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Chapter 10: Problem 21

What is the formal charge of the indicated atom in each of the following structures? (a) the central \(\mathrm{O}\) atom in \(\mathrm{O}_{3}\) (b) \(\mathrm{Al}\) in \(\mathrm{AlH}_{4}^{-}\) (c) \(\mathrm{Cl}\) in \(\mathrm{ClO}_{3}^{-}\) (d) \(\sin \sin ^{2} \theta^{-}\) (e) \(\mathrm{Cl}\) in \(\mathrm{Cl} \mathrm{F}_{3}\)

Short Answer

Expert verified
The formal charges for the atoms in the given molecules and ions are: (a) 0, (b) -1, (c) -1, (d) -2, (e) 0.

Step by step solution

01

Calculate the formal charge of O in O3

Oxygen normally has 6 valence electrons. For O3, each oxygen atom is bonded to another oxygen (that's 2 bonding electrons). Each outer oxygen has six non-bonding electrons and the central one has four. So for the central oxygen, the formal charge = 6 - 4 - 1/2*2 = 0.
02

Calculate the formal charge of Al in AlH4-

Aluminum normally has 3 valence electrons, but in AlH4-, it's bonded to 4 hydrogen atoms (4 bonding electrons) and there are no non-bonding electrons. So, the formal charge = 3 - 0 - 1/2*4 = -1
03

Calculate the formal charge of Cl in ClO3-

Chlorine normally has 7 valence electrons, but in ClO3-, it's bonded to 3 oxygen atoms (6 bonding electrons as each bond is a double bond) and has one pair of non-bonding electrons. So the formal charge = 7 - 2 - 1/2*6 = -1.
04

Calculate the formal charge of S in S2-

Sulfur normally has 6 valence electrons. In S^2-, these electrons are all non-bonding. So, the formal charge = 6 - 6 - 1/2*0 = 0. (We add 2 for the extra electrons that sulfur has gained, giving a formal charge of -2.)
05

Calculate the formal charge of Cl in ClF3

Chlorine normally has 7 valence electrons, but in ClF3, it's bonded to 3 fluorine atoms (6 bonding electrons) and has one pair of non-bonding electrons. So, the formal charge = 7 - 2 - 1/2*6 = 0.

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Most popular questions from this chapter

Both oxidation state and formal charge involve conventions for assigning valence electrons to bonded atoms in compounds, but clearly they are not the same. Describe several ways in which these concepts differ.

Draw a plausible Lewis structure for the following series of molecules and ions: \((a) \operatorname{SiF}_{6}^{2-} ;\) (b) \(\mathrm{PF}_{5} ;\) (c) \(\mathrm{SF}_{4}\); (d) \(\mathrm{XeF}_{4}\). Describe the electron group geometry and molecular structure of these species.

One of the allotropes of sulfur is a ring of eight sulfur atoms. Draw the Lewis structure for the \(S_{8}\) ring. Is the ring likely to be planar? The \(S_{8}\) ring can be oxidized to produce \(S_{8}\) O. In \(S_{8} O,\) the oxygen atom is bonded to one of the S atoms and the Sg ring is still intact. Draw the Lewis structure for \(\mathrm{S}_{8} \mathrm{O}\).

One possibility for the electron-group geometry for seven electron groups is pentagonal-bipyramidal, as found in the IF \(_{7}\) molecule. Write the VSEPR notation for this molecule. Sketch the structure of the molecule, labeling all the bond angles.

Use the VSEPR theory to predict a probable shape of the molecule \(\mathrm{F}_{4} \mathrm{SCH}_{2}\), and explain the source of any ambiguities in your prediction.
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