Louisiana Interactive Reader
Louisiana Interactive Reader
Louisiana Interactive Reader
ISBN 978-0-547-46398-8
1 2 3 4 5 6 7 8 9 10 0956 19 18 17 16 15 14 13 12 11 10
Modern Chemistry
Interactive Reader
Chapter 1 Matter and Change 2
1 Chemistry Is a Physical Science 3
2 Matter and Its Properties 6
3 Elements 16
Math Tutor Significant Figures 22
Chapter Review 23
CONTENTS iii
Chapter 4 Arrangement of Electrons in Atoms 90
1 The Development of a New Atomic Model 91
2 The Quantum Model of the Atom 99
3 Electron Configurations 107
Math Tutor Weighted Averages and Atomic Mass 120
Chapter Review 121
iv CONTENTS
Chapter 8 Chemical Equations and Reactions 246
1 Describing Chemical Reactions 247
2 Types of Chemical Reactions 261
3 Activity Series of the Elements 274
Math Tutor Balancing Chemical Equations 278
Chapter Review 279
CONTENTS v
Chapter 12 Solutions 378
1 Types of Mixtures 379
2 The Solution Process 385
3 Concentration of Solutions 396
Math Tutor Calculating Solution Concentration 406
Chapter Review 407
vi CONTENTS
Chapter 16 Reaction Energy 502
1 Thermochemistry 503
2 Driving Force of Reactions 519
Math Tutor Hess’s Law 526
Chapter Review 527
Glossary 558
Credits 570
CONTENTS vii
CHAPTER 1
2 CHAPTER 1
LOUISIANA STANDARDS
SECTION 1.1 LA.SI.12 Cite evidence that scientific
investigations are conducted for many
Physical Science
Physical science is the study of nonliving objects and VOCABULARY
materials. Physics, astronomy, and geology are examples of
chemistry chemical
physical sciences. Chemistry is another physical science.
Chemistry is the study of the structure and properties of
matter, the processes that matter undergoes, and the energy
changes that accompany these processes. A chemist studies
matter and tries to answer such questions as “What makes up
this material?”, “How does this material change when heated
or cooled?”, “What happens when I mix this material with
another material?”, and “What rules determine how materials
change in different situations?”
READING CHECK
M AT T E R A n d C H A n g E 3
Basic Research
The goal of basic research is to increase knowledge. In CONNECT
chemistry, basic research includes the study of the properties
Basic research can lead to chance
of a chemical. It also includes the study of what happens when discoveries. For example, Ray
two chemicals are mixed. Plunkett discovered the properties
of Teflon®
by accident. After
Sometimes, scientists do basic r esearch simply to satisfy removing the gas from a container,
their curiosity about a chemical and its qualities. Sometimes the container was still heavier than
basic research results in a new product or technology, though expected. Inside, he found a white
solid with nonstick properties.
such results are not the scientists’ goal.
Sometimes the uses of a material or
chemical are not realized at first. It
Applied Research
was many years later before Teflon®
The goal of applied research is to solve problems. In chemistry, was used in nonstick cookware.
applied research includes finding materials with certain
properties. For example, a chemist might want to develop a
new cooling fluid for a refrigerator that is not dangerous to
the environment. Sometimes applied research in chemistry has
the goal of learning how to prevent a certain reaction from
occurring, or how to control the reaction, or how to speed up
the reaction.
Technological Development
The goal of technological development is to come up with
new products and processes that improve the quality of life.
Sometimes new technologies are the result of determined
efforts to create a certain product. New technologies can also
result from knowledge gained in basic and applied research, or
can build on other technologies. For example, lasers were
developed from basic research on crystals and light. Scientists
looking for new ways to transmit information found that laser
pulses could be sent through plastic fibers. This led to the
The nonstick coating on this frying
technology of fiber optic cables that carry television, pan is new technology that makes use
telephone, and computer signals. of the unique properties of Teflon®
.
Critical Thinking
2. Connect Why is basic research important to
technological development?
4 CHAPTER 1
SECTION 1.1 REVIEW
VOCABULARY
1. Define chemistry.
REVIEW
3. Name six branches in the study of chemistry.
Critical Thinking
5. INFERRING RELATIONSHIPS Scientific and technological advances are
c onstantly changing how people live and work. Discuss a change that you
have observed in your lifetime that has made life easier or more enjoyable
for you.
M atter a n d C ha n g e 5
LOUISIANA STANDARDS
SECTION 1.2 LA.PS.14 Identify unknowns as elements,
compounds, or mixtures based on physical
Matter and Its Properties properties (e.g., density, melting point, boiling
point, solubility). (PS-H-C1)
LA.PS.31 Describe chemical changes and
reactions using diagrams and descriptions of
the reactants, products, and energy changes.
(PS-H-D1)
6 ChaPter 1
Properties and Changes in Matter
Every substance has characteristic properties. Chemists use
properties to distinguish between different substances.
Properties can help reveal the identity of an unknown
substance. Comparisons of several properties can be used
together to establish the identity of the unknown.
Chemists also use properties to separate different
substances that are mixed together. For example, a mixture
of iron and aluminum shavings can be separated using the
magnetic property of iron. The iron shavings are attracted to
a magnet and the aluminum shavings are not.
Extensive properties depend on the amount of matter that
is present. The volume of an object is an extensive property
because it changes when material is added to, or taken away
from, an object. Extensive properties include the volume,
mass, and amount of energy in an object.
In contrast, intensive properties do not depend on the
amount of matter present. Such properties include melting
point, boiling point, and density. Intensive properties are the
same for two samples of a substance even if the samples are
different in size.
READING CHECK
M at t e r a n d C h a n g e 7
Changes of State The three common states of matter are
solid, liquid, and gas. One type of physical change is a change
from one of these states to another, such as when ice melts
into liquid water. The change of matter from one state to
another is called a change of state.
Another example of a change of state is freezing, the
opposite of melting, in which a substance changes from a
liquid to a solid state. When matter changes state, the
movement of and distance between the particles in the matter
change, but the matter itself stays the same.
READING CHECK
8 ChaPter 1
A fourth state of matter, plasma, is found in
fluorescent bulbs and stars. A plasma has a high
temperature and its matter is made up of charged
particles. Like a gas, a plasma takes the shape of its
container. Unlike a gas, its particles can be
influenced by electrical charges. A lightning bolt is
made of air particles that have been converted into
plasma.
READING CHECK
M at t e r a n d C h a n g e 9
Iron
Properties: hard gray
metal, solid at room
temperature, easily
shaped
Iron(III) chloride
Properties: dark colored
powder, solid at room
temperature, poisonous
and corrosive
Chlorine
Properties: yellow-
green gas at room
temperature, smells Iron wool dipped in a flask of
like bleach chlorine gas ignites and forms
iron(III) chloride.
10 C HA P TER 1
Energy and Changes in Matter
Every physical change and chemical change requires energy.
This energy can take several forms, such as heat or light. The
amount of energy helps to determine what type of change takes
place. For example, heat can cause the change of state in which
water boils and becomes water vapor. But heat can also cause
water vapor to break down into oxygen gas and hydrogen gas.
This decomposition of water vapor is a chemical change.
Scientists keep track of the energy present before and after
a physical or chemical change. In every case, they have found
! Remember
A scientific law is a statement
that the total amount of energy present before the change is that summarizes how the natural
also present after the change. Energy can be absorbed or world works.
released by one of the substances involved in the change.
However, even if some of the energy has changed form, the
total energy in the system remains the same. This concept is
called the law of conservation of energy.
Classification of Matter
Any sample of matter can be classified as either a pure
substance or as a mixture. The composition of a pure
substance is the same throughout, with no variation from
sample to sample. A pure substance can be either an element
or a compound.
Mixtures, in contrast, contain more than one substance.
The properties within a mixture can vary from sample to
sample. Sometimes two samples from the same mixture will be
different depending on the composition of the mixture at each
location. For example, samples from different places on a
block of gold will have the same composition and properties
because gold is a pure substance. But, samples from different
Two samples from this mixture of salt,
places on a large rock may be completely different because a sand, poppy seeds, and iron would not
rock is usually a mixture of smaller minerals. be identical.
READING CHECK
M at t e r a n d C h a n g e 11
Mixtures
Nearly every object around you, including most things you eat
TIP Knowing the meanings of
the prefixes homo- and
and drink, and even the air you breathe, is a mixture. A hetero- can help you remember the
mixture is a blend of two or more kinds of matter, each of meanings of homogeneous and
heterogeneous. Homo- means
which keeps its own identity and properties. As a result, the “same.” Hetero- means “other” or
properties of a mixture are a combination of the properties of “different.”
its components.
A mixture is called homogeneous if it is uniform in
composition. In other words, a homogeneous mixture looks
the same throughout the entire mixture. Often it is hard to
tell that a homogeneous mixture contains more than one
substance. An example of a homogeneous mixture is salt
water. Homogeneous mixtures are also called solutions. LOOKING CLOSER
A mixture is called heterogeneous if it is not uniform 8. Use the chart below to classify
throughout. For example, in a mixture of clay and water, chocolate milk as a homogeneous
mixture, a heterogeneous mixture, a
heavier clay particles concentrate near the bottom of the compound, or an element.
container. A sample from the bottom of the container will be
different from a sample from the top of the container.
Matter
12 C HA P TER 1
Separating Mixtures The parts of a mixture can usually be
separated. Using differences in properties of the substances
making up the mixture enables their separation. For example,
passing some mixtures through a filter or sieve will separate
the components. Filters separate parts of a mixture using the
property of particle size.
Other methods of separating mixtures include using
centrifuges and chromatography. When a centrifuge spins
really fast, the solid particles tend to accumulate at the bottom
of the test tube. A centrifuge separates substances using
density. Paper chromatography can separate mixtures of dyes
or pigments. Different substances are absorbed and flow up
through the fibers of the paper at different rates.
Pure Substances
Though mixtures can be either homogeneous or
heterogeneous, any sample of a pure substance is
Barium chromate can be separated
homogeneous. A pure substance has a fixed composition from the solution in the beaker using
and differs from a mixture in the following ways: filtration.
READING CHECK
A centrifuge spins rapidly to separate
9. Is every mixture also a solution? Is every solution also a components of a solution.
mixture? Explain.
M at t e r a n d C h a n g e 13
Breaking Down Compounds Pure substances are either
compounds or elements. A compound can only be broken
down into two or more simpler compounds or elements by a
chemical change. In contrast, a mixture can be separated by
the use of its physical properties.
Critical Thinking
11. Identify Look at the two photographs of a
bottle of zinc nitrate to the right. What grade
is the chemical?
14 C HA P TER 1
SECTION 1.2 REVIEW
VOCABULARY
1. Classify each of the following as either a physical change or a
chemical change.
c. burning a log:
REVIEW
2. What is the main difference between physical properties and
chemical properties?
Critical Thinking
6. ANALYZING INFORMATION Compare the composition of sucrose purified from
sugar cane with the composition of sucrose purified from sugar beets.
Explain your answer.
M at t e r a n d C h a n g e 15
LOUISIANA STANDARDS
SECTION 1.3 LA.PS.17 Use the periodic table to compare
electronegativities and ionization energies of
As you have read, elements are pure substances that cannot VOCABULARY
be broken down by chemical changes. Each element has
group metal
characteristic properties. Chemists have organized the family nonmetal
elements into groups based on these properties. This period metalloid
organization of the elements is called the periodic table. The
periodic table on the next page uses the chemical symbol
for each element to show its position. The complete periodic
table that includes the names for all of the elements can be
found on pages 132–133.
Sr; barium, Ba; and radium, Ra. All of these elements are
metals, tend to react quickly with other elements, and bond
to other kinds of atoms in similar ways.
16 CHAPTER 1
Group 18
1
H 2
58 59 60 61 62 63 64 65 66 67 68 69 70 71
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
90 91 92 93 94 95 96 97 98 99 100 101 102 103
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
The horizontal rows of elements in the periodic table are The periodic table of the elements.
called periods. Physical and chemical properties change in a The names of the elements can be
somewhat regular pattern as you move from left to right found on the expanded periodic table
on pages 132–133.
across a period. Elements that are close to each other in the
same period tend to be more similar than elements that are
farther apart. For example, the Period-2 elements lithium, Li,
and beryllium, Be, in Groups 1 and 2, respectively, have
somewhat similar properties. However, their properties are
very different from the properties of fluorine, F, the Period-2
element in Group 17.
The two sets of elements placed below the periodic table TIP The elements in the bottom
two rows belong to none of
make up what are called the lanthanide series and the actinide
the 18 groups. The elements in the
series. These metallic elements fit into the table just after first of these rows are called
elements 57 and 89. They are placed below the table to keep lanthanides, and the elements in the
the table from being too wide. second of these rows are called the
actinides.
READING CHECK
2. What is the group and period number for the element with
the chemical symbol O, called oxygen?
M AT T E R A n d C H A n g E 17
Types of Elements
The periodic table is broadly divided into two main sections:
metals and nonmetals. The metals are at the left and in the
center of the table. The nonmetals are toward the right. Some
elements, such as boron, B, and silicon, Si, show characteristics
of both metals and nonmetals.
Metals
A metal is an element that is a good electrical conductor and LOOKING CLOSER
a good heat conductor. At room temperature, most metals are 3. List below all of the words that
solids. Most metals are also malleable, meaning that they can indicate the typical properties
be hammered or rolled into thin sheets. Metals also tend to be of metals.
ductile, making it possible for them to be drawn into a fine
wire. Metals behave this way because they have high tensile
strength, the ability to resist breaking when pulled. Most
metals also have a silvery or grayish white luster, or shine.
Though metals generally share these properties, they
are diverse. Mercury is a liquid at room temperature. The
metals in Group 1 are soft enough that they can be cut with
a knife. Some metals, such as manganese, are brittle. Instead
of being silvery, gold and copper shine yellow and reddish
brown, respectively.
CONNECT
(a) (b) Any metal becomes a better
conductor of electricity as its
temperature decreases. In 1911,
scientists discovered that when
mercury is cooled to about −269°C,
it loses all resistance and becomes
a superconductor. Today, scientists
can make superconducting
materials that only need to be
cooled to −183°C.
(a) Gold has low reactivity and is usually found in a pure form.
(b) Aluminum is malleable and can be made into a thin foil for wrapping.
18 CHAPTER 1
Copper: A Typical Metal
Copper has a characteristic reddish color and a metallic luster.
It is found naturally in minerals such as chalcopyrite and
malachite. Pure copper melts at 1083°C and boils at 2567°C. It
can be readily drawn into fine wire, pressed into thin sheets,
and formed into tubing. Copper conducts electricity with little
loss of energy.
Copper remains unchanged in pure, dry air at room
temperature. When heated, it reacts with oxygen in air. It Copper is used in electrical
also reacts with sulfur and the elements in Group 17 of the wiring because of its high
electrical conductivity.
periodic table. The green coating on a piece of weathered
copper comes from the reaction of copper with oxygen,
carbon dioxide, and sulfur compounds. Copper is an essential
mineral in the human diet.
READING CHECK
Nonmetals
A nonmetal is an element that is a poor conductor of heat and
electricity. The periodic table includes many more metals than
nonmetals. Many nonmetals are gases at room temperature.
These include nitrogen, oxygen, fluorine, and chlorine. One
nonmetal, bromine, is a liquid. The solid nonmetals include
carbon, phosphorus, selenium, sulfur, and iodine. These solids
tend to be brittle rather than malleable and ductile.
Some nonmetallic elements: (a) carbon, (b) sulfur, (c) phosphorus, and (d) iodine
M atter a n d C ha n g e 19
Phosphorus: A Typical Nonmetal
Phosphorus is one of five solid nonmetals. Pure phosphorus
is known in two common forms. Red phosphorus is a dark
red powder that melts at 597°C. White phosphorus is a waxy
solid that melts at 44°C. Because it ignites in air at room
temperature, white phosphorus is stored underwater.
Phosphorus is too reactive to exist in pure form in nature.
It is present in huge quantities in phosphate rock, where
it is combined with oxygen and calcium. All living things
contain phosphorus. White phosphorus is kept underwater
to keep it from catching fire.
Metalloids
A metalloid is an element that has some characteristics of
metals and some characteristics of nonmetals. As you look
from left to right on the periodic table, you can see that the
metalloids are found between the metals and the nonmetals.
All metalloids are solids at room temperature. They tend
to be less malleable than metals but not as brittle as
nonmetals. Some metalloids, such as antimony, have a
somewhat metallic luster.
Metalloids tend to be semiconductors of electricity. That
is, their ability to conduct electricity is intermediate between
that of metals and that of nonmetals. Metalloids are used in
the solid state circuitry found in desktop computers, digital Phosphorus is often found in
watches, televisions, and radios. match heads.
Noble Gases
The elements in Group 18 of the periodic table are the
noble gases. These elements are generally unreactive. Low
reactivity makes the noble gases very different from the other
families of elements. Group 18 elements are gases at room
temperature. Neon, argon, krypton, and xenon are all used
in lighting. Helium is used in party balloons and weather
balloons because it is less dense than air.
Critical Thinking
5. Draw Conclusions To what group of elements do the noble
gases belong: metals, nonmetals, or metalloids? Explain.
20 CHAPTER 1
SECTION 1.3 REVIEW
VOCABULARY
1. Describe the main differences among metals, nonmetals, and metalloids.
REVIEW
2. Use the periodic table on pages 132–133 to write the names for the
following elements:
a. O
b. S
c. Cu
d. Ag
3. Use the periodic table to write the symbols for the following elements:
a. iron
b. nitrogen
c. calcium
d. mercury
4. Which elements are most likely to undergo the same kinds of reactions,
those in a group or those in a period?
Critical Thinking
5. INFERRING CONCLUSIONS If you find an element in nature in its pure
elemental state, what can you infer about the element’s chemical reactivity?
How can you tell whether that element is a metal or a nonmetal?
M atter a n d C ha n g e 21
Math Tutor SIGNFICANT FIGURES
mass of paperclip
2 3
Problem-Solving TIPS
• Every nonzero digit is significant. Zeros between nonzero digits are significant.
• Zeros appearing in front of the first nonzero digit are not significant.
• If there is no decimal point, zeros that follow the last nonzero digit are not significant.
• If there is a decimal point, zeros that follow the last nonzero digit are significant.
• When measurements are added or subtracted, the result must be rounded to the same
number of decimal places that the measurement with the fewest decimal places has.
• When measurements are multiplied or divided, the result must be rounded to the
same number of significant figures that the measurement with the smallest number
of significant figures has.
SAMPLE
22 CHAPTER 1
CHAPTER 1 REVIEW
4. a. What is mass?
b. What is volume?
6. a. Define property.
M AT T E R A N D C H A N G E 23
8. Define chemical property. List two examples of chemical properties.
reactants: products:
24 CHAPTER 1
15. a. What is the significance of the vertical columns of the periodic table?
16. Compare the physical properties of metals, nonmetals, metalloids, and noble
gases, and describe where in the periodic table each of these kinds of
elements is located.
17. Suppose element X is a poor conductor of electricity and breaks when hit
with a hammer. Element Z is a good conductor of electricity and heat. In
what area of the periodic table does each element most likely belong?
18. Use the periodic table to write the names of the elements that have the
following symbols, and identify each as a metal, nonmetal, metalloid, or
noble gas.
a. K d. Na
b. Ag e. Hg
c. Si f. He
20. Use the periodic table to identify the group numbers and period numbers of
the following elements:
22. Perform the following calculations and apply the rules for significant figures.
M AT T E R A N D C H A N G E 25
CHAPTER 2
Measurements and
Calculations
LOUISIANA STANDARDS
26 CHAPTER 2
LOUISIANA STANDARDS
SECTION 2.1 LA.SI.2 Describe how investigations can be
observation, description, literature survey,
phosphorus
system is a specific portion of matter in a given region fertilizer
15
of space that has been selected for study. For
10%
example, when you observe a chemical reaction in a 10 phosphorus
fertilizer
test tube, the test tube and its contents form a system. 5 no
In a closed system, in which nothing can enter or fertilizer
M E A S U R E M E N T S A N D C A L C U L AT I O N S 27
Formulating Hypotheses
Scientists attempt to find relationships and patterns in the
observations and data they collect by studying systems.
Besides graphs, scientists use a wide array of tools and
techniques to help them analyze their data. For example, they
can organize data into tables or perform a statistical analysis
on data using a computer.
After scientists have collected and analyzed enough
data to find a relationship or pattern, they want to determine
the reason the relationship exists. Scientists construct a
These students have designed an
hypothesis, or testable statement. The hypothesis serves as a experiment to determine how to get
basis for making predictions and for carrying out further the largest volume of popcorn from
experiments. A hypothesis is often written as an “if-then” a fixed number of kernels.
Testing Hypotheses
To test a hypothesis, a scientist must make measurements to
TIP The plural of hypothesis
determine if the prediction in the hypothesis was correct. If is hypotheses.
testing reveals that the prediction was not correct, the scientist
must reject or modify the hypothesis.
A scientist must also design an experiment that eliminates
all the other factors besides the factor mentioned in the “if”
part of the hypothesis. This allows the scientist to state that
the result was a direct consequence of the factor being tested.
During testing, the conditions that remain constant, and
therefore are not factors in the result, are called controls.
Any conditions that change during the experiment are called
variables. The outcome of the experiment should rely on the
effect of the variables and not of the controls.
READING CHECK
28 CHAPTER 2
Theorizing
After a hypothesis is tested and is believed to be correct, the
next step is often to construct a model. A model is an
explanation of why an event occurs and how data and events
are related. In science, a model in more than just a physical
object. Models may be visual, verbal, or mathematical.
One important model in chemistry is the atomic model of
matter, which states that matter is composed of tiny particles
called atoms.
If a model successfully explains many observable facts, it
may become part of a theory. The atomic model is a part of
the atomic theory, which you will study in Chapter 3. A theory
is a broad generalization that explains a related group of facts
or processes. Theories are considered successful if they can
predict the results of many new experiments.The flowchart
below shows where scientific theories fit in the scheme of the
scientific method. The ability to contribute to the development
of new theories is often the reason people become scientists.
Critical Thinking
3. Infer In the diagram below, communication is listed under
every stage. Why is communication important at every
stage in the scientific method?
FORMULATING THEORIZING
HYPOTHESES
• constructing
• organizing and
models
analyzing data
• predicting
• classifying
• communicating
• inferring
• predicting
• communicating
The scientific method is not a single, fixed process. Each stage represents a
number of different activities, and stages are often repeated several times.
M E A S U R E M E N T S A N D C A L C U L AT I O N S 29
SECTION 2.1 REVIEW
VOCABULARY
1. How do hypotheses and theories differ?
REVIEW
2. What is the scientific method?
Critical Thinking
5. INTERPRETING CONCEPTS Suppose you had to test how well two types of soap
work. Describe an experiment that could perform this test. Be sure to use
the terms control and variable.
30 CHAPTER 2
LOUISIANA STANDARDS
STANDARDS
ST
SECTION 2.2 LA.PS.1
TK Text TK
Convert metric system units
involving length, mass, volume, and time
SI Measurement
Scientists have agreed on a single measurement system called TIP Numbers in this book
appear without commas to
Le Système International d’Unités, or SI. The system has seven
separate groups of digits because
base units, which are shown in the table on page 32. Each unit the comma is often used as a
is defined in terms of a standard of measurement. Some decimal point in other countries.
standards are specific objects used for comparison. Others are For example, in Europe the number
seventy-five thousand might be
amounts that do not vary and can be reproduced easily. written as 75.000. In this book,
therefore, the number seventy-five
Critical Thinking thousand will appear as 75 000.
M E A S U R E M E N T S A N D C A L C U L AT I O N S 31
SI Base Units
Quantity Symbol Name Abbreviation Defined standard
Length l meter m the distance light travels in 1/299 792 458 s
Mass m kilogram kg the mass of the international prototype
Time t second s the time it takes cesium-133 to transition
between two levels of its ground state a total
of 9 192 631 770 times
Temperature T kelvin K 1/273.16 of the temperature at which water can
exist as a solid, liquid, and gas at the same time
Amount of n mole mol the number of atoms in 0.012 kg of carbon-12
substance
Electric I ampere A the current that would produce 2 × 10–7
N/m
current of force per unit length between two parallel
conducting plates
Luminous Iv candela cd the luminous intensity of light from a
intensity 540 × 1012
Hz source that has radiant intensity
of 1/683 watt per steradian
Critical Thinking
2. Analyze Fill in the missing information in the table.
SI Prefixes
Prefix Unit abbreviation Exponential factor Meaning Example
giga G 109 1 000 000 000 1 gigameter (Gm) = 1 × 109 m
mega M 106 1 000 000
kilo k 103 1 kilometer (km) = 1000 m
hecto h 100 1 hectometer (hm) = 100 m
deka da 101 10
0
— — 10 1 1 meter (m)
deci d 1/10 1 decimeter (dm) = 0.1 m
centi c 10–2 1/100
milli m 10–
3 1 millimeter (mm) = 0.001 m
micro μ 10–
6 1/1 000 000 1 micrometer (μm) = 1 × 10–6
m
nano n 1/1 000 000 000 1 nanometer (nm) = 1 × 10–9
m
pico p 10–12 1/1 000 000 000 000
femto f 10–
15 1 femtometer (fm) = 1 × 10–15
m
32 CHAPTER 2
SI Base Units
The seven base units in the SI system are listed at the top of
page 32. All of the other SI units can be derived from these
seven units. The prefixes listed in the bottom table on page 32
are often added to the names of the base units to represent
quantities that are much larger or much smaller than the base
units. For example, the prefix kilo-, abbreviated k, represents a
factor of 103 . Therefore, a kilogram is 1000 grams. Similarly,
the prefix centi- represents a factor of 1/100. A centimeter is
1/100 of a meter, or 0.01 meter.
Mass
As you learned in Chapter 1, mass is a measure of the quantity
TIP When a mass is given as a
unit with a prefix, the prefix
of matter. The SI standard unit of mass is the kilogram, which compares the amount to grams, not
is about the weight of a small textbook. The kilogram is the kilograms. For example, a milligram
only base unit with a prefix. is 1/1000 of a gram, not 1/1000 of a
kilogram, the SI base unit.
Mass is often confused with weight because people often
express the weight of an object in kilograms. Mass is
determined by comparing the mass of an object with a set of
standard masses. Weight is a measure of the force with which
gravity pulls on matter. Weight depends on the strength of the
force of gravity, while mass does not. For example, the weight
of an object on the moon is about one-sixth of its weight on
Earth. But the object has the same mass whether it is on the
moon or on Earth. Mass is measured on instruments such as a
balance, and weight is typically measured on a spring scale.
Length
The standard unit for length is the meter. A distance of 1 m is
about the width of an average doorway. Longer distances are
often expressed in kilometers. Kilometers are used for
The base unit of length is the meter,
highway distances in most countries other than the United but the centimeter is often used to
States. A kilometer is about six-tenths of a mile. measure smaller distances.
PRACTICE
M E A S U R E M E N T S A N D C A L C U L AT I O N S 33
Derived SI Units
Many SI units are combinations of the seven base units. For
example, speed is measured in meters per second, which is a
combination of the base units for length and time. Derived
units are formed by multiplying or dividing standard units.
For example, area is a derived unit formed by multiplying
length times width. If both length and width are expressed in A speedometer measures distance
meters, then the area is expressed in meters times meters, or traveled per unit of time. This
square meters, abbreviated m2 . Prefixes can also be added to speedometer shows speed in miles per
hour and kilometers per hour.
derived units. For example, area can be expressed in c m2,
square centimeters, or mm2 , square millimeters.
Some combination units are given their own names. For
example, the unit for force is a combination of the units for
READING CHECK
mass, length, and time, given by kg•m/s2 . The name newton, N,
is given to this combination. A joule is another combined unit 3. Fill in the missing information
in the table below using the
used to measure the quantity of energy in a system or object. information given in the
A joule is a newton times a meter. other columns.
Derived SI Units
Quantity
Quantity Unit Unit abbreviation Derivation
symbol
34 CHAPTER 2
Volume
Volume is the amount of space occupied by an object. The
derived SI unit of volume is cubic meters, m 3 . One cubic meter
is equal to the volume of a cube whose edges are 1 m long.
Such a large unit is inconvenient for expressing the volume of
materials in a chemistry laboratory. Instead, a smaller unit, the
cubic centimeter, cm3 , is often used. There are 100 centimeters
in a meter, so a cubic meter contains 1 000 000 cm3 .
100 cm
1m3 × _______
100 cm
× _______
100 cm
× _______ = 1 000 000 c m3
1m 1m 1m
When chemists measure the volumes of liquids and gases,
they often use a non-SI unit called the liter. The liter, L, is
equivalent to one cubic decimeter. Because there are 10
centimeters in a decimeter, a cube with sides one decimeter
long has a volume of 10 cm times 10 cm times 10 cm, or
1000 cm3. Thus, 1 L is also equivalent to 1000 cm3 .
Another non-SI unit used for smaller volumes is the
milliliter, mL. There are 1000 mL in 1 L. Because there are
also 1000 cm3in 1 L, the two units—milliliter and cubic
centimeter—are interchangeable.
READING CHECK
1L
15 mL
1L 1 cm3 1000 cm3 15 mL
M E A S U R E M E N T S A N D C A L C U L AT I O N S 35
Density
Density is the mass of a substance divided by its volume.
Mathematically, density is written as:
mass
density = _______ m
or D = __
volume V
The quantity m is mass, V is volume, and D is density. The SI
unit for density is derived from the base units for mass and
length. Since volume is length cubed, density is mass divided
by length cubed. The SI unit for density is expressed as
kilograms per cubic meter, or kg/m3 .
Density is an intensive property, which means it is a
characteristic property of a substance. It does not depend on
the size of a sample. If a more massive sample is taken, the
volume would increase by the same proportion. Therefore,
density is a property that can be used to identify substances.
An object or substance will float on a liquid if it has a density Density is the ratio of mass to volume.
less than that of the liquid. Both water and copper shot float on
mercury because mercury is so dense.
The density of water is often used as a reference because its
density is so close to 1 when expressed in grams per cubic
READING CHECK
centimeter. Mercury is a liquid that is much denser than water.
The photograph above shows that water floats on mercury. 5. How many of these objects
would float on water: an ice cube,
Copper is a metal that has a density between those of water a bone dog toy, a sugar cube?
and mercury. The copper in the photograph floats between the
water and the mercury. Cork has a density that is much less
than that of water. If a cork cylinder were added to the
graduated cylinder above, it would float on top of the water.
36 CHAPTER 2
SAMPLE PROBLEM
A sample of aluminum metal has a mass of 8.4 g. The volume
of the sample is 3.1 cm3 . Calculate the density of aluminum.
SOLUTION
PRACTICE
Given: m =
D = __
V
Unknown:
E. What
is the volume of a sample of liquid mercury that has a
mass of 76.2 g, given that the density of mercury is 13.6 g/cm3 ?
M E A S U R E M E N T S A N D C A L C U L AT I O N S 37
Conversion Factors
A conversion factor is a ratio derived from a relationship
between two different units that can be used to convert from
one unit to the other. For example, suppose you want to know
how many quarters there are in 12 dollars. To figure out the
answer, you need to know how quarters and dollars are
related. There are four quarters in a dollar. This fact can be
expressed in many ratios, such as those shown below.
4 quarters 1 dollar
_________ = 1 _________
=1
1 dollar 4 quarters
38 CHAPTER 2
PRACTICE
PRACTICE
M E A S U R E M E N T S A N D C A L C U L AT I O N S 39
SAMPLE PROBLEM
Express a mass of 5.712 grams in milligrams and kilograms.
SOLUTION
1 kg
5.712 g = 5.712 g × ______
= 0.005 712 kg
1000 g
40 CHAPTER 2
PRACTICE
16.45 m = 16.45 m ×
M E A S U R E M E N T S A N D C A L C U L AT I O N S 41
SECTION 2.2 REVIEW
VOCABULARY
1. How does a quantity differ from a unit? Use examples to explain the
difference between the two terms.
REVIEW
2. Label each of the following measurements by the quantity each represents.
For instance, 10.6 kg/m3 represents a density.
Critical Thinking
6. INFERRING CONCLUSIONS A student converts grams to milligrams by
1g
multiplying by the conversion factor ________
. Is the student performing this
1000 mg
calculation correctly? Explain.
42 CHAPTER 2
LOUISIANA STANDARDS
SECTION 2.3 LA.PS.2 Differentiate between accuracy and
precision and evaluate percent error.
READING CHECK
M E A s u R E M E n T s A n d C A l C u l AT i o n s 43
High Accuracy In the set of dartboards shown at the right, the (a) (b)
region covered by the throws is centered on the bull’s-eye.
Both dartboards show that the thrower was accurate.
However, only the left dartboard shows that the thrower was
both accurate and precise.
In scientific experiments, it is important to be both accurate Darts within Darts within
and precise. The second dartboard shown in each of these small area large area
= High precision = Low precision
pairs of dartboards illustrate the problems that can occur if
one of the two qualities is missing. Area centered Area centered
on bull’s-eye around bull’s-eye
= High accuracy = High accuracy
Critical Thinking (on average)
2. Apply Imagine that you throw four darts at the dartboard Both of these dartboards show high
below. Your throws are both imprecise and inaccurate. accuracy, but only dartboard (a) shows
high precision.
Draw four X’s on the dartboard that show a set of darts
that you have thrown imprecisely and inaccurately.
Percentage Error
The accuracy of an experimental value can be compared with
the correct or accepted value by calculating its percentage
error. To calculate the percentage error, subtract the accepted
value from the experimental value, dividing the difference by
the accepted value, and then multiply by 100.
Valueexperimental- Valueaccepted
Percentage error = ___________________________
× 100
Valueaccepted
Critical Thinking
3. Reasoning What is the percentage error if the experimental
value is equal to the accepted value?
44 CHAPTER 2
SAMPLE PROBLEM
A student measures the mass and volume of a substance and
calculates its density as 1.40 g/mL. The correct, or accepted,
value of the density is 1.30 g/mL. What is the percentage
error of the student’s measurement?
SOLUTION
PRACTICE
-
Percentage error = ________________________
× 100 =
B. A
volume is measured experimentally as 4.26 mL.
What is the percentage error, given that the correct
value is 4.15 mL?
M ea s u reme n t s a n d ca l c u l at i o n s 45
Error in Measurement
Some error or uncertainty always exists in any measurement.
For example, the skill of the measurer places limits on
reliability. The measuring instruments limit how precisely a
value can be determined. In addition, the readings from
balances, rulers, and graduated cylinders are controlled by
how fine the markings are on the instruments.
As an example, look at the ruler above. This ruler can be
used to determine the length of an object precisely to the
tenths digit. You can tell that the nail is definitely between
6.3 cm and 6.4 cm long. However, it is hard to tell whether the
value should be read as 6.35 cm or 6.36 cm. The hundredths
place is somewhat uncertain, but a reasonable estimate of the A nail’s length is measured by a
digit can be made. You might include a plus-or-minus value to centimeter ruler.
express the range, such as 6.36 cm ± 0.01 cm.
46 CHAPTER 2
Rules for Determining Significant Zeros
Rule Examples
1. Zeros appearing between nonzero digits a. 40.7 L has three significant figures.
are significant. b. 87 009 km has five significant figures.
2. Zeros appearing in front of all nonzero a. 0.095 897 m has five significant figures.
digits are not significant. b. 0.0009 kg has one significant figure.
3. Zeros at the end of a number and to a. 85.00 g has four significant figures.
the right of a decimal point are significant. b. 9.000 000 000 mm has 10 significant figures.
4. Zeros at the end of a number but to the a. 2000 m may contain from one to four
left of a decimal point may or may not be significant figures, depending on how many
significant. If a zero has not been measured zeros are placeholders. For measurements
or estimated but is just a placeholder, it is not given in this book, assume that 2000 m has
significant. A decimal point placed after zeros one significant figure.
indicates that they are significant. b. 2000. m contains four significant figures due
to the presence of the decimal point.
PRACTICE
M ea s u reme n t s a n d ca l c u l at i o n s 47
Rounding
When you add, subtract, multiply, or divide two measurements,
the reliability of the result should be reflected by its significant
figures. How the result should be rounded is determined
partly by the rules given at the bottom of the page. A
measurement’s rounding is also determined by the operation
(addition, subtraction, multiplication, or division) that is used.
48 CHAPTER 2
Rules for Addition and Subtraction Examples
1. For decimals, the sum or difference should have the same 25.1 g + 2.03 g = 27.1 g
number of significant figures after the decimal point as 3.70 mL – 0.493 mL = 3.21 mL
the measurement with the fewest digits to the right of the 17 cm + 5.7 cm = 23 cm
decimal point.
2. For whole numbers, the final significant digit of the sum 5400 g + 365 g = 5800 g
or difference should be the same as the final significant 2710 mL – 1000 mL = 2000 mL
digit of the least precise measurement.
PRACTICE
M ea s u reme n t s a n d ca l c u l at i o n s 49
Scientific Notation
Sometimes measurements are very small or very large. It is
often awkward to give measurements with a lot of zeros, such
as 23 000 000 000 000 or 0.000 000 000 000 001 79. Instead,
measurements are often given in scientific notation.
In scientific notation, numbers are written in the form
M × 10n, where M is greater than or equal to 1 and less
than 10 and n is a whole number. When numbers are written
in scientific notation, only the significant figures are shown.
For example, to write 65 000 km in scientific notation and
show that the first two digits are significant, you would write
6.5 × 1 04km. If you wanted to show that the first three digits
were significant, you would write 6.50 × 104 km.
To write a number in scientific notation, use the following
two steps. This scientific calculator shows the
most common format for displaying
1. Determine M by moving the decimal point in the original scientific notation. The number shown
number to the left or the right so that only one nonzero is 5.44 × 107 .
digit remains to the left of the decimal point. Delete all
zeros that are not significant.
2. Determine n by counting the number of places that you
moved the decimal point. If you moved it to the left, n is
positive. If you moved it to the right, n is negative.
PRACTICE
50 CHAPTER 2
Mathematical Operations Using Scientific Notation
Numbers in scientific notation can be added, subtracted,
multiplied, and divided just as any number can. The rules for
adding, subtracting, multiplying, and dividing two numbers in
scientific notation are given below.
Addition and Subtraction
1. If the exponents of the two numbers differ, rewrite one
number so that both numbers have the same exponent.
2. To find the new value of M, find the sum or difference of
the M values of the two numbers. Keep the same value of n.
Example: 5.93 × 106kg – 4.2 × 105 kg
= 5.93 × 106kg – 0.42 × 106 kg
= 5.51 × 106kg
Multiplication and Division
1. To find the value of M in a product, multiply the two values
of M. To find the value of M in a quotient, divide the first
value of M by the second value of M.
2. To find the new value of n in a product, add the two values Remember that a calculator does not
of n. To find the value of n in a quotient, subtract the value display the correct number of
significant figures after performing
second value of n from the first value of n. an operation.
Example: (2.6 × 108s)(4.7 × 104 s) = (2.6)(4.7) × 108+4
s
12
= 12 × 10 s
= 1.2 × 1013s
Often, the result of these calculations is not in scientific
notation. If M is less than 1 or greater than or equal to 10, the
decimal point must be moved. If the decimal point is moved
one spot to the left, increase n by 1. If the decimal point is
moved one spot to the right, decrease n by 1.
PRACTICE
M ea s u reme n t s a n d ca l c u l at i o n s 51
Using Sample Problems
Learning to analyze and solve such problems requires practice
and a logical approach. In this section, you will review a
process that can help you analyze problems effectively.
Analyze
The first step is to read the problem carefully at least twice
and to analyze the information in it. Identify and list the data
! Remember
A problem may not always give you
in the problem and identify what you are being asked to find. all the data you need for a solution.
For example, you may need to look
Plan up a value on the periodic table.
The second step is to develop a plan for solving the problem.
The plan should show how the information given is to be used
to find the unknown. It is often helpful to draw a picture that
represents the problem to help you visualize the problem.
Decide which conversion factors, mathematical formulas,
or chemical principles you will need to solve the problem.
Your plan might suggest a single calculation or a series of
them involving different conversion factors. Once you
understand how you need to proceed, you may wish to sketch
out the stages of your solution in a table or a flowchart.
Compute
The third step is using the data and conversion factors to carry
out your plan. Make sure to keep track of the units and to
round the result to the correct number of significant figures.
Evaluate
Examine your answer to determine if it is reasonable. Make
sure that the units in the answer are what you would expect.
Use simpler, rounded numbers and repeat the calculations to
check the order of magnitude of your answer.
READING CHECK
52 CHAPTER 2
SAMPLE PROBLEM
Calculate the volume of a sample of aluminum that has a
mass of 3.057 kg. The density of aluminum is 2.70 g/cm3 .
SOLUTION
D = _ m
_ _ _ _ _ _ _ _ ⇒ m
V = __
V D
3.057 kg 1000 g
V = _________
× ______
3
2.70 g/cm kg
M ea s u reme n t s a n d ca l c u l at i o n s 53
PRACTICE
SOLUTION
0.020 second
t = ____________
× × × × 6 months
minute
54 CHAPTER 2
Direct Proportions
Two quantities are directly proportional to each other if
dividing one by the other gives a constant value. For example,
the quantities mass and volume are proportional. The table
below shows the measured mass and volume of five separate
samples of aluminum. As the masses of the samples increase,
their volumes increase by the same factor.
When two variables, x and y, are directly proportional to
each other, the relationship can be expressed as y ∝ x, which is
! Remember
A variable is a quantity that can
read as “y is proportional to x.” The relationship can be shown take on many values.
in these two forms using a proportionality constant k:
y
__
x = k or y = kx
The first equation shows that there is a constant ratio LOOKING CLOSER
between the values of two quantities that are directly 7. What is the constant of propor-
proportional. Note that the data in the table have ratios tionality for the data in the table
that are nearly equal, indicating a relationship that is and what does it represent?
directly proportional.
The second equation shows that a directly proportional
relationship is a straight line with a slope of k. The graph of
every directly proportional relationship also passes through
the origin, or the data point (0,0). Both the values in the third
column of the table below and the graph demonstrate that the
data in the table are directly proportional.
60
83.5 30.9 2.70
40
96.3 35.8 2.69
20
105.7 39.1 2.70
0
0 10 20 30 40 50 60
Volume (cm3)
M ea s u reme n t s a n d ca l c u l at i o n s 55
Inverse Proportions
Two quantities are inversely proportional to each other if
multiplying them gives a constant value. For example, the
amount of time it takes to travel a certain distance is inversely
proportional to the speed of travel. The greater the speed, the
less time is needed to go a certain distance.
When two variables, x and y, are inversely proportional to READING CHECK
each other, the relationship can be expressed as 8. Two quantities are
1
y ∝ __
x
proportional
350
200 250 50 000
300
250 200 50 000 250
56 CHAPTER 2
SECTION 2.3 REVIEW
VOCABULARY
1. What is the difference between a graph representing data that are directly
proportional and a graph of data that are inversely proportional?
REVIEW
2. The density of copper is 8.94 g/cm3 . Two students each measure the density
of three samples of the substance. Student A’s results are 7.3 g/mL, 9.4 g/mL,
and 8.3 g/mL. Student B’s results are 8.4 g/cm3 , 8.8 g/cm3 , and 8.0 g/cm3 .
Compare the two sets of results in terms of precision and accuracy.
5. A student measures the mass of a beaker filled with corn oil. The mass
reading averages 215.6 g. The mass of the beaker is 110.4 g. What is the
density of the corn oil if its volume is 114 cm3 .
Critical Thinking
6. APPLYING CONCEPTS The mass of a liquid is 11.50 g and its volume is
9.03 mL. How many significant figures should its density value have?
Explain the reason for your answer.
M ea s u reme n t s a n d ca l c u l at i o n s 57
Math Tutor Scientific Notation
Any value expressed in scientific notation has two parts. The first part, the first factor,
consists of a number greater than or equal to 1 but less than 10. The second part consists
of a power of 10. To write the first part, move the decimal to the right or the left so that
there is only one nonzero digit to the left of the decimal point. To write the second part,
count how many places the decimal point was moved. The exponent is positive if the
decimal point moved to the left and negative if the decimal point moved to the right.
exponent
Problem-Solving TIPS
• In addition and subtraction, all values must first be converted to numbers that have
the same exponent of 10. The result is the sum or the difference of the first factors,
multiplied by the same exponent of 10. The result should be rounded to the correct
number of significant figures and expressed in scientific notation.
• In multiplication, the first factors are multiplied and the exponents of 10 are added.
• In division, the first factors of the numbers are divided and the exponent of 10 in the
denominator is subtracted from the exponent of 10 in the numerator.
SAMPLE
= 16 × 107
= 1.6 × 108
58 CHAPTER 2
CHAPTER 2 REVIEW
3. Identify the SI unit that would be most appropriate for measuring each of
the following.
4. What is a derived unit, and what is the SI-derived unit for area?
8. Arrange in the correct order the following four basic steps for finding the
solution to a problem: compute, plan, evaluate, and analyze.
M ea s u reme n t s a n d C a l c u l at i o n s 59
9. What is the volume, in cubic meters, of a rectangular solid that is
0.25 m long, 6.1 m wide, and 4.9 m high?
10. Find the density of a material, given that a 5.03 g sample occupies 3.24 mL.
11. A sample of a substance that has a density of 0.824 g/mL has a mass
of 0.451 g. Calculate the volume of the sample.
13. The density of gold is 19.3 g/cm3 . What is the volume, in cubic centimeters,
of a sample of gold that has a mass of 0.715 kg? If this sample of gold is a
cube, what is the length of each edge in centimeters?
14. A student measures the mass of a sample as 9.67 g. Calculate the percentage
error, given that the correct mass is 9.82 g.
60 CHAPTER 2
15. A handbook gives the density of calcium as 1.54 g/cm3 . Based on lab
measurements, what is the percentage error of a density calculation
of 1.25 g/cm3 ?
16. How many significant figures are in each of the following measurements?
17. Perform the following calculations. Express the answers with the correct
number of significant digits.
18. Calculate the product of 0.002 115 m and 0.000 040 5 m. Express the answer
in scientific notation and with the correct number of significant figures.
ANALYZE:
PLAN:
COMPUTE:
EVALUATE:
M ea s u reme n t s a n d C a l c u l at i o n s 61
CHAPTER 3
62 CHAPTER 3
LOUISIANA STANDARDS
SECTION 3.1 LA.PS.7 Write a balanced symbolic equation
from a word equation. (PS-H-A2)
Philosophical Idea to
Scientific Theory
When you crush a lump of sugar, you can see that it is made VOCABULARY
up of many smaller sugar particles. You may grind these
law of conservation of mass
particles into a very fine powder, but each tiny piece is still law of definite proportions
sugar. Now suppose you mix the sugar powder into water. The law of multiple proportions
tiny particles seem to disappear. Yet if you were to taste the
solution, you’d know that the sugar is still there.
Observations like these led people to wonder about the
nature of matter. Can matter be divided into pieces forever or
is it made of miniature pieces that cannot be divided at all?
The particle theory of matter was supported as early as
400 B.C. by certain Greek thinkers, such as Democritus. He
called nature’s basic particle an atom, based on the Greek
word meaning “indivisible.” On the other hand, Aristotle
did not believe in atoms. He thought that all matter was
continuous. He lived in the generation after Democritus, and
his opinion was accepted for nearly 2000 years.
The opinions of Aristotle and Democritus were not
based on experimental evidence. The discussion remained Sugar seems to disappear
a philosophical one until the eighteenth century. Then when it is mixed with water.
scientists began to gather evidence favoring the atomic
theory of matter.
AT o m s : T H E B u i l d i n g B l o C k s o f m AT T E R 63
Evidence for Atomic Theory
Law of conservation Mass is neither created nor destroyed during ordinary chemical
of mass reactions or physical changes.
Law of definite A compound contains the same elements in exactly the same
proportions proportions regardless of sample size or source of the compound.
Law of multiple If two or more compounds are composed of the same two elements,
proportions then the ratio of the masses of the second element combined with a
certain mass of the first element is a ratio of small whole numbers.
64 CHAPTER 3
Dalton’s Atomic Theory
In 1808, an English schoolteacher named John Dalton
proposed an explanation for the three newly discovered laws.
He reasoned that elements were composed of atoms and that
only whole numbers of atoms can combine to form
compounds. His theory is summarized in the table below.
(b)
= +
READING CHECK
At o m s : T he B u i l d i n g B l o c k s o f Matter 65
(a)
(a) CO is always composed of one
+ = C atom and one O atom. (b) C
O2
molecules are always composed of
Carbon, C Oxygen, O Carbon monoxide, CO one C atom and two O atoms. The
same number of CO2molecules
contain twice as many O atoms.
(b)
+ + =
Critical Thinking
4. Infer When scientists discovered that atoms are c omposed
of smaller particles, why didn’t they reject atomic theory?
66 CHAPTER 3
SECTION 3.1 REVIEW
VOCABULARY
1. What chemical laws can be explained by Dalton’s atomic theory?
REVIEW
2. List the five main points of Dalton’s atomic theory.
1.
2.
3.
4.
5.
Critical Thinking
3. ANALYZING INFORMATION Three compounds containing potassium and
oxygen are compared. Analysis shows that for each 1.00 g of O, the
compounds have 1.22 g, 2.44 g, and 4.89 g of K, respectively. Show how
these data support the law of multiple proportions.
At o m s : T he B u i l d i n g B l o c k s o f Matter 67
LOUISIANA STANDARDS
SECTION 3.2 LA.PS.8 Analyze the development of the
modern atomic theory from a historical
the Atom
The atomic theory proposed by John Dalton stated that atoms
are indivisible. This idea was proven incorrect by the end of
the nineteenth century. It became clear that atoms are actually
composed of several basic types of particles. The number and
arrangement of these particles determine the properties of
the atom.
The atom is now defined as the smallest particle of VOCABULARY
an element that retains the chemical properties of the
atom nuclear forces
element. Atoms consist of two regions. Each region contains
different types of particles, called subatomic particles.
• The nucleus is a very small region located at the center of
the atom. The nucleus includes at least one positively
charged particle called a proton and usually one or more READING CHECK
neutral particles called neutrons. 1. Draw a model showing the two
regions of an atom in the space
• Surrounding the nucleus is a much larger region that below. Label the nucleus and the
contains negatively charged particles called electrons. electron region.
68 CHAPTER 3
Cathode Rays and Electrons
Scientists discovered that the surface of a Voltage source
cathode-ray tube glowed when electric current
was passed through the tube. They hypothesized Gas at low pressure Cathode ray
Thomson concluded that all cathode rays are (a) A simple cathode-ray tube. (b) A magnet above
a cathode-ray tube deflects the beam downward,
composed of identical negatively charged particles
showing that the particles in the beam must have a
called electrons. The atoms in the cathode-ray negative charge.
experiments above were releasing electrons. This was
evidence that atoms are divisible and that electrons
are present in different types of atoms.
AT O M S : T H E B U I L D I N G B L O C K S O F M AT T E R 69
Screen to detect
deflected
particles
Thin gold
foil
70 CHAPTER 3
Composition of the Atomic Nucleus
With one exception, all atomic nuclei are made of two kinds of T I P Many words derived from
particles, protons and neutrons. A proton has a positive charge Latin that end with -us
form plurals by changing the
equal in magnitude to the negative charge of an electron. ending to -i. Thus, the plural of
Atoms are electrically neutral because they contain equal nucleus is nuclei and the plural of
numbers of protons and electrons. A neutron has no charge, radius is radii.
and, like an atom, is electrically neutral.
The one atomic nucleus that lacks a neutron is that of the
simplest hydrogen atom. Its nucleus is a single proton with
a single electron moving around it. A proton has a mass of
1.673 × 10–27
kg, which is 1836 times greater than the mass of
an electron. Therefore, a proton has nearly all of the mass in
the simplest hydrogen atom. The mass of a neutron is
1.675 × 10–27
kg, which is slightly larger than the mass of
a proton.
The nuclei of atoms of different elements differ in the
number of protons they possess. Therefore, the number of
protons determines an atom’s identity. For every proton
an atom has in its nucleus, the same number of electrons
surrounds the nucleus. Physicists have discovered other
subatomic particles, but they have little effect on the chemical
properties of matter. The properties of electrons, protons, and
neutrons, are summarized in the table below.
READING CHECK
Proton , 1 1H
p+ +1 1 1.007 276 1.673 × 10−27
AT O M S : T H E B U I L D I N G B L O C K S O F M AT T E R 71
Forces in the Nucleus
Generally, particles that have the same electric charge repel CONNECT
one another. Therefore, you might expect a nucleus with more
than one proton in it to be unstable. However, a force exists In physics, there are four known
fundamental forces that describe
between two protons that overcomes the electric force trying
how matter interacts. These forces
to push them apart. This force only acts when two protons are are the electromagnetic force, the
very close to one another. gravitational force, the strong
nuclear force, and the weak
A similar force acts when two neutrons are very close nuclear force.
together, or when a neutron and a proton are very close
together. Together, these short-range proton-proton,
neutron-neutron, and proton-neutron forces are called nuclear
forces. These forces allow atoms with up to 83 positively-
charged protons in the same nucleus to be stable.
Critical Thinking
5. Calculate Verify the value for the density of an
atomic nucleus given above for a spherical atom
with a mass of 1 amu. Recall that the volume of
4 πr 3.
a sphere is given by V = __
3
72 CHAPTER 3
SECTION 3.2 REVIEW
VOCABULARY
1. Define each of the following:
a. atom
b. neutron
REVIEW
2. Describe one conclusion made by each scientist that led to the development
of the current atomic theory.
a. Thomson
b. Millikan
c. Rutherford
Critical Thinking
5. EVALUATING IDEAS Nuclear forces are said to hold protons and neutrons
together. What is it about the composition of the nucleus that requires the
concept of nuclear forces?
AT O M S : T H E B U I L D I N G B L O C K S O F M AT T E R 73
LOUISIANA STANDARDS
SECTION 3.3 LA.PS.4 Use scientific notation to express
large and small numbers. (PS-H-A1)
Neon gas only makes up 0.002% of the air you breathe. Yet VOCABULARY
there are 5 × 1017 atoms of neon in each breath you take. In
atomic number average atomic mass
most experiments, atoms are too small and numerous to isotope mole
track individually. Instead, chemists make calculations that mass number Avogadro’s number
nuclide molar mass
take into account the properties of large groups of atoms. atomic mass unit
Atomic Number
All atoms are composed of the same basic particles. Yet all
atoms are not the same. Atoms of different elements have
different numbers of protons. Atoms of the same element
all have the same number of protons. The atomic number
(Z) of an element is the number of protons of each atom of
that element.
Turn to the periodic table on pages 132–133. The
periodic table square for lithium is also shown at the right. 3
An element’s atomic number is indicated above its symbol.
Notice that the elements are placed in order in the periodic Li
table using the atomic number. At the top left is hydrogen, Lithium
H, with an atomic number of 1. Next in order is helium, He,
with an atomic number of 2. The next row of the periodic
6.941
table includes the elements with the atomic numbers 3, 4, 5, [He]2 s1
and so on.
This periodic table entry shows that
The atomic numbers give the number of protons in an the atomic number of lithium is 3.
element. So all atoms of hydrogen have one proton, all
atoms of helium have two protons, and so on. READING CHECK
The atomic number also identifies an element. If you 1. How many protons does every
atom of hydrogen have?
want to know which element has atomic number 47, you can
look at the periodic table for the box with a “47” at the top.
Silver, Ag, is the correct element. You then know that all
2. How many protons does every
silver atoms have 47 protons. Since atoms are electrically atom of lithium have?
neutral, you also know that all silver atoms must also have
47 electrons.
74 CHAPTER 3
Isotopes
The simplest atoms are those of hydrogen. All hydrogen
TIP The names for the types of
atoms have only one proton. However, like many naturally hydrogen atoms are
derived from the number of
occurring elements, hydrogen atoms can have different particles in the nucleus. The prefix
numbers of neutrons. proto- means “first,” deutero- means
“second,” and trito- means “third.”
Three types of hydrogen atoms are known. The most The “o” is dropped before the
common type of hydrogen is sometimes called protium. ending -ium in the names of the
It accounts for 99.9885% of the hydrogen atoms found on hydrogen atoms.
Earth. A protium atom has one electron and a nucleus with
one proton. Another form of hydrogen is called deuterium.
A deuterium atom has one electron and a nucleus with two
particles: a neutron and a proton. Finally, a tritium atom is a
hydrogen atom with one electron and a nucleus of one proton
and two neutrons. 1 Proton
Deuterium
Mass Number
2 Neutrons
An isotope is identified by its name, such as protium, or its
atomic number and mass. The mass number of an isotope is
the total number of protons and neutrons that make up its
nucleus. For example, the mass number of protium is one
because there is one particle, a proton, in its nucleus. 1 Proton
Tritium
READING CHECK
The three hydrogen isotopes are shown.
3. Use the definition of mass number to complete the table.
protium 1 0 1
deuterium 1 1
tritium 1 2
AT O M S : T H E B U I L D I N G B L O C K S O F M AT T E R 75
Designating Isotopes
That the isotopes of hydrogen have their own names is
unusual. An isotope is usually identified by specifying its mass
number. There are two methods for specifying isotopes.
• In hyphen notation, the mass number is written with a
hyphen after the name of the element. For example, in
hyphen notation, tritium would be written as hydrogen-3.
• A nuclear symbol is used to show the composition of an
isotope’s nucleus. A number to the upper left of the element
symbol indicates the mass number (protons + neutrons). A
number to the lower left of the element symbol indicates
the atomic number (number of protons). For example, the
nuclear symbol for tritium is 3 1 H
.
Nuclide is a general term for the specific isotope of an
element. For example, you could say that deuterium is a
hydrogen nuclide. You could also say that hydrogen has three
different nuclides. The composition of the three isotopes, or
nuclides, of hydrogen and the two isotopes of helium are given
in the table below.
Critical Thinking
4. Identify A particular isotope of uranium has a nucleus with
92 protons and 143 neutrons. Identify this isotope in two
different ways.
helium-3 32 He 2 1
helium-4 42 He 2
76 CHAPTER 3
SAMPLE PROBLEM
How many protons, electrons, and neutrons are there in an
atom of chlorine-37?
SOLUTION
1 ANALYZE Determine what information is given and unknown.
Given: name of isotope is chlorine-37
Unknown: number of protons, electrons, and neutrons
2 PLAN Write equations for the unknowns in terms of what is given.
number of protons = number of electrons = a tomic number
mass number = number of neutrons + number of protons, so
number of neutrons = mass number – number of protons
3 COMPUTE Substitute the known values and calculate.
Because the name of the isotope is chlorine-37, its mass number
is 37. The element chlorine is element 17 on the periodic table, so
its atomic number is 17.
number of protons = number of electrons = 17
number of neutrons = 37 – 17 = 20
An atom of chlorine-37 has 17 electrons, 17 protons, and
20 neutrons.
4 EVALUATE Determine if the answer makes sense.
he number of protons in a neutral atom equals the number of
T
electrons. The number of protons plus the number of neutrons
equals the mass number because 17 + 20 = 37.
PRACTICE
AT O M S : T H E B U I L D I N G B L O C K S O F M AT T E R 77
Relative Atomic Masses
Masses of atoms expressed in grams are very small. For
example, an atom of oxygen-16 has a mass of 2.656 × 10–23g.
It is usually more convenient to talk about the relative mass of
an atom. The relative atomic mass of an atom is the mass of
the atom as compared to the mass of a defined standard.
Scientists have defined the atomic mass unit as the LOOKING CLOSER
standard mass for use in comparisons. One atomic mass unit, 6. Define the two parts of the term
or amu, is exactly 1/12 the mass of a carbon-12 atom. In atomic mass unit separately in your
other words, one amu is the average mass of a particle in the own words:
nucleus of a carbon-12 atom. The value of amu in grams is atomic mass
1.660 540 × 10–24g.
The mass of a hydrogen-1 atom is slightly more than one
atomic mass unit—1.007 825 amu. An oxygen-16 atom has a
precise mass of 15.994 915 amu. Additional atomic masses for
the isotopes of certain elements are given in the table below.
unit
Isotopes of an element do not differ significantly in their
chemical behavior from the other isotopes of the element. So
the three isotopes of oxygen all have the same chemical
properties despite varying in mass.
The table below shows some isotopes that can be found in
nature. The natural abundance, or relative amount of each
isotope in a sample of an element, is also given in the table.
Artificial isotopes can only be created in the laboratory. They
have a natural abundance of zero.
78 CHAPTER 3
Average Atomic Masses of Elements
Chemists have found that a sample of an element will contain READING CHECK
the same percentage of each isotope no matter where on 7. Define the two parts of the term
Earth the sample is obtained. This percentage is taken into average atomic mass separately in
account when calculating the average atomic mass that is your own words:
reported on the periodic table. The average atomic mass is average
the weighted average of the atomic masses of the isotopes
of an element found in nature. The table on the bottom of
page 78 also includes the average atomic mass for each
element in the table.
Critical Thinking
8. Reasoning Why is the average atomic mass usually
a decimal number and not a whole number like the
mass number?
AT O M S : T H E B U I L D I N G B L O C K S O F M AT T E R 79
Relating Mass to Numbers of Atoms
The atomic mass unit allows scientists to compare the mass of
an atom to the mass of a standard atom. The average atomic
mass gives scientists a value for the average mass of an atom
in a sample. Another quantity that scientists also need to
determine is the number of atoms in a sample.
The Mole
The mole is the SI unit for the amount of a substance. The
abbreviation for a mole is mol. A mole is the amount of a
substance that contains as many particles as there are atoms in
exactly 12 g of carbon-12. The mole is a counting unit, just like
a dozen. If you buy two dozen ears of corn at a farm stand,
you are purchasing 2 times 12, or 24 ears of corn. Similarly, a
chemist might desire 1 mol of carbon or 2.567 mol of calcium.
READING CHECK
10. How many particles does the SI unit for the number of
particles represent?
80 CHAPTER 3
(a) (b) (c)
About one molar mass of (a) carbon (graphite), (b) iron (nails), and (c) copper
(wire) is shown on each balance.
Molar Mass
The number of particles in one mole of a substance is given by
Avogadro’s number. The mass of one mole of a substance is
called the molar mass of that substance. Molar mass is usually
written in units of g/mol. The molar mass of an element in
g/mol is equivalent to the atomic mass of the element as given
on the periodic table in amu. For example, the molar mass of
carbon is 12.01 g/mol, the molar mass of iron is 55.84 g/mol,
and the molar mass of copper is 63.55 g/mol.
Gram/Mole Conversions
Chemists use molar mass as a conversion factor in chemical
calculations. For example, to find the mass of 2 mol of a
substance, you would multiply 2 mol by the molar mass of the
substance (in grams per mole) to obtain a value in grams.
The diagram shows the relationship among mass, moles, and number of atoms.
AT O M S : T H E B U I L D I N G B L O C K S O F M AT T E R 81
SAMPLE PROBLEM
A chemist produced 11.9 g of aluminum, Al. How many
moles of aluminum were produced?
SOLUTION
= moles Al
PRACTICE
= g Fe
82 CHAPTER 3
SAMPLE PROBLEM
How many moles of silver, Ag, are in 3.01 × 1023
atoms
of silver?
SOLUTION
= moles Ag
PRACTICE
2.75 mol Al =
2.75 mol Al ×
= atoms Al
AT O M S : T H E B U I L D I N G B L O C K S O F M AT T E R 83
SAMPLE PROBLEM
What is the mass in grams of 1.20 × 108 atoms of copper, Cu?
SOLUTION
moles
Cu
Cu atoms = Cu atoms × _____________________________
Avogadro’s number of Cu atoms
grams Cu
× _________
moles Cu
= grams Cu
PRACTICE
4.00 g S = 4.00 g S × ×
= S atoms
84 CHAPTER 3
SECTION 3.3 REVIEW
VOCABULARY
1. Define the term molar mass.
REVIEW
2. Complete the table at the right.
Number of Number of Number of
Isotope
protons electrons neutrons
3. Write the nuclear symbol and hyphen
notation for each of the following isotopes. sodium-23
4. To two decimal places, what is the relative atomic mass and the molar mass
of the element potassium, K?
Critical Thinking
7. ANALYZING DATA Beaker A contains 2.06 mol of copper, and Beaker B
contains 222 g of silver.
AT O M S : T H E B U I L D I N G B L O C K S O F M AT T E R 85
Math Tutor Conversion Factors
Most calculations in chemistry require that all measurements of the same quantity
(mass, length, volume, temperature, and so on) be expressed in the same unit. To
change the units of a quantity, you can multiply the quantity by a conversion factor.
With SI units, such conversions are easy because units of the same quantity are
related by multiples of 10, 100, 1000, or 1 million.
Suppose you want to convert a given amount in milliliters to liters. You
can use the relationship 1 L = 1000 mL. From this relationship, you can 1000 mL
_______ 1 L
and _______
1L 1000 mL
derive the conversion factors shown at the right.
Problem-Solving TIPS
• Multiply the given amount by the conversion factor that allows the units from which
you are converting to cancel out and the new units to remain.
• Most conversion factors are based on exact definitions, so significant figures do not
apply to these factors. The number of significant figures in a converted measurement
depends on the certainty of the measurement you start with.
SAMPLE
86 CHAPTER 3
CHAPTER 3 REVIEW
5. What are isotopes? How are the isotopes of a particular element alike and
how are they different?
At o m s : T he B u i l d i n g B l o c k s o f Matter 87
8. What is the mass in grams of each of the following?
a. 1.00 mol Li
b. 1.00 mol Al
e. 6.022 × 1023
atoms C
f. 6.022 × 1023
atoms Ag
a. 6.022 × 1023
atoms Ne
b. 3.011 × 1023
atoms Mg
10. How many moles of atoms are there in each of the following?
12. Naturally occurring boron is 80.20% boron-11 (atomic mass of 11.01 amu)
and 19.80% of some other isotopic form of boron. What must the atomic
mass of this second isotope be in order to account for the 10.81 amu average
atomic mass of boron? (Write the answer to two decimal places.)
88 CHAPTER 3
13. What is the mass in grams of each of the following?
At o m s : T he B u i l d i n g B l o c k s o f Matter 89
CHAPTER 4
Arrangement of Electrons
in Atoms
LOUISIANA
LOUISIANA STANDARDS
STANDARDS
90 CHAPTER 4
LOUISIANA STANDARDS
SECTION 4.1 LA.PS.8 Analyze the development of the
modern atomic theory from a historical
The Development of a
perspective. (PS-H-B1)
To the beach
λ 2. Which wave in the figure
to the left, (a) or (b), has a
higher frequency?
(a)
λ
10–2 nm 10–1 nm 100 nm 101 nm 102 nm 103 nm 10–3 cm 10–2 cm 10–1 cm 100 cm 101 cm 1m 101 m 102 m 103 m 104 m
Wavelength, λ
1019 Hz 1018 Hz 1017 Hz 1016 Hz 1015 Hz 1014 Hz 1013 Hz 1012 Hz 1011 Hz 1010 Hz 109 Hz 100 MHz 10 MHz 1 MHz 100 KHz
Frequency, ν
Electromagnetic spectrum
92 CHAPTER 4
The Photoelectric Effect
Light
In the early 1900s, scientists conducted two experiments on
light that could not be explained by a wave theory of light.
One experiment involved the photoelectric effect.
Stream of electrons
In the photoelectric effect, a metal gives off electrons when Anode
light shines on it. The bulb at the right shows that light shining Cathode
(metal plate)
on the cathode causes the cathode to emit electrons. The
electrons then travel to the anode and an electric current flows Voltage source
from the cathode and the anode. The photoelectric effect only
happens if the frequency of the light is above a certain value.
In 1900 scientists knew that light is a form of energy. So the In the photoelectric effect, light strikes
ability of light to cause a metal to give off electrons made the surface of the metal, causing
electrons to be ejected and an electric
sense. However, a new theory was needed to explain why light
current to flow.
of lower frequencies did not cause the photoelectric effect.
READING CHECK
A rra n g e m e n t o f E l ectr o n s i n At o m s 93
Photons A photon is a particle of electromagnetic radiation
having no mass and carrying a quantum of energy. The energy
of a photon is given by Planck’s formula for quanta of energy.
Ephoton= hν
Einstein explained the photoelectric effect by proposing READING CHECK
that matter can only absorb a whole number of photons. In
5. Use the quantites mass and
other words, matter cannot absorb part of a photon. An energy to complete this
electron can only be ejected from an object if one photon sentence.
is absorbed that has enough energy to knock it loose. Even a The of a photon
large stream of photons striking an object will fail to knock
is zero, while the
any electrons loose if the individual photons contain too
little energy. of a photon is
The Hydrogen-Atom
Emission-Line Spectrum
In addition to the photoelectric effect experiments, a second
set of experiments forced scientists to revise their ideas about
light. These experiments focused on light emitted by gases. A firefly, or lightning bug, uses a
chemical reaction to produce light.
When electric current passes through a gas at low pressure, The spectrum of light emitted by the
the gas gives off colored light. For example, when electric chemical luciferin gives the bugs their
current passes through neon gas, a neon sign lights up. yellow-green glow.
434 nm
486 nm
656 nm
94 CHAPTER 4
Lyman series Balmer series Paschen series
(ultraviolet) (visible) (infrared)
e a dc b a c b a
d c b
0 200 400 600 800 1000 1200 1400 1600 1800 2000
Wavelength (nm)
READING CHECK
A rra n g e m e n t o f E l ectr o n s i n At o m s 95
Bohr Model of the Hydrogen Atom
By the end of the nineteenth century, scientists had a formula
that related the different wavelengths in hydrogen’s emission-
line spectrum. Now scientists needed a model of the hydrogen
atom that accounted for this relationship. They also needed to
update Rutherford’s model to explain the nature, placement,
and motion of electrons in atoms. The Danish physicist Niels
Bohr provided this model in 1913.
According to Bohr’s model, the electron in a hydrogen
atom can only circle the nucleus in certain paths, or orbits.
According to Bohr’s model of the
The orbit the electron is in determines the energy of the atom. atom, electrons travel around the
The hydrogen atom is in its ground state when the electron is nucleus in specific energy levels.
in the orbit closest to the nucleus. The electron cannot go
closer to the nucleus than this orbit. If the electron enters an
orbit that is farther from the nucleus, the atom has more
energy. Each successive orbit represents a higher energy state.
The electron orbits, or atomic energy levels, can be
compared to the rungs of a ladder. When you stand on the
first rung, you have less potential energy than you would if
you were standing on any other, higher, rung of the ladder.
Because you cannot stand between the rungs, you can only
have amounts of potential energy corresponding to each rung.
E3
When the electron in a hydrogen atom is in one orbit, it Ephoton = E2-E1
e- E
neither gains nor loses energy. However, it can change orbits E1
2
96 CHAPTER 4
E
E6
E5
E4
E3
a b c
Paschen series
E2
a b c d
Energy
Balmer series
Critical Thinking
9. Summarize Explain how Bohr’s model built on the ideas of
Einstein and Planck.
A rra n g e m e n t o f E l ectr o n s i n At o m s 97
SECTION 4.1 REVIEW
VOCABULARY
1. Define the following.
a. electromagnetic radiation
b. quantum
REVIEW
2. What was the major shortcoming of Rutherford’s model of the atom?
3. Write and label the equation that relates the speed, wavelength,
=
and frequency of electromagnetic radiation.
4. What is meant by the dual wave-particle nature of light?
Critical Thinking
6. INTERPRETING GRAPHICS Use the diagram on page 97 to answer (a) and (b):
• highest energy?
• lowest energy?
98 CHAPTER 4
LOUISIANA STANDARDS
SECTION 4.2 LA.PS.8 Analyze the development of the
modern atomic theory from a historical
of the Atom
The Bohr model of the atom seemed at first to contradict
common sense. Bohr’s model explained the experimental
evidence of the hydrogen emission-line spectrum, but Bohr
provided no reason for the atom’s orbital structure.
A R R A N G E M E N T O F E L E C T R O N S I N AT O M S 99
The Heisenberg Uncertainty Principle
The idea of electrons having a dual wave-particle nature
troubled scientists. They wanted to be able to determine the
specific locations of electrons within an atom. In 1927, German
LOOKING CLOSER
physicist Werner Heisenberg developed another principle that
changed how scientists viewed the subatomic world. 2. Define uncertainty and
principle separately.
Heisenberg hypothesized that the act of observing a
uncertainty:
particle would itself change the behavior of the particle. For
example, to detect an electron, you use particles of light, or
photons. You locate the electron by its absorption, emission,
or other interaction with a photon. However, this interaction
will change the course of the electron’s movement. As a result,
there is always uncertainty in trying to locate an electron, or principle:
any other particle.
The Heisenberg uncertainty principle states that it is
impossible to determine the exact position and the exact
z
velocity of a particle at the same time. This principle is
now a foundation of current theories of the nature of light
and matter.
y
100 CHAPTER 4
The Four Quantum Numbers
Quantum Number Symbol Description
principal quantum n main energy level of the
number electron
angular momentum l shape of the orbital
quantum number
magnetic quantum m orientation of the orbital
number around the nucleus
spin quantum 1 or – __
1 spin state of the
+ __
number 2 2 electron
A R R A N G E M E N T O F E L E C T R O N S I N AT O M S 101
Angular Momentum Quantum Number
Not all orbitals are the same shape. At the n = 1 level, there is
just one orbital and it has a spherical shape. For higher main
energy levels, the orbitals can take on multiple shapes. The
angular momentum quantum number, symbolized by l,
indicates the shape of the orbital.
For a specific main energy level n, there are n possible
shapes for the orbitals. Each shape is called a sublevel. For
example, for the n = 2 level, there are two sublevels for
the orbitals: spherical (l = 0) and dumbbell-shaped (l = 1).
The n = 3 level includes three sublevels that include
spherical (l = 0), dumbbell-shaped (l = 1), and more
complex (l = 2) orbitals.
Each sublevel is also given a letter designation. The letter Orbital Letter Designations
designations are given in the table at the right. Every main According to Values of l
energy level includes an s orbital. Every energy level for n = 2 l Letter
and higher also includes p orbitals. Every energy level for
0 s
n = 3 and higher includes d orbitals. So n = 3 includes
s orbitals, p orbitals, and d orbitals. 1 p
2 d
Every atomic orbital has a designation that includes a
number followed by a letter. The number is the main energy 3 f
level, or principal quantum number. The letter is the sublevel.
For example, the 1s orbital is the only orbital on the main
energy level n = 1. A 4d orbital is any one of the d orbitals on
energy level n = 4. The information given by these first two
quantum numbers for the first four main energy levels of an
atom is summarized in the table on the page 103.
READING CHECK
z z z
y y y
x x x
102 CHAPTER 4
Orbital Types for the First Four Main Energy Levels
Prinicipal quantum Number of Possible values of Possible
number, n (main orbital shapes angular momentum orbital Orbital
energy level) possible quantum number, l types designations
n=1 1 l=0 s 1s
n=2 2 l = 0, 1 s, p 2s, 2p
n=3 3 l = 0, 1, 2 s, p, d 3s, 3p, 3d
n=4 4 l = 0, 1, 2, 3 s, p, d, f 4s, 4p, 4d, 4f
A R R A N G E M E N T O F E L E C T R O N S I N AT O M S 103
z z z
y y y
x x x
z z
y y
x x
PRACTICE
Designation 4s 4p 4p 4p 4d 4d 4d 4d 4d 4f 4f 4f 4f 4f 4f 4f
angular momentum
quantum number, l
magnetic quantum
number, m
104 CHAPTER 4
Spin Quantum Number
An electron in an orbital behaves in some ways like Earth LOOKING CLOSER
spinning on its axis. Earth’s spinning generates a magnetic 6. Which quantum numbers
field. An electron exists in one of two possible spin states. define the properties of electrons
Each spin state creates a different magnetic field. To account in an orbital?
for the magnetic properties of the electron, scientists assign
electrons a spin quantum number.
The spin quantum number has only two possible values,
+1/2 and –1/2, which indicate the two possible spin states of
an electron in an orbital. Each orbital of an atom can contain
7. Which quantum numbers define
up to two electrons. However, the electrons in the orbital must the properties of the orbitals?
have opposite spin states.
For example, examine the n = 2 main energy level. It has
two sublevels. The s sublevel includes one s orbital. The
p sublevel includes three p orbitals. Each one of these four
orbitals can contain two electrons if the electrons have
opposite spin states. Therefore, the n = 2 level of an atom can
hold up to 8 electrons. The structure of the atom given by
these quantum numbers is summarized in the table below.
READING CHECK
A R R A N G E M E N T O F E L E C T R O N S I N AT O M S 105
SECTION 4.2 REVIEW
VOCABULARY
1. Define each of the following.
b. quantum number
REVIEW
2. Identify the four quantum numbers by name and symbol.
Critical Thinking
5. INFERRING RELATIONSHIPS
106 CHAPTER 4
LOUISIANA STANDARDS
SECTION 4.3 LA.PS.15 Predict the physical and chemical
properties of an element based only on its
READING CHECK
6d
1. Name two ways the quantum model of the atom is an 7s 5f 6d
6p 5f
7s
improvement over Bohr’s model of the atom 5d 4f 6p
6s 5d
5p 6s 4f
4d 5p
5s
4p 5s 4d
3d 4p
4s
3d
Energy
3p 4s
3s 3p
3s
Rules Governing Electron 2p
2s 2p
Configurations 2s
relationship between the energy levels of the orbitals is The energy of each atomic sublevel is
given in the diagram above. Next, place the electrons in shown on the vertical axis. Each
orbitals, one by one, according to three basic rules. individual box represents one orbital.
A R R A n g E m E n T o f E l E C T R o n s i n AT o m s 107
Aufbau Principle According to the Aufbau principle, an
TIP In this book, electrons are
electron occupies the lowest-energy orbital that can receive it. represented by up and
down arrows, such as ↑ and ↓. The
The 1s orbital has the lowest energy. A hydrogen atom in the up arrow represents the +1/2 spin
ground state has an electron in this orbital. state, and the down arrow repre-
sents the –1/2 spin state.
The next four sublevels are 2s, 2p, 3s, and 3p. However, the
4s sublevel has a lower energy than the 3d sublevel, so the 4s
orbital is filled with electrons before any of the five 3d orbitals
are filled.
Pauli Exclusion Principle According to the Pauli exclusion
principle, no two electrons in the same atom can have the
same set of four quantum numbers. The principal, angular
momentum, and magnetic quantum numbers specify the
energy, shape, and orientation of an electron’s orbital. The
fourth quantum number specifies the spin of the electron.
Because there are only two possible values for the spin of
an electron, only two electrons can exist in the same orbital.
1s orbital
As first stated in Section 2 of this chapter, two electrons in the
same orbital must have opposite spin states. The box shows the electron configura-
tion of a helium atom. According to
Hund’s Rule According to Hund’s rule, orbitals of equal the Pauli exclusion principle, this
orbital can only contain two electrons
energy are each occupied by one electron before any are and they must have opposite spin
occupied by a second electron. In addition, all electrons in states.
orbitals with just one electron must have the same spin state.
Hund’s rule reflects the fact that electrons are arranged in
the lowest energy state possible. Electrons are not repelled as
strongly by other negatively-charged electrons if they are in
different orbitals, and therefore different regions, of the atom.
PRACTICE
108 CHAPTER 4
Representing Electron Configurations
Three types of notation are used to describe electron
configurations. This page discusses two types of notation, and
another part of this section will deal with noble-gas notation.
The three notations are summarized as follows:
• Orbital notation represents an orbital by showing its number
and letter below a horizontal line. Arrows above each line
represent electrons.
• Electron-configuration notation eliminates the lines and
arrows and represents the electrons by numbers listed to the
top right of an orbital’s name.
• Noble-gas notation uses the symbol for a noble gas in place
of part of the electron-configuration notation.
Orbital Notation
Orbital notation uses a line with the orbital’s name
Critical Thinking
underneath it. An unoccupied orbital is represented as . 2. Reasoning Why do you think
orbital notation is rarely used for
An orbital containing one electron is represented as ↑ , and atoms with more than ten electrons.
an orbital containing two electrons is represented as ↑↓ . The
table above shows the orbital notation for the elements in the
first period of the periodic table: hydrogen and helium.
Electron-Configuration Notation
Electron-configuration notation eliminates the lines and
arrows of orbital notation. Instead, this notation shows the
number of electrons in a sublevel by placing a number to
the top right of the sublevel’s designation. For example, the
electron configuration for sulfur is 1s2 2s2 2p6 3s2 3p4 . The
s sublevels of the sulfur atom all contain 2 electrons. The
2p sublevel contains 6 electrons, and the 3p sublevel contains
4 electrons. The table above also shows the electron
configuration notations for hydrogen and helium.
SOLUTION
2p
Next, fill in the electrons one at a time. Complete one sublevel
before beginning the next sublevel.
⇅ ⇅ ↑
1s 2s
}
2p
4 EVALUATE Determine if the answers make sense.
The number of arrows in each orbital match the numbers to
the top right of the orbital in the electron configuration.
PRACTICE
110 CHAPTER 4
Elements of the Second Period
The elements in the second period of the periodic table can
have up to five orbitals. The 1s orbital will be filled first,
followed by the 2s orbital, and then the 2p orbitals. Thus, the
first element of the second period, lithium, has a configuration
of 1s2 2s1 .
The electron occupying the 2s sublevel of a lithium atom is
in the atom’s highest-occupied energy level. For the first period
elements, the highest-occupied energy level is n = 1. For the
second period elements, the highest-occupied energy level is
n = 2. In these elements, the elements in the n = 1 level are ! Remember
The electrons within the same main
inner-shell electrons. These electrons are at an energy level energy level of an atom are said to
below the highest-occupied energy level. be in the same electron shell.
Noble-Gas Notation
The Group 18 elements in the periodic table are called the
noble gases. Neon is an example of a noble gas. The first ten
electrons in a sodium atom have the same configuration as a
neon atom. The same is true for all the third period elements.
Scientists used this fact to develop a shortened version of
electron-configuration notation called noble-gas notation. The
symbol for neon is enclosed in brackets and substituted for
that portion of the configuration. So, [Ne] = 1s2 2s2 2p6 , and the
noble-gas notation for sodium is [Ne]3s1 .
The last element in the third period is a noble gas called
argon. Argon, like neon, has an octet in its highest-occupied READING CHECK
energy level. In fact, all noble gases other than helium have
4. Use the information in the
such an octet. A noble-gas configuration refers to an outer other rows of the table to complete
main energy level occupied, in most cases, by eight electrons. the table.
112 CHAPTER 4
Elements of the Fourth Period
The fourth period of the periodic table contains elements in
which the highest-occupied energy level is the n = 4 level. 1s
Potassium, K, is the first element in the fourth period. It has
2s 2p
an atomic number of 19. Its first 18 electrons are placed in the
same way as the 18 electrons of the argon atom. These 3s 3p 3d
electrons fill the 1s, 2s, 2p, 3s, and 3p sublevels. However, 4s 4p 4d 4f
according to the diagram on page 107, the 4s sublevel has a 5s 5p 5d 5f
lower energy than the 3d sublevel. Therefore, the 19th electron
6s 6p 6d
fills the 4s sublevel, not the 3d sublevel. Potassium atoms have
7s 7p
the electron configuration [Ar]4s1 .
The next element is calcium, Ca, which has the electron Follow the diagonal arrows from the
top to get the order in which atomic
configuration [Ar]4s2 and an atomic number of 20. The next orbitals are filled according to the
element after that is scandium, Sc, with an atomic number Aufbau principle.
of 21. Its first 20 electrons are placed in the same way as the
calcium atom. The 21st electron is placed at the next lowest
energy sublevel, which is the 3d sublevel.
Scandium atoms have the electron configuration [Ar]3d1 4s2 .
Note that the 3d sublevel is written first, even though it has a
higher energy than the 4s sublevel. Electron-configuration
notation and noble-gas notation always show the sublevels in
order from lowest to highest main energy level.
The next nine elements fill the other remaining positions
within the 3d sublevel. The 4p sublevel is then filled starting
with gallium, Ga, element number 31. Therefore, the electron
configuration of a gallium atom is [Ar]3d10
4s2 4p1 .
The diagram at the top right provides another way to
remember the order in which atomic orbits are filled. Note
that according to the diagram, the d sublevel for a given
main energy level is never filled until the s sublevel of the
next-highest main energy level is filled.
READING CHECK
Calcium Ca 20 2 6 2 [Ar]4s2
Iron Fe 26 2 6 6 2
Copper Cu 29 2 6 10 1 [Ar]3d10
4s 1
Zinc Zn 30 2 6 10 2 [Ar]3d10
4s 2
Gallium Ga 31 2 6 10 2 1
Germanium Ge 32 [Ar]3d10
4s 2 4p2
Arsenic As 33 2 6 10 2 3
Selenium Se 34 [Ar]3d10
4s 2 4p 4
Bromine Br 35
Krypton Kr 36 2 6 10 2 6 [Ar]3d10
4s 2 4p6
2 2 6 2 6
*[Ar] = 1s 2s 2p 3s 3p
The table above shows all 18 elements in the fourth period READING CHECK
of the periodic table. There are two exceptions to the normal 6. Use the information in the
other rows of the table to complete
rules for placing electrons in orbitals that are reflected in the the table.
table. In each case, the configuration listed has the lowest
possible energy.The first exception is chromium, in which one
of the electrons in the 4s orbital switches to the 3d orbital.
The second exception is copper, in which the same switch
occurs. There is no simple explanation for this departure from
the pattern.
114 CHAPTER 4
Elements of the Fifth Period Critical Thinking
7. Explain Describe why
The patterns seen in the first four periods of the periodic table ruthenium, Ru, does not have
continue with the fifth period. There are 18 elements in the a noble-gas configuration.
fifth period. The sublevels are filled in the order 5s, 4d, and
finally 5p. All of the elements in the fifth period of the
periodic table have a highest-occupied energy level of n = 5.
The table at the bottom of the page shows the elements in the
fifth period. This table also includes configurations that differ
from those predicted by the rules on page 108.
Strontium Sr 38 2 6 2 [Kr]5s2
Palladium Pd 46 2 6 10 [Kr]4d10
Silver Ag 47 2 6 10 1 [Kr]4d10
5s1
Cadmium Cd 48 2 6 10 2 [Kr]4d10
5s2
Indium In 49 2 6 10 2 1 [Kr]4d10
5s2 5p1
Tin Sn 50 2 6 10 2 2 [Kr]4d10
5s2 5p 2
Antimony Sb 51 2 6 10 2 3 [Kr]4d10
5s2 5p3
Tellurium Te 52 2 6 10 2 4 [Kr]4d10
5s2 5p 4
Iodine I 53 2 6 10 2 5 [Kr]4d10
5s2 5p5
Xenon Xe 54 2 6 10 2 6 [Kr]4d10
5s2 5p6
2 2 6 2 6 10 2 6
*[Kr] = 1s 2s 2p 3s 3p 3d 4s 4p
SOLUTION
116 CHAPTER 4
PRACTICE
C. Write
both the complete electron-configuration
notation and the noble-gas notation for iodine, I.
E. Write
the complete electron configuration for the
element with atomic number 18.
READING CHECK
118 CHAPTER 4
SECTION 4.3 REVIEW
VOCABULARY
1. a. What is an atom’s electron configuration?
REVIEW
2. What three methods are used to represent the arrangement of electrons
in atoms?
b. [Ar]4s1
c. contains four electrons in its third and outer main energy level
Critical Thinking
5. RELATING IDEAS Write the electron configuration for the following
third-period elements.
a. aluminum, Al
b. silicon, Si
c. phosphorus, P
d. sulfur, S
e. chlorine, Cl
The atomic masses listed on the periodic table are not whole numbers. Instead they are
decimals that represent average atomic masses. The atomic masses are averages
because most elements occur in nature as a specific mixture of isotopes. For example,
75.76% of chlorine atoms have a mass of 34.969 amu, and 24.24% have a mass of
36.966 amu. If the isotopes were in a 1:1 ratio, you could simply add the masses of the
two isotopes together and divide by 2. However, to account for the differing abundance
of the isotopes, you must calculate a weighted average. For chlorine, the weighted
average is 35.45 amu. The following two examples demonstrate how weighted averages
are calculated.
Problem-Solving TIPS
• To find an average atomic mass, convert the abundance of each isotope from a
percentage to a decimal equivalent.
• Multiply each decimal equivalent with the atomic mass of each isotope. The result
is the contribution of the isotope to the weighted average.
• Add the contributions of each isotope. The sum is the average atomic mass.
SAMPLE
120 CHAPTER 4
CHAPTER 4 REVIEW
2. In the early twentieth century, what two experiments involving light and
matter could not be explained by the wave theory of light?
a. n = 1 c. n = 3
b. n = 2 d. n = 4
b. Hund’s rule
13. Write both the complete electron-configuration notation and the noble-gas
notation for each of the elements below.
a. sodium, Na
b. strontium, Sr
122 CHAPTER 4
14. The element silicon occurs as a mixture of three isotopes: 92.22%
Si-28, 4.69% Si-29, and 3.09% Si-30. The atomic masses of these
three isotopes are Si-28 = 27.976 926 amu, Si-29 = 28.976 495 amu,
and Si-30 = 29.973 770 amu. Find the average atomic mass of silicon.
16. How long would it take a radio wave whose frequency is 7.25 × 105 Hz
to travel from Mars to Earth if the distance between the two planets is
approximately 8.00 × 107 km?
17. How does a 2 s orbital differ from a 1s orbital? How does a 2px orbital differ
from a 2pyorbital?
18. List the order in which the orbitals generally fill, from 1s to 7p.
19. How do the electron configurations of chromium and copper contradict the
Aufbau principle?
124 CHAPTER 5
LOUISIANA STANDARDS
SECTION 5.1 LA.PS.15 Predict the physical and chemical
properties of an element based only on its
Periodic Table
By 1860, more than 60 elements had been discovered.
Vocabulary
Chemists had to learn the properties of these elements as
periodic law lanthanides
well as those of the many compounds that the elements periodic table actinides
formed. The elements were not organized into any patterns
and the factors that determined the properties of the
elements were unknown.
In addition, there was no method for determining an Critical Thinking
element’s atomic mass or the number of atoms of an 1. Infer How could organizing the
element in a compound. Different chemists used different elements according to trends or
patterns help chemists study the
atomic masses for the same elements, resulting in different properties of the elements?
compositions being listed for the same compounds. This
made communication of results between chemists difficult.
In 1860, Italian chemist Stanislao Cannizzaro presented a
method for accurately measuring the relative masses of
atoms. This allowed chemists to agree on standard values of
atomic mass. It also allowed them to search for relationships
between atomic mass and other properties of the elements.
T h e P e r i o d i c L aw 125
Mendeleev’s First Periodic Table
Ti = 50 Zr = 90 ? = 180
V = 51 Nb = 94 Ta = 182
Cr = 52 Mo = 96 W = 186
Mn = 55 Rh = 104.4 Pt = 197.4
Fe = 56 Ru = 104.4 Ir = 198
Ni = Co = 59 Pl = 106.6 Os = 199
H=1 Cu = 63.4 Ag = 108 Hg = 200
Be = 9.4 Mg = 24 Zn = 65.2 Cd = 112
B = 11 Al = 27.4 ? = 68 Ur = 116 Au = 197?
C = 12 Si = 28 ? = 70 Su = 118
N = 14 P = 31 As = 75 Sb = 112 Bi = 210
O = 16 S = 32 Se = 79.4 Te = 128?
F = 19 Cl = 35.5 Br = 80 I = 127
Li = 7 Na = 23 K = 39 Rb = 85.4 Cs = 133 Tl = 204
Ca = 40 Sr = 87.6 Ba = 137 Pb = 207
? = 45 Ce = 92
?Er = 56 La = 94
?Yt = 60 Di = 95
?In = 75.6 Th = 118?
126 CHA P T ER 5
Moseley and the Periodic Law
The periodic law states that the physical and chemical
properties of the elements are periodic functions of their
atomic numbers. In other words, when the elements are
arranged in order of increasing atomic number, elements with
similar properties appear at regular intervals.
Mendeleev’s periodic table ordered the elements by atomic
mass, not atomic number. In 1911, English scientist Henry
Moseley found that the elements fit into patterns better
when they were ordered by nuclear charge, or the amount of
positive charge in the nucleus. Moseley’s work led to the
modern definition of the atomic number and also to the
reorganization of the elements by atomic number instead of
atomic mass.
T h e P e r i o d i c L aw 127
The Lanthanides
The next step in the development of the periodic table
was completed in the early 1900s. The lanthanides are the
14 elements with atomic numbers from 58 (cerium, Ce) to
71 (lutetium, Lu). These elements have very similar properties.
The process of separating and identifying these elements
required a large effort by many chemists. However, finally it
was understood that the elements were best located in the
sixth period, between Group 3 and Group 4.
The Actinides
Another major step in the development of the periodic table
was the discovery of the actinides. The actinides are the
14 elements with atomic numbers from 90 (thorium, Th) to
103 (lawrencium, Lr). The actinides belong between Group 3
and Group 4 directly below the lanthanides. They are located Element Difference
and atomic in atomic
in the seventh period of the table. To save space, the number numbers
lanthanides and the actinides are often broken off and He 2
displayed below the rest of the periodic table. The periodic 8
Ne 10
table on pages 132–133 is one example of this form of display. 8
Group 18
Ar 18
18
Periodicity Kr 36
18
The periodicity of the atomic numbers is one of the most Xe 54
32
important features of the periodic table. Consider the noble Rn 86
gases in Group 18. The first noble gas is helium, He. It has an
atomic number of 2. The next element that has similar Li 3
properties to helium is neon, Ne, with atomic number 10. 8
Na 11
The rest of the noble gases are argon (Ar, atomic number 18), 8
Group 1
K 19
krypton (Kr, atomic number 36), xenon (Xe, atomic 18
number 54), and radon (Rn, atomic number 86). The Rb 37
18
progression of the differences between the atomic numbers Cs 55
32
of these elements is shown at the right. Fr 87
Critical Thinking
4. Calculate What is the pattern of differences between
the atomic numbers of Groups 2 and 13–17 of the
periodic table?
128 CHA P T ER 5
SECTION 5.1 REVIEW
VOCABULARY
1. State the periodic law.
REVIEW
2. a. Who is credited with developing a method that led to the
determination of standard atomic masses?
3. Name three sets of elements that have been added to the periodic table
after Mendeleev’s time.
4. How do the atomic numbers of the elements within each of Groups 1, 2, and
13–18 of the periodic table vary?
Critical Thinking
5. RELATING IDEAS Why are the atomic masses of the elements not strictly in
increasing order in the periodic table, even though the properties of the
elements are similar? For example, by atomic mass, tellurium, Te, should be
in Group 17 and iodine, I, should be in Group 16, but grouping by properties
has Te in Group 16 and I in Group 17.
T h e P e r i o d i c L aw 129
LOUISIANA STANDARDS
SECTION 5.2 LA.PS.15 Predict the physical and chemical
properties of an element based only on its
130 CHAPTER 5
Group 18
1 s-block elements
2
H
p-block elements Group 13 Group 14 Group 15 Group 16 Group 17
He
Group 1 Group 2
3 4 d-block elements 5 6 7 8 9 10
Li Be f-block elements B C N O F Ne
11 12 13 14 15 16 17 18
Na Mg Group 3 Group 4 Group 5 Group 6 Group 7 Group 8 Group 9 Group 10 Group 11 Group 12
Al Si P S Cl Ar
19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
87 88 89 104 105 106 107 108 109 110 111
Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Rg
58 59 60 61 62 63 64 65 66 67 68 69 70 71
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
90 91 92 93 94 95 96 97 98 99 100 101 102 103
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
READING CHECK
T he P er i o d i c L aw 131
Periodic Table of the Elements
1
H
1
Hydrogen Key:
1.007 94
1s 1
Group 1 Group 2
Atomic number 6
3 4 Symbol
C
2
Li Be Name Carbon
Lithium Beryllium
6.941 9.012 182
Average atomic mass 12.0107
[He]2s 1 [He]2s 2 Electron configuration [He]2s22p2
11 12
3
Na Mg
Sodium Magnesium
22.989 769 28 24.3050
[Ne]3s1
Period
4
K Ca Sc Ti V Cr Mn Fe Co
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt
39.0983 40.078 44.955 912 47.867 50.9415 51.9961 54.938 045 55.845 58.933 195
[Ar]4s 1 [Ar]4s 2 [Ar]3d14s 2 [Ar]3d24s 2 [Ar]3d34s 2 [Ar]3d54s1 [Ar]3d54s 2 [Ar]3d64s 2 [Ar]3d74s 2
37 38 39 40 41 42 43 44 45
5
Rb Sr Y Zr Nb Mo Tc Ru Rh
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium
85.4678 87.62 88.905 85 91.224 92.906 38 95.94 (98) 101.07 102.905 50
[Kr]5s1 [Kr]5s 2 [Kr]4d15s 2 [Kr]4d25s 2 [Kr]4d45s1 [Kr]4d55s 1 [Kr]4d65s1 [Kr]4d75s1 [Kr]4d85s1
55 56 57 72 73 74 75 76 77
6 Cs Ba La Hf Ta W Re Os Ir
Cesium Barium Lanthanum Hafnium Tantalum Tungsten Rhenium Osmium Iridium
132.905 4519 137.327 138.905 47 178.49 180.947 88 183.84 186.207 190.23 192.217
[Xe]6s1 [Xe]6s 2 [Xe]5d16s 2 [Xe]4f145d26s 2 [Xe]4f145d36s 2 [Xe]4f145d46s 2 [Xe]4f145d56s 2 [Xe]4f145d66s 2 [Xe]4f145d76s2
7 Fr Ra Ac Rf Db Sg Bh Hs Mt
Francium Radium Actinium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium
(223) (226) (227) (261) (262) (266) (264) (277) (268)
[Rn]7s1 [Rn]7s 2 [Rn]6d17s 2 [Rn]5f146d27s2 [Rn]5f146d37s 2 [Rn]5f146d47s 2 [Rn]5f146d57s 2 [Rn]5f146d67s 2 [Rn]5f146d77s2
90 91 92 93 94
Th Pa U Np Pu
Thorium Protactinium Uranium Neptunium Plutonium
232.038 06 231.035 88 238.028 91 (237) (244)
[Rn]6d27s 2 [Rn]5f 26d17s 2 [Rn]5f 36d17s 2 [Rn]5f 46d17s 2 [Rn]5f67s 2
132 CHAPTER 5
Hydrogen
Semiconductors
(also known as metalloids) Group 18
2
Metals
Alkali metals
He
Helium
Alkaline-earth metals 4.002 602
1s 2
Transition metals Group 13 Group 14 Group 15 Group 16 Group 17
Other metals 5 6 7 8 9 10
Nonmetals B C N O F Ne
Halogens Boron Carbon Nitrogen Oxygen Fluorine Neon
10.811 12.0107 14.0067 15.9994 18.998 4032 20.1797
Noble gases [He]2s 22p 1 [He]2s 22p 2 [He]2s 22p 3 [He]2s 22p 4 [He]2s 22p 5 [He]2s 22p 6
Other nonmetals
13 14 15 16 17 18
Al Si P S Cl Ar
Aluminum Silicon Phosphorus Sulfur Chlorine Argon
26.981 5386 28.0855 30.973 762 32.065 35.453 39.948
[Ne]3s 23p 1 [Ne]3s 23p 2 [Ne]3s 23p 3 [Ne]3s 23p 4 [Ne]3s 23p 5 [Ne]3s 23p 6
Group 10 Group 11 Group 12
28 29 30 31 32 33 34 35 36
Ni Cu Zn Ga Ge As Se Br Kr
Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
58.6934 63.546 65.409 69.723 72.64 74.921 60 78.96 79.904 83.798
[Ar]3d 84s 2 [Ar]3d 104s 1 [Ar]3d 104s 2 [Ar]3d 104s 24p 1 [Ar]3d 104s 24p 2 [Ar]3d 104s 24p 3 [Ar]3d 104s 24p 4 [Ar]3d 104s 24p 5 [Ar]3d 104s 24p 6
46 47 48 49 50 51 52 53 54
Pd Ag Cd In Sn Sb Te I Xe
Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
106.42 107.8682 112.411 114.818 118.710 121.760 127.60 126.904 47 131.293
[Kr]4d 105s 0 [Kr]4d 105s 1 [Kr]4d 105s 2 [Kr]4d 105s 25p 1 [Kr]4d 105s 25p 2 [Kr]4d 105s 25p 3 [Kr]4d 105s 25p 4 [Kr]4d 105s 25p 5 [Kr]4d 105s 25p 6
78 79 80 81 82 83 84 85 86
Pt Au Hg Tl Pb Bi Po At Rn
Platinum Gold Mercury Thallium Lead Bismuth Polonium Astatine Radon
195.084 196.966 569 200.59 204.3833 207.2 208.980 40 (209) (210) (222)
[Xe]4f 145d 96s 1 [Xe]4f 145d 106s 1 [Xe]4f 145d 106s 2 [Xe]4f 145d 106s 26p 1 [Xe]4f 145d 106s 26p 2 [Xe]4f 145d 106s 26p 3 [Xe]4f 145d 106s 26p 4 [Xe]4f 145d 106s 26p 5 [Xe]4f 145d 106s 26p 6
The discoveries of elements with atomic numbers 112, 114, and 116 have been reported but not fully confirmed.
63 64 65 66 67 68 69 70 71
Eu Gd Tb Dy Ho Er Tm Yb Lu
Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium
151.964 157.25 158.925 35 162.500 164.930 32 167.259 168.934 21 173.04 174.967
[Xe]4f 76s 2 [Xe]4f 75d 16s 2 [Xe]4f 96s 2 [Xe]4f 106s 2 [Xe]4f 116s 2 [Xe]4f 126s 2 [Xe]4f 136s 2 [Xe]4f 146s 2 [Xe]4f 145d 16s 2
The atomic masses listed in this table reflect the precision of current measurements. (Each value listed in
parentheses is the mass number of that radioactive element’s most stable or most common isotope.)
T he P er i o d i c L aw 133
The s-Block Elements: Groups 1 and 2
The elements in the s-block are chemically reactive metals.
TIP Some groups in the
The Group 1 metals are more reactive than the Group 2 periodic table are
also represented by their own
metals. As you will learn in Section 3, Group 1 metals are configuration. This notation reflects
reactive because the single electron in their highest-occupied the fact that each element in the
energy level is easily lost. group has the same configuration in
its highest-occupied energy level.
The elements in Group 1 (lithium, sodium, potassium, For example, Group 1 has a group
rubidium, cesium, and francium) are known as the alkali configuration of ns1 , and Group 2
has a group configuration of ns2 .
metals. In their pure state, all of the alkali metals have a
silvery appearance and are soft enough to cut with a knife.
As you move down the column in the periodic table, the pure
forms of these elements melt at lower and lower temperatures. READING CHECK
3. Which is more reactive,
potassium or calcium?
(a) (b)
134 CHAPTER 5
SAMPLE PROBLEM
a. Without looking at the periodic table, identify the group, period,
and block in which the element that has the electron configuration
[Xe]6s2 is located.
b. Without looking at the periodic table, write the electron
configuration for the Group 1 element in the third period.
SOLUTION
A. Identify
the following for the element with the electron
configuration [Kr]5s1.
Group number Period number
Block designation
B. Write
the noble-gas configuration for the Group 2
element in the fourth period.
T he P er i o d i c L aw 135
Hydrogen and Helium
Hydrogen has the electron configuration 1s1 , but does not
share the same properties as the Group 1 elements. It is a
unique element, whose properties do not closely resemble
those of any group. It is usually placed above the Group 1
elements with a gap between it and lithium.
Helium has the same ns2 configuration as the Group 2
elements. However, it has the stability of a noble gas because
its outermost energy level is filled. It is placed in Group 18,
instead of with the reactive, unstable Group 2 elements. The transition element copper is
found in nature as part of copper
The d-Block Elements: Groups 3–12 ore and can be obtained from
surface mines.
The d-block elements are metals with typical metallic
properties and are often referred to as transition elements or
transition metals. They are good conductors of electricity and
have a high luster. They are usually less reactive than the
alkali metals and alkaline-earth metals. Some transition
elements, such as gold, palladium, and platinum, are so
unreactive that they exist in nature mostly as free elements.
The last electron placed in an atom of a d-block element
is in a d orbital. The d block spans ten groups on the periodic
table. The group configuration of Group 3 elements is (a)
(n – 1)d1 ns2 . For example, scandium has the configuration
[Ar]3d14s2 , as shown below. The group configuration of
Group 12 elements is (n – 1)d10ns2 . One example is mercury,
with the configuration [Xe]4f 145d10 6s2 . Groups 4–11 do not
have regular configurations, but the sum of the electrons in the
highest s and d sublevels is always equal to the group number.
(b)
Energy
4s 3d
3s 3p
The pure forms of (a) tungsten
and (b) vanadium are shown.
2p
2s
1s
READING CHECK
136 CHAPTER 5
SAMPLE PROBLEM
An element has the electron configuration [Kr]4d5 5s1 . Without
looking at the periodic table, identify the period, block, and
group in which this element is located. Then, consult the
periodic table to identify this element.
SOLUTION
C. Identify
the following for the element with the electron
configuration [Ar]3d8 4s2 .
Period number Block designation
Group number
D. Write
the outer electron configuration for the Group 12
element in the fifth period.
T he P er i o d i c L aw 137
Relationships Among Group Numbers, Blocks, and Electron Configurations
3–12 (n - 1)d1–10ns0–2
d Sum of electrons in ns and
(n - 1)d levels equals group number
! Remember
The properties of the elements in the p block vary greatly.
The right side of the p block contains all of the nonmetals
except hydrogen and helium. All six of the metalloids (boron, A nonmetal is a poor conductor
of electricity and heat. A metalloid,
silicon, germanium, arsenic, antimony, and tellurium) are or semiconductor, has some
also in the p block. The bottom left of the p block contains properties of metals and some
eight metals. properties of nonmetals.
READING CHECK
138 CHAPTER 5
(a)
Fluorine
T he P er i o d i c L aw 139
SAMPLE PROBLEM
Without looking at the periodic table, write the outer electron
configuration for the Group 14 element in the second period.
Then, name the element and identify it as a metal, nonmetal,
or metalloid.
SOLUTION
E. Write
the outer electron configuration for the Group 17
element in the third period.
F. Identify
the following for the element with the electron
configuration [Ar]3d10
4s2 4p3.
Period number Block designation
Group number
140 CHAPTER 5
SAMPLE PROBLEM
Use the given electron configuration to determine the period, group,
and block for these four elements. Then use the periodic table to
name each element, identify it as a metal, nonmetal, or metalloid,
and describe its relative reactivity.
SOLUTION
5d 9 6s 1
a. [Xe]4f14 6 d 10 platinum metal low
PRACTICE Complete the table below. Fill in the first three columns without
consulting the periodic table.
H. [Ar]3d10
4s 1
T he P er i o d i c L aw 141
SECTION 5.2 REVIEW
VOCABULARY
1. What name is given to each of the following groups of elements in the
periodic table?
REVIEW
2. Into what four blocks can the periodic table be divided to illustrate the
relationship between the elements’ electron configurations and their
placement in the periodic table?
3. What are the relationships between group configuration and group number
for elements in the s, p, and d blocks?
4. Without looking at the periodic table, write the outer electron configuration
5. Without looking at the periodic table, identify the period, block, and group
of the element that has the electron configuration [Ar]3d7 4s2 .
Critical Thinking
6. APPLYING MODELS Period 7 contains elements in the s, p, d, and f blocks.
Suppose there were a Period 8 and it contained elements in the “g” block,
where “g” had the angular momentum quantum number of 4. If a
hypothetical element in Period 8 had an atomic number of 120, into what
group in the periodic table would the element fit, and what properties might
it have (assuming it does not radioactively decay)?
142 CHAPTER 5
LOUISIANA STANDARDS
SECTION 5.3 LA.PS.13 Identify the number of bonds an
atom can form given the number of valence
Atomic Radii
Ideally, the size of an atom is defined by the edge of its Chlorine
nucleus
outermost orbital. However, the location of this boundary
Atomic
is difficult to determine, and can vary under different radius
conditions. Therefore, to estimate the size of an atom, the 99 pm
T H E P E R i o d i C L Aw 143
Period Trends
The table below lists the atomic radii of all of the elements in
the periodic table in picometers. The graph on the next page
plots the atomic radii of the elements versus their atomic
number. By examining the table and the graph, several trends
in the atomic radius can be observed.
Note that there is a gradual decrease in atomic radius as
you move from left to right across the second period of the
periodic table. The atomic radius of lithium in Group 1 is
152 pm, and the values decrease through neon in Group 18,
which has an atomic radius of 71 pm.
The trend to smaller atoms across a period is caused by the
increasing positive charge of the nucleus. As electrons are
added to the s and p levels in the same main energy level, they READING CHECK
are gradually pulled closer to the positively charged nucleus 2. If two elements are from the
that contains more and more protons. This increased pull same period, which element is most
results in a decrease in atomic radius. Because the increased likely to have the largest atomic
radius, the element from Group 13
pull is offset somewhat by the repulsion of more and more or the element from Group 17?
electrons in the same main energy level, atomic radii decrease
more slowly on the right side of the periodic table.
K 19 Ca 20 Sc 21 Ti 22 V 23 Cr 24 Mn 25 Fe 26 Co 27 Ni 28 Cu 29 Zn 30 Ga 31 Ge 32 As 33 Se 34 Br 35 Kr 36
Period
4 4
227 197 162 147 134 128 127 126 125 124 128 134 135 122 120 119 114 112
Rb 37 Sr 38 Y 39 Zr 40 Nb 41 Mo 42 Tc 43 Ru 44 Rh 45 Pd 46 Ag 47 Cd 48 In 49 Sn 50 Sb 51 Te 52 I 53 Xe 54
5 5
248 215 180 160 146 139 136 134 134 137 144 149 167 140 140 142 133 131
Cs 55 Ba 56 La 57 Hf 72 Ta 73 W 74 Re 75 Os 76 Ir 77 Pt 78 Au 79 Hg 80 Tl 81 Pb 82 Bi 83 Po 84 At 85 Rn 86
6 6
265 222 183 159 146 139 137 135 136 139 144 151 170 175 150 168 140 141
Fr 87 Ra 88 Ac 89 Rf 104 Db 105 Sg 106 Bh 107 Hs 108 Mt 109 Ds 110 Rg 111 112 113 114 115 116 117 118
7 7
270 220 188 — — — — — — — —
Lanthanide series
Ce 58 Pr 59 Nd 60 Pm 61 Sm 62 Eu 63 Gd 64 Tb 65 Dy 66 Ho 67 Er 68 Tm 69 Yb 70 Lu 71
182 182 181 183 180 208 180 177 178 176 176 176 — 174
Th 90 Pa 91 U 92 Np 93 Pu 94 Am 95 Cm 96 Bk 97 Cf 98 Es 99 Fm 100 Md101 No 102 Lr 103
In general, atomic radii decrease from left to right across a period and increase
from top to bottom down a group.
144 CHAPTER 5
Atomic Radius vs. Atomic Number
300
Period Period Period Period Period Period
2 3 4 5 6 Fr 7
Cs
Rb
250
K
200
Es
Atomic radius (pm)
Na
Li
150
Rn
Xe
Kr
100 Ar
Ne
50
H
He
0
0 10 20 30 40 50 60 70 80 90 100
Atomic number
PRACTICE
B. hich of Ca, Be, Ba, and Sr has the largest atomic radius?
W
Explain your answer in terms of trends in the periodic table.
T he P er i o d i c L aw 145
Ionization Energy
If an atom absorbs enough energy, it can lose an electron.
Using A as a symbol for an atom of any element, the process
can be expressed as follows.
+
A + energy → A + e-
The symbol A+represents an ion of element A with a
single positive charge. For example, Na+ refers to a sodium
ion. Such an ion is also referred to as a 1+ ion. An ion is an
atom or group of bonded atoms that has a positive or negative
charge. An atom gains positive charge when an electron is lost.
Any process during which an ion is formed is referred to as
ionization. The energy required to remove one electron from a
neutral atom of an element is the ionization energy, or IE.
Scientists calculate the ionization energies of atoms using
READING CHECK
isolated atoms in the gas phase. This removes the influence of
3. What does the chemical symbol
other atoms from the process. The table below lists ionization Li+
refer to?
energies in kJ/mol. The graph on the next page plots the
ionization energies versus atomic number.
C
Group 1 Group 2 Group 13 Group 14 Group 15 Group 16 Group 17 2372
Symbol
3 4 5 6 7 8 9 10
Li Be B C N O F Ne 2
1086
2
520 900 801 1086 1402 1314 1681 2081
First ionization
11 12 energy 13 14 15 16 17 18
3 Na Mg Al Si P S Cl Ar 3
496 738 Group 3 Group 4 Group 5 Group 6 Group 7 Group 8 Group 9 Group 10 Group 11 Group 12 578 787 1012 1000 1251 1521
19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36
Period
Period
4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 4
419 590 633 659 651 653 717 762 760 737 746 906 579 762 947 941 1140 1351
37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54
5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 5
403 550 600 640 652 684 702 710 720 804 731 868 558 709 834 869 1008 1170
55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86
6 Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 6
376 503 538 659 761 770 760 839 878 868 890 1007 589 716 703 812 — 1038
87 88 89 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118
7 Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Rg 7
— 509 490 — — — — — — — —
Lanthanide series
58 59 60 61 62 63 64 65 66 67 68 69 70 71
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
534 527 533 536 545 547 592 566 573 581 589 597 603 523
90 91 92 93 94 95 96 97 98 99 100 101 102 103
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
587 570 598 600 585 578 581 601 608 619 627 635 642 —
Actinide series
In general, first ionization energies increase from left to right across a period
and decrease from top to bottom down a group.
146 CHAPTER 5
First Ionization Energy vs. Atomic Number
3000
Period Period Period Period Period Period
2 3 4 5 6 7
2500
He
First ionization energy (kJ/mol)
Ne
2000
Ar
1500
Kr
H Xe
Rn
1000
No
Ra
Li
500 Na
K Rb Cs
0
0 10 20 30 40 50 60 70 80 90 100
Atomic number
The plot of first ionization energy versus atomic number shows both period trends
and group trends.
Period Trends
By examining the table on page 146 and the graph above,
TIP In the graph above, note
several trends in the ionization energy can be observed. that the period trends and
group trends are less pronounced as
Note that there is a gradual increase in ionization energy as the atomic number increases.
you move from left to right across a period in the periodic
table. Another trend is that nonmetals usually have higher
ionization energies than metals.
The Group 1 metals have the lowest ionization energies in
their respective periods. The ease with which these metals lose
electrons is a major reason for their high reactivity. On the
opposite side of the periodic table, the Group 18 elements
have the highest ionization energies. This is a major reason
READING CHECK
why noble gases are so unreactive.
4. Suppose you have an alkaline-
The general increase in ionization energies moving across a earth metal and a halogen.
Which is more likely to have a
period is caused by the increasing charge in the nucleus. A higher ionization energy?
higher positive charge in the nucleus attracts the negatively
charged electrons more strongly.
T he P er i o d i c L aw 147
Group Trends
The ionization energy tends to decrease when moving down a
group. This is related to the increase in the atomic radii of
elements with larger atomic numbers. The force that binds an
electron in the outermost energy level to the atom is weaker
as the electron moves farther from the positively charged
nucleus. These electrons are removed more easily.
In addition to electrons being farther from the center of
positive charge, another effect contributes to the group trend
in ionization energies. As atomic number increases, more and
more electrons are located between the nucleus and the
outermost electrons. These inner shell electrons exert a
repulsive force on the outermost electrons, further weakening
their bond with the atom.
PRACTICE
148 CHAPTER 5
Ionization Energies (in kJ/mol) for Elements of Periods 1–3
Period 1 Period 2
H He Li Be B C N O F Ne
IE1 1312 2372 520 900 801 1086 1402 1314 1681 2081
IE2 5250 7298 1757 2427 2353 2856 3388 3374 3952
IE3 11 815 14 849 3660 4621 4578 5300 6050 6122
IE4 21 007 25 026 6223 7475 7469 8408 9370
IE5 32 827 37 830 9445 10 990 11 023 12 1 78
Period 3
Na Mg Al Si P S Cl Ar
IE1 496 738 578 787 1012 1000 1251 1521
IE2 4562 1451 1817 1577 1903 2251 2297 2666
IE3 6912 7733 2745 3232 2912 3361 3822 3931
IE4 9544 10 540 11 578 4356 4957 4564 5158 5771
IE5 13 353 13 628 14 831 16 091 6274 7013 6540 7238
The table above shows the first five ionization energies for Critical Thinking
the elements in the first three periods of the periodic table.
6. Infer Suppose that the table
For each element, IE2 is always higher than IE1 , IE3 is always above was extended to list the
higher than IE2 , and so on. This is because each successive sixth ionization energies for
electron removed from an ion feels a stronger and stronger the elements. Which element in the
second period, and which element
effective nuclear charge from the nucleus. In other words, as in the third period, would have the
each electron is removed, fewer electrons remain within the highest value for IE6 ?
atom to shield the attractive force of the nucleus.
Removing the first electron from a noble gas is more
difficult than removing the first electron from any other
element. This is due to the stability of the octet in the highest-
occupied energy level of the noble-gas atom. This concept also
applies to ions that have a noble-gas configuration.
For example, when sodium, Na, loses a single electron, it
has a noble-gas configuration. You can see why when you look
at sodium’s electron configuration, [Ne]3s1 . So, there is a large
increase in sodium’s IE2 compared to its IE1 . In the same way,
magnesium has a large jump in IE3from IE2 . Magnesium’s
electron configuration, [Ne]3s2, shows that it has a noble-gas
configuration after it loses two electrons. The table uses
highlighting to emphasize the large jumps in ionization
energies caused by the noble-gas configurations of some ions.
T he P er i o d i c L aw 149
Electron Affinity
Not only can neutral atoms lose electrons; they can gain
TIP In a chemical equation,
electrons. The energy change that occurs when a neutral atom energy is released if it is
shown on the right side of the
gains an electron is called the atom’s electron affinity. Most equation. Energy is absorbed if it
atoms release energy when they gain an electron. is shown on the left side of a
chemical equation.
A + e-→ A-
+ energy
On the other hand, some atoms must absorb energy before
they will gain an electron. An ion produced in this way is
unstable and will usually lose the added electron quickly.
A + e -
+ energy → A-
Period Trends
The table below shows the electron affinities for most READING CHECK
elements. The graph on the next page shows a plot of these 7. Why are most of the
data. The values are negative to show that the atom loses electron affinities in the table
below negative?
energy when it gains an electron. The electron affinities for
atoms that gain energy when gaining an electron are difficult
to determine, so they are given as zeros in parentheses.
The Group 17 elements, the halogens, have the electron
affinities that are most negative. Their ability to attract
electrons helps to explain why halogens are so reactive.
The electron affinities generally become more negative when
moving left to right across the p block.
Period
5 –50.1 (0) –18.8 –7.9 –52.5 –66.6 (0) –16.3 –66.1 –115.6 –122.8 (0) –30 –135 –81 –202.1 –336.5 (0)
5
37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
6 –48.6 (0) –30.7 –42.6 –89.3 –74.6 –55 –105 –113.7 –55.7 –130.2 (0) –30 –120 –107 –197.1 –305.9 (0)
6
55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
7 –47.2 (0) –50 (0) –32.2 –81.5 –15 –110 –156.5 –212.8 –230.9 (0) –20 –36 –94.6 –190 –280 (0)
7
87 88 89 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118
Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Rg
–47.0 (0) — — — — — — — — —
This table shows the electron affinities for the s-, p-, and d-block elements. Values
in parentheses are approximate. All of the lanthanides have an estimated electron
affinity of – 50 kJ/mol. All of the actinides have an estimated electron affinity
of 0 kJ/mol.
150 CHAPTER 5
Electron Affinity vs. Atomic Number
0
He Ne Ar Kr Xe Rn Ra
K Rb Cs Fr
Na
Li
H
–100
Electron affinity (kJ/mol)
–200
–300
One exception to the trend is that it is easier to add an The plot of electron affinity versus
atomic number shows that most atoms
electron into the empty third p orbital of Group 14 elements
release energy when they acquire an
than it is to add an electron into one of the partially filled electron because the electron affinity
p orbitals in the Group 15 elements. This makes the Group 14 is negative.
electron affinities more negative than those in Group 15.
Group Trends
Electron affinities become slightly less negative as you move
down a group. Because atomic radii increase moving down a
group, electrons are less likely to be attracted to the nucleus.
The strength of the positively charged nucleus relative to the
electrons around the nucleus also decreases.
T he P er i o d i c L aw 151
Ionic Radii
The radius of an ion is different from the radius of the atom
from which it formed. There are two types of ions, depending
on whether the ion has a positive charge or a negative charge.
A cation is a positive ion. A cation forms when an atom
TIP The number next to the
loses one or more electrons. A cation is always smaller than charge in an ion gives the
strength of the charge. For example,
the atom from which it was formed because of the removal of an Fe2+ion is a 2+ ion formed from
the electrons from the highest energy level. Also, the nucleus an iron atom.
draws the remaining electrons closer because its positive
charge is greater than the negative charge of the electrons.
An anion is a negative ion. An anion forms when an atom
gains one or more electrons. Since an anion has more negative
charge in the electron cloud than positive charge in the
nucleus, the electron cloud is not pulled as close to the
nucleus. The electron cloud also spreads out because of the
repulsion between the increased number of electrons. These
two factors make an anion larger than the original atom.
READING CHECK
–
b. F e. Te2–
c. Fe2+ f. C4–
19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36
4 4
K+ 138 Ca2+ 100 Sc3+ 75 Ti2+ 86 V2+ 79 Cr2+ 80 Mn2+ 83 Fe2+ 78 Co2+ 65 Ni2+ 69 Cu2+ 73 Zn2+ 74 Ga3+ 62 Ge — As — Se2– 198 Br– 196 Kr —
37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54
5 5
Rb+ 152 Sr2+ 118 Y3+ 90 Zr — Nb — Mo — Tc — Ru — Rh3+ 67 Pd2+ 86 Ag+ 115 Cd2+ 95 In3+ 80 Sn2+ 118 Sb3+ 76 Te2– 221 I– 220 Xe —
55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86
6 6
Cs+ 167 Ba2+ 136 La3+ 116 Hf — Ta — W — Re — Os — Ir — Pt2+ 80 Au+ 137 Hg2+ 119 Tl3+ 89 Pb2+ 119 Bi3+ 103 Po — At — Rn —
87 88 89 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118
7 7
Fr+ 180 Ra2+ 148 Ac3+ 111 Rf — Db — Sg — Bh — Hs — Mt — Ds — Rg —
This table shows the ionic radii for the most common ions of the s-, p-,
and d- block elements. Cations are smaller and anions are larger than the
original atoms.
152 CHAPTER 5
Period Trends
The metals at the left of the periodic table tend to form
cations. This is because it is easier for one of these atoms to
lose the few electrons in its highest-occupied energy level than
it is for one of them to gain electrons to form an octet. On the
other hand, the nonmetals at the upper right of the periodic
table tend to form anions. It is easier for one of these atoms to
form a stable octet by gaining a few electrons than to lose all
of the electrons in the outermost energy level.
The ionic radii of cations decrease as the charge on the
cation becomes stronger. For example, the lithium ion, Li+
is
2+
much larger than the beryllium ion, Be . The more the
charges are unbalanced, the stronger the force with which
READING CHECK
the nucleus can pull on the electron cloud.
10. Describe whether the following
The ionic radii of anions increase as the charge on the pairs of particles repel each other
anion becomes stronger. The more negative the charge on or attract each other.
the ion, the less force with which the nucleus can pull on the a. anion and anion
electrons. Since the charge becomes less negative, moving
from Group 15 to Group 17, the ionic radii of anions tend to
decrease moving across a period from left to right. b. cation and anion
Group Trends
There is a gradual increase in ionic radii moving down a group
for the same reason as there is a gradual increase in atomic
radii moving down a group. The highest-occupied energy level
increases when moving down a group, so the electron cloud
extends farther from the nucleus of the ion.
Critical Thinking
11. Infer Whyare there no common ions or ionic radii listed
for the noble-gas elements in Group 18 of the periodic
table on page 152?
T he P er i o d i c L aw 153
PRACTICE
Q: block T: block
R: block X: block
b. Which of these elements are in the same period?
d. Which
would you expect to have the highest IE1 ? Which
would have the lowest IE1 ?
k. Using
the letters Q, R, T, and X, write the symbol that represents the most
common ion that each hypothetical element would form.
154 CHAPTER 5
Valence Electrons
Chemical compounds form when atoms of two or more
TIP If you look back to the
elements bond together. These bonds form when electrons are table of ionic radii on
page 152, you will see the most
lost, gained, or shared between atoms. The electrons involved common ions listed for each main
in the process of forming chemical bonds are the electrons in group element. For Groups 1, 2, and
the outer energy levels. These electrons are the most likely to 13, the most common ion is a cation
with a charge equal to the number
be influenced by other atoms or ions.
of valence electrons in the atom.
The electrons available to be lost, gained, or shared in the For Groups 15, 16, and 17, the most
common ion is an anion with a
formation of chemical compounds are called valence electrons. charge equal to the number of
Valence electrons are usually located in main energy levels valence electrons needed to form
that are incompletely filled. For example, a sodium atom has an octet of eight electrons.
an electron configuration of [Ne]3s1 . The n = 3 level of the
sodium atom only has a single electron. That one valence
a+ion forms.
electron is lost when the N
For main group elements, the valence electrons are the
electrons in the outermost s and p sublevels. The other main
energy levels are completely filled. The attraction between the
nucleus and the inner-shell electrons is too strong for these
electrons to become involved in forming chemical bonds.
The number of valence electrons in the atoms of a main-
group element is determined by the group configuration. This
configuration shows the number of electrons in the outermost
s and p sublevels. For example, a Group 13 element has two READING CHECK
electrons in its outermost s sublevel and one electron in its
12. Use the information in the
outermost p sublevel. Therefore, a Group 13 element has three other rows of the table to complete
valence electrons. the table.
T he P er i o d i c L aw 155
Electronegativity
Valence electrons form the bonds that hold compounds Critical Thinking
together. In many compounds, the negative charge of the 13. Apply Potassium has a
valence electrons is concentrated closer to one atom than to relatively low electronegativity of
another. This uneven charge distribution helps to determine 0.8. If potassium forms a compound
with a Group 17 element, where is
the chemical properties of a compound. the electron most likely to be
Electronegativity is a measure of the ability of an atom in a located within that compound?
Period Trends
The table below and the graph on the next page show the
electronegativities of the elements. The graph shows that the
electronegativities increase when moving across a period.
The alkali metals and alkaline-earth metals are the least
electronegative elements. These elements tend to lose
electrons when forming ions. Nitrogen, oxygen, and the This table shows the electronegativi-
halogens are the most electronegative elements. These atoms ties for all of the elements. The highest
attract electrons strongly when in compounds. values are in the upper right of the
p block. The lowest values are in the
lower part of the s block.
Periodic Table of Electronegativities
Atomic
number Group 18
1
1 H 6 2
2.1 He 1
C
Group 1 Group 2 Group 13 Group 14 Group 15 Group 16 Group 17 —
Symbol
3 4 5 6 7 8 9 10
2 Li Be B C N O F Ne 2
1.0 1.5 2.5 Electronegativity
2.0 2.5 3.0 3.5 4.0 —
11 12 13 14 15 16 17 18
3 Na Mg Al Si P S Cl Ar 3
0.9 1.2 Group 3 Group 4 Group 5 Group 6 Group 7 Group 8 Group 9 Group 10 Group 11 Group 12 1.5 1.8 2.1 2.5 3.0 —
19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36
Period
Period
4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 4
0.8 1.0 1.3 1.5 1.6 1.6 1.5 1.8 1.8 1.8 1.9 1.6 1.6 1.8 2.0 2.4 2.8 3.0
37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54
5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 5
0.8 1.0 1.2 1.4 1.6 1.8 1.9 2.2 2.2 2.2 1.9 1.7 1.7 1.8 1.9 2.1 2.5 2.6
55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86
6 Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 6
0.7 0.9 1.1 1.3 1.5 1.7 1.9 2.2 2.2 2.2 2.4 1.9 1.8 1.8 1.9 2.0 2.2 2.4
87 88 89 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118
7 Fr Ra Ac Rf Db Sg Bh Hs Mt Ds Rg 7
0.7 0.9 1.1 — — — — — — — —
Lanthanide series
58 59 60 61 62 63 64 65 66 67 68 69 70 71
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
1.1 1.1 1.1 1.1 1.2 1.1 1.2 1.1 1.2 1.2 1.2 1.3 1.1 1.3
90 91 92 93 94 95 96 97 98 99 100 101 102 103
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
1.3 1.5 1.4 1.4 1.3 1.3 1.3 1.3 1.3 1.3 1.3 1.3 1.3 —
Actinide series
156 CHAPTER 5
Electronegativity vs. Atomic Number
F
4.0
Period Period Period Period Period
2 3 4 5 6
3.5
Cl Kr
3.0
Xe
Electronegativity
2.5 Rn
H
2.0
1.5
1.0 Li
Na
K Rb
Cs
0.5
0 10 20 30 40 50 60 70 80
Atomic number
PRACTICE
E. Of the elements Ga, Br, and Ca, which has the highest
electronegativity? Explain in terms of periodic trends.
T he P er i o d i c L aw 157
Periodic Properties of the
d- and f-Block Elements
The properties of d- and f-block elements vary less and with
less regularity than those of main group elements. Electrons in
the outermost d sublevel are responsible for many of the
characteristics of d-block elements.
Atomic Radii
The graph on page 145 shows that, moving from left to right
across a period, atomic radii for d-block elements generally
decrease. This is the same trend as for main-group elements.
However, the graph also shows that the radii dip to a low and
then increase slightly at the right end of the d block. This rise
results from the increased repulsion of electrons within the
Iron is often found in iron ore in
d sublevel.
compounds formed from 2+ ions.
In the sixth period, d-block elements have similar atomic
radii to the fifth period elements above them. The addition of
32 protons in the nucleus pulls the electron cloud closer to the
atom’s center despite the extra electron shell that is filled.
Ionization Energy
The periodic table on page 146 shows that in the main-group
elements, IE1 increases moving across a period and decreases
moving down a group. In the d-block, IE1 increases when
moving down a group. Because the d sublevels are incomplete
in these elements, the outermost electrons are shielded less
from the positively charged nucleus.
158 CHAPTER 5
SECTION 5.3 REVIEW
VOCABULARY
1. What is electron affinity? What signs are associated with electron affinities
and what is the significance of each sign?
REVIEW
2. State the general period and group trends among main-group elements with
respect to each of the following properties:
a. atomic radii
c. electronegativity
3. For each main-group element, what is the relationship between its group
number and the number of valence electrons that the group members have?
Critical Thinking
4. RELATING IDEAS Graph the general trends (left to
right and top to bottom) in the second ionization
energy (IE2 ) of an element as a function of its
atomic number (Z), over the range Z = 1 – 20.
Label the minima and maxima on the graph with
the appropriate element symbol.
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Math Tutor Writing Electron Configurations
ends with an atom that has a filled set of p orbitals (with the 3s 3p 3d
exception of the first period). 4s 4p 4d 4f
SAMPLE
160 CHAPTER 5
CHAPTER 5 REVIEW
a. Stanislao Cannizzaro
b. Dmitri Mendeleev
c. Henry Moseley
7. For each group, indicate whether electrons are likely to be lost or gained to
form compounds and give the typical number of electrons involved.
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8. Write the noble-gas notation for the electron configuration of each of the
following elements, and indicate the period in which each belongs.
9. Without looking at the periodic table, identify the period, block, and group
in which the elements with the following electron configurations are located.
b. [Kr]4d10
5s2 5p2 Period: Block: Group:
c. [Xe]4f14
5d10
6s2 6p5 Period: Block: Group:
10. Based on the information given below, give the group, period, block, and
identity of each element described.
11. Without looking at the periodic table, write the expected outer electron
configuration for each of the following elements.
a. [Ne]3s2 3p 1
b. [Ar]3d10
4s2 4p6
c. [Kr]4d10
5s 1
162 CHAPTER 5
13. Of cesium, Cs, hafnium, Hf, and gold, Au, which element has the smallest
atomic radius? Explain your answer in terms of trends in the periodic table.
14. a. What is the difference between the first, second, and third ionization
energies of an atom?
15. Without looking at the electron affinity table, arrange the following
elements in order of decreasing electron affinities: C, O, Li, Na, Rb, and F.
16. a. Without looking at the ionization energy table, arrange these elements
in order of decreasing first ionization energies: Li, O, C, K, Ne, and F.
b. Which of the elements listed in (a) would you expect to have the
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