1. Chapter 11a
Modern Atomic Theory

2. 11.1 Rutherford’s Atom
11.2 Electromagnetic Radiation
11.3 Emission of Energy by Atoms
11.4 The Energy Levels of Hydrogen
11.5 The Bohr Model of the Atom

3. Nuclear Model of the Atom
The atom has a small dense nucleus which
is positively charged.
contains protons (+1 charge).
contains neutrons (no charge).
The remainder of the atom
is mostly empty space.
contains electrons (–1 charge). 

4. What are the electrons doing?
The nuclear charge (n+) is balanced by the presence of n electrons moving in some way around the nucleus.
What are the electrons doing?
How are the electrons arranged and how do they move?
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5. Wavelength ( ) – distance between two peaks or troughs in a wave.
Characteristics
Wavelength ( ) – distance between two peaks or troughs in a wave. 

6. Different Wavelengths Carry Different Amounts of Energy

7. One of the ways that energy travels through space.

8. Speed (c) – speed of light (2.9979×108 m/s)
Characteristics
Frequency ( ) – number of waves (cycles) per second that pass a given point in space
Speed (c) – speed of light (2.9979×108 m/s)
186,000 miles/s

9. Photon – packet of energy
Dual Nature of Light
Wave
Photon – packet of energy

10. Characteristics
Energy of a photon of light = Planck’s constant (h) (6.626 x 10-34Js) times the speed of light (c) (2.9979×108 m/s) divided by the wavelength in meters (λ) or the wavelength in nanometers (nm) times ten to the -9 power (550nm = 550 x 10-9m).
OR:
Ephoton = (hc) / λ

11. Let’s Practice! E=hc/λ =6.626 x 10-34Js(3.00 x 108m/s)/(535 x 10-9m)
What is the energy of light with a wavelength of 535 nm?
E=hc/λ
=6.626 x 10-34Js(3.00 x 108m/s)/(535 x 10-9m)
=3.70 x J

12. Seeing the Light-A New Model of the Atom
Maxwell Planck-Black Body Radiation
1900—Nobel Prize in 1918
Found that blackbody radiation was quantized.

13. Quantized Energy Levels
The energy levels of all atoms are quantized.

14. Einstein’s Photoelectric Effect (1905--Nobel Prize in 1921)
Only light from a certain color (energy) could eject electrons. Intensity of the light had no effect. Energy is absorbed only at quantized energies!

15. Atoms can give off light.
They first must receive energy and become excited.
The energy is released in the form of a photon.
The energy of the photon corresponds exactly to the energy change experienced by the emitting atom.
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16. Excited state – atom with excess energy
Atomic states
Excited state – atom with excess energy
Ground state – atom in the lowest possible state
When an H atom absorbs energy from an outside source it enters an excited state.
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17. Energy Level Diagram
Energy in the photon corresponds to the energy used by the atom to get to the excited state.
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18. Stokes Shift-Absorb high energy (UV) and emit low energy (visible).

19. Only certain types of photons are produced when H atoms release energy
Only certain types of photons are produced when H atoms release energy. Why?

20. Line Spectra

21. The word laser comes from light amplification by stimulated emission of radiation.

22. Lasers
Lasers used to remove blood clots. Laser light transmitted in fiber optics. Cataract Removal Light Shows

23. Holograms
3D pictures made by Lasers using the interference pattern between reflected laser light from the surface of an object and the undisturbed laser light reflected from a mirror. The Interference pattern is recorded on film. The developed film can then be used by a laser to recreate the image in 3D.

24. Holograms
Holograms are made from laser light without using an image forming device. Tupac holographic concert and a holographic fashion display.

25. The Doppler Effect
The doppler effect is the apparent change in frequency of a wave due to the relative motion of the listener and the source of the sound.
The doppler effect also occurs in light waves and is used by astronomers to calculate the speed at which stars are approaching or receding.

26. Bohr Model
Line Spectra in Stars and the red shift indicating movement away or towards us.
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27. Quantized Energy Levels
Since only certain energy changes occur the H atom must contain discrete energy levels.
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28. Concept Check
Why is it significant that the color emitted from the hydrogen emission spectrum is not white?
How does the emission spectrum support the idea of quantized energy levels?
If the levels were not quantized, we’d probably see white light. This is because all possible value of energy could be released, meaning all possible colors would be emitted. All the colors combined make white light. Since only certain colors are observed, this means that only certain energy levels are allowed. An electron can exist at one level or another, and there are regions of zero probability in between.

29. When an electron is excited in an atom or ion
Concept Check
When an electron is excited in an atom or ion
a) only specific quantities of energy are released in order for the electron to return to its ground state.
b) white light is never observed when the electron returns to its ground state.
c) the electron is only excited to certain energy levels.
d) All of the above statements are true when an electron is excited.
The correct answer is d.
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30. Niels Bohr 1913—Nobel Prize in 1922
Niels Bohr hypothesized that electrons orbit the nucleus just as the planets orbit the sun (planetary model).
Niels Bohr 1913—Nobel Prize in 1922
Quantized energy levels
Electron moves in a circular orbit.
Electron jumps between levels by absorbing or emitting a photon of a particular wavelength.
Actually electrons do not move in a circular orbit.

31. D
Chapter 11b
Modern Atomic Theory

32. Modern Model of the Atom
Chapter 11
Modern Model of the Atom
11.6 The Wave Mechanical Model of the Atom
11.7 The Hydrogen Orbitals
11.8 The Wave Mechanical Model: Further Development
11.9 Electron Arrangements in the First Eighteen Atoms on the Periodic Table
11.10 Electron Configurations and the Periodic Table
11.11 Atomic Properties and the Periodic Table

33. Orbitals
Nothing like orbits
Probability of finding the electron within a certain space
This model gives no information about when the electron occupies a certain point in space or how it moves.

34. Orbitals
Orbitals do not have sharp boundaries and are represented by probability distributions or where the electron is likely to be found without regards to movement of the electrons.
Chemists arbitrarily define an orbital’s size as the sphere that contains 90% of the total electron probability.
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35. Scanning Tunneling Microscope

36. Louis DeBroglie
1924 – Nobel Prize in 1929
He found that matter (electrons) moved in waves. Just as light behaved like particles and waves, so did matter.
An 18-wheeler moving down Hwy 99 at 60mph has a wavelength smaller than an atom.
However, an electron (very light) moves much faster and its wavelength is much larger than its size.

37. Erwin Schrödinger 1926 -Nobel Prize in 1933
Found the probability of finding an electron in an atom.

38. Erwin Schrödinger 1926 -Nobel Prize in 1933
Found the probability of finding an electron in an atom.

39. Hydrogen Energy Levels
Hydrogen has discrete energy levels.
Called principal energy levels
Labeled with whole numbers

40. Hydrogen Energy Levels
Each principal energy level is divided into sublevels.
Labeled with numbers and letters
Indicate the shape of the orbital
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41. Hydrogen Energy Levels
The s and p types of sublevel
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42. d Orbitals
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43. f-orbtals

44. Why the different shapes?
3py 3d
2py 1s 2s 3s
2px
3px
2pz
3pz

45. The number tells the principal energy level.
Orbital Labels
The number tells the principal energy level.
The letter tells the shape.
The letter s means a spherical orbital or shape of the probability distribution of the electron.
The letter p means the orientation. The x, y, or z subscript on a p orbital label tells along which of the coordinate axes the two lobes lie.

46. Why does an H atom have so many orbitals and only 1 electron?
Hydrogen Orbitals
Why does an H atom have so many orbitals and only 1 electron?
An orbital is a potential space for an electron.
Atoms can have many potential orbitals.

47. Atoms Beyond Hydrogen
The Bohr model was discarded because it does not apply to all atoms. It did not consider the different energy sublevels or suborbitals within each orbital.
Atoms beyond hydrogen have multiple electrons that distorts the energy levels due to electron-electron interactions.
Need one more property to determine how the electrons are arranged: Spin – electrons spin like a top causing a magnetic field. Opposite magnetic fields can attract allowing electrons to occur in pairs if their spin or magnetic field is opposite.

48. Atoms Beyond Hydrogen
Pauli Exclusion Principle – an atomic orbital can hold a maximum of 2 electrons and those 2 electrons must have opposite spins.

49. Principal Components of the Wave Mechanical Model of the Atom
Atoms have a series of energy levels called principal energy levels (n = 1, 2, 3, etc.).
The energy of the level increases as the value of n increases.
Each principal energy level contains one or more types of orbitals, called sublevels.
The number of sublevels present in a given principal energy level equals n.
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50. Principal Components of the Wave Mechanical Model of the Atom
5. The n value is always used to label the orbitals of a given principal level and is followed by a letter that indicates the type (shape) of the orbital (1s, 3p, etc.).
6. An orbital can be empty or it can contain one or two electrons, but never more than two. If two electrons occupy the same orbital, they must have opposite spins.
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51. Principal Components of the Wave Mechanical Model of the Atom
7. The shape of an orbital does not indicate the details of electron movement. It indicates the probability distribution for an electron residing in that orbital.

52. Concept Check
Which of the following statements best describes the movement of electrons in a p orbital?
a) The electron movement cannot be exactly determined.
b) The electrons move within the two lobes of the p orbital, but never beyond the outside surface of the orbital.
c) The electrons are concentrated at the center (node) of the two lobes.
d) The electrons move along the outer surface of the p orbital, similar to a “figure 8” type of movement.
The correct answer is a.
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53. Energy Level Diagram for Carbon

54. Electron configuration – electron arrangement 1s1
H Atom
Electron configuration – electron arrangement 1s1
Orbital diagram – orbital is a box grouped by sublevel containing arrow(s) to represent electrons
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55. Electron configuration 1s2 2s1 Orbital diagram
Li Atom
Electron configuration
1s2 2s1
Orbital diagram
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56. O Atom
The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons in a particular set of degenerate (same energy) orbitals.
Oxygen: 1s 2s p
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57. The electron configurations in the sublevel last occupied for the first eighteen elements.
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58. Classifying Electrons
Core electrons – inner electrons
Valence electrons – electrons in the outermost (highest) principal energy level of an atom
1s22s22p6 (valence electrons = 8)
The elements in the same group on the periodic table have the same valence electron configuration.
Elements with the same valence electron arrangement show very similar chemical behavior.

59. 3d suborbitals Concept Check
How many unpaired electrons does the element cobalt (Co) have in its lowest energy state?
a) 0
b) 2
c) 3
d) 7
3d suborbitals
The correct answer is c. There are 7 electrons in the d orbitals, with two of the d orbitals containing paired electrons and three of the d orbitals containing unpaired electrons.
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60. a) Yes. An electron can be excited into a 3d orbital.
Concept Check
Can an electron in a phosphorus atom ever be in a 3d orbital? Choose the best answer.
a) Yes. An electron can be excited into a 3d orbital.
b) Yes. A ground-state electron in phosphorus is located in a 3d orbital.
c) No. Only transition metal atoms can have electrons located in the d orbitals.
d) No. This would not correspond to phosphorus’ electron arrangement in its ground state.
The correct answer is a.
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61. Quantum #’s are like an Address.
What do you need to know to find out where you live?
State
City
Street
House
Magnetic Quantum # (ml)
Spin Quantum # (ms)
Principle Quantum # (n)
Angular Quantum # (l)

62. Look at electron configurations for K through Kr.
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63. Orbital Filling and the Periodic Table
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64. Orbital Filling
In a principal energy level that has d orbitals, the s orbital from the next level fills before the d orbitals in the current level.
After lanthanum, which has the electron configuration [Xe]6s25d1, a group of fourteen elements called the lanthanide series, or the lanthanides, occurs. This series of elements corresponds to the filling of the seven 4f orbitals.
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65. Orbital Filling
3. After actinum, which has the configuration [Rn]7s26d1,a group of fourteen elements called the actinide series, or actinides, occurs. This series corresponds to the filling of the seven 5f orbitals.

66. Orbital Filling
4. Except for helium, the group numbers indicate the sum of electrons in the ns and np orbitals in the highest principal energy level that contains electrons (where n is the number that indicates a particular principal energy level). These electrons are the valence electrons.

67. Exercise
Determine the expected electron configurations for each of the following.
a) S
1s22s22p63s23p4 or [Ne]3s23p4
b) Ba
[Xe]6s2
c) Eu
[Xe]6s24f7
a) 16 electrons total; 1s22s22p63s23p4 or [Ne]3s23p4
b) 56 electrons total; [Xe]6s2
c) 63 electrons total; [Xe]6s24f7
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68. Write electron configurations for the following: Al Sc K Br Zn Hg
1s22s22p63s23p1
1s22s22p63s23p64s23d1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p5
1s22s22p63s23p64s23d10
1s22s22p63s23p64s23d104p65s24d105p66s2 4f145d10
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69. Write the abbreviated electron configuration for the following:
Magnesium
Carbon
Boron
Chlorine
Selenium
– [Ne] 3s2
– [He] 2s22p2
– [He] 2s22p1
– [Ne] 3s23p5
– [Ar] 4s23d104p4
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70. [Kr] isoelectronic with Kr [Ne] isoelectronic with Ne
Write the electron configuration in long and abbreviated notation for the following ions.
Br-
N3- K+
Sr2+
S2-
Ni2+
[Kr] isoelectronic with Kr
[Ne] isoelectronic with Ne
[Ar] isoelectronic with Ar
[Kr] isoelectronic with Kr
[Ar] isoelectronic with Ar
[Ar]4s23d6 isoelectronic with Fe
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72. Metals tend to lose electrons to form positive ions.
Metals and Nonmetals
Metals tend to lose electrons to form positive ions.
Nonmetals tend to gain electrons to form negative ions.
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73. Energy required to remove an electron from a gaseous atom or ion.
Ionization Energy
Energy required to remove an electron from a gaseous atom or ion.
X(g) → X+(g) + e–
Mg → Mg+ + e– I1 = 735 kJ/mol (1st IE)
Mg+ → Mg2+ + e– I2 = 1445 kJ/mol (2nd IE)
Mg2+ → Mg3+ + e– I3 = 7730 kJ/mol *(3rd IE)
*Core electrons are bound much more tightly than valence electrons.
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74. Ionization Energy
In general, as we go across a period from left to right, the first ionization energy increases.
Why?
Electrons added in the same principal quantum level do not completely shield the increasing nuclear charge caused by the added protons.
Electrons in the same principal quantum level are generally more strongly bound from left to right on the periodic table.
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75. In general, as we go across a period the ionization energy increases.
As we go up a group from top to bottom, the first ionization energy increases.

76. Which atom would require more energy to remove an electron? Why? Na Cl
Concept Check
Which atom would require more energy to remove an electron? Why?
Na Cl
Cl would require more energy to remove an electron because the electron is more tightly bound due to the increase in effective nuclear charge.
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77. Which atom would require more energy to remove an electron? Why? Li Cs
Concept Check
Which atom would require more energy to remove an electron? Why?
Li Cs
Li would require more energy to remove an electron because the outer electron is on average closer to the nucleus (so more tightly bound).
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78. In general atomic radius increases in going down a group.
Atomic Size
In general as we go across a period from left to right, the atomic radius decreases.
Effective nuclear charge increases, therefore the valence electrons are drawn closer to the nucleus, decreasing the size of the atom.
In general atomic radius increases in going down a group.
Orbital sizes increase in successive principal quantum levels.
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79. Relative Atomic Sizes for Selected Atoms (Fig. 11-36)
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80. Which should be the larger atom? Why? Na Cl
Concept Check
Which should be the larger atom? Why?
Na Cl
Na should be the larger atom because the electrons are not bound as tightly due to a smaller effective nuclear charge.

81. Which should be the larger atom? Why? Li Cs
Concept Check
Which should be the larger atom? Why?
Li Cs
Cs should be the larger atom because of the increase in orbital sizes in successive principal quantum levels (to accommodate more electrons).
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82. Which is lower in energy?
Concept Check
Which is larger?
The hydrogen 1s orbital
The lithium 1s orbital
Which is lower in energy?
The hydrogen 1s orbital is larger because the electrons are not as tightly bound as the lithium 1s orbital (lithium has a higher effective nuclear charge and will thus draw in the inner electrons more closely).
The lithium 1s orbital is lower in energy because the electrons are closer to the nucleus.
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83. Exercise
Arrange the elements oxygen, fluorine, and sulfur according to increasing:
Ionization energy
S, O, F
Atomic size
F, O, S
Ionization Energy: S, O, F (IE increases as you move up a column and to the right across a period.)
Atomic Size: F, O, S (Atomic radius increases as you move to the left across a period and down a column.)
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