2. Electrons in Atoms

3. Basic Properties of Waves

4. Amplitude l Amplitude is the maximum distance the particles of the medium carrying the wave move away from their rest positions.

5. Amplitude l Greater amplitude corresponds with greater energy.

6. Wavelength Wavelength ( λ ) is the distance between two corresponding parts of a wave.

7. Frequency l Frequency equals the number of complete waves that pass a given point in a certain amount of time.

8. Frequency and Wavelength l Frequency and wavelength are inversely related, which means that as one goes up the other goes down.

9. Frequency and Wavelength l Different wavelengths (and frequencies) of light correspond to different colors of light.

10. The Quantum Concept l In 1900, the German physicist Max Planck began searching for an explanation as he studied the light emitted from heated objects.

12. The Quantum Concept l Matter can gain or lose energy only in small, specific amounts called quanta. l That is, a quantum is the minimum amount of energy that can be gained or lost by an atom.

13. The Quantum Concept

14. l While a beam of light has many wavelike characteristics, it also can be thought of as a stream of tiny particles, or bundles of energy, called photons. l Thus, a photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.

15. Energy and Frequency l Planck showed that the energy of radiation increases as the radiation’s frequency increases. l Energy and frequency are directly proportional.

16. Wave-Particle Duality of Electrons

17. The Bohr Model of the Atom l Niels Bohr produced a model of the hydrogen atom based on experimental observations.

18. The Bohr Model of the Atom l This model indicated that 1.An electron circles the nucleus only in fixed energy ranges called orbits.

19. The Bohr Model of the Atom 2.An electron can neither gain nor lose energy inside this orbit, but could move up or down to another orbit.

20. The Bohr Model of the Atom 3.The lowest energy orbit is closest to the nucleus.

21. Modern View l The atom has two regions and is 3- dimensional. l The nucleus is at the center and contains the protons and neutrons.

22. Modern View l The electron cloud is the region where you might find an electron (the region of high probability) and most of the volume of an atom.

23. Wave-Particle Duality of Electrons l Recall that electrons have wave- like properties, although they are normally thought of as particles.

24. Particle Properties of Electrons l There is always an integer number of electrons orbiting the nucleus. l Electrons jump between orbitals (orbits) in a particle-like fashion.

25. Wave Properties of Electrons l The electrons do not orbit the nucleus in the sense of a planet orbiting the sun, but instead exist as standing waves.

26. Wave Properties of Electrons l The electrons are never in a single point location, although there is an area of high probability (the electron cloud) of interacting with the electron at a single point.

27. The Quantum Mechanical Model l The lowest allowable energy state of an atom is called its ground state.

28. The Quantum Mechanical Model l When an electron gains an amount of energy equivalent to the energy difference, it is said to be in an excited state. l The electron moves from its ground state to a higher energy level.

29. The Quantum Mechanical Model

30. l When the electron moves to a lower energy level, it releases an amount of energy equal to the energy difference in these levels as electromagnetic radiation (emissions spectrum).

31. The Quantum Mechanical Model l This electromagnetic radiation is given off as photons.

33. l The spectrum of light released from excited atoms of an element is called the emission spectrum of that element.

34. These are called line spectra. Each is unique to an element.

36. l The further electrons fall, the greater the energy. l This results in a higher frequency because energy and frequency are directly proportional.

37. l Bohr’s model of the atom helped to explain spectral lines.

38. Question 1. Use the Chemistry Reference Tables to answer the following: a) An electron falls from energy level 5 to energy level 3. What is the wavelength of the light emitted? (1282 nm)

39. Question b) An electron falls from energy level 6 to energy level 2. What is the wavelength of the light emitted? (410 nm)

40. Question c) An electron falls from energy level 3 to energy level 1. What type of electromagnetic radiation is emitted (infrared, visible or ultraviolet)? (UV)

41. Question d) An electron falls from energy level 4 to energy level 2. What type of electromagnetic radiation is emitted (infrared, visible or ultraviolet)? (visible)

42. Question e) An electron falls from energy level 5 to energy level 2. What color of visible light is emitted? (blue)

43. Question f) An electron falls from energy level 3 to energy level 2. What color of visible light is emitted? (red)

44. The Quantum Mechanical Model l The space around the nucleus of an atom where the atom’s electrons are found is called the electron cloud.

45. l A three-dimensional region around the nucleus called an atomic orbital describes the electron’s probable location.

46. l In general, electrons reside in principal energy levels. Energy Levels

47. l As the energy level number (n) increases, the orbital becomes larger, the electron spends more time farther from the nucleus, and the atom’s energy level increases.

48. Sublevels l Principal energy levels contain energy sublevels. l Principal energy level 1 consists of a single sublevel, principal energy level 2 consists of two sublevels, principal energy level 3 consists of three sublevels, and so on.

49. Sublevels l Sublevels are labeled s, p, d, or f. l The s sublevel can hold 2 electrons, the p sublevel can hold 6 electrons, the d sublevel can hold 10 electrons, and the f sublevel can hold 14 electrons.

50. p block d block s block f block

51. Orbitals l Sublevels contain orbitals. l Each orbital may contain at most two electrons.

52. l The s sublevel starts at energy level 1. l There is one s orbital forevery energy level l The s orbital is spherical shaped. l The s orbital can only hold 2 electrons. s sublevel and orbital

53. p sublevel and orbitals l The p sublevel starts at the second energy level l p orbitals have 3 dumbbell shapes l The p sublevel can hold 6 electrons

54. d sublevel and orbitals l d orbitals start at the third energy level l d orbitals have 5 different shapes l The d sublevel can hold 10 electrons.

55. f sublevel and orbitals l f orbitals start at the fourth energy level l f orbitals have seven different shapes l The f sublevel can hold 14 electrons.

56. f orbitals

57. Summary s p d f # of shapes (orbitals) Max # of electrons Starts at energy level 1 21 3 62 5103 7144

58. Maximum Number of Electrons in Energy Levels l The maximum number of electrons in an energy level (n) equals 2n 2. l For example, the maximum number of electrons found in the 4 th energy level is 2 (4) 2 = 2 (16) = 32.

59. p block d block s block f block

60. l Really have to include helium. l Helium has the properties of the noble gases. s2s2 s1s1 s- block s2s2

61. Transition Metals - d block d1d1 d2d2 d3d3 d4d4 d5d5 d6d6 d7d7 d8d8 d9d9 d 10

62. The p-block p1p1 p2p2 p3p3 p4p4 p5p5 p6p6

63. f - block l inner transition elements

64. l d orbitals fill up after previous energy level so first d is 3d even though it’s on row 4. 12345671234567 3d

65. l f orbitals start filling at 4f 12345671234567 4f 5f

66. l Each row (or period) is the energy level for s and p orbitals. 12345671234567

68. Electron Configurations l Electron configurations represent the way electrons are arranged in atoms. l Electrons enter the lowest energy level first.

69. Electron Configurations 2. Determine the electron configuration for phosphorus (P). l The atomic number of phosphorus is 15, so we need to account for 15 electrons.

70. Electron Configurations 1s 2 2 electrons 2s 2 4 electrons 2p 6 10 electrons 3p 3 15 electrons The 3p sublevel can hold 6 electrons, but only 3 are needed to get the 15 electrons. 3s 2 12 electrons

71. Electron Configurations 3. Determine the electron configuration for chromium (Cr). l The atomic number of chromium is 24, so we need to account for 24 electrons.

72. Electron Configurations 1s 2 2 electrons 2s 2 4 electrons 2p 6 10 electrons 3p 6 18 electrons The 3d orbital does not need all 10 electrons to get the 24 electrons. 3s 2 12 electrons 4s 2 20 electrons 3d 4 24 electrons

73. Question 4. Write the electron configuration for aluminum (Al). (1s 2 2s 2 2p 6 3s 2 3p 1 )

74. Question 5. Write the electron configuration for neon (Ne). (1s 2 2s 2 2p 6 )

75. Question 6. Write the electron configuration for calcium (Ca). (1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 )

76. Question 7. Write the electron configuration for iron (Fe). (1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 )

77. Question 8. Write the electron configuration for bromine (Br). (1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 )

78. Electron Configuration and the Bohr Model

79. Electron configuration was first conceived of under the Bohr model of the atom.

80. Electron Configuration and the Bohr Model Consider the electron configuration and Bohr model image for silicon (Si). 1s 2 2s 2 2p 6 3s 2 3p 2

81. Electron Configuration and the Bohr Model There are 2 electrons in the first energy level (orbit). 1s 2 2s 2 2p 6 3s 2 3p 2

82. Electron Configuration and the Bohr Model There are 8 electrons in the second energy level (orbit). 1s 2 2s 2 2p 6 3s 2 3p 2

83. Electron Configuration and the Bohr Model There are 4 electrons in the third energy level (orbit). 1s 2 2s 2 2p 6 3s 2 3p 2

84. STOP HERE

85. Bohr Model Review An electron falls from energy level 6 to energy level 2. What is the wavelength AND color of visible light is emitted? (410 nm, violet) Use the Chemistry Reference Tables to answer the following:

86. ELECTRON CONFIGURATION REVIEW 1. Write the electron configuration for oxygen (O). (1s 2 2s 2 2p 4 )

87. ELECTRON CONFIGURATION REVIEW 2. Write the electron configuration for titanium (Ti). (1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 )

88. Orbital Diagrams l Remember sublevels contain orbitals, and each orbital may contain at most two electrons. l The 2 electrons in an orbital must have opposite “spins.”

89. Orbital Diagrams l In addition, when electrons occupy orbitals of equal energy, they don’t pair up with an electron of opposite spin until they have to.

90. Orbital Diagrams Ne

91. Example 9. Draw the orbital diagram for oxygen (O). The electron configuration for oxygen is: 1s 2 2s 2 2p 4 The orbital diagram is: 1s2s2p

92. Problem 10. Draw the orbital diagram for silicon (Si). The electron configuration for silicon is: 1s 2 2s 2 2p 6 3s 2 3p 2 The orbital diagram is: 1s2s2p3s3p

93. Problem 11. Draw the orbital diagram for potassium (K). The electron configuration for is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 The orbital diagram is: 1s2s 2p3s3p4s

94. Using Electron Configurations for Element Identification l To identify an element with a given electron configuration, add the superscript numbers together and find the element with that atomic number.

95. Example 12. Identify the element with the following electron configuration: 1s 2 2s 2 2p 6 3s 1

96. Example l 1s 2 2s 2 2p 6 3s 1 l 2 + 2 + 6 + 1 = 11 l Element 11 is sodium (Na).

97. Question 13. Identify the element with the following electron configuration: 1s 2 2s 2 2p 6 3s 2 3p 4 (sulfur - S)

98. Question 14. Identify the element with the following electron configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 (copper - Cu)

99. Question 15. Identify the element with the following electron configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 2 (germanium - Ge)

100. Electron Configuration Using a Noble Gas Abbreviation l In order to write this type of configuration, find the noble gas (from Group 8A) that comes before the element in question.

101. Electron Configuration Using a Noble Gas Abbreviation l Put the symbol for the noble gas in brackets and then write the part of the configuration that follows to reach the desired element.

102. Example 16. Write the electron configuration using a noble gas abbreviation for magnesium (Mg). l Neon is the noble gas that proceeds magnesium.

103. Example l Put neon’s symbol in brackets. l [Ne] l Now use the periodic table to determine the rest of the configuration.

104. p block d block s block f block 1 2 3 4 5 6 7 Neon is in light blue; magnesium is in bright yellow. The additional configuration is 3s 2. The noble gas electron configuration for magnesium is: [Ne] 3s 2

105. Example 17. Write the electron configuration using a noble gas abbreviation for nickel (Ni). l Argon is the noble gas that proceeds nickel.

106. Example l Put argon’s symbol in brackets. l [Ar] l Now use the periodic table to determine the rest of the configuration.

107. p block d block s block f block 1 2 3 4 5 6 7 Argon is in light blue; nickel is in bright yellow. The additional configuration is 4s 2 3d 8. (Remember you subtract 1 from the d sublevel row number.) The noble gas electron configuration for nickel is: [Ar] 4s 2 3d 8

108. Question 18. Write the electron configuration using a noble gas abbreviation for fluorine (F). ([He] 2s 2 2p 5 )

109. Question 19. Write the electron configuration using a noble gas abbreviation for silicon (Si). ([Ne] 3s 2 3p 2 )

110. Question 20. Write the electron configuration using a noble gas abbreviation for zirconium (Zr). ([Kr] 5s 2 4d 2 )

111. Valence Electrons You can also use the periodic table as a tool to predict the number of valence electrons in any atom in Groups 1, 2, 13, 14, 15, 16, 17, and 18. The electrons in the outermost energy level are called valence electrons.

112. All atoms in Group 1, like hydrogen, have one valence electron. Likewise, atoms in Group 2 have two valence electrons. Valence Electrons

113. All atoms in Group 13 have three valence electrons. All atoms in Group 14 have four valence electrons. All atoms in Group 15 have five valence electrons. Valence Electrons

114. All atoms in Group 16 have six valence electrons. All atoms in Group 17 have seven valence electrons. All atoms in Group 18 have eight valence electrons, except helium which only has two. Valence Electrons

115. All atoms in Groups 3 – 12 (sublevels d and f ) have 2 valence electrons. Valence Electrons

116. Question 21. How many valence electrons does each of the following elements have? a) carbon (C) (4) b) bromine (Br) (7)

117. Question c) iron (Fe) (2) d) potassium (K) (1) e) aluminum (Al) (3)

118. Question 22. How many valence electrons does each have based on the following electron configuration? a) [Kr] 5s 2 4d 2 (2) b) [He] 2s 2 2p 5 (7)

119. Question c) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 (2) d) 1s 2 2s 2 2p 6 3s 2 3p 1 (3) e) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 (2)

120. Lewis Dot Diagrams Because valence electrons are so important to the behavior of an atom, it is useful to represent them with symbols.

121. Lewis Dot Diagrams A Lewis dot diagram illustrates valence electrons as dots (or other small symbols) around the chemical symbol of an element.

122. Each dot represents one valence electron. In the dot diagram, the element’s symbol represents the core of the atom—the nucleus plus all the inner electrons.

123. Electron (Lewis) Dot Diagrams l Write the symbol. X l Put one dot for each valence electron. l Don’t pair electrons until you have to.

124. 23. Write a Lewis dot diagram for chlorine. Question

125. 24. Write a Lewis dot diagram for calcium. Question

126. 25. Write a Lewis dot diagram for potassium. Question