LESSON 6.3
Key Objectives
6.3.1 DESCRIBE trends among elements for
atomic size.
6.3 Periodic Trends
6.3.2 EXPLAIN how ions form.
6.3.3 DESCRIBE trends for first ionization
energy, ionic size, and electronegativity.
CHEMISTRY
Additional Resources
Q: How are trends in the weather similar to trends in the properties of
elements? Although the weather changes from day to day. The weather
you experience is related to your location on the globe. For example,
Florida has an average temperature that is higher than Minnesota’s.
Similarly, a rain forest receives more rain than a desert. These differences are attributable to trends in the weather. In this lesson, you will
learn how a property such as atomic size is related to the location of
an element in the periodic table.
Reading and Study Workbook, Lesson 6.3
Available Online or on Digital Media:
• Teaching Resources, Lesson 6.3 Review
• Small-Scale Chemistry Laboratory Manual, Lab 9
Trends in Atomic Size
Key Questions
What are the trends among the elements for atomic size?
What are the trends among the
elements for atomic size?
One way to think about atomic size is to look at the units that form
when atoms of the same element are joined to one another. These
units are called molecules. Figure 6.14 shows models of molecules
(molecular models) for seven nonmetals. Because the atoms in each
molecule are identical, the distance between the nuclei of these atoms
can be used to estimate the size of the atoms. This size is expressed
as an atomic radius. The atomic radius is one half of the distance
between the nuclei of two atoms of the same element when the atoms
are joined.
The distances between atoms in a molecule are extremely small.
So the atomic radius is often measured in picometers (pm). Recall that
there are one trillion, or 1012, picometers in a meter. The molecular
model of iodine in Figure 6.14 is the largest. The distance between the
nuclei in an iodine molecule is 280 pm. Because the atomic radius is
one half the distance between the nuclei, a value of 140 pm (280/2) is
In general, atomic size
assigned as the radius of the iodine atom.
increases from top to bottom within a group and decreases from
left to right across a period.
How do ions form?
Engage
&
CHEMISTRY
Y
YO
YOU
U Have students read the
opening paragraph. As a class, discuss the seasonal
trends in weather typical for your region of the
country. Ask How is knowing the trends in weather
for a specific region helpful? (Sample answer: It can
help in determining the type of plants that will grow
well in your yard.) Have students consider how
knowing trends in elemental properties might be
helpful to scientists.
Access Prior Knowledge
Ask student volunteers to summarize what they have
learned about the organization of the periodic table.
Have students compare and contrast properties and
characteristics of periods with those of groups.
What are the trends among the
elements for first ionization energy,
ionic size, and electronegativity?
Vocabulary
tBUPNJDSBEJVT
tJPO
tDBUJPO
tBOJPO
tJPOJ[BUJPOFOFSHZ
tFMFDUSPOFHBUJWJUZ
Figure 6.14 Atomic Radii
This diagram compares the atomic radii
of seven nonmetals.
Distance between nuclei
Nucleus
Hydrogen (H2)
30 pm
Atomic radius
National Science Education Standards
Y
&YOU
Fluorine (F2)
62 pm
Oxygen (O2)
66 pm
Chlorine (Cl2)
102 pm
Nitrogen (N2)
70 pm
Bromine (Br2)
120 pm
Iodine (I2)
140 pm
174 $IBQUFSt-FTTPO
A-1, A-2, B-1, B-2, B-3
Focus on ELL
1 CONTENT AND LANGUAGE Have students write the lesson title Periodic Trends in their
vocabulary notebook. Have students determine the common definitions of each word
and write them in their notebook. Have students predict the lesson content based on
these definitions and what they have learned previously about the periodic table.
2 FRONTLOAD THE LESSON Provide four sets of photos of fashion trends from the
1950s, the 1970s, the 1990s, and the current decade. Ask students to identify the
fashion trends for each decade. Explain that a trend is a pattern over time. Then ask
students if they notice any similarities between clothes in the pictures. Explain that in
fashion, and in element properties, trends repeat periodically.
3 COMPREHENSIBLE INPUT Play “The Elements” song written by Tom Lehler. Use
174
Chapter 6 • Lesson 3
this as a tool for engaging students’ interest in the content. Point out that, even
though the element names are never repeated, sections of the music are repeated.
Similarly, trends in the periodic table are repeated in each period.
Foundations for Reading
Atomic Radius vs. Atomic Number
300
Period 2
200
Period 1
Atomic radius (pm)
250
Period 4
Period 3
Figure 6.15 This graph plots
Period 5
atomic radius versus atomic
number for elements with
atomic numbers from 1 to 55.
a. Read Graphs Which
alkali metal has an atomic
radius of 238 pm?
Cs
Rb
K
Na
Sc
Li
150
b. Draw Conclusions
Cd
Based on the data for alkali
metals and noble gases, how
does atomic size change within
a group?
c. Predict Is an atom of
barium, atomic number 56,
smaller or larger than an atom
of cesium (Cs)?
Zn
Xe
100
50
0
Kr
Ar
Ne
He
10
20
30
40
50
READING STRATEGY Tell students that they will
be learning about trends related to the location of
elements in the periodic table. Guide students to
read the visuals throughout the lesson closely, as the
visuals summarize the trends described in the text.
Explain
Trends in Atomic Size
USE VISUALS Guide students’ attention to Figure
Group Trends in Atomic Size Look at the data for the alkali metals and
noble gases in Figure 6.15. The atomic radius within these groups increases as
the atomic number increases. This increase is an example of a trend.
As the atomic number increases within a group, the charge on the
nucleus increases and the number of occupied energy levels increases. These
variables affect atomic size in opposite ways. The increase in positive charge
draws electrons closer to the nucleus. The increase in the number of occupied
orbitals shields electrons in the highest occupied energy level from the attraction of protons in the nucleus. The shielding effect is greater than the effect of
the increase in nuclear charge, so the atomic size increases.
6.14. Ask What is the main reason why a scientist
cannot measure the diameter of a single atom?
(because an atom does not have a sharply defined
border) Discuss how measuring the distance
between nuclei solves this problem. (NOTE: In
Chapter 8 there are formal definitions of molecule
and diatomic molecule. The operational definition
of a molecule should be sufficient for a discussion
of atomic radii.)
Period Trends in Atomic Size Look again at Figure 6.15. With increasing atomic number, each element has one more proton and one more electron than the preceding element. Across a period, the electrons are added to
the same principal energy level. The shielding effect is constant for all the
elements in a period. The increasing nuclear charge pulls the electrons in
the highest occupied energy level closer to the nucleus, and the atomic size
decreases. Figure 6.16 summarizes the group and period trends in atomic size.
Size generally increases
a compare/contrast table for cation and anion, and
use the table to decide which type of ion an element
is likely to form.
60
Atomic number
Size generally decreases
BUILD VOCABULARY Have students‘ attention make
Figure 6.16 Trends in Atomic Size
The size of atoms tends to decrease
from left to right across a period and
increase from top to bottom within
a group.
Predict If a halogen and an alkali
metal are in the same period, which
one will have the larger radius?
See periodic trends
animated online.
ET
KIN IC
ART
5IF1FSJPEJD5BCMF 175
CRITICAL THINKING Emphasize the key roles
electrical attraction and repulsion play within atoms
and ions. Review the effects of increasing nuclear
charge and changes in the shielding effect of
electrons on the size of an atom: nuclear charge
increases within groups and across periods; the
shielding effect increases within groups, but it is
constant across periods. Have students use these
effects to describe the trends for atomic size within
a period and within groups.
USE AN ANALOGY As an analogy to positions and
trends in properties of elements in the periodic
table, use seating charts and pricing data from
local theaters or sports venues to discover trends.
Instruct students to determine patterns that relate
the position of a seat to its price. Students should
discover that variables such as distance from the
stage or field, location relative to the center of the
action, and whether the view will be obstructed,
all affect price.
Elements and the Big Bang
At the time of the Big Bang, the temperature was many billions of degrees.
Neutrons, protons, and electrons may have formed within 10–4 second after the Big
Bang, and the lightest nuclei formed within 3 minutes. Matter was in the form of
plasma, a sea of positive nuclei and negative electrons. It took an estimated 500,000
years for electrons and nuclei to cool enough to form atoms. According to the Big
Bang theory, Earth, with its wealth of chemical elements, formed from the debris of
supernova explosions.
Answers
FIGURE 6.15
a. potassium
b. It increases with increasing atomic number.
c.
smaller
FIGURE 6.16 the alkali metal
The Periodic Table
175
LESSON 6.3
Interpret(SBQIT
LESSON 6.3
Explore
Figure 6.17 Cation Formation
When a sodium atom loses an
electron, it becomes a positively
charged ion.
Lose one electron
ź1eź
Ions
Class Activity
Nucleus
11 pá
12 n0
10 eź
11 eź
Sodium ion (Naá)
Sodium atom (Na)
PURPOSE To give students practice identifying
Nucleus
11 pá
12 n0
positive and negative ions
Ions
PROCEDURE Give students a list of elements.
Ask them to locate each element in the periodic
table, and decide whether its atoms are likely to
form positive or negative ions. Have students make
a list of elements that are likely to form positive ions
and another list of elements that are likely to form
negative ions.
How do ions form?
Misconception Alert
Many students will associate the words “losing” and
“gaining” with subtraction and addition, respectively.
Make sure they understand that when an atom loses
an electron, its charge becomes more positive, rather
than more negative. Similarly, when an element gains
an electron, it becomes more negative rather than
more positive. It may be helpful to remind students
that they are adding or subtracting the total charge
of the electrons gained or lost, rather than the total
number of electrons. For example, the elemental
form has a charge of 0, and it loses a single electron,
which has a charge of −1. The charge can be
calculated as 0 – (−1) = 0 + 1 = +1. For chlorine,
which gains an electron to become negative, the
calculation would be 0 + (−1) = −1.
Figure 6.18 Anion Formation
When a chlorine atom gains an
electron, it becomes a negatively
charged ion.
Interpret Diagrams What
happens to the protons and
neutrons during this change?
Nucleus
17 pá
18 n0
Some compounds are composed of particles called ions. An ion is an atom
or group of atoms that has a positive or negative charge. An atom is electrically neutral because it has equal numbers of protons and electrons. For
example, an atom of sodium (Na) has 11 positively charged protons and
11 negatively charged electrons. The net charge on a sodium atom is zero
[(à11) à (Ź11) â 0].
Positive and negative ions form when electrons are transferred
between atoms. Atoms of metals, such as sodium, tend to form ions by losing
one or more electrons from their highest occupied energy levels. Figure 6.17
compares the atomic structure of a sodium atom and a sodium ion. In the
sodium ion, the number of electrons (10) is not equal to the number of protons (11). Because there are more positively charged protons than negatively
charged electrons, the sodium ion has a net positive charge. An ion with a
positive charge is called a cation. The charge for a cation is written as a number followed by a plus sign. If the charge is 1, the number in 1à is usually
omitted from the symbol for the ion. For example, Na1à is written as Naà.
Atoms of nonmetals, such as chlorine, tend to form ions by gaining one or
more electrons. Figure 6.18 compares the atomic structure of a chlorine atom
and a chloride ion. In a chloride ion, the number of electrons (18) is not equal
to the number of protons (17). Because there are more negatively charged
electrons than positively charged protons, the chloride ion has a net negative
charge. An ion with a negative charge is called an anion. The charge for an
anion is written as a number followed by a minus sign.
Gain one electron
á1eź
17 eź
18 eź
Chlorine atom (Cl)
Nucleus
17 pá
18 n0
Chloride ion (Clź)
176 $IBQUFSt-FTTPO
Check for Understanding
How do ions form?
Assess students’ knowledge about the formation of ions by asking them the
following questions:
a. What occurs when an atom in Group 2 becomes an ion? (It loses two electrons.)
b. Is the ion that forms from a Group 2 atom called an anion or a cation? How do
you know? (It is a cation because it has a positive charge, +2.)
ADJUST INSTRUCTION If students are confused, have them use a copy of the
periodic table as a reference as they read the Ions lesson or as they review. Review
the information provided in the Misconception Alert with students and show them
how to use the group numbers shown in red in Figure 6.9 to help them determine
the number of electrons lost or gained in groups 1A–7A. Note that determining the
charge on transition metal ions will not be covered until Chapter 20.
176
Chapter 6 • Lesson 3
Explain
What are the trends among the elements for first ionization energy?
Recall that electrons can move to higher energy levels when atoms absorb
energy. Sometimes the electron has enough energy to overcome the attraction of the protons in the nucleus. The energy required to remove an electron
from an atom is called ionization energy. This energy is measured when an
element is in its gaseous state. The energy required to remove the first electron from an atom is called the first ionization energy. The cation produced
First ionization energy tends to decrease from top to
has a 1à charge.
bottom within a group and increase from left to right across a period.
Ionization energies can help you predict what ions an element will form.
Look at the data in Table 6.1 for lithium (Li), sodium (Na), and potassium (K).
The increase in energy between the first and second ionization energies is
large. It is relatively easy to remove one electron from a Group 1A metal atom,
but it is difficult to remove a second electron. This difference indicates that
Group 1A metals tend to form ions with a 1à charge.
Trends in Ionization Energy
MAKE A CONNECTION Explain that ionization energy
is a measure of the difficulty in removing an electron
from the highest occupied energy level. Ask Why
is the first ionization energy of a nonmetal much
higher than that of an alkali metal? (Because the
nuclear charge increases from left to right across a
period and the shielding effect stays the same, it is
more difficult to remove an electron.)
Misconception Alert
Interpret Data
Table 6.1 The table
Ionization Energies of First 20 Elements (kJ/mol*)
Symbol
First
H
He (noble gas)
1312
2372
Second
5247
Li
Be
B
C
N
O
F
Ne (noble gas)
Na
Mg
Al
Si
P
S
Cl
Ar (noble gas)
K
Ca
520
899
801
1086
1402
1314
1681
2080
496
738
578
786
1012
999
1256
1520
419
590
7297
1757
2430
2352
2857
3391
3375
3963
4565
1450
1816
1577
1896
2260
2297
2665
3069
1146
Third
11,810
14,840
3659
4619
4577
5301
6045
6276
6912
7732
2744
3229
2910
3380
3850
3947
4600
4941
*An amount of matter equal to the atomic mass in grams
compares ionization energies
for elements with atomic
numbers 1 through 20.
a. Read Tables What are
the values for the first, second,
and third ionization energies
for sodium and aluminum?
b. Compare Is it easier to
remove an electron from a
sodium (Na) or aluminum (Al)
atom? From Naá or Alá? From
Na2á or Al2á?
c. Draw Conclusions
Which ion is more common—
Na3á or Al3á?
Explore
Teacher Demo
Note: The second
ionization energy is the
energy needed to remove
an electron from an ion with
a 1à charge. This produces
an ion with a 2à charge. The
third ionization energy is the
energy needed to remove an
electron from an ion with a
2à charge. This produces an
ion with a 3à charge.
The Periodic Table 177
Differentiated Instruction
L1 LESS PROFICIENT READERS Have students refer back to their KWL charts to
review their notes on atomic structure. Have them use this information to identify
three of the factors that affect ionization energy: nuclear charge, number of energy
levels, and shielding.
ELL ENGLISH LANGUAGE LEARNERS Use student volunteers to set up a tug-of-war
game that to demonstrate why the amount of energy needed to remove successive
electrons increases. Assign one team to be “protons” and the other team to be
“electrons.” Show students how the amount of force exerted by the “protons” on
the “electrons” increases as each “electron” is removed from the game, and that
the “electrons” then have to pull harder against the “protons.”
L3
Some students may be tempted to place a negative
sign on the ionization energy value. Students
may confuse the phrase “losing an electron”
with a negative energy value. Explain that the
ionization energy is the energy change associated
with removing one electron from a neutral atom.
Explain that this energy value is always a positive
number since energy must be added to the atom
system to remove an electron. (NOTE: The unit for
ionization energy is kJ/mol. The footnote in Table 6.1
supplies an operational definition of mole, which is
introduced in Lesson 10.1.)
ADVANCED STUDENTS Have students create a three-dimensional tactile or
technological model that depicts the energy needed to remove an atom from an
electron.
PURPOSE To help students understand the concepts
of effective nuclear charge and electron shielding
PURPOSE Choose four students to be “protons” and
four students to be “electrons.” Construct a lithium
“nucleus” by having three protons stand together at
the front of the room. Note that for purposes of this
demo, you are ignoring the neutrons.
• Place two electrons together at a short distance
from the nucleus to represent the 1s electrons.
• Place the third electron a bit farther away to
represent the 2s electron. You should be able
to draw a line from the nucleus through the 1s
electrons to the 2s electron.
• Point out that there are no other electrons
between the 1s electrons and the nucleus.
Explain that these electrons experience the full
impact of the 3+ charge because the third electron’s
“view” of the nucleus is partially blocked. Convey
that this means the nucleus it is shielded somewhat
from the full force of the 3+ charge.
Answers
FIGURE 6.18 nothing
TABLE 6.1
a. sodium: 496, 4565, 6912 kJ/mol; aluminum:
b.
c.
578, 1816, 2744 kJ/mol
Na; Al+; Al2+
Al3+
The Periodic Table
177
LESSON 6.3
Trends in Ionization Energy
Explain
USE VISUALS Direct students to Table 6.1,
Figure 6.19, and their copy of the periodic table.
Show students how to use Table 6.1 and the
periodic table to create ordered pairs that are then
plotted in the graph in Figure 6.19. Call out various
atomic numbers and have students estimate the
ionization energy, and vice versa. When students
are comfortable reading the graph, direct them to
answer the questions.
CRITICAL THINKING Challenge students to explain
Ne
2000
Ar
Kr
N
1500
Xe
H
1000
500
0
P
Be
Zn As
Mg
Li
Na
10
20
30
40
Cs
50
60
Atomic number
Group Trends in Ionization Energy Figure 6.19 is a graph of first ionization energy versus atomic number. Look at the data for the noble gases and
the alkali metals. In general, first ionization energy decreases from top to
bottom within a group. Recall that the atomic size increases as the atomic
number increases within a group. As the size of the atom increases, nuclear
charge has a smaller effect on the electrons in the highest occupied energy
level. Less energy is required to remove an electron from this energy level,
and the first ionization energy is lower.
Explore
Teacher Demo
Period Trends in Ionization Energy In general, the first ionization energy
of representative elements tends to increase from left to right across a period.
This trend can be explained by the nuclear charge and the shielding effect.
The nuclear charge increases across the period, but the shielding effect
remains constant. As a result, there is an increase in the attraction of the
nucleus for an electron. Thus, it takes more energy to remove an electron
from an atom. Figure 6.20 summarizes the group and period trends for first
ionization energy.
PURPOSE Students observe the relative reactivities
of magnesium and calcium and predict relative
reactivities for other pairs of elements.
MATERIALS 20 mL 1M HCl, two 50-mL beakers,
overhead projector, 20 cm magnesium ribbon,
1 g calcium
PROCEDURE Pour 20 mL HCl into each beaker.
Set the beakers on an overhead projector. Coil the
magnesium ribbon and drop it into one beaker.
Drop 1 g calcium into the other beaker. Compare
the reaction rates in the two beakers. Point out
the positions of the two elements in the periodic
table, and relate the difference in reactivity to their
first and second ionization energies. Ask students
to predict the relative reactivities of other pairs of
elements in Groups 1A and 2A.
Cd
Rb
K
why the portion of the graph for Periods 4 and 5 is
different from the portion of the graph for Periods
2 and 3. (Periods 4 and 5 include transition metals,
whose atoms have electrons in d orbitals.)
SAFETY Wear goggles for this demo.
reveals group and period
trends for ionization energy.
a. Read Graphs Which
element in Period 2 has the
lowest first ionization energy?
In Period 3?
b. Describe What are
the group trends for first
ionization energy for noble
gases and alkali metals?
c. Predict If you drew a
graph for second ionization
energy, which element would
you have to omit? Explain.
He
Figure 6.20
Trends in First Ionization Energy
First ionization energy tends to
increase from left to right across a
period and decrease from top to
bottom within a group.
Predict Which element would have
the larger first ionization energy—
an alkali metal in Period 2 or an
alkali metal in Period 4?
Energy generally increases
Energy generally decreases
Trends in Ionization Energy
Figure 6.19 This graph
First Ionization Energy vs. Atomic Number
2500
First ionization energy (kJ/mol)
LESSON 6.3
Interpret(SBQIT
178 $IBQUFSt-FTTPO
EXPECTED OUTCOME The calcium fizzes in the HCl.
The magnesium reacts more slowly with the HCl.
Check for Understanding
What are the trends among the elements for first ionization energy
and ionic size and electronegativity?
Assess students’ knowledge about the trends among the elements for first ionization
energy by having students use arm gestures to answer the following questions.
Ask In which direction on the periodic table does the first ionization energy generally
increase in value? (Accept arm gestures from left to right, and upward.) Ask In which
direction on the periodic table does the first ionization energy generally decrease in
value? (Accept arm gestures downward and from right to left.)
ADJUST INSTRUCTION If students are having trouble answering, have them write
some of the values in Table 6.1 on the element squares in a copy of the periodic
table and examine their table for ionization patterns.
178
Chapter 6 • Lesson 3
What are the trends among the elements for ionic size?
During reactions between metals and nonmetals, metal atoms
tend to lose electrons and nonmetal atoms tend to gain electrons.
This transfer of electron has a predictable effect on the size of the
ions that form. Cations are always smaller than the atoms from
which they form. Anions are always larger than the atoms from
Ionic size tends to increase from top to
which they form.
bottom within a group. Generally, the size of cations and anions
decrease from left to right across a period.
Figure 6.21
Comparing Atomic and Ionic Sizes
This diagram compares the relative sizes of
atoms and ions for selected alkali metals
(Group 1A) and halogens (Group 7A). The
numbers are measurements of the radii
given in picometers (pm).
Group Trends in Ionic Size Figure 6.21 compares the relative
sizes of the atoms and ions for three metals in Group 1A—lithium
(Li), sodium (Na), and potassium (K). For each of these elements,
the ion is much smaller than the atom. For example, the radius of
a sodium ion (95 pm) is about half the radius of a sodium atom
(191 pm). When a sodium atom loses an electron, the attraction
between the remaining electrons and the nucleus is increased. As
a result, the electrons are drawn closer to the nucleus. Metals that
are representative elements tend to lose all their outermost electrons during ionization. Therefore, the ion has one fewer occupied
energy level.
The trend is the opposite for nonmetals, like the halogens in
Group 7A. Look at Figure 6.21, and compare the relative sizes of
the atoms and ions for fluorine (F), chlorine (Cl), and bromine
(Br). For each of these elements, the ion is much larger than the
atom. For example, the radius of a fluoride ion (133 pm) is more
than twice the radius of a fluorine atom (62 pm). As the number
of electrons increases, the attraction of the nucleus for any one
electron decreases.
Period Trends in Ionic Size Look ahead at Figure 6.23. From left
to right across a period, two trends are visible—a gradual decrease
in the size of the positive ions (cations), followed by a gradual
decrease in the size of the negative ions (anions). Figure 6.22 summarizes the group and period trends in ionic size.
Group 1A
eź
Li
Liá
60
156
eź
Naá
Na
95
191
eź
Ká
K
133
238
Group 7A
eź
F
62
eź
Cl
102
eź
Fź
133
Clź
181
Br
Brź
120
Explain
Trends in Ionic Size
MAKING CONNECTIONS Relate the periodic trends
in ionic size to those discussed earlier for atomic
size. Explain that the effective nuclear charge
experienced by an electron in the highest occupied
orbital of an atom or ion is equal to the total
nuclear charge (the number of protons) minus the
shielding effect due to electrons in lower energy
levels. Point out that the effective nuclear charge
determines the atomic and ionic radii. Explain
that as you proceed from left to right in any given
period, the principal quantum number, n, of the
highest occupied energy level remains constant, but
the effective nuclear charge increases. Therefore,
atomic and ionic radii decrease as you move to the
right in a period. Convey that in contrast, within
any group, as you proceed from top to bottom, the
effective nuclear charge remains nearly constant,
but the principal quantum number, n, increases.
Consequently, point out that atomic and ionic radii
increase from top to bottom within a group.
Explore
196
Teacher Demo
Size of anions decreases
Size generally increases
Size of cations decreases
Figure 6.22 Trends in Ionic Size
The ionic radii for cations and
anions decrease from left to right
across periods and increase from
top to bottom within groups.
PURPOSE Students observe an analogy for the effect
of adding or removing electrons from an atom.
MATERIALS washers or other small circular items,
smaller item (such as a button) to represent the
nucleus, overhead projector
PROCEDURE On the overhead projector, make a
5IF1FSJPEJD5BCMF 179
circle of washers to represent an electron cloud in
a neutral atom. The washers should be touching.
Place the “nucleus” in the center of the circle. Add
or subtract washers to mimic ion formation. With
each change, adjust the circle so that the washers
are still touching. Explain that the change in the
diameter of the circle is analogous to the change
in the effective attraction of the nuclear charge for
electrons.
Differentiated Instruction
L1 LESS PROFICIENT READERS Have students examine Table 6.2. Act out the
meaning of electronegativity. Then have students predict if there are any elements
other than noble gases that do not have a value for electronegativity. Guide students
in researching their prediction. (Predictions will vary; students will most likely find
electronegativity tables lacking values for manmade elements with atomic numbers
greater than 102.) Note these elements rarely, if ever, have been demonstrated to
form compounds because their atoms exist only momentarily due to their instability.
ELL SPECIAL NEEDS STUDENTS Provide tactile spherical models of various atoms.
Have students arrange the models on a copy of the periodic table to visually convey
the trend in ionic size.
L3 ADVANCED STUDENTS Have students research and describe the phenomenon
of the lanthanide contraction. Ask them to discuss how the lanthanide contraction
accounts for the fact that zirconium and hafnium have virtually the same atomic
radius even though hafnium is below zirconium in Group 4B of the periodic table.
Answers
FIGURE 6.19
a. lithium; sodium
b. First ionization energy decreases as atomic
c.
number increases.
Hydrogen; it has only one electron.
FIGURE 6.20 an alkali metal in Period 2
The Periodic Table
179
LESSON 6.3
Trends in Ionic Size
Explore
Periodic Trends in Ionic Radii
Purpose To use a graph to identify period and group trends
Trends in Ionic Size
Quick Lab
SKILLS FOCUS Using tables and graphs, predicting,
PREP TIME none
CLASS TIME 40 minutes
TEACHING TIPS If time is too limited for students
to make the graph, use Figure 6.23 to answer
Questions 1, 2, 4, and 5. You may want to reference
the radii diagrams in the Elements Handbook on R3,
R7, R11, R15, R21, R25, and R29.
EXPECTED OUTCOME Ionic radii increase from top
to bottom within a group. The radii of cations and
anions decrease from left to right across a period.
ANALYZE AND CONCLUDE
1.
2.
3.
4.
5.
Cations are smaller than their atoms; anions are
larger than their atoms.
Ionic radii increase from top to bottom within
a group of metals or within a group of
nonmetals.
Two portions of the curve slope down from left
to right.
The trend is similar for the periods.
The radii increase within a group because the
number of occupied energy levels increases.
The radii of cations decrease across a period
because the nuclear charge increases, the
shielding effect is constant, and the number
of electrons decreases. (The effect is smaller
with anions because the number of electrons
increases.)
FOR ENRICHMENT Have students use the graph on
page R37 to describe the periodic trend in atomic
size for transition metals. Have students examine
how the trend for transition metals compares to the
trend for representative elements.
Use the data presented in Figure 6.23 to
plot ionic radius versus atomic number.
rgraph paper
rpencil
OBJECTIVE Af
After completing this activity, students
will be able to identify periodic trends in ionic size.
drawing conclusions
Procedure
Materials
Analyze and Conclude
Ionic Radius vs.
Atomic Number
250
Ionic radius (pm)
LESSON 6.3
Quick Lab
1. Compare How does the size change when an atom forms a cation
and when an atom forms an anion?
2. Describe How do the ionic radii vary within a group of metals?
How do they vary within a group of nonmetals?
3. Describe What is the shape of a portion of the graph that corresponds to one period?
4. Compare and Contrast Is the trend across a period similar or
different for Periods 2, 3, 4, and 5?
5. Explain Propose explanations for the trends you have described for
ionic radii within groups and across periods.
200
150
100
50
0
0
10
20
30
40
50
60
Atomic number
Figure 6.23 Atomic and Ionic Radii
Atomic and ionic radii are an
indication of the relative size of atoms
and ions. The data listed are reported
in picometers (pm).
156
60
Li
Atomic radius
Metal atom
Cation
Metalloid atom
Ionic radius
Anion
Nonmetal atom
1A
8A
30
50
H
He
2A
1á
60
44
Be 2
191
95
Na1
á
K
1á
Mg2
66
148
á
99
Ca 2
Cs 1
á
á
á
112
á
51
Al 3
Ba 2
á
15
á
62
Ga3
á
41
81
á
53
Tl
3á
146
á
Ge 4
á
71
Sn 4
á
P
212
Pb 4
á
140
3ź
As 3
222
ź
S
62
á
170
74
Bi 5
á
2ź
184
Se 2
198
ź
221
ź
168
Po
70
Ne
102
Cl 1
181
ź
94
Ar
111
120
Br 1
196
ź
Kr
140
139
Te 2
1ź
133
120
137
Sb 5
ź
F
105
122
175
84
ź
62
66
O2
109
139
172
95
Si 4
7A
70
122
166
In 3
4á
6A
N3
109
141
224
134
3á
5A
77
C
143
215
Sr 2
273
169
23
197
255
Rb 1
B
160
238
133
á
4A
83
113
156
Li
3A
220
I
1ź
130
Xe
140
At
140
Rn
180 $IBQUFSt-FTTPO
Focus on ELL
4 LANGUAGE PRODUCTION Have students work in groups or pairs to complete the
lab. Review plotting data on a coordinate grid with the class. Pair students with
higher proficiency in English with students of lower proficiency.
BEGINNING: LOW/HIGH Help students create a set of ordered pairs to plot on the
graph. Use gestures to visually show how to plot an ordered pair on a coordinate grid.
INTERMEDIATE: LOW/HIGH Paraphrase the questions in the Analyze and Conclude
section. Guide students to answer one question at a time.
ADVANCED: LOW/HIGH Direct students to read the graph aloud, pointing out trends
and making predictions prior to answering the questions.
180
Chapter 6 • Lesson 3
Explore
What are the trends among the elements for electronegativity?
In Chapters 7 and 8, you will study two types of bonds that can exist in
compounds. Electrons are involved in both types of bonds. There is a property that can be used to predict the type of bond that will form during a
reaction. This property is called electronegativity. Electronegativity is the
ability of an atom of an element to attract electrons when the atom is in a
compound. Scientists use factors such as ionization energy to calculate values
for electronegativity.
Table 6.2 lists electronegativity values for representative elements in
Groups 1A through 7A. The elements are arranged in the same order as in
the periodic table. The noble gases are omitted because they do not form
many compounds. The data in Table 6.2 is expressed in Pauling units. Linus
Pauling won a Nobel Prize in Chemistry for his work on chemical bonds. He
was the first to define electronegativity.
In general, electronegativity values decrease from top to bottom
within a group. For representative elements, the values tend to increase
from left to right across a period. Metals at the far left of the periodic table
have low values. By contrast, nonmetals at the far right (excluding noble
gases) have high values. The electronegativity values among the transition
metals are not as regular.
The least electronegative element in the table is cesium, with an electronegativity value of 0.7. It has the least tendency to attract electrons. When it
reacts, it tends to lose electrons and form cations. The most electronegative
element is fluorine, with a value of 4.0. Because fluorine has such a strong
tendency to attract electrons, when it is bonded to any other element it either
attracts the shared electrons or forms an anion.
Figure 9.24, on the next page, summarizes several trends that exist
among the elements. Refer to this figure as you study the periodic trends
presented in this chapter.
Trends in Electronegativity
START A CONVERSATION Lead a class discussion
on periodic and group trends in electronegativities.
Point out that electronegativity values help chemists
predict the type of bonding that exists between
atoms in compounds. Ask Why are the noble gases
not included in a discussion on electronegativity?
(because they form very few compounds) Ask
Which element represented in Table 6.2 is the
most electronegative and which is the least
electronegative? (fluorine; cesium) Stress that
electronegativity is a calculated value rather than a
measured quantity.
APPLY CONCEPTS Explain that the values for
electronegativity are often based on values for
ionization energy and electron affinity. Explain that
ionization energy is a measure of an atom’s ability
to lose electrons and electron affinity is a measure
of an atom’s ability to gain electrons.
Misconception Alert
Students often confuse the meanings of
electronegativity and ionization energy. As a
class, brainstorm ideas for how to remember
the meanings of each term. Then have students
pictorially illustrate both concepts.
Table 6.2
Electronegativity Values for Selected Elements
H
2.1
Li
Be
B
C
N
O
F
1.0
1.5
2.0
2.5
3.0
3.5
4.0
Na
Mg
Al
Si
P
S
Cl
0.9
1.2
1.5
1.8
2.1
2.5
3.0
K
Ca
Ga
Ge
As
Se
Br
0.8
1.0
1.6
1.8
2.0
2.4
2.8
Rb
Sr
In
Sn
Sb
Te
I
0.8
1.0
1.7
1.8
1.9
2.1
2.5
Cs
Ba
Tl
Pb
Bi
0.7
0.9
1.8
1.9
1.9
The Periodic Table 181
Focus on ELL
4 ENABLE LANGUAGE PRODUCTION Have students work in small groups to
complete the Small-Scale Lab on page 184. Make sure each group has ELLs of varied
language proficiencies, so that more proficient students can help less proficient ones.
Have students work according to their proficiency level.
BEGINNING
LOW Model the procedure and have students mimic you. Show students how to
read measures from measuring tools.
HIGH Rephrase steps 1 and 3 in the procedure as several single-direction steps.
INTERMEDIATE: LOW/HIGH Restate the Analyze and Conclude questions in simpler
terms. Allow students to orally respond to the Analyze and Conclude questions.
ADVANCED: LOW/HIGH Have students paraphrase the questions in Analyze and
Conclude and You’re the Chemist and read them aloud to students with lower
English proficiencies.
The Periodic Table
181
LESSON 6.3
Trends in Electronegativity
Ionization energy increases
Electronegativity increases
Shielding increases
Nuclear charge increases
Electronegativity decreases
6.24. Point out that this diagram incorporates
information from several diagrams earlier in the
chapter. Lead a discussion as to whether students
find this or earlier diagrams more helpful, and why.
Ionic size increases
USE VISUALS Direct students’ attention to Figure
Ionization energy decreases
Atomic size increases
Explain
Nuclear charge increases
Shielding is constant
1A
8A
2A
3A
4A
5A
6A
7A
&
CHEMISTRY
Y
YO
YOU
U Trends in atomic size,
ionization energy, ionic size, and electronegativity
can be idendified with the help of the periodic table.
Evaluate
Size of cations decreases
Informal Assessment
Assign each student two elements in the same group
and have the student compare the elements in terms
of atomic radius, ionic radius, ionization energy, and
electronegativity. For successful students, repeat the
exercise with a metal and nonmetal from the same
period. Have students write general statements to
summarize the trends revealed by these comparisons.
Then, have students complete the 6.3 Lesson Check.
CHEMISTRY
Size of anions decreases
&YYOU
Q: You are familiar with using
a weather map to identify trends
in the weather. For example, certain areas are typically warmer
than other areas. What trends
in the properties of elements can
you identify with the help of the
periodic table?
Figure 6.24 Summary of Periodic Trends
Trends for atomic size, ionization energy, ionic size, and electronegativity vary
within groups and across periods. The trends that exist among these properties
can be explained by variations in atomic structure. The increase in nuclear charge
within groups and across periods explains many trends. Within groups, an increase
in the number of occupied energy levels and an increase in shielding both have a
significant effect on each trend.
Interpret Diagrams Which properties tend to decrease across a period? Which
properties tend to decrease down a group?
Reteach
S
E
NLIN
PR
OBLE
18.
19.
M
Review the terms used in Figure 6.24. Then, use the
periodic table and the terms to play a version of “I’m
thinking of . . . .” For example, choose fluorine and
say you are thinking of an element that has a very
small atomic size and a very high electronegativity. Let
students guess, and then discuss the correct answer.
Have students continue the game in small groups.
O
LESSON 6.3
Atomic size decreases
6.3
LessonCheck
Review How does atomic size change
within groups and across periods?
Explain When do ions form?
20.
Summarize How do first ionization energies vary within groups and across periods?
21.
Describe Compare the size of ions to the
size of the atoms from which they form.
22.
Review How do electronegativity values
vary within groups and across periods?
23. Explain In general, how can the periodic trends
displayed by elements be explained?
24. Sequence Arrange these elements in order of
decreasing atomic size: sulfur, chlorine, aluminum, and sodium. Does your arrangement
demonstrate a periodic trend or a group trend?
25. Identify Which element in each pair has the
larger first ionization energy?
a. sodium, potassium
b. magnesium, phosphorus
182 $IBQUFSt-FTTPO
Lesson Check Answers
18. Atomic size generally increases
within a group and decreases from
left to right across a period.
19. Ions form when electrons are
transferred between atoms.
20. First ionization energy generally
decreases within a group and
increases from left to right across a
period.
21. Anions are larger and cations are
smaller than the atoms from which
they form.
182
Chapter 6 • Lesson 3
22. Electronegativity values generally
decrease from top to bottom within
a group and increase from left to
right across a period.
23. The trends can be explained by
variations in atomic structure.
24. sodium, aluminum, sulfur, chlorine;
periodic trend
25. a. sodium b. phosphorus
YOU:
O EVERYDAY
V
MATTER
&YOU:
&
CHEMISTRY
Y
YO
YOU
U Have students look at
the photos and read about the Elements of Life.
Engage students in a conversation of ways the four
main elements of life—hydrogen, oxygen, carbon,
and nitrogen—are a part of their everyday lives.
Students should realize that everything they do,
every moment of the day, involves these elements.
Pose the following question to students: How many
different encounters have you had with these four
elements today? You may need to assist students in
the following ways:
• activities involving eating, drinking,
or performing personal hygiene tasks
• activities involved in getting to school
• activities the body carries out to live
• objects encountered in nature
Elements of Life
Like everything else in the universe, your body is made up of elements.
Your body uses these elements for different functions. Roughly
97 percent of the human body consists of just four elements: oxygen,
carbon, hydrogen, and nitrogen. The remaining 3 percent contains
about 20 other elements that are essential to life.
CIRCULATORY SYSTEM
M Iro
Iron
on and
oxygen are critical to the circulatory
c
system—the system that
hat carries
ca
arries blood
throughout the body. Iron, which is
contained in red blood
cells,
od ce
ells, helps
transport oxygen from
m the lungs to
other cells in your body.
dy. Two
Tw other
elements—copper and
cobalt—are
d co
balt—are
necessary for the formation
matio
on of red
blood cells.
Explain
NERVOUS SYSTEM Sodium
and potassium are essential
to the nervous system, in
particular the nerve cells. These
elements allow your brain to
communicate with other tissues
in your body. Other elements
that are important for proper
nervous system function include
calcium, chlorine, zinc, and
magnesium.
SKELETAL SYSTEM Your bones
and teeth—two components
of the skeletal system—are
largely comprised of calcium
and phosphorus, which
give bones and teeth their
strength. Fluorine, boron,
magnesium, and silicon
are also important
for bone growth and
for maintaining bone
strength.
START A CONVERSATION Explain to students that
about 97% of the atoms in the body are either
hydrogen, oxygen, carbon, or nitrogen. Encourage
them to think about what properties might make
these elements so special. Explain that in later
chapters they will learn about how atoms combine
by forming chemical bonds. Atoms of these four
elements are small and light. Also, the arrangement
of electrons in the atoms allows them to form
bonds in such a way that the atoms can combine
into the large, stable molecules necessary to carry
out life functions.
Take It Further
1. Describe Use the information provided on
page R1 to estimate the composition of the human
body in terms of metals, nonmetals, and metalloids.
2. Predict The elements sodium, magnesium,
potassium, and calcium are the most abundant
metals in the human body and are present as ions.
What is the charge
g of each of these ions?
3. Sequence Use Figure 6.23 to list the ions in
Question 2 from smallest to largest.
Chemistry & You 183
Extend
Connect to
PHYSIOLOGY
Point out to students that even though hydrogen,
oxygen, carbon, and nitrogen are by far the most
common atoms in organisms, they are not the only
atoms that are necessary for life. Have students
research the relative percentages of all the different
elements present in the human body and to present
their findings to the class.
Differentiated Instruction
L1 STRUGGLING STUDENTS Help students make a circle graph of the distribution of
elements in living organisms as a percentage of body weight. Write these numbers
on the board for them to use: oxygen, 61%; carbon, 23%; nitrogen, 2.6%;
hydrogen, 10%, and other elements, 3.4%. You may wish to have students use a
graphing program to make the graphs, or have them make the graph on the board
as a class project.
LPR
LESS PROFICIENT READERS Have students read the feature together with a
partner. First, have students read a paragraph independently, and then have them
discuss what they have read with their partners. Afterwards, discuss the feature
together as a class.
L3 ADVANCED STUDENTS Ask students to research and write a report, including
tables, that distinguishes between the percentages by mass and percentages by
number of atoms of each of the four main elements in the body.
Answers
FIGURE 6.24 sizes of atoms and ions; ionization
energy and electronegativity
TAKE IT FURTHER
1.
2.
3.
metals: 2%; nonmetals: 98%; metalloids: 0%
sodium: 1+; magnesium: 2+; potassium: 1+;
calcium: 2+
magnesium, sodium, calcium, potassium
Chemistry & You
183
CHEMISTRY & YOU
CHEMISTRY
Y
SMALL-SCALE LAB
Small-Scale Lab
Explore
Periodicity in Three Dimensions
Small-Scale Lab
OBJECTIVE Aft
After completing
l
this activity, students
should be able to build concrete models to reinforce
periodic trends. They should also be able to apply a
procedure to a new variable and design a model on
their own.
PREP TIME 10 minutes
Purpose
To build three-dimensional models for
periodic trends
Materials
r 96-well spot plate
r straws
r scissors
r metric ruler
r permanent fine-line
marker
CLASS TIME 40 minutes
MATERIALS 96-well spot plates, straws, scissors,
metric rulers, permanent fine-line markers
Procedure
ADVANCE PREPARATION Straws with a 1/4-inch
1. Measure the depth of a well in the spot plate by
inserting a straw into a well and holding the straw
upright as shown in the photograph. Make a mark
on the straw at the point where the straw meets the
surface of the plate. Measure the distance from the
end of the straw to the mark in centimeters (cm).
Record this distance as well depth.
2. Cut the straw to a length that is 4.0 cm plus well
depth. The straw will extend exactly 4.0 cm above
the surface of the plate.
3. Fluorine has an electronegativity value of 4.0. On
a scale of 1.0 cm equals 1.0 unit of electronegativity, the portion of the straw that extends above the
surface of the plate represents the electronegativity
value for fluorine. Using the same scale, cut straws
to represent the electronegativity values for all the
elements listed in Table 6.2. Remember to add the
well depth to the electronegativity value before cutting a straw. As you cut the straws, mark each straw
with the chemical symbol of the element that the
straw represents.
4. Arrange the straws in the spot plate in rows and
columns to match the locations of the elements in
the periodic table.
diameter fit snugly in the wells.
TEACHING TIPS Students can use colored straws to
color code groups or periods. If you do not have
spot plates, press a lump of clay the size of a golf
ball flat on a table with a block of wood. Students
can mark out a 1-cm square grid and insert the
straws in the clay.
EXPECTED OUTCOME Students produce 3-D models
for periodic trends.
ANALYZE AND CONCLUDE
1.
2.
3.
4.
5.
fluorine
Electronegativity generally increases from left to
right along a period.
Metals, which are on the left side of the
table, have lower electronegativity values than
nonmetals, which are on the right.
Electronegativity generally increases from
bottom to top within a group. Except for boron,
the rest of Group 3A shows a reverse in this
trend.
Hydrogen is placed in Group 1A based on its
electron configuration, but is classified as a
nonmetal.
Analyze and Conclude
1. Use Models Which element represented in your
model is the most electronegative?
2. Use Models Based on your model, what is the
general trend in electronegativity from left to right
across a period?
FOR ENRICHMENT Have students use the data on
page R37 to make a 3-D model of trends in atomic
size for transition metals.
YOU’RE THE CHEMIST
1.
2.
3.
Students divide the values of first ionization
energies by 300 and measure the appropriate
length of straws.
Students must determine their own scale
before they begin. Students often use two
wells to represent both ionic and atomic radii.
Other students cut a straw to a length that
represents the larger radius of an atom and
mark the straw to show the smaller radius of
the corresponding cation.
The value for xenon is similar to iodine, which
is consistent with the general trend. Based on
this value, xenon appears to have the ability to
attract electrons and form compounds.
184
Chapter 6 • Small-Scale Lab
184 $IBQUFSt4NBMM4DBMF-BC
3. Interpret Diagrams Relate the trend in electronegativity across a period to the location of metals
and nonmetals in the periodic table.
4. Use Models Based on your model, what is the
general trend in electronegativity within a group?
Are there any notable exceptions?
5. Explain Why do you think that the electronegativity value for hydrogen is so high given its location
in the periodic table?
You’re the Chemist
1. Design an Experiment Construct a similar
three-dimensional model for first ionization energies. Use the data in Table 6.1 to construct the
model. Use a scale of 1.0 cm equals 300 kJ/mol.
2. Design an Experiment Design and construct
a three-dimensional model that shows trends in
atomic and ionic radii for the elements in Groups 1A
and 7A. Devise a way to display both ionic and
atomic radii in the same model.
3. Analyze Data Xenon has an electronegativity value of 2.6. Cut and place a straw in your first
model to represent xenon. Does xenon support the
trend for electronegativity across a period? Is xenon
likely to form compounds? Explain your answers.