The Periodic Table
The how and why
Atomic Size
 First
problem where do you start
measuring.
 The electron cloud doesn’t have a
definite edge.
 Chemists get around this by measuring
more than 1 atom at a time.
2
Atomic Size
}
Radius
Atomic
Radius = half the distance between two
nuclei of a diatomic molecule (homo-nuclear
molecule).
3
Trends in Atomic Size
Influenced
by two factors.
Energy Level
Higher energy level is further away.
The effective charge from the
nucleus
The greater the nuclear charge
reaching the valence electrons the
closer these electrons are pulled in.
4
Group trends
 As
we go down a group
 Each atom has another
energy level, so the
atoms get bigger.
 There are more levels in
the kernel and therefore
greater shielding of
valence electrons
(weaker attraction).
H
Li
Na
K
Rb
5
Periodic Trends






As you go across a period the radius gets smaller.
Valence electrons are in the same energy level.
Greater nuclear charge gets to the valence electrons –
same shielding e- but greater nuclear charge.
Outermost electrons are closer.
Row 3 elements all have 10 shielding electrons;
+11
(nuclear charges)
+18
Na
Mg
Al
Si
P
S Cl Ar
6
Rb
K
Atomic Radius (nm)
Overall
Na
Li
Kr
Ar
Ne
H
10
Atomic Number
7
Ionization Energy
 The
amount of energy required to
completely remove an electron from
a gaseous atom.
 Removing one electron makes a +1
ion.
 The energy required is called the first
ionization energy.
8
Ionization Energy
 The
second ionization energy is the
energy required to remove the
second electron.
 Always greater than first IE.
 The third IE is the energy required to
remove a third electron.
 Greater than 1st of 2nd IE.
9
Symbol First
H
He
Li
Be
B
C
N
O
F
Ne
1312
2731
520
900
800
1086
1402
1314
1681
2080
Second
Third
5247
7297
1757
2430
2352
2857
3391
3375
3963
11749
14840
3569
4619
4577
5301
6045
6276
10
Symbol First
H
He
Li
Be
B
C
N
O
F
Ne
1312
2731
520
900
800
1086
1402
1314
1681
2080
Second
5247
7297
1757
2430
2352
2857
3391
3375
3963
Third
11749
14840
3569
4619
4577
5301
6045
6276
11
What determines IE
 The
greater the effective nuclear
charge the greater IE.
 Less distance from nucleus
increases IE
 Filled and half filled sublevels have
lower energy, so removing them
raises the IE.
 Shielding: effective blocking of
nuclear charge weakens the
attraction of valence electrons.
12
Shielding
 The
electron on the
outside energy level
has to look through
all the other energy
levels to see the
nucleus
13
Shielding
 The
electron on the
outside energy level
has to look through
all the other energy
levels to see the
nucleus.
 A second electron
has the same
shielding.
14
Group trends
As
you go down a group first IE
decreases because
The electron is further away.
More effective shielding.
15
Periodic trends
 All
the atoms in the same period have
the same energy level.
 Same shielding.
 Increasing nuclear charge
 So IE generally increases from left to
right.
 Exceptions at full and 1/2 fill sublevels.
16
First Ionization energy
He
 He
has a greater IE than H.
 no shielding
 greater nuclear charge
H
Atomic number
17
First Ionization energy
He  Li has lower IE than H
more shielding
 further away (outweighs greater
nuclear charge)

H
Li
Atomic number
18
First Ionization energy
He
Be has higher IE than Li
 same shielding
 greater nuclear charge
 removing an electron
from a full s sublevel

H
Be
Li
Atomic number
19
First Ionization energy
He
B has lower IE than Be
 same shielding
 greater nuclear charge
 removing an electron from a
Be
partially filled p sublevel

H
B
Li
Atomic number
20
First Ionization energy
He
H
C
Be
B
Li
Atomic number
21
First Ionization energy
He
N
H
C
Be
B
Li
Atomic number
22
First Ionization energy
He
N
H
C O
Be
 Breaks
B
Li
the pattern
because in N the electron
gets removed from a ½
filled p sublevel
Atomic number
23
First Ionization energy
He
N F
H
C O
Be
B
Li
Atomic number
24
First Ionization energy
He
Ne
 Ne
N F
H
C O
Be
B
has a lower IE
than He
 Both have full
valence shells
 Ne has more
shielding
 Greater distance
Li
Atomic number
25
Ne
First Ionization energy
He

N F
Na has a lower
IE than Li
Both are s1
 Na has more
shielding
 Greater distance

H
C O
Be
B
Li
Na
Atomic number
26
Atomic number
27
First Ionization energy
Driving Force
 Full
Energy Levels are very low
energy.
 Noble Gases have full energy levels.
 Atoms behave in ways to achieve
noble gas configuration (become
isoelectronic).
28
nd
2
Ionization Energy
 For
elements that reach a filled or
half filled sublevel by removing 2
electrons 2nd IE is lower than
expected.
 True for s2
 Alkaline earth metals form +2 ions.
29
3rd IE
the same logic s2p1 atoms have
an low 3rd IE.
 Atoms in the aluminum family form +3
ions.
 2nd IE and 3rd IE are always higher than
1st IE!!!
 Using
30
Electron Affinity
The energy change associated with
adding an electron to a gaseous atom.
 Easiest to add to group 17 (7A).
 Gets them to full energy level becomes
stable – releases a large amount of
energy.
 Increase from left to right atoms –metals
are losers so to gain an electron will only
bring about a small increase in stability (if
any) – small energy change.
 Decrease as we go down a group.

31
Ionic Size
 Cations
form by losing electrons.
 Cations are smaller than the atom
they come from.
 Metals form cations.
 Cations of representative elements
have noble gas configuration.
32
Ionic size
 Anions
form by gaining electrons.
 Anions are bigger than the atom they
come from.
 Nonmetals form anions.
 Anions of representative elements
have noble gas configuration.
33
Configuration of Ions
 Ions
always have noble gas
configuration.
 Na is 1s22s22p63s1
 Forms a +1 ion - 1s22s22p6
 Same configuration as neon.
 Metals form ions with the
configuration of the noble gas before
them - they lose electrons.
34
Configuration of Ions
 Non-metals
form ions by gaining
electrons to achieve noble gas
configuration.
 They end up with the configuration of
the noble gas after them.
35
Group trends
 Adding
energy level
 Ions get bigger as
you go down.
Li+1
Na+1
K+1
Rb+1
Cs+1
36
Periodic Trends
 Across
the period nuclear charge
increases so they get smaller.
 Energy level changes between
anions and cations.
Li+1
B+3
Be+2
N-3
O-2
F-1
C+4
37
Radii vs atomic #
atomic radii
ionic radii
51
33
0.21
19
52
37
15
34 35
38
16 17
20
Radii (nonometers)
7
0.16
11
49
37
8
9
12
13
3
31 32 33
14
0.11
15
11
4
5
16
17
34 35
38
20
49
6
7
0.06
50 51 52
19
8
3
9
50
12
31
32
13
14
4
5
6
0.01
3
4
5
6
7
8
9 11 12 13 14 15 16 17 19 20 31 32 33 34 35 37 38 49 50 51 52
Atomic #
38
Size of Isoelectronic ions
 Iso
- same
 Iso electronic ions have the same #
of electrons (same configuration)
 Al+3
Mg+2 Na+1 (Ne) F-1 O-2 and N-3
 all have 10 electrons
 all have the configuration 1s12s22p6
39
Size of Isoelectronic ions
 Positive
ions have more protons than
electrons so they are smaller (strong
attraction for electrons).
Al+3
Na+1
Ne
F-1
O-2
N-3
Mg+2
40
Electronegativity
Electronegativity
 The
tendency for an atom to attract
electrons to itself when it is
chemically combined with another
element.
 How fair it shares.
 Big electronegativity means it pulls
the electron toward it.
 Atoms with large negative electron
affinity have larger electronegativity.
42
Group Trend
 The
further down a group the farther
the electron is away and the more
electrons an atom has – the lower the
electronegativity value.
43
Periodic Trend
 Metals
are at the left end.
 They let their electrons go easily
 Low electronegativity
 At the right end are the nonmetals.
 They want more electrons.
 Try to take them away.
 High electronegativity.
44
Ionization energy, electronegativity
Electron affinity INCREASE
45
Atomic size decreases,
shielding constant
Atomic
size
increases
Shielding
increases
Ionic size decreases
46