1. Periodic Groups and Trends

2. Periodic Table
Periodicity: regular variations (or patterns) of properties with increasing atomic weight. Both chemical and physical properties vary in a periodic (repeating pattern).
Group: vertical column of elements (“family”)
Period: horizontal row of elements

3. Activity: get out your black and white copy of the periodic table.

4. On one side, color and label the metals, nonmetals, and metalloids.
Another name for “metalloid” is “semi-metal”.

5. Color and label the groups/families of elements on the other side of
your paper. Remember to create a legend.
noble gases
Transition metals
alkali metals
halogens
alkaline earth metals
lanthanides
actinides

6. PERIODIC GROUPS alkali metals alkaline earth metals transition metals
halogens
noble gases
lanthanides
actinides

7. Alkali Metals Group 1 on periodic table Very reactive Soft solids
Readily combine with halogens (ex. NaCl)
Tendency to lose one electron (to become an ion)

8. Alkaline Earth Metals Group 2 on periodic table
Abundant metals in the earth
Not as reactive as alkali metals
Higher density and melting point than alkali metals

9. Transition Metals
Groups 3-12 on periodic table Important for living organisms (think vitamins: multivitamins (they have iron, zinc, chromium, etc.)

10. Halogens Group 7A on periodic table (***or GROUP 17)
“Salt former” – combines with groups 1 and 2 to form salts (ionic bonds)

11. Noble Gases Group 8 on periodic table (** or GROUP 18)
Relatively inert, or nonreactive
Gases at room temperature

12. Lanthanides Part of the “inner transition metals” Soft silvery metals
Tarnish readily in air
React slowly with water

13. Actinides
Radioactive elements
Part of the “inner transition metals”

14. PERIODIC TRENDS Atomic radii Ionization energy Ionic radii
Electronegativity

15. Atomic Radii

16. Atomic Radii Trend: increases down a group WHY???
The atomic radius gets bigger because electrons are added to energy levels farther away from the nucleus.
Plus, the inner electrons shield the outer electrons from the positive charge (“pull”) of the nucleus; known as the SHIELDING EFFECT

17. Atomic Radii Trend: decreases across a period WHY???
As the # of protons in the nucleus increases, the positive charge increases and as a result, the “pull” on the electrons increases.

18. Ionization Energy
Definition: energy required to remove outer electrons

19. Ionization Energy
Definition: energy required to remove outer electrons

20. Ionization Energy Trend: decreases down a group WHY???
Electrons are in higher energy levels as you move down a group; they are further away from the positive “pull” of the nucleus and therefore easier to remove.

21. Ionization Energy Trend: increases across a period WHY???
The increasing charge in the nucleus as you move across a period exerts greater “pull” on the electrons; it requires more energy to remove an electron.

22. Ionic Radii
Cations are always smaller than the metal atoms from which they are formed. (fewer electrons because metals LOSE electrons)
Anions are always larger than the nonmetal atoms from which they are formed. (more electrons because nonmetals GAIN electrons)

23. Electronegativity
Definition: the tendency of an atom to attract electrons to itself when chemically combined with another element

24. Electronegativity Trend: decreases down a group WHY???
Although the nuclear charge is increasing, the larger size produced by the added energy levels means the electrons are farther away from the nucleus; decreased attraction, so decreased electronegativity; plus, shielding effect

25. Electronegativity Trend: increases across a period
(noble gases excluded!)
WHY???
Nuclear charge is increasing, atomic radius is decreasing; attractive force that the nucleus can exert on another electron increases.

26. Summing Up Periodic Trends