1. Modern Atomic Theory and the Periodic Table

2. Outline
A Brief History Electromagnetic Radiation and the Electromagnetic Spectrum The Bohr Atom Energy Levels of Electrons Atomic Structure of the First 18 Elements Electron Structures and the Periodic Table

3. A Brief History

5. Electromagnetic Radiation

6. Electromagnetic Radiation
travels at the speed of light: 3 X 108 m/s
carries energy
does not have mass
exhibits behavior of both waves and particles
has electrical and magnetic components
Sound waves are NOT electromagnetic radiation and travel much slower: 344m/s in dry air.

7. Examples of Electromagnetic Radiation
light, all the colors of the rainbow
radio and TV waves
microwaves
X-rays
radiant heat

8. Characteristics of a wave:
wavelength
(measured from peak to peak)
wavelength
(measured from trough to trough)

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

10. Speed is how fast a wave moves through space.

11. EMR also exhibits the properties of a particle
EMR also exhibits the properties of a particle. EMR particles are called photons. Both the wave model and the particle model are used to explain the properties of EMR.

12. The Electromagnetic Spectrum

14. Equations
ν=c/λ
Tells us the higher the frequency the shorter the wavelength.
E=hv
Tells us the higher the frequency the higher the energy.

15. Which color of light has more energy?
Blue Orange
The answer is blue.

16. The Bohr Atom

17. 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.

18. 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.

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

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

21. 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.
The color of the light emitted corresponds to one of the lines of the hydrogen spectrum.

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

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

24. 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.

25. 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.

26. 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.

27. 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.

28. Energy Levels of Electrons

29. The wave-mechanical model of the atom predicts discrete principal energy levels within the atom

30. 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.

31. Each principal energy level is subdivided into sublevels.

32. 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.

33. 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.

34. 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.

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

36. 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.

37. Number of Orbitals in a Sublevel

38. 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

39. The Hydrogen Atom
The diameter of hydrogen’s electron cloud is about 100,000 times greater than the diameter of its nucleus.
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 nucleus is about cm.
The diameter of hydrogen’s electron cloud is about 10-8 cm.

40. Atomic Structure of the First 18 Elements

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

42. No more than two electrons can occupy one orbital

43. 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.

44. 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.

45. 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

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

47. Orbital Filling

48. 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.

49. Filling the 1s Sublevel

50. 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

51. Filling the 2s Sublevel

52. 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

53. Filling the 2p Sublevel

54. ↑ ↓ ↑ ↓ ↑ 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

55. ↑ ↓ 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

56. 2p Ne 1s 2s ↑ ↓
There are 6 electrons in the 2p sublevel of neon, which fills the sublevel. ↑ ↓ ↑ ↓ ↑ ↓
1s22s22p6

57. Filling the 3s Sublevel

58. ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ 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

61. Electron Structures and the Periodic Table

62. 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.

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

64. Elements with similar properties are organized in groups or families.
10.14

65. Elements in the A groups are designated representative elements
10.14

66. Elements in the B groups are designated
transition elements
10.14

67. 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.

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

69. B 1s22s22p1 [He]2s22p1 Na 1s22s22p63s1 [Ne]3s1 Cl 1s22s22p63s23p5
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

70. 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

71. The number of a d orbital is 1 less than its period number
d orbital filling
Arrangement of electrons according to sublevel being filled.

72. The number of an f orbital is 2 less than its period number
f orbital filling
Arrangement of electrons according to sublevel being filled.

73. A period number corresponds to the highest energy level occupied by electrons in the period.

74. 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.