Periodic trends
Vivian Hensley-Noe
Ck12 Science
Say Thanks to the Authors
Click http://www.ck12.org/saythanks
(No sign in required)
www.ck12.org
To access a customizable version of this book, as well as other
interactive content, visit www.ck12.org
AUTHORS
Vivian Hensley-Noe
Ck12 Science
CK-12 Foundation is a non-profit organization with a mission to
reduce the cost of textbook materials for the K-12 market both
in the U.S. and worldwide. Using an open-content, web-based
collaborative model termed the FlexBook®, CK-12 intends to
pioneer the generation and distribution of high-quality educational
content that will serve both as core text as well as provide an
adaptive environment for learning, powered through the FlexBook
Platform®.
Copyright © 2014 CK-12 Foundation, www.ck12.org
The names “CK-12” and “CK12” and associated logos and the
terms “FlexBook®” and “FlexBook Platform®” (collectively
“CK-12 Marks”) are trademarks and service marks of CK-12
Foundation and are protected by federal, state, and international
laws.
Any form of reproduction of this book in any format or medium,
in whole or in sections must include the referral attribution link
http://www.ck12.org/saythanks (placed in a visible location) in
addition to the following terms.
Except as otherwise noted, all CK-12 Content (including CK-12
Curriculum Material) is made available to Users in accordance
with the Creative Commons Attribution-Non-Commercial 3.0
Unported (CC BY-NC 3.0) License (http://creativecommons.org/
licenses/by-nc/3.0/), as amended and updated by Creative Commons from time to time (the “CC License”), which is incorporated
herein by this reference.
Complete terms can be found at http://www.ck12.org/terms.
Printed: March 27, 2014
iii
Contents
www.ck12.org
Contents
1
Periodic Trends: Atomic Radius
1
2
Periodic Trends in Ionization Energy
5
3
Electron Affinity
11
4
Periodic Trends: Electronegativity
14
iv
www.ck12.org
C ONCEPT
Concept 1. Periodic Trends: Atomic Radius
1
Periodic Trends: Atomic
Radius
• Define atomic radius.
• Describe how the atomic changes within a period.
• Describe how the atomic radius changes within a group.
FIGURE 1.1
crowd at an outdoor event
Events draw large numbers of people to them. Evan an outdoor event can fill up so that there is no room for more
people. The crowd capacity depends on the amount of space in the venue, and the amount of space depends on
the size of the objects filling it. We can get more people into a given space than we can elephants, because the
elephants are larger than people. We can get more squirrels into that same space than we can people for the same
reason. Knowing the sizes of objects we are dealing with can be important in deciding how much space is needed.
Atomic Radius
The size of atoms is important when trying to explain the behavior of atoms or compounds. One of the ways we
can express the size of atoms is with the atomic radius . This data helps us understand why some molecules fit
together and why other molecules have parts that get too crowded under certain conditions.
The size of an atom is defined by the edge of its orbital. However, orbital boundaries are fuzzy and in fact are
variable under different conditions. In order to standardize the measurement of atomic radii, the distance between
the nuclei of two identical atoms bonded together is measured. The atomic radius is defined as one-half the distance
between the nuclei of identical atoms that are bonded together.
Atomic radii have been measured for elements. The units for atomic radii are picometers, equal to 10−12 meters .
As an example, the internuclear distance between the two hydrogen atoms in an H2 molecule is measured to be 74
pm. Therefore, the atomic radius of a hydrogen atom is 74
2 = 37 pm .
Periodic Trend
The atomic radius of atoms generally decreases from left to right across a period. There are some small exceptions,
such as the oxygen radius being slightly greater than the nitrogen radius. Within a period, protons are added to the
1
www.ck12.org
FIGURE 1.2
The atomic radius (r) of an atom can be
defined as one half the distance (d) between two nuclei in a diatomic molecule.
FIGURE 1.3
Atomic radii of the representative elements measured in picometers.
nucleus as electrons are being added to the same principal energy level. These electrons are gradually pulled closer
to the nucleus because of its increased positive charge. Since the force of attraction between nuclei and electrons
increases, the size of the atoms decreases. The effect lessens as one moves further to the right in a period because of
electron-electron repulsions that would otherwise cause the atom’s size to increase.
2
www.ck12.org
Concept 1. Periodic Trends: Atomic Radius
Group Trend
The atomic radius of atoms generally increases from top to bottom within a group. As the atomic number increases
down a group, there is again an increase in the positive nuclear charge. However, there is also an increase in the
number of occupied principle energy levels. Higher principal energy levels consist of orbitals which are larger in
size than the orbitals from lower energy levels. The effect of the greater number of principal energy levels outweighs
the increase in nuclear charge and so atomic radius increases down a group.
FIGURE 1.4
a graph of atomic radius plotted versus
atomic number. Each successive period
is shown in a different color.
As the
atomic number increases within a period,
the atomic radius decreases.
Summary
• Atomic radius is determined as the distance between the nuclei of two identical atoms bonded together.
• The atomic radius of atoms generally decreases from left to right across a period.
• The atomic radius of atoms generally increases from top to bottom within a group.
Practice
Use the link below to answer the following questions:
http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Descriptive_Chemistry/Periodic_Table_of_the_Elements/Atomic
_Radi
1. What influences the atomic size of an atom?
2. What is a covalent radius?
3. What is an ionic radius?
Review
1. Define “atomic radius”
2. What are the units for measurement of atomic radius?
3. How does the atomic radius change across a period?
3
www.ck12.org
4. How does atomic radius change from top to bottom within a group?
5. Explain why the atomic radius of hydrogen is so much smaller that the atomic radius for potassium.
• atomic radius: The atomic radius is defined as one-half the distance between the nuclei of identical atoms that
are bonded together.
References
1.
2.
3.
4.
4
CK-12 Foundation.
CK-12 Foundation.
CK-12 Foundation.
CK-12 Foundation.
.
.
.
.
CCSA
CCSA
CCSA
CCSA
www.ck12.org
Concept 2. Periodic Trends in Ionization Energy
C ONCEPT
2
Periodic Trends in
Ionization Energy
Lesson Objectives
The student will:
• define ionization energy.
• describe the trends that exist in the periodic table for ionization energy.
• use the general trends to predict the relative ionization energies of atoms.
Introduction
Atoms are capable of forming ions by either losing or gaining electrons. Since the electrons are attracted to the
positively charged nucleus, energy is needed to pull the electron away from the nucleus. In this lesson, we will gain
an understanding of the energy required to remove an electron and recognize its trend on the periodic table.
Vocabulary
• effective nuclear charge
• ionization energy
Ionization Energy Defined
Consider lithium, which has an electron configuration of 1s2 2s1 and has one electron in its outermost energy level.
In order to remove this electron, energy must be added to the system. Look at the equation below:
energy +
−
Li(g) → Li+
(g) + e
1s2 2s1
1s2
With the addition of energy, a lithium atom can lose one electron and form a lithium ion. This energy is known as
the ionization energy. The ionization energy is the energy required to remove the most loosely held electron from
a gaseous atom or ion. The higher the value of the ionization energy, the harder it is to remove that electron. In
the equation above, the subscript “g” indicates that the element is in the form of a gas. The definition for ionization
energy specifies “in the gaseous phase” because when the atom or ion is in the liquid or solid phases, other factors
are involved. The general equation for the ionization energy is as follows.
− IE
A(g) + energy → A+
1
(g) + e
If a second electron is to be removed from an atom, the general equation for the ionization energy is as follows:
5
www.ck12.org
2+
− IE
A+
2
(g) + energy → A(g) + e
After the first electron is removed, there are a greater number of protons than electrons. As a result, when a second
electron is being removed, the energy required for the second ionization (IE2 ) will be greater than the energy required
for the first ionization (IE1 ) . In other words, IE1 < IE2 < IE3 < IE4 .
Group and Period Trends in Ionization Energy
We can see a trend when we look at the ionization energies for the elements in period 2. Table 2.1 summarizes the
electron configuration and the ionization energies for the elements in the second period.
TABLE 2.1: First Ionization Energies for Period 2 Main Group Elements
Element
Lithium (Li)
Beryllium (Be)
Boron (B)
Carbon (C)
Nitrogen (N)
Oxygen (O)
Fluorine (F)
Electron Configuration
[He]2s1
[He]2s2
[He]2s2 2p1
[He]2s2 2p2
[He]2s2 2p3
[He]2s2 2p4
[He]2s2 2p5
First Ionization Energy, IE1
520 kJ/mol
899 kJ/mol
801 kJ/mol
1086 kJ/mol
1400 kJ/mol
1314 kJ/mol
1680 kJ/mol
We can see that as we move across the period from left to right, in general the ionization energy increases. At the
beginning of the period with the alkali metals and the alkaline earth metals, losing one or two electrons allows these
atoms to become ions.
energy +
energy +
Li(g)
[He]2s1
→
Li+
(g)
[He]
+ e−
−
Mg(g)
→ Mg2+
(g) + 2e
[Ne]3s1 2
[Ne]
As we move across the period, the atoms become smaller, which causes the nucleus to have greater attraction for the
valence electrons. Therefore, the electrons are more difficult to remove.
A similar trend can be seen for the elements within a family. Table 2.2 shows the electron configuration and the
first ionization energies (IE1 ) for some of the elements in the first group, the alkali metals.
TABLE 2.2: First Ionization Energies for Some Period 1 Elements
Element
Lithium (Li)
Sodium (Na)
Potassium (K)
Electron Configuration
[He]2s1
[Ne]3s1
[Ar]4s1
First Ionization Energy, IE1
520 kJ/mol
495.5 kJ/mol
418.7 kJ/mol
By comparing the electron configurations of lithium to potassium, we know that the valence electron is further away
from the nucleus. We know this because the n value is larger, which means that the energy level holding the valence
electron is larger. It is easier to remove the most loosely held electron when the electron is further away from the
nucleus, because the attractive pull between the nucleus and an electron decreases as the distance between the two
increases. Therefore, IE1 for potassium is less than IE1 for lithium.
6
www.ck12.org
Concept 2. Periodic Trends in Ionization Energy
Why does the ionization energy increase going across a period? It has to do with two factors. One factor is that the
atomic size decreases. The second factor is that the effective nuclear charge increases. The effective nuclear charge
is the charge experienced by a specific electron within an atom. Recall that the nuclear charge was used to describe
why the atomic size decreased going across a period. Table 2.3 shows the effective nuclear charge along with the
ionization energy for the elements in period 2.
TABLE 2.3: Effective Nuclear Charge for Period 2 Main Group Elements
Element
Electron Configuration
Number
Protons
Lithium (Li)
Beryllium (Be)
Boron (B)
Carbon (C)
Nitrogen (N)
Oxygen (O)
Fluorine (F)
[He]2s1
[He]2s2
[He]2s2 2p1
[He]2s2 2p2
[He]2s2 2p3
[He]2s2 2p4
[He]2s2 2p5
3
4
5
6
7
8
9
of
Number
of
Core Electrons
2
2
2
2
2
2
2
Effective
Nuclear
Charge
1
2
3
4
5
6
7
Ionization Energy
520 kJ/mol
899 kJ/mol
801 kJ/mol
1086 kJ/mol
1400 kJ/mol
1314 kJ/mol
1680 kJ/mol
The electrons that are shielding the nuclear charge are the core electrons, which are the 1s2 electrons for period 2.
The effective nuclear charge is approximately the difference between the total nuclear charge and the number of core
electrons. Notice that as the effective nuclear charge increases, the ionization energy also increases. Overall, the
general trend for ionization energy is summarized in the diagram below.
Example:
What would be the effective nuclear charge for chlorine? Would you predict the ionization energy to be higher or
lower than the ionization energy for fluorine?
Solution:
Chlorine has the electron configuration: Cl = [Ne]3s2 3p5 . The effective nuclear charge is 7, which is the same as
the nuclear charge for fluorine. Predicting the ionization energy with just this information would be difficult. The
atomic size, however, is larger for chlorine than it is for fluorine because chlorine has three energy levels (chlorine
is in period 3). Now we can conclude that the ionization energy for chlorine should be lower than that of fluorine
7
www.ck12.org
because the electron would be easier to pull off when it is further away from the nucleus. (Indeed, the value for the
first ionization energy of chlorine is 1251 kJ/mol , compared to 1680 kJ/mol for fluorine.)
A few anomalies exist with respect to the ionization energy trends. Going across a period, there are two ways in
which the ionization energy may be affected by the electron configuration. When we look at period 3, we can see
that there is an anomaly as we move from the 3s sublevels to the 3p sublevel. The table below shows the electron
configurations and first ionization energy for the main group elements in period 3.
In the table, we see that when we compare magnesium to aluminum, the IE1 decreases instead of increases. Why is
this? Magnesium has its outermost electrons in the 3s sub-level. The aluminum atom has its outermost electron in
the 3p sublevel. Since p electrons have just slightly more energy than s electrons, it takes a little less energy to
remove that electron from aluminum. One other factor is that the electrons in 3s2 shield the electron in 3p1 . These
two factors allow the IE1 for aluminum to be less than IE1 for magnesium.
When we look again at the table, we can see that the ionization energy for nitrogen also does not follow the general
trend.
While nitrogen has one electron occupying each of the three p orbitals in the second sub-level, oxygen has an
additional electron in one of the three 2p orbitals. The presence of two electrons in an orbital lead to greater
8
www.ck12.org
Concept 2. Periodic Trends in Ionization Energy
electron-electron repulsion experienced by these 2p electrons, which lowers the amount of energy needed to remove
one of these electrons. Therefore, IE1 for oxygen is less than that for nitrogen.
This video discusses the ionization energy trends in the periodic table (1c) : http://www.youtube.com/watch?v=
xE9YOBXdTSo (9:25).
MEDIA
Click image to the left for more content.
Lesson Summary
• Ionization energy is the energy required to remove the most loosely held electron from a gaseous atom or ion.
Ionization energy generally increases across a period and decreases down a group.
• Once one electron has been removed, a second electron can be removed, but IE1 < IE2 . If a third electron is
removed, IE1 < IE2 < IE3 , and so on.
• The effective nuclear charge is the charge of the nucleus felt by the valence electrons.
• The effective nuclear charge and the atomic size help explain the trend of ionization energy. Going down a
group, the atomic size gets larger and the electrons can be more readily removed. Therefore, ionization energy
decreases down a group. Going across a period, both the effective nuclear charge and the ionization energy
increases, because the electrons are harder to remove.
Review Questions
1. Define ionization energy and write the general ionization equation.
2. Which of the following would have the largest ionization energy?
a.
b.
c.
d.
Na
Al
H
He
3. Which of the following would have the smallest ionization energy?
a.
b.
c.
d.
K
P
S
Ca
4. Place the following elements in order of increasing ionization energy: Na, O, Mg, Ne, K.
5. Place the following elements in order of decreasing ionization energy: N, Si, P, Mg, He.
6. Using experimental data, the first ionization energy for an element was found be 600 kJ/mol . The second
ionization energy was found to be 1800 kJ/mol . The third, fourth, and fifth ionization energies were found to
be, respectively, 2700 kJ/mol , 11, 600 kJ/mol , and 15, 000 kJ/mol . To which family of elements does this
element belong? Explain.
7. Using electron configurations and your understanding of ionization energy, which would you predict to have
a higher second ionization energy: Na or Mg?
9
www.ck12.org
8. Comparing the first ionization energies of Ca and Mg,
a.
b.
c.
d.
e.
calcium has a higher ionization energy because its radius is smaller.
magnesium has a higher ionization energy because its radius is smaller.
calcium has a higher ionization energy because it outermost sub-energy level is full.
magnesium has a higher ionization energy because it outermost sub-energy level is full.
they have the same ionization energy because they have the same number of valence electrons.
9. Comparing the first ionization energies of Be and B,
a.
b.
c.
d.
e.
10
beryllium has a higher ionization energy because its radius is smaller.
boron has a higher ionization energy because its radius is smaller.
beryllium has a higher ionization energy because it outermost sub-energy level is full.
boron has a higher ionization energy because it outermost sub-energy level is full.
they have the same ionization energy because boron only has one extra valence electron.
www.ck12.org
C ONCEPT
Concept 3. Electron Affinity
3
Electron Affinity
• Define electron affinity.
• Describe trends in electron affinity in the periodic table.
Do you tend to overpack before going on trips?
Packing for a trip can be very challenging. What do you take with you? Where will you be going and what will
you need? We usually pack too much (like the suitcase above) and then find it hard to close the suitcase. When the
suitcase is over-full, there is stress on the system and forces pushing the suitcase open. When electrons are added
to an atom, the increased negative charge puts stress on the electrons already there, causing energy to be released.
When electrons are removed from an atom, that process requires energy to pull the electron away from the nucleus. Addition of an electron releases energy from the process.
Electron Affinity
In most cases, the formation of an anion by the addition of an electron to a neutral atom releases energy. This can be
shown for the chloride ion formation below:
11
www.ck12.org
Cl + e
−
→ Cl
−
+ energy
The energy change that occurs when a neutral atom gains an electron is called its electron affinity . When energy
is released in a chemical reaction or process, that energy is expressed as a negative number. The figure below
shows electron affinities in kJ/mole for the representative elements. Electron affinities are measured on atoms in the
gaseous state and are very difficult to measure accurately.
FIGURE 3.1
Electron affinities (in kJ/mol) of the first
five periods of the representative elements.
Electron affinities are negative
numbers because energy is released.
The elements of the halogen group (Group 17) gain electrons most readily, as can be seen from their large negative
electron affinities. This means that more energy is released in the formation of a halide ion than for the anions of any
other elements. Considering electron configuration, it is easy to see why. The outer configuration of all halogens is
ns 2 np 5 . The addition of one more electron gives the halide ions the same electron configuration as a noble
gas, which we have seen is particularly stable.
Period and group trends for electron affinities are not nearly as regular as for ionization energy. In general, electron
affinities increase (become more negative) from left to right across a period and decrease (become less negative)
from top to bottom down a group. However, there are many exceptions, owing in part to inherent difficulties in
accurately measuring electron affinities.
Summary
•
•
•
•
Electron affinity is a measure of the energy released when an extra electron is added to an atom.
Electron affinities are measured in the gaseous state.
In general, electron affinities become more negative as we move from left to right on the periodic table.
In general, electron affinities become less negative from top to bottom of a group.
Practice
Use the link below to answer the following questions:
http://www.chemguide.co.uk/atoms/properties/eas.html
1. Which groups (using the old Roman numeral system) of elements are primarily involved with issues of electron
affinity?
12
www.ck12.org
Concept 3. Electron Affinity
2. What does a negative energy imply?
3. Why is the electron affinity value for fluorine less than that of chlorine?
Review
1.
2.
3.
4.
Define “electron affinity."
Does addition of an electron to a neutral atom require energy or release energy?
Describe the general trend for electron affinity values moving from left to right on the periodic table.
Describe the general trend for electron affinity values moving from top to bottom in a group on the periodic
table.
5. Why is more energy released in the formation of a halide ion than with other elements?
• electron affinity: The energy change that occurs when a neutral atom gains an electron.
References
1. Image copyright Africa Studio, 2013. Suitcase . Used under license from Shutterstock.com
2. CK-12 Foundation - Christopher Auyeung. . CC-BY-NC-SA 3.0
13
www.ck12.org
C ONCEPT
4
Periodic Trends:
Electronegativity
• Define electronegativity.
• Describe trends in electronegativity in the periodic table.
FIGURE 4.1
many friends
FIGURE 4.2
few friends
Have you ever noticed how some people attract others to them? Whether it is their personality, attractiveness, or
athletic skills – something pulls people toward them, while others have a smaller group of friends and acquaintances. Atoms do the same thing. One atom may pull electrons strongly to it, while a second type of atom has
much less “pulling power.”
14
www.ck12.org
Concept 4. Periodic Trends: Electronegativity
Electronegativity
Valence electrons of both atoms are always involved when those two atoms come together to form a chemical bond.
Chemical bonds are the basis for how elements combine with one another to form compounds. When these chemical
bonds form, atoms of some elements have a greater ability to attract the valence electrons involved in the bond than
other elements.
Electronegativity is a measure of the ability of an atom to attract the electrons when the atom is part of a compound.
Electronegativity differs from electron affinity because electron affinity is the actual energy released when an atom
gains an electron. Electronegativity is not measured in energy units, but is rather a relative scale. All elements are
compared to one another, with the most electronegative element, fluorine, being assigned an electronegativity value
of 3.98. Fluorine attracts electrons better than any other element. The table below shows the electronegativity values
for elements.
FIGURE 4.3
The electronegativity scale was developed by Nobel Prize winning American
chemist Linus Pauling. The largest electronegativity (3.98) is assigned to fluorine
and all other electronegativities measurements are on a relative scale.
http://authors.ck12.org/wiki/index.php/File:IntCh-06-26-Electronegativity.png
Since metals have few valence electrons, they tend to increase their stability by losing electrons to become cations.
Consequently, the electronegativities of metals are generally low. Nonmetals have more valence electrons and
increase their stability by gaining electrons to become anions. The electronegativities of nonmetals are generally
high.
Trends
Electronegativities generally increase from left to right across a period. This is due to an increase in nuclear charge.
Alkali metals have the lowest electronegativities, while halogens have the highest. Because most noble gases do not
form compounds, they do not have electronegativities. Note that there is little variation among the transition metals.
Electronegativities generally decrease from top to bottom within a group due to the larger atomic size.
Of the main group elements, fluorine has the highest electronegativity (EN = 4.0) and cesium the lowest (EN =
0.79) . This indicates that fluorine has a high tendency to gain electrons from other elements with lower electronegativities. We can use these values to predict what happens when certain elements combine. The following video
shows this.
http://www.youtube.com/watch?v=Kj3o0XvhVqQ
When the difference between electronegativities is greater than ∼ 1.7 , then a complete exchange of electrons occurs.
Typically this exchange is between a metal and a nonmetal. For instance, sodium and chlorine will typically combine
to form a new compound and each ion becomes isoelectric with its nearest noble gas. When we compare the EN
values, we see that the electronegativity for Na is 0.93 and the value for Cl is 3.2. The absolute difference between
15
www.ck12.org
ENs is |0.93 − 3.2|= 2.27 . This value is greater than 1.7, and therefore indicates a complete electron exchange
occurs.
Summary
• Electronegativity is a measure of the ability of an atom to attract the electrons when the atom is part of a
compound.
• Electronegativity values generally increase from left to right across the periodic table.
• Electronegativities generally decrease from top to bottom of a group.
• The highest electronegativity value is for fluorine.
Practice
Use the link below to answer the following questions:
http://www.chemguide.co.uk/atoms/bonding/electroneg.html
1. What are the least electronegative elements?
2. What is a polar bond?
3. What happens if atom A in a bond has much more electronegativity that atom B ?
Review
1.
2.
3.
4.
5.
Define “electronegativity.”
How does electronegativity differ from electron affinity?
Why are the electronegativity values of metals generally low?
Describe the trend in electronegativities across the periodic table.
Describe the trends in electronegativities in a group of the periodic table.
• electronegativity
References
1. CK-12 Foundation. . CCSA
2. CK-12 Foundation. . CCSA
3. CK-12 Foundation. . CCSA
16