1. 1 Chapter 7 Atomic Structure

2. 2 Light n Made up of electromagnetic radiation n Waves of electric and magnetic fields at right angles to each other.

3. 3 Parts of a wave  Wavelength Frequency = number of cycles in one second Measured in hertz 1 hertz = 1 cycle/second

4. 4 Frequency = ν

5. 5 Kinds of EM waves n There are many Different λ and ν Different λ and ν n Radio waves, microwaves, x rays and gamma rays are all examples n Light is only the part our eyes can detect Gamma Rays Radio waves

6. 6 The speed of light n in a vacuum is 2.998 x 10 8 m/s n = c c =  c =  n What is the wavelength of light with a frequency 5.89 x 10 5 Hz? n What is the frequency of blue light with a wavelength of 484 nm?

7. 7 In 1900 n Matter and energy were seen as different from each other in fundamental ways n Matter was particles n Energy could come in waves, with any frequency. n Max Planck found that the cooling of hot objects couldn’t be explained by viewing energy as a wave.

8. 8 Energy is Quantized Planck found ΔE came in chunks with size h Planck found ΔE came in chunks with size h ΔE = nh ν ΔE = nh ν n where n is an integer. n and h is Planck’s constant n h = 6.626 x 10 -34 J s these packets of h ν are called quantum these packets of h ν are called quantum

9. 9 EinsteinEinstein is next Einstein n Said electromagnetic radiation is quantized in particles called photons Each photon has energy = h ν = hc/λ Each photon has energy = h ν = hc/λ n Combine this with E = mc 2 n you get the apparent mass of a photon n m = h / (λc)

10. 10 Which is it? n Is energy a wave like light, or a particle? n Yes n Concept is called the Wave -Particle duality. n What about the other way, is matter a wave? n Yes

11. 11 Matter as a wave Using the velocity v instead of the frequency ν we get Using the velocity v instead of the frequency ν we get De Broglie’s equation  = h/mν De Broglie’s equation  = h/mν n can calculate the wavelength of an object

12. 12 Examples n The laser light of a CD is 7.80 x 10 2 m. What is the frequency of this light? C=λν n What is the energy of a photon of this light? n What is the apparent mass of a photon of this light? n What is the energy of a mole of these photons? E=hc/λ mass=E/c 2 or h/λc

13. 13 What is the wavelength? n of an electron with a mass of 9.11 x 10 -31 kg traveling at 1.0 x 10 7 m/s? n Of a softball with a mass of 0.10 kg moving at 125 m/hr n Pg.296

14. 14 How do they know it is a wave? n When light passes through, or reflects off, a series of thinly spaced lines, it creates a rainbow effect n because the waves interfere with each other.

15. 15 A wave moves toward a slit.

16. 16 Comes out as a curve

17. 17 with two holes

18. 18 with two holes Two Curves

19. 19 Two Curves with two holes Interfere with each other

20. 20 Two Curves with two holes Interfere with each other crests add up

21. 21 Several waves

22. 22 Several waves Several Curves

23. 23 Several waves Interference Pattern Several Curves

24. 24 What will an electron do? n It has mass, so it is matter. n A particle can only go through one hole n A wave goes through both holes n Light shows interference patterns interference patterns interference patterns

25. Electron “gun” Electron as Particle

26. Electron “gun” Electron as wave

27. Which did it do?  It made the diffraction pattern  The electron is a wave  Led to Schrödingers equation

28. 28 What will an electron do? n An electron does go though both, and makes an interference pattern. n It behaves like a wave. n Other matter has wavelengths too short to notice. Image

29. 29

30. 30 Spectrum n The range of frequencies present in light. n White light has a continuous spectrum. n All the colors are possible. n A rainbow. H2 SpectrumH2

31. 31 Hydrogen spectrum n Emission spectrum because these are the colors it gives off or emits n Called a line spectrum. n There are just a few discrete lines showing 410 nm 434 nm 486 nm 656 nm h

32. 32 What this means n Only certain energies are allowed for the hydrogen atom. n Can only give off certain energies. Use ΔE = hν  = hc / λ Use ΔE = hν  = hc / λ n Energy in the atom is quantized

33. 33 Niels Bohr n Developed the quantum model of the hydrogen atom. n He said the atom was like a solar system n The electrons were attracted to the nucleus because of opposite charges. n Didn’t fall in to the nucleus because it was moving around

34. 34 The Bohr Ring Atom n He didn’t know why but only certain energies were allowed. n He called these allowed energies energy levels. n Putting energy into the atom moved the electron away from the nucleus n From ground state to excited state. n When it returns to ground state it gives off light of a certain energy

35. 35 The Bohr Ring Atom n = 3 n = 4 n = 2 n = 1

36. 36 The Bohr Model n n is the energy level n for each energy level the energy is n Z is the nuclear charge, which is +1 for hydrogen. n E = -2.178 x 10 -18 J (Z 2 / n 2 ) n n = 1 is called the ground state n when the electron is removed, n = ∞ n E = 0

37. 37 We are worried about the change n When the electron moves from one energy level to another. n ΔE = E final – E initial n ΔE = -2.178 x 10 -18 J Z 2 (1/ n f 2 - 1/ n i 2 )

38. 38 Examples n Calculate the energy need to move an electron from its initial level to the third energy level. n Calculate the energy released when an electron moves from n= 4 to n=2 in a hydrogen atom. n Calculate the energy released when an electron moves from n= 5 to n=3 in a He +1 ion

39. 39 When is it true? n Only for hydrogen atoms and other monoelectronic species. n Why the negative sign? n To increase the energy of the electron you make it further to the nucleus. n the maximum energy an electron can have is zero, at an infinite distance.

40. 40 The Bohr Model n Doesn’t work n only works for hydrogen atoms n electrons don’t move in circles n the quantization of energy is right, but not because they are circling like planets.

41. 41 The Quantum Mechanical Model n A totally new approach n De Broglie said matter could be like a wave. n De Broglie said they were like standing waves. n The vibrations of a stringed instrument

42. 42

43. 43 What’s possible? n You can only have a standing wave if you have complete waves. n There are only certain allowed waves. n In the atom there are certain allowed waves called electrons. n 1925 Erwin Schroedinger described the wave function of the electron n Much math, but what is important are the solutions

44. 44 Schrödinger’s Equation n The wave function is a F(x, y, z) n Actually F(r,θ,φ) n Solutions to the equation are called orbitals. n These are not Bohr orbits. n Each solution is tied to a certain energy n These are the energy levels AnimationAnimationAnimation

45. 45 There is a limit to what we can know n We can’t know how the electron is moving or how it gets from one energy level to another. n The Heisenberg Uncertainty Principle n There is a limit to how well we can know both the position and the momentum of an object.

46. 46 Mathematically  x ·  (mv) > h/4   x ·  (mv) > h/4   x is the uncertainty in the position  x is the uncertainty in the position  (mv) is the uncertainty in the momentum.  (mv) is the uncertainty in the momentum. the minimum uncertainty is h/4  the minimum uncertainty is h/4 

47. 47 Examples n What is the uncertainty in the position of an electron. mass 9.31 x 10 - 31 kg with an uncertainty in the speed of 0.100 m/s n What is the uncertainty in the position of a baseball, mass 0.145 kg with an uncertainty in the speed of 0.100 m/s

48. 48 What does the wave Function mean? n nothing. n it is not possible to visually map it. n The square of the function is the probability of finding an electron near a particular spot. n best way to visualize it is by mapping the places where the electron is likely to be found.

49. 49 Probability Distance from nucleus

50. 50 Sum of all Probabilities Distance from nucleus

51. 51 Defining the size n The nodal surface. n The size that encloses 90% to the total electron probability. n NOT at a certain distance, but a most likely distance. n For the first solution it is a a sphere.

52. 52 Quantum Numbers n There are many solutions to Schrödinger’s equation n Each solution can be described with quantum numbers that describe some aspect of the solution. n Principal quantum number (n) size and energy of an orbital n Has integer values >0

53. 53 Quantum numbers Angular momentum quantum number l Angular momentum quantum number l n shape of the orbital n integer values from 0 to n-1 l = 0 is called s l = 0 is called s l = 1 is called p l = 1 is called p l =2 is called d l =2 is called d l =3 is called f l =3 is called f l =4 is called g l =4 is called g

54. 54 S orbitals

55. 55 P orbitals

56. 56 P Orbitals

57. 57 D orbitals

58. 58 F orbitals

59. 59 F orbitals

60. 60 Quantum numbers Magnetic quantum number (m l ) Magnetic quantum number (m l ) –integer values between - l and + l –tells direction in each shape Electron spin quantum number (m s ) Electron spin quantum number (m s ) –Can have 2 values –either +1/2 or -1/2

61. 61 1. A 2. B 3. C 4. A 5. B 6. A 7. B 8. A 9. A

62. 62 Polyelectronic Atoms n More than one electron n three energy contributions n The kinetic energy of moving electrons n The potential energy of the attraction between the nucleus and the electrons. n The potential energy from repulsion of electrons

63. 63 Polyelectronic atoms n Can’t solve Schrödinger’s equation exactly n Difficulty is repulsion of other electrons. n Solution is to treat each electron as if it were effected by the net field of charge from the attraction of the nucleus and the repulsion of the electrons. n Effective nuclear charge

64. 64 +11 11 electrons e-e- Z eff Sodium Atom +11 10 other electrons e-e-

65. 65 Effective Nuclear charge n Can be calculated from E = -2.178 x 10 -18 J (Z eff 2 / n 2 ) n And n ΔE = -2.178 x 10 -18 J Z eff 2 (1/ n f 2 - 1/ n i 2 )

66. 66 The Periodic Table n Developed independently by German Julius Lothar Meyer and Russian Dmitri Mendeleev (1870”s) n Didn’t know much about atom. n Put in columns by similar properties. n Predicted properties of missing elements.

67. 67 Aufbau Principle n Aufbau is German for building up n As the protons are added one by one, the electrons fill up hydrogen- like orbitals. n Fill up in order of energy

68. 68 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f 6f Orbitals available to a Hydrogen atom

69. 69 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f With more electrons, repulsion changes the energy of the orbitals.

70. 70 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f He with 2 electrons

71. 71 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

72. 72 Details n Valence electrons- the electrons in the outermost energy levels (not d). n Core electrons- the inner electrons n Hund’s Rule- The lowest energy configuration for an atom is the one have the maximum number of unpaired electrons in the orbital. n C 1s 2 2s 2 2p 2

73. 73 Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2 electrons 2s 2 4 2p 6 3s 2 12 3p 6 4s 2 20 3d 10 4p 6 5s 2 38 4d 10 5p 6 6s 2 56

74. 74 Details n Elements in the same column have the same electron configuration. n Put in columns because of similar properties. n Similar properties because of electron configuration. n Noble gases have filled energy levels. n Transition metals are filling the d orbitals

75. 75 The Shorthand n Write the symbol of the noble gas before the element n Then the rest of the electrons. n Aluminum - full configuration n 1s 2 2s 2 2p 6 3s 2 3p 1 n Ne is 1s 2 2s 2 2p 6 n so Al is [Ne] 3s 2 3p 1

76. 76 The Shorthand Sn- 50 electrons The noble gas before it is Kr [ Kr ] Takes care of 36 Next 5s 2 5s 2 Then 4d 10 4d 10 Finally 5p 2 5p 2 [ Kr ]5s 2 4d 10 5p 2

77. 77 Exceptions n Ti = [Ar] 4s 2 3d 2 n V = [Ar] 4s 2 3d 3 n Cr = [Ar] 4s 1 3d 5 n Mn = [Ar] 4s 2 3d 5 n Half filled orbitals n Scientists aren’t certain why it happens n same for Cu [Ar] 3d 10 4s 1

78. 78 More exceptions n Lanthanum La: [Xe] 5d 1 6s 2 n Cerium Ce: [Xe] 5d 1 4f 1 6s 2 n Promethium Pr: [Xe] 4f 3 6s 2 n Gadolinium Gd: [Xe] 4f 7 5d 1 6s 2 n Lutetium Pr: [Xe] 4f 14 5d 1 6s 2 n We’ll just pretend that all except Cu and Cr follow the rules.

79. 79 More Polyelectronic n We can use Z eff to predict properties, if we determine it’s pattern on the periodic table. n Can use the amount of energy it takes to remove an electron for this. n Ionization Energy- The energy necessary to remove an electron from a gaseous atom.

80. 80 Remember this n E = -2.18 x 10 -18 J(Z 2 /n 2 ) n was true for Bohr atom. n Can be derived from quantum mechanical model as well n for a mole of electrons being removed n E =(6.02 x 10 23 /mol)2.18 x 10 -18 J(Z 2 /n 2 ) n E= 1.13 x 10 6 J/mol(Z 2 /n 2 ) n E= 1310 kJ/mol(Z 2 /n 2 )

81. 81 Example n Calculate the ionization energy of B +4

82. 82 Remember our simplified atom +11 11 e- Z eff 1 e-

83. 83 This gives us n Ionization energy = 1310 kJ/mol(Z eff 2 /n 2 ) n So we can measure Z eff n The ionization energy for a 1s electron from sodium is 1.39 x 10 5 kJ/mol. n The ionization energy for a 3s electron from sodium is 4.95 x 10 2 kJ/mol. n Demonstrates shielding

84. 84 Shielding n Electrons on the higher energy levels tend to be farther out. n Have to look through the other electrons to see the nucleus. n They are less effected by the nucleus. n lower effective nuclear charge n If shielding were completely effective, Z eff = 1 n Why isn’t it?

85. 85 Penetration n There are levels to the electron distribution for each orbital 2s

86. 86 Graphically Penetration 2s Radial Probability Distance from nucleus

87. 87 Graphically Radial Probability Distance from nucleus 3s

88. 88 Radial Probability Distance from nucleus 3p

89. 89 Radial Probability Distance from nucleus 3d

90. 90 Radial Probability Distance from nucleus 4s 3d

91. 91 Penetration effect n The outer energy levels penetrate the inner levels so the shielding of the core electrons is not totally effective. n from most penetration to least penetration the order is n ns > np > nd > nf (within the same energy level) n This is what gives us our order of filling, electrons prefer s and p

92. 92 How orbitals differ n The more positive the nucleus, the smaller the orbital. n A sodium 1s orbital is the same shape as a hydrogen 1s orbital, but it is smaller because the electron is more strongly attracted to the nucleus. n The helium 1s is smaller as well n This provides for better shielding

93. 93 Z eff 1 2 4 5 1 Atomic Number

94. 94 Z eff 1 2 4 5 1 If shielding is perfect Z= 1 Atomic Number

95. 95 Z eff 1 2 4 5 1 No shielding Z = Z eff Atomic Number

96. 96 Z eff 1 2 4 5 16 Atomic Number

97. 97 Periodic Trends n Ionization energy the energy required to remove an electron form a gaseous atom n Highest energy electron removed first. First ionization energy ( I 1 ) is that required to remove the first electron. First ionization energy ( I 1 ) is that required to remove the first electron. Second ionization energy ( I 2 ) - the second electron Second ionization energy ( I 2 ) - the second electron n etc. etc.

98. 98 Trends in ionization energy n for Mg I 1 = 735 kJ/moleI 1 = 735 kJ/mole I 2 = 1445 kJ/moleI 2 = 1445 kJ/mole I 3 = 7730 kJ/moleI 3 = 7730 kJ/mole n The effective nuclear charge increases as you remove electrons. n It takes much more energy to remove a core electron than a valence electron because there is less shielding

99. 99 Explain this trend n For Al I 1 = 580 kJ/moleI 1 = 580 kJ/mole I 2 = 1815 kJ/moleI 2 = 1815 kJ/mole I 3 = 2740 kJ/moleI 3 = 2740 kJ/mole I 4 = 11,600 kJ/moleI 4 = 11,600 kJ/mole

100. 100 Across a Period Generally from left to right, I 1 increases because Generally from left to right, I 1 increases because n there is a greater nuclear charge with the same shielding. As you go down a group I 1 decreases because electrons are further away and there is more shielding As you go down a group I 1 decreases because electrons are further away and there is more shielding

101. 101 It is not that simple Z eff changes as you go across a period, so will I 1 Z eff changes as you go across a period, so will I 1 n Half-filled and filled orbitals are harder to remove electrons from n here’s what it looks like

102. 102 First Ionization energy Atomic number

103. 103 First Ionization energy Atomic number

104. 104 First Ionization energy Atomic number

105. 105 Atomic Size n First problem where do you start measuring n The electron cloud doesn’t have a definite edge. n They get around this by measuring more than 1 atom at a time

106. 106 Atomic Size n Atomic Radius = half the distance between two nuclei of a diatomic molecule } Radius

107. 107 Trends in Atomic Size n Influenced by two factors n Shielding n More shielding is further away n Charge on nucleus n More charge pulls electrons in closer

108. 108 Group trends n As we go down a group n Each atom has another energy level n So the atoms get bigger H Li Na K Rb

109. 109 Periodic Trends n As you go across a period the radius gets smaller. n Same energy level n More nuclear charge n Outermost electrons are closer NaMgAlSiPSClAr

110. 110 Overall Atomic Number Atomic Radius (nm) H Li Ne Ar 10 Na K Kr Rb

111. 111 Electron Affinity n The energy change associated with adding an electron to a gaseous atom n High electron affinity gives you energy- n exothermic n More negative n Increase (more - ) from left to right –greater nuclear charge. n Decrease as we go down a group –More shielding

112. 112 Ionic Size n Cations form by losing electrons n Cations are smaller than the atom they come from n Metals form cations n Cations of representative elements have noble gas configuration.

113. 113 Ionic size n Anions form by gaining electrons n Anions are bigger than the atom they come from n Nonmetals form anions n Anions of representative elements have noble gas configuration.

114. 114 Configuration of Ions n Ions always have noble gas configuration n Na is 1s 2 2s 2 2p 6 3s 1 n Forms a 1+ ion - 1s 2 2s 2 2p 6 n Same configuration as neon n Metals form ions with the configuration of the noble gas before them - they lose electrons

115. 115 Configuration of Ions n Non-metals form ions by gaining electrons to achieve noble gas configuration. n They end up with the configuration of the noble gas after them.

116. 116 Group trends n Adding energy level n Ions get bigger as you go down Li +1 Na +1 K +1 Rb +1 Cs +1

117. 117 Periodic Trends n Across the period nuclear charge increases so they get smaller. n Energy level changes between anions and cations Li +1 Be +2 B +3 C +4 N -3 O -2 F -1

118. 118 Size of Isoelectronic ions n Iso - same n Iso electronic ions have the same # of electrons n Al +3 Mg +2 Na +1 Ne F -1 O -2 and N -3 n all have 10 electrons n all have the configuration 1s 2 2s 2 2p 6

119. 119 Size of Isoelectronic ions n Positive ions have more protons so they are smaller Al +3 Mg +2 Na +1 Ne F -1 O -2 N -3

120. 120 Electronegativity

121. 121 Electronegativity n The tendency for an atom to attract electrons to itself when it is chemically combined with another element. n How “greedy” n Big electronegativity means it pulls the electron toward itself. n Atoms with large negative electron affinity have larger electronegativity.

122. 122 Group Trend n The further down a group more shielding n Less attracted (Z eff ) n Low electronegativity.

123. 123 Periodic Trend n Metals are at the left end n Low ionization energy- low effective nuclear charge n Low electronegativity n At the right end are the nonmetals n More negative electron affinity n High electronegativity n Except noble gases

124. 124 Ionization energy, electronegativity Electron affinity INCREASE

125. 125 Atomic size increases, Ionic size increases

126. 126 Parts of the Periodic Table

127. 127 The information it hides n Know the special groups n It is the number and type of valence electrons that determine an atom’s chemistry. n You can get the electron configuration from it. n Metals lose electrons have the lowest IE n Non metals- gain electrons most negative electron affinities

128. 128 The Alkali Metals n Doesn’t include hydrogen- it behaves as a non-metal n decrease in IE n increase in radius n Decrease in density n decrease in melting point n Behave as reducing agents

129. 129 Reducing ability n Lower IE< better reducing agents n Cs>Rb>K>Na>Li n works for solids, but not in aqueous solutions. n In solution Li>K>Na n Why? n It’s the water -there is an energy change associated with dissolving

130. 130 Hydration Energy n Li + (g) → Li + (aq) is exothermic n for Li + -510 kJ/mol n for Na + -402 kJ/mol n for K + -314 kJ/mol n Li is so big because of it has a high charge density, a lot of charge on a small atom. n Li loses its electron more easily because of this in aqueous solutions

131. 131 The reaction with water n Na and K react explosively with water n Li doesn’t. Even though the reaction of Li has a more negative  H than that of Na and K Even though the reaction of Li has a more negative  H than that of Na and K n Na and K melt  H does not tell you speed of reaction  H does not tell you speed of reaction n More in Chapter 12.