2. PERIODICITYPERIODICITY

3. Periodic Table Dmitri Mendeleev developed the modern periodic table. Argued that element properties are periodic functions of their atomic weights.Dmitri Mendeleev developed the modern periodic table. Argued that element properties are periodic functions of their atomic weights. We now know that element properties are periodic functions of their ATOMIC NUMBERS.We now know that element properties are periodic functions of their ATOMIC NUMBERS.

4. Periods in the Periodic Table

5. Groups in the Periodic Table

6. Regions of the Periodic Table

7. Metals, Nonmetals, and Metalloids –Three classes of elements are metals, nonmetals, and metalloids. –Across a period, the properties of elements become less metallic and more nonmetallic. 6.1

8. Metals, Nonmetals, and Metalloids »Metals, Metalloids, and Nonmetals in the Periodic Table 6.1

9. Metals, Nonmetals, and Metalloids »Metals, Metalloids, and Nonmetals in the Periodic Table 6.1

10. Metals, Nonmetals, and Metalloids »Metals, Metalloids, and Nonmetals in the Periodic Table 6.1

11. Metals, Nonmetals, and Metalloids »Metals, Metalloids, and Nonmetals in the Periodic Table 6.1

12. Metals, Nonmetals, and Metalloids –Metals Metals are good conductors of heat and electric current. –80% of elements are metals. –Metals have a high luster, are ductile, and are malleable. 6.1

13. Metals, Nonmetals, and Metalloids »Uses of Iron, Copper, and Aluminum 6.1

14. Metals, Nonmetals, and Metalloids »Uses of Iron, Copper, and Aluminum 6.1

15. Metals, Nonmetals, and Metalloids »Uses of Iron, Copper, and Aluminum 6.1

16. Metals, Nonmetals, and Metalloids –Nonmetals In general, nonmetals are poor conductors of heat and electric current. –Most nonmetals are gases at room temperature. –A few nonmetals are solids, such as sulfur and phosphorus. –One nonmetal, bromine, is a dark-red liquid. 6.1

17. Metals, Nonmetals, and Metalloids –Metalloids A metalloid generally has properties that are similar to those of metals and nonmetals. The behavior of a metalloid can be controlled by changing conditions. 6.1

18. http://www.webelements.com/webelements/elements/text/Si/geol.html Element Abundance Fe C Al O Si

19. HydrogenHydrogen Shuttle main engines use H 2 and O 2 The Hindenburg crash, May 1939.

20. Group 1A: Alkali Metals Cutting sodium metal Reaction of potassium + H 2 O

21. Magnesium Magnesium oxide Group 2A: Alkaline Earth Metals

22. Calcium Carbonate— Limestone The Appian Way, Italy Champagne cave carved into chalk in France

23. Group 3A: B, Al, Ga, In, Tl Aluminum Boron halides BF 3 & BI 3

24. Gems & Minerals Sapphire: Al 2 O 3 with Fe 3+ or Ti 3+ impurity gives blue whereas V 3+ gives violet.Sapphire: Al 2 O 3 with Fe 3+ or Ti 3+ impurity gives blue whereas V 3+ gives violet. Ruby: Al 2 O 3 with Cr 3+ impurityRuby: Al 2 O 3 with Cr 3+ impurity

25. Group 4A: C, Si, Ge, Sn, Pb Quartz, SiO 2 Diamond

26. Group 5A: N, P, As, Sb, Bi White and red phosphorus Ammonia, NH 3

27. Phosphorus Phosphorus first isolated by Brandt from urine, 1669

28. Group 6A: O, S, Se, Te, Po Sulfuric acid dripping from snot-tite in cave in Mexico Sulfur from a volcano

29. Group 7A: F, Cl, Br, I, At

30. Group 8A: He, Ne, Ar, Kr, Xe, Rn Lighter than air balloons “Neon” signs XeOF 4

31. Transition Elements Lanthanides and actinides Iron in air gives iron(III) oxide

33. LithiumLithium Group 1A Atomic number = 3 1s 2 2s 1 ---> 3 total electrons

34. BerylliumBeryllium Group 2A Atomic number = 4 1s 2 2s 2 ---> 4 total electrons

35. BoronBoron Group 3A Atomic number = 5 1s 2 2s 2 2p 1 ---> 5 total electrons 5 total electrons

36. CarbonCarbon Group 4A Atomic number = 6 1s 2 2s 2 2p 2 ---> 6 total electrons 6 total electrons

37. NitrogenNitrogen Group 5A Atomic number = 7 1s 2 2s 2 2p 3 ---> 7 total electrons 7 total electrons

38. OxygenOxygen Group 6A Atomic number = 8 1s 2 2s 2 2p 4 ---> 8 total electrons 8 total electrons

39. FluorineFluorine Group 7A Atomic number = 9 1s 2 2s 2 2p 5 ---> 9 total electrons 9 total electrons

40. NeonNeon Group 8A Atomic number = 10 1s 2 2s 2 2p 6 ---> 10 total electrons 10 total electrons

41. Colors of Transition Metal Compounds Iron Cobalt Nickel CopperZinc

42. PERIODIC TRENDS

43. General Periodic Trends Atomic and ionic sizeAtomic and ionic size Ionization energyIonization energy Electron affinityElectron affinity ElectronegativityElectronegativity Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly.

44. Effective Nuclear Charge, Z* Explains why E(2s) < E(2p)Explains why E(2s) < E(2p) Z* is the nuclear charge experienced by the outermost electrons. Is the result of the nuclear attraction being blocked by the core electrons. Nuclear attraction increases with an increase in protonsZ* is the nuclear charge experienced by the outermost electrons. Is the result of the nuclear attraction being blocked by the core electrons. Nuclear attraction increases with an increase in protons Estimate Z* by --> [ Z - (no. core electrons) ]Estimate Z* by --> [ Z - (no. core electrons) ] Charge felt by 2s e- in Li Z* = 3 - 2 = 1Charge felt by 2s e- in Li Z* = 3 - 2 = 1 Be Z* = 4 - 2 = 2Be Z* = 4 - 2 = 2 B Z* = 5 - 2 = 3 and so on!B Z* = 5 - 2 = 3 and so on!

45. Effective Nuclear Charge, Z* Shielding effect remains constant across a period. As the nuclear attraction increases across the shielding effect is less effective.Shielding effect remains constant across a period. As the nuclear attraction increases across the shielding effect is less effective. Shielding effect increases down a group thus effectively blocking any increase in nuclear attraction.Shielding effect increases down a group thus effectively blocking any increase in nuclear attraction. Electrons with a higher quantum number have more kinetic energy and thus are less affected by the nuclear charge.Electrons with a higher quantum number have more kinetic energy and thus are less affected by the nuclear charge.  Each of these forces need to be accounted for in each trend.

46. Effective Nuclear Charge, Z* AtomZ* Experienced by Electrons in Valence Orbitals Li+1.28 Be------- B+2.58 C+3.22 N+3.85 O+4.49 F+5.13 Increase in Z* across a period

47. Periodic Trend in the Reactivity of Alkali Metals with Water Lithium SodiumPotassium

48. Atomic Size Size goes UP on going down a group.Size goes UP on going down a group. Because electrons are added further from the nucleus, there is less attraction, due to an increase in sheilding effectiveness and in increase in kinetic energy.Because electrons are added further from the nucleus, there is less attraction, due to an increase in sheilding effectiveness and in increase in kinetic energy. Size goes UP on going down a group.Size goes UP on going down a group. Because electrons are added further from the nucleus, there is less attraction, due to an increase in sheilding effectiveness and in increase in kinetic energy.Because electrons are added further from the nucleus, there is less attraction, due to an increase in sheilding effectiveness and in increase in kinetic energy.

49. Atomic Size Size goes UP on going down a group.Size goes UP on going down a group. Because electrons are added further from the nucleus, there is less attraction, due to an increase in sheilding effectiveness and in increase in kinetic energy.Because electrons are added further from the nucleus, there is less attraction, due to an increase in sheilding effectiveness and in increase in kinetic energy. Size goes UP on going down a group.Size goes UP on going down a group. Because electrons are added further from the nucleus, there is less attraction, due to an increase in sheilding effectiveness and in increase in kinetic energy.Because electrons are added further from the nucleus, there is less attraction, due to an increase in sheilding effectiveness and in increase in kinetic energy.

50. General Outline for Trends Trend-define Down a group Nuclear attraction-define once –Trend, effect –Shielding effect-define once Trend, effect –Kinetic energy-define once Trend, effect Across a period Nuclear attraction –Trend, effect –Shielding effect Trend, effect –Kinetic energy Trend, effect

51. Atomic Radius Atomic radius is the distance from the nucleus to the valance electrons. –Nuclear attraction (the attraction of the protons in the nucleus on valance electrons) increases going down a group. This should pull the electrons in closer to the nucleus. –Shielding effect (the blocking of nuclear attractions by core electrons) Shielding effect increases down a group offsetting the increase in nuclear attraction. – Kinetic energy (the energy of valance electrons associated with principle energy levels) increases down a group allowing the valance electrons to orbit farther from the nucleus increasing atomic radius.

52. Atomic Radius –Nuclear attraction increases across a period. This should pull the electrons in closer to the nucleus decreasing atomic radius. –Shielding effect remains constant across a period not offsetting nuclear attraction. –Kinetic energy remains constant across a period so effective nuclear attraction is greater and the atomic radius decreases.

53. Atomic Radii Figure 8.9

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

55. 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?

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

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

58. Ion Sizes ANIONS are LARGER than the atoms from which they come.ANIONS are LARGER than the atoms from which they come. The electron/proton attraction has gone DOWN and so size INCREASES.The electron/proton attraction has gone DOWN and so size INCREASES. Trends in ion sizes are the same as atom sizes.Trends in ion sizes are the same as atom sizes. Forming an anion. F, 71 pm 9e and 9p F -, 133 pm 10 e and 9 p -

59. Trends in Ion Sizes Figure 8.13

60. Ionic Size Ionic size is the distance from the nucleus to the valence electrons after an atom has lost or gained electrons. Cations form when an atom loses one or more electrons. Cations are smaller than the atoms from which they form Ionic size - Cations –Effective nuclear charge increases dramatically when electrons are removed. –Shielding effect decreases compared to the atom because valence electrons are lost and some or all of the core electrons become valence electrons. –Kinetic energy the new valence electrons have less kinetic energy to resist the pull of the nucleus.

61. Ionic Size Anions form when an atom gains one or more electrons Ionic size - Anions –After the addition of valence electron(s) the nuclear attraction is diluted. –Shielding effect remains still resisting the nuclear attraction. –Kinetic energy increases because of additional repulsion due to more electrons in the valence shell increasing the anions size. –Anions are larger than the atoms from which they form

62. Ionization Energy IE = energy required to remove an electron from an atom in the gas phase. Mg (g) + 738 kJ ---> Mg + (g) + e-

63. Mg + (g) + 1451 kJ ---> Mg 2+ (g) + e- Mg + has 12 protons and only 11 electrons. Therefore, IE for Mg + > Mg. IE = energy required to remove an electron from an atom in the gas phase. Ionization Energy

64. 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. Ionization Energy

65. General Periodic Trends Atomic and ionic sizeAtomic and ionic size Ionization energyIonization energy Electron affinityElectron affinity Higher Z*. Electrons held more tightly. Larger orbitals. Electrons held less tightly.

66. Trends in Ionization Energy

67. IE increases across a period because Z* increases.IE increases across a period because Z* increases. Metals lose electrons more easily than nonmetals.Metals lose electrons more easily than nonmetals. Metals are good reducing agents.Metals are good reducing agents. Nonmetals lose electrons with difficulty.Nonmetals lose electrons with difficulty.

68. Trends in Ionization Energy IE decreases down a groupIE decreases down a group Because size increases.Because size increases. Reducing ability generally increases down the periodic table.Reducing ability generally increases down the periodic table. See reactions of Li, Na, KSee reactions of Li, Na, K

69. 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. Ionization Energy See Screen 8.12

70. Ionization Energy Ionization energy is the energy needed to remove an electron from an atom. Nuclear attraction –Increases down a group holding the electrons tighter. –Shielding effect increases down a group offsetting the increase in nuclear attraction. –Kinetic energy increases down a group giving the electrons greater initial energy. This reduces the additional energy needed to remove an electron. Nuclear attraction increases across a period holding the electrons tighter. Shielding effect is constant across and does not offset the increase in nuclear attraction. Kinetic energy remains constant across. With the same initial energy valence electrons are increasingly harder to remove due to the greater effective nuclear charge.

71. Electron Affinity A few elements GAIN electrons to form anions. Electron affinity is the energy involved when an atom gains an electron to form an anion. A(g) + e- ---> A - (g) E.A. = ∆E A(g) + e- ---> A - (g) E.A. = ∆E

72. 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

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

74. See Figure 8.12 and Appendix FSee Figure 8.12 and Appendix F Affinity for electron increases across a period (EA becomes more positive).Affinity for electron increases across a period (EA becomes more positive). Affinity decreases down a group (EA becomes less positive).Affinity decreases down a group (EA becomes less positive). 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

76. Electronegativity Values See page 177 in text

77. Electronegativity Electronegativity is the tendency of an atom to remove an electron from another atom when forming a compound. Nuclear attraction –Increases down a group attracting the electrons more. –Shielding effect increases down a group offsetting the increase in nuclear attraction. –Kinetic energy increases down a group giving the atom a larger radius and increasing the proximity of the nucleus to adjacent electrons decreasing electronegativity Nuclear attraction increases across a period attracting electrons more. Shielding effect and kinetic energy are constant across. This increases effective nuclear charge allowing the atom to remove electron from other atoms with lesser electronegativity.