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Chapter 10: Problem 24

Show that the idea of minimizing the formal charges in a structure is at times in conflict with the observation that compact, symmetrical structures are more commonly observed than elongated ones with many central atoms. Use \(\mathrm{ClO}_{4}^{-}\) as an illustrative example.

Short Answer

Expert verified
In the given exercise, it's shown that two principles, minimizing formal charges and achieving symmetric structures, could produce different Lewis structures for the same molecule. While non-symmetrical structures tend to minimize formal charges, observed structures in nature often prioritize symmetry over charge minimization, as seen in the perchlorate ion example, \(\mathrm{ClO}_{4}^{-}\).

Step by step solution

01

Constructing a Lewis structure minimizing formal charges

If we want to minimize formal charges, we would add all the extra electrons in the outer layers of oxygen instead of the central chlorine atom. To create the Lewis structure, the single bond should be created between Chlorine (Cl) and each Oxygen (O) atom. Since each oxygen atom is considered to be having a formal charge of -1, the structure indicates -4 overall charge which is consistent with its ionic form. Now the chlorine atom, which is in the center, becomes +3. So, we have a structure with a minimized formal charge, but it is not symmetrical.
02

Constructing a Lewis structure focusing on symmetry

For more symmetrical structure, we distribute the charges evenly throughout all the atoms. We add a double bond with each of the oxygen atoms which reduces the charge on oxygen to zero since there are now 6 electrons distributed on each oxygen atom’s outer layer. As a result the central chlorine atom will get a charge of -1 which will provide us a structure of \(\mathrm{ClO}_{4}^{-}\), the negative charge being on the central atom. This structure represents a symmetric structure, however, it contains non-minimal formal charges.
03

Discussing conflicting principles in terms of observed structures

While the structure built under the principle of minimizing formal charges has a formal charge distributed throughout the molecule, the symmetrically built structure converges this charge in the center by establishing double bonds with each of the peripheral atoms. The latter, symmetric structure, is observed more frequently. Therefore, although we aim to minimize formal charges when building structures in chemistry, the more often observed structures suggests that a symmetrical format is often realized in nature, which is contrary to the principle of charge minimization.

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Most popular questions from this chapter

Each of the following ionic compounds consists of a combination of monatomic and polyatomic ions. Represent these compounds with Lewis structures. (a) \(\mathrm{Al}(\mathrm{OH})_{3}\); (c) \(\mathrm{NH}_{4} \mathrm{F}\); (d) \(\mathrm{KClO}_{3}\); (b) \(\mathrm{Ca}(\mathrm{CN})_{2}\); (e) \(\mathrm{Ba}_{3}\left(\mathrm{PO}_{4}\right)_{2}\).

In your own words, define the following terms: (a) valence electrons; (b) electronegativity; (c) bonddissociation energy; (d) double covalent bond;(e) coordinate covalent bond.

Under appropriate conditions, both hydrogen and nitrogen can form monatomic anions. What are the Lewis symbols for these ions? What are the Lewis structures of the compounds (a) lithium hydride; (b) calcium hydride; (c) magnesium nitride?

Alternative strategies to the one used in this chapter have been proposed for applying the VSEPR theory to molecules or ions with a single central atom. In general, these strategies do not require writing Lewis structures. In one strategy, we write (1) the total number of electron pairs \(=[\) (number of valence electrons) \(\pm\) (electrons required for ionic charge) \(] / 2\) (2) the number of bonding electron pairs \(=\) (number of atoms) -1 (3) the number of electron pairs around central atom \(=\) total number of electron pairs \(-3 \times[\) number of terminal atoms (excluding \(\mathrm{H}\) )] (4) the number of lone-pair electrons = number of central atom pairs - number of bonding pairs After evaluating items \(2,3,\) and \(4,\) establish the VSEPR notation and determine the molecular shape. Use this method to predict the geometrical shapes of the following: (a) \(\mathrm{PCl}_{5} ;\) (b) \(\mathrm{NH}_{3} ;\) (c) \(\mathrm{ClF}_{3} ;\) (d) \(\mathrm{SO}_{2} ;\) (e) \(\mathrm{ClF}_{4}^{-}\); (f) \(\mathrm{PCl}_{4}^{+}\). Justify each of the steps in the strategy, and explain why it yields the same results as the VSEPR method based on Lewis structures. How does the strategy deal with multiple bonds?

Hydrogen azide, \(\mathrm{HN}_{3}\), is a liquid that explodes violently when subjected to physical shock. In the \(\mathrm{HN}_{3}\) molecule, one nitrogen- to-nitrogen bond length is \(113 \mathrm{pm},\) and the other is \(124 \mathrm{pm} .\) The \(\mathrm{H}-\mathrm{N}-\mathrm{N}\) bond angle is \(112^{\circ} .\) Draw Lewis structures and a sketch of the molecule consistent with these facts.
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