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Chemistry-140
Lecture 20
Chapter 8:
Electron Configurations & Chemical Periodicity
 Chapter Highlights
 electron spin
 Pauli Exclusion Principle
 atomic electron configurations

atomic properties

periodic trends

chemical reactions & periodic properties
Chemistry-140
Lecture 20
Magnetism & Electron Spin
North pole of a magnet
producing a magnetic field
N
N
Electron spinning
counterclockwise
Electron spinning
clockwise
S
South pole of a magnet
producing a magnetic field
S
Chemistry-140
Lecture 20
Magnetism in Chemical Substances
o.ox g
 Paramagnetism
a substance is
attracted to a
magnetic field
Sample sealed
in a glass tube
Electronic
balance
 Diamagnetism:
a substance is
repelled by a
magnetic field
Electromagnet to
provide magnetic field
Chemistry-140
Lecture 20
Quantum Numbers
 The electron spin magnetic quantum number (ms)
has possible values of:
1
1
ms = + or 2
2
It describes the electron character of an electron in an orbital.
 Electron spins are either paired or unpaired
 If there is only one electron the value of ms is arbitrary
Chemistry-140
Lecture 20
Pauli Exclusion Principle
 Pauli exclusion principle: no two electrons can have the
same set of four quantum numbers ( n
l ml
ms ).
Only two electrons can be in the same orbital (sharing the
same values of n
l ml ).
 They must have opposite spins, ms = + 1/2 and - 1/2.
Chemistry-140
Lecture 20
Pauli Exclusion Principle
 Represent the 1s orbital of an H atom as This orbital has quantum numbers
n=1
l
ml = 0
=0
 One electron can occupy this 1s orbital with a value of
ms = + 1/2
H: 1s
or
ms = - 1/2
H: 1s
Chemistry-140
Lecture 20
Pauli Exclusion Principle
 One electron can occupy the 1s orbital of He with a value of
ms = + 1/2
He: 1s
or
ms = - 1/2
He: 1s
 The second electron must occupy the 1s orbital of He such
that the spins are paired
He: 1s
Chemistry-140
Lecture 20
Electron Configurations in Multi-Electron Atoms
 Electron configuration: an arrangement of
electrons in atomic orbitals
 The most stable electron configurations are those that
minimize energy.
 Each element has a ground state electron configuration in
which all the electrons are in the lowest-energy orbitals.
 Any other configuration is an excited state
Chemistry-140
Lecture 20
Determining Atomic Electron Configurations
 The lowest energy configuration is most favourable
 The Pauli exclusion principle must be obeyed.
 Follow Hund's rule. Electrons in a set of degenerate orbitals
(those with equal energy) will go into separate orbitals if
that option is available.
Chemistry-140
Lecture 20
Orbital Energies in Multi-Electron Atoms
 Recall for Hydrogen
1
En = -RH 2
n
 Bohr's equation can be extended to multi-electron atoms
since the energy of an electron is proportional to the square
of the nuclear charge, Z:
 Z2 
En = -RH  2 
n 
Chemistry-140
Lecture 20
Orbital Energies in Multi-Electron Atoms
 Nuclear charge "felt" by an electron is called the effective
nuclear charge, Z*.
Z* = Z - S
(S = screening by closer electrons)
 Due to effective nuclear charge, the energies of orbital
subshells vary for a value of n. For a given value of n, Z*
decreases as l increases. So the energies of the subshells
increase as l increases,
ns < np < nd < nf
Chemistry-140
Lecture 20
Orbital Energies in Multi-Electron Atoms
Chemistry-140
Lecture 20
Orbital Energies for a Multi-Electron Atom
Chemistry-140
Lecture 20
Electron Configurations in Multi-Electron Atoms
 For the simplest case (the Hydrogen atom), one electron
occupies the 1s orbital
H: 1s
 This can be represented as:
 and for Helium:
He: 1s
H: 1s1
becomes
He: 1s2
Chemistry-140
Lecture 20
Valence Electrons
 Valence electrons: the outer-shell electrons - those electrons
present beyond the last full subshell or preceding noble gas.
For example: Aluminum has the
ground state electron configuration
Al: 1s22s22p63s23p1
three valence electrons;
3s23p1
Chemistry-140
Lecture 20
Valence Electrons
 Core electrons: electrons that are not valence electrons. Those
which comprise the preceding noble gas electron configuration.
Al: 1s22s22p63s23p1
Al: [Ne]3s23p1
This emphasizes the valence electrons which
are used in chemical reactions and bonding.
Chemistry-140
Lecture 20
Electron Configurations & the Periodic Table
Chemistry-140
Lecture 20
Electron Configurations for Transition Metals
 Some anomalies occur for the transition elements.
V: [Ar]4s23d3
Cr: [Ar]4s13d5
.
.
.
Ni: [Ar]4s23d8
Cu: [Ar]4s13d10
Chemistry-140
Lecture 20
Electron Configurations for Ions
Cl:[Ne]3s23p5
Cl-:[Ne]3s23p6
Ca:[Ar]4s2
Ca2+:[Ar]
Ni:[Ar]4s23d8
Ni2+:[Ar]3d8
Cu:[Ar]4s13d10
Cu+:[Ar]3d10 cuprous ion
Cu2+:[Ar]3d9 cupric ion
Cl:[Ar]
Chemistry-140
Lecture 21
Chapter 8:
Electron Configurations & Chemical Periodicity
 Chapter Highlights

electron spin

Pauli Exclusion Principle

atomic electron configurations
 atomic properties
 periodic trends
 chemical reactions & periodic properties
Chemistry-140
Lecture 21
Atomic Properties & Periodic Trends
 Members of the same Group have similar chemical
properties and the same valence shell electron
configurations
Example Group 2 Elements:
M
M
M2+
+ 1/2 O2
Be: [He]2s2
Mg: [Ne]3s2
Ca: [Ar]4s2
Sr: [Kr]5s2
+
2 e-
MO
Chemistry-140
Lecture 21
Periodic Trends
A Comparison of Mendeleev's eka-Silicon to Germanium
Property
Prediction for
Observed for
eka-Silicon (1871) Germanium
Atomic weight
72
72.59
Density (g/cm3)
5.5
5.35
Specific Heat (J/g-K)
0.305
0.309
Melting Point (oC)
High
947
Colour
Dark grey
Greyish White
Formula of Oxide
XO2
GeO2
Density of Oxide (g/cm3)
4.7
4.70
Formula of Chloride
XCl4
GeCl4
Boiling Point of Chloride (oC) A little under 100
84
Chemistry-140
Lecture 21
Sizes of Atoms
 The effective nuclear charge, Z*, increases from left to
right in a period and from bottom to top in a group.
Z* = Z - S
(S = screening by core electrons)
 The greater Z*, the greater the attractive force between the
nucleus and its electrons
 The greater Z*, the smaller the average distance between
the nucleus and its valence electrons (the atomic radius)
Chemistry-140
Lecture 21
Atomic Radii
C
154 pm
Cl
200 pm
176 pm
Chemistry-140
Lecture 21
Atomic Radii
1
2
3
4
5
6
7
8
Chemistry-140
Lecture 21
Ionization Energies
 Ionization energy, IE: The energy required to remove the
outermost electron from a gaseous atom or ion. The first
ionization energy, I1 is the energy for the removal of an
electron from a neutral, gaseous atom:
Atom in ground state(g)
Atom(g)+ + e-
DE = ionization energy, IE
Chemistry-140
Lecture 21
Ionization Energies
 The second ionization energy, I2, is the energy of the removal
of a second electron from a gaseous ion with a 1+ charge:
M(g)+
M(g)2+ + e-
 As Z* increases, it requires more energy to remove the
outermost electron. Consequently, ionization energy is
related to atomic radius, with ionization energy increasing as
atomic radius decreases
Chemistry-140
Lecture 21
Trends in Ionization Energies
 Ionization energies can be correlated to electron configuration:
Mg(g)
Mg: [Ne]3s2
Mg(g)+ + eMg+: [Ne]3s1
IE(1) = 738 kJ/mol
Mg(g)+
Mg+: [Ne]3s1
Mg(g)2+ + eMg2+: [Ne]
IE(2) = 1451 kJ/mol
Mg(g)2+
Mg2+: [Ne]
Mg(g)3+ + eMg3+: [He]2s22p5
IE(3) = 7733 kJ/mol
Chemistry-140
Lecture 21
Trends in Ionization Energies
Chemistry-140
Lecture 21
First Ionization Energies
Chemistry-140
Lecture 21
Electron Affinities
 Electron affinity, EA: the energy change of the reaction of
adding an electron to a gaseous atom or ion:
Atom in ground state(g)
+ e-
Atom(g)-
DE = electron affinity, EA
 These reactions tend to be exothermic and so the values of
EA are generally negative.
 In general, electron affinity tends to decrease from left to
right. Going down a group, there is little change
Chemistry-140
Lecture 21
Electron Affinities
Chemistry-140
Lecture 21
Ion Size (Cations)
Li atom (radius = 152 pm)
Li+ cation (radius = 90 pm)
Li
Li+
152
pm
1s
2s
90
pm
1s
2s
Chemistry-140
Lecture 21
Ion Size (Anions)
F atom (radius = 72 pm)
F- anion (radius = 119 pm)
F-
F
72 pm
2s
2p
119 pm
2s
2p
Chemistry-140
Lecture 21
Trends in Ion Size
Chemistry-140
Lecture 21
Ion Sizes in an Isoelectronic Series
Ion
Ionic radius (pm)
2-
O
-
F
+
Na Mg
126 119 116
2+
86
Number of protons
8
9
11
12
Number of electrons
10
10
10
10
Chemistry-140
Lecture 21
Textbook Questions From Chapter # 8
Electronic configurations:
10, 12, 16, 23, 26, 32
Quantum numbers:
36, 41, 42
Periodic Properties:
45, 46, 52, 54
General & Conceptual
60, 61, 73, 75