1. Pre-AP 11/1
Pick up the Alien periodic table packet from the side table and the building the periodic table paper (this is your HW).
Take out a blank piece of paper.
We will review groups, properties, and history of the periodic table today at the beginning of class.
You will then be given an activity that will introduce you to the trends from the periodic table.
HW: Building the periodic table

2. Review (without looking at notes)
Name the group that are highly reactive nonmetals.
What is another name for the D-block?
Which group reacts with water and must be stored under oil?
List TWO metalloids.
Which group is stable, unreactive, and contains a full outer shell?
Where are the nonmetals located? Which is a liquid at room temperature?
What is the name of the row of elements that are radioactive and most are made in a lab?
How did Mendeleev organize the periodic table?
Who changed Mendeleev table? How did he organize it?
What is the name of the group that has two valence electrons?

3. Periodic Trends

4. Today you will need.. A blank piece of paper
The packet from the side table.
Then remainder of your papers have been graded unless turned in late. Please quickly pass those out.
From the packet, answer the following: effective nuclear charge and isoelectronic
NO QUESTIONS ABOUT ANYTHING UNTIL AFTER I GET YOU STARTED

5. On the clean piece of paper that you picked up, draw the following…
Trend
Definition
Why trend for group
Why trend for period/row
Exceptions Explained
Atomic Radius
N/A
Ionization Energy (I.E.)
*Group 2 to 13
*Group 15 to 16
Electron Affinity
Electronegativity
Ionic Radius:
Cation/Anion

6. Trends Tricks:
“A trend is an observation NOT an explanation”. Say it over and over until it sticks in your head.
Talk about BOTH atoms involved on the AP exam or you will lose credit.
Mention Coulombic attraction where it pertains to a trend.
Emphasize energy
Use effective nuclear charge (Zeff) when in a period/row AND use size, distance, and SHIELDING when in a group as an explanation

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

8. Atomic Radius
½ the distance between the nuclei of identical atoms that are bonded together
Increases to the LEFT and DOWN

9. Atomic Radius Why bigger across a period/row: Why larger down a group:
Effective nuclear charge decrease attraction of nucleus and therefore the pull of the nucleus results in larger radius
Why larger down a group:
Increase number of energy levels down a group, increase distance over which nucleus must pull electrons and therefore because of shielding of inner core, reduces attraction for electrons

10. Atomic Radius K Na Li Ar Ne

11. Atomic Radius Why larger going down? Why smaller to the right?
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

12. Atomic Radius Why bigger across a period/row: Why larger down a group:
Effective nuclear charge decrease attraction of nucleus and therefore the pull of the nucleus results in larger radius
Why larger down a group:
Increase number of energy levels down a group, increase distance over which nucleus must pull electrons and therefore because of shielding of inner core, reduces attraction for electrons

13. Ionization Energy
First Ionization Energy-energy required to remove one electron from a neutral atom
Increases UP and to the RIGHT

14. Ionization Energy
First Ionization Energy He Ne Ar Li Na K

15. Ionization Energy Why more energy across a period?
Effective nuclear charge increases the attraction of the nucleus therefore holds electrons more tightly (WITH EXCEPTIONS from GROUP 2 to 13 AND from GROUP 15 to 16)
Why more energy up a group?
Decreased number of energy levels, decreases the distance over which nucleus must pull and therefore reduces attraction for electrons
Full energy levels provide some shielding between nucleus and valence electrons so closer to the top of group, less energy levels and less shielding

16. Ionization Energy Successive Ionization Energies 2nd I.E. 1,445 kJ
Large jump in I.E. occurs when a CORE e- is removed.
Mg 1st I.E kJ
2nd I.E. 1,445 kJ
Core e- 3rd I.E. 7,730 kJ

17. Ionization Energy Successive Ionization Energies 2nd I.E. 1,815 kJ
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

18. Electron Affinity
Energy change that occurs when an electron is acquired by a neutral atom (addition of electron in gaseous atom or ion).
Tends to become less negative (less energy released) DOWN and to the LEFT

19. Electron Affinity Explained:
Why down a group:
Change little moving down a group.
Why more negative across a row/period towards noble gases:
Become increasingly negative from left to right. More positive, the less attractive to electrons around them. Careful: addition or subtraction can be exothermic (-) or endothermic (+). As you more toward the noble gases, the affinities become more negative.
Trend explained because of octet rule. (atoms close to full valence will tend to gain electrons and have very negative affinities (give off great deal energy when gaining electrons) NOBLE GASES DO NOT CONFORM TO THIS. They have very positive values.

20. Ionic Radius Ionic Radius (WHY) Cations (+) Anions (–) lose e- smaller
gain e-
larger
© 2002 Prentice-Hall, Inc.

21. Electronegativity
A measure of the ability of an atom in a chemical compound to attract electrons
Most electronegative element is fluorine
Given arbitrary value of 4; all others relative

22. Electronegativity Explained
Why higher up a group:
Decrease number of energy levels, distance over which the nucleus must pull and therefore increases attraction for electron.
Full energy levels AT THE bottom of a group provides shielding for valence electrons
Effect nuclear charge increases attraction of the nucleus increases and therefore strengthens the attraction of the electrons

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

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

25. Examples K or Li Al or Cl Li Cl
Which has the greater electonegativity?
K or Li
Al or Cl Li Cl

26. Examples S or S2- Al or Al3+ S2- Al
Which particle has the larger radius?
S or S2-
Al or Al3+
S2- Al

27. Add to review: Define isoelectronic and give a set of examples.
Place the following in order of increasing first ionization energy: Na, Al, Sn, Ca
Place the following in order of increasing atomic radii: Na, Mg, Rb, K
Why does the He atom have a smaller radius than the H atom? Why is the He atom smaller than the Ne atom. EXPLAIN.
List all atoms and common ions of representative elements that are isoelectronic with the aluminum ion. Which is the smallest?
Why does Mg atom require a larger amount of ionization energy than the Ba atom? Why does the Cl atom require a larger amount of ionization energy than the Mg atom. EXPLAIN.