1. The Periodic Table 1. Number the groups and periods on your table.
2. Mark the following:
X - solid
O - liquid
* - gas
3. Use colors to indicate the following (include a key!!):
Alkali metals Halogens Alkaline earth metals Noble gases
Transition metals Nonmetals
Metalloids Lanthanides
(indicate the staircase!!!) Actinides
Other metals
4. Mark the 4 different blocks of the periodic table.

2. The Periodic Law
Unit 2
Topic 2
Chapter 5

3. History of the Periodic Table
Dmitri Mendeleev
- Arranged all the known elements by atomic mass.
- Noticed that the similarities in their chemical properties appeared at regular intervals (periodic).
- Moved elements into different groups if they had similar properties.
- Left gaps for elements not yet discovered.
Henry Moseley
- Concluded that atomic number not atomic mass should be used to arrange the periodic table.
Periodic Law
- The physical and chemical properties of the elements are periodic functions of their atomic numbers.

4. Atomic Radius
- One-half the distance between the nuclei of identical atoms that are bonded together.
- Gives a more reliable indicator for size since the edges of orbitals are fuzzy.
Period trends:
- atomic radius decreases across a period.
Why?
The attraction of the nucleus.
Group trends:
- atomic radius increases down a group
Why?
More main energy levels.

5. Trends for Atomic Radius↓
Atomic radius
Atomic
radius ↑

6. Ionization Energy
Electrons can be removed from an element provided there is enough energy available.
A + energy → A+ + e
IE is the energy required to remove one electron from a neutral atom of an element. (also called first ionization energy IE1).
In general, metals lose electrons and nonmetals gain electrons.
So, metals are LOSERS!!!

7. Trends for Ionization Energy
Period trends:
- ionization energy increases across a period.
Why?
The increase of nuclear charge.
A higher charge more strongly attracts electrons.
Group trends:
- ionization energy decreases down a group
Why?
More main energy levels.
More electrons lie between the nucleus and electrons in the highest energy levels. This partially shields the outer electrons from the effect of nuclear charge. This is called
Shielding Effect.

8. Trends for Ionization Energy↑
Ionization
energy ↓

9. Electron Affinity
Electrons can be acquired by a neutral atom of an element.
A + e → A- + energy
(-)
Some atoms are “forced” to gain an electron by adding energy. A + e + energy → A-
(+)

10. Trends for Electron Affinity
Period trends:
- electron affinity increases across a period.
- Halogens gain electrons easily.
- This explains their high reactivity.
Group trends:
- electron affinity decreases down a group
- Electrons add with greater difficulty down a group.
Why?
Increased atomic radius decreases the effect of increased nuclear charge.

11. Trends for Electron Affinity↑
Electron
Affinity ↓

12. Valence Electrons
Valence electrons are found in s and p sublevels of the highest main energy level of an element.
The periodic table can be used to find valence electrons.

13. Electronegativity
The ability of an atom in a chemical compound to attract electrons.
Period trends:
- electronegativity increases across a period.
Group trends:
- electronegativity decreases or stays about the same down a group.

14. Trends for Electronegativity↑
Electro-
negativity ↓