1. Periodic Trends Part Deux Mrs. Wilson
Mon and Tues 10-27
Periodic Trends Part Deux
Mrs. Wilson

2. Turn in your Flame Test Labs NOW. Please follow the LASA honor code.
Daily Quiz
Turn in your Flame Test Labs NOW.
Please follow the LASA honor code.
No calculator required, but you need a white laminated periodic table.

3. - electron configuration - ionic radius
Objectives
Describe, understand and explain periodic trends down a group and across a period of the periodic table, including exceptions:
- electron configuration - ionic radius
- atomic radius ionization energy
- electronegativity
Homework: Periodic Trends Activity due
Mon 11-02; Hwk 2.8; Unit 2 Exam next class

4. **EXTREMELY IMPORTANT TEST INFO**
The Unit 2 Test is on Wed and Thurs
You can use one 3x5 inch handwritten index card (both sides) just like last time.
Both Unit 1 and Unit 2 material will be on the test’s multiple choice section.
A list of Unit 2 Exam resources is now online at the class website and will grow as more are added.

5. Four other periodic trends:
Atomic Radius: half the diameter (in pm) across the atom
Ionic Radius: half the diameter (in pm) across the ion
Ionization Energy: the amount of energy (in J) absorbed by the atom in order to remove valence electrons.
The first ionization energy is the energy absorbed to remove the first, loosest valence electron
Electronegativity: measure of tendency of an atom already in a chemical bond to attract other atoms/electrons
Reactivity: How quickly an element reacts (loses/gains valence e)

6. To understand these explanations, understand:
1. Electrons repel each other.
More electrons = more repulsion (esp. in the same orbital)
More repulsion = electrons will stand as far apart as possible
More energy levels = more electron repulsion
More electron repulsion = easier to remove electrons from atom
2. The nucleus attracts electrons.
More protons in nucleus = more nuclear attraction / pull
(= greater effective nuclear charge or Zeff)
Electrons closer to nucleus will shield it from its valence electrons.
(= greater shielding effect)

7. Explaining the Atomic Radius Trend
Decreases across a period
Increased nuclear charge (more #protons) pulls electrons closer to nucleus, decreasing radius (greater effective nuclear charge = Zeff )
Increases down a group
New electrons are placed in higher energy levels
More energy levels makes atom larger
More electrons = more repulsion = atom larger

8. Explaining the Ionic Radius Trend
Ions – atom or bonded group of atoms that has a positive or negative charge due to a loss/gain of electrons
Positive charge  lost e’s
Cation is always smaller than atom
Fewer electrons in cation = less electron repulsion
= more nuclear attraction or Zeff = smaller size
Negative charge  gained e’s
Anion is always larger than atom
More electrons in anion = more electron repulsion = less nuclear attraction or Zeff = increased size

9. What if the # of electrons is equal?
Which is larger, S2- or Ar?
Both of them have the same number of electrons (18) and same e. config
Solution: The one with greater Zeff (more #protons) is smaller S2- is larger
Species that have the same # of electrons = isoelectronic.
Which is largest? Br -, Kr or Rb+?

10. Explaining Electronegativity Trends
Increase across a period **F is the highest EN on PT
Zeff increases across period (b/c # protons increases).
Metals have low EN – so, lose valence e’s
Nuclear attraction increases across period – nonmetals gain e’s
Decrease down a group
Shielding effect increases
Core (inner) electrons surround nucleus and shield it from valence electrons; more shielding = less attraction on electrons
- Greater shielding = less “pull” on valence e = less attraction for other electrons

11. Explaining Ionization Energy Trends
Increases across a period
Zeff increases (# protons in nucleus increases)
Nuclear attraction increases; greater pull on e’s = more energy needed to remove valence e’s
Decreases down a group
Shielding effect increases down a group
Greater shielding of nucleus = less Zeff on valence e’s
Less Zeff means less energy needed to remove valence electrons

12. Exceptions to the Ionization Energy Trend
Between Group 2 (ex. Be) and Group 13 (ex. B): IE drops - Why?
Be = 1s2 2s2 B = 1s2 2s2 2p1
The 2p orbital is further from the nucleus; less Zeff
The 2p1 electron in B is shielded from the nucleus; less Zeff occurs
despite increase in #protons
Between Group 15 (ex. N) and Group 14 (ex. O): IE drops – why?
N = 1s2 2s2 2p O = 1s2 2s2 2p4
The 2p4 electron in O feels repulsion from the electron already in the orbital
Repulsion in that orbital = easier to remove that last electron

13. 1st Ionization Energy

14. Periodic Trends Activity
Groups of 3 – 4 for the activity portion
Each chart represents an element. Some have groups, and some do not.
Cut each “chart” apart. You should only have to cut once or twice. Do not waste time trimming each chart.
Put elements with the same group together.
Use your knowledge of periodic trends, move the charts around to construct a “Periodic Table” on your white construction paper, using BOTH the ones with groups and the ones without.
Glue or tape down the charts once you have checked your order with me. The chart and questions are due Mon

15. Quizizz: Explaining Periodic Trends
On your phone, open a web browser
Go to: join.quizizz.com
I will show you the code. Enter the code and tap “Join”.
Enter your REAL first name.