1. Modern Atomic Theory
Chapter 11

2. The ELECTRON: Wave – Particle Duality
“No familiar conceptions can be woven around the electron. Something unknown is doing we don’t know what.”
-Sir Arthur Eddington
The Nature of the Physical World (1934)

3. The Dilemma of the Atom
Electrons outside the nucleus are attracted to the protons in the nucleus
Charged particles moving in curved paths lose energy
What keeps the atom from collapsing?

4. Wave-Particle Duality
JJ Thomson won the Nobel prize for describing the electron as a particle.
His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.
The electron is a particle!
The electron is an energy wave!

6. Electromagnetic Radiation
Most subatomic particles behave as PARTICLES (photons) and obey the physics of waves.

7. The Wave-like Electron
The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves.
Louis deBroglie

8. Electromagnetic radiation propagates through space as a wave moving at the speed of light.
c = speed of light, a constant (3.00 x 108 m/s)
 = frequency, in units of hertz (hz, sec-1)
 = wavelength, in meters

9. The energy (E ) of electromagnetic radiation is directly proportional to the frequency () of the radiation.
E = h
E = Energy, in units of Joules (kg·m2/s2)
h = Planck’s constant (6.626 x J·s)
 = frequency, in units of hertz (hz, sec-1)
Often times in calculations, you will use the following formula (remember c=vλ):
E = hc λ

10. Long
Wavelength =
Low Frequency
Low ENERGY
Short
Wavelength =
High Frequency
High ENERGY

11. Answering the Dilemma of the Atom
Treat electrons as waves
As the electron moves toward the nucleus, the wavelength shortens
Shorter wavelength = higher energy
Higher energy = greater distance from the nucleus

12. Electromagnetic radiation (light) is divided into various classes according to wavelength.

13. Tell me what you know… Which color has the greatest wavelength?
Which color has the shortest wavelength?

14. Electromagnetic Spectrum
Long wavelength --> small frequency --> LOW ENERGY
Short wavelength --> high frequency --> HIGH ENERGY
increasing frequency
increasing wavelength

15. Electromagnetic Spectrum
In increasing energy (from low to high), ROY G BIV

16. Back to Light as energy…
Photons carry energy
E=hc λ
This energy can mean different things

17. Excited State – atom with excess energy
Ground State – lowest possible energy state
Wavelengths of light carry different amounts of energy per photon
Only certain types of photons are produced (see only certain colors)
Quantized – only certain energy levels (and therefore colors) are allowed

18. Emission of Energy by Atoms
When atoms receive energy from some source, they become excited, and they can release this energy by emitting light by releasing a photon.
The energy of the photon corresponds exactly to the energy change experienced by the emitting atom
Red=lower energy=longer wavelength
Violet=higher energy=shorter wavelength

19. Some Definitions Ground State: Lowest possible energy state of an atom
Excited States: Possible Higher energy states. Only certain states are allowed for certain atoms.

20. Figure 11.8: An excited lithium atom emitting a photon of red light to drop to a lower energy state.
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21. Flame Tests
Many elements give off characteristic light which can be used to help identify them.
strontium
sodium
lithium
potassium
copper

23. Emission Spectra
Atoms in the excited state are unstable. The electrons will fall back to the ground state.
When an electron moves from a higher energy level to a lower energy level, energy is released and light is emitted.
This energy is emitted in a form of light. The energy (color) of the light depends on how far the electron falls.

24. As e- ‘fall’ from higher orbitals energy is given off.
Amount of energy given off = to the distance of the fall 6 5 4 3 2 1
nucleus

25. ‘Atomic Fingerprints’
Each element has a unique
atomic line emission spectrum.
Used to identify the element

26. Continuous and Line Spectra
5000
6000
7000
Visible Light
Spectrum Na H Ca Hg
Spectroscopy is a method of identifying unknown substances from their spectra. Because all materials have unique spectra, you can think of these spectra as being molecular fingerprints.
Line spectra of selected elements showing the emission spectra or fingerprint of that element.
The unique set of lines produced is due to the fact that electrons are falling from different excited states in the atoms to the ground state.
Objects at high temperature emit a continuous spectrum of electromagnetic radiation.
A different kind of spectrum is observed when pure samples of individual elements are heated.
When the emitted light is passed through a prism, only a few narrow lines, called a line spectrum, are seen rather than a continuous range of colors.
Using Planck’s equation, the observation of only a few values of  (or ) in the line spectrum meant that only a few values of E were possible — only states that had certain values of energy were possible or allowed.
Any given element has both a characteristic emission spectrum and a characteristic absorption spectrum, which are complementary images.
– Emission spectrum: emission of light by atoms in excited states
– Absorption spectrum: absorption of light by ground-state atoms to produce an excited state
• Emission and absorption spectra form the basis of spectroscopy, which uses spectra to provide information about the structure and composition of a substance or an object

27. Flame Emission Spectra
Photographs of flame tests of burning wooden splints soaked in different salts.
methane gas
wooden splint
strontium ion
copper ion
sodium ion
calcium ion
This technique is called emission spectroscopy.

29. Let’s look more at hydrogen…

30. Figure 11.9: A sample of H atoms receives energy from an external source.
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31. Figure 11.9: The excited atoms release energy by emitting photons.
Excited atom can release some or all of its energy by emitting a photon (electromagnetic radiation “particle”)
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32. Figure 11.10: An excited H atom returns to a lower energy level.
Energy contained in photon = change in energy of atom
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33. Figure 11.11: Colors and wavelengths of photons in the visible region.
Visible light photons emitted by Hydrogen – always the same
Because only certain photons are emitted, only certain energy changes are occurring
Hydrogen atom has certain discrete energy levels
Energy levels of Hydrogen are quantized – only certain values allowed
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34. Figure 11.12: The color of the photon emitted depends on the energy change that produces it.
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35. Figure 11.13: Each photon emitted corresponds to a particular energy change.
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36. Figure 11.14: Continuous (a) and discrete (b) energy levels.
Quantized nature of energy surprised scientists (b)
Previously assumed atom could exist at any energy level (a)
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37. Figure 11.15: The difference between continuous (a) and quantized (b) energy levels.
Staircase – can move from one step to another or even skip, but must be on a step
Ramp – can be at any elevation
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38. Let’s look at Old Atom Models and Compare them to Newer Ones
Rutherford
Bohr
NEW
de Broglie
Schrodinger

39. Figure 11.1: The Rutherford atom.
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40. Rutherford Atom Review
Alpha particle/Gold foil experiment
Nuclear Atom
Nucleus composed of protons & neutrons
Nucleus small compared to atomic size
Electrons account for rest of atom
Unanswered questions:
What are electrons doing? – How are they arranged & how do they move?
Thought electrons revolved around nucleus like planets orbit the sun
Couldn’t explain why electrons aren’t attracted to protons causing atom to collapse
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41. Limitations of Rutherford’s Atomic Model
It explained only a few simple properties of atoms.
It could not explain the chemical properties of elements.
For example, Rutherford’s model could not explain why an object such as the iron scroll shown here first glows dull red, then yellow, and then white when heated to higher and higher temperatures. 41

42. The Bohr Model
In 1913, Niels Bohr (1885–1962), a young Danish physicist and a student of Rutherford, developed a new atomic model.
He changed Rutherford’s model to incorporate newer discoveries about how the energy of an atom changes when the atom absorbs or emits light. 42

43. The Bohr Model The energy levels are also called “electron shells”.
According to Bohr’s atomic model, electrons move in definite orbits around the nucleus, much like planets circle the sun.
These orbits, or energy levels, are located at certain distances from the nucleus.
The energy levels are also called “electron shells”. 43

44. Figure 11.17: The Bohr model of the hydrogen atom.
Electrons moved in circular orbits like planets
Electrons could jump from one orbit to another by emitting/absorbing a photon
Only worked for H, didn’t work for other atoms
Showed experimentally to be incorrect
Paved way for other theories
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45. The Modern Model of the Atom
Accredited to Schrodinger,
According to the wave mechanic model, electrons do not move about an atom in a definite path, like the planets around the sun.
Electrons are NOT in fixed paths
Electrons are in orbitals which are nothing like Bohr’s orbits
The electrons move constantly throughout the energy levels forming an electron cloud

46. Electron Cloud
The space in which electrons are most likely to be found.
The cloud is more dense where the probability of finding the electron is high.
We can’t know exactly where an electron is (Heisenberg Uncertainty Principle)
Electron cloud 46

47. Heisenberg Uncertainty Principle
“One cannot simultaneously determine both the position and momentum of an electron.”
You can find out where the electron is, but not where it is going.
OR…
You can find out where the electron is going, but not where it is!
Werner
Heisenberg

48. Figure 11.18: A representation of the photo of the firefly experiment (lightning bugs).
Shows probability (or likelihood) of where firefly will be found
Usually near the center, but can be found in any of the shaded areas at any time
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49. Figure 11.19: The orbital that describes the hydrogen electron in its lowest possible energy state. Darker pink = greater probability
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50. Drawbacks of wave mechanical model:
Gives no information about when the electron occupies a certain point in space or how it moves
We will probably never know the details of electron motion
Confident that Bohr model is incorrect
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51. Quantum Mechanical Model

52. Quantum Mechanical Model

53. Parts of the Wave Mechanical Model
1. Principle Energy Level (n) – energy level designated by numbers 1-7.
-called principle quantum numbers 1 2 3 4 5 6 7
2. Sublevel – exist within each principle energy level
-the energy within an energy level is slightly different
-each electron in a given sublevel has the same energy
-lowest sublevel = s, then p,
then d, then f s p d f

54. Parts of the Wave Mechanical Model cont.
3. Orbital – region within a sublevel or energy level where electrons can be found
s sublevel – 1 orbital
p sublevel – 3 orbitals
d sublevel – 5 orbitals
f sublevel – 7 orbitals
- ** No more than two electrons can occupy an orbital**
-an orbital can be empty, half-filled, filled

55. Shapes of orbitals All s orbitals are spherical
as the principle energy level
increases the diameter increases.
All p orbitals are dumbbell or figure-8 shaped – all have the same size and shape within an energy level

56. 4 of the d orbitals are 4-leaf clover shaped and the last is a figure-8 with a donut – all have the same size and shape within an energy level

57. f orbitals are complicated!!!!!

58. Principle Energy Level
Summary
Principle Energy Level
# of sublevels
# of orbitals present
s p d f
Total # of orbitals
Maximum # of electrons 1 1 1 2 2 2 4 8 3 3 9 18 4 4 16 32

59. Sublevels in Each Energy Level
Principal Energy Level # of Sublevels sublevels (s)
(s,p)
(s,p,d)
(s,p,d,f)
Orbitals in Each Sublevel
Sublevel
No. of Orbitals
No. of Electrons s 1 2 p 3 6 d 5 10 f 7 14

60. Sulfur = 1s2 2s2 2p6 3s2 3p4 Cd = 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
Electron Configuration – arrangement of the electrons among the various orbitals of the atom Ex: 1s22s22p6 = Neon
Sulfur =
1s2
2s2
2p6
3s2
3p4
Cd =
1s2
2s2
2p6
3s2
3p6
4s2
3d10
4p6
5s2
4d10
Na =
1s2
2s2
2p6
3s1 Ne Na

61. Orbital filling table

62. Electron configuration of the elements of the first three series

63. *Examples: Write the electron configuration using sublevels.
*Examples: Write the electron configuration using sublevels.
Hydrogen: ______________________
Carbon: _________________________
Phosphorous: ______________________________________
Potassium: _______________________________________
5. Calcium:___________________________________________
Iodine: __________________________________________
**NOTE 1: Superscripts should ALWAYS add up to the # of electrons the atom has! (This is a good way of checking yourself!!!**
**NOTE 2: When electrons are in an excited state, they will jump to a higher energy level before all the orbitals in the lower energy level are completely filled.

64. Sample Problems
Which is the electron configuration of an atom in the excited state?
1s22s22p2
1s22s22p1
1s22s22p53s2
1s22s22p63s1

65. An atom in the excited state can have an electron configuration of?
1s22s2
1s22p1
1s22s22p5
1s22s22p6
Which electron configuration represents a potassium atom in the excited state?
1s22s22p63s23p3
1s22s22p63s13p4
1s22s22p63s23p64s1
1s22s22p63s23p54s2

66. Electron Spin Spin – motion that resembles earth rotating
on its axis– clockwise or counterclockwise
Pauli Exclusion Principle – two electrons in the same orbital must have opposite spins
Hund’s Rule – All orbitals within a sublevel must contain at least one electron before any orbital can have two
Orbital Diagram – describes the placement of electrons in orbitals
use arrows to represent electrons with spin
line represents orbital (s=1, p=3, d=5, f=7)
____ full ____ half-full ____ empty

67. Aufbau Order
Aufbau Order – Tool to predict the order in which sublevels will fill.
Aufbau Principle: Must fill lower energy levels first
OR use order on Periodic Table

69. Orbital Diagrams Neon = 1s__ 2s__ 2p__ __ __ Carbon = 1s__ 2s__
Zinc =
1s__
2s__
2p__ __ __
3s__
3p__ __ __
4s__
3d__ __ __ __ __
Gallium =
1s__
2s__
2p__ __ __
3s__
3p__ __ __
4s__
3d__ __ __ __ __
4p__ __ __

70. 10 Ne Atomic # Symbol Orbital Filling Diagram Electron Configuration
(= # of electrons) 10
Symbol Ne
Orbital
Filling
Diagram
Electron
Configuration

71. 5 B Atomic # Symbol Orbital Filling Diagram Electron Configuration
(= # of electrons) 5
Symbol B
Orbital
Filling
Diagram
Electron
Configuration

72. 7 N Atomic # Symbol Orbital Filling Diagram Electron Configuration
(= # of electrons) 7
Symbol N
Orbital
Filling
Diagram
Electron
Configuration

73. 8 O Atomic # Symbol Orbital Filling Diagram Electron Configuration
(= # of electrons) 8
Symbol O
Orbital
Filling
Diagram
Electron
Configuration

74. Valence Configuration – shows just the valence electrons
Noble Gas Configuration – Shorthand configuration that substitutes a noble gas for electrons.
Ex:
Valence Electrons – Electrons in the outermost (highest) principle energy level in an atom, electrons use in bonding (outside of brackets)
Core Electrons – innermost electrons – not involved in bonding (inside brackets)
Valence Configuration – shows just the valence electrons
Na = 1s22s22p63s or [Ne]3s1
Sn = 1s22s22p63s23p64s23d104p65s24d105p2 or [Kr]5s24d105p2
Na = 3s rd Shell/1valence electron
Sn = 5s25p2 5th Shell/4 valence electrons

75. Element Configuration notation Orbital notation Noble gas
Lithium
1s22s1
____ ____ ____ ____ ____
1s s p
[He]2s1
Beryllium
1s22s2
[He]2s2
Boron
1s22s22p1
[He]2s2p1
Carbon
1s22s22p2
[He]2s2p2
Nitrogen
1s22s22p3
1s s p
[He]2s2p3
Oxygen
1s22s22p4
[He]2s2p4
Fluorine
1s22s22p5
[He]2s2p5
Neon
1s22s22p6
[He]2s2p6

76. Figure 11.34: Periodic table with atomic symbols, atomic numbers, and partial electron configurations.
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77. Ion Configurations
To form anions from elements, add 1 or more e- from the highest sublevel.
P [Ne] 3s2 3p3 + 3e- ---> P3- [Ne] 3s2 3p6 or [Ar]

78. Ions
Cations: Take away that many electrons as the number on the charge (+1 means take away highest energy electron in configuration)
Anions: Add that many electrons as the number on the charge. Abide by Aufbau’s Principle