1. Section 5.3 – Electron Configuration and Periodic Properties
Honors chemistry

2. Atomic Radius
Determined by the distance from the nucleus to the edge of the outer orbital.
Edge of outer orbital not well defined
Use identical bonded atoms – ½ of the distance between the nuclei

3. Trends in Atomic Radii

4. Trends in Atomic Radii Decrease across the period
Due to increasing positive charge of the nucleus
Increase as you go down the group
Exception Ga to Al – Ga smaller due to increased nuclear charge (first addition of d electrons)

5. Problems
Which of these elements; Li, Rb, K or Na has the smallest radius? Largest?
Which of these elements; O, Se, S and Po has the smallest atomic radius? Largest?

6. Ionization Energy Atom + energy → Atom+ + e-
First electron removed – First Ionization Energy (IE1)
Second electron removed – Second Ionization Energy (IE2) etc.
Group 1 – lowest ionization energy
Group 18 - highest ionization energy
Ionization energies increase across the period due to increased nuclear charge.
Ionization energies decrease down a group due to further distance from nucleus and electron shielding

7. Ionization Energies

8. Ionization Energies

9. Successive Ionization Energies

10. Why? Each successive electron feels a stronger nuclear attraction
This information lead to the understanding of the stability of the noble gas configuration+

11. Practice Choose the element with the higher IE1: Ca or Ba Ca or Br
Ca or K
Ca or Mg

12. Electron Affinity Atom + e- → Atom- + (- energy) IMPORTANT!!!!!
Negative energy means energy lost by system
Positive energy means energy gained by system
Sign indicates direction not numerical value!!!!

13. Electron Affinity Values

14. Electron Affinity Trends
Generally become larger (look at as absolute value) as you move across the period.
Exception – Group 15 due to half filled p orbitals
Generally become smaller as you move down a group due to:
Greater Nuclear Attraction
Greater Atomic Radius

15. Second Electron Affinities
Very difficult to add an electron to an anion (negative ion)
Second Electron Affinities are all positive

16. Ionic Radii Cation (positive ion) Anion
Smaller atomic radius than atom
Due to:
electrons being removed
increased effective nuclear charge
Anion
Larger atomic radius than atom
Due to
electrons being added
decreased effective nuclear charge
Greater repulsion of electrons

17. Ionic Radii

18. Valence Electrons
Available to be lost, gained or shared when compounds are formed.
In outer main energy levels
For Main Group Elements – s and p orbitals

19. Bonded Atoms Very rarely are electrons shared equally
Usually attracted more to one atom
This will effect the chemical properties of the compound!!!
Measure of attraction – called electronegativity
Based on a 4.0 scale – F = 4.0.
Developed by Linus Pauling

20. Electronegativity

21. Electronegativity Trends
Increase across a period.
Tend to decrease or stay the same down a group.
If a noble gas does not form compounds – it does not have an electronegativity
If a noble gas does form compounds – it will have a high electronegativity

22. Summary of Trends

23. Summary of Trends

24. D-Block
These elements tend to vary less and with less regularity than Main Group Elements.
Still electrons in d orbitals are often responsible for characteristics of elements in the d-block
Atomic radius tends to decrease across the block
Ionization energies generally increase across both the d and f-blocks

25. D and F Blocks Tend to lose electrons from outer shell!!!!
That means the valence electrons come from the ns shell not the (n-1)d shell
Generally these elements from 2+ ions.
Electronegativities
D-block - between 1.1 and 2.54 (Only groups 1 and 2 are lower)
F-block – between 1.1 and 1.5