Electron Configuration and Chemical Periodicity
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
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Figure 8.1
Observing the Effect of Electron Spin
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Figure 8.7
Order for filling energy sublevels with
electrons
Illustrating Orbital Occupancies
The electron configuration
n l
#
of electrons in the sublevel
as s,p,d,f
The orbital diagram (box or circle)
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Figure 8.8
A vertical
orbital diagram
for the Li
ground state
no color-empty
light - half-filled
dark - filled, spin-paired
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SAMPLE PROBLEM 8.1
Determining Quantum Numbers from Orbital
Diagrams
PROBLEM: Write a set of quantum numbers for the third electron and a set
for the eighth electron of the F atom.
PLAN: Use the orbital diagram to find the third and eighth electrons.
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
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l= 1
ml = -1
ms= -1/2
Figure 8.9
Orbital occupancy for the first 10 elements, H through Ne.
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8
8
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Figure 8.26
The Period 4
crossover in
sublevel energies
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Figure 8.14
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Defining metallic and covalent radii
Figure 8.15
Atomic radii of the maingroup and transition
elements.
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Atomic Size
• As n increases, electrons spend more time
away from the nucleus : atoms become larger
down a group ( increasing n)
• Across a period, Zeff draws electrons closer,
atoms shrink as nuclear attraction increases
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Figure 8.16
Periodicity of atomic radius
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SAMPLE PROBLEM 8.3
Ranking Elements by Atomic Size
PROBLEM: Using only the periodic table (not Figure 8.15)m rank each set of
main group elements in order of decreasing atomic size:
(a) Ca, Mg, Sr
PLAN:
(b) K, Ga, Ca
(c) Br, Rb, Kr
(d) Sr, Ca, Rb
Elements in the same group increase in size and you go down;
elements decrease in size as you go across a period.
SOLUTION:
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(a) Sr > Ca > Mg
These elements are in Group 2A(2).
(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.
Figure 8.17
Periodicity of first ionization
energy (IE1)
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Figure 8.18
First ionization
energies of the
main-group
elements
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SAMPLE PROBLEM 8.4
Ranking Elements by First Ionization Energy
PROBLEM: Using the periodic table only, rank the elements in each of the
following sets in order of decreasing IE1:
(a) Kr, He, Ar
PLAN:
(b) Sb, Te, Sn
(c) K, Ca, Rb
(d) I, Xe, Cs
IE decreases as you proceed down in a group; IE increases as
you go across a period.
SOLUTION:
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(a) He > Ar > Kr
Group 8A(18) - 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 furtther to the left and
down one period.
Figure 8.20
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Electron affinities of the main-group elements
Main-group ions and the noble
gas configurations
Figure 8.25
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Figure 8.29
Ionic vs.
atomic radius
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SAMPLE PROBLEM 8.8
Ranking Ions by Size
PROBLEM: 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+
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These are members of the same Group (2A/2) and
therefore decrease in size going up the group.
The 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 + charge, the smaller the ion.
Table 8.2 Summary of Quantum Numbers of Electrons in Atoms
Name
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Symbol
Permitted Values
Property
principal
n
positive integers(1,2,3,…) orbital energy (size)
angular
momentum
l
integers from 0 to n-1
orbital shape (The l values
0, 1, 2, and 3 correspond to
s, p, d, and f orbitals,
respectively.)
magnetic
ml
integers from -l to 0 to +l
orbital orientation
spin
ms
+1/2 or -1/2
direction of e- spin
Factors Affecting Atomic Orbital Energies
The Effect of Nuclear Charge (Zeffective)
Higher nuclear charge lowers orbital energy (stabilizes the
system) by increasing nucleus-electron attractions.
The Effect of Electron Repulsions (Shielding)
Additional electron in the same orbital
An additional electron raises the orbital energy through
electron-electron repulsions.
Additional electrons in inner orbitals
Inner electrons shield outer electrons more effectively than
do electrons in the same sublevel.
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Development of Periodic Table
• Elements in the same
group generally have
similar chemical
properties.
• Properties are not
identical, however.
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Development of Periodic Table
Dmitri Mendeleev
and Lothar Meyer
independently
came to the same
conclusion about
how elements
should be grouped.
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Development of Periodic Table
Mendeleev, for instance, predicted the discovery of
germanium (which he called eka-silicon) as an
element with an atomic weight between that of zinc
and arsenic, but with chemical properties similar to
those of silicon.
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Periodic Trends
• In this chapter, we will rationalize observed
trends in
– Sizes of atoms and ions.
– Ionization energy.
– Electron affinity.
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Effective Nuclear Charge
• In a many-electron atom,
electrons are both
attracted to the nucleus
and repelled by other
electrons.
• The nuclear charge that
an electron experiences
depends on both factors.
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Effective Nuclear Charge
The effective nuclear
charge, Zeff, is found this
way:
Zeff = Z − S
where Z is the atomic
number and S is a
screening constant,
usually close to the
number of inner
electrons.
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The effect of
another electron
in the same
orbital
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The effect of other
electrons in inner
orbitals
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The effect of orbital shape
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Interpenetration of s electrons
• Penetration increases attraction to nucleus
• Shields outer electrons more effectively
• Causes orbitals to split in energy sublevels
• lower l value - more penetration
• Results in separation of energy:
s<p<d<f
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