Periodic Properties
■  Atomic
& Ionic Radius
■  Ionization
■  Electron
Energy
Affinity
☛ We
want to understand the variations in
these properties in terms of electron
configurations.
Valence Electrons
The Periodic Table
■  Elements
in a column of the periodic
table have very similar electron
configurations.
■  Elements
in a column of the periodic
table also have similar chemical
properties.
➡  Valence
electrons determine the
chemical behavior of an atom.
Can you explain this graph?
■  Valence
electrons are those with the
highest n value, plus any in partially
filled d or f shells.
■  These
electrons are the farthest from
the nucleus, and they have the highest
energies. Thus they are the most
“accessible” to other atoms.
Valence Electrons
■  Valence
electrons are those with the
highest n value, plus any in partially
filled d or f shells.
■  These
electrons are the farthest from
the nucleus, and they have the highest
energies. Thus they are the most
“accessible” to other atoms.
☛ Valence electrons determine the
chemical properties of an atom.
Paramagnetism
■  Electrons
have magnetic properties.
■  Two electrons with opposite spins have
opposite magnetic properties, so they
“cancel out” each other’s magnetism.
■  Atoms with unpaired electrons are
attracted or repelled by magnetic fields,
and are said to be “paramagnetic.”
■  This is a way to “verify” e– configuration.
Atomic Radius
■  Size
of an atom is determined mainly by
valence electrons. (Why?)
■  Hard
to define and measure
■  Which
would you expect to be larger:
Na or K?
N or F?
Atomic Radius Variations
■  Moving
across a row, Z increases while
valence electrons are added to the
same n-shell → size decreases
■  Moving
down a column, the n quantum
number of the valence electrons
increases → size increases
Ionic Radii
■  Think
about these the same way as for
atoms: electron configurations.
■  Be
Atomic
Radius
Variations
■  Which
Ionic Radii
■  Anions
are always bigger than the
corresponding neutral atom.
■  Cations
are always smaller than the
corresponding neutral atom.
■  For
sure to use correct # of electrons.
isoelectronic ions, the larger the
nuclear charge, the smaller the ion.
should be larger: Mg or Mg2+?
F or F–?
Ionization Energy
■  IE
= amount of energy needed to
remove an electron from a free neutral
atom.
Z + energy → Z+ + e–
■  Easy
to measure (experiment similar to
photoelectric effect)
Ionization Energy
■  How
would you expect IE to vary as you
go down a column of the Periodic Table?
■  ... as you go across a row?
■  Why?
■  Tell
us the relative stabilities of different
orbitals
Ionization Energies
Ionization Energies
■  IE
decreases going down a column of
the periodic table.
■  IE increases going across a row of the
periodic table.
■  Some exceptions:
➡ Filled shells or subshells are especially
stable.
➡ Half-filled subshells are also fairly
stable.
Ionization
Energies
Can you explain the
general trends
and the variations?
kJ/mol
kJ/mol
Ionization
Energies
Again?
Again?
2nd Ionization Energies
Higher Ionization Energies
■  Can
also define and measure higher
IE’s:
2nd IE:
Z+ + energy → Z2+ + e–
■  Can
predict and understand these
based on electron configuration of the
ions involved.
5000
4500
4000
3500
3000
2500
2000
1500
1000
500
0
2nd Ionization
1st Ionization
Energy
Energy
1st Ionization
Energy
Li Be B C N O F Ne NaMg Al Si P S Cl Ar K
Electron Attachment Enthalpy
(Affinity)
Ionization
Energies
■  Measures
atom’s tendency to form anions
= energy released upon adding an
electron to a neutral atom,
Z(g) + e– → Z– (g) + energy (=EA)
■  Older books called this the electron
affinity, but is now called the Electron
Attachment Enthalpy. This is opposite in
sign from old definition.
■  Note: EA can be positive or negative
kJ/mol
■  EA
Transition
Metals - some
complications
Periodic Table: Metals & Non-metals
Electron Affinity
■  If
EA is negative, the atom “wants” to
add an electron and form an anion.
■  If EA is positive, the atom does not
“want” to add an electron to form an
anion. i.e., the anion is unstable.
■  Which elements will have the most
negative EA’s?
Why?
■  What
makes an element a metal or a
non-metal?
Properties?
Electron configuration?
■  How
are metals & non-metals grouped
in the periodic table?
Comments on the Periodic Table
2.1
1
H
3
1.5
4
Li Be
6.941
1.0
11
9.012
1.2
22.99
0.9
19
37
20 1.3
40.08
1.0
21
85.47
55
87.62
1.0
44.96
38 1.2
Rb Sr
0.8
132.9
87
39
Y
88.91
56 1.1
137.3
1.0
13
1.8
1.4
22 1.5
57
47.88
1.3
23
1.6
41
1.6
V
C
10.81
1.6
25 1.7
26
1.7
27 1.8
138.9
88 1.1
89
28
1.8
29
1.6
30 1.7
31
3.5
15
2.5
N
12.01
33
16
3.0
34
F Ne
91.22
52.00
92.91
72 1.4
178.5
104
73
42
54.94
1.7
95.94
1.5
180.9
2.1
20.18
17
18
35.45
2.8
55.85
43 1.8
44
58.93
1.8
58.69
45 1.8
46
63.55
1.6
47
65.39
1.6
69.72
48 1.6
49
72.61
1.8
74.92
50 1.9
51
78.96
2.1
52
39.95
35
36
79.90
2.5
Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
1.3
10
Cl Ar
32.07
2.4
9
19.00
S
30.97
32 2.1
4.0
16.00
P
28.09
8
O
14.01
14 2.1
1.9
7
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
50.94
40 1.5
24
6 3.0
Al Si
74
(98)
1.7
101.1
75 1.9
76
102.9
1.9
106.4
77 1.8
Cs Ba La Hf Ta W Re Os Ir
0.8
2.5
26.98
K Ca Sc Ti
39.10
0.9
5
B
1.5
24.30
1.0
He
4.003
2.0
Metals
12
Na Mg
2
Nonmetals
Zintl Line
(fuzzy)
1.008
1.0
183.8
105
186.2
106
190.2
78
107.9
1.9
79
112.4
1.7
114.8
80 1.6
81
118.7
1.7
121.8
82 1.8
83
I
127.6
1.9
84
83.80
53
126.9
2.1
54
Xe
131.3
85
86
Pt Au Hg Tl Pb Bi Po At Rn
192.2
195.1
197.0
200.6
204.4
207.2
209.0
(209)
(210)
(222)
107
Fr Ra Ac UnqUnpUnhUns
(223)
(226)
(227)
(261)
(262)
1.1
Lanthanide series
140.1
1.2
Actinide series
(263)
58 1.1
59
(262)
1.1
60 1.1
61
1.1
62
1.1
63 1.1
64
1.1
65 1.1
66
1.1
67
1.1
68 1.1
69
1.0
70 1.2
71
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
140.9
90 1.3
91
Th Pa
232.0
231.0
144.2
1.5
(145)
92 1.3
93
150.4
1.3
94
152.0
1.3
157.2
95 1.3
96
158.9
1.3
162.5
97 1.3
98
164.9
1.3
99
167.3
1.3
168.9
100 1.3
173.0
101 1.3
175.0
102 1.5
103
U Np Pu Am Cm Bk Cf Es Fm Md No Lr
238.0
(237)
(244)
(243)
(247)
(247)
(251)
(252)
(257)
(258)
(259)
(260)