2. THE PERIODIC TABLE Objectives: Examine the progression of periodicity

3. Alkali metals alkaline earth metals “s” group Nonmetals “p” block Metalloids (semimetals) Transition metals “d” block Inner transition metals “f” group Noble gases Halogens

4. Periodic Patterns l ALKALI METALS (part of the “s” group of elements) ~ all are in shiny solid form but are quite soft ~ form the 1st group of metals on the periodic table ~ highly reactive elements based upon their electron configurations (ns 1 ) Other Characteristics:  malleable and ductile;low density and melting points  good conductors of electricity; very soluble as comps.

5. Alkaline Earth Metals  Belong to the second group of metals on the periodic table.  Harder, more dense, and stronger than there group 1 counterparts  Not as reactive as the Alkali metals due to the metals in this group having 2 electrons in their valence shell.  This gives them the configuration of “ns 2 ” for these metals. PART OF THE “S” GROUP JUST LIKE ALKALI METALS Be Sr Ba Mg Ra Ca

6. Transition Metals Z Transition metals begin in the 4th period after the alkaline earth metals. Z Metallic elements with varying properties. Z Not nearly as reactive as group 1 and 2 elements. Z Fill their sublevels differently than do the Main group elements. The “d” - block elements Valuable as structurally useful materials! Important in living organisms!

7. Lanthanoids & Actinoids LANTHANOIDS C Composed of the elements with atomic numbers 58 through 71 C Electrons are being added to the “ 4f “ sublevel C Shiny reactive metals with practical uses ie. dots in TV tubes ACTINOIDS l Composed of the elements with the atomic numbers 90 through 103 l They fill the “ 5f “ sublevel l All are radioactive with an unstable nucleus “Y” The ‘f’-group is broken into two classifications

8. Nonmetals & Metalloids (semi-metals ) NONMETALS! 4 Generally are gases at room temperature (or brittle solids) 4 Poor conductors of heat and electricity 4 Have more electrons in their outer level than metals METALLOIDS! 4 Properties of both metals and nonmetals 4 Will give up (electron donor) electron(s) when reacted with a nonmetal, and will accept (electron acceptor) electron(s) when reacted with a metal 4 In general, more like nonmetals than metals 4 Considered semiconductors

9. Periodic Trends l Trends (we will study) – atomic radius (ionic radius), ionization energy, electronegativity, electron affinity l Trends are looked at from top to bottom of a column and from left to right in a period (row) l Trends show patterns of atoms properties (relationships among elements)

10. ATOMIC RADIUS Li Na K Rb Cs Fr  Atomic radius is the half the distance between the nuclei of two like atoms. Trend Number 1  the trend for atomic radius shows us the size of the atom will increase as we move down a column WHY: more levels and orbitals, greater distances from the nucleus TREND:

11. l the atomic radii will decrease from the left to the right in a period WHY: Effective Nuclear Charge (also applies to what takes place from top to bottom of a column) l positive charge felt by the outermost electrons of an atom l atomic # - # of inner complete level electrons l The larger the ENC, the greater the attraction of electrons to the nucleus Shielding - the ability of other electrons,especially inner electrons, to lessen the nuclear charge of the outer electron(s) ATOMIC RADIUS

12. THE TREND!  The trend shows the increase of radii down a group and decrease of radii across a period.

13. IONIZATION ENERGY IONIC BOND 4 bond formed between two ions by the transfer of electrons Ions: How do they form? 4 In certain types of bonding, the atom will “lose” or “gain” an electron(s) N When an atom loses or gains electrons, it is called an ion Magnesium NUMBER TWO!

14. N Atoms that lose electrons have a positive charge N Atoms that gain electrons have a negative charge Magnesium BOINK!

15. 4 For the most part, the metals will lose electrons and the nonmetals will accept the electrons 4 The atoms gain or lose electrons to reach outer shell (valence) stability CHLORINE Electron from magnesium

17. Ionic Bonds: One Big Greedy Thief Dog!

18. Ion Sizes Does the size go up or down when losing an electron to form a cation?

19. Ion Sizes l CATIONS are SMALLER than the atoms from which they come. l 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.

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

21. Ion Sizes l ANIONS are LARGER than the atoms from which they come. l The electron/proton attraction has gone DOWN and so size INCREASES. Forming an anion. F, 71 pm 9e and 9p F -, 133 pm 10 e and 9 p -

22. Trends in Ion Sizes Figure 8.13

23. IONIZATION ENERGY v the energy required to remove the most loosely held electron from an atom v ionization energy decreases as the size of the atom increases (top to bottom of a column) “Y”? Because the outer most electron is farther from the nucleus and the electrical attraction to the protons.

24. More Details!  Energy is absorbed by the atom to free the electron(s) endothermic  Ionization is endothermic, meaning that the atom or molecule increases its internal energy ( takes energy from an outside source) A + energy A + + e -

25. Ionization Energy is affected by three factors: 1. Effective Nuclear Charge 2. Number of Energy Levels 3. Shielding

26. Ionization Energies l The first ionization energy, I 1, is the energy needed to remove the first electron from the atom: Mg  Mg + + 1e -

27. l The second ionization energy, I 2, is the energy needed to remove the next (i.e. the second) electron from the atom Mg +  Mg 2+ + 1e - The higher the value of the ionization energy, the more difficult it is to remove the electronThe higher the value of the ionization energy, the more difficult it is to remove the electron

28. 1 st IE2 nd IE3 rd IE4 th IE5 th IE6 th IE7 th IE Na4964,560 Mg7381,4507,730 Al5771,8162,88111,600 Si7861,5773,2284,35416,100 P1,0601,8902,9054,9506,27021,200 S999.62,2603,3754,5656,9508,49027,107 Cl1,2562,2953,8505,1606,5609,36011,000 Ar1,5202,6653,9455,7707,2308,78012,000 Ionization Energies in kJ/mol

29. Within each period ( row) the ionization energy increases with atomic number. Y?Y? -Electrons are being added to the same energy level (ENC) - increasing valence electrons as approaching the nonmetals NaMgAlSiPSClAr

30. The Trend

31. Electronegativity l The tendency for an atom to attract electrons to itself when in combination with another atom l Defined differences in electronegativity determine the bonding character of a compound Ionic or Covalent bonds Linus Pauling scale is used to determine electronegativity differences

32. COVALENT BOND  bond formed by the sharing of electron clouds COVALENT BOND  bond formed by the sharing of electron clouds Between nonmetallic elements of similar electronegativity. Formed by sharing electron pairs

34. when electron clouds are shared equally NONPOLAR COVALENT BONDS H 2 or Cl 2

35. 2. Covalent bonds- Two atoms share one or more pairs of outer-shell electrons. Oxygen Atom Oxygen Molecule (O 2 ) Oxygen Molecule (O 2 )

36. when electron clouds are shared but shared unequally when electron clouds are shared but shared unequally POLAR COVALENT BONDS H2OH2OH2OH2O

37. Polar Covalent Bonds: Unevenly matched, but willing to share.

38. - water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

40. Electronegativity Differences and Bond Type nonpolar covalent polar covalent ionic If the electronegativity difference is less than 0.2 then the bond is a nonpolar covalent If the difference is between 0.2 and 1.6, the bond is polar covalent If the difference is greater than 2, the bond is ionic ? ?  between 1.6 and 2, if a metal is involved, the bond is ionic. If only nonmetals are involved the bond is polar covalent 0 4 2 1.6 0.2

41. Trend of EN decrease increase

42. Electron Affinity elements GAIN electrons to form anions. Electron affinity is the energy change when an electron is added: A(g) + e- ---> A - (g) E.A. = ∆E A(g) + e- ---> A - (g) E.A. = ∆E

43. 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. l Affinity for electron increases across a period (EA becomes more negative). l 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

46. Practice with Comparing Ionization Energies For each of the following sets of atoms, decide which has the highest and lowest ionization energies and why. a. Mg, Si, S b. Mg, Ca, Ba c. F, Cl, Br d. Ba, Cu, Ne e. Si, P, N

47. Answers to Comparing Ionization Energies Here are answers to the exercises above. a. Mg, Si, S All are in the same period and use the same number of energy levels. Mg has the lowest I.E. because it has the lowest effective nuclear charge. S has the highest I.E. because it has the highest effective nuclear charge. b. Mg, Ca, Ba All are in the same group and have the same effective nuclear charge. Mg has the highest I.E. because it uses the smallest number of energy levels. Ba has the lowest I.E. because it uses the largest number of energy levels.

48. c. F, Cl, Br All are in the same group and have the same effective nuclear charge. F has the highest I.E. because it uses the smallest number of energy levels. Br has the lowest I.E. because it uses the largest number of energy levels. d. Ba, Cu, Ne All are in different groups and periods, so both factors must be considered. Fortunately both factors reinforce one another. Ba has the lowest I.E. because it has the lowest effective nuclear charge and uses the highest number of energy levels. Ne has the highest I.E. because it has the highest effective nuclear charge and uses the lowest number of energy levels.

49. e. Si, P, N Si has the lowest I.E. because it has the lowest effective nuclear charge and is tied (with P) for using the most energy levels. N has the highest I.E. because it uses the fewest energy levels and is tied (with P) for having the highest effective nuclear charge.

50. BECAUSE... The relative stability of an atom can be predicted by its electron configuration

51. Rule of Thumb As a general rule, elements with three or fewer electrons in their outer level are considered to be metals.

52. Lets Review! 1. What is the periodic Law? 2. How is an element’s outer electron configuration related to its position in the periodic table? 3 Indicate which element in each of the following pairs has the greater atomic radius. a. sodium & lithiumb. strontium & magnesium c. carbon & germaniumd. selenium & oxygen 4. In general, would you expect metals or nonmetals to have higher ionization energies?

53. More review! 5. Arrange the following elements in order of increasing ionization energies. a. Be, Mg, Srb. Bi, Cs, Bac. Na, Al, S 6. How does the ionic radius of a typical metallic atom compare to its atomic radius? 7. Explain why it takes more energy to remove a 4s electron from an atom of zinc than from than from an atom of calcium. 8. Give the symbol of the element found at each of the following locations in the periodic table. a. group 1, period 4b. group 13, period 3 c. group 2, period 6d. group 10, period 2 Test Friday

54. Even more review! 9. What was Newland’s Law of Octaves all about? 10. How was Mendeleev’s periodic table of elements better than the previous attempts by others? 11. What property do the noble gases share? How does this property relate to the electron configuration of the noble gases? 12. How do the electron configurations of the transition metals differ form the electron configurations of the metals in groups 1 and 2? 13. What group numbers make up the main-block elements? This test will be a bear if you forget to study!

55. Are you kidding me, more good stuff! 14. Define what ionization energy and electron affinity are. 15. What periodic trends exist for ionization energy? How about for electron affinity? What about atomic radius and its trend? 16. Why does the first period of the periodic table contain only two elements while all the other periods have eight or more element in them? 17 What feature of electron configuration is unique to actinoids and lanthanoids?

56. That should just about cover it! TO REINFORCE YOUR ALREADY EXTENSIVE KNOWLEDGE OF THE PERIODIC TABLE, YOU CAN READ THROUGH THE PAGES IN YOUR BOOK OF CHAPTER 14. Lastly, if you need to out any last minute problems you can show up at 7:00 in room 224 for some last minute brushing up.

57. Transition cont. IT’S ALL ABOUT ELECTRONS! a Transition elements fill their sublevels differently than do the Main group elements. a For the most part, there are a few exceptions, these “d” block metals will place the 2 electrons into a higher s-sublevel before the electrons go into a “d” energy sublevel.

58. Inner Transition Metals the “ f ”-group l The ‘f’-group is broken into two classifications ~lanthanides ~actinides