Periodic Relationships Among the Elements Chapter 8 Copyright © The McGraw-Hill Companies, Inc.  Permission required for reproduction or display.

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

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

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

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

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

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

Effective Nuclear Charge (Zeff) increasing Zeff increasing Zeff 8.3

8.3

8.3

Atomic Radii 8.3

Comparison of Atomic Radii with Ionic Radii 8.3

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

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

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

8.4

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

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

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)

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.

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.

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.

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.

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

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 15-16-17 – 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

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.

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