1. Periodic Table Chapter 8 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

3. Non metal Metal

4. Outermost subshell being filled with electrons 7.8

5. 8.2 ns 1 ns 2 ns 2 np 1 ns 2 np 2 ns 2 np 3 ns 2 np 4 ns 2 np 5 ns 2 np 6 d1d1 d5d5 d 10 4f 5f Ground State Electron Configurations of the Elements

6. 8.2 Classification of the Elements

7. Electron Configurations of Cations and Anions Na [Ne]3s 1 Na + [Ne] Ca [Ar]4s 2 Ca 2+ [Ar] Al [Ne]3s 2 3p 1 Al 3+ [Ne] Atoms lose electrons so that cation has a noble-gas outer electron configuration. H 1s 1 H - 1s 2 or [He] F 1s 2 2s 2 2p 5 F - 1s 2 2s 2 2p 6 or [Ne] O 1s 2 2s 2 2p 4 O 2- 1s 2 2s 2 2p 6 or [Ne] N 1s 2 2s 2 2p 3 N 3- 1s 2 2s 2 2p 6 or [Ne] Atoms gain electrons so that anion has a noble-gas outer electron configuration. Of Representative Elements 8.2

8. +1+2+3 -2-3 Cations and Anions Of Representative Elements 8.2

9. Na + : [Ne]Al 3+ : [Ne]F - : 1s 2 2s 2 2p 6 or [Ne] O 2- : 1s 2 2s 2 2p 6 or [Ne]N 3- : 1s 2 2s 2 2p 6 or [Ne] Na +, Al 3+, F -, O 2-, and N 3- are all isoelectronic with Ne What neutral atom is isoelectronic with H - ? H - : 1s 2 same electron configuration as He 8.2

10. Electron Configurations of Cations of Transition Metals 8.2 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]4s 2 3d 6 Fe 2+ : [Ar]4s 0 3d 6 or [Ar]3d 6 Fe 3+ : [Ar]4s 0 3d 5 or [Ar]3d 5 Mn: [Ar]4s 2 3d 5 Mn 2+ : [Ar]4s 0 3d 5 or [Ar]3d 5

11. Effective nuclear charge (Z eff ) is the “positive charge” felt by an electron. Na Mg Al Si 11 12 13 14 10 1 2 3 4 186 160 143 132 Z eff Core Z Radius Z eff = Z -  0 <  < Z (  = shielding constant) Z eff  Z – number of inner or core electrons Within a Period as Z eff increases radius decreases 8.3

13. Atomic Radii 8.3 Alkaline Metals (1A) Halides (7A)

14. 8.3 Anions Neutral Atoms Neutral Atoms Cations Cations are smaller Anions are larger

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

16. isoelectronic Smaller Cations Big Anions

17. Ionization energy is the minimum energy (kJ/mol) required to remove one electron each from a mole of gaseous atoms/ions in the ground state. I 1 + X (g) X + (g) + e - I 2 + X + (g) X 2 + (g) + e - I 3 + X 2+ (g) X 3 + (g) + e - I 1 first ionization energy I 2 second ionization energy I 3 third ionization energy 8.4 I 1 < I 2 < I 3 Endothermic: always > 0

18. Filled n=1 shell Filled n=2 shell Filled n=3 shell Filled n=4 shell Filled n=5 shell 8.4 Noble Gases (8A) Alkaline Metals (1A) “High IE  hard to remove Electrons” “Low IE  easy to remove Electrons  cations formed”

19. General Trend in First Ionization Energies 8.4 Increasing First Ionization Energy LOW High

20. Electron affinity is the negative of the energy change that occurs when one electron each is accepted by a mole of atoms in the gaseous state to form a mole of anions. X (g) + e - X - (g) 8.5 F (g) + e - F - (g) O (g) + e - O - (g)  H = -328 kJ/mol EA = +328 kJ/mol  H = -141 kJ/mol EA = +141 kJ/mol Exothermic

21. 8.5 Halides (7A) Alkaline Metals (1A) Anions Cations

22. 8.6 Remember size variations

23. Group 1A Elements (ns 1, n  2) M M +1 + 1e - 2M (s) + 2H 2 O (l) 2MOH (aq) + H 2(g) 4M (s) + O 2(g) 2M 2 O (s) Increasing reactivity 8.6 Cations (+1) Basic Stable oxides (strongly ionic) Hydrogen displacement in cold water Alkaline Metals

24. Group 2A Elements (ns 2, n  2) M M +2 + 2e - Be (s) + 2H 2 O (l) No Reaction Increasing reactivity 8.6 Mg (s) + 2H 2 O (g) Mg(OH) 2(aq) + H 2(g) M (s) + 2H 2 O (l) M(OH) 2(aq) + H 2(g) M = Ca, Sr, or Ba Alkaline Earth Metals Cations (+2) Basic Hydrogen displacement Many metals react with acid

25. Group 3A Elements (ns 2 np 1, n  2) 8.6 4Al (s) + 3O 2(g) 2Al 2 O 3(s) 2Al (s) + 6H + (aq) 2Al 3+ (aq) + 3H 2(g) Amphoteric

26. Group 7A Elements (ns 1 np 5, n  2) X + 1e - X - 1 X 2(g) + H 2(g) 2HX (g) Increasing reactivity 8.6 Halides Acidic amphotericBasic

27. Oxides can be classified as acids, bases, or both (amphoteric) Na 2 O (s) + H 2 O (l) 2NaOH (aq) MgO (s) + 2HCl (aq) MgCl 2(aq) + H 2 O (l) Al 2 O 3(s) + 6HCl (aq) 2AlCl 3(aq) + 3H 2 O (l) Al 2 O 3(s) + 2NaOH (aq) + 3H 2 O (l) 2NaAl(OH) 4(aq) acid base acid base salt Al 2 O 3 is amphoteric!

28. P 4 O 10(s) + 6H 2 O (l) 4H 3 PO 4(aq) base salt acid SiO 2(s) + 2NaOH (aq) Na 2 SiO 3(aq) + H 2 O (l) SO 3(g) + H 2 O (l) H 2 SO 4(aq) Cl 2 O 7(s) + H 2 O (l) 2HClO 4(aq) acid Normal metallic oxides are basic, and most non-metallic oxides are acidic. As the metallic character of elements increase, so does the basicness of their oxides.