2. Periodic property
A property whose resemblance shows
among a group of elements when the elements
are arranged according to some particular order.
Periodicity
The recurrence of the resemblance of physical
chemical properties of elements in a particular
group when the elements are arranged
to a given order.
4. In Dimitri Mendeleev period table (1869)
elements were arranged according to atomic
masses .
• The Nucleus had not been discovered
Atomic number is a fundamental property of
the periodic table (it indicates the number of
protons in the nucleus and this determines the
physical and chemical properties.
6. Effective Nuclear Charge (Z*)
Electrons in an atom are attracted by the nucleus but are repelled
by other electrons in the atom.
Effective nuclear charge increase from left to right (Z increases by
unit, Z* increases by 0.65 units); Atomic Radius decreases from
to right.
Effective nuclear charge decreases as you go down in the group;
increase in atomic size reduces effective attraction on the valence
electrons.
Effective nuclear charge increases with increase in atomic
7. Effective Nuclear Charge:
Effective nuclear charge slowly
increases down the group.
Effective nuclear charge increases across the
period.
H 1.0
Li 1.3
Na 2.2
K 2.2
Rb 2.2
Cs 2.2
Li Be B C N P F Ne
1.3 1.95 2.6 3.3 3.9 4.6 5.2 5.9
8. Ionization Energy (IE)
• The minimum energy required to remove an electron
from a gas phase atom.
• IE decreases as the size of the atom increases.
• IE increases with increase in nuclear charge.
• IE2 > IE1 ; Energy for removing the second electron
(IE2). Nuclear pull increases upon removal of the
first electron.
• Nobel gases have the highest IE because of the octet
electronic structure.
9. The principal quantum number of the
orbital holding the outermost electron
becomes larger as we go down a
column of the periodic table.
As the number of protons in the
nucleus increases, the electrons in
smaller shells and subshells tend to
screen the outermost electron from
some of the force of attraction of the
nucleus.
The electron removed first is further
from the nucleus, needs less energy to
remove this electron from the atom.
10. Electronegativity:
The ability of an atom to attract electrons to
itself.
• Affinity is caused by the strong desire for the atom to fill its outer most
shell.
• Metals tend to be electropositive and non-metal tend to be electronegative.
• Electronegativity increases as you up a group in the periodic table.
• Electronegativity increases to the right of the period in the periodic table.
• Fluorine is the most electronegative element (4.0); Noble gases are not
electronegative and Francium the least electronegative (0.7).
• Electropositivity, on the other hand is the element’s ability to donate
electrons.
11. Electronegativity:
Across the period:
• Number of shells remains the same.
• Atomic Radius decreases.
• Nuclear force of attraction on the valence electron increases.
• Z increases, Z* increases.
Down the group:
• Number of shells increase.
• Atomic Radius increases.
• Nuclear force of attraction on the valence electron decreases.
• Z increases, Z* nearly remains constant.
12. Group 1 2 3 4 5 6 7
Element Li Be B C N O F
Electronegativity 1.0 1.6 2.0 2.6 3.0 3.4 4.0
Element Na Mg Al Si P S CL
Electronegativity 0.9 1.3 1.6 1.9 2.2 2.6 3.2
Low High
Electronegativity
Low
Electronegativity difference
X-Y (0-0.4 Covalent Bond); X-Y (0.5-1.7 Polar-Covalent); X-Y (>1.7 Ionic Bond)
13. Electron Affinity (EA):
The amount of energy released when a neutral atom in
gaseous state gains an electron to form an anion.
• The greater the energy released in the process of taking
up the extra electron, greater is the EA
• The EA of an atom measures the tightness with which it
binds an additional electron to itself.
• Electron Affinity increases to the right of a period
(increased attraction, smaller radii) and decreases down
the group (less attraction of outer most electrons).
