CHE 106 Prof. J. T. Spencer
Chapter Seven: Periodicity
 Periodic
Table - Organizes many
“facts” and trends in chemical
reactivity
 Based upon electronic
configurations of the elements
 Similarities in electronic properties
leads to similarities in chemical
reactivity and structure
Copyright J. T. Spencer 1995 - 1997
1
CHE 106 Prof. J. T. Spencer
2
Chapter Seven: Periodicity
 Element
- a substance that cannot be
decomposed into simpler substances by
chemical or physical means.
 Most elements not found in nature in
“elemental” form but in combination with
other elements (particularly H or O) in the
form of minerals
 Prior to 1800, 31 elements known (mostly
those that were found in elemental form
in nature (i.e., gold, silver, nitrogen, etc...).
CHE 106 Prof. J. T. Spencer
3
Periodic Table Development
 Between
1800 and 1865, 32 more were isolated by
improved chemical techniques.
 Dmitri Mendeleev - noted that chemical and
physical properties recur periodically when the
elements are arranged in order of increasing
atomic weight. (atomic number not known then
although atomic weights generally follow atomic
number sequence)
 When he created the table, he was forced to leave
blank spaces for “undiscovered” elements but he
correctly predicted many of their properties by
use of his table.
CHE 106 Prof. J. T. Spencer
4
Mendeleev’s Guesses
Guesses for Ge based upon Periodic Trends
Table in
TEXT
Mendeleev’s guesses
for Ge (1871)
Atomic Weight
Density
Specific Heat (J/gK)
Melting Point (°C)
Oxide Formula
Density of Oxide
Chloride Formula
Chloride b.p. (°C)
72
5.5
0.305
High
XO2
4.7 g/cm3
XCl4
<100
Observed
properties
for Ge (1886)
72.59
5.35
0.309
947
GeO2
4.70 g/cm3
GeCl4
84
CHE 106 Prof. J. T. Spencer
5
Mendeleev’s Guesses
Guesses for Ge based upon Periodic Trends
Table in
TEXT
Mendeleev’s guesses
for Ge (1871)
Atomic Weight
Density
Specific Heat (J/gK)
Melting Point (°C)
Oxide Formula
Density of Oxide
Chloride Formula
Chloride b.p. (°C)
72
5.5
0.305
High
XO2
4.7 g/cm3
XCl4
<100
Observed
properties
for Ge (1886)
72.59
5.35
0.309
947
GeO2
4.70 g/cm3
GeCl4
84
CHE 106 Prof. J. T. Spencer
Periodic Development
Moseley (1887-1915) discovered atomic number
concept (atomic number = number or protons
and electrons in an atom) from assigning X-ray
from the elements.
 Periodic Table:
– column or groups (increasing n, principal
quantum number, on descending
G
r
o
u
p

Period
6
CHE 106 Prof. J. T. Spencer
7
Periodic Table
1
2
3
1H
3 Li
4
5
6
s orbitals
7
p orbitals
2s Be
9
10
11
f orbitals
8
12
d orbitals
12 M g
19 K
20 Ca
21 Sc
22 Ti
23 V
24 Cr
25 M n
37 Rb
38 Sr
39 Y
40 Zr
41 Nb
42 M o
43 Tc
44
55
56 Ba
57 La
72 Hf
73 Ta
74 W
75 Re
76
89 Ac
104 Unq 105 Unp
106 Unh
107 Ns
58 Ce
59 Pr
60 Nd
90 Th
91 Pa
92 U
3s
7s
87 Fr
88 Ra
15
16
17
closed shell
11 Na
5s
Cs
6s
14
18
2 He
4
4s
13
5B
6C
7N
2p
8O
9F
10 Ne
13 Al
14 Si
15 P
3p
16 S
17 Cl
18 Ar
3dFe
26
27 Co
28 Ni
29 Cu
30 Zn
31 Ga
32 Ge
33 As
4p
34 Se
35 Br
36 Kr
4dRu
45 Rh
46 Pd
47 Ag
48 Cd
49 In
50 Sn
51 Sb
52 Te
53 I
54 Xe
5dOs
77 Ir
78 Pt
79 Au
80 Hg
81 Tl
82 Pb
5p
Bi
6p
83
84 Po
85 At
86 Rn
6d Hs
108
109 M t
61 Pm
62 Sm
63 Eu
64 Gd
4f
65 Tb
66 Dy
67 Ho
68 Er
69 Tm
70 Yb
71 Lu
93 Np
94 Pu
95 Am 96 Cm
5f
97 Bk
98 Cf
99 Es
100 Fm
101 Md 102 No
103 Lr
CHE 106 Prof. J. T. Spencer
8
“Electron Shells”
Electrons with same principal energy value are in the same
“shell”.
 Why is the Ar n = 1 closer than n = 1 Ne or He? SCREENING!

2
(rad)
Helium showed 1 shell (n = 1)
Ar Neon showed 2 shells (n = 1,2)
Argon showed 3 shells ( n= 1,2,3)
Ne
He
Distance from Nucleus
CHE 106 Prof. J. T. Spencer
9
Screening (Shielding)
In many electron atoms, electron-electron repulsions
become more important than electron-nuclear
attractions.
 Estimate the energy of an electron in an orbital by
considering how it, on the average, interacts with its
electronic environment (treat electrons individually).
 The net attractive force that an electron will feel is the
effective nuclear charge (Zeff).
Z = nuclear charge
Zeff = Z - S
S = screening value
 Screening is the average number of other electrons that
are between the electron and the nucleus.

CHE 106 Prof. J. T. Spencer
10
Screening (Shielding)
Zeff = Z - S
Average electronic
charge (S) between
the nucleus and the
electron of interest
r
Z
The larger the
Zeff an electron
feels leads to a
lower energy
for the electron
Electrons outside of sphere of
radius r have very little effect on
the effective nuclear charge
experienced by the electron at
radius r
CHE 106 Prof. J. T. Spencer
11
Atomic Sizes
 Atoms
as hard spheres vs. quantum
mechanical picture without sharply defined
boundaries. (electron density does not
abruptly stop)
 Estimate atomic radii by assuming that
atoms are spheres that “touch” when
bonded together in molecules.
 Atomic radii are approx. constant when an
atom is bonded in many compounds.
CHE 106 Prof. J. T. Spencer
12
Atomic Radii
Molecule
H2
F2
C-C
CH
CF
2r
Radius Increases
Periodic Table
Bond Distance (pm)
74 pm
272 pm
154 pm
114 pm
213 pm
1. Radii increase
from top to
bottom within a
group
2. Radii increase in
moving R to L
within a period
CHE 106 Prof. J. T. Spencer
13
Atomic Sizes

Why does radius decrease across a group?
– Nuclear charge increases while screening
does not increase.
For Outer electron:
Boron; Zeff (approx.) = 5 - 4 = 1 nuclear charge
Carbon; Zeff (approx.) = 6 - 4 = 2 nuclear charges
Nitrogen; Zeff (approx.) = 7 - 4 = 3 nuclear charges
Fluorine; Zeff (approx.) = 9 - 4 = 5 nuclear charges
Radius Increases
Periodic Table
2p electrons DO NOT
screen each other while the
2s very effectively screen
the 2p, also each element to
right adds +1 to nucleus.
14
CHE 106 Prof. J. T. Spencer
Atomic Sizes
Li < Na < K < Rb < Cs, due to increasing n values.
 Li > F because of Zeff changes.

Na
Atomic
Radius
(Å)
Rb
K
Transition
Metals
Li
F
I
Br
Cl
10
Transition
Metals
20
30
Atomic Number
40
50
CHE 106 Prof. J. T. Spencer
15
Atomic Sizes: Charge Effects
Fe
(232 pm)
F
(128 pm)
Fe+2
(152 pm)
F -1
(272 pm)
Fe+3
(128 pm)
Li
(304 pm)
Li+
(120 pm)
CHE 106 Prof. J. T. Spencer
Atomic Sizes: Charge Effects
Sample exercise: Predict which will be
greater, the P-Br bond length in PBr3 or
the As-Cl bond length in AsCl3.
16
CHE 106 Prof. J. T. Spencer
17
Atomic Sizes: Charge Effects
Sample exercise: Predict which will be
greater, the P-Br bond length in PBr3 or the
As-Cl bond length in AsCl3.
PBr3
P - 1.06
Br - 1.14
Length: 2.20
CHE 106 Prof. J. T. Spencer
18
Atomic Sizes: Charge Effects
Sample exercise: Predict which will be
greater, the P-Br bond length in PBr3 or the
As-Cl bond length in AsCl3.
PBr3
P - 1.06
Br - 1.14
Length: 2.20
AsCl3
As - 1.19
Cl - 0.99
Length: 2.18
CHE 106 Prof. J. T. Spencer
19
Atomic Sizes: Charge Effects
Sample exercise: Arrange the following
atoms in order of increasing atomic radius:
Na, Be, Mg
CHE 106 Prof. J. T. Spencer
20
Atomic Sizes: Charge Effects
Sample exercise: Arrange the following
atoms in order of increasing atomic radius:
Na, Be, Mg
- radius decreases as you go across the
period, increases as you go down the
group.
CHE 106 Prof. J. T. Spencer
21
Atomic Sizes: Charge Effects
Sample exercise: Arrange the following
atoms in order of increasing atomic radius:
Na, Be, Mg
- radius decreases as you go across the
period, increases as you go down the
group.
CHE 106 Prof. J. T. Spencer
22
Atomic Sizes: Charge Effects
Sample exercise: Arrange the following
atoms in order of increasing atomic radius:
Na, Be, Mg
- radius decreases as you go across the
period, increases as you go down the
group.
Be < Mg < Na
CHE 106 Prof. J. T. Spencer
23
Ionization Energy (IE)
Electrons are “exchanged” in many chemical reactions.
[i.e., A gives electron(s) to B]
 Ionization Energy (IE) measures how strongly an atom
holds on to its electrons (the minimum energy necessary
to remove an electron from the ground state of an isolated
gaseous atom).

E(g)

E+1(g) + 1 e-
Higher ionization processes are possible:
– first IE - remove 1st electron;
A
– second IE - remove 2nd electron; A+1
– third IE - remove 3rd electron;
A+2
A+1 + eA+2 + eA+3 + e-
CHE 106 Prof. J. T. Spencer
24
Ionization Energies
 I1
< I2 < I3 (increasing positive charge on atom).
 Very sharp increase in IE when an inner shell (core)
electron is removed.
 Within each period, I1, generally increases with increasing
atomic number (increasing Zeff).
 Within each group, IE generally decreases with increasing
atomic number (“bigger” atoms).
1st IE Increases
Periodic Table
CHE 106 Prof. J. T. Spencer
Ionization Energies (IE)
He (1s2)
1st IE
(kJ/mol) H (1s1)
Atomic Number
25
CHE 106 Prof. J. T. Spencer
Ionization Energies (IE)
He (1s2)
1st IE
(kJ/mol) H (1s1)
Li (1s22s1)
Atomic Number
26
CHE 106 Prof. J. T. Spencer
Ionization Energies (IE)
He (1s2)
1st IE
(kJ/mol) H (1s1)
Be (1s22s2)
Li (1s22s1)
Atomic Number
27
CHE 106 Prof. J. T. Spencer
28
Ionization Energies (IE)
1s
2s
He (1s2)
2p
2p
2p
(1s22s22p6) Ne
(1s22s22p3) N
1st IE
(kJ/mol) H (1s1)
Be
(1s22s2)
F (1s22s22p5)
O (1s22s22p4)
C (1s22s22p2)
B (1s22s22p1)
Li (1s22s1)
Atomic Number
CHE 106 Prof. J. T. Spencer
29
Ionization Energies (IE)
Screening and
increasing Zeff
He
Ne
Zeff = 1 to 2
Zeff = 2 to 3
1st IE
(kJ/mol) H
F
Zeff = 1 to 2
N
Be
C
B
Li
2s to 2p
Zeff = 1 to 2
Atomic Number
O
Zeff = 5 to 6
Zeff = 4 to 5
CHE 106 Prof. J. T. Spencer
30
Ionization Energies (IE)
8O
1s
2s
2p
2p
2p
He
Ne
pairing energy
F
1st IE
(kJ/mol) H
N
Be
n = 1 to 2
C
B
Li
O
increasing n and
pairing energy
Atomic Number
CHE 106 Prof. J. T. Spencer
31
Ionization Energies (IE)
Trends recur throughout periodic table
He
2400
Ne
1800
Ar
Kr
1st IE
(kJ/mol) H
1200
600
Li
Na
K
Atomic Number
Rb
CHE 106 Prof. J. T. Spencer
32
Ionization Energies (IE)
Elem. Elec. Config.
I1
I2
I3
I4
I5
I6
Na
[Ne]3s1
496
4560
Mg
[Ne]3s2
738
1450
7730
Al
[Ne]3s23p1
577
1816
2744
11600
Si
[Ne]3s23p2
786
1577
3228
4354
16100
P
[Ne]3s23p3
1060
1890
2905
4950
6270 21200
S
[Ne]3s23p4
999
2260 3375 4565
Core Electrons
Being Removed
6950 8490
CHE 106 Prof. J. T. Spencer
33
Ionization Energies (IE)
Sample exercise: Based on the trends
discussed in this section, predict which of the
following atoms, B, Al, C, or Si - has the
lowest first ionization energy.
CHE 106 Prof. J. T. Spencer
34
Ionization Energies (IE)
Sample exercise: Based on the trends
discussed in this section, predict which of the
following atoms, B, Al, C, or Si - has the
lowest first ionization energy.
As atoms get smaller, the ionization energy gets larger, so the
lowest ionization energy belongs to the largest atom.
CHE 106 Prof. J. T. Spencer
35
Ionization Energies (IE)
Sample exercise: Based on the trends
discussed in this section, predict which of the
following atoms, B, Al, C, or Si - has the
lowest first ionization energy.
As atoms get smaller, the ionization energy gets larger, so the
lowest ionization energy belongs to the largest atom.
Size decreases as you go across the period, so the left-most
atom is the largest.
CHE 106 Prof. J. T. Spencer
36
Ionization Energies (IE)
Sample exercise: Based on the trends
discussed in this section, predict which of the
following atoms, B, Al, C, or Si - has the
lowest first ionization energy.
As atoms get smaller, the ionization energy gets larger, so the
lowest ionization energy belongs to the largest atom.
Size decreases as you go across the period, so the left-most
atom is the largest.
Al
CHE 106 Prof. J. T. Spencer
37
Electron Affinities (EA)

Electron Affinity - the energy associated with adding
an electron to a gaseous atom.
E(g) + e-1

E-1(g)
Signs;
– Negative (exothermic) when energy is released
upon adding an electron.
– Positive (endothermic) when energy is required to
add an electron.
– Most neutral atoms and all ions (cations) have
negative (exothermic) EA’s.
CHE 106 Prof. J. T. Spencer
38
Electron Affinities (EA)
Be
EA
(kJ/mol)
Mg
He
H
Li
C
Ar
Ne
N
B
Cs
Na
100
0
P
O
200
K -100
Al
Si
-200
S
-300
F
Atomic Number
Cl
-400
CHE 106 Prof. J. T. Spencer
39
Electron Affinities (EA)
Be
Filled
Shell
Already
EA
(kJ/mol)
ns to np
Mg
Ne
He
Ar
0
Li
np5 to
Filled
Shell
Cs
K
Na
ns1 to
ns2
F
Atomic Number
Cl
CHE 106 Prof. J. T. Spencer
40
Electron Affinities (EA)

Summary of General Observations:
– Neutral atoms and cations, EA exothermic (neg).
– EA becomes more negative across period (Zeff
increases).
– Group 2 EA positive because of ns2 to ns2np1
addition.
– Gp 18 has filled shell (no need to gain electrons).
– Gp 15 has significant electron-electron repulsions
(pairing electrons going from ns2np3 to ns2np4).
– EA’s do not range significantly down group (because
decreased nuclear attraction (Zeff) is offset by
decreased electron-electron repulsions).
CHE 106 Prof. J. T. Spencer
41
Trends and Group Properties
Trends of radius, IE, EA, etc. are useful in predicting
chemical behavior
 radius, IE, etc.. are ATOMIC properties (individual atoms)
which only noble gases exist as isolated atoms in nature.
 Trends and properties of groups of atoms (and atoms in
compounds) are also useful (i.e., metal, non-metal,
conductor, insulator, etc...).

Non metal character
increases
Periodic Table
CHE 106 Prof. J. T. Spencer
Metals, Nonmetals and Metalloids
metals
Gold
non-metals
Graphite
42
CHE 106 Prof. J. T. Spencer
43
Metals, Nonmetals and Metalloids
metals
non-metals
conductors
shiny
high thermal conductivity
solids at RT (except Hg)
ductile and malleable
insulators
dull
thermal insulators
freq. non-solids at RT
brittle
Metalloids (along line in table) have properties
between metals and non-metals
CHE 106 Prof. J. T. Spencer
44
Metals, Nonmetals and Metalloids
metals
non-metals
Low Ionization Energies
High IE
 tend to form cations
 tend to form anions
TM’s form multiple + states Tend to form single (-)
states
Basic Oxides
Acidic Oxides
Metalloids (along line in table) have properties
between metals and non-metals
CHE 106 Prof. J. T. Spencer
45
Metals, Nonmetals and Metalloids
1
2
3
4
5
6
1H
7
metals
metalloids
non-metals
8
9
10
11
12
13
14
15
16
Nonmetallic
Properties
Increase
18
17
2 He
3 Li
4 Be
5B
6C
7N
8O
9F
10 Ne
11 Na
12 M g
13 Al
14 Si
15 P
16 S
17 Cl
18 Ar
19 K
20 Ca
21 Sc
22 Ti
23 V
24 Cr
25 M n
26 Fe
27 Co
28 Ni
29 Cu
30 Zn
31 Ga
32 Ge
33 As
34 Se
35 Br
36 Kr
37 Rb
38 Sr
39 Y
40 Zr
41 Nb
42 M o
43 Tc
44 Ru
45 Rh
46 Pd
47 Ag
48 Cd
49 In
50 Sn
51 Sb
52 Te
53 I
54 Xe
55 Cs
56 Ba
57 La
72 Hf
73 Ta
74 W
75 Re
76 Os
77 Ir
78 Pt
79 Au
80 Hg
81 Tl
82 Pb
83 Bi
84 Po
85 At
86 Rn
87 Fr
88 Ra
89 Ac
104 Unq 105 Unp 106 Unh 107 Ns
108 Hs
109 M t
69 Tm
70 Yb
71 Lu
Metallic
Properties
Increase
58 Ce
59 Pr
60 Nd
61 Pm
62 Sm
63 Eu
64 Gd
65 Tb
66 Dy
67 Ho
68 Er
90 Th
91 Pa
92 U
93 Np
94 Pu
95 Am 96 Cm
97 Bk
98 Cf
99 Es
100 Fm 101 Md 102 No
103 Lr
CHE 106 Prof. J. T. Spencer
46
Metal Compounds : Metal Oxides
Compounds of metals and nonmetals tend to be ionic.
 Metal Oxides are basic;

Metal Oxide + Water
Metal Hydroxide
Na2O(s) + H2O(l)
K2O(s) + H2O(l)
MgO(s) + H2O(l)
2 NaOH(aq)
2 KOH(aq)
Mg(OH)2(aq)
Metal Oxide + Acid
Salt + Water
MgO(s) + 2 HCl(aq)
MgCl2(aq) +
H2O(l)
Fe2O3(s) + 3HNO3(aq)
Fe(NO3)3(aq) + 3H2O(l) CuO(s) + H2SO4(aq)
CuSO (aq) + H O(l)
CHE 106 Prof. J. T. Spencer
47
Non-metal Compounds
Nonmetals reacting with metals give salts (electron
receivers)
 Non metal oxides are acidic;

Nonmetal oxide + water
P4O10(s) + 6 H2O(l)
B2O3(s) + 3 H2O(l)
SO3(g) + H2O(l)
Nonmetal oxide + base
CO2(g) + 2NaOH(aq)
H2O(l) B2O3(s) + 6NaOH
2NaBO3(aq) + 6H2O(l)
acid
4 H3PO4(aq)
2 H3BO3(aq)
H2SO4(aq)
salt + water
Na2CO3(aq) +
CHE 106 Prof. J. T. Spencer
48
CO2 Nonmetal Chemistry
Nonmetal oxides are acidic

Carbon dioxide (CO2) is a “typical” non-metal and
dissolves in water to form an acidic solution:
CO2(s) + H2O(l)
H2CO3(aq) + H2O(l)
HCO3-(aq) + H2O(l)
H2CO3(aq)
H3O+(aq) + HCO3-(aq)
H3O+(aq) + CO3-2(aq)

pH - measure of the acidity of a solution.

Indicators - display different colors depending upon
the pH of the solution.
Demonstration 6.2
CHE 106 Prof. J. T. Spencer
Metalloids
 Along
Diagonal Line in Periodic Table
 Have
properties between metals and
nonmetals
 Metalloid
oxides - AMPHOTERIC -
– As bases:
Al2O3 + 6H+
– As Acids:
Al2O3 + 2OH- + 3 H2O
2Al+3 + 3 H2O
2 Al(OH)4-
49
CHE 106 Prof. J. T. Spencer
Group Trends: Group I (1A)
Alkali Metals
3Li
7
11Na
23
19K
39
37Rb
85
55Cs
133
87Fr
223
Li
MOST REACTIVE
Na
K
50
CHE 106 Prof. J. T. Spencer
51
Group Trends: Group I (1A)
3Li
7
1st IE
decreases
11Na
23
19K
39
37Rb
85
55Cs
133
87Fr
223
MOST REACTIVE
Alkali Metals
mp/bp
Radius
increases decreases
electron sharing
dec. (covalency)
CHE 106 Prof. J. T. Spencer
52
Group Trends: Group I (1A)
 Soft,
3Li
7
11Na
23
19K
39
37Rb
85
55Cs
133
87Fr
223
metallic solids.
 Low first ionization energies (electron
donors); form +1 ions. Form pure metals by
electrolysis (passing electrical current
through a molten salt).
2Na+ + 2e2Cl Reactions
2 Na
Cl2 + 2 e-
dominated by 1 electron loss (to
1+) [i.e., Rx with hydrogen = MH; Rx with
S = M2S,etc...].
 MH are hydride compounds (H-1 not H+).
CHE 106 Prof. J. T. Spencer
53
Group Trends: Group I (1A)
 Oxides:
3Li
7
11Na
23
19K
39
37Rb
85
55Cs
133
87Fr
223
Li + O2
2Li2O
lithium oxide
[O]22Na + O2
Na2O2
sodium
peroxide [O2]2-K + O2
KO2
pot.
superoxide [O2] React with water to form hydroxides;
2M + 2 H2O
2 MOH(aq) + H2
H2O reactivity; Li - v. slowly; Na vigorously; K - inflames; Rb & Cs - explode
 React Flame Tests - elements excited to
higher state by the flame and then emit
light as they return to the ground state (Na yellow, 3p to 3s).
CHE 106 Prof. J. T. Spencer
Flame Tests
54
CHE 106 Prof. J. T. Spencer
Group Trends: Group 2 (2A)
Alkaline Earth Metals
4Be
9
12Mg
24
20Ca
40
38Sr
87
56Ba
137
88Ra
226
MOST REACTIVE
Mg
Ca
55
CHE 106 Prof. J. T. Spencer
56
Group Trends: Group 2 (2A)
4Be
9
1st IE
decreases
12Mg
24
20Ca
40
38Sr
87
56Ba
137
88Ra
226
MOST REACTIVE
Alkaline Earth Metals
electron sharing
mp/bp
Radius
increases decreases decr. (covalency)
CHE 106 Prof. J. T. Spencer
57
Group Trends: Group 2 (2A)
Alkaline Earth Metals
4Be
9
12Mg
24
20Ca
40
38Sr
87
56Ba
137
88Ra
226
 Tendency
to lose two electrons to form
M+2 cations (achieves noble gas electron
config.).
 Mg(s) + Cl2(g)
MgCl2(s) and
MgCl(s)
 Flame Tests: Ca - brick red; strontium crimson red (in fireworks); barium green.
 In nature;
beryl
Be3Al2(SiO3)6
dolomite CaCO3MgCO3
limestone CaCO3
CHE 106 Prof. J. T. Spencer
58
Group Trends: Group 13 (3A)
5B
11
13Al
27
1st IE
decreases
Radius
increases
mp/bp
decreases
electron sharing
decr. (covalency)
+3 Ox. State
31Ga
70
49In
115
81Tl
204
+1 Ox. State
CHE 106 Prof. J. T. Spencer
59
Group Trends: Group 13 (3A)
5B
11
13Al
27
31Ga
70
49In
115
81Tl
204
Boron Compounds;
Numerous Polyhedral Compounds (covalent)
Boron Neutron Capture Therapy
 Aluminum;
2Al + Fe2O3
Al2O3 + 2Fe
THERMITE REACTION

Gemstones; Al2O3 (Alumina)
trace Cr+3 = ruby
trace Fe+2, Fe+3, Ti+3 = blue sapphire
CHE 106 Prof. J. T. Spencer
60
Clusters??
• Dictionary - “A number of things held together”.
• Earliest Man-made Polyhedra: Neolithic Scots
• Plato : Five “Platonic” Bodies: All Triangular Face
(Its not what you discover but who publishes First)
• tetrahedron
trigonal
bipyramid
octahedron
dodecahedron
icosahedron
• Archimedes : Thirteen Semi-Regular Polyhedra
Derived from the Platonic Solids
CHE 106 Prof. J. T. Spencer
Boron Neutron Capture Therapy
(BNCT)
10B
11B*
7Li
+  + 2.4 MeV
thermal neutron
to Cells
C
E
L
L
61
thermal
neutron
CHE 106 Prof. J. T. Spencer
62
Group Trends: Group 16 (6A)
Chalcogens
8O
16
16S
32
34Se
79
52Te
128
84Po
209
1st IE
decreases
Radius
increases
mp/bp
increases
electron sharing
decr. (covalency)
CHE 106 Prof. J. T. Spencer
63
Group Trends: Group 16 (6A)
Chalcogens
8O
16
16S
32
34Se
79
52Te
128
84Po
209
 Oxygen
- two allotropes (allotropes
- different form of the same
element) - O2 and O3.
 Great tendency to gain electrons
(oxidize other elements)
 Most common O-2.
 S reacts similarly to oxygen
CHE 106 Prof. J. T. Spencer
64
Group Trends: Group 17 (7A)
Halogens
9F
19
17Cl
35
35Br
80
53I
127
85At
210
1st IE
decreases
Radius
increases
mp/bp
increases
electron sharing
decr. (covalency)
CHE 106 Prof. J. T. Spencer
65
Group Trends: Group 17 (7A)
 Greek
9F
19
17Cl
35
35Br
80
53I
127
85At
210
for “salt formers”.
 Diatomic in elemental state (F2, Cl2, Br2,
etc...).
 Highest electron affinities - form E-1
anions.
– F removes electrons from almost
everything else.
– Reacts with most metals directly to
form salts.
– Reacts with hydrogen to form hydrogen
halides which dissolve in water to form
acids (all except HF are strong acids).
CHE 106 Prof. J. T. Spencer
66
Group Trends: Group I8 (8A)
2He
4
10Ne
20
18Ar
40
36Kr
83
54Xe
131
86Rn
222
1st IE
decreases
Noble Gases
mp/bp
Radius
increases increases
electron sharing
decr. (covalency)
CHE 106 Prof. J. T. Spencer
67
Group Trends: Group I8 (8A)
Noble Gases
2He
4
10Ne
20
18Ar
40
36Kr
83
54Xe
131
86Rn
222
Very unreactive and have closed shell electronic
configurations [called noble or inert gases].
 All are monoatomic in native state.
 1962, Bartlett noticed that Xe has a similar
ionization energy to oxygen and could possible
form compounds with Xe.
– Reacted Xe with very strong oxidizer
(remover of electrons) F and O to form XeF2,
XeF4, XeF6, XeO3, etc...
– No compounds of He, Ne or Ar are known
and only one compound is known for Kr
(KrF2).

CHE 106 Prof. J. T. Spencer
Setting up ionic compounds,
molecular compounds, and
chemical equations
Sample exercise: Predict the formula
of the compound formed by Rb and
Se.
68
CHE 106 Prof. J. T. Spencer
Setting up ionic compounds,
molecular compounds, and
chemical equations
Sample exercise: Predict the formula
of the compound formed by Rb and
Se.
Rb1+
Se2-
69
CHE 106 Prof. J. T. Spencer
Setting up ionic compounds,
molecular compounds, and
chemical equations
Sample exercise: Predict the formula
of the compound formed by Rb and
Se.
Rb1+
Se2Rb2Se
70
CHE 106 Prof. J. T. Spencer
Setting up ionic compounds,
molecular compounds, and
chemical equations
Sample exercise: Write the balanced
chemical equation for the reaction
between copper II oxide and
sulfuric acid.
71
CHE 106 Prof. J. T. Spencer
Setting up ionic compounds,
molecular compounds, and
chemical equations
Sample exercise: Write the balanced
chemical equation for the reaction
between copper II oxide and
sulfuric acid.
Cu2+ O2-
72
CHE 106 Prof. J. T. Spencer
Setting up ionic compounds,
molecular compounds, and
chemical equations
Sample exercise: Write the balanced
chemical equation for the reaction
between copper II oxide and
sulfuric acid.
Cu2+ O2CuO + H2SO4 
73
CHE 106 Prof. J. T. Spencer
Setting up ionic compounds,
molecular compounds, and
chemical equations
Sample exercise: Write the balanced
chemical equation for the reaction
between copper II oxide and
sulfuric acid.
Cu2+ O2CuO + H2SO4  CuSO4 + H2O
74
CHE 106 Prof. J. T. Spencer
Chapter Seven
 Periodic
Table Trends and
Generalizations
 Electron Shells
 Atomic Radii and screening (shielding)
 Ionization Energy
 Electron Affinities
 Metals, Nonmetals and Metalloids
 Group I, 2, 16, 17 and 18 Chemistry Examples of above concepts and trends
75