2. 1 PERIODIC TRENDS

3. 2 Electron Filling Order Figure 8.5

4. 3 Electron Configurations and the Periodic Table Figure 8.7

5. 4 PhosphorusPhosphorus All Group 5A elements have [core ] ns 2 np 3 configurations where n is the period number. Group 5A Atomic number = 15 1s 2 2s 2 2p 6 3s 2 3p 3 [Ne] 3s 2 3p 3

6. 5 LithiumLithium Group 1A Atomic number = 3 1s 2 2s 1 ---> 3 total electrons

7. 6 Electron Properties Diamagnetic : NOT attracted to a magnetic field Paramagnetic : substance is attracted to a magnetic field. Substance has unpaired electrons. Diamagnetic : NOT attracted to a magnetic field Paramagnetic : substance is attracted to a magnetic field. Substance has unpaired electrons.

8. 7 NeonNeon Group 8A Atomic number = 10 1s 2 2s 2 2p 6 ---> Diamagnetic

9. 8 BerylliumBeryllium Group 2A Atomic number = 4 1s 2 2s 2 Diamagnetic

10. 9 BoronBoron Group 3A Atomic number = 5 1s 2 2s 2 2p 1 Paramagnetic

11. 10 CarbonCarbon Group 4A Atomic number = 6 1s 2 2s 2 2p 2 Paramagnetic

12. 11 FluorineFluorine Group 7A Atomic number = 9 1s 2 2s 2 2p 5 ---> Paramagnetic

13. 12 Ion Configurations How do we know the configurations of ions? Determine the magnetic properties of ions. Ions with UNPAIRED ELECTRONS are PARAMAGNETIC. Without unpaired electrons DIAMAGNETIC.

14. 13 transition metal ions Fe [Ar] 4s 2 3d 6 loses 2 electrons ---> Fe 2+ [Ar] 4s 0 3d 6 loses 3 electrons ---> Fe 3+ [Ar] 4s 0 3d 5

15. 14 Transition Metals How do they fill? How can we determine? Copper Iron Chromium

16. 15 Ion Configurations Mn Mn [Ar] 4s 2 3d 5 ---> Mn 5+ [Ar] 4s 0 3d 2 loses 5 electrons ---> Mn 5+ [Ar] 4s 2 3d 0 D P

17. 16 PERIODIC TRENDS

18. 17 PERIODICITYPERIODICITY Period Law- - physical and chemical properties of elements are a periodic function of their atomic numbers

19. 18 General Periodic Trends Atomic and ionic size Ionization energy Electron affinity, electronegativity

20. 19 Effective Nuclear Charge Z*

21. 20 Effective Nuclear Charge, Z* Z* is the nuclear charge experienced by the outermost electrons. See p. 295 and Screen 8.6.Z* is the nuclear charge experienced by the outermost electrons. See p. 295 and Screen 8.6. Explains why E(2s) < E(2p)Explains why E(2s) < E(2p) Z* increases across a period owing to incomplete shielding by inner electrons.Z* increases across a period owing to incomplete shielding by inner electrons. Estimate Z* by --> [ Z - (no. inner electrons) ]Estimate Z* by --> [ Z - (no. inner 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 = 3and so on!B Z* = 5 - 2 = 3and so on!

22. 21 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

23. 22 General Periodic Trends Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly.

24. 23 Periodic Trend in the Reactivity of Metals Reactivity Periodic Trend in the Reactivity of Metals ReactivityLithium Sodium Potassium MOST

25. 2. Reactivity for MetalsReactivity As you go down a group for metals the number of energy levels increase. Because of this, reactivity increases because the atom is more willing to give away its electron (react).

26. 3.Nonmetalic Trends: Gain electrons Nonmetals on right side, form anions Going right elements are more nonmetallic (better gainers of electrons) Going UP elements become more nonmetallic (want to gain)

27. 8. Reactivity nonmetals: Gain e The reason Across = fill the energy level Going UP a group, nonmetals have same valence but fewer total electrons Flourine is the most reactive nonmetal.

28. 27 Atomic Radii Figure 8.9

29. 28 Atomic Size Size increases, down a group.Size increases, down a group. Because electrons are added into additional energy levels, there is less attraction.Because electrons are added into additional energy levels, there is less attraction. Size decreases across a period.Size decreases across a period. Because, increased effective nuclear charge.Because, increased effective nuclear charge. Size increases, down a group.Size increases, down a group. Because electrons are added into additional energy levels, there is less attraction.Because electrons are added into additional energy levels, there is less attraction. Size decreases across a period.Size decreases across a period. Because, increased effective nuclear charge.Because, increased effective nuclear charge.

30. 29 Atomic Size Size decreases across a period owing to increase in Z*. Each added electron feels a greater and greater + charge. Large Small

31. 30 Trends in Atomic Size See Figures 8.9 & 8.10

32. 31 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.

33. 32 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 -

34. 33 Trends in Ion Sizes Figure 8.13

35. 34 Ionization Energy IE = energy required to remove an electron from an atom in the gas phase. Mg (g) + 738 kJ ---> Mg + (g) + e-

36. 35 Mg (g) + 738 kJ ---> Mg + (g) + e- 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

37. 36 1 st IE: Mg (g) + 735 kJ ---> Mg + (g) + e- 2 nd IE: Mg + (g) + 1451 kJ ---> Mg 2+ (g) + e- 3 rd IE: Mg 2+ (g) + 7733 kJ ---> Mg 3+ (g) + e- Energy cost is very high to dip into a shell of lower n (core electrons). This is why ox. no. = Group no. Ionization Energy

38. 37 Trends in Ionization Energy

39. 38 Trends in Ionization Energy IE decreases down a group Because size increases. IE increases across a period Because effective nuclear charge increases IE increases across a period Because effective nuclear charge increases

40. 39 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. X(g) + e- ---> X - (g) E.A. = ∆E

41. 40 Trends in Electron Affinity

42. 41 Affinity for electron increases across a period (EA becomes more negative). 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

43. 42 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

44. 43 Electron Affinity of Nitrogen ∆E is zero for N - due to electron- electron repulsions. EA = 0 kJ [He]     Natom  [He]    N - ion  + electron

45. 44 Electronegativity So how is this different from electron affinity? Electron Affinity – is rating of how well an atom wants to gain an electron Electronegativity – is rating of how well an atom keeps the electron once it is bonded to another atom

46. 45 Electronegativity

47. 46 Electron Configurations and the Periodic Trends

48. 47 “Your Best Friend” Periodic table