1. Modern Atomic Theory and the Periodic Table Chapter 10
Hein and Arena
Eugene Passer Chemistry Department Bronx Community College
© John Wiley and Sons, Inc.
Version 1.1

2. Chapter Outline 10.1 A Brief History
10.5 Atomic Structures of the First Elements
10.2 Electromagnetic Radiation
10.6 Electron Structures and the Periodic Table
10.3 The Bohr Atom
10.4 Energy Levels of Electrons

3. A Brief History

5. Electromagnetic Radiation

6. Energy can travel through space as electromagnetic radiation.
Examples
Energy can travel through space as electromagnetic radiation.

7. light from the sun x-rays microwaves radio waves television waves
radiant heat
All show wavelike behavior.
Each travels at the same speed in a vacuum.
3.00 x 108 m/s

8. Characteristics of a Wave

9. Wavelength (λ)

10. Light has the properties of a wave.
wavelength
(measured from peak to peak)
wavelength
(measured from trough to trough)
10.1

11. Frequency (ν)

12. Frequency is the number of wavelengths that pass a particular point per second.
10.1

13. Speed (v)

14. Speed is how fast a wave moves through space.
10.1

15. Light also exhibits the properties of a particle
Light also exhibits the properties of a particle. Light particles are called photons.
Both the wave model and the particle model are used to explain the properties of light.

16. The Electromagnetic Spectrum

17. X-rays are part of the electromagnetic spectrum
visible light is part of the electromagnetic spectrum
Infrared light is part of the electromagnetic spectrum
10.2

18. The Bohr Atom

19. At high temperatures or voltages, elements in the gaseous state emit light of different colors.
When the light is passed through a prism or diffraction grating a line spectrum results.

20. Each element has its own unique set of spectral emission lines that distinguish it from other elements.
These colored lines indicate that light is being emitted only at certain wavelengths.
Line spectrum of hydrogen. Each line corresponds to the wavelength of the energy emitted when the electron of a hydrogen atom, which has absorbed energy falls back to a lower principal energy level.
10.3

21. Niels Bohr

22. Niels Bohr, a Danish physicist, in 1912-1913 carried out research on the hydrogen atom.

23. The Bohr Atom

24. An electron has a discrete energy when it occupies an orbit.
Electrons revolve around the nucleus in orbits that are located at fixed distances from the nucleus.
10.4

25. The color of the light emitted corresponds to one of the lines of the hydrogen spectrum.
When an electron falls from a higher energy level to a lower energy level a quantum of energy in the form of light is emitted by the atom.
10.4

26. Different lines of the hydrogen spectrum correspond to different electron energy level shifts.
10.4

27. Light is not emitted continuously
Light is not emitted continuously. It is emitted in discrete packets called quanta.
10.4

28. E1 E2 E3
An electron can have one of several possible energies depending on its orbit.
10.4

29. Bohr’s calculations succeeded very well in correlating the experimentally observed spectral lines with electron energy levels for the hydrogen atom.
Bohr’s methods did not succeed for heavier atoms.
More theoretical work on atomic structure was needed.

30. In 1924 Louis De Broglie suggested that all objects have wave properties.
De Broglie showed that the wavelength of ordinary sized objects, such as a baseball, are too small to be observed.
For objects the size of an electron the wavelength can be detected.

31. In 1926 Erwin Schröedinger created a mathematical model that showed electrons as waves.
Schröedinger’s work led to a new branch of physics called wave or quantum mechanics.
Using Schröedinger’s wave mechanics, the probability of finding an electron in a certain region around the atom can be determined.
The actual location of an electron within an atom cannot be determined.

32. Based on wave mechanics it is clear that electrons are not revolving around the nucleus in orbits.
Instead of being located in orbits, the electrons are located in orbitals.
An orbital is a region around the nucleus where there is a high probability of finding an electron.

33. Energy Levels of Electrons

34. According to Bohr the energies of electrons in an atom are quantized.
The wave-mechanical model of the atom also predicts discrete principal energy levels within the atom

35. As n increases, the energy of the electron increases.
The first four principal energy levels of the hydrogen atom.
Each level is assigned a principal quantum number n.
10.7

36. Each principal energy level is subdivided into sublevels.
10.7, 10.8

37. Within sublevels the electrons are found in orbitals.
An s orbital is spherical in shape.
The spherical surface encloses a space where there is a 90% probability that the electron may be found.
10.10

38. An atomic orbital can hold a maximum of two electrons.
An electron can spin in one of two possible directions represented by ↑ or ↓.
The two electrons that occupy an atomic orbital must have opposite spins.
This is known as the Pauli Exclusion Principal.
10.10

39. A p sublevel is made up of three orbitals.
Each p orbital has two lobes.
Each p orbital can hold a maximum of two electrons.
A p sublevel can hold a maximum of 6 electrons.
10.10

40. The three p orbitals share a common center.pz
The three p orbitals share a common center.
The three p orbitals point in different directions. px py
10.10

41. A d sublevel is made up of five orbitals.
The five d orbitals all point in different directions.
Each d orbital can hold a maximum of two electrons.
A d sublevel can hold a maximum of 10 electrons.
10.11

42. Number of Orbitals in a Sublevel

43. Distribution of Subshells by Principal Energy Level
2p 2p 2p
n = 3 3s
3p 3p 3p
3d 3d 3d 3d 3d
n = 4 4s
4p 4p 4p
4d 4d 4d 4d 4d
4f 4f 4f 4f 4f 4f 4f

44. The Hydrogen Atom
The diameter of hydrogen’s nucleus is about cm.
The diameter of hydrogen’s electron cloud is about 10-8 cm.
In the ground state hydrogen’s single electron lies in the 1s orbital.
Hydrogen can absorb energy and the electron will move to excited states.
The diameter of hydrogen’s electron cloud is about 100,000 times greater than the diameter of its nucleus.
10.12

45. Atomic Structure of the First 18 Elements

46. To determine the electronic structures of atoms, the following guidelines are used.

47. No more than two electrons can occupy one orbital
10.10

48. 2 s orbital
1 s orbital
Electrons occupy the lowest energy orbitals available. They enter a higher energy orbital only after the lower orbitals are filled.
For the atoms beyond hydrogen, orbital energies vary as s<p<d<f for a given value of n.
10.10

49. Each orbital in a sublevel is occupied by a single electron before a second electron enters. For example, all three p orbitals must contain one electron before a second electron enters a p orbital.
10.10

50. Nuclear makeup and electronic structure of each principal energy level of an atom.
number of protons and neutrons in the nucleus
number of electrons in each sublevel
10.13

51. Electron Configuration
Number of electrons in sublevel orbitals
Arrangement of electrons within their respective sublevels.
2p6
Principal energy level
Type of orbital

52. Orbital Filling

53. In the following diagrams boxes represent orbitals.
Electrons are indicated by arrows: ↑ or ↓.
Each arrow direction represents one of the two possible electron spin states.

54. Filling the 1s Sublevel

55. H ↑
1s1
Hydrogen has 1 electron. It will occupy the orbital of lowest energy which is the 1s. He
Helium has two electrons. Both helium electrons occupy the 1s orbital with opposite spins. ↑ ↓
1s2

56. Filling the 2s Sublevel

57. Li ↑ ↓ ↑
1s22s1 1s 2s
The 1s orbital is filled. Lithium’s third electron will enter the 2s orbital. Be ↑ ↓
The 2s orbital fills upon the addition of beryllium’s third and fourth electrons. 1s 2s ↑ ↓
1s22s2

58. Filling the 2p Sublevel

59. ↑ ↓ ↑ ↓ ↑ B 1s22s22p1 C ↑ ↓ ↑ ↑ 1s22s22p2 N ↑ ↓ ↑ ↑ ↑ 1s22s22p3 1s 2s
Boron has the first p electron. The three 2p orbitals have the same energy. It does not matter which orbital fills first. C 1s 2s 2p ↑ ↓
The second p electron of carbon enters a different p orbital than the first p electron so as to give carbon the lowest possible energy. ↑ ↑
1s22s22p2 N 1s 2s 2p ↑ ↓
The third p electron of nitrogen enters a different p orbital than its first two p electrons to give nitrogen the lowest possible energy. ↑ ↑ ↑
1s22s22p3

60. ↑ ↓ O ↑ ↓ ↑ ↑ 1s22s22p4 F ↑ ↓ ↑ ↓ ↑ ↓ ↑ 1s22s22p5 1s 2s 2p 2p 1s 2s
There are four electrons in the 2p sublevel of oxygen. One of the 2p orbitals is now occupied by a second electron, which has a spin opposite to that of the first electron already in the orbital. ↑ ↓ ↑ ↑
1s22s22p4 2p F 1s 2s ↑ ↓
There are five electrons in the 2p sublevel of fluorine. Two of the 2p orbitals are now occupied by a second electron, which has a spin opposite to that of the first electron already in the orbital. ↑ ↓ ↑ ↓ ↑
1s22s22p5

61. 2p Ne 1s 2s ↑ ↓
There are 6 electrons in the 2p sublevel of neon, which fills the sublevel. ↑ ↓ ↑ ↓ ↑ ↓
1s22s22p6

62. Filling the 3s Sublevel

63. ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ Na 1s22s22p63s1 Mg ↑ ↓ ↓ 1s22s22p63s2 1s 2s 2p
The 2s and 2p sublevels are filled. The next electron enters the 3s sublevel of sodium. Mg 1s 2s 2p 3s ↑ ↓
The 3s orbital fills upon the addition of magnesium’s twelfth electron. ↓
1s22s22p63s2

66. Electron Structures and the Periodic Table

67. Mendeleev’s arrangement is the precursor to the modern periodic table.
In 1869 Dimitri Mendeleev of Russia and Lothar Meyer of Germany independently published periodic arrangements of the elements based on increasing atomic masses.
Mendeleev’s arrangement is the precursor to the modern periodic table.

68. Horizontal rows are called periods.
Period numbers correspond to the highest occupied energy level.
10.14

69. Groups are numbered with Roman numerals.
Elements with similar properties are organized in groups or families.
Elements in the B groups are designated transition elements.
Elements in the A groups are designated representative elements.
10.14

70. The chemical behavior and properties of elements in a family are associated with the electron configuration of its elements.
For A family elements the valence electron configuration is the same in each column.
10.15

71. With the exception of helium which has a filled s orbital, the nobles gases have filled p orbitals.
10.15

72. The electron configuration of any of the noble gas elements can be represented by the symbol of the element enclosed in square brackets. B
1s22s22p1
[He]2s22p1 Na
1s22s22p63s1
[Ne]3s1 Cl
1s22s22p63s23p5
[Ne]3s23p5

73. The electron configuration of argon is
1s22s22p63s23p6
The elements after argon are potassium and calcium Instead of entering a 3d orbital, the valence electrons of these elements enter the 4s orbital. K
1s22s22p63s23p64s1
[Ar]4s1 Ca
1s22s22p63s23p6 4s2
[Ar]4s2

74. d orbital numbers are 1 less than the period number d orbital filling
Arrangement of electrons according to sublevel being filled.
10.16

75. f orbital numbers are 2 less than the period number f orbital filling
Arrangement of electrons according to sublevel being filled.
10.16

76. Period number corresponds with the highest energy level occupied by electrons in that period.
10.17

77. The elements of a family have the same outermost electron configuration except that the electrons are in different energy levels.
The group numbers for the representative elements are equal to the total number of outermost electrons in the atoms of the group.
10.17

78. The End