1. Lecture 0802 Trends on the Periodic Table

2. PERIODIC TRENDS Li Na K

3. Effective Nuclear Charge Z* The 2s electron PENETRATES the region occupied by the 1s electron. 2s electron experiences a higher positive charge than expected.

4. Effective Nuclear Charge, Z* u AtomZ* e - in Valence Orbitals estmeasured u Li1+1.28 u Be2------- u B3+2.58 u C4+3.22 u N5+3.85 u O6+4.49 u F7+5.13 Increase in Z* across a period

5. General Periodic Trends u Atomic and ionic size u Ionization energy u Electron affinity Higher effective nuclear charge Electrons held more tightly Larger shells. Electrons held less tightly.

6. Atomic Radius u Is taken as the covalent radius for non-metallic elements and as the metallic radius for metals

7. Atomic Radius u Covalent radius is one-half the distance between the nuclei of two identical atoms that are singly bonded to one another. Chlorine Bond Length

8. Atomic Radius u Covalent radii for elements whose atoms do not bond to one another can be estimated by combining radii of those that do with the distances between unlike atoms in various molecules.

10. Atomic Radius u Metallic radius is one-half the closest internuclear distance in a metallic crystal.

11. Prediction!

14. Atomic Size u Size goes UP on going down a group. u Because electrons are added further from the nucleus, there is less attraction. u Size goes DOWN on going across a period. u Size goes UP on going down a group. u Because electrons are added further from the nucleus, there is less attraction. u Size goes DOWN on going across a period.

15. Atomic Size Size decreases across a period owing to increase in Z*. Each added electron feels a greater and greater + charge. Large Small

16. Trends in Atomic Size

17. Sizes of Transition Elements u 3d subshell is inside the 4s subshell. u 4s electrons feel a more or less constant Z*. u Sizes stay about the same and chemistries are similar!

18. General Periodic Trends u Atomic and ionic size u Ionization energy u Electron affinity Higher effective nuclear charge Electrons held more tightly Larger shells. Electrons held less tightly.

19. Ion Sizes Does the size go up or down when losing an electron to form a cation? Does the size go up or down when losing an electron to form a cation?

20. Ion Sizes u CATIONS are SMALLER than the atoms from which they come. u The proton/electron attraction has gone UP and so size DECREASES. Li,152 pm 3e and 3p Li +, 78 pm 2e and 3 p + Forming a cation.

21. Ion Sizes Does the size go up or down when gaining an electron to form an anion?

22. Ion Sizes u ANIONS are LARGER than the atoms from which they come. u The proton/electron attraction has gone DOWN and so size INCREASES. Forming an anion. F, 71 pm 9e and 9p F -, 133 pm 10 e and 9 p -

23. Trends in Ion Sizes Trends in ion sizes are the same as atom sizes.

24. Redox Reactions Why do metals lose electrons in their reactions? Why does Mg form Mg 2+ ions and not Mg 3+ ? Why do nonmetals take on electrons? Why do metals lose electrons in their reactions? Why does Mg form Mg 2+ ions and not Mg 3+ ? Why do nonmetals take on electrons?

25. Ionization Energy (General) u Is the energy required to remove the outermost electron from an atom or a positive ion in the ground state.

26. First Ionization Energy u Energy required to remove the first electron from a neutral atom in the gaseous state.

27. Ionization Energy Mg (g) + 738 kJ  Mg + (g) + e -

28. Prediction!

30. Trends in Ionization Energy

32. Atomic Radii

34. Trends in Ionization Energy u IE increases across a period because Z* increases. u Metals lose electrons more easily than nonmetals. u Metals are good reducing agents. u Nonmetals lose electrons with difficulty.

35. Trends in Ionization Energy u IE decreases down a group u Because size increases. u Reducing ability generally increases down the periodic table. u Remember Li, Na, K

36. Second Ionization Energy u Energy needed to remove the outermost electron from a +1 ion. u Energy needed to remove the second electron from a neutral atom.

37. Ionization Energy Mg (g) + 738 kJ  Mg + (g) + e - Mg + (g) + 1451 kJ  Mg 2+ (g) + e Mg + (g) + 1451 kJ  Mg 2+ (g) + e - Mg + has 12 protons and only 11 electrons. Therefore, IE for Mg + > Mg.

39. Ionization Energy Mg (g) + 735 kJ  Mg + (g) + e - Mg + (g) + 1451 kJ  Mg 2+ (g) + e - Mg 2+ (g) + 7733 kJ  Mg 3+ (g) + e - Energy cost is very high to dip into a shell of lower n. This is why ox. no. = Group no.

40. General Periodic Trends u Atomic and ionic size u Ionization energy u Electron affinity Higher effective nuclear charge Electrons held more tightly Larger shells. Electrons held less tightly.

41. Electron Affinity u A few elements GAIN electrons to form anions. u E.A. is the energy released or absorbed when an electron is added to the valence level of a gas-phase atom. u A(g) + e -  A - (g) E.A. = ∆E

42. Prediction!

43. Trends in Electron Affinity

44. Electron Affinity of Oxygen ∆E is EXOthermic because O has an affinity for an e -. [He]      O atom EA = - 141 kJ + electron O [He]       - ion

45. Electron Affinity of Nitrogen ∆E is zero for N - due to electron- electron repulsions. EA = 0 kJ [He]     Natom  [He]    N - ion  + electron

46. u See Figure 8.12 and Appendix F u Affinity for electron increases across a period (EA becomes more negative). u Affinity decreases down a group (EA becomes less negative). Atom EA F-328 kJ Cl-349 kJ Br-325 kJ I-295 kJ Atom EA F-328 kJ Cl-349 kJ Br-325 kJ I-295 kJ Trends in Electron Affinity

48. General Periodic Trends u Atomic and ionic size u Ionization energy u Electron affinity Higher effective nuclear charge Electrons held more tightly Larger shells. Electrons held less tightly.