CK-12-Chemistry Intermediate

http://www.ck12.org Chapter 5. Electrons in Atoms

TABLE5.8:(continued)

Element Name Symbol Atomic Number Electron Configuration

Lawrencium Lr 103 [Rn] 5 f^147 s^27 p^1

Rutherfordium Rf 104 [Rn] 5 f^146 d^27 s^2

Dubnium Db 105 [Rn] 5 f^146 d^37 s^2

Seaborgium Sg 106 [Rn] 5 f^146 d^47 s^2

Bohrium Bh 107 [Rn] 5 f^146 d^57 s^2

Hassium Hs 108 [Rn] 5 f^146 d^67 s^2

Meitnerium Mt 109 [Rn] 5 f^146 d^77 s^2

Lesson Summary

• The ground state electron configurations of most atoms can be predicted based on the Aufbau principle, the
  Pauli exclusion principle, and Hund’s rule.


• Orbital filling diagrams are drawn to show how electrons fill up various energy levels. In these diagrams,
  orbitals are arranged from lowest to highest energy.


• The electron configuration is unique for each element.


• Valence electrons are the electrons in the outermost principal energy level. An atom can have a maximum of
  eight valence electrons.


• Unpaired electrons influence the magnetic properties of an atom.


• Some electron configurations, such as those of chromium and copper, do not strictly follow the Aufbau
  principle.

Lesson Review Questions

Reviewing Concepts

1. Electron configurations:

a. What do the superscripts in an electron configuration represent?

b. What is the atomic number of an element with the electron configuration 1s^22 s^22 p^63 s^23 p^63 d^104 s^24 p^5?

2. Arrange these sublevels in order of increasing energy: 3p, 5s, 4f, 2s, 3d.


3. Which rule is violated by the electron configuration 1s^22 s^22 p^63 s^23 p^63 d^2? Explain.


4. Which rule is violated by each of the orbital filling diagrams below? Explain.

a.

b.