2. Modern Atomic Theory

3. Please select the appropriate Team. 1.Girls 2.Guys

4. A_O_ 1.W H 2.M N 3.M T 4.G Z

5. R_T_E_F_R_ 1.H R U O D 2.N I V W O 3.U V M E R 4.Q O R O H

6. The fact that carbon dioxide always contains 73 percent oxygen by mass is an illustration of 1.the ideas of Democritus. 2.the law of conservation of matter. 3.the law of definite composition. 4.Dalton's atomic theory of matter. 10

7. Which of the following is not one of the three fundamental particles that makeup atoms? 1.electron 2.proton 3.neutron 4.alpha particle 10

8. The atomic number of an atom is defined as its 1.mass in amu. 2.number of electrons. 3.number of neutrons. 4.number of protons. 10

9. An ion always contains an 1.unequal number of protons and electrons. 2.equal number of protons and electrons. 3.unequal number of protons and neutrons. 4.equal number of protons and neutrons. 10

10. Two atoms are isotopes if they contain 1.different numbers of electrons. 2.different numbers of protons and different numbers of neutrons. 3.the same number of protons but different numbers of neutrons. 4.the same number of neutrons but different numbers of protons. 10

11. Team Scores 212.17Girls 166Guys

12. In the last 200 years, vast amounts of data have been accumulated to support atomic theory. When atoms were originally suggested by the early Greeks, no physical evidence existed to support their ideas. Early chemists did a variety of experiments, which culminated in Dalton’s model of the atom. Because of the limitations of his model, modifications were proposed first by Thomson and then Rutherford, which eventually led to our modern concept of the nuclear atom.

13. As with the mystery box you cannot directly see what is inside, we cannot easily see what is inside an atom. Scientists have studied energy and light for centuries, and several models have been proposed to explain how energy is transferred from place to place. One is through the electromagnetic spectrum.

15. Three Characteristics of Waves Wavelength: λ, distance between consecutive peaks or troughs in a wave Frequency: f, how many waves pass a particular point per second. Speed: v, how fast a wave moves through space (m/s)

16. We have evidence for the wavelike nature of light. We know that a beam of light behaves like a stream of tiny packets of energy called photons.Flinn Spectra SlidesFlinn Spectra Slides

17. Which of the following is not a form of electromagnetic radiation? 1.X-rays 2.gamma rays 3.sound waves 4.visible light 10

18. Which of the following has the longest wavelength? 1.ultraviolet radiation 2.infrared radiation 3.X-rays 4.gamma rays 10

19. What does the wavelength of the substance tell you? 1.Energy of the atom 2.Identity of the atom 3.Color of light given off from the atom 4.All of the above 5.None of the above

20. Where does this connect to the atom? At high temperatures or when subjected to high voltages, elements in the gaseous state give off colored light. Brightly colored neon signs illustrate this property well. When the light is emitted by a gas is passed through a prism or diffraction grating, a set of brightly colored lines called a line spectrum results.

21. Continuous and Line Spectra

22. Fireworks

23. Composition of Fireworks Gunpowder –Sulfur, charcoal, potassium nitrate (saltpeter) Salts (to give color) –Red = lithium –Green = copper

24. Neils Bohr proposed a model of the atom, where electrons are found in energy levels and they jump from one energy level to another by adding or losing a quantum of energy.

25. Neils Bohr studied the line spectrum of hydrogen which led him to believe that electrons exist in specific regions at various distances from the nucleus. He visualized that electrons revolved in orbits around the nucleus like the planets revolve around the sun.

26. Bohr applied Planck’s concept of energy quanta. Quanta: A small discrete package of energy, energy is not emitted in a continuous stream.

27. Bohr’s Ideas Electrons have several possible energies These energies correspond to several possible orbits at different distances from the nucleus Therefore an electron has to be in a specific energy level, not between levels. The electron is quantized When a hydrogen atom absorbed one or more quanta of energy, it jumped to a higher energy level or orbital.

28. The color of the light emitted corresponds to one of the lines of the hydrogen spectrum. 10.4 When an electron falls from a higher energy level to a lower energy level a quantum of energy in the form of light is emitted by the atom.

29. Principal Quantum Number (n): tells the shell of the electron a.Ground state (n=1) lowest possible energy level, closest to the nucleus b.Excited state (n=2, 3, 4, etc) when the electron absorbs the appropriate amount of energy, it jumps to a level of higher energy. Radiation is emitted when the electron falls back from a higher energy level to a lower one.

30. In Bohr’s model of hydrogen, the energy levels that the electrons occupy are similar to lanes on a freeway. The lowest energy level, ground state, is like the slow lane on the freeway. This is where electrons are normally found. When the electron receives more energy from an outside source, it has to change to the next highest energy level. Similarly, a car in the slow lane that increases its speed will move to a faster lane. Both electrons and cars will return to the ground state when their energy or speed decreases…we can’t drive between the lanes, therefore electrons cannot stay between levels. This is basically known as Aufbau’s principle.

31. The maximum number of electrons in a energy level = 2n 2 Energy level 1 = 2(1) 2 = 2 electrons Energy level 2 = 2(2) 2 = 8 electrons Energy level 3 = 2(3) 2 = 18 electrons Energy level 4 = 2(4) 2 = 32 electrons

32. Bohr Model: theoretical model where electrons orbit the nucleus in defined energy levels. 1.Gather information – number of protons, neutrons, and electrons 2.Build the nucleus – protons and neutrons 3.Draw the energy levels 4.Place the electrons Practice: H, He, Cl, Cl +2

33. Modern Theory (1925) Bohr’s calculations succeeded very well in correlating the experimentally observed spectral lines with electron energy levels for the hydrogen atom. His methods did not succeed for heavier atoms. Bohr’s Model no longer explained all observations. New theory proposed that atoms did not travel in definite paths but behaved more like waves on a vibrating string.

34. Electrons as Waves Louis de Broglie (1924) –Applied wave-particle theory to electrons –electrons exhibit wave properties QUANTIZED WAVELENGTHS Adapted from work by Christy Johannesson www.nisd.net/communicationsarts/pages/chem Standing Wave 200 150 100 50 0 - 50 -100 -150 -200 0 50 100 150 200 Second Harmonic or First Overtone 200 150 100 50 0 - 50 -100 -150 -200 0 50 100 150 200 Fundamental mode 200 150 100 50 0 - 50 -100 -150 -200 0 50 100 150 200 Louis de Broglie ~1924

35. Electrons as Waves QUANTIZED WAVELENGTHS Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

36. Electrons as Waves Evidence: DIFFRACTION PATTERNS ELECTRONS VISIBLE LIGHT Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Davis, Frey, Sarquis, Sarquis, Modern Chemistry  2006, page 105

37. Quantum Mechanics (1926) Erwin Schrödinger created a mathematical model of electrons as waves.(1926) Erwin Schrödinger created a mathematical model of electrons as waves. Schrödinger’s work led to a new branch of physics called wave or quantum mechanics.Schrödinger’s work led to a new branch of physics called wave or quantum mechanics. Using Schrödinger’s wave mechanics, the probability of finding an electron in a certain region around the atom can be determined.Using Schrödinger’s wave mechanics, the probability of finding an electron in a certain region around the atom can be determined. The actual location of an electron within an atom cannot be determined.The actual location of an electron within an atom cannot be determined. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Erwin Schrodinger ~1926

38. Quantum Mechanics Schrödinger Wave EquationSchrödinger Wave Equation quantized –finite # of solutions  quantized energy levels probability –defines probability of finding an electron Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Erwin Schrodinger ~1926

39. Orbitals This makes it quite difficult to pinpoint an exact location, but a region can be predicted. An orbital is a region in an atom where there is a high probability of finding electrons.

40. Quantum Mechanics Orbital (“electron cloud”) –Region in space where there is 90% probability of finding an electron Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Electron Probability vs. Distance Electron Probability (%) Distance from the Nucleus (pm) 100150200250500 0 10 20 30 40 Orbital 90% probability of finding the electron

41. 5 points on tomorrow’s test for attending the cancer assistance ballgame tonight. You must check in with me to get credit.

42. The specific location of an electron can be determined. 1.True 2.False

43. Electrons cannot be located between energy levels or orbitals. 1.True 2.False

46. Heisenberg’s Uncertainty Principle The more precisely the position is determined, the less precisely the momentum is known in this instant, and vice versa. --Heisenberg, uncertainty paper, 1927 http://youtu.be/noZWLPpj3to ?list=PL07E6A22017705261http://youtu.be/noZWLPpj3to ?list=PL07E6A22017705261

47. Quantum Mechanics Heisenberg Uncertainty PrincipleHeisenberg Uncertainty Principle –Impossible to know both the velocity and position of an electron at the same time Microscope Electron  Werner Heisenberg ~1926

50. Electron Configuration UPPER LEVEL Allows us to map the location of all electrons Specify the “address” of each electron’s region in an atom Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

51. An Atom’s Address An electron subshell is a region of space within an electron shell that contains electrons that have the same energy.

52. grape : raisin :: plum : 1.peach 2.fig 3.apricot 4.prune

53. Clark Kent : Superman :: Bruce Wayne : 1.Hulk 2.Batman 3.Spiderman 4.Wolverine

54. Electron Subshell : energy level ::.. : region of space 1.shell 2.orbital 3.level 4.path

55. Team Scores

57. Figure 3.7 The number of subshells within a shell is equal to the shell number, as shown here for the first four shells. Each individual subshell is denoted with both a number (its shell) and a letter (the type of subshell it is in).

58. Maximum Number of Electrons In Each Sublevel Maximum Number of Electrons In Each Sublevel Sublevel orMaximum Number Subshell Number of Orbitals of Electrons S (sharp) 1 2 P (principal) 3 6 D (diffuse) 5 10 F (fundamental) 7 14 LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World, 1996, page 146

59. Figure 3.8 An s orbital has a spherical shape, a p orbital has two lobes, a d orbital has four lobes, and an f orbital has eight lobes. S orbitals: hold a maximum of 2 electrons

60. p-Orbitals Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 335 pxpx pypy pzpz P orbitals: hold a maximum of 6 electrons

61. d-orbitals Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 336 D orbitals: hold a maximum of 10 electrons

62. f-orbitals Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 336 F orbitals: hold a maximum of 14 electrons

63. Electrons in an orbital each have an intrinsic spin. One has spin up ↑ and one has spin down ↓. Electrons in the same orbital cannot have the same spin. There must be one of each ↑↓..\ElectronSpin.exe Electron Spin

64. Electron Configuration Filling-Order of Electrons in an Atom

65. 4f4f 4d4d 4p4p 4s4s n = 4 3d3d 3p3p 3s3s n = 3 2p2p 2s2s n = 2 1s1s n = 1 Energy Sublevels

66. 4f4f 4d4d 4p4p 4s4s n = 4 3d3d 3p3p 3s3s n = 3 2p2p 2s2s n = 2 1s1s n = 1 Energy Sublevels s s s s p p p d df 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 …

67. Filling Rules for Electron Orbitals Aufbau Principle: Electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for. Pauli Exclusion Principle: An orbital can hold a maximum of two electrons. To occupy the same orbital, two electrons must spin in opposite directions. Hund’s Rule: Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results. *Aufbau is German for “building up”

68. Energy Level Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N H = 1s 1 Hydrogen H He Li C N Al Ar F Fe LaHeLiCNAlArFFeLa

69. Energy Level Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N He = 1s 2 Helium HH He Li C N Al Ar F Fe LaLiCNAlArFFeLa

70. Energy Level Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N Li = 1s 2 2s 1 Lithium HH He Li C N Al Ar F Fe LaHeCNAlArFFeLa

71. Energy Level Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N C = 1s 2 2s 2 2p 2 Carbon HH He Li C N Al Ar F Fe LaHeLiNAlArFFeLa

72. Energy Level Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N N = 1s 2 2s 2 2p 3 Bohr Model Nitrogen Hund’s Rule “maximum number of unpaired orbitals”. HH He Li C N Al Ar F Fe LaHeLiCAlArFFeLa

73. Energy Level Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N F = 1s 2 2s 2 2p 5 Fluorine HH He Li C N Al Ar F Fe LaHeLiCNAlArFeLa

74. Energy Level Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N Al = 1s 2 2s 2 2p 6 3s 2 3p 1 Aluminum HH He Li C N Al Ar F Fe LaHeLiCNArFFeLa

75. Energy Level Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N Ar = 1s 2 2s 2 2p 6 3s 2 3p 6 Bohr Model Argon HH He Li C N Al Ar F Fe LaHeLiCNAlFFeLa

76. Energy Level Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS CLICK ON ELEMENT TO FILL IN CHARTS Fe = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 N HH He Li C N Al Ar F Fe LaHeLiCNAlArFLa Bohr Model Iron Electron Configuration

77. Energy Level Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS CLICK ON ELEMENT TO FILL IN CHARTS La = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 N HH He Li C N Al Ar F Fe LaHeLiCNAlArFFe Bohr Model Lanthanum Electron Configuration

79. An orbital diagram is a statement of how many electrons an atom has in each of its electron orbitals. (arrow diagrams) An electron configuration is a statement of how many electrons an atom has in each of its electron subshells. A ground state electron configuration is the location of electron distributions in energy levels when at ground state. (MOST STABLE)

80. Electron Configurations and Orbital Diagrams There are many orbitals about the nucleus of an atom. Electrons do not occupy these orbitals in a random, haphazard fashion; a very predictable pattern exists for electron orbital occupancy. 1.Electron subshells are filled in order of increasing energy. 2.Electrons occupy the orbitals of a subshell such that each orbital acquires one electron before any orbital acquires a second electron. 3.No more than two electrons may exist in a given orbital-and then only if they have opposite spins.

81. Orbital Diagrams Element # of electrons 1s2s2px2py2pz3s H He B N F Na

82. Orbital Diagrams Element # of electrons 1s2s2px2py2pz3s Electron Configuration H 1↑ 1s 1 He 2↑↓ 1s 2 B 5↑↓ ↑ 1s 2 2s 2 2p 1 N 7↑↓ ↑↑↑ 1s 2 2s 2 2p 3 F 9↑↓ ↑ 1s 2 2s 2 2p 5 Na 11↑↓ ↑ 1s 2 2s 2 2p 6 3s 1

83. A summary of the interrelationships among electron shells, electron subshells, and electron orbitals for the first four shells. Similar relationship patterns exist for higher-numbered shells.

85. Electron Configurations and the Periodic Table The distinguishing electron is the last electron added to the electron configuration for an element when electron subshells are filled in order of increasing energy, thus the purpose of writing electron configurations. Steps in writing ground state electron configurations: 1.Determine the number of electrons 2.Using the diagonal rule, write the abbreviations until the superscripts add up to the total of electrons.

86. Quanta are fundamental "pieces" of 1.energy. 2.matter. 3.nuclei. 4.electrons. 10

87. Quanta of light are called 1.electrons. 2.protons. 3.photons. 4.joules. 10

88. The probability of finding electrons in certain regions of an atom is described by 1.orbits. 2.orbitals. 3.quanta. 4.photons. 10

89. Under what conditions can two electrons occupy the same orbital? 1.never 2.if they have opposite spins 3.if they have parallel spins 4.if they have different principal quantum numbers 10

90. What information about electrons is given by the electron configuration of an atom? 1.paths within the principal quantum level 2.density of the electron cloud 3.average angular momentum 4.distribution among orbitals 10

91. When electrons are in the lowest- energy orbitals available, the atom is 1.unstable. 2.in an excited state. 3.in the ground state. 4.chemically unreactive. 10

92. Determine the electron configuration for phosphorus (P, atomic number 15). 1.1s 2 2s 2 2p 6 3p 5 2.1s 2 2s 2 2p 6 3s 2 3p 3 3.1s 2 2s 2 2p 6 3s 5 4.1s 2 2s 2 2p 6 3s 1 3p 4 10

93. 1.2. 3.4. Determine the orbital diagram for the p sublevel of the next element, sulfur (S, atomic number 16). 10

94. Pick up a new periodic table and cut it apart and reattach as shown. 1 2 1s

95. F 9e’s – 1s 2 2s 2 2p 5 1 st 2 nd 2e-7e- Cl 17e’s – 1s 2 2s 2 2p 6 3s 2 3p 5 1 st 2 nd 3 rd 2e-8e-7e- Fe 26e’s – 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 1 st 2 nd 3 rd 4 th 2e-8e-14e-2e- Au 79e’s – 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 9 1 st 2 nd 3 rd 4 th 5 th 6 th 2e-8e-18e-32e-17e-2e-

96. Quantum Model of the Atom Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

97. Figure 3.8 An s orbital has a spherical shape, a p orbital has two lobes, a d orbital has four lobes, and an f orbital has eight lobes. S orbitals: hold a maximum of 2 electrons

98. p-Orbitals Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 335 pxpx pypy pzpz P orbitals: hold a maximum of 6 electrons

99. d-orbitals Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 336 D orbitals: hold a maximum of 10 electrons

100. f-orbitals Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 336 F orbitals: hold a maximum of 14 electrons

101. Electrons in an orbital each have an intrinsic spin. One has spin up ↑ and one has spin down ↓. Electrons in the same orbital cannot have the same spin. There must be one of each ↑↓..\ElectronSpin.exe Electron Spin

102. Quantum Numbers Principal Quantum Number n Principal Quantum Number ( n ) Angular Momentum Quantum # l Angular Momentum Quantum # ( l ) Magnetic Quantum Number m l Magnetic Quantum Number ( m l ) Spin Quantum Number Spin Quantum Number ( m s )

103. Quantum Numbers Principal Quantum Number n 1. Principal Quantum Number ( n ) –Energy level –Size of the orbital –n 2 = # of orbitals in the energy level –1, 2, 3, 4, 5, 6, 7 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem 1s1s 2s2s 3s3s

104. Quantum Numbers s p d f Angular Momentum Quantum # l 2. Angular Momentum Quantum # ( l ) –Energy sublevel –Shape of the orbital spdfspdf 01230123 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

105. Quantum Numbers Orbitals combine to form a spherical shape. 2s 2p z 2p y 2p x Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

106. Quantum Numbers Magnetic Quantum Number m l 3. Magnetic Quantum Number ( m l ) –Orientation of orbital –Specifies the exact orbital within each sublevel Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

107. Shapes of s, p, and d-Orbitals

108. Atomic Orbitals

109. s p d f □ □□□ □□□□□ □□□□ □□ □ 0 1 0 -1 2 1 0 -1 -2 3 2 1 0 -1 -2 -3

110. Quantum Numbers Magnetic Spin Quantum Number 4. Magnetic Spin Quantum Number ( m s ) –Electron spin  +½ or -½ –An orbital can hold 2 electrons that spin in opposite directions. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

111. Quantum Numbers 1. Principal #  2. Ang. Mom. #  3. Magnetic #  4. Spin #  energy level sublevel (s,p,d,f) orbital electron Pauli Exclusion PrinciplePauli Exclusion Principle –No two electrons in an atom can have the same 4 quantum numbers. –Each electron has a unique “address”: Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

112. Feeling overwhelmed? Read Section 4-2! Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem "Teacher, may I be excused? My brain is full." Chemistry

113. According to Heisenberg’s uncertainty principle, you cannot determine simultaneously an electron’s 1.position and velocity 2.mass and velocity 3.position and mass 4.size and position 10

114. The Pauli exclusion principle indicates that the 3s orbital can hold a maximum of 1.2 electrons with opposite spin. 2.2 electrons with identical spin. 3.3 electrons with identical spin. 4.3 electrons with variable spin. 10

115. According to the wave mechanical model, which of the following statements is not correct? 1.Four quantum numbers describe an electron in an atom. 2.Two electrons in an atom can have all four identical quantum numbers. 3.The energy of an electron increases with the principal quantum number. 4.An electron can have one of two kinds of spin. 10

116. Examples: Be 1s 2 2s 2 2 00 -1/2 Fe 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 3 2 2 -1/2