1. Periodic Trends
Glencoe: Section 6.3

2. Atomic Radius
Remember that the electron cloud around a nucleus is based on a probability of 90% and does not have a definite edge.
Atomic size varies from substance to substance. The radius of a metal atom is ½ the distance between two adjacent atoms. The radius of a nonmetal atom is often determined from a diatom of an element.

3. Trends within periods
There is a decrease across periods due to the increasing positive charge in the nucleus and because the principal energy level stays the same.
The valence electrons are not shielded from the increased nuclear charge, and thus that charge pulls the valence electrons closer to the nucleus.

4. Trends within groups
Atomic radii increases as you move down groups. The nuclear charge increases and electrons are added to successively higher principal energy levels. The outermost energy level increases in size along with the p.e.l. which makes the atom larger. The increased distance offsets the greater pull of the nuclear charge.

5. Ionic Radius
Atoms can gain or lose electrons to form ions. When this happens, atoms acquire a net charge. An ion is an atom or a bonded group of atoms.
When atoms lose electrons and form positively charged ions (cations), they always become smaller.

6. The electron lost will always be a valence electron, which may leave a completely empty outer shell and results in a smaller radius. The electrostatic repulsion between the remaining electrons decreases, allowing them to be pulled closer to the nucleus.

7. When atoms gain electrons to form negatively charged ions (anions), they always become larger. The addition of an electron to an atom increases the electrostatic repulsion between the atom’s outer electrons, forcing them further apart. The increased distance results in a larger radius.

8. Trends within periods
Elements in the s-block form smaller cations and elements in the p-block form larger anions. As you move left to right, the size of the cations decrease, then the anions’ size also decreases.

9. Trends within groups
As you move down a group, an ion’s outer electrons are in higher principal energy levels, resulting in a gradual increase in ionic size. Thus, as you move down groups, ionic radii of both cations and anions increase.

10. Ionization Energy
To form a cation, an electron must be removed from a neutral atom. This requires energy to overcome the positively charged nucleus and the negatively charged electrons.
Ionization energy is the energy required to remove an electron from a gaseous atom.

11. The energy required to remove the first electron from an atom is called the first ionization energy.
This is an indication of how strongly an atom’s nucleus holds onto its valence electrons. A high IE indicates the atom has a strong hold on its electrons. Likewise, a low IE indicates a weak hold.

12. After removing the first electron, it is possible to remove additional electrons.
From left to right across periods, IE increases. The increased nuclear charge produces an increased hold on valence electrons.
From top to bottom down groups, IE decreases. Atomic size increases and the valence electrons are further away from the nucleus.

13. Electronegativity
This is the ability of an element’s atoms to attract electrons in a chemical bond.
Values are of 4.0 or less, in units of paulings.
Fluorine is the most electronegative element, with a value of 3.98 paulings.
Francium and cesium are the least electronegative, with values of .79 and .7 paulings, respectively.

14. Trends
Electronegativity generally decreases as you move down a group and increases as you move across a period.
The lowest electronegativities are in the lower left corner, and the highest are in the upper right.

15. The Octet Rule
Atoms tend to gain, lose or share electrons in order to acquire a full set of eight valence electrons.
Elements on the right side gain electrons and elements on the left side lose electrons.