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Unformatted text preview: Chapter Eight PERIODIC RELATIONSHIPS AMONG THE
• Periodic Classification of the Elements
Periodic Variation in Physical Properties
Ionization Energy and Electron Affinity
Types of Elements PERIODIC CLASSIFICATION OF THE ELEMENTS
4. Discuss the basis for the arrangement of elements in the periodic table.
Write the electron configuration of an element and determine to what group of the periodic table it belongs.
Give the locations in the periodic table of the representative, transition, and inner transition elements.
Write the electron configurations of cations and anions. The Periodic Table. A modern version of the periodic table is shown inside the cover of the text. The
table organizes a vast array of information about the chemical and physical properties of the elements. In the
periodic table the elements are arranged by atomic number. The horizontal rows of elements with atomic number
increasing by one from left to right are called p eriods . The vertical columns with atomic number and atomic
mass increasing from top to bottom are called g roups or f amilies. Elements within a group have similar,
but not identical, properties. Table 8.1 shows how properties of the halogens vary gradually within a group.
Table 8.1 Properties of the Halogens
Initiated by light
Slow at 25°C
The chemical similarities are that the free elements are all diatomic molecules, they form salts with sodium
that have the same formula, that is, NaX, and they all react with hydrogen. The ease of reaction with hydrogen,
and the physical properties of melting point and boiling point all vary in a regular way as you read down in the
group. 1 45
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Element groups are often referred to collectively by their group number: Group 1A, 2A, and so on. Some
groups have acquired special names which are given below:
Group 1A elements are called a lkali metals.
Group 2A elements are called a lkaline earth metals.
Group 7A elements are called h alogens.
Group 8A elements are called n oble gases. Element Groups and Electron Configurations. The modern theory of atomic structure suggests an
explanation as to why elements in a group have similar chemical properties. First let's examine the electron
configurations of elements in any vertical column in the periodic table. Table 8.2 shows the electron
configurations of atoms of the Group 7A elements.
Table 8.2 Group 7A
1s2 2s2 2p5
[Ar]4s2 3d10 4p5
[Kr]5s2 4d10 5p5
The table shows that the electron configuration of the outermost principal energy levels of all Group 7A
elements contain seven electrons arranged ns2 np5 , where n is the principal quantum number of the outermost
energy level. Furthermore, the halogen group is not unique in this regard. The electron configuration of the
outermost electrons of elements in any group is the same. See Figure 8.2 in the text.
The chemical similarities of the Group 7A elements can be explained quite simply by their electron
configurations. All the elements in a group have the same electron configuration in their outermost principal
energy level. The electrons in the outermost energy level are called v alence electrons. The outer electrons
are the ones involved in chemical bonding and so deserve a special name. Representative Elements. The modern periodic table is arranged according to the type of subshells being
filled with electrons (Figure 8.2 textbook). Groups 1A through 7A of the periodic table include elements that
have an incompletely filled set of s or p orbitals. These are called r epresentative elements , or m ain
group elements. The groups are given numerals from one to seven, followed by the letter A, which stands for a
group of representative elements. The n oble gases have completely filled outer ns and np subshells, ns2 np6 ,
except for helium.
Electron Configurations. The chemical properties of an atom are related to the configuration of the
atom's valence electrons. For representative elements the group number gives the number of electrons in the
outermost principal energy level (valence electrons), as shown in Table 8.3.
Table 8.3 Valence Electron Configurations of Representative Elements
2 np1 ns2 np2
2 np3 ns2 np4
__________________________________________________________________________ Transition Elements. The elements in the center of the periodic table, whose groups are labeled B, are
the t ransition metals. Transition metals have incompletely filled d subshells or readily form cations having
incompletely filled d subshells. In the fourth period, either the atoms or the ions of the elements Sc through Cu
have incompletely filled 3d subshells. In the fifth period, the atoms or the ions of the elements Y to Cd have Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Periodic Relationships among the Elements / 1 47 Representative
Transition Elements Noble Gas Elements Representative
Elements incompletely filled 4d subshells. Because their d subshells are completed, the elements Zn, Cd, and Hg (Group
2B) are often not considered true transition elements.
Below the main part of the periodic table, there are two sets of 14 elements each depicted as horizontal
rows. These groups are generally known as i nner transition elements, but are also referred to as
lanthanides (atomic numbers 58–71) and a ctinides (atomic numbers 90–103). For these elements an inner f
subshell is incompletely filled. Figure 8.1 shows the positions in the periodic table of the groups of elements
discussed above. Inner Transition
Figure 8.1 General groups of elements. Electron Configurations of Cations and Anions. The ions of the representative elements have
noble gas electron configurations, ns2 np6 . The representative metals form positive ions by losing one or more
electrons until a noble gas core is achieved.
Na [Ne]3s1 → Na+ [Ne] + e–
Mg [Ne]3s2 → Mg 2+ [Ne] + 2e–
Al [Ne]3s2 3p1 → Al 3+ [Ne] + 3e–
The electron configurations of the above ions (Na+, Mg2+ , and Al3+ ) are all the same, namely 1s 2 2s2 2p6 or
[Ne]. All three have 8 electrons in the outer shell (n = 2), and 10 electrons all together. Atoms and ions that
have the same number of electrons and hence the same electron configurations are called i soelectronic
species. All three ions are isoelectronic with the noble gas neon. Representative metal ions will have electron
configurations that are isoelectronic with the previous noble gas atom
The nonmetal atoms form negative ions by acquiring electrons until a noble gas core is achieved.
S [Ne]3s2 3p4 + 2e – → S 2– 3s2 3p6 or [Ar]
Cl [Ne]3s2 3p5 + e – → C l – 3s2 3p6 or [Ar]
Sulfide and chloride ions then have net charges of –2 and –1 because the atoms needed 2e– and 1e– , respectively,
to complete an octet. Negative ions of the representaive nonmetals will have electron configurations that are
isoelectronic with the next noble gas atom. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 1 48 / Periodic Relationships among the Elements
Transition Metal Cations. When forming a cation from an atom of a transition metal, recall that
most of the transition metal atoms do not acquire a noble gas electron configuration for their cations. Atoms of
transition metal elements generally have 2 electrons in the s orbital of the highest principle energy level n, and a
partially filled d subshell of the n – 1 principle energy level (see Table 7.3 of the text). The electrons most
easily lost are those in the outermost principle energy level, the ns. Indeed many of the transition metals form
+2 ions. Loss of additional electrons from the (n – 1)d subshell yields ions with charges greater than a +2
charge. See Example 8.5, parts c and d.
EXAMPLE 8.1 Element Groups
Given the following formulas of compounds:
Sodium sulfate Na2 SO4
Magnesium oxide MgO
Aluminum chloride AlCl 3
Write the formulas of the following compounds using the periodic table:
a. Calcium oxide
b. Potassium sulfate
c. Gallium bromide
•Method of Solution
c. Notice that Ca is in the same chemical group as Mg. If one Mg ion combines with one oxide ion as in
MgO, then one Ca ion will combine with one oxide ion. Answer: The formula is CaO.
If sodium sulfate is Na2 SO4 , then potassium, being in the same chemical group as sodium, should form a
sulfate with the same ratio of potassium atoms to sulfate ions. Answer: The formula is K2 SO4 .
Gallium is in Group 3A along with Al, and bromine is a halogen, as is Cl. Therefore, since aluminum
chloride has the formula AlCl3 , gallium will combine with any three halogen atoms. Answer: The formula
of gallium bromide is GaBr3 . _______________________________________________________________________________
EXAMPLE 8.2 Electron Configurations and the Periodic Table
Referring only to the periodic table:
a. Write the electron configuration for the outermost principal energy level (the valence shell) for any Group
b. Identify the group whose outermost electrons have the configuration ns2 np1 .
•Method of Solution
a. The elements in Group 5A are representative elements and so have all inner subshells filled, and the
outermost electrons occupy s and p orbitals. The numeral 5 corresponds to five outermost electrons. The
first two occupy the ns subshell, and the next three the np subshell. Answer: The valence electron
configuration of any Group 5A element is ns 2 np3 .
b. Among the representative elements, those in the first group to the right of the transition metals, Group 3A,
have the electron configuration ns2 np1 .
EXAMPLE 8.3 Electron Configurations of Ions
Write the electron configurations of the following ions:
a. Mg 2+
b . S 2–
c . Fe 2+ d. Fe3+ Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Periodic Relationships among the Elements / 1 49
•Method of Solution
a. A Mg atom has Z = 12, so the atom has 12 electrons, but the Mg2+ ion contains only 10 electrons. Write
the configuration for Mg.
Mg 1s2 2s2 2p6 3s2 or [Ne]3s 2
Then remove the two electrons in the outermost energy level to form the 2+ ion.
Mg2+ 1s2 2s2 2p6 or [Ne]
Mg2+ ions are isoelectronic with a Ne atom. That is, they have the same number of electrons. b. A S atom has 16 electrons and has the electron configuration 1s 2 2s2 2p6 3s2 3p4 . The sulfur atom becomes
the negative S 2– ion by gaining two electrons. Adding two electrons to the outermost subshell gives
S 2– 1s 2 2s2 2p6 3s2 3p6 or [Ar] c. As before, write the configuration of the atom.
Fe [Ar]4s2 3d6
When forming a cation from an atom of a transition metal, electrons are always removed first from the ns
orbital, and then from the (n – 1)d orbital. The first two electrons lost are removed from the outermost
principal energy level (highest n value), which is the 4s orbital. This gives:
Fe2+ [Ar]3d6 d. Next, remove one more electron from Fe2+ . This gives:
Fe3+ [Ar]3d5 •Comment
The stable ions of all but a few representative elements are isoelectronic to a noble gas. Keep in mind that
most transition metals can form more than one cation, and that for the most part these ions are not
isoelectronic with the preceding noble gases.
1. What is a representative element? Give three examples of representative elements.
2. What is a transition element? Give three examples of transition elements.
3. What is an inner transition element? Give two examples of inner transition elements.
4. How many valence electrons does an atom of phosphorus have?
5. Write the outer electron configurations for:
a. the Group 3A elements. b. the Group 6A elements.
6. Write the symbol of an element with 4 valence electrons. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 1 50 / Periodic Relationships among the Elements
7. What group of elements has the outer electron configuration ns2 np5 ?
8. Write the electron configuration of scandium (Sc). Classify the element as a representative, transition, or
inner transition element. Classify the element as a metal or nonmetal.
9. Name the groups of the periodic table in which each of the following elements are found.
a. [Ar] 4s2 b. [Ar]4s2 4p3 c. [Kr]5s2 4d7
10. Which of the following species are isoelectronic with each other?
a. Kr b. Cl– c. Rb+ d. K+ e. Cu2+ f. Se2– g. Zn2+
11. Write the ground state electron configurations of the following ions:
a. K + b . O 2– c. Sr2+ d. I– e. Li+ f. Mg2+
12. Write the ground state electron configurations of the following transition metal ions.
a. Fe 2+ b . Fe 3+ c. Cu2+ d. Cu+ e. Ti2+ f. Ti4+
13. Write the ground state electron configurations of the following ions.
a. Sc 3+ b . V 5+ c . Pb 2+ d. Pb4+ PERIODIC VARIATION IN PHYSICAL PROPERTIES
3. Account for the trends in the magnitudes of the atomic radius for the elements in a period and within a
group of the periodic table.
Predict the relative sizes of atomic and ionic radii within an isoelectronic series.
Estimate certain physical properties of an element given properties of elements above and below the given
element in a group of elements. Atomic Radius. Figure 8.5 of the textbook shows that the atomic radius decreases, in general, as the
atomic number increases across any period of the periodic table. Thus, the alkali metals have the largest atoms
and the noble gases the smallest. Furthermore, within a group the atomic radius increases as the atomic number
Another trend in chemical behavior is the d iagonal relationship. This refers to similarities in some
properties between a second row element, and the third row element that is one space to the right in the table.
Two such elements are on a diagonal (Figure 8.13 in the text). For example, Li and Mg are on a diagonal, and
they have similar atomic radii. If you start at Li and go first to Be, the atomic radius decreases. Next, on moving
down to Mg, the radius will increase again. If the decrease and increase are about the same, then Li and Mg will
have approximately the same radius (155 pm for Li versus 160 pm for Mg). This relationship does not hold for
all properties of two diagonally related elements. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Periodic Relationships among the Elements / 1 51 Increasing Atomic Radius Decreasing Atomic Radius Figure 8.2 Periodic trends in the atomic radius. Ionic Radii . Because atoms and ions of the same element have different numbers of electrons, we should
expect atomic radii and ionic radii to have different values. The radii of cations are smaller than those of the
corresponding neutral atoms. This should be expected because positive ions are formed by removing one or more
electrons from the outermost shell. Since these electrons are furthest from the nucleus, their absence will make
the cation significantly smaller. In addition, loss of an electron causes a decrease in the amount of electronelectron repulsion which also causes the cation to be smaller than the neutral atom.
In contrast, the radii of anions are larger than those of the corresponding neutral atom. When an electron is
added to an atom to form an anion, there is an increase in the electron-electron repulsion. This causes electrons
to spread out as much as possible, and so anions have a larger radius than the corresponding atoms. The radii of
several metal and nonmetal ions are compared with their atomic radii in Table 8.4.
Table 8.4 Atomic and Ionic Radii of Elements in Groups 1A and 7A
We can also compare radii within an isoelectronic series. In the series in Table 8.5 the radii decrease steadily
as atomic number increases.
Table 8.5 Atomic and Ionic Radii in an Isoelectronic Series
C a 2+
_______________________________________________ Each species has 18 electrons arranged in an argon configuration, and so the electron-electron repulsion is about
the same in each ion. The reason for the decrease in radius is that the nuclear charge increases steadily within Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 1 52 / Periodic Relationships among the Elements
this series. This causes the electrons to be attracted more strongly toward the nucleus, and the ionic radius to
contract as the nuclear charge increases. Effective Nuclear Charge. The explanation of the trends in atomic size and ionization energy (to be
discussed) is based on the concept of the effective nuclear charge (Zeff). The outermost electrons in an atom do
not feel the full positive charge of the nucleus. Electrons in inner levels, lying between the nucleus and the
outermost electrons, tend to shield the outermost electrons from the nuclear charge. In a sodium atom, for
example, the inner 10 electrons (1s2 2s2 2p6 ) shield the outer 3s electron from the positive charge of 11 protons.
The Zeff experienced by the 3s electron is only about +1. The Zeff is equal to the nuclear charge Z minus the
shielding constant σ. The shielding constant is essentially equal to the number of inner shell electrons. Thus,
Zeff = Z – σ.
Atomic Radii. The Z eff affects the atomic radius in the following way. In proceeding across a period, one
proton at a time is added to the nucleus, and one electron is added to the outermost orbital. Electrons within the
same energy level do not effectively shield each other from the nucleus. Thus, for magnesium the inner 10
electrons shield the electrons in the 3s orbitals from the 12+ charge of the nucleus, resulting in an effective
nuclear charge of +2. One of the 3s electrons does not effectively shield the other 3s electron because they are in
the same orbital. Continuing from left to right across the third period, all the outermost electrons are in the n =
3 energy level, and the Zeff increases by about one unit per element.
Increasing Z eff
__________________________________________→ Na(+1) Mg(+2) Al(+3) Si(+4) P(+5) S(+6) Cl(+7) Ar(+8)
This means a greater force of attraction is experienced by the outermost electrons of atoms of elements on
the right side of the periodic table than for those on the left side. Consequently, the atoms gradually decrease in
size as we proceed from left to right across a period.
In proceeding from one element to another down a group, each successive element has its outer electrons in
a principal energy level with a larger n value. The effective nuclear charge experienced by the outermost
electrons is essentially the same for all elements within a group. The result of this is that the size of the outer
orbital is affected most by the value of n . As you will recall, the size of an orbital increases as n increases.
Therefore, atomic radius increases from top to bottom in a group. Transition Elements. Within a row of transition elements, in contrast to the representative elements,
there is only a slight decrease in atomic radius when reading from left to right, as shown in Table 8.6.
This is so because electrons are being added to an inner d subshell. In period four, for instance, the
outermost electrons occupy a 4s orbital, but each successive electron is added to an inner 3d subshell. As we read
across the period, the increasing nuclear charge is effectively shielded by the increase in the number of 3d
electrons. Thus the outer 4s electrons within a series of the transition elements experience almost a constant
Table 8.6 Atomic Radius (pm)
EXAMPLE 8.4 Trends in Atomic Radius
Which one of the following has the smallest atomic radius?
d. Mg Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Periodic Relationships among the Elements / 1 53
•Method of Solution
Atomic radii increase from top to bottom in a group, therefore Li atoms are smaller than Na atoms, and Be
atoms are smaller than atoms of Mg. Next compare Be and Li. The atomic radius decreases from left to right
within a period of the periodic table, thus Be atoms are smaller than Li atoms, as well as the other choices
EXAMPLE 8.5 Trends in Ionic Radii
In each of the following pairs, choose the ion with the largest ionic radius.
a. K+ or Na+
b. K+ or Ca 2+
c. K+ or Cl–
•Method of Solution
c. K+ and Na+ are in the same group in the periodic table. The outer electrons in K+ occupy the third principal
energy level, and those in Na+ occupy the second. Answer: K + has the greater ionic radius.
K+ and Ca2+ belong to an isoelectronic series; both have 18 electrons. Calcium ions with their greater
nuclear charge attract their electrons more strongly than K+ ions and so are smaller. Answer: K + is larger.
Again K+ and Cl– are isoelectronic species. The ion with the greater nuclear charge (K+) will be smaller.
Answer: C l – has the greater ionic radius. _______________________________________________________________________________
EXAMPLE 8.6 Predicting Properties of Elements
Predict the missing value for the following:
Density (g/cm3 )
I Density (g/cm3 )
? •Method of Solution
Within a group of the periodic table the physical properties vary in a regular fashion as you read down in the
a. From the positions of these elements in Group 2A of the periodic table, the density of Sr could be
estimated to be half-way between the values for Ca and Ba. The average of 1.55 and 3.5 g/cm3 is 2.5
g/cm3 . The observed value is 2.6 g/cm3 .
b. In this case we must extrapolate rather than interpolate as above because I is farther down in Group 7A than
Cl and Br. We could assume the change in the radius from Br to I will be the same as the change from Cl to
Br. The difference between Cl and Br is 0.015 nm. Adding 0.015 to 0.114 gives the estimated radius for I as
0.129 nm. The observed radius of I is 0.133 nm. _______________________________________________________________________________ Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 1 54 / Periodic Relationships among the Elements
14. Which of the following has the largest radius?
a. Na, Cl, Ne, S, Li
b. C, Si, Sn, Pb
15. Identify any isoelectronic pairs among the following.
Na+, Ar, Cl– , Ne, Se2–
16. Which of the following has the largest radius?
S 2– , Ar, Se2– , O 2– , Al 3+
17. Which of the following has the largest radius?
Na+, Mg2+ , Al 3+ , Ne
18. Which species is larger: a. Co2+ or Co 3+ b . S or S 2– 19. Estimate the melting point of Br2 (s) given that the melting points of Cl2 (s) and I2 (s) are
–101°C and 114°C, respectively. IONIZATION ENERGY AND ELECTRON AFFINITY
4. Describe the terms ionization energy and electron affinity.
Describe the trends in magnitudes of the ionization energy and electron affinity while reading across a period
and down a group in the periodic table.
Know what types of elements have low ionization energies and high electron affinities.
Account for the observed trends in ionization energy within a period and within a group of the periodic
table. Ionization Energy. The minimum energy required to remove an electron from the ground state of an
atom in the gas phase is its ionization energy. The magnitude of the ionization energy is a measure of how
strongly the outermost electron is held by an atom. The greater the ionization energy, the stronger the electron
For a many-electron atom, the energies required to remove a second and a third electron are called the second
ionization energy and the third ionization energy, respectively. The third ionization energy is always greater than
the second, which in turn is always greater than the first ionization energy. The explanation of this trend is
related to the electron repulsion among the remaining electrons. The first electron removed comes from a neutral
atom. With an electron missing the repulsion among the remaining electrons decreases, and they move closer to
the nucleus. This means that to remove the second electron requires a higher ionization energy. The third
ionization energy is higher yet due to a further reduction in electron repulsion.
Table 8.11 in the textbook shows that the ionization energy increases when moving from left to right
across a period. Thus, the alkali metals have the lowest ionization energy and the noble gases the highest.
Within a group the ionization energy decreases as atomic number increases (Figure 8.3). Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Periodic Relationships among the Elements / 1 55 Decreasing Ionization Energy Increasing Ionization Energy Figure 8.3 Periodic trends in ionization energy. Ionization Energy Trends. The importance of the ionization energy is that it correlates closely with
electron configuration and chemical properties. When considering ionization energy, it helps to recall that the
first electron removed is the one farthest from the nucleus. Within a group of elements the ionization energy
decreases with increasing atomic number. This trend is explained in the following way. When reading down a
group of elements the effective nuclear charge is essentially constant. However, the outermost electron resides in
increasingly higher energy levels. Thus, as atomic number increases within a group, the outermost electron is
held more weakly and the ionization energy decreases.
The trend within a period should also be related to the effective nuclear charge. In moving across the
periodic table from left to right, the size of the atoms decreases due to the increase in effective nuclear charge.
The outer electrons become more tightly held as we move from left to right, and the ionization energy must
One characteristic of metals is the relative ease with which electrons can be removed from their atoms.
Thus, the lower the ionization energy, the more metallic the element. The metals are located on the left hand
side of the periodic table. Also, elements in all groups exhibit increased metallic behavior with increased atomic
Electron Affinity. Another property of atoms is one that relates to their ease in forming negative ions.
The e lectron affinity of an element is the energy given off when a gaseous atom gains an additional electron
to form a negatively charged ion. A fluorine atom, for example, gives off energy when it gains an electron to
form a fluoride ion.
F(g) + e – → F – (g) ∆ H = –328 kJ/mol The electron affinity of fluorine is +328 kJ/mol because that is the amount of energy given off.
F(g) + e – → F – (g) + 328 kJ/mol
The greater the attraction of an atom for an electron, the greater the electron affinity value. The more energy
released, the more stable the ion is compared to the atom.
Table 8.4 of the text shows the electron affinity values of elements arranged according to their positions in
the periodic table. Figure 8.12 (textbook) is a plot of electron affinity values for the first 20 elements. A clear
trend in electron affinity values of elements within a period is not evident. Some trends in electron affinity can
4. Back The electron affinities of the Group 6A and 7A elements are much greater than those of other elements.
The electron affinities of the noble gases are slightly negative. Therefore the anions of these gases would be
The Group 1A metals have small positive electron affinities. The Group 2A metals have very small
positive values and some negative electron affinities.
The electron affinities of metals are generally lower than those of nonmetals. Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 1 56 / Periodic Relationships among the Elements
Electron Affinity Trend. The trend in electron affinities is also related to effective nuclear charge, and to
the energy of the orbital that the electron will enter. The Group 6A and 7A elements have the highest electron
affinities. These elements have high Zeff and the electrons enter the valence shell. The noble gases with even
higher Zeff have no affinity for an additional electron. With filled s and p subshells, noble gas atoms have no
tendency to add an electron because the next available orbital is in the next higher energy level beyond the
EXAMPLE 8.7 Trends in Ionization Energy
Which one of the following has the highest ionization energy?
•Method of Solution
Ionization energy increases from left to right within a period. Thus, the value for Cl is greater than for S, and
the value for Br is greater than for K. In comparing Cl and Br, Cl has the higher ionization energy value because
ionization energy decreases from top to bottom within a group.
EXAMPLE 8.8 Ionization Energy
Why does sodium have a lower first ionization energy than lithium?
•Method of Solution
When a lithium atom is ionized the electron most easily removed comes from a 2s orbital, whereas in sodium
the electron most easily removed comes form a 3s orbital. Because of shielding effects, the effective nuclear
charge Zeff experienced by these electrons is about +1 in both atoms. The main reason for a difference in
ionization energy in these atoms is the larger distance of separation between the electron and the nucleus in the
case of Na atoms. As the principal quantum number n increases, so does the average distance of the electron
from the nucleus. Consequently, the electron becomes easier to remove as you read down a group. Answer: Na
will have a lower ionization energy than lithium.
EXAMPLE 8.9 Ionization Energies
The first, second, and third ionization energies for calcium are:
I1 = 590 kJ/mol
I2 = 1145 kJ/mol
I3 = 4900 kJ/mol
Explain why so much more energy is required to remove the third electron from Ca, as compared to removal of
the first and second electrons.
Answer: The successive ionization energy values follow the usual trend I3 > I2 > I1 . However, we must
explain the very large difference between I2 and I3 . The first two electrons are valence electrons and are removed
from the 3s orbital. But the third electron must be removed from the inner 2p subshell. The n = 2 principal
energy level lies much closer to the nucleus than the n = 3 energy level; therefore, its electrons are held much
more strongly. Consequently, I3 >> I2 .
_______________________________________________________________________________ Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Periodic Relationships among the Elements / 1 57
20. What group of elements in the periodic table have:
a. the highest ionization energies b. the lowest ionization energies
21. Based on periodic trends, which one of the following elements has the lowest ionization energy?
K, S, Se, Li, Br
22. Based on periodic trends, which one of the following elements has the largest ionization energy?
Cl, K, S, Se, Br
23. For silicon atoms, which ionization energy value (I) will show an exceptionally large increase over the
preceding ionization energy value?
I2 I 3 I 4 I 5 I 6
24. Define electron affinity. What group of elements have negative electron affinities?
25. Which one of the following elements has both a low ionization energy and a positive electron affinity?
K, Ne, Br, Fe, N
26. Which of the following atoms has both a high ionization energy and a large negative electron affinity?
K, Ne, Br, Fe, N TYPES OF ELEMENTS
2. Describe typical properties of metals, nonmetals, and metalloids.
Identify the following chemical groups in the periodic table: alkali metals, alkaline earths, halogens, noble
gases, and coinage metals. Metals. All elements are classified in three broad groups according to their chemical and physical
properties. These are the metals, nonmetals, and metalloids. Within this classification, certain groups of
elements are known by common names, such as the alkali metals, alkaline earth metals, coinage metals,
halogens, and noble gases. In the section below, the three broad groups and their characteristics are described first
followed by the groups with common names and their characteristics.
Most of the elements are metals. In general, they appear in the left and center of the periodic table. All
metallic elements except mercury are solids at 25°C. They have a lustrous appearance, are good conductors of
heat and electricity, can be hammered or rolled into sheets (a property referred to as malleability), and can be
drawn into wires (a property referred to as ductility). Metals generally have high densities and high melting
Metals have low ionization energies, low electron affinities, and relatively large atomic radii. All metallic
elements combine with nonmetals such as oxygen and chlorine to form salts. The most reactive metals are at
the left of the periodic table. The transition metals are less reactive than the Group 1A and 2A metals. Elements
within a chemical group are more metallic as atomic mass increases. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 1 58 / Periodic Relationships among the Elements
Nonmetals. Eighteen of the elements are nonmetals. These elements are on the right side of the periodic
table. At 25°C, eleven are gases, one (bromine) is a liquid, and the rest are brittle solids. Typically, their
densities and melting points are low.
The atoms of nonmetal elements have high ionization energies and high electron affinities. Nonmetals
combine with metals to form ionic compounds, and with other nonmetal elements to form molecular
compounds. The nonmetals, except hydrogen, are located on the upper right hand side of the periodic table.
Recall that hydrogen is a gas and exists as diatomic molecules, as do a number of nonmetals.
Metalloids. Several elements have some properties that are characteristic of metals and some that are like
those of nonmetals. These elements are called metalloids. Many periodic tables show a zig-zag line separating
the metals from nonmetals. The elements that border this line on both sides are metalloids (except for Al, which
is a metal). They include boron, silicon, germanium, arsenic, antimony, and tellurium. The metalloids have
ionization energies and electron affinity values intermediate between metals and nonmetals.
Alkali Metals. Certain groups of elements are known by common names (Figure 8.4). Some of the
characteristics of the alkali metals, alkaline earth metals, coinage metals, halogens, and noble gases are described
The elements in Group 1A, with the exception of hydrogen, are called alkali metals. The ns electron in the
highest principal energy level is well shielded from the nucleus and is easily lost. Their low ionization energies
make the alkali metals the most active family of metals. These elements are found in nature as +1 ions in
chemical combination with nonmetal ions and polyatomic ions. The densities of these elements are low, in part,
because of their large radii. Li, Na, and K are even less dense than water. Noble Gases Transition Metals Halogens Coinage Metals Alkaline Earth Metals Alkali Metals H Lanthanide Series
Figure 8.4 Common names of several groups of elements. Alkaline Earths. Group 2A elements are called alkaline earth metals. The outermost electron
configuration is ns 2 . Although these two are not as easily lost as the ns1 electron in alkali metal atoms, they
are readily given up. Alkaline earth metals exist in nature as +2 ions in chemical combination with negative
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Halogens. Group 7A elements are called the halogens. None of these elements is ever found free in
nature. All have high electron affinities and so tend to acquire one electron to form –1 ions. Ionic compounds
containing these anions are salts. Indeed, the name halogen means "salt-former." These nonmetals exist in the
elemental form as diatomic molecules. At 25˚C, F2 and Cl2 are gases, Br2 is a liquid, and I2 is a solid.
Noble Gases. Group 8A elements make up the noble gas elements. the term noble here means
nonreactive. He, Ne, and Ar do not form any chemical compounds. Since 1960, a number of Kr and Xe
compounds have been synthesized. The chemical inactivity of the noble gases is the result of the ns 2 np6
electron configuration of the outermost electrons in their atoms. The noble gas elements have negative electron
affinities and high ionization energies.
Coinage Metals. The elements of Group 1B, copper, silver, and gold, are generally nonreactive and are
called the noble metals or coinage metals. These metals are excellent conductors of heat and electricity.
Atoms of the coinage metals have one electron in the outer s subshell and 10 electrons in the underlying d
subshell. The electron configurations of their outermost electrons are ns1 (n – 1)d10 . Note that this is an
exception to the configuration expected from the Aufbau principle, which predicts ns2 (n – 1)d9 . It is as if the
(n – 1)d orbitals "borrowed" an electron from the higher energy ns orbital. For the coinage metal atoms the
"borrowed" electron is used to complete an inner subshell. Apparently the completed d subshell has an enhanced
stability that corresponds to a lowering of the energy of the atom. Thus ns1 (n – 1)d10 is a lower energy
configuration than ns2 (n – 1)d9 .
In copper, silver, and gold this one outer electron is held much more tightly than the outermost electron in
the alkali metal atoms. This can be most easily seen by comparing ionization energies as shown in Table 8.7.
Table 8.7 Comparison of Ionization Energies
of Coinage Metals and Alkali Metals
EXAMPLE 8.10 Properties of Group 4A Elements
Write the electron configurations for the atoms of all elements in Group 4A. Specify whether the element is a
metal, a nonmetal, or a metalloid.
•Method of Solution
The outermost principal energy level of atoms in Group 4A contains four electrons. The general electron
configuration for elements in this group is ns2 np2 . Their electron configurations are
C 1s 2 2s2 2p2
Ge [Ar]3d10 4s2 4p2
10 5s2 5p2
Pb [Xe]4f14 5d10 6s2 6p2
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EXAMPLE 8.11 Properties of Potassium and Chlorine
Compare the magnitudes of electron affinity and ionization energy for the alkali metal potassium with the
halogen chlorine. Comment on their relative abilities to form ions.
•Method of Solution
From Table 8.4 (text) the electron affinities of K and Cl are 48 and 349 kJ/mol, respectively. Table 8.3 (text)
gives the ionization energies of K and Cl as 419 and 1251 kJ/mol, respectively. These values suggest that the
metal atom K can lose an electron and form a K+ ion much more easily than can the nonmetal Cl. And the
nonmetal atom Cl can attract an electron to form a Cl– much more readily than can the metallic K atom.
27. Of the following, which is the most metallic element?
V, Ge, Se, As, Zn
28. Which of the following is the least metallic element?
V, Ge, Al, As, Ca
29. The first ionization energies of boron and silicon are 801 and 786 kJ/mol , respectively.
The similarities of these values is an example of what relationship?
30. What charges do you expect for the ions of the alkali metals and the alkaline earth metals?
31. What charges do you expect for the ions of the Group 6A and the halogen elements?
_______________________________________________________________________________ CONCEPTUAL QUESTIONS
1. Explain why the sizes of atoms increases when proceeding down a group of the periodic table. 2. Explain why the ionization energy of atoms increases when proceeding from left to right across a period of
the periodic table. 3. Why is an emphasis placed on the valence electrons in an atom when discussing its atomic properties. PRACTICE TEST
1. The element francium is extremely rare, and very little is known about its chemical and physical properties.
Use the following data to estimate its density and ionization energy.
Density (g/cm3 )
Ionization energy (kJ/mol)419
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2. The element technetium (Z = 43) does not occur on earth. The densities of Mo and Ru are 10.2 and 12.4
g/cm3 , respectively. Estimate the density of Tc.
3. The periodic table has been extended to include the transuranium elements (Z > 92). What elements would
element number 104 be similar to in chemical properties?
4. Which is the general electron configuration for the outermost electrons of elements in Group 4A?
a. ns 1
c. ns 2 np4
d. ns2 np2
e. ns 2 np6 nd7
5. In what group of the periodic table is each of the following elements found?
a. 1s 2 2s2 2p6
c. [Xe]6s2 4f14 5d5
d. [Ne]3s2 3p5
6. Use the periodic table to write the electron configurations of Cr, Sb, and Pb.
7. How many valence electrons does an arsenic atom have?
8. Successive ionization energies—1st, 2nd, 3rd, etc.—always show an increasing trend. I1 < I2 < I3 < In . For
aluminum atoms, which ionization energy value will show an exceptionally large increase over the
preceding ionization energy value?
a. 2nd b. 3d c. 4th
9. Write the electron configurations for the following ions.
a. Ca 2+ b . Se 2–
c . Cl – d. Mn2+ e . Co 3+ f . Sc 3+
10. An Ar atom is isoelectronic with which one of the following?
a. Ne b. K c. Sc3+ d. Cl2 – e. Na+
11. Which atom should have the largest radius?
12. Which is the larger ion or atom in each pair?
a. I– or Cs+
b. Ne or K+
c. Mg or Mg 2+
13. Which atom should have the greatest ionization energy?
14. Why does atomic radius decrease in going from left to right across a row in the periodic table?
15. Which two from the following would be most likely to have similar ionization energies?
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1. Elements that have an incompletely filled set of s or p orbitals. Pick any from Groups 1A through 7A
2. Transition metals have incompletely filled d subshells or readily form cations having incompletely filled d
subshells. Pick any from Groups 2B through 8B in the periodic table.
3. These elements have an incompletely filled inner f subshell. Pick any between atomic numbers 58 and 71.
5. a. ns 2 np1 b . ns 2 np4
6. Pick any element from Group 4A.
7. The halogen elements. Group 7A.
8. [Ar] 4s2 3d1 Sc is a metal.
9. a. Group 7A b. Group 5A
c. Group 8B
10. a, c and f have 36 electrons. b and d have 18 electrons.
11. a. K + [Ar] b. O2– [Ne] c. Sr2+ [Kr] d. I – [Xe] e. Li+ [He] f. Mg2+ [Ne]
12. a. [Ar] 3d6 b. [Ar] 3d5 c. [Ar] 3d 9 d. [Ar] 3d10 e. [Ar] 3d 2 f. [Ar]
13. a. [Ar] b. [Ar] c. [Xe] 6s2 5d10 d. [Xe] 5d10
14. a. Na b. Pb
15. Ar and Cl– are one isoelectronic pair, and Na+ and Ne are another.
18. a. Co 2+
b . S 2–
20. a. Noble gases b. Alkai metals
24. Noble gases
29. Diagonal relationship
30. +1 and +2, respectively
31. –2 and –1, respectively Conceptual Questions
1. 2. Back As the principal quantum number n increases, the size of the orbital increases, and the farther on average
electrons will be from the nucleus. Since the outermost electrons have an increasingly higher n value in
going down a group the outermost electrons are farther from the nucleus and the atomic radius increases
When going across a period from left to right protons are being added one at a time to the nucleus and
electrons are added similarly to the outermost principle energy level. The first ionization energy increases
across the period because the electrons in the same principle energy level do not completely shield the
increasing nuclear charge due to the added protons. As the effective nuclear charge increases from left to
right so will the ionization energy. Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website Periodic Relationships among the Elements / 1 63
3. The number of outermost electrons and the type of orbital occupied is what determines the chemical
behavior of an element. Practice Test
14. 15. I ≅ 350 kJ; density ≅ 2.0 g/cm3
Density ≅ 11.3 g/cm3
Ti, Zr, and Hf
d. ns2 np2
a. Group 8A, noble gases
b. Group 1A, alkali metals
c. A transition metal
d. Group 7A, halogens
Cr [Ar]4s2 3d4 ; the observed configuration is Cr [Ar]4s1 3d5 , which is an exception to the rules we have
Sb [Kr]5s 2 4d10 5p3
Pb [Xe]6s2 4f14 5d10 6p2
a. Ca 2+ [Ar]b. Se2– [Kr]
c. Cl– [Ar]
d. Mn2+ [Ar]3d5
e. Co 3+ [Ar]3d6 f. Sc3+ [Ar]
b. K +
The effective nuclear charge increases from left to right because the nuclear charge increases by one proton
at a time, and as electrons are added to the atom they go into the same principal energy level where they do
not shield each other effectively. As Zeff increases, the electrons are pulled closer to the nucleus and the
atomic radius decreases correspondingly.
B(I = 801 kJ) and Si(I = 786 kJ) diagonal relationship.
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