1. Chapter 5 Electrons In Atoms 5.1 Revising the Atomic Model
5.2 Electron Arrangement in Atoms
5.3 Atomic Emission Spectra and
the Quantum Mechanical Model
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2. Energy Levels in Atoms
The Bohr Model
Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus.
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3. Each possible electron orbit in Bohr’s model has a fixed energy.
Energy Levels in Atoms
The Bohr Model
Each possible electron orbit in Bohr’s model has a fixed energy.
The fixed energies an electron can have are called energy levels.
A quantum of energy is the amount of energy required to move an electron from one energy level to another energy level.
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4. Energy Levels in Atoms
The Bohr Model
The rungs on this ladder are somewhat like the energy levels in Bohr’s model of the atom.
A person on a ladder cannot stand between the rungs. Similarly, the electrons in an atom cannot exist between energy levels.
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5. Energy Levels in Atoms
The Bohr Model
The rungs on this ladder are somewhat like the energy levels in Bohr’s model of the atom.
The energy levels in atoms are unequally spaced, like the rungs in this unusual ladder. The higher energy levels are closer together.
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6. The Quantum Mechanical Model
The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus of an atom.
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7. The Quantum Mechanical Model
In the quantum mechanical model, the probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloudlike region.
The cloud is more dense where the probability of finding the electron is high.
Electron cloud
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8. These numbers are assigned the values n = 1, 2, 3, 4, and so forth.
Atomic Orbitals
The energy levels of electrons in the quantum mechanical model are labeled by principal quantum numbers (n).
These numbers are assigned the values n = 1, 2, 3, 4, and so forth.
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9. Atomic Orbitals
Each energy sublevel corresponds to one or more orbitals of different shapes. The orbitals describe where an electron is likely to be found.
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10. Different atomic orbitals are denoted by letters.
The s orbitals are spherical.
The p orbitals are dumbbell-shaped.
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11. Atomic Orbitals
For a given principal energy level greater than 1, there is one s orbital, 3 p orbitals...
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12. Atomic Orbitals
For a given principal energy level greater than 1, there is one s orbital, 3 p orbitals, and 5 d orbitals.
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13. Atomic Orbitals
The numbers and types of atomic orbitals depend on the principal energy level.
Summary of Principal Energy Levels and Sublevels
Principal energy level
Number of sublevels
Type of sublevel
Maximum number of electrons
n = 1 1
1s (1 orbital) 2
n = 2
2s (1 orbital), 2p (3 orbitals) 8
n = 3 3
3s (1 orbital), 3p (3 orbitals),
3d (5 orbitals) 18
n = 4 4
4s (1 orbital), 4p (3 orbitals),
4d (5 orbitals), 4f (7 orbitals) 32
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14. Atomic Orbitals
The principal quantum number, n, always equals the number of sublevels within that principal energy level.
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15. The number of orbitals in a principal energy level is equal to n2.
Atomic Orbitals
The principal quantum number, n, always equals the number of sublevels within that principal energy level.
The number of orbitals in a principal energy level is equal to n2.
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16. The number of orbitals in a principal energy level is equal to n2.
Atomic Orbitals
The principal quantum number, n, always equals the number of sublevels within that principal energy level.
The number of orbitals in a principal energy level is equal to n2.
A maximum of two electrons can occupy an orbital.
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17. The number of orbitals in a principal energy level is equal to n2.
Atomic Orbitals
The principal quantum number, n, always equals the number of sublevels within that principal energy level.
The number of orbitals in a principal energy level is equal to n2.
A maximum of two electrons can occupy an orbital.
Therefore, the maximum number of electrons that can occupy a principal energy level is given by the formula 2n2.
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18. Calculate the maximum number of electrons in the 5th principal energy level (n = 5).
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19. Calculate the maximum number of electrons in the 5th principal energy level (n = 5).
The maximum number of electrons that can occupy a principal energy level is given by the formula 2n2. If n = 5, 2n2 = 50.
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20. CHEMISTRY & YOU
Why do scientists no longer use physical models to describe the motion of electrons?
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21. Chapter 5 Electrons In Atoms 5.2 Electron Arrangement in Atoms
5.1 Revising the Atomic Model
5.2 Electron Arrangement
in Atoms
5.3 Atomic Emission Spectra and
the Quantum Mechanical Model
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22. Electron Configurations
The ways in which electrons are arranged in various orbitals around the nuclei of atoms are called electron configurations.
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23. Electron Configurations
Three rules—the aufbau principle, the Pauli exclusion principle, and Hund’s rule—tell you how to find the electron configurations of atoms.
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24. Electron Configurations
Aufbau Principle
According to the aufbau principle, electrons occupy the orbitals of lowest energy first. In the aufbau diagram, each box represents an atomic orbital.
Increasing energy 6s 5s 4s 3s 2s 1s 6p 5p 5d 4p 4d 4f 3p 3d 2p
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25. Electron Configurations
Aufbau Principle
Increasing energy 6s 5s 4s 3s 2s 1s 6p 5p 5d 4p 4d 4f 3p 3d 2p
The aufbau diagram shows the relative energy levels of the various atomic orbitals. Orbitals of greater energy are higher on the diagram.
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26. Electron Configurations
Aufbau Principle
Increasing energy 6s 5s 4s 3s 2s 1s 6p 5p 5d 4p 4d 4f 3p 3d 2p
The range of energy levels within a principal energy level can overlap the energy levels of another principal level.
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27. Electron Configurations
Pauli Exclusion Principle
According to the Pauli exclusion principle, an atomic orbital may describe at most two electrons.
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28. Electron Configurations
Pauli Exclusion Principle
According to the Pauli exclusion principle, an atomic orbital may describe at most two electrons.
To occupy the same orbital, two electrons must have opposite spins; that is, the electron spins must be paired.
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29. Electron Configurations
Pauli Exclusion Principle
Spin is a quantum mechanical property of electrons and may be thought of as clockwise or counterclockwise.
A vertical arrow indicates an electron and its direction of spin ( or ).
An orbital containing paired electrons is written as .
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30. Electron Configurations
Hund’s Rule
According to Hund’s rule, electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible.
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31. Electron Configurations
Hund’s Rule
Three electrons would occupy three orbitals of equal energy as follows.
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32. Electron Configurations
Hund’s Rule
Three electrons would occupy three orbitals of equal energy as follows.
Electrons then occupy each orbital so that their spins are paired with the first electron in the orbital.
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33. Electron Configurations
Let’s look at the orbital filling diagram of the oxygen atom.
Electron Configurations of Selected Elements
Element 1s 2s
2px 2py 2pz 3s
Electron configuration H
1s1 He
1s2 Li
1s22s1 C
1s22s22p2 N
1s22s22p3 O
1s22s22p4 F
1s22s22p5 Ne
1s22s22p6 Na
1s22s22p63s1
An oxygen atom contains eight electrons.
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34. Electron Configurations
Look at the orbital filling diagram of the oxygen atom.
Electron Configurations of Selected Elements
Element 1s 2s
2px 2py 2pz 3s
Electron configuration H
1s1 He
1s2 Li
1s22s1 C
1s22s22p2 N
1s22s22p3 O
1s22s22p4 F
1s22s22p5 Ne
1s22s22p6 Na
1s22s22p63s1
Each of the three 2p orbitals has one electron. The remaining electron now pairs with an electron occupying one of the 2p orbitals.
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35. Electron Configurations
A convenient shorthand method for showing the electron configuration of an atom involves writing the energy level and the symbol for every sublevel occupied by an electron.
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36. Electron Configurations
A convenient shorthand method for showing the electron configuration of an atom involves writing the energy level and the symbol for every sublevel occupied by an electron.
You indicate the number of electrons occupying that sublevel with a superscript.
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37. Electron Configurations
For hydrogen, with one electron in a 1s orbital, the electron configuration is written 1s1.
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38. Electron Configurations
For hydrogen, with one electron in a 1s orbital, the electron configuration is written 1s1.
For oxygen, with two electrons in a 1s orbital, two electrons in a 2s orbital, and four electrons in 2p orbitals, the electron configuration is 1s22s22p4.
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39. Electron Configurations
For hydrogen, with one electron in a 1s orbital, the electron configuration is written 1s1.
For oxygen, with two electrons in a 1s orbital, two electrons in a 2s orbital, and four electrons in 2p orbitals, the electron configuration is 1s22s22p4.
Note that the sum of the superscripts equals the number of electrons in the atom.
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40. CHEMISTRY & YOU
Explain why the correct electron configuration of oxygen is 1s22s22p4 and not 1s22s22p33s1.
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41. CHEMISTRY & YOU
Explain why the correct electron configuration of oxygen is 1s22s22p4 and not 1s22s22p33s1.
The 2p orbitals are lower in energy than the 3s orbital, so they will be completely filled before any electrons will be found in the 3s orbital.
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42. Writing Electron Configurations
Sample Problem 5.1
Writing Electron Configurations
The atomic number of phosphorus is 15. Write the electron configuration of a phosphorus atom.
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43. Solve Apply the concepts to this problem.
Sample Problem 5.1
Solve Apply the concepts to this problem. 2
Write the electron configuration.
The electron configuration of phosphorus is 1s22s22p63s23p3.
The superscripts add up to the number of electrons.
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44. Electron Configurations
Exceptional Electron Configurations
You can obtain correct electron configurations for the elements up to vanadium (atomic number 23) by following the aufbau diagram for orbital filling.
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45. Electron Configurations
Exceptional Electron Configurations
You can obtain correct electron configurations for the elements up to vanadium (atomic number 23) by following the aufbau diagram for orbital filling.
If you were to continue in that fashion, however, you would assign chromium and copper the following incorrect configurations.
Cr 1s22s22p63s23p63d44s2
Cu 1s22s22p63s23p63d94s2
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46. Electron Configurations
Exceptional Electron Configurations
The correct electron configurations are as follows:
Cr 1s22s22p63s23p63d54s1
Cu 1s22s22p63s23p63d104s1
These arrangements give chromium a half-filled d sublevel and copper a filled d sublevel.
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47. Electron Configurations
Exceptional Electron Configurations
Some actual electron configurations differ from those assigned using the aufbau principle because although half-filled sublevels are not as stable as filled sublevels, they are more stable than other configurations.
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48. What is the correct electron configuration of a sulfur atom?
A. 1s22s22p43s23p6
B. 1s22s22p63s23p3
C. 1s22s22p63s23p4
D. 1s22s22p63s63p2
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49. What is the correct electron configuration of a sulfur atom?
A. 1s22s22p43s23p6
B. 1s22s22p63s23p3
C. 1s22s22p63s23p4
D. 1s22s22p63s63p2
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50. What gives gas-filled lights their colors?
CHEMISTRY & YOU
What gives gas-filled lights their colors?
An electric current passing through the gas in each glass tube makes the gas glow with its own characteristic color.

51. Light and Atomic Emission Spectra
What causes atomic emission spectra?

52. Light and Atomic Emission Spectra
The Nature of Light
Light consists of electromagnetic waves.
Electromagnetic radiation includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays.

53. Light and Atomic Emission Spectra
The Nature of Light
White light - sun and incandescent bulbs - emit, light with a range of wavelengths and frequencies.
When passed through a prism, the different wavelengths separate into a spectrum of colors.
Red – longest wavelength, lowest frequency
Violet – shortest wavelength, highest frequency

54. Light and Atomic Emission Spectra
Low energy
( = 700 nm)
High energy
( = 380 nm)
Frequency  (s-1)
3 x 106
3 x 1012
3 x 1022
102
10-8
10-14
Wavelength  (m)

55. Light and Atomic Emission Spectra
When an electron has its lowest possible energy, the atom is in its ground state.
After absorbing energy – excited state
These electrons lose energy by emitting light when they return to lower energy levels.

56. Light and Atomic Emission Spectra
A prism separates light into the colors it contains. White light produces a rainbow of colors.
Screen
Slit
Prism
Light
bulb

57. Light and Atomic Emission Spectra
Light from a helium lamp produces discrete lines.
Screen
Slit
Prism
Helium
lamp

58. Light and Atomic Emission Spectra
The wavelengths of the spectral lines are characteristic of the element, and they make up the atomic emission spectrum of the element.
No two elements have the same emission spectrum.

59. Light also behaves as a particle
Wave-particle duality
Photoelectric effect - electrons are ejected when light shines on a metal.
These light quanta are called photons.

60. The Quantum Concept and Photons
The Photoelectric Effect
No electrons are ejected because the frequency of the light is below the threshold frequency.
If the light is at or above the threshold frequency, electrons are ejected.
If the frequency is increased, the ejected electrons will travel faster.

61. CHEMISTRY & YOU
The glass tubes in lighted signs contain helium, neon, argon, krypton, or xenon gas, or a mixture of these gases. Why do the colors of the light depend on the gases that are used?

62. CHEMISTRY & YOU
The glass tubes in lighted signs contain helium, neon, argon, krypton, or xenon gas, or a mixture of these gases. Why do the colors of the light depend on the gases that are used?
Each different gas has its own characteristic emission spectrum, creating different colors of light when excited electrons return to the ground state.

63. Quantum Mechanics
Given that light behaves as waves and particles, can particles of matter behave as waves?
Yes!

64. The Wavelike Nature of Matter
Quantum Mechanics
The Wavelike Nature of Matter
Today, the wavelike properties of beams of electrons are useful in viewing objects that cannot be viewed with an optical microscope.
The electrons in an electron microscope have much smaller wavelengths than visible light.
These smaller wavelengths allow a much clearer enlarged image of a very small object, such as this pollen grain, than is possible with an ordinary microscope.

65. Quantum Mechanics
Classical mechanics adequately describes the motions of bodies much larger than atoms, while quantum mechanics describes the motions of subatomic particles and atoms as waves.

66. Quantum Mechanics The Heisenberg Uncertainty Principle
The Heisenberg uncertainty principle states that it is impossible to know both the velocity and the position of a particle at the same time.

67. To locate an electron, you might strike it with a photon.
Quantum Mechanics
To locate an electron, you might strike it with a photon.
The electron has such a small mass that striking it with a photon affects its motion in a way that cannot be predicted accurately.
The very act of measuring the position of the electron changes its velocity, making its velocity uncertain.
After collision: The impact changes the electron’s velocity, making it uncertain.
Before collision: A photon strikes an electron during an attempt to observe the electron’s position.