1. Lecture 29 Periodic trends Ozgur Unal
NIS – CHEMISTRY
Lecture 29
Periodic trends
Ozgur Unal

2. Periodic Trends – Atomic Radius
How do we determine the size and shape of atomic orbitals?

3. Periodic Trends – Atomic Radius
Atomic size is a periodic trend influenced by electron configuration.
In order to determine the atomic size of elements we first separate metals and nonmetals.
For metals, the atomic radius is half the distance between adjacent nuclei in a crystal of the element.
For nonmetals, the atomic radius is half the distance between nuclei of
identical atoms that
are chemically
bonded together.
Figure 6.10.

4. Atomic Radius – Trends within Periods
In general the atomic radius
decrease as you move from left
to right across a period.
This is because of
the increasing positive charge
in the nucleus
the fact that the principal
energy level within a period
remains the same.

5. Atomic Radius – Trends within Groups
Atomic radii generally increase as you move down a group.
As you move down a group, the nuclear charge increases adn new energy levels are added.
However, the increased nuclear charge does not pull the outer electrons toward the nucleus to
make atoms smaller.

6. Atomic Radius

7. Atomic Radius
Example: Which has the largest atomic radius: C, F, Be or Li? Answer without referring to Figure Explain your answer in terms of trends in atomic radii.
Example: Which has the largest atomic radius: Mg, Si, S or Na? The smallest?

8. Ionic Radius
Atoms can gain or lose one or more electrons to form ions.
Since electrons have charge, atoms that gain or lose electrons acquire a net charge.
An ion is an atom that has a positive or negative charge.
How does the formation of ion affect the size of an atom?
When atoms lose electrons and form positively charged ions, they always become smaller.
When atoms gain electrons and form negatively charged ions, the become larger.
Figure 6.13

9. Ionic Radius – Trends within Periods
Elements on the left side of the periodic table form smaller positive ions and elements on the right side of the table form larger negative ions.
In general, as you move from left to
right across a period, the size of the
positive ions gradually decreases.
Then, beginning in group 15 or 16,
the size of the much-larger negative
ions also gradually decreases.

10. Ionic Radius – Trends within Groups
As you move down a group, an ion’s outer electrons are in orbitals corresponding to higher principal energy levels, resulting in a gradual increase in
ionic size.
Thus, the ionic radii of both
positive and negative ions increase
as you move down a group.

12. Lecture 30 Periodic Trends - Continued Ozgur Unal
NIS – CHEMISTRY
Lecture 30
Periodic Trends - Continued
Ozgur Unal

13. Ionization Energy
In order to form a positive ion, an electron must be removed fro ma neutral atom.
Removing an electron requires energy.
Ionization energy is defined as the energy required to remove an electron from a gaseous atom.
The energy required to remove the first electron from an atom is called the first ionization energy.
Example: Lithium, Li
The first ionization energy of Li is 8.64 x J.
The loss of an electron from Li results in Li+.
Check out Figure 6.16

14. Ionization Energy

15. Ionization Energy Removing more than one electron:
After removing the first electron from an atom, it is possible to remove additional electrons.
The energy required to remove the second electron from a +1 ion is called the second ionization energy.
Third, fourth fifth ionization energies and so on.
Check out Table 6.5
After removing all valence electrons, the ionization energy of the next electron increases in a great amount.
Atoms hold onto their inner core electrons much more strongly than they hold onto their valence electrons.

16. Ionization Energy Trends Within Periods:
First ionization energies generally increase as you move from left to right across a period.
Check out Figure 6.16
Trends Within Groups:
First ionization energies
generally decrease as you
move down a group.
Check out Figure 6.17

17. Octet Rule Sodium atom: Na  1s2 2s2 2p6 3s1
Sodium ion: Na+1  1s2 2s2 2p6 the same electron configuration with Ne.
The octet rule states that atoms tend to gain, lose or share electrons in order to acquire a full set of 8 valence electrons.
This is why noble gases are stable, because they already have 8 valence electrons in s and p orbitals.
The octet rule is useful in determining the type of ions likely to form.
Elements on the right side of the periodic table gain electrons (negative ions) and on the elements on the left side of the table lose electrons (positive ions) to be more stable.

18. Electronegativity
The electronegativity of an element indicates the relative ability of its atoms to attract electrons in a chemical bond.
Electronegativity generally decreases as you move down a group and increases as you move from left to right across a period.
Electronegativity values are expressed in terms of a numerical value of 3.98 or less.
The higher the electronegativity value, the stronger the atom attracts electrons in a chemical bond.
Fluorine is the most electronegative element.
Noble gases do not have electronegativity values. Why not?

19. Electronegativity