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Chapter 10: Problem 51

Draw Lewis structures for the following species, indicating formal charges and resonance where applicable: (a) \(\mathrm{HCO}_{2}=\) (b) \(\mathrm{HCO}_{3}^{-}\) (c) \(\mathrm{FSO}_{3}^{-}\) (d) \(\mathrm{N}_{2} \mathrm{O}_{3}^{2-}\) (the nitrogen atoms are joined centrally with one oxygen atom on one \(\mathrm{N}\) and two on the other)

Short Answer

Expert verified
The Lewis structures can be drawn by following these steps: counting the total number of valence electrons, arranging the atoms with the less electronegative atom usually in the center, placing a single bond between the central atom and the surrounding atoms, distributing the remaining electrons in the form of lone pairs, and creating double or triple bonds if necessary to fulfill the octet rule. Resonance structures may be present when there are several valid ways to place the remaining electrons.

Step by step solution

01

Calculate Total Valence Electrons

This is determined by adding up the number of valence electrons each atom in the molecule/ion has. Hydrogen has 1, oxygen has 6, carbon has 4, nitrogen has 5, sulfur has 6, and fluorine has 7 valence electrons. For ions, remember to add electrons if the molecule/ion carries a negative charge and subtract electrons if it carries a positive charge.
02

Arrange Atoms

The less electronegative atom is typically the central atom (hydrogen never goes in the center). The arrangement of atoms in \( \mathrm{N}_{2} \mathrm{O}_{3}^{2-} \) is mentioned in the question.
03

Place Four Two-Electron Bonds

Place a single bond (two electrons) between the central atom and surrounding atoms.
04

Spread Out Remaining Electrons

Starting from the outer atoms, remaining electrons are placed in the form of lone pairs. Each atom, except hydrogen, should be surrounded by 8 electrons (Octet Rule). If we've not used all the electrons after the outer atoms have obtained octets, place left electrons around the central atom.
05

Create Double or Triple Bonds

If any atoms don’t have a full octet, move electrons from lone pairs on the outer atoms to create double or triple bonds. You may see different resonance forms, if applicable, since those are alternate ways the same molecule could look.

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Most popular questions from this chapter

The species \(\mathrm{PBr}_{4}^{-}\) has been synthesized and has been described as a tetrahedral anion. Comment on this description.

One of the isomers of chloromethanol has the formula $\mathrm{ClCH}_{2} \mathrm{OH} .$ Sketch, by using the dash and wedge symbolism, this isomer of chloromethanol, and indicate the various bond angles.

The formal charges on the \(\mathrm{O}\) atoms in the ion \([\mathrm{ONO}]^{+}\) is \((\mathrm{a})-2 ;(\mathrm{b})-1 ;(\mathrm{c}) 0 ;(\mathrm{d})+1\).

The following polyatomic anions involve covalent bonds between O atoms and the central nonmetal atom. Propose an acceptable Lewis structure for each. (a) \(\mathrm{SO}_{3}^{2-} ;\) (b) \(\mathrm{NO}_{2}^{-} ;\) (c) \(\mathrm{CO}_{3}^{2-} ;\) (d) \(\mathrm{HO}_{2}^{-}\).

Assign formal charges to the atoms in the following species, and then select the more likely skeletal structure. (a) \(\mathrm{H}_{2} \mathrm{NOH}\) or \(\mathrm{H}_{2} \mathrm{ONH}\) (b) SCS or CSS (c) NFO or FNO (d) \(\mathrm{SOCl}_{2}\) or \(\mathrm{OSCl}_{2}\) or \(\mathrm{OCl}_{2} \mathrm{S}\) (e) \(\mathrm{F}_{3} \mathrm{SN}\) and \(\mathrm{F}_{3} \mathrm{NS}\)
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