1. Section 4.3 “Electron Configurations”
Pre-AP Chemistry

2. Chapter 4 Vocabulary
20. Electron Configuration 21. Aufbau principle 22. Pauli exclusion principle 23. Hund’s rule 24. Noble gas 25. Noble-gas configuration

3. Chapter 4 Notes
Electron Configuration Rules
J. Scientists

4. Electron Configuration Blocks

5. Chapter 4 Lesson Starter
Section 3 Electron Configurations
Chapter 4
Lesson Starter
The electron configuration of carbon is 1s22s22p2.
An electron configuration describes the arrangement of electrons in an atom.
The integers indicate the main energy level of each orbital occupied by electrons.
The letters indicate the shape of the occupied orbitals.
The superscripts identify the number of electrons in each sublevel.

6. Section 3 Electron Configurations
Chapter 4
Objectives
List the total number of electrons needed to fully occupy each main energy level.
State the Aufbau principle, the Pauli exclusion principle, and Hund’s rule.
Describe the electron configurations for the atoms of any element using orbital notation, electron-configuration notation, and, when appropriate, noble-gas notation.

7. Section 4.3 Electron Configurations
OBJECTIVES:
List the total number of electrons needed to fully occupy each main energy level

8. Electron Configurations
Section 3 Electron Configurations
Chapter 4
Electron Configurations
The arrangement of electrons in an atom is known as the atom’s electron configuration.
The lowest-energy arrangement of the electrons for each element is called the element’s ground- state electron configuration.

9. Orbital Location on Periodic Table

10. Orbital Location on Periodic Table

12. Relative Energies of Orbitals
Section 3 Electron Configurations
Chapter 4
Relative Energies of Orbitals

13. Electron Configuration
Section 3 Electron Configurations
Chapter 4
Electron Configuration
Click below to watch the Visual Concept.
Visual Concept

14. Section 4.3 Electron Configurations
OBJECTIVES:
State the Aufbau principle, the Pauli exclusion principle, and Hund’s rule.

15. Section 5.2 Electron Arrangement in Atoms p. 133
Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f
aufbau diagram - page 105
Aufbau is German for “building up”

16. Rules Governing Electron Configurations
Section 3 Electron Configurations
Chapter 4
Rules Governing Electron Configurations
According to the Aufbau principle, an electron occupies the lowest-energy orbital that can receive it.
According to the Pauli exclusion principle, no two electrons in the same atom can have the same set of four quantum numbers.
According to Hund’s rule, orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin state.

17. Chapter 4 Aufbau Principle Section 3 Electron Configurations
Click below to watch the Visual Concept.
Visual Concept

18. Chapter 4 Notes
I. Electron Configuration Rules 1. Aufbau Principle 2. Pauli’s Exclusion Principle 3. Hund’s Rule

19. Electron Configurations…
…are the way electrons are arranged in various orbitals around the nuclei of atoms. Three rules tell us how:
Aufbau principle – e- enter the lowest energy first.
This causes difficulties because of the overlap of orbitals of different energies – follow the diagram!
Pauli Exclusion Principle - at most 2 e- per orbital - different spins

20. Pauli Exclusion Principle
No two electrons in an atom can have the same four quantum numbers.
To show the different direction of spin, a pair in the same orbital is written as:
Wolfgang Pauli

21. Quantum Numbers
Each electron in an atom has a unique set of 4 quantum numbers which describe it.
Principal quantum number
Angular momentum quantum number
Magnetic quantum number
Spin quantum number

22. Electron Configurations
Hund’s Rule- When e- occupy orbitals of equal energy, they don’t pair up until they have to.
Let’s write the electron configuration for Phosphorus
We need to account for all 15 electrons in phosphorus

23. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
The first 2 e-’s go into the 1s orbital
Notice opposite direction of spins

24. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s
The next e-’s go in 2s orbital

25. Increasing energy The next e-’s go in 2p orbital 7p 6d 5f 7s 6p 5d 6s

26. Increasing energy The next e-’s go in 3s orbital 7p 6d 5f 7s 6p 5d 6s

27. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
The last 3 e-’s go in 3p orbitals
They each go into separate shapes (Hund’s)
3 unpaired e-’s
= 1s22s22p63s23p3
Orbital notation

28. Section 5.2 Electron Configurations
OBJECTIVES:
Describe the electron configurations for the atoms of any element using orbital notation, electron-configuration notation, and, when appropriate, noble-gas notation.

29. Representing Electron Configurations
Section 3 Electron Configurations
Chapter 4
Representing Electron Configurations
Orbital Notation
An unoccupied orbital is represented by a line, with the orbital’s name written underneath the line.
An orbital containing one electron is represented as:

30. Representing Electron Configurations, continued
Section 3 Electron Configurations
Chapter 4
Representing Electron Configurations, continued
Orbital Notation
An orbital containing two electrons is represented as:
The lines are labeled with the principal quantum number and sublevel letter. For example, the orbital notation for helium is written as follows: 1s He

31. Chapter 4 Orbital Notation Section 3 Electron Configurations
Click below to watch the Visual Concept.
Visual Concept

32. Representing Electron Configurations, continued
Section 3 Electron Configurations
Chapter 4
Representing Electron Configurations, continued
Electron-Configuration Notation
Electron-configuration notation eliminates the lines and arrows of orbital notation.
Instead, the number of electrons in a sublevel is shown by adding a superscript to the sublevel designation.
The helium configuration is represented by 1s2.
The superscript indicates that there are two electrons in helium’s 1s orbital.

33. Reading Electron-Configuration Notation
Section 3 Electron Configurations
Chapter 4
Reading Electron-Configuration Notation
Click below to watch the Visual Concept.
Visual Concept

34. Sec 4.3 Practice Problems
Textbook: A. 1-2 (p 107)

35. Representing Electron Configurations, continued
Section 3 Electron Configurations
Chapter 4
Representing Electron Configurations, continued
Sample Problem A
The electron configuration of boron is 1s22s22p1. How many electrons are present in an atom of boron? What is the atomic number for boron? Write the orbital notation for boron.

36. Section 3 Electron Configurations
Chapter 4
Representing Electron Configurations, continued Sample Problem A Solution
The number of electrons in a boron atom is equal to the sum of the superscripts in its electron-configuration notation: = 5 electrons. The number of protons equals the number of electrons in a neutral atom. So we know that boron has 5 protons and thus has an atomic number of 5. To write the orbital notation, first draw the lines representing orbitals. 1s 2s 2p

37. Section 3 Electron Configurations
Chapter 4
Representing Electron Configurations, continued Sample Problem A Solution, continued
Next, add arrows showing the electron locations. The first two electrons occupy n = 1 energy level and fill the 1s orbital. 1s 2s 2p

38. Section 3 Electron Configurations
Chapter 4
Representing Electron Configurations, continued Sample Problem A Solution, continued
The next three electrons occupy the n = 2 main energy level. Two of these occupy the lower-energy 2s orbital. The third occupies a higher-energy p orbital. 1s 2s 2p

39. Elements of the Second Period
Section 3 Electron Configurations
Chapter 4
Elements of the Second Period
In the first-period elements, hydrogen and helium, electrons occupy the orbital of the first main energy level.
According to the Aufbau principle, after the 1s orbital is filled, the next electron occupies the s sublevel in the second main energy level.

40. Elements of the Second Period, continued
Section 3 Electron Configurations
Chapter 4
Elements of the Second Period, continued
The highest-occupied energy level is the electron- containing main energy level with the highest principal quantum number.
Inner-shell electrons are electrons that are not in the highest-occupied energy level.

41. Writing Electron Configurations
Section 3 Electron Configurations
Chapter 4
Writing Electron Configurations

42. Notecard: Aufbau Diagram
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p6

43. Elements of the Third Period
Section 3 Electron Configurations
Chapter 4
Elements of the Third Period
After the outer octet is filled in neon, the next electron enters the s sublevel in the n = 3 main energy level.
Noble-Gas Notation
The Group 18 elements (helium, neon, argon, krypton, xenon, and radon) are called the noble gases.
A noble-gas configuration refers to an outer main energy level occupied, in most cases, by eight electrons.

44. Orbital Notation for Three Noble Gases
Section 3 Electron Configurations
Chapter 4
Orbital Notation for Three Noble Gases

45. Elements of the Fourth Period
Section 3 Electron Configurations
Chapter 4
Elements of the Fourth Period
The period begins by filling the 4s orbital, the empty orbital of lowest energy.
With the 4s sublevel filled, the 4p and 3d sublevels are the next available vacant orbitals.
The 3d sublevel is lower in energy than the 4p sublevel. Therefore, the five 3d orbitals are next to be filled.

46. Orbital Notation for Argon and Potassium
Section 3 Electron Configurations
Chapter 4
Orbital Notation for Argon and Potassium

47. Elements of the Fifth Period
Section 3 Electron Configurations
Chapter 4
Elements of the Fifth Period
In the 18 elements of the fifth period, sublevels fill in a similar manner as in elements of the fourth period.
Successive electrons are added first to the 5s orbital, then to the 4d orbitals, and finally to the 5p orbitals.

48. Sec 4.3 Practice Problems
Textbook: A. 1-2 (p 107) B. 1-4 (p. 115)

49. Chapter 4 Sample Problem B
Section 3 Electron Configurations
Chapter 4
Sample Problem B
a. Write both the complete electron-configuration notation and the noble-gas notation for iron, Fe.
b. How many electron-containing orbitals are in an atom of iron? How many of these orbitals are completely filled? How many unpaired electrons are there in an atom of iron? In which sublevel are the unpaired electrons located?

50. Sample Problem B Solution
Section 3 Electron Configurations
Chapter 4
Sample Problem B Solution
a. The complete electron-configuration notation of iron is 1s22s22p63s23p63d64s2. Iron’s noble-gas notation is [Ar]3d64s2.
b. An iron atom has 15 orbitals that contain electrons.
They consist of one 1s orbital, one 2s orbital, three 2p orbitals, one 3s orbital, three 3p orbitals, five 3d orbitals, and one 4s orbital.
Eleven of these orbitals are filled, and there are four unpaired electrons.
They are located in the 3d sublevel.
The notation 3d6 represents 3d

51. Sec 4.3 Practice Problems
Textbook: A. 1-2 (p 107) B. 1-4 (p. 115) C. 1-2 (p. 116)

52. Chapter 4 Sample Problem C
Section 3 Electron Configurations
Chapter 4
Sample Problem C
a. Write both the complete electron-configuration notation and the noble-gas notation for a rubidium atom.
b. Identify the elements in the second, third, and fourth periods that have the same number of highest-energy-level electrons as rubidium.

53. Sample Problem C Solution
Section 3 Electron Configurations
Chapter 4
Sample Problem C Solution
a. 1s22s22p63s23p63d104s24p65s1, [Kr]5s1
b. Rubidium has one electron in its highest energy level (the fifth). The elements with the same outermost configuration are,
in the second period, lithium, Li;
in the third period, sodium, Na;
and in the fourth period, potassium, K.

54. Exceptions to the Aufbau Principle
Lowest energy to higher energy.
Adding electrons can change the energy of the orbital. Full orbitals are the absolute best situation.
However, half filled orbitals have a lower energy, and are next best
Makes them more stable.
Changes the filling order

55. Write the electron configurations for these elements:
Titanium - 22 electrons
1s22s22p63s23p64s23d2
Vanadium - 23 electrons
1s22s22p63s23p64s23d3
Chromium - 24 electrons
1s22s22p63s23p64s23d4 (expected)
But this is not what happens!!

56. Chromium is actually: 1s22s22p63s23p64s13d5 Why?
This gives us two half filled orbitals (the others are all still full)
Half full is slightly lower in energy.
The same principal applies to copper.

57. Copper’s electron configuration
Copper has 29 electrons so we expect: 1s22s22p63s23p64s23d9
But the actual configuration is:
1s22s22p63s23p64s13d10
This change gives one more filled orbital and one that is half filled.
Remember these exceptions: d4, d9

58. Irregular configurations of Cr and Cu
Chromium steals a 4s electron to make its 3d sublevel HALF FULL
Copper steals a 4s electron to FILL its 3d sublevel

59. Sec 4.3 Practice Problems
Textbook: A. 1-2 (p 107) B. 1-4 (p. 115) C. 1-2 (p. 116)

60. Chapter 4 Notes J. Scientists Niels Bohr Max Planck Albert Einstein
Louis de Broglie
Werner Heisenberg
Erwin Schrödinger

61. End of Chapter 4