1. III. Periodic Trends (p. 140 - 154)
Ch. 5 - The Periodic Table
III. Periodic Trends (p )

2. Periodic Law
When elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals.

3. Three rules apply to all periodic trends
1. Electrons are attracted to the protons in the nucleus
2. Electrons are repelled by other electrons in an atom
3. Completed shells are very stable

4. Atomic Radius
Atomic Radius K Na Li Ar Ne

5. Atomic Radius Decreases to the RIGHT across a period
Increases going down a group

6. Atomic Radius Why larger going down?
Higher energy levels have larger orbitals
Shielding - core e- block the attraction between the nucleus and the valence e-
Why smaller to the right?
Increased nuclear charge without additional shielding pulls e- in tighter

7. First Ionization EnergyHe Ne Ar Li Na K

8. First Ionization Energy
Increases UP and to the RIGHT

9. Ionization Energy Why does it increase as we move right?
As more protons are added to nucleus, the negatively charged e- are more strongly attracted which increases the energy required to remove them
Why decrease down a group?
Inner shells shield the outer valence electrons from the pulling power of the nucleus, making them easier to remove

10. Ionization Energy Successive Ionization Energies
Large jump in I.E. occurs when a CORE e- is removed.
Al 1st I.E kJ
2nd I.E. 1,815 kJ
3rd I.E. 2,740 kJ
Core e- 4th I.E. 11,600 kJ

11. Electronegativity
Moving across a period from left to right, electronegativity INCREASES.
Moving down a group, electronegativity DECREASES.

12. Electron Affinity
Electron Affinity – The energy change that occurs when an electron is acquired by a neutral atom.
Most atoms release energy (negative values)
If the electron causes the atom to have a more stable configuration then it releases a lot of energy.

13. Ionic Radii
Cations – shrink because the nucleus is attracting fewer electrons
Anions – expand because there is increased electron/electron repulsive forces
Isoelectronic – Ions containing same # of electrons. Protons determine size

14. Melting/Boiling Point
Highest in the middle of a period.

15. Examples
Which atom has the larger radius?
Be or Ba
Ca or Br Ba Ca

16. Examples
Which atom has the higher 1st I.E.?
N or Bi
Ba or Ne N Ne

17. Ch. 5 - The Periodic Table
I. History (p )

18. A. Mendeleev Dmitri Mendeleev (1869, Russian)
Organized elements by increasing atomic mass.
Elements with similar properties were grouped together.

19. A. Mendeleev Dmitri Mendeleev (1869, Russian)
Predicted properties of undiscovered elements.

20. B. Moseley
Henry Mosely (1913, British)
Organized elements by increasing atomic number.
Resolved discrepancies in Mendeleev’s arrangement.

21. Quick Review
Who organized the elements according to atomic mass?
Dimitri Mendeleev
Correctly predicted properties of new unknown elements

22. What are Chemical Families?
Elements with similar chemical properties
Alkali Metals
Alkaline Earth Metals
Transition Metals
Halogens
Noble Gases

23. B. Chemical Reactivity Families
Similar valence e- within a group result in similar chemical properties

24. B. Chemical Reactivity Alkali Metals Alkaline Earth Metals
Transition Metals
Halogens
Noble Gases

25. Color Code the Following:
-Group #’s ions -Period #’s ions -Lanthanides ions -Actinides Transition -Noble Gases Elements -Halogens

26. Homework:
Complete