Periodic Trends from the Periodic Table.

Periodicity

1. Periodic Trends
Mrs. Landgraf
Honors Chemistry

2. Development of the
Periodic Table
Mendeleev developed periodic table
to group elements in terms of chemical
properties.
Blank spots where elements should be
were observed.
Periodic law = elements arranged by
atomic number gives physical and
chemical properties varying
periodically.

3. Part I – Atomic Size

4. Atomic Radius
Measures as distance from
nucleus to nucleus and divided by
2.
Unit commonly used is pm
1 picometer = 10
-12
meter
Example: iodine atomic radius
=140pm

5. How does atomic radius
change across a period?
It is smaller to the right.
Why?
More protons in the nucleus
higher electrical force pulls
electrons closer to nucleus.

6. How does atomic radius
change down a group?
It is larger down the group.
Why?
Valence electrons are at higher
energy levels and are not bound as
tightly to the nucleus because they are
screened or shielded ( pushed away)
by other electrons in inner levels.

7. Note
There are some exceptions.
Example column 13.

8. The Periodic Table and
Atomic Radius

9. Example:
Which is larger: a lithium atom or a
fluorine atom?
A lithium atom

10. Example:
Which is larger: an arsenic atom
or a sulfur atom?
An arsenic atom

11. Ion
Positive ion- removal of
electron
Negative ion - addition of
electron

12. Ionic size
Metallic elements easily lose
electrons.
Non-metals more readily gain
electrons.
How does losing or gaining an
electron effect the size of the
atom (ion) ?

13. Positive ions
Positive ions are always smaller
that the neutral atom. Loss of
outer shell electrons.

14. Negative Ions
Negative ions are always larger
than the neutral atom. Gaining
electrons.

15. Ion size trends in periods.
Going from left to right there is a
decrease in size of positive ions.
Starting with group 5, there is
sharp increase followed by a
decrease in the size of the anion
as you move from left to right.

16. Ion size trends in columns.
Ion size increases as you move
down a column for both positive
and negative ions

17. Part II – Ionization Energy

18. Ionization energy
Ionization energy is the amount
of energy needed to remove an
electron from a gaseous atom.
First ionization energy – 1+
Second ionization energy – 2+

19. How does ionization energy
change down a group?
The first ionization energy
decreases as you move down a
group.
Why?
The size of the atom increases.
Electron is further from the
nucleus.

20. How does ionization energy
change across a period?
The first ionization energy
increases as you move from left
to right across a period.
Why?
Nuclear charge increases while
shielding is constant.
Attraction of the electron to
the nucleus increases.

22. Part III – Electronegativity

23. Electronegativity: the ability of an
atom in a bond to pull on the electron.
(Linus Pauling)

24. Electronegativity
When electrons are shared by two atoms
a covalent bond is formed.
When the atoms are the same they pull
on the electrons equally. Example, H-H.
When the atoms are different, the atoms
pull on the electrons unevenly. Example,
HCl

25. Trends in Electronegativity
Electronegativity generally decreases as
you move down a group.
Electronegativity of the representative
elements (Group A elements) increases as
you move across a period.

26. Electronegativities of Some
Elements
Element Pauling scale
F 4.0
Cl 3.0
O 3.5
N 3.0
S 2.5
C 2.5
H 2.1
Na 0.9
Cs 0.7

27. Note
Most electronegative element is F
(EN 4.0)
Least electronegative stable element is
Cs
(EN 0.7)

28. Summary
Shielding is constant
Atomic Radius decreases
Ionization energy increases
Electronegativity increases
Nuclear charge increases
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