Chapter 7 Periodic Properties of the Elements

1. Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 7Periodic Propertiesof the Elements John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc.

2. November 9 • For Monday finish chapter 7 and do all the questions in the worksheet. • Test on chapter 6 and 7 on Tuesday. • Equation book review Chapter 2 and 3

3. Anomalies • Cr 3d5 4s1 • Mo • Cu • Ag • Au • The energies between 3d and 4s are close, and to have a sublevel half filled or completely filled gives stability to the atom.

4. Paramagnetism • Elements and compounds that have unpaired electrons are attracted to a magnet. The effect is weak but it can be observed. • Liquid O2 attracted to a magnet.

5. Diagmagnetism • Substances with no unpaired electrons experiment a slight repulsion when subjected to a magnetic field.

6. http://books.google.com/books?id=jcn6sgt7RpoC&lpg=PA405&ots=0AV83UNudy&dq=paramagnetism%20images%20kotz&pg=PA293#v=onepage&q&f=falsehttp://books.google.com/books?id=jcn6sgt7RpoC&lpg=PA405&ots=0AV83UNudy&dq=paramagnetism%20images%20kotz&pg=PA293#v=onepage&q&f=false

7. Development of Periodic Table • Elements in the same group generally have similar chemical properties. • Properties are not identical, however.

8. Development of Periodic Table Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped in 1869.

9. PERIODIC ARRANGEMENT • Mendeleev arranged the elements according to their atomic masses. • In 1913 Henry Moseley developed the concept of ATOMIC NUMBERS which lead to the current arrangement according to their ATOMIC NUMBER.

10. Effective Nuclear Charge • In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. • The nuclear charge that an electron experiences depends on both factors.

11. Effective Nuclear Charge The effective nuclear charge, Zeff, is found this way: Zeff = Z−S where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons.

12. Effective Nuclear Charge • The effective nuclear charge is smaller than the actual nuclear charge because the effective nuclear charge Zeff accounts for the repulsion of the electron by the other electrons in the atom. • Zeff < Z

13. Core electrons shield or screen the outer electrons from the attraction of the nucleus.

14. Sizes of Atoms The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

15. Sizes of Atoms Bonding atomic radius tends to… …decrease from left to right across a row due to increasing Zeff. …increase from top to bottom of a column due to increasing value of n

16. Examples – Place each group of elements in order of increasing atomic radius: • S, Al, Cl, Mg, Ar, Na • K, Li, Cs, Na, H • Ca, As, F, Rb, O, K, S, Ga

17. Examples – Place each group of elements in order of increasing atomic radius: • S, Al, Cl, Mg, Ar, Na • Ar < Cl < S < Al < Mg < Na • K, Li, Cs, Na, H • H < Li < Na < K < Cs • Ca, F, As, Rb, O, K, S, Ga • F < O < S < As < Ga < Ca < K < Rb

18. Electron Configurations of Ions • Cations: electrons removed from orbital with highest principle quantum number, n, first: • Li (1s2 2s1)  Li+ (1s2) • Fe ([Ar]3d6 4s2)  Fe3+ ([Ar]3d5) • Anions: electrons added to the orbital with highest n: • F (1s2 2s2 2p5)  F-(1s2 2s2 2p6)

19. Write electron configurations for the following ions: • Al3+ • S2- • Li+ • Br- • Fe2+ • Fe3+

20. Write electron configurations for the following ions: • Al3+ 1s22s22p6 • S2-[Ne]3s23p6 • Li+ 1s2 • Br- [Ar]4s23d104p6 • Fe2+[Ar]3d6 • Fe3+[Ar]3d5

21. Sizes of Ions • Ionic size depends upon: • Nuclear charge. • Number of electrons. • Orbitals in which electrons reside.

22. Sizes of Ions • Cations are smaller than their parent atoms. • The outermost electron is removed and repulsions are reduced.

23. Sizes of Ions • Anions are larger than their parent atoms. • Electrons are added and repulsions are increased.

24. Sizes of Ions • Ions increase in size as you go down a column. • Due to increasing value of n.

25. November 10 • DO NOW • Write the formula for • Potassium Permanganate and Sodium hypochlorite • Write the electronic configuration of both cations. Indicate to what element are each isoelectronic. • Which one has the largest radius?

26. Homework • New Atomic Theory Exams Questions is a file in the John Bowne site. Has similar questions to AP test. Practice them. The more you do them the better you’ll do in the test. Test Unit 6 and 7 Tuesday

27. Sizes of Ions • In an isoelectronic series, ions have the same number of electrons. • Ionic size decreases with an increasing nuclear charge.

28. Examples – Choose the larger species in each case: • Na or Na+ • Br or Br- • N or N3- • O- or O2- • Mg2+ or Sr2+ • Mg2+ or O2- • Fe2+ or Fe3+

29. Examples – Choose the larger species in each case: • Na or Na+ • Br or Br- • N or N3- • O- or O2- • Mg2+ or Sr2+ • Mg2+ or O2- • Fe2+ or Fe3+

30. Ionization Energy • Amount of energy required to remove an electron from the ground state of a gaseous atom or ion. • First ionization energy is that energy required to remove first electron. • Second ionization energy is that energy required to remove second electron, etc.

31. IONIZATION ENERGY • THE GREATER THE IONIZATION ENERGY THE MORE DIFFICULT IT IS TO REMOVE AN ELECTRON! • It requires more energy to remove each successive electron. • When all valence electrons have been removed, the ionization energy takes a quantum leap (removing core electrons require huge amounts of energy).

33. Trends in First Ionization Energies • As one goes down a column, less energy is required to remove the first electron. • For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.

34. Trends in First Ionization Energies • Generally, as one goes across a row, it gets harder to remove an electron. • As you go from left to right, Zeff increases.

35. N 1402 [He] 2s2 2p3 2p3 this configuration is more stable because half filled orbitals minimize repulsions between electrons O 1314 [He] 2s2 2p4 The decrease is due to repulsions of paired electrons in 2p4 Irregularities

36. Trend across a periodIE icreases • The energy needed to remove an electron increases because the effective nuclear charge increases and the atomic radius decreases. • Irregularities • Be 899 B 801 • [He] 2s2 [He] 2s2 2p1

37. Trends in First Ionization Energies However, there are two apparent discontinuities in this trend.

38. Trends in First Ionization Energies • The first occurs between Groups IIA and IIIA. • Electron removed from p-orbital rather than s-orbital • Electron farther from nucleus • Small amount of repulsion by s electrons.

39. Trends in First Ionization Energies • The second occurs between Groups VA and VIA. • Electron removed comes from doubly occupied orbital. • Repulsion from other electron in orbital helps in its removal.

40. Examples – Put each set in order of increasing first ionization energy: • P, Cl, Al, Na, S, Mg • Ca, Be, Ba, Mg, Sr • Ca, F, As, Rb, O, K, S, Ga

41. Examples – Put each set in order of increasing first ionization energy: • P, Cl, Al, Na, S, Mg • Ca, Be, Ba, Mg, Sr • Ca, F, As, Rb, O, K, S, Ga • 1. Na < Al < Mg < S < P < Cl • 2. Ba < Sr < Ca < Mg < Be • 3. Rb < K < Ga < Ca < As < S < O < F

42. Electron Affinity • It measures the ease with which and atom gains an electron

43. Electron Affinity Energy change accompanying addition of electron to gaseous atom: Cl + e− Cl− For most atoms energy is released when an electron is added. The change is exothermic. The more exothermic the most stable the product.

44. Trends in Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row.

45. Trends in Electron Affinity There are again, however, two discontinuities in this trend.

46. Trends in Electron Affinity • The first occurs between Groups IA and IIA. • Added electron must go in p-orbital, not s-orbital. • Electron is farther from nucleus and feels repulsion from s-electrons.

47. Trends in Electron Affinity • The second occurs between Groups IVA and VA. • Group VA has no empty orbitals. • Extra electron must go into occupied orbital, creating repulsion.

48. November 11 • Do now • Find the Oxidation Number for Oxygen in • KO2 • K2O2 • KO2 • Name the compounds

49. Electronegativity: The ability of an atom in a molecule to attract electrons to itself. • Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F). Values are calculated from ionization energies and electron affinities. • Electronegativity increases • across a period and • Up a group.

50. Electronegativity