1. Modern Atomic Theory and the Periodic Table Chapter 10

2. Chapter 10 - Modern Atomic Theory and the Periodic Table
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

4. Electromagnetic Radiation
Examples
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

5. Characteristics of a Wave
Wavelength (λ)
Light has the properties of a wave.
wavelength
(measured from peak to peak)
wavelength
(measured from trough to trough)
10.1

6. Frequency (n) is the number of wavelengths that pass a particular point per second.
10.1

7. Speed (v) is how fast a wave moves through space.
10.1

8. 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.

9. The Electromagnetic Spectrum
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

10. The Bohr Atom
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.

11. These colored lines indicate that light is being emitted only at certain wavelengths.
Each element has its own unique set of spectral emission lines that distinguish it from other elements.
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

12. Niels Bohr
Niels Bohr, a Danish physicist, in carried out research on the hydrogen atom.

13. The Bohr Atom
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

14. The Bohr Atom
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

15. The Bohr Atom
Different lines of the hydrogen spectrum correspond to different electron energy level shifts.
10.4

16. The Bohr Atom
Light is not emitted continuously. It is emitted in discrete packets called quanta.
10.4

17. The Bohr Atom E1 E2 E3
An electron can have one of several possible energies depending on its orbit.
10.4

18. The Bohr Atom Bohr’s methods did not succeed for heavier atoms.
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.

19. For objects the size of an electron the wavelength can be detected.
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.

20. 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.

21. 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.

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

24. 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

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

26. 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

27. 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

28. 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

29. The P orbitals 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

30. 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

31. Number of Orbitals in a Sublevel

32. 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

33. 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

34. Atomic Structure of the First 18 Elements
To determine the electronic structures of atoms, the following guidelines are used.

35. Pauli exclusion principle
No more than two electrons can occupy one orbital
10.10

36. 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

37. 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

38. 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

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

40. Orbital Filling Electrons are indicated by arrows: ↑ or ↓.
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.

41. H ↑ 1s1 He ↑ ↓ 1s2 Filling the 1s Sublevel
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

42. Li ↑ ↓ ↑ 1s22s1 Be ↑ ↓ ↑ ↓ 1s22s2 Filling the 2s Sublevel 1s 2s 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

43. ↑ ↓ B ↑ ↓ ↑ 1s22s22p1 C ↑ ↓ ↑ ↑ 1s22s22p2 Filling the 2p Sublevel 1s
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

44. 1s22s22p3 N ↑ ↓ ↑ ↑ ↑ ↑ ↓ O ↑ ↓ ↑ ↑ 1s22s22p4 1s 2s 2p 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. ↑ ↑ ↑ 1s 2s 2p ↑ ↓ O
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

45. F ↑ ↓ ↑ ↓ ↑ ↓ ↑ 1s22s22p5 Ne ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ 1s22s22p6 2p 1s 2s 2p 1s
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 2p Ne 1s 2s ↑ ↓
There are 6 electrons in the 2p sublevel of neon, which fills the sublevel. ↑ ↓ ↑ ↓ ↑ ↓
1s22s22p6

46. ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ Na 1s22s22p63s1 Mg ↑ ↓ ↓ 1s22s22p63s2
Filling the 3s Sublevel ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ Na
1s22s22p63s1 1s 2s 2p 3s
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

49. Electron Filling Order

50. Sublevel energy level order:
1s < 2s < 2p < 3s < 3p <
4s < 3d < 4p < 5s < 4d <
5p < 6s < 4f < 5d < 6p <
7s < 5f < 6d
You can memorize this sequence or....

52. Mendeleev’s arrangement is the precursor to the modern periodic table.
Electron Structures and the 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.

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

54. 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

55. Elements in the U.S. B-groups
Main group elements
Elements in the U.S.
A-groups
Transition elements
Elements in the U.S. B-groups
Metals
Elements on the left of the stair-step line
Nonmetals
Elements on the right of the stair-step line
Metalloids
Elements on the stair-step line

56. Elemental Symbols and the Periodic Table

57. 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

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

59. To write an electron configuration using a noble-gas core:
1. Find the highest atomic-numbered
noble gas (Group 8A element)
less than the atomic number
of the element for which the
configuration is being written
2. Write the elemental symbol of the
noble gas in square brackets, followed
by the remaining configuration

60. 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

61. 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

62. Exceptions to the conventional filling order:
1. d4 configurations generally do not exist
Chromium (Z = 24):
Systematic prediction:
Cr: [Ar]4s23d4
But d4 is not likely,
so promote an electron from the 4s sublevel:
Cr: [Ar]4s13d5

63. 2. d9 configurations generally do not exist
Copper (Z = 29):
Systematic prediction:
Cu: [Ar]4s23d9
But d9 is not likely,
so promote an electron from the 4s sublevel:
Cu: [Ar]4s13d10

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

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

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

67. 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

68. Chapter 10 - Modern Atomic Theory and the Periodic Table
10.1 A Brief History
10.2 Electromagnetic Radiation
10.3 The Bohr Atom – Niels Bohr description of the atom (electron orbitals).
10.4 Energy Levels of Electrons – Electron configuration (from the periodic table), s, p, d, and f orbitals.
10.5 Atomic structures of the First 18 Elements – Valence electrons, Representatives and Transition elements, Families names.
10.6 Electron Structures and the Periodic table – Relationship between group number and valence electrons.