Structures and Properties Unit Test
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For multiple choice, select the best answer (1 mark each)
1. Which of the following elements could have the valence orbital configuration of s2p4?
a. Carbon
d. sulfur
b. Nitrogen
e. neon
c. Boron
2. First ionization energy is
a. The amount of energy an electron must gain to escape from a neutral atom
b. The amount of energy an electron must lose to escape from a neutral atom
c. The amount of energy an atom gains when it first becomes an ion
d. Two of these are correct
e. All of these are correct
3. In which period can you start to use condensed electron configurations?
a. 1
d. 4
b. 2
e. 5
c. 3
4. When comparing the properties of carbon atoms and oxygen atoms, which statements are
correct?
a. Carbon and oxygen have the same spin quantum number, ms.
b. Carbon and oxygen have the same magnetic quantum number, ml
c. Carbon and oxygen have the same principal quantum number, n.
d. Carbon has a greater orbital-shape quantum number, l, than oxygen.
e. Oxygen has a greater orbital-shape quantum number, l, than carbon.
5. Chemical bonding can best be predicted and explained by the properties of
a. Valence electrons
d. the nuclear charge
b. Inner core electrons
e. the atomic radius
c. The effective nuclear charge
6. Which of the following is not a property of ionic compounds?
a. Conducts electric current in solution
d. ductile
b. Does not conduct electric current as a solid
e. brittle
c. None of these answers
7. Which type of hybrid orbitals is likely to exist for the sulfur atom in the bonds for the molecule
sulfur hexafluoride, SF6?
a. sp3d2
d. sp2
b. sp3d
e. sp
c. sp3
8. Electrons are said to be “delocalized” in metallic bonds.
a. In your own words, describe what this means, and how it is different from the locations
of valence electrons in ionic and covalent bonds. (3 marks)
b. How does the “delocalization” of electrons in metallic solids explain “malleability”, the
property unique to all metals? (2 marks)
9. Explain why ionic solids are brittle. Use a diagram to assist your answer. (3 marks)
10. Determine whether each of the following compounds will be polar or non-polar. Explain (2 )
a. CO2
b. PCl3
11. Use VSEPR theory and Lewis structures to predict the number of bonding pairs and lone pairs
around the central atom so that you can identify the VSEPR shape for: (6 marks)
a. XeF2
c. SF5+
b. SO2
12. For XeF2, predict the orbital hybridization of the Xe atom. (2 marks)
13. Which compound in each of the following pairs has the higher boiling point? Explain your
choice in each case. (4 marks)
a. NH3 or PH3
b. C2H6 or C4H10
14. If the last electron to be added to an atom was the seventh electron in the third energy level
A. what element would that be? (1 mark)
B. What would the 4 digit quantum ‘address’ of that electron be? (4 marks)
15. Answer the following questions about the element antimony.
a. Write the condensed electron configuration for antimony. (1 mark)
b. Antimony can have a valence of -3, +3, or +5. Use condensed electron configurations to
show how each of these valences must form. In each case, explain why that change
makes sense. (3 marks)
c. Antimony is a metalloid. Defend this claim using your answers to part ‘b’ of this
question. (1 mark)
16. Your friend is having a hard time understanding condensed electron configurations. Write an email, giving one example of a condensed electron configuration, explaining why they are used,
and helping to avoid any possible misconceptions. (2 marks)
17. The success of a theory or model is judged by how well it can explain and predict experimental
findings. The Bohr model was at the same time highly successful and highly unsuccessful.
Explain 2 successes and 2 shortcomings of the model. (4 marks)