2. Chapter 5 Preview Questions are Due on Monday!

3. Section 5.1 The Development of a New Atomic Model Previously, Rutherford reshaped our thoughts of the atom by showing the protons were located in the nucleus of the atom, but he could not model for us where the electrons were, other than outside the nucleus somewhere. Fortunately, studies into the properties of light and the effects of light on matter soon gave clues to where electrons actually are.

4. Light is a small part of all the radiation (something that spreads from a source) called electromagnetic radiation. Electromagnetic radiation is energy in the form of waves (of electric and magnetic fields). Electromagnetic radiation includes radio waves, microwaves, infrared, visible light, X-rays, and Gamma rays. All these together are considered the Electromagnetic Spectrum.

6. As all the forms of electromagnetic radiation are waves, they all have similar properties. All electromagnetic radiation travels at the speed of light (c), 299,792,458 m/s (2.998 x 10 8 ) in a vacuum

8. The crest is the top of the waves, the trough is the bottom of the waves, and the amplitude is a measurement from the rest or zero line to a crest or trough

10. The wavelength (λ – lambda) is the distance between successive crests/troughs and is measured in meters (often nm, and 1 x 10 9 nm = 1 m)

11. The wavelength (λ – lambda) is the distance between successive crests/troughs and is measured in meters (often nm = 1 x 10 -9 m) The frequency (ν – nu) is the number of waves that pass a point in one second and is measured in (per second – can be written as s -1 ) or Hz (Hertz) 1 s

12. How many hertz is the first wave? How many hertz is the second wave?

13. How many hertz is the first wave? 1 wave per second = 1 Hz How many hertz is the second wave? 2 waves per second = 2 Hz

14. The speed of a wave is directly proportional to the wavelength and the frequency; c = λν is the formula c λ ν

15. Example. A certain violet light has a wavelength of 413 nm. What is the frequency of the light?

16. ν = c λ

17. Example. A certain violet light has a wavelength of 413 nm. What is the frequency of the light? ν = c λ 2.998 × 10 8 m/s 413 nm

18. Example. A certain violet light has a wavelength of 413 nm. What is the frequency of the light? ν = c λ 2.998 × 10 8 m/s 413 nm WAIT, This won’t work!

19. Example. A certain violet light has a wavelength of 413 nm. What is the frequency of the light? ν = c λ 2.998 × 10 8 m/s 413 nm 1 m =

20. Example. A certain violet light has a wavelength of 413 nm. What is the frequency of the light? ν = c λ 2.998 × 10 8 m/s 413 nm1 m = 1 x 10 9 nm

21. Example. A certain violet light has a wavelength of 413 nm. What is the frequency of the light? ν = c λ 2.998 × 10 8 m/s 4.13 x 10 -7 m 413 nm1 m = 4.13 x 10 -7 m 1 x 10 9 nm

22. Example. A certain violet light has a wavelength of 413 nm. What is the frequency of the light? ν = c λ 2.998 × 10 8 m/s 4.13 × 10 -7 m

23. Unfortunately, thinking of light as waves lead to a problem. It was noticed that if light strikes a metal, then sometimes it could cause electrons to be emitted (leave the atoms entirely – like in a solar panel); called the photoelectric effect. If light was a wave, then all amounts of light energy should cause this to happen, but this was not the case. It always took some minimum amount of energy to get the electrons to be emitted.

25. This lead Max Planck to theorize that light must carry energy in basic minimum amounts that he called quanta. Like a delivery person cannot correctly deliver half a box, the electrons in atoms cannot gain a fraction of a quantum of energy (it has to be in whole numbers).

26. He proposed that this energy was directly proportional to the frequency of the electromagnetic radiation and a constant, now called Planck’s constant. E = h ν E = energy in Joules (J) h = Planck’s constant = 6.626 × 10 -34 Js ν = frequency in Hz or 1/s E hν

27. Example. What is the energy content of one quantum of the light with a wavelength of 413 nm?

28. Note: wavelength is not in the energy equation, but frequency is. So first, you must solve for the frequency.

29. Example. What is the energy content of one quantum of the light with a wavelength of 413 nm? Note: wavelength is not in the energy equation, but frequency is. So first, you must solve for the frequency. As seen in the earlier example, a wavelength of 413 nm gives a ν = 7.26 × 10 14 Hz.

30. Example. What is the energy content of one quantum of the light with a wavelength of 413 nm? ν = 7.26 × 10 14 Hz E = h × ν E = 6.626 × 10 -34 Js × 7.26 × 10 14 1/s

31. Example. What is the energy content of one quantum of the light with a wavelength of 413 nm? ν = 7.26 × 10 14 Hz E = h × ν E = 6.626 × 10 -34 Js × 7.26 × 10 14 1/s

32. Example. What is the energy content of one quantum of the light with a wavelength of 413 nm? ν = 7.26 × 10 14 Hz E = h × ν E = 6.626 × 10 -34 Js × 7.26 × 10 14 1/s E = 4.81 × 10 -19 J

33. In 1905 Einstein used Plancks work to propose that electromagnetic radiation had a dual wave-particle nature. As a particle, electromagnetic radiation carries a quantum of energy of energy, has no mass, and is called a photon.

34. So to get an electron to emit from a metal, it must be struck with a photon having quantum energy big enough, or nothing will happen. Each metal requires a different quantum energy, thus each metal can be identified by the frequency of light needed to emit electron.

35. This idea was expanded upon to develop an idea of where the electrons were in an atom. It was found that low pressure gases could be trapped in a tube and electrified, and would then glow a color particular to the gas inside.

36. Furthermore this light could be passed into a prism, and instead of getting the entire spectrum (rainbow) of colors, only certain wavelengths of light would be seen as small bars of color, called a line-emission spectrum.

37. This would indicated that the electrons in an atom were only absorbing specific amounts of energy from the electricity, causing the electrons to move from their ground state (normal position close to the nucleus) to an excited state (higher energy position further away from the nucleus). The electrons do not stay in the excited state for long and fall back to their ground state, losing the energy equal to what they gained.

38. Quanta Video

39. Niels Bohr used this to develop a model of the atom where the electrons could only be in certain, specific energy level (n) orbits around the nucleus. Just as you cannot go up half a rung on a ladder, the electron could not go up a partial energy level. The electrons gained or lost enough energy to move a whole number amount of energy levels (n) away from or closer to the nucleus, or it did not move.

40. He calculated the amount of energy needed for an electron of hydrogen to move between each energy level (n) (which was not constant) and his calculations agreed with experimental results.

41. The Balmer series of hydrogen spectral lines refer to the four lines seen in the visible light region (the four colored bars). If the electron was excited to energy level (n) 6, 5, 4, or 3 and fell to energy level (n) 2, the resulting energy given off would have a frequency in the visible region of electromagnetic radiation. (One line for dropping from 6 to 2, one for 5 to 2, one for 4 to 2, and one for 3 to 2).

43. However, there are other possibilities. If the electrons drop from n=6, 5, or 4 to n=3, then the energy given off is not big enough to be seen as it is in the infrared region. These three lines in the infrared region are referred to as the Paschen series. If the electrons drop to n=1, then the five lines given off are too high in energy to be seen, as they are in the ultraviolet region. These lines are referred to as the Lyman series.

46. Turn your lab in to the white box, by tomorrow. You need destructible pens for tomorrow! Grab 4 different colors of pencils (or use your own)

47. Model of Atom Review:

48. 1.Thomson’s Plum Pudding Model – the atom is a ball of evenly spread positive stuff with random negative particles (electrons).

49. 2.Rutherford’s Nuclear Model – the atom has a central nucleus containing the positive particles (protons) with the electrons outside.

50. 1.Thomson’s Plum Pudding Model – the atom is a ball of evenly spread positive stuff with random negative particles (electrons). 2.Rutherford’s Nuclear Model – the atom has a central nucleus containing the positive particles (protons) with the electrons outside. 3.Bohr’s Orbital Model – The electrons circle the nucleus in specific energy orbits, like the planets orbit the sun. Unfortunately this only works for atoms with one electron…

51. 4.Quantum Mechanical Model – electrons are found in specific regions around the nucleus, but the exact location of the electrons inside the regions cannot be determined

53. The quantum mechanical model starts with a Principal Quantum Number (n), which is the basic energy level of an electron, and often matches the period number. Possible values (currently) are 1-7.

54. Inside the principal quantum energy level are sublevels that correspond to different cloud shapes. The sublevels are designated as s (sharp), p (principal), d (diffuse), and f (fundamental).

55. Inside the sublevels are orbitals, specific regions with a 90% probability of finding electrons. s –orbitals are spherically shaped clouds around the nucleus p -orbitals are bar-bell shaped clouds with the nucleus between the lobes d and f are much more complex in shape

56. Each sublevel has room for a different amount of electrons, because an orbital can hold two electrons, then each sublevel has a different amount of orbitals

57. s –sublevel can hold 2 electrons, so it has 1 orbital (shape) p –sublevel can hold 6 electrons, so it has 3 orbitals (shapes) d –sublevel can hold 10 electrons, so it has 5 orbitals (shapes) f –sublevel can hold 14 electrons, so it has 7 orbitals (shapes)

58. The s sublevel is simply a sphere centered on the nucleus.

59. The p sublevel has three orbitals. These are often referred to a dumbbell shape.

60. The d sublevel has five orbitals:

61. The f sublevel has seven orbitals

62. To know the maximum amount of electrons that could be in any principal quantum level (and the number of elements that could be represented) use the formula 2n 2 if n=1, then

63. To know the maximum amount of electrons that could be in any principal quantum level (and the number of elements that could be represented) use the formula 2n 2 if n=1, then 2 electrons will fit if n=4,

64. To know the maximum amount of electrons that could be in any principal quantum level (and the number of elements that could be represented) use the formula 2n 2 if n=1, then 2 electrons will fit if n=4, then 32 electrons will fit

65. Color Now

66. Section 4.3 Electron Configurations We don’t want to draw this, do we?

67. Section 4.3 Electron Configurations In order to show on paper where electrons are likely to be located in an atom, orbital filling diagrams and electron configurations are drawn or written. When this is done, three rules must be followed:

68. 1.Aufbau principle – electrons fill lower energy levels first, thus 1 before 2 and s before p, etc. a.orbitals within a sublevel are equal in energy (called degenerate) b.the principal energy levels often overlap, making them seem a little out of order c.boxes are used to represent orbitals

69. Another way of writing the aufbau principle diagram: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p

70. 2.Pauli Exclusion principle – an orbital (box) can hold a maximum of two electrons (arrows) a.for two electrons to fit, they have to have opposite spins b. for one electron in the orbital c. for two electrons in the orbital (opposite spins)

71. 3.Hund’s Rule – when electrons occupy degenerate orbitals, one electron is placed into each orbital with parallel spins before doubling up Ex. _____ _____ _____ NOT _____ _____ _____ 3p 3p

72. 1s 2s 2p 3d 4f 3s3s 4s4s 5s5s 6s6s 3p3p 4p4p 5p5p 4d4d 5d5d Energy and Distance from the Nucleus

73. 2s 2p 3d 4f 3s3s 4s4s 5s5s 6s6s 3p3p 4p4p 5p5p 4d4d 5d5d H

74. 2s 2p 3d 4f 3s3s 4s4s 5s5s 6s6s 3p3p 4p4p 5p5p 4d4d 5d5d He

75. 2p 3d 4f 3s3s 4s4s 5s5s 6s6s 3p3p 4p4p 5p5p 4d4d 5d5d Li

76. 2p 3d 4f 3s3s 4s4s 5s5s 6s6s 3p3p 4p4p 5p5p 4d4d 5d5d Be

77. 3d 4f 3s3s 4s4s 5s5s 6s6s 3p3p 4p4p 5p5p 4d4d 5d5d B

78. 3d 4f 3s3s 4s4s 5s5s 6s6s 3p3p 4p4p 5p5p 4d4d 5d5d C

79. 3d 4f 3s3s 4s4s 5s5s 6s6s 3p3p 4p4p 5p5p 4d4d 5d5d N

80. 3d 4f 3s3s 4s4s 5s5s 6s6s 3p3p 4p4p 5p5p 4d4d 5d5d O

81. 3d 4f 3s3s 4s4s 5s5s 6s6s 3p3p 4p4p 5p5p 4d4d 5d5d F

82. 3d 4f 3s3s 4s4s 5s5s 6s6s 3p3p 4p4p 5p5p 4d4d 5d5d Ne

83. 3d 4f 4s4s 5s5s 6s6s 3p3p 4p4p 5p5p 4d4d 5d5d Na

84. 3d 4f 4s4s 5s5s 6s6s 3p3p 4p4p 5p5p 4d4d 5d5d Mg

85. 3d 4f 4s4s 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d Al

86. 3d 4f 4s4s 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d Si

87. 3d 4f 4s4s 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d P

88. 3d 4f 4s4s 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d S

89. 3d 4f 4s4s 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d Cl

90. 3d 4f 4s4s 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d Ar

91. 3d 4f 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d K

92. 3d 4f 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d Ca

93. 4f 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d Sc

94. 4f 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d Ti

95. 4f 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d V

96. 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d Cr

97. 4f 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d Mn

98. 4f 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d Fe

99. 4f 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d Co

100. 4f 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d Ni

101. 4f 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d Cu

102. 4f 5s5s 6s6s 4p4p 5p5p 4d4d 5d5d Zn

103. 4f 5s5s 6s6s 5p5p 4d4d 5d5d Ga

104. 4f 5s5s 6s6s 5p5p 4d4d 5d5d Ge

105. 4f 5s5s 6s6s 5p5p 4d4d 5d5d As

106. 4f 5s5s 6s6s 5p5p 4d4d 5d5d Se

107. 4f 5s5s 6s6s 5p5p 4d4d 5d5d Br

108. 4f 5s5s 6s6s 5p5p 4d4d 5d5d Kr

109. 4f 6s6s 5p5p 4d4d 5d5d Rb

110. 4f 6s6s 5p5p 5d5d Sr

111. 4f 6s6s 5p5p 5d5d Y

112. 6s6s 5p5p 5d5d Zr

113. 4f 6s6s 5p5p 5d5d Nb

114. 4f 6s6s 5p5p 5d5d Mo

115. 4f 6s6s 5p5p 5d5d Tc

116. 4f 6s6s 5p5p 5d5d Ru

117. 4f 6s6s 5p5p 5d5d Rh

118. 4f 6s6s 5p5p 5d5d Pd

119. 4f 6s6s 5p5p 5d5d Ag

120. 4f 6s6s 5p5p 5d5d Cd

121. 4f 6s6s 5d5d In

122. 4f 6s6s 5d5d Sn

123. 4f 6s6s 5d5d Sb

124. 4f 6s6s 5d5d Te

125. 4f 6s6s 5d5d I

126. 6s6s 5d5d Xe

127. 4f 5d5d Cs

128. 4f 5d5d Ba

129. 5d5d La

130. 5d5d Ce

131. 5d5d Pr

132. 5d5d Nd

133. 5d5d Pm

134. 5d5d Sm

135. 5d5d Eu

136. 5d5d Gd

137. 5d5d Tb

138. 5d5d Dy

139. 5d5d Ho

140. 5d5d Er

141. 5d5d Tm

142. 5d5d Yb

143. Lu

144. Hf

145. Ta

146. W

147. Re

148. Os

149. Ir

150. Pt

151. Au

152. Hg

154. Orbital Notation shows the arrows in the boxes to represent the electrons in an atom. To shorten this process, an electron configuration can be written. It leaves out the information about the number of orbitals in each sublevel, so it will be expect you remember that information.

155. It has the general form n l ° n = principal quantum number (1-7…) l = sublevel letter (s, p, d, or f) ° = number of e - in that orbital (1-14)

156. Ni = 28 e -

157. 1s 2

158. Ni = 28 e - 1s 2 2s 2

159. Ni = 28 e - 1s 2 2s 2 2p 6

160. Ni = 28 e - 1s 2 2s 2 2p 6 3s 2

161. Ni = 28 e - 1s 2 2s 2 2p 6 3s 2 3p 6

162. Ni = 28 e - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2

163. Ni = 28 e - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8

164. Ni = 28 e - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 2 + 2 + 6 + 2 + 6 + 2 + 8 = 28

165. Sn = 50 e - Your Turn to Try

166. Sn = 50 e - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 2

167. If writing out the entire electron configuration is too much, we can use the previous (in the periodic table) noble gas to take the place of part of the electron configuration: Polonium: 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 4 Xenon: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 Polonium:[Xe] 6s 2 4f 14 5d 10 6p 4

168. When the electron configuration is written for an element using the noble gas configuration the electrons written after the noble gas are the ones that appear on the outside of the atom, called valence electrons..

169. When elements bond to form compounds, it is these electrons that are involved. The amount of valence electrons makes a big difference in how the element will bond, so to make it easy to predict, we draw electron dot diagrams. A) In an electron dot diagram, we use the symbol of the element and dots to represent the number of valence electrons. B) Only s and p electrons with the highest quantum number count for dot diagrams, even if there are d and f electrons after the noble gas.

170. Lithium = 1s 2 2s 1 So Li

171. Beryllium = 1s 2 2s 2 So Be

172. Boron = 1s 2 2s 2 2p 1 So B

173. Carbon = 1s 2 2s 2 2p 2 So C

174. Nitrogen = 1s 2 2s 2 2p 3 So N

175. Oxygen = 1s 2 2s 2 2p 4 So O

176. Fluorine = 1s 2 2s 2 2p 5 So F

177. Neon = 1s 2 2s 2 2p 6 So Ne

178. QUIZ TIME!

179. Strontium = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2

180. SoSr

181. Cobalt = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7

182. So Co

183. Bromine = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5

184. SoBr