Students ScoreBooster Series Videos
WAEC, SSCE, GCE, JAMB (UTME),
NECO and NABTEB
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Chemistry
The Periodic Table cont’d
Presented by
A.A.S Lateef
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The periodic table
• Learning objectives, at the end of this lecture,
 Students should be able to explain some properties of
periodic table and the trend shown across the periods as
well as down the groups
 Students should be able to predict the chemical behavior
based on these properties.
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The Periodic Table
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Properties under consideration
•
•
•
•
Ionization Energy
Atomic radius (and ionic radius)
Electron affinity
Electronegativity (and electropositivity)
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Ionization energy
 This is the energy required to remove the most loosely-held electron
from the outermost shell of the gaseous atom. Its unit is kJ/mol.
 Down the group, ionization energy decreases.
 Each element has more occupied energy levels than the one above it
has.
 The outermost electrons are farthest from the nucleus in elements
near the bottom of a group.
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 As you move down a group, each successive element contains
more electrons in the energy levels between the nucleus and
the outermost electrons.
 Reason: Electron shielding is the reduction of the attractive force
between a positively charged nucleus and its outermost electrons
due to the cancellation of some of the positive charge by the
negative charges of the inner electrons.
The first ionization energy is the energy which is required when a gaseous atom/ion loses
an electron to form a gaseous +1 valence ion.
The energy which is required for a gaseous +1 valence ion to loose an electron to form a
gaseous +2 valence ion, is called the second ionization energy of an element.
The second ionization energy is generally higher than the first.
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 Ionization energy tends to increase as you move from left to right
across a period.
 From one element to the next in a period, the number of protons
and the number of electrons increase by one each.
 The additional proton increases the nuclear charge.
 A higher nuclear charge more strongly attracts the outer electrons
in the same energy level, but the electron-shielding effect from
inner-level electrons remains the same.
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Atomic radius
 The exact size of an atom is hard to determine.
 The volume the electrons occupy is thought of as an electron cloud, with no
clear-cut edge.
 In addition, the physical and chemical state of an atom can change the size
of an electron cloud.
 One method for calculating the size of an atom involves calculating the
bond radius, which is half the distance from center to center of two like
atoms that are bonded together.
 The bond radius can change slightly depending on what atoms are
involved.
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Atomic Radius Increases as You Move Down a Group
 As you proceed from one element down to the next in a group, another
principal energy level is filled.
 The addition of another level of electrons increases the size, or atomic
radius, of an atom.
 Because of electron shielding, the effective nuclear charge acting on the
outer electrons is almost constant as you move down a group,
regardless of the energy level in which the outer electrons are located.
 As you move from left to right across a period, each atom has one more
proton and one more electron than the atom before it has.
 All additional electrons go into the same principal energy level—no
electrons are being added to the inner levels.
 Electron shielding does not play a role as you move across a period.
 As the nuclear charge increases across a period, the effective nuclear
charge acting on the outer electrons also increases.
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Ionic radii
 This is the average distance from the nucleus to the outermost electron shell of
an ion.
 For metals, it is generally lower than atomic radius because when a metal loses
outermost electron, the number of shell reduces.
 However for non-metals, it is greater than atomic radii as the additional electron
added reduces the effective nuclear charge and thus, the size increases.
 Down the group, it has the same trend like the atomic radii also across the
period.
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Electronegativity
 Not all atoms in a compound share electrons equally.
 Knowing how strongly each atom attracts bonding electrons can
help us explain the physical and chemical properties of the
compound.
 Electronegativity is a measure of an ability of an atom in a chemical
compounds to attract electrons.
 The atom with the higher electronegativity value will pull electron
more strongly than the other atom will.
 Fluorine is the element whose atoms most strongly attract shared
electrons in compound. Pauling arbitrarily gave fluorine an
electronegativity value of 4.0
 Values for the other elements were calculated in relation to this
value
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Electronegativity (cont’d)
• Down the group, electronegativity decreases. Although, the more proton an
atom has, the more strongly it would attract an electron. However electron
shielding plays more significant role here.
• Across the period, electronegativity value increases.
• As you proceed across a period, each atom has one more proton and one
more electron—in the same principal energy level—than the atom before it
has.
• Electron shielding does not change across the period because no electrons
are being added to inner energy levels.
• The effective nuclear charge increases across a period.
 As this increases, electrons are added much more strongly, resulting in
an increase in electronegativity.
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Electropositivity
 There is variation in the way elements lose electrons.
 Physical and chemical properties of an element can be predicted
from the knowledge of how electrons are lost from an atom
 Thus, electro positivity is a measure of an ability of an atom in a
chemical compounds to lose electrons.
 It is the opposite of electronegativity
 The atom with the higher electropositivity value will lose electron
more readily than the other atom will.
 Francium is the most electropositive element.
 In metals, electropositivity value indicates reactivity. A more
electropositive metal is more reactive than the other with lower
value.
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Electropositivity (cont’d)
• Down the group, electro positivity increases. Electron shielding is more
pronounced here despite that nuclear charge increases.
• Across the period, electropositivity value decreases.
• As you proceed across a period, each atom has one more proton and one
more electron—in the same principal energy level—than the atom before it
has.
• Electron shielding does not change across the period because no electrons
are being added to inner energy levels.
• The effective nuclear charge increases across a period.
 As this increases, electrons are added much more strongly, resulting in
an decrease in electropositivity.
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Electron affinity
 The energy change that occurs when a neutral atom (in gaseous
state) gains an electron is called the atom’s electron affinity.
 This property of an atom is different from electronegativity.
 The electron affinity tends to decrease as you move down a group
because of the increasing effect of electron shielding.
 Electron affinity tends to increase as you move across a period
because of the increasing nuclear charge.
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ElEctron affinity (cont’d)
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Points to note:
 Down any group, the following properties increase;
 Atomic radius , Ionic radius and Electropositivity
While the following decrease;
 Electronegativity, Ionization energy and Electron affinity
 Across any group, what is being experienced is opposite of the above.
 Electron shielding effect and atomic/ionic radius are important in
explaining the trends in the properties above.
 First ionization energy is generally lower than second ionization
energy. (it becomes more difficult to remove the second electron
because of increased nuclear charge).
 Second electron affinity is always positive and in absolute terms, lower
than the first.
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Examples
•
JAMB 1985, Q1
JAMB 1986, Q8
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Examples
•
Option D
JAMB 1985, Q1
Option B
JAMB 1986, Q8
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JAMB 1986, Q9
Examples (cont’d)
JAMB 1988, Q10
JAMB 1988, Q12
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JAMB 1986, Q9
Examples (cont’d)
Ionization energy,
electronegativity and
electron affinity
JAMB 1988, Q10
Option D: Atomic radius
JAMB 1988, Q12
Option C: Electron affinity
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Exercise
JAMB 1990,Q11
JAMB 1991,Q13
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Exercise
JAMB 1990,Q11
Option D, electron affinity
[energy involved when electron
is added to gaseous atom].
JAMB 1991,Q13
Option B:
[electronegativity increases
across the period from metal to
non-metal]
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Exercises
JAMB 1998, Q10
Post UTME
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Exercises
JAMB 1998, Q10
Option A
[fewer shells indicates higher
attraction from the nucleus]
Post UTME
Option B
[azimuthal quantum number]
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