Periodic Relationships Among
the Elements
Chapter 8
Chapter 8 Topics
Periodic Relationships
•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
Dr. Ali Bumajdad
History of periodic table and periodic trends
•Lother Meyer (German, 1830-1895) and Ivanonich Mendeleev
(Russian, 1834-1907) 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
Atomic mass
Not atomic number
Mendeleev’s periodic Table
Corrected (they thought its atomic mass is 76)
When the Elements Were Discovered
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.
Classification of the Elements
•Atomic radius
•Ionization energy
•Electron affinity
•Electronegativity
•Atomic radius
•Ionization energy
•Electron affinity
•Electronegativity
-Electron is added on
different principle
quantum number
Effective nuclear
+ve charge not much
change due to the
shielding effect
 Size increase
-Electron is added on same principle
quantum number
Effective nuclear +ve charge
increase due to no shielding effect
 Size decrease
•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
•Energy required to remove e •Energy released or absorbed
distance between two
from an isolated gaseous atom, when an electron is added to
neighboring atoms in their
ion or molecules (kJ/mol)
an isolated gaseous atom,
solid state
ion (kJ/mol)
Na(g) + E
Na+ + 1e Cl(g) + 1e
Cl- + E
•As you go down a group
Release =exothermic
the atom size increase
• Endothermic process (+ve E)
(Zeff is approx. constant because e is attracted to nucleus Cl-(g) + 1e + E
Cl2or even decrease and e
• As size increase it is easier to Absorbed = endothermic
placed in a new energy
remove e
level)
(1) Atomic size
•As you go left to right
across a period atom get
(2) Ionization energy (3) Electron affinity
• All element (except H) have • All element have more than
more than 1 ionization energy 1 Electron Afffinity
smaller (Zeff increase and e
placed in the same energy
•1st Eionization < 2ed Eionization <… • 1st Eaffinity usually exothermic
Level.
because e placed into
(removing e from neutral atom
•For transition elements and is easier than from +ve ion) environment where it
experience attraction to nucleu
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 • 2st E
affinity endothermic
alkali metal is much more because the second e added
than the first one?
to a negatively charge ion
It has a noble gas
electron configuration
(1) Atomic size
Q) Why Zr have similar size
as Hf?
(2) Ionization energy (3) Electron affinity
Q) Why the 3ed Eionization for
• Some Irregularities a cross
earth alkali metal is the
a period
highest change?
A raw of Lanthanide elements It has a noble gas
electron configuration
•For ions:
1
P Eaffinity < As affinity
O Eaffinity < S Eaffinity
+ve ion < -ve ion
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
(1) Atomic size
(2) Ionization energy (3) Electron affinity
• Some Irregularities a cross a •Because the e in c inter a
vacant 2p orbital while for N e
period
must be placed in an already
Be Eionization > B Eionization
occupied orbital
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
Effective Nuclear Charge (Zeff)
increasing Zeff
increasing Zeff
Adding electron to the same principal quantum number
Adding electron on different principal quantum number
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
(a) Ne
(b) Mg (the first and second e’s are valence electrons)
Be2+ < Mg2+ < Ca2+ < Sr2+
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
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]
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
Atoms gain electrons
so that anion has a
noble-gas outer
electron configuration.
Atoms lose electrons so that
cation has a noble-gas outer
electron configuration.
H 1s1
H- 1s2 or [He]
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]
-1
-2
-3
+3
+2
+1
Cations and Anions Of Representative Elements
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
Fe2+: [Ar]4s03d6 or [Ar]3d6
Fe3+: [Ar]4s03d5 or [Ar]3d5
Mn:
[Ar]4s23d5
Mn2+: [Ar]4s03d5 or [Ar]3d5
Ions Sizes (isoelectronic ions)
•Isoelectronic ions: ions contain the same number of electrons
Na+: [Ne]
Al3+: [Ne]
O2-: 1s22s22p6 or [Ne]
F-: 1s22s22p6 or [Ne]
N3-: 1s22s22p6 or [Ne]
Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
No. proton 11 13
9 8
7
10
The ion with less protons is bigger
N3- > O2- > F- > Ne > Na+
> Al3+
Dr. Ali Bumajdad
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:
a) Li+, Na+, K+, Rb+, Cs+
b) Ba2+, Cs+, I-, Te2-
Cation is always smaller than atom from
which it is formed.
Anion is always larger than atom from
which it is formed.
Dr. Ali Bumajdad
The Radii (in pm) of Ions of Familiar Elements
Sa Ex. : Predict the trend in radius for :
Be2+ ,Mg2+ ,Ca2+ ,Sr2+
Group 1A Elements (ns1, n  2)
Group 2A Elements (ns2, n  2)
Group 3A Elements (ns2np1, n  2)
Group 4A Elements (ns2np2, n  2)
Group 5A Elements (ns2np3, n  2)
Group 6A Elements (ns2np4, n  2)
Group 7A Elements (ns2np5, n  2)
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
Properties of Oxides Across a Period
basic
acidic