Transcript Periodic Trends

ATOMIC RADIUS
IONIZATION ENERGY
ELECTRON AFFINITY
ELECTRONEGATIVITY
As you look at the periodic table and focus in on the
elements and their characteristics, you can see there
are noticeable patterns (trends) that go across a
period (horizontal row) on the periodic table.
 These quantitative characteristics that follow distinct
patterns across the periodic table are called periodic
trends.
 Remember…Mendeleev didn’t get all the fame and
fortune for nothing – he was crazy smart!
 The periodic trends that we will be studying are:





Atomic Radius
Ionization Energy
Electron Affinity
Electronegativity

Density

In the second energy level (n=2), if we look from the left of the
chart to the right we first see:
Li




Be
B
C
N
O
F
Ne
Li (metal), then Be (metal), B (metalloid), C (nonmetal), N
(nonmetal), O, F, and Ne (all nonmetals).
This pattern holds true for every energy level. As you read the
periodic table from left to right, first are the metals, followed
by the metalloids, and finally the nonmetals.
The densities of the elements vary in a regular way when plotted
against atomic numbers of the elements.
Metals have the highest density and nonmetals have the lowest
density.

Influenced by

Number of Energy Levels (Size)



Shielding Effect



Inner shell electrons have the ability to block the nuclear pull
(charge on p+) on the outer electrons
This mainly effects group trends.
Nuclear Charge




The distance between the nucleus and the outer electrons
Higher energy levels are further away from the nucleus
 Electrons in these levels are affected less by the electrostatic
pull towards the nucleus.
The number of protons in the nucleus
Higher charge pulls electrons in closer towards the nucleus
This greatly effects period trends because shielding effect is not
an issue across a period.
Electron Configuration

Filled and Half-filled sublevels are more stable, thus causing
exceptions to some trends
 The
electron on the outside
energy level has to look
through all the other energy
levels to see the nucleus
 It
is difficult to measure the radius of an atom
simply because the electron is in constant
random motion about the nucleus.

The electron cloud doesn’t have a definite edge.
 Scientists
took a diatomic molecule and
measured the distance between the two
nuclei.

Cutting the number in half gives the approximate
radius of each of the atoms.
Determining the Atomic Radius of a Nonmetal (Carbon)
Atomic radii in picometers (10-12 m)
Atomic radii in picometers (10-12 m)
Atomic radii in picometers (10-12 m)
Atomic radii in picometers (10-12 m)
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atomic number
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atomic number
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atomic number
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atomic number
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atomic number
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atomic number
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37
140 90
80 77
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1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
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37
140 90
80 77 71
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200
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110
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1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
50
37
140 90
80 77 71 66
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200
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170
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140
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110
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90
80
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0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
50
37
140 90
80 77 71 66 64
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atomic number
50
37
140 90
50
80 77 71 66 64 70
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1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157
80 77 71 66 64 70
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
80 77 71 66 64 70
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
80 77 71 66 64 70
143
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
80 77 71 66 64 70
143 118
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
80 77 71 66 64 70
143 118109
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
80 77 71 66 64 70
143 118109 103
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
80 77 71 66 64 70
143 118109 103 91
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
80 77 71 66 64 70
143 118109 103 91 94
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
196
80 77 71 66 64 70
143 118109 103 91 94
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
196 174
80 77 71 66 64 70
143 118109 103 91 94
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
196 174
80 77 71 66 64 70
143 118109 103 91 94
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
196 174
80 77 71 66 64 70
143 118109 103 91 94
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
196 174
80 77 71 66 64 70
143 118109 103 91 94
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
196 174
80 77 71 66 64 70
143 118109 103 91 94
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
140 90
50
80 77 71 66 64 70
143 118109 103 91 94
157 136
196 174
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
196 174
80 77 71 66 64 70
143 118109 103 91 94
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
196 174
80 77 71 66 64 70
143 118109 103 91 94
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
196 174
80 77 71 66 64 70
143 118109 103 91 94
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
196 174
80 77 71 66 64 70
143 118109 103 91 94
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
196 174
80 77 71 66 64 70
143 118109 103 91 94
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
140 90
157 136
196 174
80 77 71 66 64 70
143 118109 103 91 94
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
37
50
80 77 71 66 64 70
140 90
143 118 109 103 91 94
157 136
196 174
220
210
200
190
180
170
160
150
140
130
120
110
100
90
80
70
60
50
40
30
20
10
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
atomic number
 Radius
shrinks as you go from left to right
because the effective nuclear charge (how much
the electron actually feels a pull towards the
center) increases
 As
you go down the periodic table each new
period gets a new set of shielding electrons and
therefore the electrons are further away from the
nucleus and the radius expands.
Atomic radius tends to…
…decrease from left to right across a row
due to increasing Zeff.
…increase from top to bottom of a column
due to increasing value of n and increasing shielding effect
 Noble/Inert

With a full valence energy level, there is no
electron interaction and the atoms remain far
apart and larger.

Stable configurations tend to “spread out”
 Transition

gases
and Rare-Earth Elements
Only minor changes due to the filling of inner
shell (1 to 2 levels below the valence) and not
the valence energy level.
 Refer
to a periodic table and arrange the
following elements in order of increasing
atomic radius: Br, Se, Te.
34
Se
52
Te
35
Br
Te is larger than Se.
Se is larger than Br.
Br < Se < Te
 Explain
why Fluorine has a smaller atomic
radius than both Oxygen and Chlorine.

Ionic radius- radius of the ion
 Cations are smaller than the atoms from
which they are formed; The nucleus attracts
the remaining electrons more strongly


Anions are larger than the atoms from which
they are formed: The greater number of
electrons repel more strongly
Isoelectronic defines elements that all have the
same number of electrons
Iso - same
 Al+3 Mg+2 Na+1 Ne F-1 O-2 and N-3




all have 10 electrons
all have the configuration 1s22s22p6
For a series of isoelectronic species with the
same electron configuration, the greater the
nuclear charge, the smaller the species
 Why

do atoms unite to form compounds?
To become more stable.
 Consider


Na and Cl
Na: 1s22s22p63s1
Cl: 1s22s22p63s23p5
Na
Na
Cl
 When
the two elements react, the chlorine
removes the outer electron from the sodium
atom, which gives us…
Na+1
-1
Cl
Cl-1
 [Na]+1:
1s22s22p6
The positively charged nucleus is now
attracting fewer electrons (excess protons),
and with the loss of the 3s1 electron, the ion
now has 2 energy levels.
 The sodium ion is smaller than the sodium
atom.

Na+1

Na
Na
The ion now has the same outer level
configuration (isoelectronic) as Ne, which is
very stable.
 [Cl]-1:
1s22s22p63s23p6
The chlorine atom has gained an electron,
excess electrons (17 p+, 18 e-)
 The chloride ion is larger than the chlorine
atom.

Cl-1

Cl
The ion now has the same outer level
configuration (isoelectronic) as Ar, which is
very stable.
 Compared
to its atom, the magnesium ion is
even smaller.
Mg

Mg+2
In losing 2 outer electrons, the unbalanced nuclear
positive charge is larger than the negative charge on
the electron cloud. The cloud shrinks in size.
 The
nonmetals sulfur and chlorine form ions
that are larger then their respective atoms.
(Strong electron-electron repulsion in anions.)
S
S-2
 Metallic
ions on the left and middle of the table
are formed by the loss of electrons. They are
smaller than atoms from which they are formed.
 Nonmetallic ions are located on the right side of
the table. They are formed by the gain of
electrons and are larger than the atoms from
which they are formed.
 The metallic ions have an outer level
configuration that resembles that of the noble gas
at the end of the preceding period.
 Nonmetallic ions have an outer level resembling
that of the noble gas to the right in the same
period.
A. Cation (Positive ion) size
A. When an atom loses electrons to form a cation the ion will have
an excess of protons compared to electrons.
B. As a result the protons draw the remaining electrons in more
closely than in the corresponding atom and the cation is smaller
than the corresponding atom.
C. It is also true to say that when a cation is formed an atom often
loses a complete outside shell of electrons which has the effect of
decreases the size of the cationic species compared to the parent
atom.
B. Anion (Negative ion) size
A. When an atom gains electrons to form an anion the ion will have
an excess of electrons compared to the protons.
B. As a result the electrons are drawn in less than in the
corresponding atom and the anion is larger than the
corresponding atom.
C. Additionally the extra electrons that have been added to form
the anion tend to repel one another. This also has the effect of
slightly enlarging the new anionic species.
 Adding
energy level
 Ions get bigger as you go
down.
Li+1
Na+1
K+1
Rb+1
Cs+1
 Across
the period the nuclear charge increases, so
they ions become smaller.
 Energy
level changes between anions and cations.
Li+1
B+3
Be+2

n=1
C+4
N-3
O-2

n=2
F-1
-
8+
-
-
-
-
-
Oxygen atom
O
2
1s 2s22p4
+2e-
-
-
-
- 9+
- -
- - 10+
- -
- - 11+
- -
Fluorine atom
F
2
1s 2s22p5
Neon atom
Ne
2
1s 2s22p6
Sodium atom
Na
2
1s 2s22p63s1
+1e-
-1e-
-
- - 12+
- Magnesium atom
Mg
2
1s 2s22p63s2
-2e-
8+
-
- 9+
- -
- - 11+
- -
- - 12+
- -
Oxygen ion
O21s22s22p6
Fluorine ion
F11s22s22p6
Sodium ion
Na1+
1s22s22p6
Magnesium ion
Mg2+
1s22s22p6
-
Isoelectronic - all species have the same number of electrons.
p=8
n=8
e = 10
p=9
n=9
e = 10
p = 10
n = 10
e = 10
p = 11
n = 11
e = 10
p = 12
n = 12
e = 10
8+
-
- 9+
- -
- - 10+
- -
- - 11+
- -
- - 12+
- -
Oxygen ion
O21s22s22p6
Fluorine ion
F11s22s22p6
Neon atom
Ne
2
1s 2s22p6
Sodium ion
Na1+
1s22s22p6
Magnesium ion
Mg2+
1s22s22p6
-
-
Can you come up with another isoelectronic series of five elements?
The smallest particle is the one with the most protons.
Na+1 Mg+2 Al+3 N-3 O-2 F-1 Ne

11 p+
Al+3
12 p+
Mg+2
13 p+
7 p+
8 p+
9 p+
10 p+
Na+1 Ne F-1 O-2 N-3
Arrange the following species in order of increasing size.
Ar, K+, Ca2+, S2-, Cl-
 Ionic
size
depends
upon:

Nuclear
charge.

Number
of
electrons
 Cations
are
smaller than
their parent
atoms.

The
outermost
electron is
removed and
repulsions
are reduced.
 Anions
are
larger than
their parent
atoms.

Electrons are
added and
repulsions
are
increased.


Positive ions are smaller than the neutral atoms
from which they are formed
 Electrons are lost from energy levels farthest
from the nucleus
 Remaining electrons pulled closer to the
nucleus
Negative ions are larger than the neutral atoms
from which they are formed
 Electrons are gained, resulting in smaller
effective nuclear charge for the greater
number of electrons
 Repulsive forces between electrons increases
across the periods
Metals
Nonmetals
Group 1
Group 13
Group 17
e
e
Li+
Li
152
F
60
e
Na+
Na
95
Al
143
64
ee
e
136
e
Al3+
50
Cl
Cl-
99
186
181
e
K+
K
227
F-
133
Cations are smaller
than parent atoms
e
Br
114
Anions are larger
than parent atoms
Br195
Atomic size
INCREASES
Atomic
size
Ionic size
INCREASE
Ionic size increases
 Atoms
need to gain or lose electrons to become
stable.
 Use electron configurations to predict oxidation
numbers.


Of course, there are always exceptions to the rule
Look for patterns!!!!
 Have
an oxidation number that is the same
as their roman numeral group number



IA - Na: 1s22s22p63s1
IIA - Mg: 1s22s22p63s2
IIIA - Al: 1s22s22p63s23p1
 IA
Lose 1 eLose 2 eLose 3 e-
– Na+1: 1s22s22p6
 IIA – Mg+2: 1s22s22p6
 IIIA – Al+3: 1s22s22p6
 Have
an oxidation number that is their group
number minus 8



VA - P: 1s22s22p63s23p3
VIA - S: 1s22s22p63s23p4
VIIA - Cl: 1s22s22p63s23p5
 VA
Gain 3 eGain 2 eGain 1 e-
– P-3: 1s22s22p63s23p6
 VIA – S-2: 1s22s22p63s23p6
 VIIA – Cl-1: 1s22s22p63s23p6
 Transition


Metals
Can have oxidation numbers from +2 up to their group
number
Look at Titanium




Group 4
Predicted oxidation Numbers
 +2, +3, +4
Electron Configuration
 [Ar]4s23d2
Orbital Diagram
Remember...lose valence electrons first!!!!
 Ti+2

Lose both 4s electrons
 Ti+3

Lose 2 electrons from the 4s and 1 from the 3d
 Ti+4

Lose 2 electrons from 4s and 2 from 3d
 Because
the +4 configuration is the most stable,
Titanium typically forms the +4 ion.






Group 8
Predicted oxidation Numbers
 +2, +3, +4, +5, +6, +7, +8
Electron Configuration
 [Ar]4s23d6
Orbital Diagram
Fe+2
 Lose both 4s electrons
Fe+3
 Lose both 4s electrons and 1-3d electron
Because the +3 configuration is the most stable, Iron typically forms
the +3 ion
 The
elements under the staircase (that have a
“d” sublevel in their configuration) are
SPECIAL.

They can have multiple oxidation numbers, like
transition metals.
 They
can lose all of their valence electrons,
and/or
 They can lose only their p electrons.

But they must lose all of their p electrons as a
single entity. They cannot lose them one electron
at a time



Thallium
 Electron Configuration
2 14
10
1
 [Xe] 6s 4f 5d 6p
 Possible Oxidation Numbers
 Lose entire valence or Lose only p sublevel.
 Either +3, +1
Lead
 Electron Configuration
2 14
10
2
 [Xe] 6s 4f 5d 6p
 Possible Oxidation Numbers
 Lose entire valence or Lose only p sublevel.
 Either +4, +2
Bismuth
 Electron Configuration
2 14
10
3
 [Xe] 6s 4f 5d 6p
 Possible Oxidation Numbers
 Lose entire valence or Lose only p sublevel.
 Either +5, +3
• Energy is required to remove an electron from an
atom to form a cation.
• This Ionization energy (E) is the amount of energy
needed to remove an electron from the gaseous
atom in its ground state.
• Ionization energy is always positive (E > 0).
• Larger values of E mean that the electron is more
tightly bound to the atom and is harder to remove.

Amount of energy required to remove an electron
from the ground state of a gaseous atom or ion.

First ionization energy is that energy required to
remove first electron.
M(g)  M+(g) + eElement
1st
H
1312.1
He
2372.5
Li
520.3
Be
899.5
B
800.7
C
1086.5
Al
577.6

Amount of energy required to remove an electron
from the ground state of a gaseous atom or ion.



Group 1: easily loses its 1 valence electron



First ionization energy is that energy required to
remove first electron.
M(g)  M+(g) + eSecond ionization energy is that energy required to
remove second electron, etc.
M+(g)  M2+(g) + eLow first ionization energy
Second ionization energy will be very high since it is
“happy” with losing 1 electron.
Group 2: easily lose 2 valence electrons


Low first and second ionization energies
High third ionization energy
e.g. Consider the following Successive IONIZATION ENERGIES for
sodium.
1st
496
2nd
4562
3rd
6912
4th
9543
5th
6th
7th
8th
9th
10th
13353 16610 20114 25490 28933 141360
The large "jump" in the data between the first and second ionization
energies shows that it is relatively easy to remove the first electron
but extremely difficult to remove the second.
It can be assumed that the second electron is in a new shell, is closer
to the nucleus, is more difficult to remove (higher ionization energy)
and therefore the first electron was in the outside shell on its' own,
hence group I. There is another "jump" between the ninth and tenth
ionization energies. This indicates the start of the inner (1st) shell.
 It
requires more energy to remove each successive
electron, because the electron is removed from a
positive ion of increasing charge.
 When
all valence electrons have been removed,
the ionization energy takes a quantum leap.
 for
• I1
• I2
• I3
Mg
= 735 kJ/mole
= 1445 kJ/mole
= 7730 kJ/mole
 The
effective nuclear charge increases as you
remove electrons.
 It
takes much more energy to remove a core
electron than a valence electron because there
is less shielding.
 For
• I1
• I2
•
•
Al
= 580 kJ/mole
= 1815 kJ/mole
I3 = 2740 kJ/mole
I4 = 11,600 kJ/mole
Rb
Sr
Br
Kr
5s1
5s2
4s24p5
4s24p6
Which
 Which
 Which
 Which

has
has
has
has
the
the
the
the
highest IE1?
lowest IE1?
highest IE2?
lowest IE2?
 As
one goes down
a column, less
energy is required
to remove the
first electron.

For atoms in the
same group, Zeff is
essentially the
same, but the
valence electrons
are farther from
the nucleus.
 What

The greater the distance between a nucleus and the
outer electrons of an atom, the LOWER the ionization
energy.


differs among the elements in a group?
Less attraction/interaction
Inner electrons block the attraction of the nucleus for
the outer electrons, and LOWER the I.E.
As a group is descended, the
ionization energy will decrease
• The outer electrons enter
new shells further away
from the nucleus
• More shielding from the
inner electrons
•Valence electrons are held
less strongly
Li
520 kJ
Na
459.5 kJ
K
418.8 kJ
 What



differs among the elements in a period?
Increased nuclear charge = increased attraction
Harder to remove an electron = Higher I.E.
No increased shielding effect
 Watch
out for stability……
1s
1s
1s
2s
2s
2s
2p
2p
2p
An electron from a full or
half-full sublevel requires
additional energy to be
removed.
 Group
Trends: decreases from top to bottom
on the periodic table


Outermost electrons are farther from the effect of the
nuclear charge and therefore easier to remove
Shielding effect increases down the table
 Period


Trends: increases from left to right
Nuclear charge is increasing with no increase in
shielding effect
Outermost electrons are closer to the nucleus
 Refer
to a periodic table and arrange the
following elements in order of increasing
ionization energy: As, Br, Sb.
33
As
51
Sb
35
Br
Sb is larger than As.
As is larger than Br.
Ionization energies:
Sb < As < Br
He
has a greater IE than H.
 same shielding
First Ionization energy
 He
 greater
nuclear charge
H
Atomic number
He

First Ionization energy



H
Li has lower IE than H
more shielding
further away
outweighs greater nuclear
charge
Li
Atomic number
He

First Ionization energy


H
Be has higher IE than Li
same shielding
greater nuclear charge
Be
Li
Atomic number
He


First Ionization energy


H
Be
Li
B has lower IE than Be
same shielding
greater nuclear charge
By removing an electron
we make s orbital half
filled
B
Atomic number
First Ionization energy
He
H
Be
Li
C
B
Atomic number
He
First Ionization energy
N
H
C
Be
Li
B
Atomic number
He
 Breaks
the pattern
because removing an
electron gets to 1/2
filled p orbital
First Ionization energy
N
H
C O
Be
Li
B
Atomic number
He
First Ionization energy
N F
H
C O
Be
Li
B
Atomic number
Ne
He
First Ionization energy
N F
H
C O
Be
Li
 Ne
has a lower IE
than He
 Both are full,
 Ne has more
shielding
 Greater distance
B
Atomic number
Ne
He

Na has a lower
IE than Li

Both are s1
Na has more
shielding
Greater distance
First Ionization energy
N F
H
C O
Be


Li
B
Na
Atomic number


I.E. decreases as you read down the table.
In a period, I.E. increases from left to right
Atomic number
First Ionization energy
Consider the graph below.
1. Why is there a general increase in ionization energy on passing from
Li to Ne?
2. What accounts for the unusual behavior of oxygen?
3. Explain why He has the largest ionization energy of all the noble
gases.
1. Give a brief but complete account of the changes in ionization energy
as the following transitions in the periodic table are made.
•Passing from magnesium to strontium
•Passing from sodium to argon.
2. Consider the following successive ionization energies of elements A &
B.
A
B
1st
513
737
2nd
3rd
4th
5th
6th
7298 11814
1450 7732 10540 13360 17995
• Which group is A in?
• What would be the most likely charge on an ion of B?
• Why does element A only have values for the first three ionization
energies?
A. Electron affinity is defined as the energy change when gaseous atoms
gain an electron to form a gaseous ions.
X(g) + e-  X-(g)
B. the second electron affinity is defined as the energy change
accompanying;
X-(g) + e-  X2-(g)
C. Electron affinities are measured in kJ.
A. They have both positive and negative values depending on the
species being formed.
A. A negative energy change is described as being exothermic
A.If it is relatively easy to add an electron to a neutral atom
then the electron affinity of that element will be negative
B. A positive energy change is described as being endothermic.
A.If it is difficult to add an electron to a certain element,
then work must be done, and the value for electron
affinity is positive.
B. The more negative the value, the higher the E.A.
A. The higher the value, the LOWER the E.A.
In general,
electron affinity
becomes more
exothermic as
you go from left
to right across a
row.
The electron affinity is > 0, so the element must be in Group
IIA or VIIIA.
The dramatic difference in ionization energies is at the third
ionization.
The element is in Group IIA.
There are
again, however,
two
discontinuities
in this trend.
• Group Trends: decreases from top to bottom on the
periodic table
• Outermost electrons are farther from the effect of
the nuclear charge and therefore easier to remove
• Shielding effect increases down the table
• Period Trends: increases from left to right
• Nuclear charge is increasing with no increase in
shielding effect
• Outermost electrons are closer to the nucleus
Li
Be
B
C
N
-60
19
-27
-122
7
O
F
-141 -328
Ne
29

Both E.A. and I.E. are properties of isolated atoms.

They do not involve interaction between/among atoms.

Scientists need a scale to compare the relating
abilities of elements to attract electrons when their
atoms are combined.

ELECTRONEGATIVITY (E.N.) is the tendency for an atom
to attract electrons to itself when it is chemically
combined with another element

How fair it shares.
 Electronegativity is like two atoms playing tug of war in
the playground. In general the big kid wins and the little
kid goes home crying.


Big electronegativity means it pulls the electron toward itself.
Atoms with large negative electron affinity have larger
electronegativity.
 EN
is influenced by the same factors that
affect I.E. and E.A.
 The most active/reactive metals has low E.N.
 Fluorine has the highest EN of all the
elements

Any bond involving F, the electrons will be more
attracted to the Fluorine atom than to any other
atom.
 Why

are there no E.N. for VIIIA elements?
Typically, not involved in bonding
 In
the top right (F) you have a very strong
nucleus with very little shielding. You have a
very strong pull towards F. So strong it can
pull neighboring atom’s electrons towards it.
 In
the bottom left you have an atom with an
enormous amount of shielding and very little
effective nuclear charge making it to the
valence electron. There is very little pull
towards the nucleus.
Electronegativity Trends
(Similar to Ionization Energy)
• Group Trends: decreases from top to bottom on the
periodic table
• Outermost electrons are farther from the effect
of the nuclear charge and therefore more difficult
to attract
• Shielding effect increases down the table
• Period Trends: increases from left to right
• Nuclear charge is increasing with no increase in
shielding effect
• Outermost electrons are closer to the nucleus and
attracted more easily
Ionization energy
Electronegativity
Electron affinity
INCREASE
Ionization energy
Electronegativity
Electron affinity
INCREASE