1. Chapter 8: Periodic Properties Nerve cells must pump Na + ions out of the cell and K + ions into the cell to complete the transmission of a signal. Both are group 1A elements – yet the cells can distinguish them based on their size! Size of atoms or ions is an example of a periodic property.

2. History Elements like gold (Au) can be found in their pure state in nature and were known to the ancient humans. Others are more rare and occur only in compounds, which until the 19 th century were difficult to isolate. In 1800, only 31 elements had been identified. By 1865, the number was 63. www.ptable.com

3. History In 1869, Dmitri Mendeleev proposed a classification scheme based on the fact that physical and chemical properties begin to repeat when the elements are arranged in increasing weight. As many elements were still not yet discovered, Mendeleev left blank spaces where he predicted that these elements would be eventually isolated.

4. History

5. In 1913, Henry Moseley found that when an element is bombarded with high energy electrons it will produce an X-ray with a unique frequency. He then arranged them in order of frequency and assigned them a whole number, which we now know as their atomic number.

6. Orbitals Subshells are described by the combination of the n and l quantum numbers. Energy increases for each subshell according to Aufbau filling order. 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.

7. Electron Configurations Ground state = lowest possible energy state for all of the electrons. Each subshell in the Aufbau order fills completely before starting the next subshell. Remember, s = 1 orbital, p = 3 orbitals, d = 5 orbitals, and d = 7 orbitals and EACH orbital can hold 2 electrons per quantum number rules.

8. Electron Configurations Use a superscript after each subshell to indicate the number of electrons in that subshell. H: 1s 1 He: 1s 2 Li: 1s 2 2s 1 Be 1s 2 2s 2 B: 1s 2 2s 2 2p 1

9. Electron Configurations Imagine doing a configuration for Sr or W or Pb!!! Shorthand method LEP #8. Exceptions.

10. Subshells and the Periodic Table

11. Core and Valence Electrons Core electrons are very close to the nucleus and are the completed shell(s) of electrons. Valence electrons are the outermost shell of electrons. These are loosely held and provide for all of the bonding interactions. For the main group elements, the number of valence electrons = the group number.

12. Orbital Diagrams Hund’s Rule – when electrons are placed into degenerate orbitals, they will fill them singularly with parallel spins. For outermost shells, use boxes, circles, or lines to represent each orbital. Use a  for spin = +1/2 and use a  for spin = -1/2. LEP #1

13. Periodic Properties Elements in the same group tend to behave similarly. Ex) Group 1A metals – Li, Na, K, Rb, Cs Some, though, have major differences. Ex) O and S Electronic properties of valence electrons are always the same. The chemical properties of the elements are largely determined by the number of valence electrons they contain.

14. Effective Nuclear Charge The attraction of the electron to the nucleus depends on two factors. Factor one is the magnitude of the charge. Factor two is the distance between the charge. In a many electron atom, each electron is attracted to the nucleus and repelled by the other electrons.

15. Effective Nuclear Charge The core electrons versus the valence electrons. Argon, 10 core and 8 valence electrons.

16. Effective Nuclear Charge Core electrons “shield” the valence electrons from the nucleus. Z eff = Number of Protons – Number of Core Electrons Z eff for Na = +1, for Cl = +7 Trend is: ________________________

17. Size of Atoms Atoms are sometimes treated as hard spheres in models. Defining atomic size is tricky, though, because atoms do not have defined boundary layers. Non-bonding radii. Bonding radii.

18. Size of Atoms What trends do you see? Any exceptions? H 37 Atomic Radii pm (picometers) He 32 Li 152 Be 112 B 85 C 77 N 70 O 73 F 72 Ne 70 Na 186 Mg 160 Al 143 Si 118 P 110 S 103 Cl 99 Ar 98 K 227 Ca 197 Ga 135 Ge 122 As 120 Se 117 Br 114 Kr 112 Rb 248 Sr 215 In 166 Sn 140 Sb 141 Te 143 I 133 Xe 130

19. Electron Configurations of Ions When a main group element gains or loses electrons, it does so with predictability. Na (and other group 1A metals) always lose one electron. Na: [Ne] 3s 1  Na + : [Ne] + 1e - Cl (and other group 7A halogens) always gain one electron. Cl: [Ne] 3s 2 3p 5 + 1e -  Cl - : [Ne] 3s 2 3p 6

20. Electron Configurations of Ions For transition metal ions, their configurations are NOT what is expected. When filling, the 4s fills BEFORE the 3d orbital. When removing electrons, though, they do not remove in reverse order! ALWAYS remove the s electrons first. Ex) Fe +2, Co +3 Diamagnetic = no unpaired electrons, paramagnetic = unpaired electrons.

21. Size of Ions What trends do you see?

22. Isoelectronic Series Ions that contain the same number of electrons are said to be isoelectronic. Place the following in order of increasing size: Na +, F -, Mg +2, O -2, and Al +3.

23. Ionization Energy (IE) IE is the minimum amount of energy required to remove an electron from a gaseous atom or ion. IE 1 is the energy to remove the first electron from a neutral atom. Ex) Na (g)  Na + (g) + 1e - IE 2 is the energy to remove the next electron.

24. Ionization Energy What trends do you see? Any exceptions? H 1312 First Ionization Energies, kJ/moleHe 2372 Li 520 Be 899 B 801 C 1086 N 1402 O 1314 F 1681 Ne 2081 Na 496 Mg 738 Al 578 Si 786 P 1012 S 1000 Cl 1251 Ar 1521 K 419 Ca 590 Ga 579 Ge 762 As 947 Se 941 Br 1140 Kr 1351 Rb 403 Sr 549 In 558 Sn 709 Sb 834 Te 869 I 1008 Xe 1170 Cs 376 Ba 503 Tl 589 Pb 716 Bi 703 Po 812 Rn 1037

25. Successive IE’s What do you see? Why is it easier to remove the second electron from Cl than the second electron from Na? elementIE1IE2IE3IE4IE5IE6IE7 Na4964,562 Mg7381,4517,733 Al5781,8172,74511,577 Si7861,5773,2324,35616,091 P1,0121,9072,9144,9646,27421,267 S1,0002,2523,3574,5567,0048,49627,107 Cl1,2512,2983,8225,1596,5429,36211,018

26. Electron Affinity (EA) EA is the amount of energy released (exothermic) when a gaseous atom gains an electron. Ex) Cl (g) + e -  Cl -1 (g) ;  H = -349 kJ An important note about sign and size of  H. Some atoms will not accept an electron without having to absorb energy – these are said to have no measurable EA.

27. Electron Affinity Ordered by groups. Which group is the best? Why? Can we make predictions for transition metals?

28. Metals Recall from Chapter 2 that metals: Are usually solids at room temperature. Are shiny and have luster Are excellent conductors of heat and electricity. Tend to lose electrons easily (form cations). Metals oxides are said to be basic. Metal oxide + water Metal oxide + acid

29. Metals Some metal ions can be identified by a flame test. Seen below from left: Li, Na, and K.

30. Fireworks

31. Non-metals Recall from Chapter 2 that non-metals: Have a variety of appearances – some are gases, some are solids, and one is a liquid. Some are single atoms like metals, while others are diatomic or larger. That are solids are brittle and poor conductors – except carbon! Will either gain electrons (ionic) or share electrons (molecular). Non-metal oxides are said to be acidic. Non-metal oxide + water Non-metal oxide + base

32. Metallic Character

33. Groups 1A and 7A are highly reactive. Group 1A loses 1 electron quite easily. Thus, reactivity driven by I.E. Reactivity increases from top to bottom within this group. Why? Group 7A gains 1 electron quite easily. Thus, reactivity driven by E.A. Reactivity greatest for F 2 and Cl 2 because they are gases. Reactivity and Periodic Table

34. The reaction of the alkali metals with water are all the same. 2 Na (s) + 2 H 2 O (l)  2 NaOH (aq) + H 2(g) + heat. Reaction is highly exothermic. The reaction of Al with the liquid Br 2. 2 Al (s) + 3 Br 2(l)  2 AlBr 3(s) + heat. Reaction is also highly exothermic. Alkali metals and Halogens

35. Alkali metals

36. Halogens

37. Groups 2A and 6A are reactive, but less than the groups 1A and 7A. Group 2A metals will still react with water, but much more slowly. Group 6A non-metals (O 2 and S) are similar. Groups 2A and 6A

38. Calcium metal

39. Na  Na + + 1e - ;  H = +496 kJ Ca  Ca +2 + 2e - ;  H = +1,735 kJ Similarly, if you compared energy to gain one electron for group 7A versus gaining two electrons for group 6A, then the latter is much harder. Why the differences?

40. Fairly unreactive with a few exceptions like Phosphorus. Phosphorus comes in several forms including both white and red phosphorus. White is the most unstable. Groups 3A, 4A, and 5A

41. Red Phosphorus and Bromine

42. Group 8A forms no compounds at all with the exception of Xenon with Fluorine and Oxygen. XeF 2, XeF 4, XeF 6, XeO 3, and XeO 4 have been produced. This is due to the relatively low I.E. of Xe. Group 8A

43. Reactivity on the left-hand side of the periodic chart (metals!) is driven by the ease of losing electrons. Reactivity on the right-hand side of the periodic chart (non-metals) is driven by the ease of gaining electrons. Summary