Section 6.8-6.9
1. How are electron configurations made?
2. How are orbital diagrams produced?
3. How are electrons paired?
4. Describe Hund’s Rule.
5. Draw the orbital diagram for the electron configuration of oxygen, atomic number 8. How many unpaired
electrons does an oxygen atom possess?
6. How many of the elements in the second row of the periodic table (Li through Ne) will have at least one
unpaired electron in their electron configurations?
7. How are “condensed” electron configurations competed?
8. What are core electrons and what are valence electrons?
9. What should the periodic table actually look like?
10. Where are the s, p, d, and f blocks located?
11. What is the characteristic valence electron configuration of the 7A elements, the halogen?
12. (a) Based on its position in the periodic table, write the condensed electron configuration for bismuth,
element 83. (b) How many unpaired electrons does a bismuth atom have?
13. A certain atom has an [noble gas]5s24d105p4 electron configuration. Which element is it?
14. What causes the violations of the electron configurations.
15. The accompanying drawing shows part of the orbital diagram for an element. (a) As drawn, the drawing is
incorrect. Why? (b) How would you correct the drawing without changing the number of electrons? (c) To
which group in the periodic table does the element belong?
16. State where in the periodic table these elements appear:
a. elements with the valence-shell electron configuration ns2np5
b. elements that have three unpaired p electrons
c. an element whose valence electrons are 4s24p1
d. the d-block elements
17. For a given value of the principal quantum number, n, how do the energies of the s, p, d, and f subshells
vary for (a) hydrogen, (b) a many-electron atom?
18. (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?
19. (a) What experimental evidence is there for the electron having a “spin”? (b) Draw an energy-level diagram
that shows the relative energetic positions of a 1s orbital and a 2s orbital. Put two electrons in the 1s orbital.
(c) Draw an arrow showing the excitation of an electron from the 1s to the 2s orbital.
20. (a) State the Pauli exclusion principle in your own words. (b) The Pauli exclusion principle is, in an
important sense, the key to understanding the periodic table. (c) Explain.
21. What is the maximum number of electrons that can occupy each of the following subshells? (a) 3p, (b) 5d,
(c) 2s, (d) 4f.
22. What is the maximum number of electrons in an atom that can have the following quantum numbers? (a) n
= 3, ml = -2; (b) n = 4, l = 3; (c) n = 5, l = 3, ml = 2, (d) n = 4, l = 1, ml = 0.
23. (a) What are “valence electrons”? (b) What are “core electrons”? (c) What does each box in an orbital
diagram represent? (d) What quantity is represented by the half arrows in an orbital diagram?
24. For each element, indicate the number of valence electrons, core electrons, and unpaired electrons in the
ground state: (a) nitrogen, (b) silicon, (c) chlorine.
25. Write the condensed electron configurations for the following atoms, using the appropriate noble-gas core
abbreviations: (a) Cs, (b) Ni, (c) Se, (d) Cd, (e) U, (f) Pb.
26. Write the condensed electron configurations for the following atoms and indicate how many unpaired
electrons each has: (a) Mg, (b) Ge, (c) Br, (d) V, (e) Y, (f) Lu.
27. Identify the specific element that corresponds to each of the following electron configurations and indicate
the number of unpaired electrons for each: (a) 1s22s2, (b) 1s22s22p4, (c) [Ar]4s13d5, (d) [Kr]5s24d105p4.
28. What is wrong with the following electron configurations for atoms in their ground states? (a) 1s22s23s1, (b)
[Ne]2s22p3, (c) [Ne]3s23d5.