PERFORMANCE OBJECTIVES / Chem. B1A - Dr. Daniel

Chapter 8 – Electron Configurations and Periodic Properties

 At the end of this unit you should be able to:

1. Explain how the Pauli Exclusion Principle and Hund’s rule predicts whether a substance is paramagnetic or diamagnetic. Use an orbital diagram and the number of unpaired electrons to compare the degree of paramagnetism in different elements or ions.

2. How does the Pauli Exclusion Principle limit the number of electrons in an orbital. State how many electrons can occupy the different subshells (s, p, d, f, g, etc) and different shells.

3. Draw an orbital diagram showing how electrons occupy orbitals of an element or its ion.

4. Given its atomic number (any element up to 56 and any main group element), write the electron configuration (long or short version) of an element (or its ion), or be able to identify any element given its electron configuration.

5. Describe the general shape and geometry of s, p, or d orbital.

6. Compare the electron configuration of an "excited" atom to be atom in the ground state.

7. Describe the electron behavior and energy changes in an atom during transitions between ground and excited states.

8. Define core electrons, valence electrons and valence electronic configuration. Which electrons are involved in bonding and why?

9. Explain the structure of the periodic table using electron configurations or orbital diagrams. Draw a periodic table and label the following areas, representative or main group elements, transition elements, inner transition elements, s, p, d, and f blocks, Lanthanide and Actinide elements.

10. How was Mendeleev able to predict properties of elements that had not yet been discovered? How did the discovery of the number of protons in a nucleus modify the periodic law?

11. Define effective nuclear charge and calculate the approximate nuclear charge for any representative (main group) element.

12. Define the bonding atomic radius. Explain how the atomic radius varies as you go across and down the periodic table. Use effective nuclear charge and the number of electron shells to explain this variation.

13. Define ionization energy and electron affinity. Use the effective nuclear charge and radius to explain periodic trends in these two properties. Explain how the gain or loss of electrons changes the ionization energy of an atomic particle.

14. Given a list of elements, rank in order of increasing radius and ionization energy.

15. Explain why there is a large increase in ionization energy when you try to remove a core electron.

16. Describe how metallic character varies as you go down and to the left on the periodic table.