Periodic Table and Periodic Trends

1. Periodic Table and Periodic Trends Test 2 Material Along with Atomic Structure: 16 out of 75 MC questions Free Response every year

2. Broad Periodic Table Classifications • Representative Elements(main group): filling s and p orbitals (Na, Al, Ne, O) • Transition Elements: filling dorbitals (Fe, Co, Ni) • Lanthanide and Actinide Series(inner transition elements): filling 4f and 5f orbitals (Eu, Am, Es)

3. Periodicity • Trends in atomic radii, ionization energies, and electron affinities are all tied to two basic structural features of the atoms: • Effective nuclear charge (Zeff) • Principal energy level of valence electrons

4. Periodic Trends • Explained in terms of: • Increasing/Decreasing nuclear charge (Z) • Increasing/Decreasing distance from nucleus (n) • Increasing/Decreasing shielding by core electrons

5. Ionization Energy The quantity of energy required to remove an electron from the gaseous atom or ion—form a positive ion.

6. Low ionization energy—electron is easy to remove • High ionization energy—electron is difficult to remove

7. Periodic Trends • First ionization energy: • increasesfrom left to right across a period; higher nuclear charge • decreases going down a group; electrons farther from nucleus

8. Values of first ionization energy for the elements in the first six periods Page 326

9. First Ionization Energy • Notice that the trend is not smooth. • First ionization energy decreases between s- and p- orbital filling and just after p-orbital is half full. • Extra stability at these points

10. Trends in ionization energies (kJ/mol) for the representative elements. p. 326

11. Beyond First Ionization • Each successive electron is harder to remove—large + charge in relation to # of electrons • Large increase in ionization energy occurs when going to a lower energy level (n changes)

12. Page 325

13. Electron Affinity The energy change in energy associated with the addition of an electron to a gaseousatom. X(g) + e X(g)

14. Electron Affinity • More exothermic for atoms that easily accept electrons; energy is more negative • Generally becomes more negative from left to right across a period—several exceptions such as adding to a 2p orbital with one electron already there (N vs C)

15. Electron affinity values for atoms among the first 20 elements that form stable, isolated X- ions.

16. Trend in Electron Affinity Becomes less exothermic down a group

17. Periodic Trends • Atomic Radii: • decrease going from left to right across a period; more positive charge in nucleus • increasegoing down a group; more energy levels

18. The radius of an atom (r) = half the distance between the nuclei in a molecule consistingof identical atoms.

19. Atomic radii (picometers) for selected atoms. p. 330

20. Trends in Transition Elements Sc Ni Ti V Cr Mn Fe Co • Although there is a slight contraction at the beginning of the series, the atoms are all approximately the same size.

21. Ions Formed • Varies by group • Very predictable for representative elements • +1 +2 +3 +/-4 –3 –2 –1 0 • Lose or gain electrons to reach a Noble gas configuration

22. Sizes of Ions • Size increases down a group; more energy levels • In isoelectric ions, size decreases as positivity increases; more + nucleus holds electrons closer. • Positive ions are much smaller than negative ions—nuclear attraction; less repulsion

23. All in the Family

24. Group 1A—Alkali Metals • 1 valence electron • Very reactive—very close to stable configuration • Lose electron to form +1 ion • Donate electrons—reducing agents • Reactivity increases down the group

25. Group 2A—Alkaline Earth Metals • 2 valence electrons • Very reactive, but less than alkali metals—not as close to stable config. • Lose 2 electron to form +2 ion • Donate electrons—reducing agents • Reactivity increases down the group

26. Group 3A • 3 valence electrons (ns2, np1) • Typically forms +3 ions • Increasingly metallic down the group • Forms bonds with more covalent character than groups 1 or 2

27. Transition Metals • Show more similarity than rep. ele. both within periods and groups • Electrons added in d & f orbitals are actually inner electrons & cannot bond as easily as the outer s & p electrons—chemical properties do not change as much.

28. Transition Metals—General Properties • Typical metals: shiny, good conductors, flexible (malleable & ductile) • Melting points, hardness, and reactivity vary, so transition metals have varied uses.

29. Uses forTransition Metals • Hg—low melting point + good conductivity—thermometers and thermostats • W—high melting point + ductility—filaments in lightbulbs

30. Uses forTransition Metals • Fe & Ti—strength + hardness—structural materials • Cu, Ag & Au—soft—jewelry and art • Cu—low resistance—wiring

31. Reactions with Transition Metals • Most react with O2 to form oxides • Cr, Ni, Co—oxides protect surface • Fe—oxide flakes off exposing more metal to oxidation • Au, Ag, Pt & Pd—Noble metals—do not readily form oxides

32. Ionic Compounds with Transition Metals • Often paramagnetic (unpaired e-s) • Usually more than one oxidation state • Cations are often complex ions in which ligands (Lewis base ions or molecules) surround metal ion

33. Complex ion

34. Ionic Compounds with Transition Metals Most are colored because ions absorb specific wave-lengths of light. Electrons in split d orbitals can rearrange.

35. Wulfenite—PbMoO4

36. Rhodochrosite—MnCO3

37. Electron Configuration • Full (n)s orbital; filling (n-1)d orbital; empty (n)p orbital • Exception: Cu—4s1, 3d10

38. Breaking the Rules • Since the energy of a 3d orbital is less than that of a 4s orbital, ions formed from first row transition metals lose their 4s electrons first • Ex: Mn = [Ar] 4s2, 3d5 Mn2+ = [Ar] 3d5

39. Lanthanides & Actinides Technically part of the transition metals Filling f orbitals with electrons Rare metals with few uses All actinides are radioactive—nuclear fuels (U & Pu)

40. Bonding 8 out of 75 M/C Questions Free Response—Every year

41. In General: • All bonds occur because of electrostatic attractions. • Formation of molecules and the state of matter of a substance depends on the attractions between electron clouds of one atom and nucleus of another atom.

42. Bonding—General Rules • A metal and a nonmetal bond so that charges on the ions cancel. • When two nonmetals react to form a covalent bond, they share electrons in a way that gives both atoms a Noble gas configuration.

43. Bonding—General Rules • nonmetal + representative metal binary ionic compound: ions form to give the nonmetal the valence electron configuration of the next noble gas atom, and valence orbitals of the metal are emptied

44. Bonding—General Rules • Most bonds are combinations of ionic/covalent character • The more different the atoms bonding are, the more ionic character of the bond. • More similar—more covalent

45. Covalent Sharing Molecules Structural formula Ionic Transfer / charged Compounds Formula Unit Word Association

46. Coulomb’s Law • Describes energy of interaction between ions • E = 2.31 x 10-19 J*nm(Q1Q2 / r) • Q1 & Q2numerical ion charges • r distance between ion centers • Negative ans. means ion pair is more stable than individual ions.

47. Bond Length • Distance between bonding atoms at which energy is minimized • Atoms position themselves to minimize repulsions and maximize attraction & thus achieve lowest possible energy.

48. Predicting Formulas for Ionic Compounds • Metal—positive charge equal to # of valence electrons • Nonmetal—negative charge equal to # of electrons away from next Noble gas • Compound—charges must cancel

49. Predict Formulas for: • Potassium sulfide • Barium chloride • Aluminum oxide • Magnesium phosphide

50. Exceptions to the Rule: • Sn—forms both +2 and +4 ions • Pb—forms both +2 and +4 ions • Bi—forms +3 and +5 ions • Tl—forms +1 and +3 ions  “no simple explanation for this behavior”