1. Quantum Mechanical Atom Part III: Trends in the Periodic Table Chapter 8 Section 10 of Jespersen 6TH ed)
Dr. C. Yau
Spring 2013 1 1

2. Table of Elements
The necessity of a table of elements arose from the need to find a pattern in the various properties of the elements.
Science involves looking for patterns.
Learning about patterns allow us to place elements in categories. We would need to learn only the general trends and not try to memorize the properties of every single element. 2

3. Some Early Attempts at Organizing the Elements.

6. Very interesting video of History of the Periodic Table.

8. Mendeleev's Periodic Table
There were many other attempts at writing a table of elements.
Mendeleev's was the most successful in helping us predict phys. & chem. properties.
(1) He placed the elements in order of increasing mass (protons were not discovered yet).
(2) He then placed elements of similar properties in the same column.
He found that the properties start repeating in cycles of 8. (Every 8th element has very similar properties.)
These cycles are what we refer to as "periods" and hence, the name, “PERIODIC TABLE.” 8

9. There were discrepancies, and there were blanks in the table.9

10. The Modern Day Periodic Table
Our current periodic table is based on increasing atomic number rather than mass. Why does it work?
Atomic # = # electrons
Reactions involve addition, removal or sharing of electrons.
# electrons determines a lot of the chemistry.
Mass determines HOW MUCH is made, not WHAT is made.

11. The Modern Day Periodic Table

12. Group # = # valence electrons Period # = n of the outershell
Valence electrons of S = 3s2 3p4 (6 electrons in n = 3)
Valence electrons of iodine =
5s2 5p5 (7 electrons in n = 5) 12

13. Electron Configuration of Valence Electrons
Group IA = ns1 Group VA = ns2 np3
Group IIA = ns2 Group VIA =
Group IIIA = ns2 np1 Group VIIA =
Group IVA = ns2np2 Group VIIIA= 13

14. Electron Dot Symbols Gp IA = ns1 Gp VA = ns2 np3 Gp IIA = ns2
Gp IIIA= ns2 np1
Gp IVA= ns2 np2
Gp VA = ns2 np3
Gp VIA = ns2 np4
Gp VIIA = ns2 np5
Gp VIIIA= ns2 np6 14

15. Halogens Halides F = 2s2 2p5 F- = 2s2 2p6 Cl = 3s2 3p5 Cl- = 3s2 3p6
Br = 4s2 4p5
I = 5s2 5p5
Halides
F- = 2s2 2p6
Cl- = 3s2 3p6
Br- = 4s2 4p6
I- = 5s2 5p6 15

16. Note #57 jumps to #72, and #89 to #104. What happened? 16

17. The “f-block” refers to the lanthanum and actinum series.

18. Size Increase Fig. 11.22: Atomic radii in picometers (1 m = 1012 pm)18 18

19. We summarize the periodic trend of atomic size by drawing a diagonal arrow.
smallest atoms
largest atoms 19

20. Why are there such trends in size?
To explain why atomic size increases going down any column:
1) As n increases, outershell e- are further from the nucleus.
2) As n increases, there are also more innershells of electrons shielding the protons from the electrons. 20 20 20

21. Consider N and P radii : 74 pm 110 pm outershell: 2s2 2p3 3s2 3p3
innershell shielding of nuclear charge:
1s s2
2s2 2p6 7+
15+
extra shielding 21

22. To explain why atomic size increases going from right to left in the periodic table,
As you move right to left of any row, there is a decrease in the # of protons, which leads to a decrease in charge in the nucleus (the nuclear charge) 22 22 22

23. Outershell is the same (n = 2) Innershell is the same (1s2)
Consider N and F
radii: pm pm 7+ 9+
Outershell is the same (n = 2)
Innershell is the same (1s2)
What is different?
# protons = # nuclear charge 23

24. Higher nuclear charge pulls electrons closer to the nucleus.
Consider N and F
radii: pm pm
nuclear charge 7 p p+
Higher nuclear charge pulls electrons closer to the nucleus. 24

25. Ion vs. Atom Radii
Positive ions are always smaller than the atoms from which they are formed
due to decreased shielding effects
Negative ions always larger than the atoms from which they are formed
due to increased electron repulsion
FIG Changes in size when atoms gain or lose electrons to form ions. Adding electrons leads to an increase in the size of the particle, as illustrated for fluorine. Removing electrons leads to a decrease in the size of the particle, as shown for lithium and iron. 25

26. Ionization Energy of Hydrogen
Ionization energy refers to the energy involved in pulling an electron off a neutral atom.
For the H atom, what kind of “transition” would that electron be undergoing?
Based on Bohr’s atomic theory, what would this ionization energy be?
Experimental we find for H, the IE = 1310 kJ/mol.
Why is there a discrepancy?

27. Ionization energy (IE)
IE is the energy added to remove an electron from an isolated, gaseous atom
Successive ionizations are possible until no electrons remain
Trends in IE are the opposite of the trends in atomic size
Why?
Valence electrons are closer to the nucleus for the smaller atoms and are held tighter by the nuclear charge.
They are harder to pull off: HIGHER IE 27

28. Trends in IE
FIG Variation in first ionization energy with location in the periodic table. Elements with the largest ionization energies are in the upper
right of the periodic table. Those with the smallest ionization energies are at the lower left. 28

29. Successive Ionization Energy (IE)
Mg (g) + energy Mg+ (g) + e-
Mg+ (g) + energy Mg2+ (g) + e-
Mg2+ (g) + energy Mg3+ (g) + e-
1st IE = kJ/mol
2nd IE = 1450 kJ/mol
3rd IE = 7731 kJ/mol
2nd IE > 1st IE Why?
3rd IE >>>> 2nd IE Why?
1st IE
2nd IE
3rd IE
713 kJ/mol
6281 kJ/mol 29

30. Successive IE Why the sudden increase? 30
FIG Variations in successive ionization energies for the elements lithium through fluorine. 30

31. Why the sudden jump in IE's?31

32. Irregularities in I.E.
Easier to pull e- off B than from Be b/c it's from 2p not 2s.
Easier to pull e- off O than from N b/c it’s from orbital with 2 e- instead of 1 e- 32

33. Periodic Trend in the Ionization Energy
highest IE
Nobel gases are excluded.
lowest IE

34. Electron Affinity (EA)
Definition:
Electron affinity is the energy released when an electron is added to a neutral gaseous atom.
EA of the 1st electron is almost always negative (exothermic)
F + e F- + energy
(EA of F is -328 kJ/mole.)
Addition of subsequent electrons always requires energy (EA's are positive).
We will focus on only the 1st EA. 34 34

35. We talk about how negative the electron affinity is
We talk about how negative the electron affinity is. The more an element likes electrons the more negative is the EA.
It gets a bit confusing with the negative sign, so we sometimes refer to the absolute value of the EA. This gives us the "magnitude" of the EA. 35 35

36. For example, EA of O is -141 kJ/mol EA of F is -328 kJ/mol
EA of F is more negative than O which means it likes electrons more.
Technically, the EA of F is smaller because it is a smaller number.
It is misleading to say F has a lower electron affinity, when it actually has a higher "affinity" for electrons.
To avoid the confusion, we talk in terms of the magnitude of O is 141 kJ/mol,
and the magnitude of F is 328 kJ/mol. 36

37. Periodic Trend of the Electron Affinity
F has the largest EA in magnitude. 37
lowest EA in magnitude

38. Summary of the periodic trends
Draw in the arrows to show the trends for…
1) atomic size 2) IE 3) magnitude of EA
What generalization can you make about the IE of metals and nonmetals? 38

39. Compare the electron affinity for metals vs. nonmetals
What generalizations can you make of the metalloids concerning IE and EA? 39

40. Metals vs. Nonmetals Metals have low ionization energies.
It is easier to strip off electrons from metals.
They tend to form CATIONS.
Nonmetals have high magnitude of electron affinities.
It is easier to add electrons to nonmetals.
They tend to form ANIONS.
If there is no source of electrons to be added, they tend to SHARE electrons and form COVALENT BONDS (as molecules or polyatomic ions). 40