4. Quantum Numbers At the conclusion of our time together, you should be able to:  List and define each of the 4 quantum numbers.  Relate these numbers to the state, city, street and home address for the electron.  Give the maximum number of electrons for each level and sublevel.  Draw the basic shape of the 4 sublevels.

5. Principal Quantum Number - n Symbol = n Represents the main energy level of the electron and its distance from the nucleus Equation: 2n 2 - shows how many electrons can be in each energy level (e.g. 3 rd energy level: 2(3) 2 = 18 total possible e - in this energy level) Your turn: How many electrons in the 4 th energy level? 32

6. Principal Quantum Number - n In an address analogy, this would be the state in which the electron would probably be found. Presently, we can find electrons in 7 states. Values are 1-7 Ex. = 1s 1 (the electron configuration for H) Principal Quantum number = 1

7. The First of 4 Quantum Numbers for the One Electron of Hydrogen n (energy level) 1

8. The Second Quantum Number - l This number describes sublevels, shapes of these sublevels and is the Angular momentum quantum number. The number of sublevels (shapes) in an energy level equals the value of n, the principle quantum number. The first 4 sublevels are named s, p, d, f. Each sublevel has a unique shape:

9. Shape of the “s” Orbital s for "Sphere": the simplest shape, or shape of the simplest atoms like hydrogen and helium Electrons don't interfere with, or block, each other from the pull of the nucleus - ball shape Each energy level has an "s" orbital at the lowest energy within that level

10. p for "Peanut/Petal": a more complex shape that occurs at energy levels 2 and above Shapes of the “p” Orbitals

11. d for "Double Peanut/Petal": a complex shape occurring at energy levels 3 and above The arrangement of these orbitals allows for "s" and "p" orbitals to fit closer to the middle/nucleus Shapes of the “d” Orbitals

12. f for "Flower": bizarre-shaped orbitals for electrons of very large atoms electrons filling these orbitals are weakly attached to the atom because they are so far away from the pull of the nucleus Shapes of the “f” Orbitals

13. The Second Quantum Number - l In the address analogy, this would be the city in which the electron would probably be found. In the first state there would be 1 city, in the second state there would be 2 cities... This number is from the formula: 0 to (n-1)

14. The Second Quantum Number - l The 2 nd Quantum Number = 0 to (n-1) Ex. = 1s 1, s sublevel number is 0 to (1-1) = 0 p sublevel number is, 0 to (2-1) = 1 d sublevel number is, 0 to (3-1) = 2 f sublevel number is, 0 to (4-1) = 3 Therefore: For s, l = 0 For p, l = 1 For d, l = 2 For f, l = 3

15. The Second of 4 Quantum Numbers for the One Electron of Hydrogen n (energy level) 1 l (angular momentum, sublevel shape) 0 Ex. = 1s 1, 0 to (1-1) = 0 Second Quantum number for Hydrogen = 0

16. The Third Quantum Number - m This number describes orientations of the orbitals that each sublevel can have and is the Magnetic quantum number. s has one orientation, p has three orientations, d has five orientations and f has 7 orientations. We’ll see how we got these orientation numbers in just a moment.

17. The Third Quantum Number - m In the address analogy, this would be the street on which the electron would probably be found. In the first state there is one city. In this first city there would be one street. In the second state there are two cities. The first city with its one street and a second city with its 3 streets for a total of 4. In the 3rd state would have the previous 4 streets plus 5 more streets from the 3 rd city for a total of 9.

18. The Third Quantum Number - m 3 rd Quantum Number = – l to 0 to + l We know what the shape looks like from the second Quantum Number, now we know how many orientations of these shapes each sublevel has by plugging the l value into the formula above.

19. The Third Quantum Number - m Value = – l to 0 to + l, therefore: s = ___ 0 Or just 1 orientation

20. The Third Quantum Number - m Values are – l to 0 to + l, therefore: p = ___ ___ ___ -1 0 +1 Or 3 orientations

21. The Third Quantum Number - m Values are – l to 0 to + l, therefore: d = ___ ___ ___ ___ ___ -2 -10 1 2 or 5 orientations

22. The Third Quantum Number - m Values are – l to 0 to + l, therefore: f = ___ ___ ___ ___ ___ ___ ___ -3 -2 -1 0 1 2 3 or 7 orientations

23. The Third Quantum Number - m Let’s Review the Values for m: s = ___p = ___ ___ ___ 0 -1 0 +1 d = ___ ___ ___ ___ ___ -2 -10 1 2 f = ___ ___ ___ ___ ___ ___ ___ -3 -2 -1 0 1 2 3 Ex. = 1s 1 = 0 Third Quantum number from s = 0, or 1 shape

24. The Third of the 4 Quantum Numbers for the One Electron of Hydrogen n (energy level) 1 l (angular momentum, sublevel shape) 0 m (magnetic, orientation) 0

25. The Fourth Quantum Number - s This number describes the spin of the electrons in an orbital. There can be two electrons in each orbital as long as they are spinning in opposite directions. In the address analogy, this would be the house number of the electron. There can only be 2 houses on each street Values are +1/2 = clockwise spin -1/2 = counter clockwise spin

26. The Fourth Quantum Number - s Ex. = 1s 1, spin is up or +1/2 Fourth Quantum number from 1 = +1/2

27. The Fourth of the 4 Quantum Numbers for the One Electron of Hydrogen n (energy level) 1 l (angular momentum, sublevel shape) 0 m (magnetic, orientation) 0 s (spin, clockwise or counterclockwise) +1/2

28. Quantum # Summary for Hydrogen 1s 1 Hydrogen has one electron spinning in a clockwise direction in the first energy level that has one orbital and only one orientation in its s sublevel shape which is spherical. The 4 quantum numbers for H are: 1, 0, 0, +1/2

29. Quantum Numbers Let’s see if you can:  List and define each of the 4 quantum numbers.  Relate these numbers to the state, city, street and home address for the electron.  Give the maximum number of electrons for each level and sublevel.  Draw the basic shape of the 4 sublevels.

30. How do we know when an electron has moved from an excited state to the ground state? The electron will 1. Release a photon. 2. Release a specific amount of energy. 3. Release a specific color. 4. Release a quantum of energy. 5. All of the above are correct.

31. What does the second quantum number ( l ) describe? 1. Orbital shape. 2. Energy level. 3. Electron spin. 4. Orbital orientation.

32. Which quantum number has values of +½ or –½? 1. Orbital shape. 2. Energy level. 3. Electron spin. 4. Orbital orientation.

33. Which of the following is not a possible orbital shape? 1. s 2. p 3. z 4. f

34. What is currently the highest possible principal quantum number an electron can have? 1. No limit 2. 5 3. 6 4. 7

35. How many electrons will the 4 th energy level hold? 1. No limit 2. 8 3. 18 4. 32 5. 50

36. 1. sphere 2. petal 3. double petal 4. flower 5. mess The p sublevel would look like a

37. Quantum Numbers At the conclusion of our time together, you should be able to:  Continue to relate Quantum Numbers to the state, city, street and home address for the electron.  Give the 4 quantum number for every electron of every atom on the periodic table.

38. Let’s Try Helium H is 1s 1 He has 2 electrons, can we add another electron spinning in the other direction in the first energy level of the s sublevel with its 1 spherical orbital? Yes, He is 1s 2 He has 2 electrons spinning in opposite directions in the s sublevel with its spherical shape with 1 orientation.

39. Remember the Quantum # Summary for Hydrogen 1s 1 Hydrogen has one electron spinning in a clockwise direction in the first energy level that has one spherical orbital in its s sublevel. The 4 quantum numbers for H are: 1, 0, 0, +1/2

40. The Pauli Exclusion Principle No two electrons in an atom can have the same set of four quantum numbers. Therefore, the second electron that He has must have a different set of 4 Quantum Numbers. What would they be??

41. The 4 Quantum Numbers of Helium n1n1 l0l0 m0m0 s -1/2

42. The 4 Quantum Numbers of Helium No two electrons in an atom can have the same set of four quantum numbers. What principle?? Pauli Exclusion Principle Therefore, the second electron that He has must have a different set of 4 Quantum Numbers. What would they be?? 1, 0, 0, -1/2

43. Let’s Try Lithium He is 1s 2 He has 2 electrons, can we add another electron spinning in another direction in the first energy level of the s sublevel with its 1 spherical orbital? No, the third electron must go to the 2 nd energy level which has 2 sublevels, s and p, s with its one spherical orbital and p with its 3 orientations of its petal shaped orbitals.

44. Let’s Try Lithium So the address for all 3 electrons of Li is: 1s 2 2s 1 This is called the electron configuration for Li and basically says that the first two electrons of Li are in the s sublevel with its one spherical shape with the 3 rd electron in the 2 energy level, s sublevel with its bigger spherical shape. The Quantum #’s of the 3 rd electron:

45. The 4 Quantum Numbers for the 3 rd Electron of Lithium n2n2 l0l0 m0m0 s +1/2

46. Now Beryllium: Add one more electron to Li : 1s 2 2s 1 ? Where would it go?? 1s 2 2s 2 The 2s sublevel with its one spherical orbital can hold 2 electrons spinning in opposite directions. The Quantum #’s of the 4 th electron:

47. The 4 Quantum Numbers for the 4 th Electron of Beryllium n2n2 l0l0 m0m0 s -1/2

48. What About Boron: Add one more electron to Be: 1s 2 2s 2 ? Where would it go?? 1s 2 2s 3 ? No 1s 2 2s 2 2p 1 The 2s sublevel with its one spherical orbital is full. We must now start filling the 2p sublevel with its 3 orbitals The Quantum #’s of the 5 th electron:

49. The Electrons of the 2 nd Energy Level for Boron s = ___p = ___ ___ ___ 0 -1 0 +1

50. The 4 Quantum Numbers for the 5 th Electron of Boron n2n2 l1l1 m s +1/2

51. Let’s Move to Carbon: Add one more electron to B: 1s 2 2s 2 2p 2 The 2p sublevel has 3 orbitals. However, we can’t put another electron in the first orbital of p. Why? Hund’s Rule: Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron.

52. The Electrons of the 2 nd Energy Level for Carbon s = ___p = ___ ___ ___ 0 -1 0 +1

53. The 4 Quantum Numbers for the 6 th Electron of Carbon n2n2 l1l1 m0m0 s +1/2

54. Nitrogen: Add one more electron to C: 1s 2 2s 2 2p 3 The 3 rd electron in the p sublevel must go into the 3 rd orbital. Why? Again Hund’s Rule: Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron.

55. The Electrons of the 2 nd Energy Level for Carbon s = ___p = ___ ___ ___ 0 -1 0 +1

56. The 4 Quantum Numbers for the 7 th Electron of Nitrogen n2n2 l1l1 m1m1 s +1/2

57. Oxygen: Add one more electron to N: 1s 2 2s 2 2p 4 The 4 th electron in the p sublevel must begin to double up the electrons of the three p orbitals.

58. The Electrons of the 2 nd Energy Level for Carbon s = ___p = ___ ___ ___ 0 -1 0 +1

59. The 4 Quantum Numbers for the 8 th Electron of Oxygen n2n2 l1l1 m s -1/2

60. Fluorine: Add one more electron to O: 1s 2 2s 2 2p 5 Where would the 5 th electron for the 2p orbitals go?

61. The Electrons of the 2 nd Energy Level for Fluorine s = ___p = ___ ___ ___ 0 -1 0 +1

62. The 4 Quantum Numbers for the 9 th Electron of Fluorine n2n2 l1l1 m0m0 s -1/2

63. Neon: Add one more electron to F: 1s 2 2s 2 2p 6 Where would the 6 th electron for the 2p orbitals go?

64. The Electrons of the 2 nd Energy Level for Neon s = ___p = ___ ___ ___ 0 -1 0 +1

65. The 4 Quantum Numbers for the 10 th Electron of Neon n2n2 l1l1 m1m1 s -1/2

66. What About Sodium: Add one more electron to Ne: 1s 2 2s 2 2p 7 No!! The 2p sublevel with its 3 orbitals is full. We must now go to the 3 rd energy level with its 3 sublevels, s, p, and d, and start filling electrons all over again!

67. The Electrons of the 3 rd Energy Level for Sodium s = ___p = ___ ___ ___ 0 -1 0 +1 d = ___ ___ ___ ___ ___ -2 -10 1 2

68. The 4 Quantum Numbers for the 11 th Electron of Sodium n3n3 l0l0 m0m0 s +1/2

69. What About Magnesium: Add one more electron to Na: 1s 2 2s 2 2p 6 3s 2 Remember, the s sublevel with its one spherical orbital can have one more electron.

70. The Electrons of the 3 rd Energy Level for Magnesium s = ___p = ___ ___ ___ 0 -1 0 +1 d = ___ ___ ___ ___ ___ -2 -10 1 2

71. The 4 Quantum Numbers for the 12 th Electron of Magnesium n3n3 l0l0 m0m0 s -1/2

72. Aluminum: Add one more electron to Mg: 1s 2 2s 2 2p 6 3s 2 3p 1 Remember, the s sublevel with its one spherical orbital can only have 2 electrons. The next electron must start to fill the 3p sublevel with its 3 orbitals.

73. The Electrons of the 3 rd Energy Level for Aluminum s = ___p = ___ ___ ___ 0 -1 0 +1 d = ___ ___ ___ ___ ___ -2 -10 1 2

74. The 4 Quantum Numbers for the 13 th Electron of Aluminum n3n3 l1l1 m s +1/2

75. Silicon to Argon: Keep adding one more electron to Al: 1s 2 2s 2 2p 6 3s 2 3p 2-6 Remember Hund’s rule. We will put one electron in each of the p orbitals and then come back to double up.

76. The Electrons of the 3 rd Energy Level for Silicon to Argon s = ___p = ___ ___ ___ 0 -1 0 +1 d = ___ ___ ___ ___ ___ -2 -10 1 2

77. The 4 Quantum Numbers for the 18 th Electron of Argon n3n3 l1l1 m1m1 s -1/2

78. Quantum Numbers Let’s see if you can:  Continue to relate Quantum Numbers to the state, city, street and home address for the electron.  Give the 4 quantum number for every electron of every atom on the periodic table.

79. The 5f sublevel with its orbitals would look like a 1. sphere 2. petal 3. double petal 4. flower 5. mess

80. How many orbitals are in the 5f sublevel? 1. No limit 2. 1 3. 3 4. 5 5. 7

81. How many electrons will the 4f sublevel hold? 1. 2 2. 6 3. 10 4. 14 5. 18

82. 1. 1 2. 2 3. 3 4. 4 5. 5 In the address analogy, the first quantum number is the state and the second quantum number is the city. How many cities are there in the second state?

83. 1. 1 2. 2 3. 3 4. 4 5. 5 In the address analogy, the second state can have how many total streets?

84. Quantum Numbers At the conclusion of our time together, you should be able to:  Explain the Aufbau principle and the diagonal rule.  Use Hund’s rule in an orbital filling diagram.  Give the quantum numbers for every electron of every atom on the periodic table.  Draw the electron configuration and orbital notation for every element.

85. Now Potassium: We should now start to fill the 3d sublevel with its 5 orbitals. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 1 But now we must consider the Aufbau Principle: An electron will occupy the lowest energy level orbital that can receive it. Look at the energy needed for 3d vs. 4s: ???

87. The Aufbau Principle You will need to remember this chart so that you place electrons in the correct order. An easy way to remember the filling order is to follow the diagonal rule. Check it out on the next slide:

88. Maybe You’ve Seen This Chart??? s s 3p 3d s 2p s 4p 4d 4f s 5p 5d 5f 5g? s 6p 6d 6f 6g? 6h? s 7p 7d 7f 7g? 7h? 7i? 1234567

89. It Represents the Diagonal Rule Steps: 1.Write the energy levels top to bottom. 2.Write the orbitals in s, p, d, f order. Write the same number of orbitals as the energy level. 3.Draw diagonal lines from the top right to the bottom left. 4.To get the correct order, follow the arrows!

90. Diagonal Rule s s 3p 3d s 2p s 4p 4d 4f s 5p 5d 5f 5g? s 6p 6d 6f 6g? 6h? s 7p 7d 7f 7g? 7h? 7i? 1234567 By this point, we are past the current periodic table so we can stop.

91. The Aufbau Principle But there is even an easier way to remember the filling order. Have you noticed areas in the periodic table where certain sublevels are filling?? Check it out on the next slide:

92. The 4 Blocks of the Periodic Table

93. The Aufbau Principle So, what block is always filling on the left side of the periodic table?? The s block!! What period is it?? The 4 th !! Therefore, 4s will begin filling before we fill 3d!!!

94. So is Potassium? 1s 2 2s 2 2p 6 3s 2 3p 6 3d 1 No!!! 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1

95. The Electrons of the 4 rd Energy Level for Potassium and Calcium s = ___p = ___ ___ ___ 0 -1 0 +1 d = ___ ___ ___ ___ ___ -2 -10 1 2

96. The 4 Quantum Numbers for the 20 th Electron of Calcium n4n4 l0l0 m0m0 s -1/2

97. The “D” Block Electrons The next electron for Scandium will now start to fill the “D” block. Please note that the energy level for the d sublevel is not 4. We have not filled 3d yet. Therefore, the “D” block will always be one energy level behind the current energy level.

98. d for "Double Peanut/Petal": complex shape occurring at energy levels 3 and above How many orbitals (orientations) does “d” have? 5 Remember the Shape of the “d” Orbitals

99. So for Scandium to Zinc Following Hund’s Rule, we will put one electron in each of the 3d orbitals before we begin to double up the electrons in each orbital. Remember that the “d” block is always one energy level behind.

100. The Electrons of the 3 rd Energy Level for Scandium to Zinc s = ___p = ___ ___ ___ 0 -1 0 +1 d = ___ ___ ___ ___ ___ -2 -10 1 2

101. The 4 Quantum Numbers for the Circled Electron of Zinc n3n3 l2l2 m1m1 s +1/2

102. Gallium to Krypton So, what block is always filling on the right side of the periodic table?? The p block!! What period is it?? Back to the 4 th !! Therefore, 4p will begin filling after we fill 3d!!!

104. The Electrons of the 4 rd Energy Level for Gallium to Krypton s = ___p = ___ ___ ___ 0 -1 0 +1 d = ___ ___ ___ ___ ___ -2 -10 1 2

105. The 4 Quantum Numbers for the Circled Electron of Krypton n4n4 l1l1 m s -1/2

106. Rubidium and Strontium So, what block is always filling on the left side of the periodic table?? The s block!! What period is it?? 5 th !! Remember, the 4 th energy level has 4 sublevels, s, p, d, and f. Because of the Aufbau Principle, d and f will fill later.

107. Rubidium and Strontium Let’s do the electron configuration for these two elements 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 or 5s 2

108. How About Yttrium to Cadmium? 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 1-10

109. Indium to Xenon? 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 1-6

110. Cesium and Barium? 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 1-2 Now do the 4 quantum numbers for 55 th electron of Barium.

111. The 4 Quantum Numbers for the Circled Electron (55 th ) of Barium n6n6 l0l0 m0m0 s +1/2

112. The “F” Block Elements Now we come to a confusing part of the periodic table. Note that element #57, Lanthanum, is in the “d” block but the next element #58, Cerium, drops down to the “f” block.

113. The Periodic Table

114. The “F” Block Elements Also note that according to the Aufbau Principle, after 6s should fill 4f and then 5d.

116. The “F” Block Elements Recent discoveries suggest that Lanthanum is not the first element of the 4f block as previously thought, but really is the first element of the 5d block. From there we will move to the 4f block.

117. Lanthanum? 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 Then Cerium is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 4f 1

118. The Rest of the 4f Block #58 Praseodymium to #71 Lutetium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 4f 1-14

119. The Electrons of the 4 f Sublevel 6s = ___6p = ___ ___ ___ 0 -1 0 +1 5d = ___ ___ ___ ___ ___ -2 -1 0 1 2 4f = ___ ___ ___ ___ ___ ___ ___ -3 -2 -1 0 1 2 3

120. The 4 Quantum Numbers for the Circled Electron (67 th ) of Lutetium n4n4 l3l3 m s -1/2

121. Now Back to the “D” Block Elements As we follow the numbers on the Periodic Table, you will see that element #72, Hafnium, is back in the “d” block. Therefore, Hafnium’s configuration would be: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 4f 14 5d 2 Right?? Wrong!!

122. Now Back to the “D” Block Elements Why? Count the total electrons… 73 not 72 Why? Because 5d 2 includes 5d 1 I’ve counted 5d 1 twice!! Therefore, I must do this 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 4f 14 5d 2

123. The Rest of the “D” Block Elements Tantalum to Mercury 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 4f 14 5d 3-10

124. Now to the 6p Block Elements Thallium to Radon 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 4f 14 5d 10 6p 1-6

125. What’s Next #87 - #88, Francium and Radium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 4f 14 5d 10 6p 1-6 7s 1-2

126. Now We Run into the Crazy Area #89, Actinium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 4f 14 5d 10 6p 1-6 7s 2 6d 1 #90 - #103, Thorium to Lawrencium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 4f 14 5d 10 6p 1-6 7s 2 6d 1 5f 1-14

127. Finally, We Finish Up in the D Block #104 - #112, Rutherfordium to Copernicium Don’t forget to cross out the 6d 1 electron when you come back to the “d” block 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 4f 14 5d 10 6p 1-6 7s 2 6d 1 5f 14 6d 2-10

128. Quantum Numbers Let’s see if you can:  List and define the 4 principles that are part of quantum numbers.  List and define each of the 4 quantum numbers.  Give the quantum number for every electron of every atom on the periodic table.

129. Your Turn Let’s see if you can do the last 4 orbital filling diagrams and the 4 quantum numbers for the last electron of #112, Copernicium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 4f 14 5d 10 6p 1-6 7s 2 6d 1 5f 14 6d 10

130. The Electrons of the 4 Configurations for Copernicium 7s = ___6p = ___ ___ ___ 0 -1 0 +1 6d = ___ ___ ___ ___ ___ -2 -1 0 1 2 5f = ___ ___ ___ ___ ___ ___ ___ -3 -2 -1 0 1 2 3

131. The 4 Quantum Numbers for the Circled Electron (112 th ) of Copernicium n6n6 l2l2 m2m2 s -1/2

132. It is impossible to determine both the position and the momentum of an electron at the same time.

133. An electron occupies the lowest energy level available.

134. No two electrons in the same atom can have the same set of four quantum numbers. In other words, no two electrons can be in the same place at the same time.

135. Orbitals of equal energy are each occupied by ONE electron before any orbital is occupied by a SECOND electron All electrons in a single occupied orbital must have the same spin.

136. Symbol = n Represents the main energy level of the electron Range = 1- 7 Ex. = 3s 2 Principal Quantum number = 3

137. Symbol = l (small letter L) Represents the shape of the orbital (also called sublevel) Range = 0 – n-1 (whole number) Shapes: 0 = s (sphere)1 = p (petal) 2 = d (double petal)3 = f (flower) Ex. = 3s 2 AM Quantum number = 0

138. Symbol = m Represents the orientation of the orbital around the nucleus Each line holds 2 electrons m = -l to +l; Therefore: s = 0, p = 3, d = 5… ___ = s 0 _________ = p -1 0 +1

139. _______________ = d -2 -1 0 +1 +2 _____________________= f -3 -2 -1 0 +1 +2 +3 Ex. = 3s 2 Magnetic Quantum number = 0

140. 2 Spin States Clockwise spin = +1/2 (upward arrow) Counterclockwise spin = -1/2 (downward arrow) A single orbital can hold two electrons, but they must have opposite spins Ex. = 3s 2 Spin Quantum number = -1/2

141. Congratulations!!!! You can now do the complete electron configurations, orbital notations and give 4 quantum numbers (addresses) for every electron of every element on the periodic table!!!

142. GRADING SCALE A = upright and taking air B = eyes notable open C = responding to the environment D = comatose F = decomposed