
31. Convert 2.3 atm into mmHg
a. 2300 mmHg
b. 1750 mmHg
c.
d.
2.3 mmHg
0.0030 mmHg
32. 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
33. 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
34. 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
35. 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.
36. 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
37. 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.
38. 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
39. 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.
40. 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
86
Unit 8 – Gas Laws
41. 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
42. 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
43. 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.
44. 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
b. 0.0137 oC
d. 73.3 oC
correct
45. 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.
46. 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. The question cannot be answered without knowing the temperature of the system.
87
Unit 9- Solids, Liquids and Phase Changes
SOLIDS AND LIQUIDS
States of Matter
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
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
88
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.
Vocabulary:
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
Phase Changes
Heating Curve
89
Unit 9- Solids, Liquids and 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 _______________________.
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.
Endothermic:
Exothermic:
90
Unit 9- Solids, Liquids and Phase Changes

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?

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?
91
Unit 9- Solids, Liquids and Phase Changes
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?
H
A
F
200º C
G
_____ to _______
B
60º C
3.
Define viscosity-
4.
The temperature at which the vapor
pressure of a liquid equals the
external or atmospheric pressure is
known as
the________________________
D
E
C
Energy
Directions: Using the phase diagram below,
answer questions 5-7:
5.
What does letter T represent?
Pressure
(atm)
C
_______________
6.
0.75
At 0.75 atm, what is the melting point
and boiling point of the substance?
7.
T
0.25
________and _________
At 0.25 atm, what is the freezing point?
________
125
175
Temperature (oC)
250
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.
92
melting
boiling
Unit 9- Solids, Liquids and Phase Changes
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
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.
93
metallic
ionic
Unit 9- Solids, Liquids and Phase Changes
22. The difference between crystalline and amorphous solids is determined by
a. temperature changes
c. strength of molecular forces
b. pressure when the substances are
d. the particle arrangement
formed
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.
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 of the following interactions are the strongest?
a. hydrogen bonding force
b. ion-ion interactions
c.
d.
dipole- dipole force
London-dispersion force
29. 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.
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
94
Unit 9- Solids, Liquids and Phase Changes
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
d.
e.
Condensation point
Critical point
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. Ice floats in water because:
a. Water is denser than ice
b. Ice is colder than water
c.
d.
Water has a substantial surface
tension
Ice is denser than water
38.
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
39.
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
95
d.
The one in the middle and the one
on the right
Unit 9- Solids, Liquids and Phase Changes
40.
a.
b.
Which phase of matter is depicted here?
Liquid
Gas
c.
d.
Plasma
Solid
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.
a.
b.
Which phase of matter is depicted here?
Solid
Gas
96
c.
d.
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.
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 _____________________________.
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.
Solubility:_____________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
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
97
Unit 10- Solutions and Solubility
maximum amount of solute per amount of the solution under the given conditions is called a(n)
_____________________________ 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.

How many grams of potassium chlorate (KClO3) will dissolve in 100 g of water at 30ºC?

How many grams of potassium nitrate (KNO3) will dissolve in 100 g of water at 50ºC?

At what temperature will 90 grams of Pb(NO3)2 dissolve in 100 g of water?

At what temperature will 50 grams of KCl dissolve in 100 g of water?

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.
98
Unit 10- Solutions and Solubility
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.
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
molarity?

How many moles are in 1500 mL of a 3.2 M solution of nitric acid (HNO 3)?

80.6 g of KCl are dissolved in a small amount of water then diluted to 500. mL. What is the molarity?

125 g of NaC2H3O2 are dissolved in a small amount of water then diluted to 750. mL. What is the molarity?

How many grams of CaCl2 are needed to make 625 mL of a 2.00 M 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 the Figure on page 89 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 NaNO3 solution at 10ºC?

What is the molarity of a KClO3 solution at 70ºC?
Dilution
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
99
Unit 10- Solutions and Solubility
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
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 _______________________________.
100
Unit 10- Solutions and Solubility
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?
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 on page 88, answer the questions three through five.
3.
Which is most soluble at 20ºC? _______
4.
How many grams of KClO3 can be dissolved in 100g H2O at 90ºC? _____
5.
At 40ºC, how much KCl can be dissolved in 300 g. H2O? _________
6.
In Koolaid, describe what is the solute and what is the solvent.
Determine whether, according to the solubility rules, the mixing of these
substances will make a solution or a precipitate.
7. Water and Mg(OH)2
8. Water and Na2CO3
Determine how the following conditions can affect the rate of dissolving KCl
in water.
9. Decrease the temperature of the water
10. Agitate the mixture
solution
precipitate
Faster
Slower
12. Which of these compounds are soluble in water?
a. CaBr2
b. PbCl2
c.
d.
SrS
CaCO3
13. Iron (III) sulfide is soluble in water.
a. True
b.
False
101
Unit 10- Solutions and Solubility
14. LiBr is
a. Soluble
b. Insoluble
c.
d.
can't tell the solubility
a covalent compound
15. NH4OH is insoluble.
a. True
b.
False
16. Which of these compounds is soluble?
a. Pb(OH)4
b. NaHCO3
c.
d.
BaCrO4
Mg3(PO4)2
17. Powdered NaCl will dissolve slower then NaCl crystals because there is less surface area for the reaction to
take place.
a. True
b. False
18. 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
19. 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
20. 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?
21. 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
22. 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
23. 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
24. 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
25. 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
26. 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
102
Unit 10- Solutions and Solubility
27. 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
28. 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
29. 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
30. 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
31. In a solution of sugar and water, the solvent is the:
a. sugar
b.
112 g
water
32. 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
33. 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
34. 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
35. 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
36. Calculate the molarity of a solution made when 80 grams of NaOH is dissolved in 2 L of solution
a. 40 M NaOH
b. 82M NaOH
c. 1 M NaOH
d. 160 M NaOH
103
Unit 11- Acids and Bases
Acids and Bases
Properties of Acids and Bases
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
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) Hydrocarbonic acid _______________
b) Hydroselenic acid _______________
104
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
___________________.
Monoprotic acids have only ____________ ionizable hydrogen. Some acids have more than one ionizable hydrogen
and are called ______________________________ acids.
105
Unit 11- Acids and Bases
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 _____________________________
base _________________________
b) NH3 + H2SO4 ↔ NH4+ + HSO4Acid _____________________________
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.
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?

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 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?

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
106
Unit 11- Acids and Bases
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, HClO4, 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
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
_____________________________________________________________________________________________
____________________________________________________________________________________________
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

Find the pH of the following solutions.
a) The hydronium ion concentration equals: 10–2 M. pH = _________________
b) The hydronium ion concentration equals: 1 x 10 –6 M. pH = _________________
c) The hydroxide ion concentration equals: 10 –5 M. pH = _________________
d) 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?
107
Unit 11- Acids and Bases
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
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 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 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 hydroxide ion concentration if the hydrogen ion concentration equals: 1 x 10–9 M.
Indicators
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 
108
Unit 11- Acids and Bases
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 milliliters 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.
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.
109
Unit 11- Acids and Bases
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 HClO 3 in water
5.
Explain how strong acid solutions conduct an electric current.
6.
Write balanced molecular equations for the reactions of acids and bases.
7.
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
110
Unit 11- Acids and Bases
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?
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. Compare and Contrast Arrhenius acids/bases and Brønsted-Lowry acids/bases.
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. Find [OH ] for 1.0 × 10-12 M HClO4.
17. What is the pH of 1.0 × 10-4 M HCl?
18. What is the pH of 1.5 × 10-3 M NaOH?
19. A solution of HNO3 has a pH of 4.0. What is the molarity of HNO 3?
20. What is the molarity of KOH in a solution that has a pH of 10.0?
111
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.
pOH
[OH⁻]
1 x 10⁻6
2
17.
4
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
19. In the reaction H2SO4(aq) --> 2H+(aq) + SO4-2(aq) H2SO4 is a(n)
a. Arrhenius acid
b. Arrhenius base
112
c.
Salt
Acidic or basic
Unit 11- Acids and Bases
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
21. 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
22. Which of the following can conduct an electric current?
a. Mg(OH)2 (s)
b. H2O (s)
c.
d.
NaOH (aq)
NH4Cl (s)
23. In the process of neutralization a salt and a base react to yield water and an acid.
a. True
b. False
24. 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
c. a Arrhenius acid and an electrolyte
b. a salt and a nonelectrolyte
d. a Arrhenius base and an electrolyte
25. 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
26. What type of chemical reaction is neutralization?
a. single replacement
b. double replacement
c.
d.
synthesis
decomposition
27. A titration reaction is a complete neutralization reaction where the moles of H+ equal the moles of OH-.
a. True
b. False
28. Which of the following reactants will represent a neutralization?
a. BaCl2 + CaSO4
b. HCl + F
c.
d.
Ca(OH)2 + H2SO4
NaCl + H2O
29. Titration is a process in which a volume of solution of________
a. unknown concentration is used to determine the concentration of another solution.
b. known concentration is used to determine the volume of another solution.
c. known concentration is used to determine the concentration of another solution.
d. known concentration is used to determine the curve of another solution.
30. 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
31. 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 OH32. What is the molarity of NaOH if 5 milliliters of 4M HCl (aq) neutralizes exactly 10 mL of NaOH (aq)?
a. .5M
b. 1M
c. 1.5M
d. 2M
113
Unit 11- Acids and Bases
33. 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
34. 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
35. One acid-base theory states that an acid is an H +
a. Acceptor
b. Eliminator
c.
d.
Dissolver
Donor
36. 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
37. The acidity or alkalinity of a solution can be measured by its pH value.
a. True
b. False
38. 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.
39. Which of these 1 M solutions will have the highest pH?
a. H3PO4
b. HCl
c.
d.
NaCl
NaOH
40. Which pH indicates an acidic solution?
a. 1
b. 7
c.
9
d.
12
41. Which of these pH numbers indicates the lowest level of acidity?
a. 1
b. 3
c.
8
d.
12
42. Which formula represents a salt?
a. KOH
b. KCl
c.
d.
CH3OH
CH3COOH
43. Which substance can be classified as an Arrhenius acid?
a. HCl
b. NaCl
c.
d.
LiOH
KOH
44. An acidic solution could have a pH of
a. 7
b.
c.
3
d.
14
c.
5
10
45. What is the pH of a 0.00001 molar HCl solution?
a. 1
b. 9
46. What is the pH of a solution with a hydronium ion concentration of 0.01 moles per liter?
a. 1
b. 2
c. 10
47. Given the equation: H+ + OH- ↔ H2O
a. esterification
b. decomposition
Which type of reaction does the equation represent?
c. hydrolysis
d. neutralization
114
d.
d.
4
14
Unit 11- Acids and Bases
48. 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
49. 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. adding 50 mL of 12 M HCl to 50 mL of water while stirring the mixture steadily.
b. adding 50 mL of 12 M HCl to 50 mL of water and then stirring the mixture steadily.
c. adding 50 mL of water to 50 mL of 12 M HCl while stirring the mixture steadily.
d. adding 50 mL of water to 50 mL of 12 M HCl and then stirring the mixture steadily.
50. 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
115
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.

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).
116
Unit 12- Kinetics and Thermochemistry
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
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. ___________________________________________________________________________________.
2. ___________________________________________________________________________________.
3.____________________________________________________________________________________
____________________________________________________________________________________.
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) ___________________________________. As you know, some substances react more readily than others.
The more reactive a substance is, the _____________________________ the reaction rate.
2) ____________________________________. The rate of a reaction is _________________ when the
concentrations of reacting particles are increased. Increasing the number of reactant particles increases
probability of collisions. The rate of gaseous reactions can be ___________ by pumping more gas into the
reaction container.
3) ______________________ the surface area of reactants provides more opportunity for collisions with other
reactants, thereby ______________________ the reaction rate.
117
Unit 12- Kinetics and Thermochemistry
4) Generally, increasing the ____________________________ at which a reaction occurs
_____________________ the reaction rate. Raising the temperature ___________________ both the
collision frequency and the collision energy.
5) Adding a ___________________ affects the rate of a chemical reaction. A catalyst is a substance that
increases the rate of a chemical reaction without itself being _____________________in the reaction.
6) ________________________ gases increases the rate of a chemical 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.
118
Unit 12- Kinetics and Thermochemistry

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.

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?
119
Unit 12- Kinetics and Thermochemistry
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)
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?
120
Unit 12- Kinetics and Thermochemistry
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.
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.
If a reaction is exothermic, are the products or the reactants more stable?
4. List 4 factors that can speed up a chemical reaction.
5. 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)
121
Unit 12- Kinetics and Thermochemistry
6. 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) a catalyst is introduced
(b) the temperature of the system is lowered
7. Convert from one unit to the other:
(a) 1.69 Joules to calories
(d) 20.0 calories to Joules
(b) 820.1 J to kilocalories
(e) 252 cal to J
(c) 423 calories to kilocalories
(f) 2.45 kilocalories to calories
8.
When 15.0 g of steam drops in temperature from 275.0 °C to 250.0 °C, how much heat energy is released?
9.
How much heat (in kJ) is given out when 85.0 g of lead cools from 200.0 °C to 10.0 °C?
10. 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.
11. 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.
12. Determine the specific heat of a 150.0 gram object that requires 62.0 cal of energy to raise its temperature
12.0 °C.
13. Determine the energy required to raise the temperature of 46.2 grams of aluminum from 35.8 °C to 78.1 °C.
14. Determine the energy required to:
(a) melt 5.62 moles of ice at 0 °C.
(b) boil 43.89 grams of water at 100.0 °C.
(c) Convert 16.2 grams of ice to liquid water.
(d) Convert 5.8 grams of water to steam
(e) Convert 98.2 grams of water to ice.
(f) Convert 52.6 grams of steam to water
(g) Convert 125.0 grams of ice at 0.0 °C to steam at 100.0 °C.
(h) Convert 25.9 grams of steam at 100.0 °C.to ice at 0.0 °C.
15. 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
122
Unit 12- Kinetics and Thermochemistry
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
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.
7.
Which of the following species has the highest entropy at 25°C? Explain your answer.
a.
CH3OH(l)
d.
H2O(l)
b.
CO(g)
e.
Ni(s)
c.
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
13.
How does a catalyst speed a reaction?
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
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
E represent?
123
Unit 12- Kinetics and Thermochemistry
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
c. heat of reaction
b. 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.
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. heat of reaction
c.
d.
potential energy of the reactants
activation energy
Catalysts increase the rate of reaction by being consumed.
a. True
b.
False
25.
26.
The following reaction coordinate diagram represents...
a.
b.
c.
d.
an endothermic reaction
an exothermic reaction
a reaction that is neither endothermic nor
exothermic
a reaction in which a catalyst is used
124
Unit 12- Kinetics and Thermochemistry
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.
The collision theory states that a reaction is most likely to occur if reactant particles collide with the
proper
a. energy and concentration
c. concentration and orientation
b. energy and orientation
d. pressure and orientation
29.
30.
What will happen to the rate of reaction when temperature increases?
a. Increase
c. remains the same
b. Decrease
d. increase then decrease
31.
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
32.
Entropy is a measure of the randomness or disorder of a system.
a. True
b.
33.
34.
False
What does a catalyst decrease when introduced to a reaction?
a. the rate of reaction
b. the energy released during the
reaction
c.
d.
the activation energy
the kinetic energy
Which process is accompanied by a increase in entropy?
a. melting of ice
b. freezing of water
c.
condensing of water vapor
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.
What phase change represents a decrease in entropy?
a. solid to liquid
b. liquid to gas
37.
38.
c.
d.
gas to liquid
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.
125
Unit 12- Kinetics and Thermochemistry
39.
In a chemical reaction, if the products have more entropy than the reactants, the change in entropy is
negative.
a. True
b. False
40.
What is the change in entropy in the following reaction C + O 2 -> CO2
a. Increases
c. remains the same
b. Decreases
d. not enough information
41.
What is the change in entropy in the following reaction 4Al(s) + 3O2(g) -> 2Al2O3(s)
a. Increase
c. remains the same
b. Decrease
d. not enough information
42.
What is the change in entropy in the following reaction N2(g) + O2(g) -> 2NO(g)
a. Increase
c. remains the same
b. Decrease
d. not enough information
43.
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
44.
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
45.
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
46.
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
47.
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
48.
Which of the following is not an endothermic process?
a. combustion
b. melting
c. crystallization
d. vaporization
e. sublimation
126
127
128
129
130
131
132
133
134