2. Review previous chapter 2 of atomic structure

3. The Electron Cloud Model
The electron cloud is the space around the nucleus of an atom where the atom’s electrons are most likely to be found.
After much research, today we realize that E.L. are not neat, planet like orbits
Instead, they are spherical regions in space
Section 2.2

4. The Electron Cloud Model (cont.)
Each energy level can hold a limited number of electrons.
First energy level—2 electrons
Second energy level—8 electrons
Third energy level—18 electrons
Section 2.2

5. The Electron Cloud Model (cont.)
The electrons in the outermost energy level are called valence electrons.
Section 2.2

6. Click a hyperlink to view the corresponding slides.
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Section 7.1 Present-Day Atomic Theory
Section 7.2 The Periodic Table and Atomic Structure
Click a hyperlink to view the corresponding slides.
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7. Present-Day Atomic Theory
Relate emission spectra to the electron configurations of atoms.
Relate energy sublevels and orbitals within the atom.
Section 7.1

8. Present-Day Atomic Theory
electromagnetic spectrum: the entire range of electromagnetic radiation
Section 7.1

9. Present-Day Atomic Theory
sublevel
aufbau principle
Heisenberg uncertainty principle
orbital
electron configuration
Electrons are organized in energy levels and sublevels.
Section 7.1

10. Expanding the Model of the Atom
Danish physicist Niels Bohr suggested that electrons revolve around the nucleus like planets revolve around the Sun.
Section 7.1

11. Expanding the Model of the Atom (cont.)
Bohr’s model explained hydrogen’s emission spectrum, but failed to explain any other element’s emission spectrum.
The current atomic model explains electron behavior by interpreting the emission spectra of all the elements.
Video: Bohr's Model of Atom
Section 7.1

12. Expanding the Model of the Atom (cont.)
The periodic table reflects each element’s electron arrangement.
Section 7.1

13. Expanding the Model of the Atom (cont.)
To calculate the exact amount of energy released by the electrons in atoms:
the higher the frequency of a wave and the shorter wavelength, the greater the energy of the radiation
the lower the frequency and the longer the wavelength, the lower the energy released
Section 7.1

14. Expanding the Model of the Atom (cont.)
Section 7.1

15. Expanding the Model of the Atom (cont.)
By absorbing a specific amount of energy, an electron can jump to a higher energy level.
When the electron falls back to the lower energy level, it releases the same amount of energy in the form of radiation with a definite frequency.
Section 7.1

16. Expanding the Model of the Atom (cont.)
Section 7.1

17. Expanding the Model of the Atom (cont.)
Energy levels within an atom are characteristic of each element.
Section 7.1

18. Expanding the Model of the Atom (cont.)
Closely spaced lines in a spectrum suggest that sublevels—divisions within a level—exist within an energy level.
The energy sublevels are indicated as s, p, d, or f.
Section 7.1

19. Electrons Distribution in Energy Levels
Each energy level has a specific number of sublevels, which is the same number of the energy level.
Bohr & Hesinberg’s Atomic Models 9 min Lecture: Modern Structure of the Atom
Section 7.1

20. Electrons Distribution in Energy Levels (cont.)
Section 7.1

21. Expanding the Model of the Atom (cont.)
The aufbau principle states that each electron occupies the lowest energy sublevel available.
Section 7.1

23. Electrons Distribution in Energy Levels (cont.)
The Heisenburg uncertainty principle states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.
Section 7.1

24. Electrons Distribution in Energy Levels (cont.)
The only quantity that can be known is the probability for an electron to occupy a certain region around the nucleus.
Section 7.1

25. Electrons Distribution in Energy Levels (cont.)
The space in which there is a high probability of finding an electron is called an orbital.
Each energy level relates to orbitals of different sizes and shapes.
Section 7.1

26. Electrons Distribution in Energy Levels (cont.)
Section 7.1

27. Electrons Distribution in Energy Levels (cont.)
The most stable arrangement of electrons in sublevels and orbitals is called an electron configuration.
Because orbitals are mostly empty space, this space can be used by another pair of electrons and the orbitals can overlap.
Section 7.1

28. Hund’s Rule: When you get onto a bus, where do you sit?
This is what electron’s do in orbitals of a sublevel.
One electron in each orbital and then they will start to double up until the sub level is full.
Example:

29. Section Assessment
Atoms move in circular orbits in which atomic model?
A. quantum mechanical model
B. Rutherford’s model
C. Bohr’s model
D. de Broglie’s model
Section 7.1

30. Section Assessment Electron behavior is explained by interpreting ___.
A. emission spectra
B. the number of energy levels
C. how quickly the electrons revolve around the nucleus
D. the number of covalent bonds
Section 7.1

31. End of Section 7.1

32. The Periodic Table and Atomic Structure
Distinguish the s, p, d, and f blocks on the periodic table and relate them to an element’s electron configuration.
Predict the electron configurations of elements using the periodic table.
Section 7.2

33. The Periodic Table and Atomic Structure
orbital: space in which there is a high probability of finding an electron
Section 7.2

34. The Periodic Table and Atomic Structure
inner transition element
A predictable pattern can be used to determine electron arrangement in an atom.
Section 7.2

35. Patterns of Atomic Structure
The shape of the modern periodic table is a direct result of the order in which electrons fill energy sublevels and their orbitals.
Section 7.2

37. Building Electron Configurations
Chemical properties repeat when elements are arranged by atomic number because electron configurations repeat in a certain pattern.
Section 7.2

38. Electron Configurations
The electron configuration of an atom is a shorthand method of writing the location of electrons by sublevel.
The sublevel is written followed by a superscript with the number of electrons in the sublevel.
If the 2p sublevel contains 2 electrons, it is written 2p2

39. Writing Electron Configurations
First, determine how many electrons are in the atom. Iron has 26 electrons.
Arrange the energy sublevels according to increasing energy:
1s 2s 2p 3s 3p 4s 3d …
Fill each sublevel with electrons until you have used all the electrons in the atom:
Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d 6
The sum of the superscripts equals the atomic number of iron (26)

40. Building Electron Configurations (cont.)
Section 7.2

41. Noble Gas Core Electron Configurations
Recall, the electron configuration for Na is:
Na: 1s2 2s2 2p6 3s1
We can abbreviate the electron configuration by indicating the innermost electrons with the symbol of the preceding noble gas.
The preceding noble gas with an atomic number less than sodium is neon, Ne. We rewrite the electron configuration:
Na: [Ne] 3s1

42. Building Electron Configurations (cont.)
Noble gas notation uses symbols in brackets to shorten inner electron configurations of other elements.
The stable electron configurations explain the lack of reactivity of the noble gases.
Section 7.2

43. Building Electron Configurations (cont.)
Section 7.2

44. Practice

45. Building Electron Configurations (cont.)
Experimental evidence indicates that the 4s and 3d sublevels are close in energy, with the 4s sublevel having a slightly lower energy.
Section 7.2

46. Building Electron Configurations (cont.)
The transition elements are any of the elements in group 3 through 12 of the periodic table, all of which are metals.
They lose electrons to attain a more stable configuration.
Most have multiple oxidation numbers because their s and d orbitals are so close in energy that electrons can be lost from both orbitals.
Section 7.2

47. Building Electron Configurations (cont.)
The lanthanides and actinides are called the inner transition elements because their last electron occupies a 5f orbital in the seventh period.
Section 7.2

48. The Size of the Orbitals
The higher the energy, the farther the outermost electrons are from the nucleus.
As the valence electron gets farther from the nucleus, the s orbital it occupies gets larger and larger.
Section 7.2

49. Section Assessment
As the valence electron gets farther from the nucleus, its s orbital becomes:
A. positively charged
B. negatively charged
C. larger
D. smaller
Section 7.2

50. Section Assessment
The most common oxidation number for the inner transition elements is:
A. 1+
B. 2+
C. 3+
D. 4+
Section 7.2

51. End of Section 7.2

52. Standardized Test Practice Image Bank Concepts in Motion
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Study Guide
Chapter Assessment
Standardized Test Practice
Image Bank
Concepts in Motion
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53. Key Concepts
The position of an element in the periodic table reveals the number of valence electrons the element has.
The outermost valence electrons determine the properties of an element.
Electrons are found only in levels of fixed energy in an atom.
Energy levels have sublevels. Each sublevel can hold a specific number of electrons.
Sublevels can be subdivided into s, p, d, and f orbitals. Sublevels hold 2, 6, 10, and 14 electrons respectively.
Study Guide 1

54. Key Concepts
The organization of the periodic table reflects the electron configurations of the elements.
The active metals occupy the s region of the periodic table. Metals, metalloids, and nonmetals fill the p block.
Within a period of the periodic table, the number of valence electrons for main group elements increases from one to eight.
The transition elements, groups 3 to 12, occupy the d block of the periodic table. These elements can have valence electrons in both s and d sublevels.
The lanthanides and actinides, called the inner transition elements, occupy the f block of the periodic table. Their valence electrons are in s and f sublevels.
Study Guide 2

55. By absorbing a specific amount of energy, an electron will:
A. revolve in a circular orbit around the nucleus
B. fall back to a lower energy level
C. jump to a higher energy level
D. remain the same energy level
Chapter Assessment 1

56. Groups of fine lines in an emission spectra indicate that electrons are moving between energy levels that are:
A. close in energy
B. far apart
C. on the same level
Chapter Assessment 2

57. The aufbau principle states that each electron occupies the ___ energy orbital available.
A. lowest
B. highest
C. same
Chapter Assessment 3

58. Spherical orbitals belong to which sublevel? A. s B. p C. d D. f
Chapter Assessment 4

59. What is the maximum number of electrons the 1s orbital can hold? A. 10
Chapter Assessment 5

60. Which atomic suborbitals have a “dumbbell” shape? A. s B. f C. p D. d
STP 1

61. A. the Pauli exclusion principle B. the aufbau principle
It is impossible to know precisely both the location and velocity of an electron at the same time because:
A. the Pauli exclusion principle
B. the aufbau principle
C. electrons travel in waves
D. the Heisenberg uncertainty principle
STP 2

62. Orbitals can overlap because: A. atoms are spherical in shape
B. most of the orbital is empty space
C. the electron configuration of the atom
D. the aufbau principle
STP 3

63. Which group of elements is more likely to have multiple oxidation numbers?
A. group 1
B. group 2
C. group 12
D. group 18
STP 4

64. Who suggested that electrons revolve around an atom’s nucleus much like planets revolve around the Sun?
A. Dalton
B. Rutherford
C. Bohr
D. Heisenberg
STP 5

65. Click on an image to enlarge.
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66. IB 1

67. IB 2

68. IB 3

69. IB 4

70. IB 5

71. IB 6

72. IB 7

73. IB 8

74. IB 9

75. IB 10

76. IB 11

77. IB 12

78. IB 13

79. IB 14

80. IB 15

81. Table 7.3 The Electron Configurations of the Noble Gases
Figure 7.2 Structure of the atom
CIM

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