Our understanding of the structure of an atom has come a long way in that last century. Now that we have connected the wave/particle duality of matter, we can better understand the organization and position of electrons using quantum atomic theory. Today we will begin that organization with orbital diagrams and electron configurations.
- Draw an energy-level diagram for the orbitals in a many-electron atom and describe how electrons populate the orbitals in the ground state of an atom, using the Pauli exclusion principle and Hund’s rule. (6.8)
- Use the periodic table to write condensed electron configurations and determine the number of unpaired electrons in an atom. (6.9)
As we begin this topic, make sure to skim the text and take down notes. The key ideas are presented in the slides below, which correspond to sections 6.8-6.9 of the text.
Now that you've seen the content, work out the sample and practice problems below (sample problems 6.7 - 6.9 of the text).
To wrap up this mission, work out the mastery problems below on a separate sheet of paper, then turn in all of your stamped work on OneNote. These problems come from the back of chapter 6 in your book.
(a) The average distance from the nucleus of a 3s electron in a chlorine atom is smaller than that for a 3p electron. In light of this fact, which orbital is higher in energy?
(b) Would you expect it to require more or less energy to remove a 3s electron from the chlorine atom, as compared with a 2p electron?
6.66 For each element, count the number of valence electrons, core electrons, and unpaired electrons in the ground state:
6.74 The following electron configurations represent excited states. Identify the element, and write its ground-state condensed electron configuration.
6.97 Using only a periodic table as a guide, write the condensed electron configurations for the following atoms:
2007b FRQ #2
Electron configurations take a while to get the hang of. Keep reviewing this material though. It's worth it when things click!