Chapter 8
Electron Configuration and Chemical Periodicity
8-1
Electron Configuration and Chemical Periodicity
8.1 Development of the Periodic Table
8.2 Characteristics of Many-Electron Atoms
8.3 The Quantum-Mechanical Model and the Periodic Table
8.4 Trends in Some Key Periodic Atomic Properties
8.5 The Connection Between Atomic Structure and Chemical
Reactivity
8-2
Table 8.1 Mendeleev’s Predicted Properties of Germanium (“eka
Silicon”) and Its Actual Properties
Property
atomic mass
appearance
density
molar volume
specific heat capacity
oxide formula
oxide density
sulfide formula
and solubility
chloride formula
(boiling point)
chloride density
element preparation
8-3
Predicted Properties of
eka Silicon(E)
Actual Properties of
Germanium (Ge)
72 amu
gray metal
5.5 g/cm3
13 cm3/mol
0.31 J/g.K
EO2
4.7 g/cm3
ES2; insoluble in H2O;
soluble in aqueous (NH4)2S
ECl4; (< 100 oC)
72.61 amu
gray metal
5.32 g/cm3
13.65 cm3/mol
0.32 J/g.K
GeO2
4.23 g/cm3
GeS2; insoluble in H2O;
soluble in aqueous (NH4)2S
GeCl4; (84 oC)
1.9 g/cm3
reduction of K2EF6 with
sodium
1.844 g/cm3
reduction of K2GeF6 with
sodium
Observing the Effect of Electron Spin
Figure 8.1
8-4
Table 8.2
Name
Summary of Quantum Numbers of Electrons in Atoms
Symbol
Allowed Values
Property
principal
n
positive integers (1, 2,
3,…)
orbital energy (size)
angular
momentum
l
integers from 0 to n-1
magnetic
ml
integers from -l to 0 to +l
orbital shape (l values of 0,
1, 2 and 3 correspond to s,
p, d and f orbitals,
respectively.)
orbital orientation
spin
ms
+1/2 or -1/2
direction of e- spin
Each electron in an atom has its own unique set of four (4) quantum numbers.
8-5
The Pauli Exclusion Principle
No two electrons in the same atom can have
the same four quantum numbers
An atomic orbital can hold a maximum of two electrons
and they must have opposite spins (paired spins)
8-6
Spectral evidence of energy-level splitting in
many-electron systems
Figure 8.2
Many-electron atoms: have nucleus-electron and electron-electron
interactions
Leads to the splitting of energy levels into sublevels of differing energies:
the energy of an orbital depends mostly on its n value (size) and somewhat
on its l value (shape)
8-7
Factors Affecting Atomic Orbital Energies
Effect of nuclear charge (Zeffective)
Higher nuclear charge lowers orbital energy (stabilizes the
system) by increasing nucleus-electron attractions.
Effect of electron repulsions (shielding)
1. Additional electron in the same orbital
An additional electron raises the orbital energy through
electron-electron repulsions.
2. Additional electrons in inner orbitals
Inner electrons shield outer electrons more effectively than
do electrons in the same sublevel.
8-8
The effect of nuclear
charge
Greater nuclear charge
lowers orbital energy
Figure 8.3
8-9
The effect of another
electron in the same
orbital
Each electron shields
the other from the full
nuclear charge, thus
raising orbital energy
Figure 8.4
8-10
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The effect of other
electrons in inner
orbitals
Inner electrons shield
outer electrons very well
and raise orbital energy
greatly
Figure 8.5
8-11
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The effect of orbital shape
Figure 8.6
2s electron farther from nucleus
than 2p electron, but penetrates
near nucleus; increased attraction
results in lower orbital energy
8-12
General Rule for Predicting Relative Sublevel Energies
For a given n value, the lower the l value, the lower the sublevel
energy; thus….
s < p < d < f
8-13
Figure 8.7
Order for filling energy sublevels with
electrons
Illustrating Orbital Occupancies
A. The electron configuration
nl#
of electrons in the sublevel
as s, p, d or f
B. The orbital diagram (box or circle)
8-14
A vertical orbital
diagram for the Li
ground state
no color = empty
light = half-filled
Sublevel energy
increases from
bottom to top
1s22s1
Figure 8.8
8-15
dark = filled, spin-paired
Hund’s Rule
When orbitals of equal energy are available, the electron
configuration of lowest energy has the maximum number
of unpaired electrons with parallel spins.
8-16
Determining Quantum Numbers from Orbital
Diagrams
Sample Problem 8.1
PROBLEM:
PLAN:
Write a set of quantum numbers for the third electron and a set
for the eighth electron of the fluorine (F) atom.
Use the orbital diagram to find the third and eighth electrons.
Up arrow = +1/2
Down arrow = -1/2
9F
1s
2s
2p
SOLUTION: The third electron is in the 2s orbital. Its quantum numbers are:
n= 2
l= 0
ml = 0
ms= +1/2
The eighth electron is in a 2p orbital. Its quantum numbers are:
n= 2
8-17
l= 1
ml = -1
ms= -1/2
Orbital occupancy for the first 10 elements, H through Ne
Figure 8.9
He and Ne have filled outer shells: confers chemical inertness
8-18
8-19
Condensed ground-state electron configurations in
the first three periods
Figure 8.10
Similar outer electron configurations correlate with similar
chemical behavior.
8-20
Similar Reactivities within
A Group
Orbitals are filled in order of increasing
energy, which leads to outer electron
configurations that recur periodically,
which leads to chemical properties that
recur periodically.
Figure 8.11
8-21
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Cr and Cu: Half-filled and filled sublevels are unexpectedly stable!
8-22
8-23
A periodic table of partial ground-state electron configurations
Figure 8.12
8-24
The relation between orbital filling and the Periodic Table
Figure 8.13
8-25
Common tool used to
predict the filling order
Of sublevels
n values are constant horizontally
l values are constant vertically
combined values of n+1 are
constant diagonally
p. 302
8-26
Categories of Electrons
Inner (core) electrons: fill all the lower energy levels of an atom
Outer electrons: those electrons in the highest energy level
(highest n value) of an atom
Valence electrons: those involved in forming compounds; the bonding
electrons; among the main-group elements, the valence electrons
are the outer electrons
8-27
General Observations about the Periodic Table
A. The group number equals the number of outer electrons (those with
the highest value of n) (main-group elements only)
B. The period number is the n value of the highest energy level.
C. The n value squared (n2) gives the total number of orbitals in that
energy level; 2n2 gives the maximum number of electrons
in the energy level.
8-28
SAMPLE PROBLEM 8.2
PROBLEM:
Determining Electron Configuration
Using the periodic table, give the full and condensed electron
configurations, partial orbital diagrams showing valence electrons,
and number of inner electrons for the following elements:
(a) potassium (K: Z = 19) (b) molybdenum (Mo: Z = 42)
PLAN:
(c) lead (Pb: Z = 82)
Use the atomic number for the number of electrons and the periodic table
for the order of filling of the electron orbitals. Condensed configurations
consist of the preceding noble gas plus the outer electrons.
SOLUTION:
(a) for K (Z = 19)
full configuration:
condensed configuration:
1s22s22p63s23p64s1
[Ar] 4s1
K has 18 inner electrons.
partial orbital diagram:
4s1
8-29
SAMPLE PROBLEM 8.2:
(continued)
(b) for Mo (Z = 42)
full configuration:
condensed configuration:
partial orbital diagram:
5s1
1s22s22p63s23p64s23d104p65s14d5
[Kr] 5s14d5
Mo has 36 inner electrons
and 6 valence electrons.
4d5
(c) for Pb (Z = 82)
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p2
full configuration:
condensed configuration: [Xe] 6s24f145d106p2
partial orbital diagram:
Pb has 78 inner electrons
and 4 valence electrons.
8-30
6s2
6p2
KEY PRINCIPLE
All physical and chemical properties of the elements
are based on the electronic configurations of their atoms.
8-31
Defining metallic and covalent radii
A. Metallic radius: 1/2 the
distance between adjacent
nuclei in a crystal
B. Covalent radius: 1/2 the
distance between bonded nuclei
in a molecule
Figure 8.14
8-32
Atomic radii of main-group and
transition elements
Opposing forces: Changes in n and
changes in Zeff
Overall Trends
(A) n dominates within a group; atomic
radius generally increases in a group
from top to bottom
(B) Zeff dominates within a period; atomic
radius generally decreases in a period
from left to right
Figure 8.15
8-33
Periodicity of atomic radius
Large size shifts when
moving from one
period to the next
Figure 8.16
8-34
SAMPLE PROBLEM 8.3
PROBLEM:
Using only the periodic table, rank each set of main group
elements in order of decreasing atomic size.
(a) Ca, Mg, Sr
PLAN:
Ranking Elements by Atomic Size
(b) K, Ga, Ca
(c) Br, Rb, Kr
(d) Sr, Ca, Rb
Size increases down a group; size decreases across a period.
SOLUTION:
(a) Sr > Ca > Mg
These elements are in Group 2A.
(b) K > Ca > Ga
These elements are in Period 4.
(c) Rb > Br > Kr
Rb has a higher energy level and is far to the left.
Br is to the left of Kr.
(d) Rb > Sr > Ca
Ca is one energy level smaller than Rb and Sr.
Rb is to the left of Sr.
8-35
Ionization Energy
The amount of energy required for the complete removal of 1 mol
of electrons from 1 mol of gaseous atoms or ions; an energyrequiring process; value is positive in sign
IE1 = first ionization energy: removes an outermost electron from
the gaseous atom: atom(g)
ion+(g) + e∆E = IE1 > 0
IE2 = second ionization energy: removes a second electron from the
gaseous ion: ion+(g)
ion+2(g) + e∆E = IE2 > IE1
Atoms with a low IE1 tend to form cations during reactions, whereas those
with a high IE1 (except noble gases) often form anions.
8-36
Ionization Energies: Correlations with Atomic Size
1. As size decreases, it take more energy to remove an electron.
2. Ionization energy generally decreases down a group.
3. Ionization generally increases across a period.
8-37
Periodicity of first ionization
energy (IE1)
Figure 8.17
Lowest values for alkali metals; highest values for noble gases
8-38
First ionization
energies of the
main-group
elements
Increase within a
period and decrease
within a group
Figure 8.18
8-39
SAMPLE PROBLEM 8.4
PROBLEM:
Using the periodic table, rank the elements in each of the following
sets in order of decreasing IE1:
(a) Kr, He, Ar
PLAN:
Ranking Elements by First Ionization Energy
(b) Sb, Te, Sn
(c) K, Ca, Rb
(d) I, Xe, Cs
IE decreases down in a group; IE increases across a period.
SOLUTION:
(a) He > Ar > Kr
Group 8A elements- IE decreases down a group.
(b) Te > Sb > Sn
Period 5 elements - IE increases across a period.
(c) Ca > K > Rb
Ca is to the right of K; Rb is below K.
(d) Xe > I > Cs
I is to the left of Xe; Cs is further to the left and
down one period.
8-40
The first three ionization energies
of beryllium (in MJ/mol)
Successive IEs increase, but a large
increase is observed to remove the first
core electron (for Be, IE3)
Figure 8.19
8-41
8-42
SAMPLE PROBLEM 8.5
PROBLEM:
PLAN:
Identifying an Element from Successive
Ionization Energies
Name the Period 3 element with the following ionization energies
(in kJ/mol) and write its electron configuration:
IE1
IE2
IE3
IE4
IE5
1012
1903
2910
4956
6278
IE6
22,230
Look for a large increase in energy that indicates that all of the
valence electrons have been removed.
SOLUTION:
The largest increase occurs at IE6, that is, after the 5th valence
electron has been removed. The element must have five valence
electrons with a valence configuration of 3s23p3, The element must
be phosphorus. P (Z = 15).
The complete electronic configuration is: 1s22s22p63s23p3.
8-43
Electron Affinity (EA)
The energy change accompanying the addition of 1 mol of electrons
to 1 mol of gaseous atoms or ions.
atom(g) + e-
ion-(g)
∆E = EA1 (usually negative)
EA2 is always positive (adding negative charge to negatively
charged ion).
8-44
Electron affinities of the main-group elements
Negative values =
energy is released when
the ion forms
Positive values =
energy is absorbed
to form the anion
Figure 8.20
8-45
General Trends Involving IEs and EAs
Reactive non-metals: Groups 6A and 7A; in their ionic compounds
they form negative ions (have high IEs and
very (-) EAs)
Reactive metals: Group 1A; in their ionic compounds, they form
positive ions (have low IEs and slightly
(-) EAs)
Noble gases: Group 8A; they do not lose or gain electrons (have very
high IEs and slightly (+) EAs)
8-46
Trends in three atomic properties
Figure 8.21
8-47
Metallic Behavior
Metals: shiny solids; tend to lose electrons in reactions with
non-metals (left and lower 3/4 of periodic table)
Non-metals: tend to gain electrons in reactions with metals;
upper right-hand quarter of periodic table
Metalloids: have intermediate properties; located between the
metals and non-metals in the periodic table
Metallic behavior decreases left to right and
increases top to bottom in the periodic table
8-48
Trends in metallic behavior
Figure 8.22
8-49
Moving down a GROUP: elements at the top tend to form
anions and those at the bottom tend to form cations
Moving across a PERIOD: elements at the left tend to form
cations and those at the right lend to form anions
8-50
The change in metallic behavior in Group 5A(15) and Period 3
Figure 8.23
8-51
Acid-Base Properties
Main-group metals: transfer electrons to oxygen; their oxides
are ionic; in water these oxides act as bases (produce OH-)
Nonmetals: share electrons with oxygen; their oxides are
covalent; in water these oxides act as acids (produce H+)
Some metals and many metalloids form oxides that are
amphoteric (can act as an acid or a base in water).
8-52
The trend in acid-base behavior of element oxides
Figure 8.24
red = oxides are acidic
8-53
blue = oxides are basic
Examples
Na2O(s) + H2O(l)
N2O5(s) + H2O(l)
P4O10(s) + 6H2O(l)
Bi2O3(s) + 6HNO3(aq)
2NaOH(aq)
2HNO3(aq)
4H3PO4(aq)
2Bi(NO3)3(aq) + 3H2O(l)
Amphoteric Behavior
Al2O3(s) + 6HCl(aq)
2AlCl3(aq) + 3H2O(l)
Al2O3(s) + 2NaOH(aq) + 3H2O(l)
8-54
2NaAl(OH)4(aq)
Monatomic Ions
Main Group
Elements in Groups 1A, 2A, 6A and 7A that readily form ions either lose
or gain electrons to attain a filled outer level and thus a noble gas configuration.
Their ions are said to be isoelectronic with the nearest noble gas.
Elements in Groups 3A, 4A and 5A form cations via a different process;
they attain pseudo-noble gas configurations.
8-55
Sn ([Kr]5s24d105p2)
Sn+4 ([Kr]4d10) + 4e-
Sn ([Kr]5s24d105p2)
Sn+2 ([Kr]5s24d10) + 2e-
Main-group ions and the noble
gas configurations
Figure 8.25
8-56
SAMPLE PROBLEM 8.6
PROBLEM:
Writing Electron Configurations of Main-Group Ions
Using condensed electron configurations, write reactions for the
formation of the common ions of the following elements:
(a) iodine (Z = 53)
(b) potassium (Z = 19)
(c) indium (Z = 49)
PLAN: Ions of elements in Groups 1A, 2A, 6A and 7A are usually isoelectronic
with the nearest noble gas.
Metals in Groups 3A to 5A can lose the np, or ns and np, electrons.
SOLUTION:
(a) Iodine (Z = 53) is in Group 7A and will gain one e- to be isoelectronic with Xe:
I([Kr]5s24d105p5) + eI- ([Kr]5s24d105p6)
(b) Potassium (Z = 19) is in Group 1A and will lose one e- to be isoelectronic with Ar:
K ([Ar]4s1)
K+ ([Ar]) + e(c) Indium (Z = 49) is in Group 3A(13) and can lose either one electron or three
electrons: In ([Kr]5s24d105p1)
In+ ([Kr]5s24d10) + eIn ([Kr]5s24d105p1)
In3+([Kr]4d10) + 3e-
8-57
Monatomic Ions
Transition Metal Ions
Rarely attain a noble gas configuration
Form more than one cation by losing all of their ns and
some of their (n-1)d electrons
For Period 4 transition metals, the 4s orbital is more
stable that the 3d orbitals; thus the rule “first in, first out” applies.
8-58
The Period 4
crossover in
sublevel energies
Figure 8.26
8-59
General Rules For Ion Formation
Main group, s-block metals: remove all electrons with highest
n value
Main group, p-block metals: remove np electrons before ns
electrons
Transition (d-block) metals: remove ns electrons before
(n-1)d electrons
Non-metals: add electrons to the p orbital of highest n value
8-60
Magnetic Properties of Transition Metal Ions
Chemical species (atoms, ions, molecules) with one or more
unpaired electrons are affected by external magnetic fields.
Ag (Z=47)
Cd (Z=48)
[Kr]5s14d10
[Kr]5s24d10
Species with unpaired electrons exhibit paramagnetism (attracted by
an external magnetic field).
Species with all electrons paired exhibit diamagnetism (not
attracted by an external magnetic field).
8-61
Apparatus for measuring the magnetic behavior of a sample
Figure 8.27
8-62
Some Examples
Fe+3 exhibits greater paramagnetism than Fe.
Fe ([Ar]4s23d6)
Fe+3 ([Ar]3d5) + 3e-
Zn, Zn+2 and Cu+ are diamagnetic, but Cu is paramagnetic.
8-63
Cu ([Ar]4s13d10)
Cu+ ([Ar]3d10) +
Zn ([Ar]4s23d10)
Zn+2 ([Ar]3d10)
e+
2e-
SAMPLE PROBLEM 8.7
PROBLEM:
Use condensed electron configurations, write the reaction for the
formation of each transition metal ion and predict whether the ion is
paramagnetic.
(a) Mn2+(Z = 25)
PLAN:
Writing Electron Configurations and Predicting
Magnetic Behavior of Transition Metal Ions
(b) Cr3+(Z = 24)
(c) Hg2+(Z = 80)
Write the electron configuration and remove electrons starting with
the ns electrons to attain the ion charge. If the remaining
configuration has unpaired electrons, the ion is paramagnetic.
SOLUTION:
(a) Mn2+(Z = 25) Mn([Ar]4s23d5)
(b) Cr3+(Z = 24) Cr([Ar]4s13d5)
(c) Hg2+(Z = 80) Hg([Xe]6s24f145d10)
Mn2+([Ar]3d5) + 2e-
paramagnetic
Cr3+([Ar] 3d3) + 3e-
paramagnetic
Hg2+([Xe] 4f145d10) + 2e-
not paramagnetic (is diamagnetic)
8-64
Ionic Size vs Atomic Size
Ionic radius: an estimate of the size of an ion in a crystalline
ionic compound
General Observations
Cations are smaller than their parent atoms (decrease in electron-electron
repulsions).
Anions are larger than their parent atoms (increase in electron-electron
repulsions).
8-65
Depicting ionic
radii
Figure 8.28
8-66
Ionic vs
atomic
radius
Ionic size increases
down a group
Trends in periods
are complex
For atoms that form more
than one cation: the
greater the ionic charge,
the smaller the ionic radius
Figure 8.29
8-67
Summary on Ionic Size
Ionic size increases down a group.
Ionic size decreases across a period but increases from cation to anion.
Ionic size decreases with increasing (+) (or decreasing (-)) charge
in an isoelectronic series
Ionic size decreases as charge increases for different cations of a
given element
8-68
SAMPLE PROBLEM 8.8
PROBLEM:
Ranking Ions by Size
Rank each set of ions in order of decreasing size, and explain your
ranking:
(a) Ca2+, Sr2+, Mg2+
PLAN:
(b) K+, S2-, Cl-
(c) Au+, Au3+
Compare positions in the periodic table, formation of positive and
negative ions and changes in size due to gain or loss of electrons.
SOLUTION:
(a) Sr2+ > Ca2+ > Mg2+
(b) S2- > Cl- > K+
(c) Au+ > Au3+
8-69
These are members of the same Group (2A) and
therefore decrease in size going up the group.
These ions are isoelectronic; S2- has the smallest Zeff and
therefore is the largest while K+ is a cation with a large Zeff
and is the smallest.
The higher the positive charge, the smaller the ion.