1) The document discusses trends in ionization potential and electron affinity across the periodic table. Ionization potential generally increases from left to right in a period as nuclear charge increases and atomic radius decreases. It also decreases down a group as atomic size increases. 2) Electron affinity generally increases from left to right as nuclear charge increases but decreases down a group as atomic size increases. Elements in the lower left of the periodic table tend to have lower ionization potentials and are more metallic. 3) Factors that influence ionization potential and electron affinity include effective nuclear charge, atomic size, shielding effects, and stability of electron configurations.

1. Ionization Potential
And Electron Affinity
Subject: Chemistry
Course code: Chem-101
Submitted By: Aqsa Manzoor
Discipline: Zoology
Semester: I
Roll No: 2180
Submitted To: Mrs. Uzma Imtiaz

2. Periodic TrendsPeriodic Trends
 Ionization Energy, Electron
Affinity and Atomic Radii.
The valence electron structure
of atoms can be used to
explain various properties of
atoms.
In general, properties correlate
down a group of elements.
In this Assignment I will
discuss the trends of Ionization
potential and electron affinity
in the periodic table.

3. Ionization EnergyIonization Energy
 Ionization : The process of removing
an electron from an isolated atom to
form a positive ion.
First ionization energy (I1): amount
of energy required to remove the
most loosely bound electron from an
isolated gaseous atom to form a
cation.
Second ionization energy (I2):
amount of energy required to remove
a second electron from the gaseous
monopositive cation to form
dipositive cation.

4.  Ionization energies are usually
expressed in electron volts (eV)
per atom or in kilojoules per mol
(kJ/mol)
1eV/atom=96.48 kJ/mol
 Value of each ionization energy
will increase with each removed
electron, since the attractive
influence of the nucleus
increases and will and will
require more energy for the
removal of an electron from
more positive charges.
 Ionization energies measure hoe
tightly
electrons are bound to atoms.
Low energies indicate ease of
removal of electrons and vice
versa.

5. Factors affecting the magnitude
of Ionization Potential:
1. Effective nuclear charge
2. Atomic size i.e. atomic radius
3. Principle quantum number
4. Shielding effect
5. Half filled and completely filled orbitals
6. Nature of orbitals
7. The extent of penetration of valence
electrons
Effective Nuclear Charge:
Greater the
magnitude of effective nuclear charge,
higher is the amount of energy needed
to remove the outermost shell
electron. Thus with the increase of the
magnitude of effective nuclear charge, the
magnitude of ionization potential also
increases. The effective nuclear charge
increases from left to right in a period.

6. Atomic size:
Greater is the atomic size of
an atom, more far is the outermost
shell electron from the nucleus and
hence lesser will be the force of
attraction exerted by the nucleus on
the outermost shell electron. Thus
higher the value of atomic radius of an
atom, lower will be the ionization
energy.
Principal Quantum Number (n):
Greater is the
value of n for the valence shell
electron of an atom, further away
this electron will be from the nucleus
and hence lesser will be the force of
attraction exerted by the nucleus on
it so lesser energy will be required to
remove the valence shell electron.
Thus with the increase ot the principal
quantum number of the orbital from
which the electron is to be removed, the
magnitude of ionization potential
decreases.

7. Shielding Affect:
The magnitude of
shielding effect determines the
magnitude of the force of attraction
between the nucleus and the valence-
shell electron. Greater is the magnitude
of shielding effect working on the
valence shell electron. Thus with the
increase of shielding effect ionization
potential increases.
Half-filled and completely-filled
orbitals:
According to Hund’s rule, half-
filled (ns1, np3, nd5) or completely-
filled (ns2. ns6, nd10) orbitals are
comparatively more stable and hence
more energy is needed to remove an
electron from such orbitals. Thus the
ionization potential of an atom having half-
filled or completely-filled orbitals in its
electronic configuration is relatively higher
than that expected normally from its
position in the periodic table .

8. Nature Of Orbitals:
The nature of
orbitals of the valence-shell from which
the electron is to be removed also
influences the magnitude of ionization
potential. The relative order of energy
of s, p, d and f orbitals of a given nth
shell is as:
ns < np < nd < nf
This order clearly shows that to remove
an electron from f-orbital will be easiest
while to remove the same from s-orbital
will be the most difficult.
The extent of penetration of
valence electrons:
The degree of
penetration of valence electrons in a
given principal energy level decreases
in the order s>p>d>f, since ns electron is
more tightly bound than any np
electron, which in turn is more tightly
bound than any nd electron etc.

9. Trends in Ionization
Potential:
 Ionization energy generally
increases from left to right in a
period because of the increase in
nuclear charge and decrease in
atomic radius.
 Ionization energies generally
decrease down a group due to the
shielding effect and increase in
atomic size.
 Departures from these trends can
usually be traced to repulsion
between electrons, particularly
electrons occupying the same
orbitals.

10.  It requires more energy to
remove each successive
electron.
When all valence electrons have
been removed, the ionization
energy takes a quantum leap.

11. Trends in FirstTrends in First
Ionization EnergiesIonization Energies
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.

12. Generally, as one
goes across a
row, it gets
harder to remove
an electron.
As you go from
left to right, Zeff
increases.
The second
occurs between
Groups VA and
VIA.
Electron removed
comes from
doubly occupied
orbital.
Repulsion from
other electron in
orbital helps in its
removal.

13. Elements in the lower left of
the periodic table tend to
have lower ionization
energies than those in the
upper right.
These are the elements in the
s block, d block, f block and
the lower left of the p block -
metallic solids

14. Electron AffinityElectron Affinity
Electron Affinity: energy that occurs when
an electron is added to a neutral atom
in the gaseous state to form a negative
ion.
X(g) + e-
X-
(g)
∆E = electron attachment energy
Electron affinity = - ∆E (electron
attachment)
Second Electron Affinity: Second electron
affinity of n element M(g) is defined as
the amount of energy required to add
one more electron to its mononegative
anion, to form dinegative anion.
The addition of second electron to uni-
negative ion is an endothermic process.

15. Factors Affecting Electron
Affinity:
Nuclear charge:
More the nuclear charge
of the atom more strongly will it attract
additional electron. Therefore, electron
affinity increases as the nuclear charge
increases.
Atomic size:
The smaller the size of atom
smaller will be the distance between
the extra electron and the nucleus.
Therefore, electrostatic force of attraction
will be more and the electron affinity will
be higher.
Electronic Configuration:
Atoms having
stable electronic configuration (i.e.
those having completely filled or half
filled outer orbitals) do not show much
tendency to add extra electron, so have
either zero or very low electron
affinities.

16. Periodic Trends:
Electron affinity values generally increase
on moving left to right in a period.
 Electron affinities undergo a general
decrease down a group.
Electron affinity is usually exothermic.
Electron affinities can be obtained by using
Born-Haber cycle.

17.  In general,
electron
affinity
becomes more
exothermic as
you go from
left to right
across a row.
 The first
occurs between
Groups IA and
IIA.
◦ Added
electron must
go in p-
orbital, not s-
orbital.
◦ Electron is
farther from
nucleus and
feels
repulsion
from s-
electrons.

18. Main Group ElementsMain Group Elements
(s and p)(s and p)
 An s-block element has low
ionization energy; outer electrons
can easily be lost.
 Group 1 form +1 ions;
Group 2 form +2 ions
An s block element is likely to be a
reactive metal
 p- block
Elements on the left have low
enough ionization energies to be
metallic, but higher than the s-
block elements and so are less
reactive
Elements on the right have high
electron affinities (tend to gain
electrons to form closed shell ions)