Development of the Periodic Table, Periodic property i.e atomic radius , ionisation energy , electron affinity.

1. PERIODICITY
Mr. A. D. Badar (Assistant Professor)
Late Pushpadevi Patil Arts & Science College,
Risod Dist. Washim
Email- badaradinath@gmail.com

2. Development of the Periodic Table
 Mendeleev developed periodic table to group
elements in terms of chemical properties.
 Alkali metals develop +1 charge, alkaline earth metals
+ 2
 Nonmetals usually develop negative charge (1 for
halides, 2 for group 6A, etc.)
 Blank spots where elements should be were observed.
Discovery of elements with correct properties.

3. Periodic Properties
 Periodic law = elements arranged by atomic
number gives physical and chemical properties
varying periodically.
 We will study the following periodic trends:
 Atomic radii
 Ionization energy
 Electron affinity

4. TRENDS IN THE PERIODIC TABLE -
USE OF WHAT WE HAVE JUST
LEARNED
A)Atomic radius - a number of physical
properties of elements are related to the
size of an atom, but with our probability
picture where does an atom end?

5. The atomic radius is ½ the distance
between the 2 nuclei of the adjacent atoms.

6. Atomic Radius
 Atomic radii actually
decrease across a row
in the periodic table.
Due to an increase in
the effective nuclear
charge.
(The more protons you have, the
harder they pull the e- to them.)
 Within each group
(vertical column),
the atomic radius
tends to increase
with the period
number.
Chapter 8-6
Fig. 8.17 Atomic Radii for
Main Group Elements (s,p)

7. Atomic Radius 2
 If positively charged the radius decreases.
(lost e-)
 If the charge is negative the radius increases
(gained e-).
 When substances have the same number of
electrons (isoelectronic), the radius will depend
upon which has the largest number of protons.

8. C) Atomic radius, in general, decreases as
we move from left to right in a row of the
periodic table.
D) Atomic radius increases from top to
bottom in a family or group.
E) These 2 trends are the result of 3
influences on size.

9. 1) As the number of the principal energy
level “n” increases, the size increases,
they extend further from the nucleus
and the covalent radius increases.
(bigger outer level, higher floor in the
motel—the larger the radius)
2) As nuclear charge (number of protons)
increases across a row, the positive charge
on the nucleus increases to electrons are
pulled closer to the nucleus…radius get
smaller.

10. 3) The shielding effect is The attraction for
electrons in the outermost shell by the
nucleus is shielded by electrons in lower
energy levels. *As you gain electrons it
becomes harder to pull in the farther ones.
a) The smaller size of atoms going across a row
can be attributed to minimum shielding.

11. b) Electrons in the same shell are attracted
more strongly as the nuclear charge (# of
protons) increases, because the shielding
effect remains the same.
c) If the shielding effect remains the same, the
Effective Nuclear Charge increases.
d) The ENC is the positive charge that an
electron experiences from the nucleus and is
equal to the nuclear charge minus the
number of shielding electrons.
For example: Li has 3protons in the nucleus, 2e in the 1s orbital (shielding)
and 1e in the 2s orbital. ENC = 3 - 2 = 1. The outermost electron "feels" a
net attraction by the inside of +1.

12. a) Ionization Energy = energy necessary to
remove an electron---is always
endothermic and positive +.
b) M(g) + h  M+ + e.
b) Electron Affinity = energy change upon
the addition of an electron can be either
endothermic or exothermic depending on the
element (for a gaseous atom)
A(g) + e A-
(g)
*An exothermic Energy = - value
Ionization Energy/Electron Affinity

13. IONIZATION ENERGY
 Ionization energy, Ei: minimum energy
required to remove an electron from the
ground state of atom (molecule) in the gas
phase.
 M(g) + h  M+ + e.
 Sign of the ionization energy is always
positive, for example, it requires energy for
ionization to occur.


14. Ionization Energy: Periodic
table Fig. 8.18 Ionization Energy vs atomic Size

15. A(g) + energy  A+ + 1e
A(g) + energy A+ + 1e H = + kJ/mol
2) THIS IS A VERY IMPORTANT
CONCEPT because the chemical properties
of any atom are determined by the
configuration of an atom's valence electrons,
those electrons in the outermost shell.
IONIZATION ENERGY

16. 6) The trend across from left to right is
accounted for by a) the increasing nuclear
charge.

17. The electrons in the outermost shell are
more strongly bound to the nucleus due to
the increasing effective nuclear charge.
a) as we go across a row of the periodic
table energy is larger nuclear charge
becomes larger as the number of protons
in the nucleus of the atom becomes
larger.
W
H
Y?
IONIZATION ENERGY

18. b) With an electron already in the orbital
there is repulsion between the two in the
same orbital and it comes out with less
energy input.
c) The trend from top to bottom of a
column shows a decrease in the FIE which
corresponds to an increase in the atomic
radius.
9) The 2nd, 3rd, and 4th ionization energies
are those required to remove the 2nd, 3rd,
and 4th electrons.

20. 1) Electron affinity is the energy change
which occurs when an electron is accepted
by an atom in the gaseous state.
A(g) + e A-
(g)
2) In contrast to ionization energy, what do
we observe on the following graph of EA's?
Electron Affinity

21. The greater the negative value of the
electron affinity, the greater the tendency of
an atom to accept an electron.
d) A +value indicates that energy must be
absorbed for an atom to gain an electron.
e) left to right on the periodic chart,
general increasing tendency to form
negative ions. However, there are more
exceptions than with Ionization Energy.
Electron Affinity

22. Electron Affinity

23. Electron affinities generally become smaller as
we go down a column of the periodic table for
two reasons.
• First, the electron being added to the atom is
placed in larger orbitals, where it spends
less time near the nucleus of the atom.
• Second, the number of electrons on an atom
increases as we go down a column, so the
force of repulsion between the electron
being added and the electrons already
present on a neutral atom becomes larger.

24. Electron Affinity

25. Electron Affinity

26. ELECTRON AFFINITY
 Electron Affinity, Eea, is the
energy change that occurs when
an isolated atom in the gas phase
gains an electron.
E.g. Cl + e  Cl Eea = 348.6
kJ/mol
 Energy is often released during
the process.
 Magnitude of released energy
indicates the tendency of the
atom to gain an electron.
 From the data in the table the
halogens clearly have a strong
tendency to become negatively
charged
 Inert gases and group I & II
elements have a very small Eea.

27. c) What should you be able to do as a
result of this???
I should be able to give you a list of
elements and you should be able to put
them in order of size from smallest to
largest by just looking at their positions
on the chart.
You should be able to tell me the reasons
why they are smaller or larger.