1. Chapter 6 Electronic Structure of Atoms
CHEMISTRY The Central Science 9th Edition
Chapter 6 Electronic Structure of Atoms
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Chapter 6

2. 6.1: The Wave Nature of Light
Electromagnetic Radiation carries energy through space
Speed of EMR through a vacuum: 3.00x108 m/s
wavelike characteristics:
wavelength, l (distance from crest to crest)
amplitude, A (height)
frequency, n (# of cycles which pass a point in 1 second)
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3. The # of times the cork bobs up and down is the frequency
The wave nature is due to periodic oscillations of the intensities of electronic and magnetic forces associated with the radiation
The # of times the cork bobs up and down is the frequency
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4. Frequency and wavelength are inversely related
Text, P. 200
Frequency and wavelength are inversely related

5. For light, speed = c (that is 3.00 x 108m/s)
The speed of a wave, c, is given by its frequency multiplied by its wavelength:
For light, speed = c (that is 3.00 x 108m/s)
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6. Example: the visible spectrum
Modern atomic theory arose out of studies of the interaction of radiation with matter
Example: the visible spectrum
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7. The Electromagnetic Spectrum, P. 201
(or Hz)

8. Text, P. 201
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9. Sample Problem #7
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10. 6.2: Quantized Energy and Photons
Problems with wave theory:
Emission of light from hot objects
“black body radiation”: stove burner, light bulb filament
The photoelectric effect
Emission spectra
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11. where h is Planck’s constant (6.626  10-34 J-s)
Planck: energy can only be absorbed or released from atoms in certain amounts called quanta
The relationship between energy and frequency is
where h is Planck’s constant (6.626  J-s)
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12. Quantization Analogy:
Consider walking up a ramp versus walking up stairs:
For the ramp, there is a continuous change in height
Movement up stairs is a quantized change in height
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13. The Photoelectric Effect
Evidence for the particle nature of light -- “quantization”
Light shines on the surface of a metal: electrons are ejected from the metal
threshold frequency must be reached
Below this, no electrons are ejected
Above this, the # of electrons ejected depends on the intensity of the light
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14. Text, P. 204

15. Einstein assumed that light traveled in energy packets called photons
The energy of one photon:
Energy and frequency are directly proportional
Therefore radiant
energy must be quantized!
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16. Flame Tests: Pretty colors!
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17. Sample Problems # 13, 17, 19
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18. 6.3: Line Spectra and the Bohr Model
Radiation that spans an array of different wavelengths: continuous
White light through a prism: continuous spectrum of colors
Spectra tube emits light unique to the element in it
Looking at it through a prism, only lines of a few wavelengths are seen
Black regions correspond to wavelengths that are absent
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19. Continuous Visible Spectrum, P. 206

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21. Therefore the atom should be unstable
Bohr Model
Rutherford model of the atom: electrons orbited the nucleus like planets around the sun
Physics: a charged particle moving in a circular path should lose energy
Therefore the atom should be unstable
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22. 3 Postulates for Bohr’s theory:
Electrons move in orbits that have defined energies
An electron in an orbit has a specific energy
Energy is only emitted or absorbed by an electron as it changes from one allowed energy state to another (E=hν)
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23. Since the energy states are quantized,
the light emitted from excited atoms must be quantized and appear as line spectra
Bohr showed that
where n is the principal quantum number (i.e., n = 1, 2, 3, … and nothing else)
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24. The furthest orbit in the Bohr model has n close to infinity
The first orbit in the Bohr model has n = 1, is closest to the nucleus, and has the lowest energy (ground state)
The furthest orbit in the Bohr model has n close to infinity
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25. Electrons in the Bohr model can only move between orbits by absorbing and emitting energy in quanta (hν)
The amount of energy absorbed or emitted on movement between states is given by
Sample Problems # 21 and 27
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26. nf values for other regions of the EMS: Lyman (UV): nf = 1
Line Spectra, P. 207
The Balmer series for Hydrogen (nf = 2, which is in the visible region)
nf values for other regions of the EMS:
Lyman (UV): nf = 1
Paschen (IR): nf = 3
Brackett (IR): nf = 4
Pfund (IR): nf = 5
The Rydberg Equation (P. 206) allows for the calculation of wavelengths for all the spectral lines
n = energy level
Note the units
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27. Balmer Series Energy Level Diagram
Lyman Series
Paschen Series
Balmer Series Energy Level Diagram
Figure 6.14, P. 208
Using the Rydberg Equation, the existence of spectral lines can be attributed to the quantized jumps of electrons between energy levels (therefore justifying Bohr’s model of the atom)
Sample Problems # 29 and 31

28. The Electromagnetic Spectrum, P. 201
(or Hz)
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29. Limitations of the Bohr Model
It can only explain the line spectrum of hydrogen adequately
Electrons are not completely described as small particles
It doesn’t account for the wave properties of electrons
DEMOS!
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30. 6.4:The Wave Behavior of Matter
Using Einstein’s and Planck’s equations, de Broglie showed:
The momentum, mv, is a particle property, but  is a wave property
Sample Exercise 6.5, P. 210
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31. The Uncertainty Principle
Heisenberg’s Uncertainty Principle: we cannot determine exactly the position, direction of motion, and speed of an electron simultaneously
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32. 6.5: Quantum Mechanics and Atomic Orbitals
Erwin Schrödinger proposed an equation that contains both wave and particle terms
Solving the equation leads to wave functions (shape of the electronic orbital)
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33. Text, P. 213
The electron density distribution in the ground state of the hydrogen atom
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34. Orbitals and Quantum Numbers
If we solve the Schrödinger equation, we get wave functions and energies for the wave functions
We call wave functions orbitals (regions of highly probable electron location)
Schrödinger’s equation requires 3 quantum numbers:
Principal Quantum Number, n
This is the same as Bohr’s n
As n becomes larger, the atom becomes larger and the electron is further from the nucleus
Orbitals and Quantum Numbers
If we solve the Schrödinger equation, we get wave functions and energies for the wave functions.
We call wave functions orbitals (regions of highly probable electron location).
Schrödinger’s equation requires 3 quantum numbers:
Principal Quantum Number, n
This is the same as Bohr’s n.
As n becomes larger, the atom becomes larger and the electron is further from the nucleus.
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35. 2. Azimuthal Quantum Number, l
This quantum number depends on the value of n
The values of l begin at 0 and increase to (n - 1)
We usually use letters for l (s, p, d and f for l = 0, 1, 2, and 3)
Usually we refer to the s, p, d and f-orbitals (the shape of the orbital) (AKA “subsidiary quantum number”)
3. Magnetic Quantum Number, ml
This quantum number depends on l
The magnetic quantum number has integral values between -l and +l
Magnetic quantum numbers give the 3D orientation of each orbital
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36. “Aufbau Diagram”: electrons fill low energy orbitals first
As n increases, note that the spacing between energy levels becomes smaller
“Aufbau Diagram”: electrons fill low energy orbitals first
Figure 6.17, P. 215
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37. Sample Problems # 41, 43, & 45
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38. 6.6: Representations of Orbitals
The s-Orbitals
All s-orbitals are spherical
As n increases, the s-orbitals get larger
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39. the s orbitals
Text, P. 216
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40. There are three p-orbitals px, py, and pz
The p-Orbitals
There are three p-orbitals px, py, and pz
The three p-orbitals lie along the x-, y- and z- axes of a Cartesian system
The letters correspond to allowed values of ml of -1, 0, and +1
The orbitals are dumbbell shaped
As n increases, the p-orbitals get larger
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41. the p orbitals
Text, P. 216
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42. There are five d and seven f-orbitals FYI for the d-orbitals:
The d and f-Orbitals
There are five d and seven f-orbitals
FYI for the d-orbitals:
Three of the d-orbitals lie in a plane bisecting the x-, y- and z-axes
Two of the d-orbitals lie in a plane aligned along the x-, y- and z-axes
Four of the d-orbitals have four lobes each
One d-orbital has two lobes and a collar
Text, P. 217
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43. the d orbitals
Text, P. 217

44. the f orbitals (not in text)
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45. http://www.uky.edu/~holler/html/orbitals_2.html Prentice Hall © 2003
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46. Text, P. 214 (-l to +l) (n2) Max = (n-1) l  letter
(# of orientations)
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47. 6.7: Many-Electron Atoms Orbitals and Their Energies
Orbitals of the same energy are said to be degenerate
For n  2, the s- and p-orbitals are no longer degenerate because the electrons interact with each other
Therefore, the Aufbau diagram looks slightly different for many-electron systems
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48. BOHR
Text, P. 215
MODERN
Text, P. 218

49. Electron Spin and the Pauli Exclusion Principle
Since electron spin is also quantized, we define
ms = spin quantum # =  ½
Text, P. 219

50. Pauli’s Exclusion Principle: no two electrons can have the same set of 4 quantum numbers
Therefore, two electrons in the same orbital must have opposite spins
Electron capacity of sublevel = 4l + 2
Electron capacity of energy level = 2n2
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51. Sample Problems # 51 & 55
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52. 6.8: Electron Configurations
Hund’s Rule
Electron configurations tell us in which orbitals the electrons for an element are located
Three rules are applied:
Aufbau Principle
Pauli’s Exclusion Principle
Hund’s Rule: for degenerate orbitals, electrons fill each orbital singly before any orbital gets a second electron
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53. Aufbau Diagram
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54. Types of electron configurations
Electron configuration notation: energy level, subshell, # of electrons per orbital
Orbital notation: each ml value is represented by a line, electrons are also shown
Noble gas configuration: “condensed configuration”
[Proceeding noble gas] electron configuration notation for outer shell electrons
Inner shell e-
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55. 6.9: Electron Configurations and the Periodic Table
The periodic table can be used as a guide for electron configurations
The period number is the value of n
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56. # of columns is related to the number of electrons that can fit in the subshells
Regular trends
Text, P. 225

57. Irregularities: half filled and completely filled subshells are stable
Text, P. 227

58. Sample Problems # 59, 61, 63 & 65
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59. End of Chapter 6
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