1. OR Why we call it the PERIODIC table
PERIODIC TRENDS OR
Why we call it the PERIODIC table

2. PERIODIC What does Periodic mean?
What are examples of things that are periodic?

3. Periodic Trends
Across the periodic table: nucleus becomes more positive, more attraction occurs, elements become more non-metallic
Down a group: the elements become more positive, but there is a larger nucleus due to the increased number of protons and neutrons. Therefore, there is less attraction between protons and outer electron levels and the elements have more metallic properties.

4. Periodic Trends
Both the position and the properties of the elements arise from their electron configuration (which comes from their atomic number).
Same column = similar outer level e- configuration
Properties that are periodic:
metals, metalloids, nonmetals, boiling point, density, atomic radii, ionic radii, oxidation numbers, ionization energies, electron affinities, electronegativity

5. Boiling Point of the Elements

6. Alkali Metals Video
Looking at a periodic trend of the alkali metal family.

7. Alkali Metals Video
What is the chemical equation for the reaction for each of these metals in water?
2Li + H2O 2LiOH + H2
2Na + H2O  2NaOH + H2
2K + H2O  2KOH + H2

8. A Trend for Groups SHIELDING
The higher the principle quantum number, the more energy levels there are. The orbitals are further out. Inner electrons shield the positive charge of the protons and the electrons in the outer energy level come off more easily.

9. A Trend for Periods PROTON PULL (nuclear charge)
The greater the positive charge (due to increased atomic number), the greater the attraction (pull) within the energy level. More protons can be attracted to more electrons.

10. ATOMIC RADII
page 188
As the energy levels increase, the principle quantum number increases , and so does the atomic radii.
As the atomic number increases across a period, the positive charge also increases within the same energy level.
TREND: atomic size increases down a group, atomic size decreases across a period.
Why: within a group, inner level electrons shield outer level e- from the positive nucleus; distance from nucleus increases; within a period, size of nuclear charge increases and the attraction between protons and electrons pulls the atom in tighter.

11. ATOMIC RADII

12. IONIC RADII
page 190
When atoms form a compound, the compound is more stable than the uncombined atoms were. The outer energy levels are full (like the noble gases).
TREND: Metallic ions are smaller than the atoms they come from. Nonmetallic ions are larger than the atoms they come from.
Why: metal lose electrons and become the noble gas configuration on the energy level below. Nonmetals gain electrons and become the noble gas configuration at the end of their period.

13. IONIC RADII

14. Sodium + Chlorine = Salt

15. OXIDATION NUMBERS
This is the charge number an atom would have if the valence electrons were completely transferred/lost.
Group 1 (1A) 1+ oxidation (H can have 1+ or 1-)
Group 2 (IIA) 2+ oxidation
Group 3-12 (IIIB) can have 1+ to 8+ (lose “d” and then “s”, one at a time)
Group 13 (IIIA) 3+
Group 14 (IVA) 2+ or 4+
Group gain electrons (-3, -2, -1)
Group 18 noble gases; 0

16. OXIDATION NUMBERS

17. IONIZATION ENERGY page 191-192
This is the energy required to remove an electron. The first ionization energy is the energy required to remove the most loosely held electron.
TREND: Ionization energy increases as atomic number increases across a period. In a group, there is a gradual decrease in the first ionization energy as the atomic number increases. Metals have low first ionization energies and nonmetals have high first ionization energies.
Why? Across a period, the larger the nuclear charge, the greater the ionization energy; Down a group, the greater shielding effect leads to less ionization energy; With a bigger electron cloud (more energy levels), the ionization energy decreases; an electron in a full or half-full sublevel requires additional energy to be removed.

18. FIRST IONIZATION ENERGY

19. IONIZATION ENERGIES (kilojoules per mole) 1st 2nd 3rd 4th 5th 6thHe Li Be B C Al Ga

20. ELECTRON AFFINITY
The attraction of an atom for an electron; the energy change that occurs when an atom gains an extra electron. The same factors that affect ionization energies affect electron affinities.
TREND: Nonmetals have large electron affinities (except for Noble gases). The more stable an atom is, the less the tendency to have an affinity for electrons. Metals have low electron affinities and Nonmetals have high electron affinities. Noble gases (p6) and Alkaline Earth Metals (s2) have negative affinities (extremely low).

21. ELECTRON AFFINITY H He 72.766 (-21) Li Be B C N O F Ne
(-21)
Li Be B C N O F Ne
59.8 (-241) (-29)
Na Mg Al Si P S Cl Ar
52.9 (-230) (-34)
K Ca Ga Ge As Se Br Kr
46.36 (-156) (-39)
Rb Sr In Sn Sb Te I Xe
46.88 (-167) (-40)
Cs Ba Tl Pb Bi Po At Rn
45.5 (-52) (170) (270) (-41)

22. ELECTRONEGATIVITY page 194
A tug of war between atoms for electrons in a chemical bond. How well the electrons “tug” is electronegativity. The better the atom ‘tugs’ an electron from another element, the higher the electronegativity
TREND: increases from left to right across a period as the number of valence electrons increases and the size of the atom decreases. A low ionization energy and a low electron affinity means low electronegativity.
Fluorine is the most electronegative element.

23. ELECTRONEGATIVITY

24. ELECTRONEGATIVITIES

25. SUMMARY Noble Gases: - valence electrons s2p6 - ionization energy HIGH
- Electron Affinity LOW
- Electronegativity LOW

26. SUMMARY Halogens: - valence electrons s2p5 - ionization energy HIGH
- Electron Affinity HIGH
- Electronegativity HIGH

27. SUMMARY Alkali metals and Alkaline Earth metals:
- valence electrons s1 and s2
- ionization energy LOW
- Electron Affinity LOW
- Electronegativity LOW

28. SUMMARY - STABILITY