1. Periodic Relationships Among the Elements
Chapter 8
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2. Ground State Electron Configurations of the Elements
ns2np6
Ground State Electron Configurations of the Elements
ns1
ns2np1
ns2np2
ns2np3
ns2np4
ns2np5
ns2
d10 d1 d5 4f 5f
8.2

3. Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Atoms lose electrons so that cation has a noble-gas outer electron configuration.
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
H 1s1
H- 1s2 or [He]
Atoms gain electrons so that anion has a noble-gas outer electron configuration.
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]
8.2

4. Cations and Anions Of Representative Elements+1 +2 +3 -3 -2 -1
8.2

5. Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
Na+: [Ne]
Al3+: [Ne]
F-: 1s22s22p6 or [Ne]
O2-: 1s22s22p6 or [Ne]
N3-: 1s22s22p6 or [Ne]
Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
What neutral atom is isoelectronic with H- ?
H-: 1s2
same electron configuration as He
8.2

6. Electron Configurations of Cations of Transition Metals
When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals.
Fe: [Ar]4s23d6
Mn: [Ar]4s23d5
Fe2+: [Ar]4s03d6 or [Ar]3d6
Mn2+: [Ar]4s03d5 or [Ar]3d5
Fe3+: [Ar]4s03d5 or [Ar]3d5
8.2

7. 0 < s < Z (s = shielding constant)
Effective nuclear charge (Zeff) is the “positive charge” felt by an electron.
Zeff = Z - s
0 < s < Z (s = shielding constant)
Zeff  Z – number of inner or core electrons
Zeff
Core Z
Radius (pm) Na Mg Al Si 11 12 13 14 10 1 2 3 4
186
160
143
132
8.3

8. Effective Nuclear Charge (Zeff)
increasing Zeff
increasing Zeff
8.3

9. 8.3

10. 8.3

11. Atomic Radii
8.3

12. Comparison of Atomic Radii with Ionic Radii
8.3

13. Cation is always smaller than atom from which it is formed.
Anion is always larger than atom from which it is formed.
8.3

14. The Radii (in pm) of Ions of Familiar Elements
8.3

15. I1 first ionization energy
Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state.
I1 + X (g) X+(g) + e-
I1 first ionization energy
I2 + X+(g) X2+(g) + e-
I2 second ionization energy
I3 + X2+(g) X3+(g) + e-
I3 third ionization energy
I1 < I2 < I3
8.4

16. 8.4

17. Variation of the First Ionization Energy with Atomic Number
Filled n=1 shell
Filled n=2 shell
Filled n=3 shell
Filled n=4 shell
Filled n=5 shell
8.4

18. General Trend in First Ionization Energies
Increasing First Ionization Energy
Increasing First Ionization Energy
8.4

19. Ionization Energy Trends DOWN a Group
Ionization Energy decreases as you move down a group because the electrons that are being lost are in successively higher energy levels – further from the nucleus – requiring less energy to remove them.
(Note – although the nuclear charge is greater (more protons), there is more shielding because of the additional electrons… so the overall effective nuclear charge doesn’t really change)

20. Ionization Energy Trends DOWN a Group
Which has a higher ionization energy, Na or Rb? Explain why.
Which has a higher ionization energy, Br or Cl? Explain why.

21. Ionization Energy Trends ACROSS a Period
Ionization Energy increases as you move across a period because there is an increase in the charge on the nucleus, which increases the effective nuclear charge – requiring more energy to remove electrons from the same energy level.
Note: The increase is not totally smooth. Between group 2 and 13 there is a ‘dip’ (decrease) in ionization energy because the electron removed from the element in group 13 is in a p-subshell which is slightly higher in energy than the s-subshells of group 2 and 1. This electron is slightly higher in energy already, so it takes less energy to remove it than it does to remove an electron from the s-subshell.
A similar glitch occurs between groups 15 and 16. In group 15 the element has only 1 electron in each p-subshell. In group 16, one of the p-subshells has a pair of electrons. The resulting greater electron-electron repulsion counteracts the effect of the increased nuclear charge... which lowers the activation energy.

22. Ionization Energy Trends ACROSS a Period
Which has a lower ionization energy, Na or Cl? Explain why.
Which has a higher ionization energy, B or Be? Explain why.
Which has a higher ionization energy, O or N? Explain why.

23. Electronegativity Trends – Why?
Down a Group – EN decreases
Valence electrons further from nucleus
Effective nuclear charge isn’t changing
Atom’s ability to attract an electron further away from the nucleus is diminishing
Across a Period – EN increases
Charge on nucleus increases (more protons)
Electrons in valence shell are more attracted to nucleus – closer
Atom’s ability to attract an electron is increased due to closer distance to nucleus and increase effective nuclear charge.

24. Electronegativity Trends – Why?
Which has a higher electronegativity, Na or Rb? Explain why.
Which has a higher electronegativity, Br or Cl? Explain why.

25. Melting Point Trends – Why?
Down a Metal Group – MPt decreases slightly
Valence electrons further from nucleus
Effective nuclear charge isn’t changing
Atom’s ability to attract an electron further away from the nucleus is diminishing
Across a Period – MPt increases… then decreases (MPt reflects strength of forces between particles in solid vs liquid states.
Gps 1&2 - Strength of metallic bond increases with inc nuclear charge, inc # mobile valence electrons, atomic radius decreases
Gps 14 – giant covalent structures… v. strong covalent bonding in all directions
Gps – molecular structures… nonpolar… only weak van der Waals forces… mpt depends on mass and size of molecules
Gp 18 – single atoms, v wk vdWaals… very low mpt

26. Melting Point Trends – Why?
Which has a higher melting point, Na or Rb? Explain why.
Which has a higher melting point, Na or Cl2? Explain why.
Which has a higher melting point, C or O2? Explain why.

27. Ionization Energy, Electronegativity, Atomic Radius, Ionic Radius, Melting Point
Periodic Trends