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

2. 1. Explain the basis of the periodic table as described by Mendeleev and Meyer and indicate the shortcomings of their method. 2. Explain the basis of the periodic table as described by Moseley and how it predicted properties of “ missing ” elements. 3. Identify elements that correspond to each of the following groups: · * representative elements · * noble gases · * transition metals · * lanthanides · * actinides 4. Describe the electron configuration of cations and anions and identity ions and atoms that are isoelectronic.

3. 5. Apply the concept of effective nuclear charge and shielding constants (screening constants) to justify why the first ionization energy is always smaller than the second ionization energy of a given atom. 6. Predict the trends from left to right and top to bottom of the periodic table for each of the following: · atomic radius · ionic radius · ionization energy · electron affinity · metallic character 7. Relate why hydrogen could be placed in a class by itself when reviewing its chemical properties. 8. Provide examples of Group 1A elements reacting with oxygen to form oxides, peroxides, and superoxides. 9. Predict the reaction of alkali metals with water. 10. Describe the reactivity of alkaline earth metals with water.

4. 11. Relate how strontium-90 could lead to human illness. 12. Compare the reactivity of boron, a metalloid, to aluminum. 13. Identify the metals, nonmetal, and metalloids of Group 4A. 14. Recall the reactions that form nitric acid, phosphoric acid and sulfuric acid. 15. List the halides (halogens) 16. Indicate the three hydrohalic acids that are strong acids and the one hydrohalic acid that is a weak acid. 17. Explain why the name for Group 8A has changed from inert gases to noble gases. 18. List the three “ coinage ” metals and explain their relative inertness.

5. 19. Rationalize the characteristics of the properties of oxides of the third period elements. 20. Classify oxides as acidic, basic, or amphoteric. 21. Explain why concentrated bases such as NaOH should not be stored in glass containers

7. 8.1

9. 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 4f4f 5f5f Ground State Electron Configurations of the Elements

11. 8.2

12. 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

13. +1+2+3 -2-3 Cations and Anions Of Representative Elements 8.2

14. 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

15. 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

16. 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 For a Period as Z eff increases radius decreases 8.3

20. Atomic Radii 8.3

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

25. Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state. I 1 + X (g) X + (g) + e - I 2 + X (g) X 2 + (g) + e - I 3 + X (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

32. Filled n=1 shell Filled n=2 shell Filled n=3 shell Filled n=4 shell Filled n=5 shell 8.4

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

34. Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion. X (g) + e - X - (g) 8.5 F (g) + e - X - (g) O (g) + e - O - (g)  H = -328 kJ/mol EA = +328 kJ/mol  H = -141 kJ/mol EA = +141 kJ/mol

35. 8.5

37. 8.6

38. Group 1A Elements (ns 1, n  2) M M +1 + 1e - 2M (s) + 2H 2 O (l) 2MOH (aq) + H 2(g) 2M (s) + O 2(g) M 2 O (s) Increasing reactivity 8.6

40. 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

42. 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)

47. 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