1) Atomic radius decreases across a period as atomic number increases, leading to higher effective nuclear charge and stronger attraction of the nucleus. Atomic radius increases down a group as shielding electrons increase. 2) Ionization energy increases across a period as effective nuclear charge increases, making it harder to remove electrons. Ionization energy decreases down a group as shielding increases and atomic radius gets larger. 3) Electronegativity increases across a period for the same reasons as ionization energy, and decreases down a group as shielding increases and attraction decreases.

1. Periodic Table
Trends in periodic atomic properties
Sharifah Mona Abdul Aziz Abdullah
Abdul Al-Hafiz Ismail
Centre for Pre-University Studies
Universiti Malaysia Sarawak
This OpenCourseWare@UNIMAS and its related course materials are licensed under
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2. Atomic radius
Atomic radius is defined as one-half of the
average distance between the two covalently
bonded atoms.
1.99 Å
Atomic radius of chlorine

3. Factor affecting the size of atoms:
a) Effective nuclear charge
The charge felt by the valence electrons after
you have taken into account the number of
shielding electrons that surround the nucleus.
Atomic radius

4. Atomic radius
b) Shielding electron/effect
Shielding electrons are the electrons in the
energy levels between the nucleus and the
valence electrons
They are called "shielding" electrons because
they "shield" the valence electrons from the
force of attraction exerted by the positive charge
in the nucleus.

5. Atomic radius
Trend in atomic radius across a period and down the
group in periodic table:
Atomic radius decreases
Atomicradiusincreases

6. Atomic radius
Decrease across a period
• Proton number increases
• Effective nuclear charge increases
• Stronger attraction between nucleus and valence electron
• Thus, atomic radius decreases
Increases down a group
• Energy level increases
• Shielding effect increases
• Weaker attraction between nucleus and valence electrons
• Thus, atomic radius increases.

7. Example:
Arrange the following atoms in order of
increasing atomic size.
Na, K, Cl, Ba, B, C, Ag
Atomic radius

8. Atomic radius
Solution:
Ba is at Period 6, the lowest position among all the elements, should
be the largest atom.
Next should be Ag (Period 5) and then K (Period 4)
Both Na and Cl are at Period 3, but because of Na is at the left side of
the table, Na is larger than Cl.
Same goes between B and C at Period 2.
Therefore,
C < B < Cl < Na < K < Ag < Ba

9. Ionic Radius
Positive ion always has a smaller ionic radius
than the original atom. Why?
The loss of electron(s) means that the remaining
electrons each have a greater share of the positive
charge of the nucleus so are more tightly bound. When
the ion is formed a whole electron shell is usually lost.

10. Ionic Radius
Negative ion has a larger radius than that of the
original atom. Why?
Even though the additional electrons are in the same
shell as existing electrons, the addition of extra
negative charge means that the electrons are less
tightly bound to the nucleus and so the radius is larger.

11. Ionic Radius
Neutral atoms or ions that have same number of electrons and
the same electronic configuration are said to be isoelectronic.
Examples:
Na atom (1s2 2s2 2p6 3s1) contain one extra shell compare to Na+
ions (1s2 2s2 2p6)
F- ion (1s2 2s2 2p6) contain an extra negative electron compare to
F atom (1s2 2s2 2p5)
Therefore, Na+ and F- are isoelectronic.

12. Ionization Energy (IE)
Ionisation energy is the energy required to
remove the valence electron from an atomic
species.
First IE can be represented by an equation as:
M(g)  M+(g) + e-
Second IE can be represented by an equation as:
M+ (g)  M2+(g) + e-

13. Ionization Energy (IE)
Factor affecting the magnitude of ionisation
energy:
a) Atomic radius
• The valence electron of an atom with a larger
radius experience less attraction to the
nucleus. Hence, the atom have lower IE.

14. b) Effective nuclear charge
• Atom with higher effective nuclear charge
holds its electron closer to the nucleus than
the atom of a lower effective nuclear charge.
• Hence, the IE increases as the effective
nuclear charge become larger.
Ionization Energy (IE)

15. Ionization Energy (IE)
c) Shielding effect
• The shielding effect of inner electrons causes
the valence electrons to be less strongly
attracted to the nucleus and thus results in
lower IE.

16. Ionization Energy (IE)
General trend in first ionisation energy across a
period and down the group:
Ionisation energy increases
Ionisationenergydecreases

17. Increases across a period
• Proton number increases
• Effective nuclear charge increases
• Stronger attraction between nucleus and valence
electron
• Atomic radius decreases
• More difficult to remove valence electron
• Thus, IE increases.
Ionization Energy (IE)

18. Decreases down a group
• Energy level increases
• Shielding effect increases
• Weaker attraction between nucleus and valence
electron
• Atomic radius increases
• Easier to remove valence electron
• Thus, IE decreases
Ionization Energy (IE)

19. Electronegativity
• Electronegativity is a measure of an atom’s
attraction for another atom’s electrons.
Electronegativity increases
Electronegativityincreases

20. Trend in electronegativity across a period and down the
group:
Increases across a period
• Proton number increases
• Effective nuclear charge increases
• Stronger attraction between nucleus and valence
electron
• Atomic radius decreases
• Greater ability to attract the bonding electrons to
itself
• Thus, electronegativity increases
Electronegativity

21. Electronegativity
Decreases down a group
• Energy level increases
• Shielding effect increases
• Weaker attraction between nucleus and valence
electron
• Atomic radius increases
• Weaker ability to attract the bonding electrons to
itself.
• Thus, the electronegativity decreases