1. Periodic Relationships Among the Elements
Chapter 8

2. Periodic Relationships
Chapter 8 Topics
History of periodic table and periodic trends
Periodic Classification of element
Periodic trend of Physical properties
(Atomic size, Ionization energy, Electron
Affinity, and Electronegativity )
Ions electron configuration
Ions Sizes (isoelectronic ions)
Group elements and its oxides
Periodic Relationships
Dr. Ali Bumajdad

3. History of periodic table and periodic trends
Lother Meyer (German, ) and Ivanonich Mendeleev
(Russian, ) put the form of periodic table
Mendeleev given most credit because he predict the existence
and properties of unknown elements and he corrected several
values for atomic masses

4. Mendeleev’s periodic Table
Atomic mass
Not atomic number
Corrected (they thought its atomic mass is 76)

5. When the Elements Were Discovered

6. The current periodic table is different than Mendeleev by:
1) Element order by atomic number not by atomic mass
2) Contain many more elements discovered after Mendeleev.

7. Classification of the Elements

8. Atomic radius
Ionization energy
Electron affinity
Electronegativity
Atomic radius
Ionization energy
Electron affinity
Electronegativity

9. -Electron is added on same principle quantum number
Effective nuclear +ve charge increase due to no shielding effect
 Size decrease
-Electron is added on
different principle
quantum number
Effective nuclear +ve charge not much change due to the shielding effect
 Size increase

10. (1) Atomic size (2) Ionization energy (3) Electron affinity
Effective nuclear charge (Zeff): the “positive charge” felt by an electron
(1) Atomic size
(2) Ionization energy
(3) Electron affinity
Atomic radius is half the
distance between two
neighboring atoms in their
solid state
Energy required to remove e
from an isolated gaseous atom,
ion or molecules (kJ/mol)
Energy released or absorbed
when an electron is added to
an isolated gaseous atom,
ion (kJ/mol)
Na(g) + E Na+ + 1e
Cl(g) + 1e Cl- + E
As you go down a group
the atom size increase
(Zeff is approx. constant
or even decrease and e
placed in a new energy
level)
Release =exothermic
Endothermic process (+ve E)
because e is attracted to nucleus
Cl-(g) + 1e + E Cl2-
As size increase it is easier to
remove e
Absorbed = endothermic

11. (1) Atomic size (2) Ionization energy (3) Electron affinity
As you go left to right
across a period atom get
smaller (Zeff increase and e
placed in the same energy
Level.
All element (except H) have
more than 1 ionization energy
All element have more than
1 Electron Afffinity
1st Eionization < 2ed Eionization <…
(removing e from neutral atom
is easier than from +ve ion)
1st Eaffinity usually exothermic
because e placed into
environment where it
experience attraction to nucleus
For transition elements and
inner transition elements,
the last e,s are added onto
an inner shell (3d) not the
valence shell. This cause
the Zeff to increase but not
much (there is a slight
shielding effect) atomic
size decrease but slowly.
Q) Why the 2nd Eionization for
alkali metal is much more
than the first one?
It has a noble gas
electron configuration
2st Eaffinity endothermic
because the second e added
to a negatively charge ion

12. (1) Atomic size (2) Ionization energy (3) Electron affinity 1 Why? 2
Q) Why the 3ed Eionization for
earth alkali metal is the
highest change?
It has a noble gas
electron configuration
Q) Why Zr have similar size
as Hf?
A raw of Lanthanide elements
Some Irregularities a cross
a period 1
P Eaffinity < As affinity
For ions:
+ve ion < -ve ion
O Eaffinity < S Eaffinity
Why?
Because e are crowding in shell 2 more than in shell 3
more repulsion in shell2
less energy evolve
less Eaffinity 2
C has -ve Eaffinity while N has +ve Eaffinity

13. (1) Atomic size (2) Ionization energy (3) Electron affinity Why?
Some Irregularities a cross a
period
Because the e in c inter a vacant 2p orbital while for N e must be placed in an already occupied orbital
Be Eionization > B Eionization
N Eionization > O Eionization
P Eionization > S Eionization
Why?
For Be and B, it is easier to remove e from 2p1 than from 2s2 because it is further apart from nucleus
For N and O, it is easier to remove e when there is e-e repulsion in an orbital

14. Effective Nuclear Charge (Zeff)
increasing Zeff
increasing Zeff

15. Adding electron to the same principal quantum number
Adding electron on different principal quantum number

16. Sa Ex. : Eionization of P = 1060kJ/mol and for S = 1005 kJ/mol, explain this irregularity.
The valence electron configuration for S contains one filled orbital
and hence there is e-e repulsion and hence easier to be removed
Sa Ex. : Among the following elements (Ne, Na, Mg)
(a) Which atom has the largest 1st ionization
(b) Which atom has the smallest 2nd ionization energy Ne
Mg (the first and second e’s are valence electrons)
Be2+ < Mg2+ < Ca2+ < Sr2+

17. The Alkali Metals Li Cs
More reactive (why)
Easier to loss electron (less effective nuclear charge)
In aqueous system less reducing ability (why)
Difficult to be solubilized in water due to its large
size (less effective nuclear charge), less polar

18. Ions electron configuration
When two nonmetals react they share electrons, so that the valence
electron configuration completes for both atoms
When nonmetal react with metal, the ions form so that the valence electron
configuration of the nonmetal achieves the electron configuration of the next
noble gas atom and the valence orbital of the metal are emptied
Na [Ne]3s1
Na+ [Ne]
Atoms lose electrons so that cation has a noble-gas outer electron configuration.
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
H 1s1
H- 1s2 or [He]
Atoms gain electrons so that anion has a noble-gas outer electron configuration.
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]

19. Cations and Anions Of Representative Elements+1 +2 +3 -3 -2 -1

20. Electron Configurations of Cations of Transition Metals
When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals.
Fe: [Ar]4s23d6
Mn: [Ar]4s23d5
Fe2+: [Ar]4s03d6 or [Ar]3d6
Mn2+: [Ar]4s03d5 or [Ar]3d5
Fe3+: [Ar]4s03d5 or [Ar]3d5

21. Ions Sizes (isoelectronic ions)
Isoelectronic ions: ions contain the same number of electrons
Na+: [Ne]
Al3+: [Ne]
F-: 1s22s22p6 or [Ne]
O2-: 1s22s22p6 or [Ne]
N3-: 1s22s22p6 or [Ne]
Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
No. proton
The ion with less protons is bigger
N3- > O2- > F- > Ne > Na+ > Al3+
Dr. Ali Bumajdad

22. Se2- > Br- > Rb+ > Sr2+
Q) What neutral atom is isoelectronic with H- ?
H-: 1s2
same electron configuration as He
Sa. Ex. : Arrange the ions in order of decreasing size ?
Se2-, Br-, Rb+, Sr2+
Se2- > Br- > Rb+ > Sr2+
Sa. Ex. : Choose the largest ion in each of the following
Groups:
Li+, Na+, K+, Rb+, Cs+
Ba2+, Cs+, I-, Te2-

23. Cation is always smaller than atom from which it is formed.
Anion is always larger than atom from which it is formed.

24. Dr. Ali Bumajdad

25. The Radii (in pm) of Ions of Familiar Elements

26. Sa Ex. : Predict the trend in radius for :
Be2+ ,Mg2+ ,Ca2+ ,Sr2+

27. Group 1A Elements (ns1, n  2)

28. Group 2A Elements (ns2, n  2)

29. Group 3A Elements (ns2np1, n  2)

30. Group 4A Elements (ns2np2, n  2)

31. Group 5A Elements (ns2np3, n  2)

32. Group 6A Elements (ns2np4, n  2)

33. Group 7A Elements (ns2np5, n  2)

34. Group 8A Elements (ns2np6, n  2)
Completely filled ns and np subshells. Highest ionization energy of all elements.
No tendency to accept extra electrons.
8.6

35. Properties of Oxides Across a Period
basic
acidic