12. Draw Lewis-dot structures for NH3, H2O, Cl2O, C2H4, and SiO2. Many compounds may be used for this question. Avoid "problem" molecules such as those with an odd number of valence electrons that cannot follow the octet rule (such as NO), at least initially. When teaching students how to draw Lewis-dot structures, a useful technique is to develop a helpful set of rules, such as these:

a. Count the total number of valence electrons in the structure.

b. For ions, add one electron for each negative charge and subtract one electron for each positive charge.

c. Draw the skeleton using dashes to represent electron pairs joining two atoms together until the skeleton is complete.

d. Add dot-pairs until all valence electrons are accounted for and each atom has an octet of electrons (duet for hydrogen).

e. If Step d is impossible when N, C, O, or S are involved, try double or triple bonds (two pairs or three pairs of dots) to form octets.

A simple, quick method for drawing Lewis-dot structures is presented in Tips for the Teacher under problem solving. It offers as an alternative to the rules given above or when students have developed some facility drawing these structures.

13. Compare the covalent bonds formed between elements of similar electronegativity such as carbon and hydrogen and covalent bonds formed between elements with significant difference in electronegativity, such as hydrogen and chlorine. [Student answers will probably vary considerably. At minimum, some reference to equal or unequal electron sharing should be made. Students should note that unequal sharing produces a charge separation. In any case, students should point out that polar covalent bonds, depending on molecular geometry, often give a molecule properties that affect its behavior.]

14. Describe the bonding trend expected when fluorine bonds with each element in the second row of the periodic table, including itself-F2, OF2, NF3, CF4, BF3, BeF2, and LiF. [Student answers will vary. At minimum, students should recognize that a variation in bond type from homonuclear covalent to essentially ionic takes place. In their answers, students may note the arbitrary nature of deciding at which point polar covalent bonds are better regarded as ionic.]

15. What happens to the system's total potential energy when two isolated atoms capable of bonding come into close proximity? [As atoms approach, the atomic nuclei are attracted to the valence electrons. If both atoms have half-filled or empty valence orbitals, then bonding may occur, lowering the potential energy of the system due to attractive forces reducing the separation between atoms. See Figure 11.]

Figure 11. Potential Energy Diagram.

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