1. Unit 3: Electrons and the Periodic Table
Periodic Table Trends
Unit 3: Electrons and the Periodic Table

2. What patterns do you notice?
Periodic Law
When elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals.
What patterns do you notice?
Atomic #s
3, 11, 19 are all alkali metals

3. Atomic Radius Atomic Radius size of atom Atomic Radius
© 1998 LOGAL
Atomic Radius
Average distance in an atom between the nucleus and the outermost electron

4. Atomic Radius Atomic Radius Increases to the LEFT and DOWN Smallest Fr
biggest

5. Atomic Size Trend Atomic Size increases down a group
Why larger going down?
Adding more energy levels.
Atomic Size decreases across a period
Why smaller across?
Increased nuclear charge (more protons) without additional energy levels pulls e- in closer.
Greater Coulombic Attraction

6. Atomic Size & Radius Examples
The closer you are to Francium, the larger you will be!
Which is larger:
Rb or Li Rb
N or Ne N

7. Ionization Energy
Ionization energy is the amount of energy needed to remove an electron.
M + energy  M e-
Electrons that are close to the nucleus are hard to remove because they are under a strong force of attraction

8. Ionization Energy Trend
Ionization Energy Increases across a period
Why? Valence electrons experience a greater nuclear force because they are closer to the nucleus.
Smaller atoms have higher Ionization energy.
Ionization Energy Decreases down a group.
Why? Valence electrons removed are farther from the nucleus because they are in higher energy levels.
Bigger atoms have lower Ionization energy.

9. Ionization Energy Trends
Why opposite of atomic radius?
In small atoms, e- are close to the nucleus where the attraction is stronger
Small atoms have High IE
Big Atoms have Low IE

10. Ionization Energy Lowest as you go DOWN and to the LEFT High IE Low IEFr
Low IE
Bottom left elements (Metals) WANT to lose an electron to become more stable.

11. Which would have a higher Ionization energy, Sodium or Chlorine?
Chlorine has higher IE.
Chlorine is smaller and has a higher nuclear charge (more protons) = stronger hold on electron = higher energy to take it away.
Also, remember – Na wants to lose an electron (it is a metal) and Cl wants to gain an electron (non-metal)

12. E. Ionization Energy
First Ionization Energy He Ne Ar Li Na K

13. Ionization Energy Successive Ionization Energies
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

14. Ionization Energy Successive Ionization Energies
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

15. Electronegativity The ability of an atom to attract an electron.
The smaller the atom, the more electronegative it is because of a greater nuclear force.

16. Electronegativity Trends
Electronegativity Increases across a period.
Why? Non-metals such as F, O and N want more electrons to complete their valence shell.
Smaller atoms have greater nuclear charge and thus, more force to attract electrons.
Exception: Noble gases are not included because they generally do not want to gain electrons. They are already stable.

17. Electronegativity Trends
Electronegativity Decreases Down a Group
Why? Atomic size increases and valence electrons are farther from the nucleus.
More energy levels increases shielding. So the pull from the positive nuclear charge is less.
In General:
Non-Metals have high Electronegativities
Metals have low Electronegativities

18. Electronegativity Trends
Highest as you go UP and to the RIGHT towards Fluorine F
Remember- Noble gases not included in this trend!

19. Ionic Radius Na  Na+ Ionic Radius
Cations (+ ions) the ionic radius is smaller than the original atom.
Why? There is an increased attraction for the fewer electrons that remain.
Na  Na+

20. Ionic Radius
For Anions (– ions) the ionic radius is larger than the original atom.
Why? The nuclear attraction is less for an increased number of electrons.
Extra electrons repel each other and spread out – larger!)
© 2002 Prentice-Hall, Inc.
Cl  Cl-1

21. Practice Which atom is larger H or He?
Which atom has a greater ionization energy, Ca or Sr?
Which atom is more electronegative, F or Cl?
Hydrogen – Smaller nuclear charge
Ca – smaller, less shielding, lower effective nuclear charge
Fluorine – Smaller, less shielding with less energy levels, so easier to attract electron

22. Ca – lower nuclear charge
Examples
Which atom has the larger radius?
Be or Ba
Ca or Br
Ba –more energy levels
Ca – lower nuclear charge

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

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

25. Alkali Metal Reactivity