1. Unit 3 Lecture 4: Periodic Table Trends

2. PERIODIC TABLE AEROBICS!

3. Electronegativity (electron affinity)
Periodic Table Trends
Atomic Size/Radius
Ionization Energy
Electronegativity (electron affinity)

4. Graphing Periodic Trends Activity
Use the information in the tables to complete the graph.
Using colored pencil or pen, list under the symbol (in this order!) the Atomic Radius, First Ionization Energy, and Electronegativity.
Use a different color for each property. Example: Write all of the Atomic radius values in red, all of the First ionization energies in green, and all of the Electronegativities in blue.
Of course, you can pick any colors that are available, as long as you are consistent.

5. Graphing Periodic Trends Activity
Observe the trends in each property as you go down the Alkali metal group, and as you go across Period 3.
Write out each of the statements written on the opposite side, completing each statement with the observed trend (increase or decrease)

6. Effective Nuclear Charge (Zeff)
(-) Electrons attracted to (+) nucleus
3 factors that effect this
The more protons in the nucleus, the greater the Zeff
The more distance between the nucleus and electrons the smaller the Zeff
The more repulsion between electrons the smaller the Zeff

7. Zeff = Z - S Zeff = effective nuclear charge
Z = # of protons (nuclear charge)
S= screening electrons (core electrons)
Ex. Oxygen: 1s22s22p4
Zeff = 8 – 2 = + 6
Ex. Fluorine: 1s22s22p5
Zeff = 9 – 2 = + 7
Fluorine has a greater effective nuclear change

8. The Analogy with Atomic Radius
Trey Songz = Nucleus
His Charm = Effective Nuclear Charge
Ladies = Electrons

9. Atomic Radius (size)
How large an atom is

10. Atomic Radius (notice any patterns?)

11. Atomic Size/ Atomic radius
down groups  Increases
Electrons in increased energy levels (shells)
across periods  Decreases.
There are more valence electrons in the same shell (energy level) so they pull towards the nucleus which has an increased effective nuclear charge with more protons.

13. First Ionization Energy
Energy required to remove one electron from an atom in the gas phase.

14. Ionization Energies (notice any patterns?)

15. Ionization Energies (notice any patterns?)

16. Ionization Energy Down groups  Decreases Across periods  increases
as electrons are further from the nucleus there is more screening of core electrons so less effective nuclear charge.
Across periods  increases
Increased effective nuclear charge holds electrons more tightly.

18. Electronegativity-
Measure of an attraction of an atom for a shared electron

19. Trend for electronegativity

20. Trends in Electronegativity:
down groups  decreases
The electrons are further from the nucleus, so less effective nuclear charge
across periods  increases
as atoms have more valence electrons in the same energy level and more protons in nucleus they have a higher effective nuclear charge.