2. Chapter 5 Electrons in Atoms

3. Greek Idea l Democritus and Leucippus l Matter is made up of indivisible particles l Dalton - one type of atom for each element

5. 1 Summarize the development of the atomic theory in terms of the following: l Thompson - atom was a positive ball with electrons in it l Rutherford - atom had nucleus with electrons around it

6. Thomson’s Model l Discovered electrons l Atoms were made of positive stuff l Negative electron floating around l “Plum-Pudding” model

7. Rutherford’s Model l Discovered dense positive piece at the center of the atom Nucleus Electrons moved around Mostly empty space

8. Bohr’s Model l Why don’t the electrons fall into the nucleus? Move like planets around the sun. l In circular orbits at different levels. Amounts of energy separate one level from another.

9. l Quantum mechanical model - atom has no definite shape #1. Summarize the development of the atomic theory in terms of the following: l Bohr - atom had electrons in orbits around nucleus

10. Bohr’s Model Nucleus Electron Orbit Energy Levels

11. Bohr’s Model Increasing energy Nucleus First Second Third Fourth Fifth } l Further away from the nucleus means more energy. l There is no “in between” energy l Energy Levels

12. #2. Define energy level and relate the energy level to the rungs of a ladder. Define quantum of energy. There are energy levels around the nucleus. Think of a ladder nucleus 1 st 2 nd 3 rd 4 th These are the regions where an electron is likely to be found. Electrons in levels farther from the nucleus contain more energy Quantum of energy - the amount of energy that moves an electron up one level Etc.

13. Waves vs. particles l Are electrons energy waves or particles? Both. They show properties of each and must be described as such: - as particles - they have mass (1/1840 of a proton) - as energy waves - they can be emitted from an atom as a measurable wave.

14. nucleus 1 st 2 nd 3 rd 4 th Etc. Note: the electron NEVER appears between energy levels. It only exists at the energy levels

15. #3. Explain what is meant by the phrase “the energies of electrons are said to be quantized”. l Quantized energies means that each electron contains a specific amount of energy and that these amounts are not continuous, but units of energy. The electrons can’t “slide” from one energy level to another - they immediately appear in the next level..

16. #4. Explain the significance of quantized energies and the quantum mechanical model of the atom. l It is significant because it changes the way we look at energy. Instead of thinking that there are an infinite number of temperatures between 50 º C and 51ºC, we are forced to admit that there are a finite number of temperatures between these two degrees. And as something heats up, its temperature goes up by jumps rather than flowing smoothly.

17. #5. Explain why it is difficult to build a model of a quantum mechanical model of an atom l To build a model, you must know the where things are and their shape. l The electrons can’t be pinpointed and we only know where they are most likely to be found. l Sometimes electrons are even a long ways from the nucleus. l You can’t model anything this vague. l Electrons also behave as waves and not particles

18. The Bohr Model of the Atom Neils Bohr I pictured electrons orbiting the nucleus much like planets orbiting the sun. But I was wrong! They’re more like bees around a hive. WRONG!!!

19. Quantum Mechanical Model of the Atom Mathematical laws can identify the regions outside of the nucleus where electrons are most likely to be found. These laws are beyond the scope of this class..... But why? Hmmm lets see an example

20. Schrodinger Wave Equation probability Equation for probability of a single electron being found along a single axis (x-axis) Erwin Schrodinger

21. Heisenberg Uncertainty Principle You can find out where the electron is, but not where it is going. OR… You can find out where the electron is going, but not where it is! “One cannot simultaneously determine both the position and momentum of an electron.” Werner Heisenberg

22. The Quantum Mechanical Model – Some Points l Energy is quantized. It comes in chunks. l A quanta is the amount of energy needed to move from one energy level to another. l Since the energy of an atom is never “in between” there must be a quantum leap in energy. l Schrodinger derived an equation that described the energy and position of the electrons in an atom.

23. l Things that are very small behave differently from things big enough to see. l The quantum mechanical model is a mathematical solution. l It is not like anything you can see. The Quantum Mechanical Model – Some Points cont.

24. l Has energy levels for electrons. l Orbits are not circular. l It can only tell us the probability of finding an electron a certain distance from the nucleus. The Quantum Mechanical Model

25. l The atom is found inside a blurry “electron cloud” l A area where there is a chance of finding an electron. l Draw a line at 90 % The Quantum Mechanical Model

26. #6. Distinguish among principal energy level, energy sublevel, and atomic orbital. l Principal energy level Any of the major energy levels that contain electrons - the rungs on the ladder designated as 1, 2, 3, 4, etc. l Energy sublevel Locations within the principal energy level where specific electrons are found. l Atomic orbital An orbital describes the shape of the energy sublevel and designated as s, p, d, & f orbitals Old

27. Electron Energy Level (Shell) Generally symbolized by n, it denotes the probable distance of the electron from the nucleus. Number of electrons that can fit in a shell: 2n 2

28. Orbital shapes are defined as the surface that contains 90% of the total electron probability. An orbital is a region within an energy level where there is a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level…

29. 1 energy sublevel: s orbital 2 energy sublevels: s & p orbitals 3 energy sublevels: s, p & d orbitals 4 energy sublevels: s, p, d & f orbitals #7. State how the principal energy levels are designated and identify the number of sublevels within each principal energy level. Old nucleus

30. Atomic Orbitals l Principal Quantum Number (n) = the energy level of the electron. l Within each energy level the complex math of Schrodinger’s equation describes several shapes. l These are called atomic orbitals l Regions where there is a high probability of finding an electron.

31. l 1 s orbital for every energy level l Spherical shaped l Each s orbital can hold 2 electrons l Called the 1s, 2s, 3s, etc.. orbitals. S orbitals

32. Orbitals of the same shape (s, for instance) grow larger as n increases… Nodes are regions of low probability within an orbital. Sizes of s orbitals

33. P orbitals l Start at the second energy level l 3 different directions l 3 different shapes l Each can hold 2 electrons (opposite spins)

34. - three dumbbell-shaped p orbitals in each energy level above n = 1, - each assigned to its own axis (x, y and z) in space. P orbital shape

35. P Orbitals - all together

36. D orbitals l Start at the third energy level l 5 different shapes l Each can hold 2 electrons

37. Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of “double dumbells ” … and a “dumbell with a donut”!

38. F orbitals fourth Start at the fourth energy level Have seven different shapes 2 electrons per shape

40. Yes - There are no atoms big enough to require orbitals beyond F.

42. Summary s p d f # of shapes Max electrons Starts at energy level 121 362 5103 7144

43. #7. State how the principal energy levels are designated and identify the number of sublevels within each principal energy level. l The principal energy levels are designated as n = 1, n = 2, n = 3, n = 4, n = 5, etc. l There may be as many as four sublevels (orbitals) in each principal energy level 1 has one sublevel called “1s” with 1 orbital 2 has two sublevels 2s, 2p 2p has 3 orbitals 3 has three sublevels 3s, 3p, 3d 3d has 5 orbitals 4 has four sublevels 4s, 4p, 4d, 4f

44. 1s 2s, 2p 3s, 3p, 3d 4s, 4p, 4d, 4f The number of the principal energy level tells you how many sub energy levels (orbitals) are in that level. spdf The letters “s”, “p”, “d” and “f” tell you the shape of the orbital

45. By Energy Level l First Energy Level l only s orbital l only 2 electrons l 1s 2 l Second Energy Level l s and p orbitals are available l 2 in s, 6 in p l 2s 2 2p 6 l 8 total electrons

46. By Energy Level l Third energy level l s, p, and d orbitals l 2 in s, 6 in p, and 10 in d l 3s 2 3p 6 3d 10 l 18 total electrons l Fourth energy level l s,p,d, and f orbitals l 2 in s, 6 in p, 10 in d, ahd 14 in f l 4s 2 4p 6 4d 10 4f 14 l 32 total electrons

47. Orbital filling table

48. 5.2 Electron Configurations

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

50. Electron Configurations (The way electrons are arranged in atoms) l Aufbau principle- electrons enter the lowest energy first. This causes difficulties because of the overlap of orbitals of different energies. l Pauli Exclusion Principle- at most 2 electrons per orbital - different spins

51. #11. State the Pauli exclusion principle. l An electron orbital will contain, at most, two electrons. If two electrons are in the orbital they must have opposite spins.

52. Electron Configuration l Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to. l Let’s determine the electron configuration for Phosphorus l Need to account for 15 electrons

53. #12. State Hund’s rule. l When electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with spins parallel. E.g. the p orbital

54. Writing electron configuration l H = 1s 1 l He = 1s 2 l Li = 1s 2 2s 1 l B = 1s 2 2s 2 2p 1

55. Writing electron configuration l C = 1s 2 2s 2 2p 2 1s 2s 2p’s l N = 1s 2 2s 2 2p 3

56. Write an electron configuration for fluorine l 1s 2 2s 2 2p 5  

57. Write an electron configuration for boron l 1s 2 2s 2 2p 1  

58. l The first to electrons go into the 1s orbital l Notice the opposite spins l only 13 more Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

59. l The next electrons go into the 2s orbital l only 11 more Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

60. The next electrons go into the 2p orbital only 5 more Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

61. The next electrons go into the 3s orbital only 3 more Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

62. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f The last three electrons go into the 3p orbitals. They each go into seperate shapes 3 upaired electrons 1s 2 2s 2 2p 6 3s 2 3p 3

63. The easy way to remember 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

64. 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 2s 2 4 electrons

65. 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 2s 2 2p 6 3s 2 12 electrons

66. 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 2s 2 2p 6 3s 2 3p 6 4s 2 20 electrons

67. 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 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 38 electrons

68. 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 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 56 electrons

69. 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 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 10 6p 6 7s 2 88 electrons

70. 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 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 10 6p 6 7s 2 5f 14 6d 10 7p 6 108 electrons

71. The periodic table is arranged to show which orbital is being filled s orbitalp orbitald orbitalf orbital

72. State the maximum number of electrons that can go into each of the following sublevels: l 2s 2 l 3p 6 l 4s 2 l 3d 10 l 4p 6 l 5s 2 l 4f 14 l 5p 6

73. State the number of electrons in the highest occupied energy level of these atoms. l Barium 2p 1 one l Sodium 3s 1 one l Aluminum 3p 1 one

74. Write electron configurations for atoms of these elements: l selenium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 4 [Ar] 4s 2 3d 10 4p 4 l vanadium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 [Ar] 4s 2 3d 3 l nickel 1s 2 2s 2 2p 6 3s 2 3p 6 3s 3 3p 6 [Ar] 4s 2 3d 8 l calcium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 [Ar] 4s 2 l oxygen 1s 2 2s 2 2p 4 [He] 2s 2 2p 4

75. s orbital is being filled s orbital is being filled p orbital is being filled d orbital is being filled

76. Exceptions to Electron Configuration

77. Orbitals fill in order l Lowest energy to higher energy. l Adding electrons can change the energy of the orbital. l Half filled orbitals have a lower energy. l Makes them more stable. l Changes the filling order

78. Write these electron configurations l Titanium - 22 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 l Vanadium - 23 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 l Chromium - 24 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 is expected l But this is wrong!!

79. Chromium is actually l 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 l Why? l This gives us two half filled orbitals. l Slightly lower in energy. l The same principal applies to copper.

80. Copper’s electron configuration l Copper has 29 electrons so we expect 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 l But the actual configuration is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 l This gives one filled orbital and one half filled orbital. l Remember these exceptions

81. Light/ Waves A Quick Review

82. Light l The study of light led to the development of the quantum mechanical model. l Light is a kind of electromagnetic radiation. l Electromagnetic radiation includes many kinds of waves l All move at 3.00 x 10 8 m/s ( c)

83. Parts of a wave Wavelength Amplitude Orgin Crest Trough

84. Parts of Wave l Origin - the base line of the energy. l Crest - high point on a wave l Trough - low point on a wave l Amplitude - distance from origin to crest l Wavelength - distance from crest to crest is abbreviated  Greek letter lambda.

85. Frequency l The number of waves that pass a given point per second. l Units are cycles/sec or hertz (hz) Abbreviated  the Greek letter nu c =

86. Frequency and wavelength l Are inversely related l As one goes up the other goes down. l Different frequencies of light is different colors of light. l There is a wide variety of frequencies l The whole range is called a spectrum

87. Radio waves Micro waves Infrared. Ultra- violet X- Rays Gamma Rays Low energy High energy Low Frequency High Frequency Long Wavelength Short Wavelength Visible Light

88. Atomic Spectrum Sec 5.3 How color tells us about atoms

89. Prism l White light is made up of all the colors of the visible spectrum. l Passing it through a prism separates it.

90. If the light is not white l By heating a gas with electricity we can get it to give off colors. l Passing this light through a prism does something different.

91. Atomic Spectrum l Each element gives off its own characteristic colors. l Can be used to identify the atom. l How we know what stars are made of.

92. Atoms can emit or absorb only specific frequencies of light

93. These are called discontinuous spectra Or line spectra unique to each element. These are emission spectra The light is emitted given off.

94. Light is a Particle l Energy is quantized. l Light is energy l Light must be quantized l These smallest pieces of light are called photons. l Energy and frequency are directly related.

95. Energy and frequency E = h x n l E is the energy of the photon n is the frequency l h is Planck’s constant h = 6.6262 x 10 -34 Joules sec. l joule is the metric unit of Energy

96. The Math in Chapter 5 l Only 2 equations c = ln E = h n l Plug and chug.

97. Examples l What is the wavelength of blue light with a frequency of 8.3 x 10 15 hz? l What is the frequency of red light with a wavelength of 4.2 x 10 -5 m? l What is the energy of a photon of each of the above?

98. An explanation of Atomic Spectra So Where Does It Come From Anyway?

99. Where the electron starts l When we write electron configurations we are writing the lowest energy. l The energy level and electron starts from is called its ground state.

100. So……… l Do the electrons ever change orbitals? l YES!

101. For Electrons to Change Orbitals they lMlMust absorb energy to move up to a higher orbital (energy level) lMlMust release energy to move down to a lower orbital (energy level)

102. Energy in the Atom is Quantified l Max Plank theorized that electrons can absorb or emit only specific amounts of energy. l Quantum amounts – no more no less, must be exact.

103. Many Things are Quantified

104. Some are not

105. The Energy Involved is Light l A photon is a quantum of light energy.

106. Lets Watch What Happens l http://spiff.rit.edu/classes/phys301/lec tures/spec_lines/Atoms_Nav.swf http://spiff.rit.edu/classes/phys301/lec tures/spec_lines/Atoms_Nav.swf

107. Changing the energy l Let’s look at a hydrogen atom

108. Changing the energy l Heat or electricity or light can move the electron up energy levels

109. Changing the energy l As the electron falls back to ground state it gives the energy back as light

110. l May fall down in steps l Each with a different energy Changing the energy

111. { { { UltravioletVisible Infrared

112. l Further they fall, more energy, higher frequency. l This is simplified l the orbitals also have different energies inside energy levels l All the electrons can move around. Ultraviolet Visible Infrared

113. Emission Spectra of an Atom l Acts like a fingerprint l Can Identify the type of atom from its spectral lines. l How helium was first discovered, by observing the Sun’s light spectra

114. Part of the Sun’s Spectrum Do you think the sun is made of more than one thing? Why?

115. ALL the Light You SEE lWlWas at one point emitted by an excited atom.

116. So What is light? l Light is a particle - it comes in chunks. l Light is a wave - we can measure its wave length and it behaves as a wave If we combine E=mc 2, c=ln, E = 1/2 mv 2 and E = hn We can get l = h/mv l The wavelength of a particle.

117. Matter is a Wave l Does not apply to large objects l Things bigger that an atom l A baseball has a wavelength of about 10 - 32 m when moving 30 m/s l An electron at the same speed has a wavelength of 10 - 3 cm l Big enough to measure.

118. The physics of the very small l Quantum mechanics explains how the very small behaves. l Classic physics is what you get when you add up the effects of millions of packages. l Quantum mechanics is based on probability because

119. Heisenberg Uncertainty Principle l It is impossible to know exactly the speed and velocity of a particle. l The better we know one, the less we know the other. l The act of measuring changes the properties.

120. More obvious with the very small l To measure where a electron is, we use light. l But the light moves the electron l And hitting the electron changes the frequency of the light.

121. Moving Electron Photon Before Electron Changes velocity Photon changes wavelength After