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Chapter 9: Problem 103

Oxygen forms three types of ionic compounds in which the anions are oxide \(\left(\mathrm{O}^{2-}\right),\) peroxide \(\left(\mathrm{O}_{2}^{2-}\right),\) and superoxide \(\left(\mathrm{O}_{2}^{-}\right) .\) Draw Lewis structures of these ions.

Short Answer

Expert verified
The Lewis structures are drawn by writing 'O' for each oxygen atom, representing its six valence electrons as dots around it and finally adding more dots for additional electrons it gains in each ion.

Step by step solution

01

Lewis Structure for Oxide Ion (\(\mathrm{O}^{2-}\))

The oxidation number of Oxygen is -2. Oxygen has 6 valence electrons and it needs 2 more electrons to complete its octet. So, in the Oxide ion, the Oxygen gains 2 extra electrons. One can draw the Lewis structure by representing these 8 electrons as dots around the symbol 'O'.
02

Lewis Structure for Peroxide Ion (\(\mathrm{O}_{2}^{2-}\))

The peroxide ion has a -2 charge. It consists of two oxygen atoms and each oxygen atom effectively 'owns' eight electrons in their valence shell. Deciding the connectivity of atoms within the molecule, we can see two oxygen atoms are bonded together. And you draw it as an O with six dots around it, attached by a bond (as a line) to another O also with six dots around it.
03

Lewis Structure for Superoxide Ion (\(\mathrm{O}_{2}^{-}\))

In the superoxide anion, one of the oxygen atoms donates one of its electrons to another oxygen atom. The resulting ion has an uneven distribution of electrons, giving it an overall charge of -1. You can draw the Lewis structure by representing these electrons as dots around the 'O' symbols.

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Most popular questions from this chapter

The amide ion, \(\mathrm{NH}_{2}^{-}\), is a Bronsted base. Represent the reaction between the amide ion and water in terms of Lewis structures.

Summarize the essential features of the Lewis octet rule. The octet rule applies mainly to the second period elements. Explain.

Using this information: $$ \begin{array}{rr} \mathrm{C}(s) \longrightarrow \mathrm{C}(g) & \Delta H_{\mathrm{rxn}}^{\circ}=716 \mathrm{~kJ} / \mathrm{mol} \\ 2 \mathrm{H}_{2}(g) \longrightarrow 4 \mathrm{H}(g) & \Delta H_{\mathrm{rxn}}^{\circ}=872.8 \mathrm{~kJ} / \mathrm{mol} \end{array} $$ and the fact that the average \(\mathrm{C}-\mathrm{H}\) bond enthalpy is \(414 \mathrm{~kJ} / \mathrm{mol}\), estimate the standard enthalpy of formation of methane \(\left(\mathrm{CH}_{4}\right)\).

The term "molar mass" was introduced in Chapter 3 What is the advantage of using the term "molar mass" when we discuss ionic compounds?

What is a coordinate covalent bond? Is it different from a normal covalent bond?
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