1. Modern Atomic Theory

2. Light shows us the structure of the atom The Bohr model of the atom
Section 1
Light shows us the structure of the atom
The Bohr model of the atom

3. In the Rutherford model electrons traveled about the nucleus in an orbit.
What did Rutherford discover about the structure of the atom, from his gold foil experiment?

4. There was a problem with the Rutherford model though.
Scientists know that just like an orbiting satellite eventually crashes to ground, an electron will slow down and crash into the nucleus.
Something else must be going on.

5. #1 Electromagnetic Spectrum
To understand the next model we must review what we know about light.
#1 Electromagnetic Spectrum
Light energy travels as a series of waves

6. #1 Electromagnetic Spectrum
To understand the next model we must review what we know about light.
#1 Electromagnetic Spectrum
Waves have a wavelength
(distance from one peak to the next peak)

7. #1 Electromagnetic Spectrum
To understand the next model we must review what we know about light.
#1 Electromagnetic Spectrum
…and a frequency
(number of waves passing per second)

8. #1 Electromagnetic Spectrum
You can think of frequency as wave density.
…more waves passing per second.
…means more energy carried.

9. #1 Electromagnetic Spectrum
Light at the blue end of the spectrum is made of waves with shorter wavelengths (400 nanometers)
…but higher frequency:
(1015 waves per second!)
That’s a lot of energy
(relatively speaking )

10. #1 Electromagnetic Spectrum
While red light is made up of waves with longer wavelengths
(700 nanometers)
…that’s a lower frequency, lower energy wave.

11. #2 Continuous spectrum of White Light
Passing light through a prism separates the colors to form the visible spectrum of light.
White light has all the different colors of light waves

12. #3 Bright line spectrum of Excited Hydrogen Gas
When electricity is forced through hydrogen gas (think “neon” light) only certain colors of light are emitted.
Only four colors
(frequencies) of light
This spectrum
has holes!

13. #4 Line Spectra of Other Elements
hydrogen
mercury
Neon
Each element has its own distinct line spectrum
Like a fingerprint – it can be used to identify the elements in an unknown sample!

14. During an eclipse the sun’s corona gives off a bright line spectrum.

15. Among other elements, what two elements would you likely see spectra of?

16. 1868: A French astronomer spots an unknown element, now known as helium, in the spectrum of the sun during a much-anticipated total eclipse. The event marks the first discovery of an “extraterrestrial” element, as helium had not yet been found on Earth.

17. Astronomers had been eagerly awaiting a total solar eclipse since 1859, when German physicist Gustav Kirchhoff figured out how to use the analysis of light to deduce the chemical composition of the sun and the stars. Scientists wanted to study the bright red flames that appeared to shoot out from the sun, now known to be dense clouds of gas called solar prominences. But until 1868, they thought the sun’s spectrum could only be observed during an eclipse.

18. French astronomer Pierre Jules César Janssen camped out in Guntoor, India, to watch as the moon passed in front of the sun and revealed the solar prominences. Like other sun-gazers that morning, Janssen discovered that the prominences were mostly made of super-hot hydrogen gas. But he also noticed something extra: Using a special prism instrument called a spectroscope, he determined that the line of yellow light everyone had assumed to be sodium didn’t match up to the wavelength of any known element.

19. Janssen wanted to keep studying the mysterious line, and he was so impressed by the brightness of the sun’s emission lines that he felt sure they could be seen without an eclipse, if he could just figure out how to block other wavelengths of visible light. Working feverishly over the next few weeks, Janssen built the first “spectrohelioscope,” a device specifically designed to examine the spectrum of the sun

20. Unbeknownst to Janssen, a second scientist was also working on the same problem 5,000 miles away. English astronomer Joseph Norman Lockyer succeeded in viewing the solar prominences in regular daylight in October In stunning scientific synchronicity, the two scientists’ papers arrived at the French Academy of Sciences on the same day, and today both men are credited with the first sighting of helium

21. At the time, however, Lockyer and Janssen got ridicule rather than accolades for their discovery. Other scientists didn’t believe the astronomers’ account of a new element … until 30 years later, when Scottish chemist William Ramsay discovered a perplexing earthly gas hidden inside a chunk of uranium ore.

22. Ramsay sent the sample to Lockyer for confirmation
Ramsay sent the sample to Lockyer for confirmation. The scientist was thrilled by the element’s “glorious yellow effulgence,” which he described in the Proceedings of the Royal Society of London in Finally vindicated, Janssen and Lockyer were honored by the French government with a gold medal bearing both their faces.

23. We use this idea when we do Flame tests.
Different ions emit unique colors of light when samples are held in the flame of a Bunsen burner
This is used in identifying an unknown salt.

24. Sodium
Lithium
What is the identity of the two unknown salts above?

25. No hydrogen or helium. ? ?
Barium lines and….?

26. How can the bright line spectrum be explained?
Excited electron states Electrons orbit the nucleus in "shells” close to the nucleus: They sit close, called the “Ground” state

27. How can the bright line spectrum be explained?
Excited electron states Electrons orbit the nucleus in "shells” close to the nucleus:
Electrons absorb energy and move to higher shells.
We say their electrons become “excited”

28. How can the bright line spectrum be explained?
Excited = unstable And electrons fall back to ground state Electrons release energy as light Color of light depends on magnitude of the drop Ex: Small drop = red light Larger drop = blue

30. How can the bright line spectrum be explained?
Drops from various shells back to ground state creates a spectrum of different colored lines.
The colors correspond to the difference in energy between the two levels

31. How can the bright line spectrum be explained?
Small drops = lower energy colors like red
bigger drops = higher energy colors like blue and purple
But why not all the colors? Why only certain colors?
We need to explore
QUANTUM THEORY!

32. Click
here
For
video

33. Simulation
How atoms emit light

34. Neon lights produce light through excited electrons
As does the Aurora borealis – the northern lights
These are examples of excited state gases!

35. 1) How do we know that electrons orbit in shells?
If Rutherford was right electrons can exist in any orbit:
All colors would be emitted
If Bohr is right electrons can only sit on energy levels:
Only specific colors emitted

36. “Quantum” Theory: Max Planck (1900)
2) Why must electrons orbit in shells?
“Quantum” Theory: Max Planck (1900)
Planet model
Electrons can have any orbit
Shell model
Electrons have specific orbits

37. “Quantum” Theory: Max Planck (1900)
Energy is “quantized” - Light energy IS absorbed and emitted as discrete packets of energy called “photons”
Energy of electrons is not continuous, it is gained and lost in discrete units called quanta (sing. Quantum)

38. The Bohr Model of the atom
Niels Bohr model of the atom (1913)
Electron is QUANTIZED
Electrons can ONLY orbit in certain specific Energy Levels (shells) depending on how much energy they contain

39. The Bohr Model of the atom
Higher energy levels =
electrons with higher energy
Lower energy levels = electrons with less energy
Normally electrons in an atom sit in the ground state but can become excited, jumping between levels and create their unique bright line spectra.

40. Learning Check
What is the source of the bright line spectrum of an element?
How does the ground state and excited state of an atom compare in terms of energy?
What determines the color of the light in an individual bright line?
In what way is an atomic spectrum (bright line spectrum) like a fingerprint? How is it useful?

41. Review:
Must compare unknown to a known “standard”

42. Singular “spectrum”

43. Section 2A Electron configurations and Bohr diagrams
Lewis electron dot diagrams
Excited electron state

44. Bohr Diagrams May show nucleus particles
Shows energy levels and electrons – paired or unpaired in orbitals
Ex: Al electron configuration 2-8-3
2 electrons in 1st level
8 electrons in 2nd level
3 electrons in 3rd level

45. Bohr Diagram for oxygen
Shows energy levels and electrons in orbitals
Total electrons = 8
First 2 electrons go in as orbital pair in 1st Energy level
Next two electrons 2nd level as a pair
The next three enter separate orbitals
The remaining 1 electron combines with one of the Three unpaired electrons
Orbital pair
8P+
Electron configuration:
2-6

46. 2(1) 2(1+3) 2(1+3+5) Arrangement of electrons in the atom
= Electron configuration
Rules: Each energy level can hold a certain maximum number of electrons
1st holds 2 e-
2nd holds 8 e-
3rd holds
[ 2n2 ]
Ex: Carbon: 2 – 4
2 electrons in 1st energy level
4 electrons in 2nd energy level
2(1)
2(1+3)
2(1+3+5)

47. F 9
18.998
2 - 7
30.973 P 15
Electron configurations on the periodic table
How many electrons occupy the 2nd energy level
of a phosphorus atom in the ground state?

48. Electrons are divided between Kernel and valence electrons
Aluminum Kernel = [Ne]:
+ 3 valence electrons
3 valence electrons
2 – 8 - 3
Chlorine
7 valence electrons
Properties are connected to valence electrons:
Elements in the same group have same # valence electrons

49. Try these: Draw Bohr diagrams for the following atoms:
Be F
2-2
2-7

50. Try these: Draw Bohr diagrams for the following atoms:
Use your periodic table to look up the electron configurations.
3. Na P

51. An easier way?: Lewis (Electron Dot) Structures
For Elements and Ions – dots or x’s used to represent valence electrons
Atoms
hydrogen: H Magnesium: Mg
e- configuration:
Oxygen: O Chlorine: Cl

52. Draw the Lewis diagrams for Aluminum and Nitrogen
e- configuration from the periodic table:
Al N

53. How does the number of valence (outer) electrons change within groups (vertically) on the table?
Within periods (horizontal)?

54. Learning check:
Write the electron configuration for an atom of Ca in the ground state:
How many kernel and valence electrons does calcium have?
Which noble gas has a kernel like calcium?
Draw the Bohr diagram and Lewis electron dot symbols for Lithium, sulfur and fluorine

56. Review

58. Excited state electron configurations
Atom in ground state – electrons in lowest E levels = lowest energy state
ex: Li C F
Become
Excited:

59. 1 As an electron in an atom moves from the ground state to the excited state, the electron
(1) gains energy as it moves to a higher energy level
(2) gains energy as it moves to a lower energy level
(3) loses energy as it moves to a higher energy level
(4) loses energy as it moves to a lower energy level
44 Which is an electron configuration for an atom of chlorine in the excited state? HDYK?
(1) 2–8–7 (3) 2–8–6–1
(2) 2–8–8 (4) 2–8–7–1

60. The periodic table revisited
Section 2B
The periodic table revisited

61. Groups with same # valence electrons
The Periodic table is organized by electron arrangement
Groups with same # valence electrons 8e 1e
1 energy
level 2e 6e 7e 3e 4e 5e 1 2
Periods = # energy levels
2 energy
levels
2-2
2-3
2-4
2-5
2-8
2-1
2-6
2-7
3 energy
levels
2-8-3
2-8-4
2-8-1
2-8-2
2-8-5
2-8-6
2-8-7
2-8-8
Etc.
Etc.

62. Why can’t the outer shell hold more than 8 electrons?K
(its hard to pack more electrons into
an already crowded space)

63. Learning Check 1. A neutral Sulfur atom is in the ground state
Without looking directly at the electron configuration determine: HDYK?
a. the number of valence electrons
b. the number of occupied energy levels
c. the electron configuration
2. A neutral atom has the ground state electron configuration of
a. What group and period is it located in? HDYK?
b. What properties will it have?

64. The wave mechanical model of the atom Periodic table revisited…again.
Section 3
The wave mechanical model of the atom
Periodic table revisited…again.

65. Modern view of the atom Wave Mechanical model
Electrons as standing waves

66. Only standing waves are allowed

67. Arrangement of Electrons in Atoms
Electrons in atoms are arranged as
PRINCIPLE ENERGY LEVELS (shell):
made up of SUBLEVELS (of different shapes):
where electron move in ORBITALS (the region of space):
Orbital = Region where electron
is most likely to be found
Region of “highest probability”

68. The first shell has only the S sublevel
S orbitals are shaped like a sphere ….and can hold only 2 electrons

69. The first shell has only the S sublevel
The first shell can only hold two electrons

70. The second shell has an S orbital
But also has 3 P sublevel orbitals
P orbitals are shaped like “dumbells
….but can hold only 2 electrons each

71. The second shell has an S orbital
But also has 3 P sublevel orbitals
The second shell can hold 2 electrons in its S orbital
PLUS…6 electrons in its three P orbitals

72. The second shell has an S orbital
But also has 3 P sublevel orbitals
The second shell can hold a total of
2 + 6 = 8 electrons!

73. The third shell can have three types of orbitals
(three sublevels) and so the pattern goes…
It gets a little crazy.

74. Just remember : electrons aren’t particles, they’re waves
Just remember : electrons aren’t particles, they’re waves. So they don’t have to behave like little orbiting planets!

75. We never know exactly where they’ll be, just the region in space
(called the orbital) where they’ll most likely spend their time.

76. Test your learning
The modern model of the atom is sometimes called the wave mechanical model. What does the name infer?
How is an orbital different from an orbit?
On modern periodic tables, the alkali metals and alkaline earth metals (groups 1 and 2) are in the “S” block, while groups 13 through 18 are in the “P” block. Write an hypothesis for why this might be so.

77. 2 In the wave-mechanical model, an orbital is a region of space in an atom where there is
(1) a high probability of finding an electron
(2) a high probability of finding a neutron
(3) a circular path in which electrons are found
(4) a circular path in which neutrons are found

79. Periodic Table II: Electron arrangement and the table

80. So why the weird shape? s s p d f “Diagonal Rule”
The sublevels don’t fill in order!

81. Shape of the table follows Filling of sublevels
3p p 6
4s1 s2

82. 1s1
1s2
2s1
2s2
2p1 p2 p3 p4 p5 p6
3s1 s2
3p p 6
4s1 s2
3d1----electrons in lower E levels d10
4p1—back outside----P6
5s1----2
4d etc d10 5p
6s1----2
5d d10
6p etc.
7s1----2
6d d10
etc.
4f electrons going two E levels inside f14
5f etc f14

83. Why are d and f orbitals always in lower energy levels?
d and f orbitals require LARGE amounts of energy due to electron repulsion
(hard to pack more electrons into
an already crowded space)
It’s better (lower in energy) to skip a sublevel that requires a large amount of energy for higher level but lower energy K

84. H He Ne Ar Hg

86. H He Hg

88. calcium
strontium
Lithium
barium
sodium
potassium
copper