1. Periodic Properties of the Elements
Chapter 7
Periodic Properties of the Elements

2. Electron Configuration exceptions
Energy cost to sticking two electrons in the same orbital
Observation: sublevel with all half-filled orbitals energetically more stable
Observation: sublevel with completely filled orbitals energetically more stable

3. Exceptions to electron configurations
Observed:
Cr: [Ar]4s13d5
Cu: [Ar]4s13d10
Only see these in transition metals
Less energy difference between 4s and 3d

4. Exceptions to electron configurations
Transition metals, when forming ions, lose s electrons first, NOT d!!
Zn = [Ar]4s23d10
Zn+2 = [Ar]3d10
Why?

5. Charge interactions in an atom
The closer + and – get, the stronger their attraction
More + = stronger attraction for –
Implications: behavior of electrons on outside of atom will depend on
How close those electrons are, relatively, to the nucleus
How strong a pull those electrons experience from the nucleus

6. Effective Nuclear Charge
Effective nuclear charge: number of p+ in nucleus (Z) minus average number of e-(S) between nucleus and e- in question
Zeff = Z – S
Positive charge acting on outer e- is less than positive charge acting on closer e-
As distance from nucleus increases, S increases and Zeff decreases
The increase in electron shells (PELs) shields the nucleus, thus decreasing the attraction between the nucleus and the valence electrons

7. How orbitals differ More positive the nucleus, smaller the orbital
Na 1s orbital = same shape as H 1s orbital,
but is smaller b/c e- more strongly attracted to nucleus!
“Greater effective nuclear charge!”

8. Atomic Size Have to measure distance between nuclei in diatomic – why?
What pattern do you see?

9. Atomic Size (Trends)
Within a group, atomic radius increases from top to bottom
Adding energy levels of electrons, electrons pulling further away from nucleus
Within a period, atomic radius decreases from left to right
Greater effective nuclear charge within the same energy level – orbitals are shrinking in comparison to one another
Stronger nucleus is pulling the electrons in closer

10. Ionic Radius For metals: For nonmetals:
Positively charged ions (cations) formed when an atom of a metal loses one or more electrons
Smaller than the corresponding atom
Ex: Na+ is smaller than Na
For nonmetals:
Negatively charged ions (anions) formed when an atom of a nonmetal gains one or more electrons
Larger than the corresponding atom
Ex: Cl- is larger than Cl

11. 1. Arrange the following atoms in order of increasing atomic radius: Na, Be, Mg.
Practice Exercise
Answer:
Be < Mg < Na
2. Arrange these atoms and ions in order of decreasing size: Mg2+, Ca2+, and Ca.
Answer:
Ca > Ca+2 > Mg+2
3. Arrange the ions K+, Cl–, Ca2+, and S2– in order of decreasing size.
Answer: S2– > Cl– > K+ > Ca2+.

12. Ionization Energy
Energy required to remove an electron from a gaseous atom
Highest energy electron removed first
Furthest away
First IE (I1) = removing first e-
Second IE (I2) = removing second e-
etc. etc.

13. IE
for Mg
I1 = 735 kJ/mole
I2 = 1445 kJ/mole
I3 = 7730 kJ/mole
Effective nuclear charge increases as you remove electrons – increase proton to electron ratio
Takes much more energy to remove core e- than valence e-
b/c core e- feeling more pull from nucleus (greater effective nuclear charge)

14. IE

15. Explain this trend in IE
For Element X – how many valence electrons?
I1 = 580 kJ/mole
I2 = 1815 kJ/mole
I3 = 2740 kJ/mole
I4 = 11,600 kJ/mole
I5 = 16,430 kJ/mol

16. IE Trends
Across a Period: Generally from left to right, I1 increases because there is a greater nuclear charge within the same energy level
Down a Group: As you go down a group, I1 decreases because electrons are farther away from nucleus

17. IE Trend

18. BUT WAIT! Some exceptions!
Explain the exceptions:
Half filled and completely filled sublevels are harder to remove electrons from!

19. IE Examples
Which will have the greater third ionization energy, Ca or S?
Answer: Ca
Arrange the following atoms in order of increasing first ionization energy: Ne, Na, P, Ar, K.
Answer: K < Na < P < Ar < Ne

20. IE and PES
PES is not in your book, pay careful attention to the notes!!!!
Photoelectric effect
Energy of an ejected electron:
Etotal = KE(kinetic energy of e) + IE(ionization energy of e)

21. What happens when an electron is ejected?
Vacuum level
Vacuum level -
I.E.
Bound State
Bound State +
Metal (initial) - M+
K.E.
Vacuum level
I.E.
Bound State + M+

22. Photoelectron Spectroscopy (PES)
Machine irradiates metal with energy
UV or X-ray
Measures the kinetic energy of electrons when they “pop off”
By measuring the total energy irradiating a metal (E) and the kinetic energy of the ejected electrons (KE), ionization energy can be determined for ANY e- (not just outermost!)
Etotal = IE + KE
Like mass spec, peaks in PES data

23. How to read PES data: B atom
Energy (MJ/mol)
Relative number of electrons 70 65 25 20 15 10 5
3 peaks = removing e- didn’t happen all at once
Implies e- in 3 locations
X-axis: energy required to remove the electrons
often inverted (highest energy on left)
BUT NOT NECESSARILY! ALWAYS CHECK AXIS!
Y-axis: height of peak corresponds with ratio of electrons
not precisely saying nof. e-, but gives ratio

24. How to read PES data: B atom
Energy (MJ/mol)
Relative number of electrons 70 65 25 20 15 10 5
Look at ratio of height of peaks on x axis
2:2:1
Which electrons hardest to remove? What does this mean?
Hardest to remove = closest to nucleus

25. How to read PES data: B atom
AND relative heights tell you 2e- in 1s, 2 e- in 2s, 1e- in 2p
Relative number of electrons
1s22s22p1 70 65 25 20 15 10 5
Energy (MJ/mol)
2p sublevel
(easiest to remove)
1s sublevel
(hardest to remove)
2s sublevel
(a bit easier to remove)

26. How to read PES data: B atom
Relative number of electrons 70 65 25 20 15 10 5
Energy (MJ/mol)
Therefore, it is Boron
5 electrons
1s22s22p1

27. Photoelectron Spectra

28. Why PES?
Look at what graph indicates: not all electrons in same energy level have same ionization energy
What conclusions can be drawn from this?
Sublevels!
PES = evidence for quantum mechanical model of the atom
Evidence for s,p,d,f sublevel

29. PES Example 4 peaks = 4 locations of electrons = 4 energy sublevels
Relative number of electrons
Energy (MJ/mol)
127
125 9 7 2
4 peaks = 4 locations of electrons = 4 energy sublevels
Relative height of peaks tells you how many e- per sublevel
Peak with highest energy = 1s, then 2s, 2p, 3s
Which element is this?
Mg: 1s22s22p63s2

30. PES Draw what you would expect the PES data for oxygen to look like
Energy (MJ/mol)
Relative number of electrons

31. Electronegativity and Electron Affinity
Electronegativity: The ability of an atom to attract an electron to itself
Scale of 0.0 – 4.0
Electron affinity: energy change associated when an electron is added to a gaseous atom
Measured in kJ/mol
The more negative, the greater the attraction of the atom for an electron
If EA > 0, REALLY does not want an e-, negative ion higher in energy than separated atom and electron

32. Electronegativity: Trend
EN increases left to right in a period
Decreases from top to bottom
F is most EN element on periodic table
EA is different – does not change down a group, but does increase (gets more negative) across the p-block

33. IE vs. EA – the final showdown
IE: energy required to take an electron AWAY!
EA: energy associated with gaining an electron!

34. Metals vs. Nonmetals

35. Alkali Metals Group 1 As you move down the group Oxidize easily
Decrease in IE
Decrease in mp
Increase in metallic character
Oxidize easily
React vigorously with water to produce metal hydroxide and hydrogen gas, highly exothermic
Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)

36. Alkaline Earth Metals Group 2
Reactive, but not as reactive as group 1 metals

37. Halogens Group 17 (or 7A) High EA High IE High eneg
F is the most reactive element on PT – high eneg, removes electrons from pretty much anything
React with H2 to make acids
H2(g) + X2 → 2HX(g)

38. Noble Gases Group 18 (8A) Stable, monatomic
Completely filled s and p sublevels
Were called “inert gases” until 1962
Have reacted under extreme circumstances
Compounds of all noble gases have been created
He was the last one – HeF2 in 1998 – why the last? Why F

39. Allotropes
Different forms of the same element in the same physical state
Oxygen: O2 (oxygen gas) and O3 (ozone)
Carbon: Diamond and Graphite
Sulfur:
Have different structures and different physical and chemical properties