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

2. 2 When the Elements Were Discovered

3. Contributors to our Periodic Table Newland – when elements were ordered according to atomic mass, every 8 th element had similar properties Mendeleev – extensive tabulation by atomic mass Mosely – used X-rays emitted from an element to order elements by atomic number 3

4. 4 Classification of the Elements

5. 5 Effective nuclear charge (Z eff ) is the “positive charge” felt by an electron. Z eff = Z -  0 <  < Z (  = shielding constant) Z eff  Z – number of inner or core electrons

6. 6 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 (pm) Z eff = Z -  0 <  < Z (  = shielding constant) Z eff  Z – number of inner or core electrons http://www.oneonta.edu/faculty/viningwj/sims/effective_nuclear_charge_s.html

7. 7 Effective Nuclear Charge (Z eff ) increasing Z eff http://www.oneonta.edu/faculty/viningwj/sims/effective_nuclear_charge_s.html

8. 8 Atomic Radii metallic radius covalent radius

9. 9

10. 10 Trends in Atomic Radii

11. 11 Comparison of Atomic Radii with Ionic Radii

12. 12 Cation is always smaller than atom from which it is formed. Anion is always larger than atom from which it is formed.

13. Put the ionic radii of the following elements in order from smallest to largest Sulfur, chlorine, potassium, calcium 13

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

15. 15 Chemistry in Action: The 3 rd Liquid Element? Liquid? 117 elements, 2 are liquids at 25 0 C – Br 2 and Hg 223 Fr, t 1/2 = 21 minutes

16. 16 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 2+ (g) X 3 + (g) + e - I 1 first ionization energy I 2 second ionization energy I 3 third ionization energy I 1 < I 2 < I 3

17. 17

18. 18 Filled n=1 shell Filled n=2 shell Filled n=3 shell Filled n=4 shell Filled n=5 shell Variation of the First Ionization Energy with Atomic Number What is the trend as we go across periods? What is the trend as we go down groups?

19. 19 General Trends in First Ionization Energies Increasing First Ionization Energy

20. 20 Variation of the First Ionization Energy with Atomic Number

21. 21 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) 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 The more positive the electron affinity (EA), the greater the tendency to accept the electron. What does this tell you about stability of the ion it forms?

22. 22

23. 23 Variation of Electron Affinity With Atomic Number (H – Ba)

24. 24 Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond. Electron Affinity - measurable, Cl is highest Electronegativity - relative, F is highest X (g) + e - X - (g)

25. 25 The Electronegativities of Common Elements

26. 26 Variation of Electronegativity with Atomic Number

27. General Trends Metals usually have low/high ionization energies Nonmetals have low/high electron affinities As we move down group 5A, what happens to ionization energies? And how does that relate to the metallic character of the elements? Groups 3A to 6A have the greatest variation within the group compared to Groups 1A/2A and Groups 7A/8A. Why? 27

28. 28 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 http://www.youtube.com/watch?v=QSZ-3wScePM

29. 29 Group 1A Elements (ns 1, n  2)

30. 30 Group 8A Elements (ns 2 np 6, n  2) Completely filled ns and np subshells. Highest ionization energy of all elements. No tendency to accept extra electrons.

31. 31 Compounds of the Noble Gases A number of xenon compounds XeF 4, XeO 3, XeO 4, XeOF 4 exist. A few krypton compounds (KrF 2, for example) have been prepared.

32. 32 The metals in these two groups have similar outer electron configurations, with one electron in the outermost s orbital. Chemical properties are quite different due to difference in the ionization energy. Comparison of Group 1A and 1B Lower I 1, more reactive

33. 33 Properties of Oxides Across a Period basicacidic

34. 34

35. 35 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

36. 36 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 valence 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 valence electrons so that anion has a noble-gas outer electron configuration. Of Representative Elements

37. 37 +1+2+3 -2-3 Cations and Anions Of Representative Elements

38. 38 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 Isoelectronic: have the same number of electrons, and hence the same ground-state electron configuration

39. 39 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]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 What are the electron configurations of Cu +2 and Cu +1 ?