1. AP Chapter 5
Structure of the Atom

2. Review Quiz Chapter 5
Net Ionic Equations

3. John Dalton (1766 – 1844)
Proposed the first scientifically supported atomic theory.

4. Dalton’s Model of the Atom
Dalton's model was that the atoms were tiny, indivisible, indestructible particles and that each one had a certain mass, size, and chemical behavior that was determined by what kind of element they were.
What made us change Dalton’s Atomic Model?

5. J. J. Thomson (1856 – 1940)
Credited for the discovery of the electron.
Invented the mass spectrometer which led to his discovery of isotopes.

6. Scientific Inquiry
Scientists use experimental results to test scientific models, such as the model of the atom.
When experimental results are not consistent with the predictions of a scientific model, the model must be revised or replaced.

7. Lord Ernest Rutherford (1871 – 1937)
Discovered the nucleus of the atom.

8. Rutherford’s Gold Foil Experiment

9. Rutherford’s Gold Foil Experiment

10. Rutherford’s Atomic Model The “Planetary Model” of the Atom
The nucleus is very small, dense, and positively charged.
Electrons surround the nucleus.
Most of the atom is empty space

11. Subatomic Particles PARTICLE SYMBOL CHARGE MASS (amu) LOCATION
electron e- -1 0
orbit nucleus
proton p+ +1 1
inside nucleus
neutron n0

12. Niels Bohr
In 1913 Bohr published a theory about the structure of the atom based on an earlier theory of Rutherford's.
Bohr expanded upon this theory by proposing that electrons travel only in certain successively larger orbits.

13. Bohr Model of the Atom
Electrons orbit the nucleus in orbits that represent specific quantities of energy.
The energies of the electrons in the atom are quantized.
Only certain electron orbits (energy levels) are allowed.
The Bohr Atom

15. Electromagnetic Waves
Properties of waves include speed, frequency, wavelength and energy
All electromagnetic waves including light travel at a speed of 3 x 108 m/s.
However the frequency, wavelength and energy of the waves vary.

16. Wavelength ()
Measured in units of length: m, nm, A

17. Frequency ()
Measured in cycles/second = hertz (Hz)

19. Visible Light

20. Electromagnetic Radiation
For all waves  •  = c
c = the speed of light = 3.00 x 108 m/s 20

21. A photon of red light has a wavelength of 665 nm
A photon of red light has a wavelength of 665 nm. What is the frequency of this light?
665 nm = 665 x 10-9 m
c =  • 
 = c ÷  = 3.00 x 108 m/s ÷ 665 x 10-9 m
 = x 1014/s or Hz
 = 4.51 x 1014/s or Hz

22. An x-ray has a frequency of 7.25 x 1020 Hz. What is the wavelength?
7.25 x 1020 Hz = 7.25 x 1020/s
c =  • 
 = c ÷  = 3.00 x 108 m/s ÷ 7.25 x 1020/s
 = x m
 = 4.14 x m

23. Energy of Electromagnetic Radiation
For all waves: E = h • 
h = 6.63 x J • s or J/Hz 23

24. A photon of red light has a wavelength of 665 nm
A photon of red light has a wavelength of 665 nm. What is the energy of this light?

25. A photon of red light has a wavelength of 665 nm
A photon of red light has a wavelength of 665 nm. What is the energy of this light?
665 nm = 665 x 10-9 m
c =  • 
 = c ÷  = 3.00 x 108 m/s ÷ 665 x 10-9 m
 = x 1014/s or Hz
 = 4.51 x 1014/s or Hz

26. A photon of red light has a wavelength of 665 nm
A photon of red light has a wavelength of 665 nm. What is the energy of this light?
665 nm = 665 x 10-9 m
c =  • 
 = c ÷  = 3.00 x 108 m/s ÷ 665 x 10-9 m
 = x 1014/s or Hz
 = 4.51 x 1014/s or Hz
E = h •  = (6.63 x J • s)(4.51 x 1014/s)
E = x J

27. A photon of red light has a wavelength of 665 nm
A photon of red light has a wavelength of 665 nm. What is the energy of this light?
665 nm = 665 x 10-9 m
c =  • 
 = c ÷  = 3.00 x 108 m/s ÷ 665 x 10-9 m
 = x 1014/s or Hz
 = 4.51 x 1014/s or Hz
E = h •  = (6.63 x J • s)(4.51 x 1014/s)
E = x J = 2.99 x J

28. An x-ray has a frequency of 7.25 x 1020 Hz. What is it’s energy?

29. An x-ray has a frequency of 7.25 x 1020 Hz. What is it’s energy?
E = h •  = (6.63 x J/Hz)(7.25 x 1020Hz)
E = x J
E = 4.81 x J

30. wavelength, frequency and energy
Red Light
 = 665 x 10-9 m
 = 4.51 x 1014 Hz
E = 2.99 x J
 = 4.14 x m
 = 7.25 x 1020 Hz
E = 4.81 x J
X-ray

31. Wavelength, frequency and energy
Wavelength and frequency have an indirect relationship.
Energy and frequency have a direct relationship.
Electromagnetic radiation of short wavelength will have high frequency and high energy.
Electromagnetic radiation of long wavelength will have low frequency and low energy.

32. Niels Bohr
Bohr also described the way atoms emit radiation by suggesting that when an electron jumps from an outer orbit to an inner one, that it emits light.

33. Bohr Model
Electrons move around the nucleus in orbits of definite energies.
The energy of the orbit is related to its distance from the nucleus. The lowest energy is found in the orbit closest to the nucleus.
Radiation is absorbed or emitted when an electron moves from one orbit to another.

34. Bohr Model of the Atom
The Bohr atom
Electrons orbit the nucleus in orbits that represent specific quantities of energy.
Electrons closer to the nucleus have less energy.
The Bohr Atom

35. Ground State
The lowest energy state of an atom.

36. Excited State
Any energy state of an atom that is of higher in energy than the ground state.

37. Energy Absorbed

38. Absorption (Dark – Line) Spectra

39. Energy Emitted
Electron jumps to a lower orbit

40. Emission (Bright – Line) Spectra

41. Emission Spectra

42. The lines present in an emission spectrum are the lines missing in an absorption spectrum.

43. The Heisenberg Uncertainty Principle

44. The Uncertainty principle
Heisenberg determined that it was impossible to experimentally determine both the position and the speed of the electron at the same time.
This became known as the Heisenberg Uncertainty Principle.
It simply means that the electron is so small and moving so fast, that the simple act of trying to measure its speed or position would change either quantity.

45. Problems with the Bohr Model
It violates the Heisenberg Uncertainty Principle because it considers electrons to have known orbits.
It makes poor predictions regarding the spectra of atoms larger than hydrogen.

46. Schrodinger
The Austrian scientist, Erwin Schrödinger, pursued the of the electron having wave properties and it seemed to him that the electron might be like a standing wave around the nucleus.

47. Schrodinger’s Model
This idea agreed very well with Bohr's idea of quantized energy levels: only certain energies and therefore, wavelengths would be allowed in the atom.
This explained why only certain colors (wavelengths) were seen in the spectrum of the hydrogen atom.

48. Schrodinger’s Model
Schrodinger set out to make a mathematical model that assumed the electron was a standing wave around the nucleus.
His solutions to that model agreed not only with the experimental evidence for hydrogen (as Bohr’s did too), but gave excellent results for all atoms when compared to their actual spectrum.

49. Quantum Mechanics: The Structure of Atoms

50. Schrodinger’s Model
This is our modern model of the atom and is known as the Quantum Mechanical Model.
Calculations involving the QM model of the atom can be approximated using computers.
The solutions to these calculations is the basis for modern software that calculates the structure and reactivity of molecules.

51. The Quantum Mechanical Model

52. The Quantum Mechanical Model
The quantum mechanical model is based on quantum theory, which says matter also has properties associated with waves.

53. The Quantum Mechanical Model
According to quantum theory, it’s impossible to know the exact position and momentum of an electron at the same time. This is known as the Uncertainty Principle.

54. The Quantum Mechanical Model
The quantum mechanical model of the atom uses complex shapes of orbitals (sometimes called electron clouds), volumes of space where an electron is likely to be found.

55. The Quantum Mechanical Model
Therefore the quantum mechanical model is based on probability rather than certainty.

56. Absorption and Emission of Energy by Molecules
A photon is a particle representing a specific amount of electromagnetic energy.
When a photon is absorbed by a molecule, the energy of the molecule is increased.
When a photon is released by a molecule, the energy of the molecule is decreased.

57. Absorption and Emission of Energy by Molecules
Different types of molecular motion lead to absorption or emission of photons in different spectral regions.

58. Absorption and Emission of Energy by Molecules
Ultraviolet and visible radiation is associated with electron transitions between energy levels and so can be used to study electronic structure.
This is the same as atomic spectroscopy.
Molecules however have bonds which allow for different types of absorption and emission.

59. Molecular Vibrations

60. Absorption and Emission of Energy by Molecules
Infrared radiation is associated with molecular vibrations and so can be used to detect the type of bonds present within the molecule.

61. Absorption and Emission of Energy by Molecules
Infrared radiation is associated with molecular vibrations and so can be used to detect the type of bonds present within the molecule.

62. Absorption and Emission of Energy by Molecules
The microwave  region of the spectrum is used for rotational spectroscopy.

63. World of Chemistry: Molecular Fingerprints
..\..\..\Videos\World of Chemistry\Molecular Fingerprints.mpg

64. Energy Levels – Sublevels - Orbitals
Electrons in an atom are within atomic orbitals which are within sublevels which are within energy levels.

65. Electron Configuration
4d7
7 electrons are in the d sublevel in the 4th energy level

67. Arrow Diagram 1 2 3 4 5 6 7 s s 2p s 3p 3d s 4p 4d 4f s 5p 5d 5f

68. Write the electron configuration for lead
(Z = 82).

69. The periodic table and electron configuration.

70. Periodic Table and Electron Configurations p 1 2 3 4 5 6 7
d (period-1) 6 7
f (period-2) 6 7
© 1998 by Harcourt Brace & Company

71. C. Periodic Patterns
Example - Germanium
[Ar]
4s2
3d10
4p2

72. Write the abbreviated electron configuration for lead (Z = 82) using the periodic table.6 7
© 1998 by Harcourt Brace & Company

73. A. General Rules Pauli Exclusion Principle
Each orbital can hold TWO electrons with opposite spins.

74. A. General Rules Incorrect Correct Hund’s Rule
Within a sublevel, place one e- per orbital before pairing them.
All electrons in singly filled orbitals have the same direction of spin.
Incorrect
Correct

75. Orbital Filling Diagrams

76. 1s2 2s2 2p4 O B. Notation 1s 2s 2p 8e- Orbital Diagram
Electron Configuration
1s2 2s2 2p4

77. Magnetism Magnetism can be a complicated concept.
We will only deal with magnetism as it is predicted by electron spin which is an oversimplified way of dealing with this subject.

78. Electron Spin & Magnetism
Opposite spins produce opposite magnetic fields.

79. Electron Spin & Magnetism
Most materials with one or more unpaired electrons are at least slightly magnetic. These substances are said to be paramagnetic.
The overall magnetic field strength of atoms with all paired electrons is zero. These substances are said to be diamagnetic.

80. Molecular Orbital Diagrams
One of the reasons that magnetism can be complicated is that often times the bonding within atoms creates situations where multiple atoms bond together creating molecular orbitals which greatly complicates the ability to predict the pairing of electrons.

81. Molecular Orbitals
A molecular orbital is an orbital within a molecule whereas an atomic orbital is an orbital within an individual atom.
Both types of orbitals can hold no more than a pair of electrons.
We can draw a molecular orbital diagram to help explain the magnetic character of some simple diatomic molecules.

82. Molecular Orbitals Diagrams
See the blank MO diagram in your notebook.

83. Molecular Orbitals Diagram for H2

84. Molecular Orbitals Diagram for C2

85. MO Diagrams Can Be Complicated.

86. Write the molecular orbital diagrams for O2 and F2.
Classify each substance as paramagnetic or diamagnetic.
Which substance will be attracted to a magnetic field? Explain.

88. Paramagnetism of O2
..\..\..\Videos\Paramagnetism of Liquid Oxygen.wmv

89. Photoelectric Effect
The photoelectric effect is the emission of electrons from substances that are exposed to light.
Ionization energy is the amount of energy it takes to remove an electron from an atom.

90. Photoelectric Effect
More energy is required to remove successive electrons from atoms.
This is due to Coulomb’s Law.

91. Coulomb’s Law
Coulomb’s Law quantifies the general rule of electrostatics that opposites charges attract and like charges repel.
The electrostatic force between two charged bodies is proportional to the product of the amount of charge on the bodies divided by the square of the distance between them.

92. Coulomb’s Law F= force of attraction or repulsion k = constant
q1 and q2 = the charges of the two bodies
r = radius between the charges

93. Coulomb’s Law
If the bodies are oppositely charged, one positive and one negative, they are attracted toward one another; if the bodies are similarly charged, both positive or both negative, the force between them is repulsive.

94. Coulomb’s Law

95. Coulomb's law helps describe the forces that bind electrons to an atomic nucleus.
Based on Coulomb’s Law, the force between two charged particles is proportional to the magnitude of each of the two charges and inversely proportional to the square of the distance (radius) between them.

96. PES Data – Hidden Slide
Read Adobe File “PES – Explained” prior to covering this with the students.

97. Photoelectric Effect
Light consists of photons of a certain energy (E = h • ).
In the photoelectric effect, light ejects electrons from a material. This requires the photon to have sufficient energy to eject the electron.
This is known as Photoelectron Spectroscopy (PES)
PES data give support to the electron configurations that we use to place electrons within atoms.

98. Photoelectron Spectroscopy (PES)
Photoelectron Spectroscopy (PES) determines the energy needed to eject electrons from a substance.

99. Photoelectron Spectroscopy (PES)
Photoelectron Spectroscopy (PES) supports the shell, subshell model of the atom.
PES data for multielectron atoms show that certain electrons have ionization energies that are relatively close to one another. We therefore group these electrons into shells and subshells as a result.

100. Photoelectron Spectroscopy (PES)
Photoelectron Spectroscopy (PES) data for atoms also helps us to confirm the number of electrons within particular shells and subshells.
The intensity (shown as the height on a graph) of the photoemission signal at a given energy is a measure of the number of electrons in that shell or subshell. In other words the height of the PES signal is proportional to the number of electrons in that sublevel.

103. Hidden Slide
Note the peak for the is of He is twice as high as the peak for H.
2 e- vs. 1e-.

107. What is the electron configuration of this element?

108. How does this PES data show the relative number of electrons within a sublevel?

109. How does this PES data show the relative energies of the sublevels?

110. Go to the website below and go through the first 10 elements:
Show different sublevels and different numbers of electrons in each sublevel based on peak intensity (height).

111. All the previous PES slides all come from this website

112. Spectrophotometer Lab: Preparing and Diluting Solutions

113. Spectrophotometer Lab: Preparing and Diluting Solutions