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Chapter 8: Problem 86

Lewis structures can be used to understand why some molecules react in certain ways. Write the Lewis structures for the reactants and products in the reactions described below. a. Nitrogen dioxide dimerizes to produce dinitrogen tetroxide. b. Boron trihydride accepts a pair of electrons from ammonia, forming \(\mathrm{BH}_{3} \mathrm{NH}_{3}\). Give a possible explanation for why these two reactions occur.

Short Answer

Expert verified
In summary, the dimerization of nitrogen dioxide (NO鈧�) to form dinitrogen tetroxide (N鈧侽鈧�) occurs because it enables nitrogen atoms to attain a full octet by sharing their unpaired electrons. The reaction between boron trihydride (BH鈧�) and ammonia (NH鈧�) to form the adduct \(\mathrm{BH}_{3} \mathrm{NH}_{3}\) occurs due to the electron-deficient boron atom in BH鈧� and the electron-rich nitrogen atom in NH鈧�. The nitrogen atom donates its lone pair of electrons to the empty orbital of boron, resulting in the formation of a dative covalent bond and allowing both boron and nitrogen atoms to achieve a full octet configuration.

Step by step solution

01

Identify reactants and products

For the given reactions: a. Nitrogen dioxide (NO鈧�) dimerizes to produce dinitrogen tetroxide (N鈧侽鈧�). b. Boron trihydride (BH鈧�) reacts with ammonia (NH鈧�) to form the adduct \(\mathrm{BH}_{3} \mathrm{NH}_{3}\).
02

Draw Lewis structures

For reaction a: - Nitrogen dioxide (NO鈧�): Lewis structure contains a nitrogen atom with a single unpaired electron, double bonded to one oxygen atom and single bonded to the other oxygen atom, which bears a formal negative charge. - Dinitrogen tetroxide (N鈧侽鈧�): Lewis structure contains two nitrogen atoms, each double bonded to two oxygen atoms. For reaction b: - Boron trihydride (BH鈧�): Lewis structure contains a boron atom with three single bonds to three hydrogen atoms. - Ammonia (NH鈧�): Lewis structure contains a nitrogen atom with a lone pair of electrons and three single bonds to three hydrogen atoms.
03

Explain the processes or behaviors that drive each reaction

a. The dimerization of NO鈧� to form N鈧侽鈧� occurs because it allows nitrogen atoms to attain a full octet. In NO鈧�, the nitrogen atom has an unpaired electron, which makes it a radical. The dimerization reaction results in the sharing of this unpaired electron, forming a N-N single bond and enabling both nitrogen atoms to achieve a full octet. b. The reaction of BH鈧� with NH鈧� to form BH鈧僋H鈧� occurs due to the electron-deficient boron atom in BH鈧� and the electron-rich nitrogen atom in NH鈧�. Boron in BH鈧� has only six valence electrons, which does not satisfy the octet rule. The nitrogen atom in NH鈧�, with a lone pair of electrons, can form a dative covalent (coordinate) bond by donating its electron pair into the empty orbital of boron. This results in the formation of the adduct BH鈧僋H鈧�, wherein both boron and nitrogen atoms attain a full octet configuration.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Dimerization
Dimerization is a process by which two identical molecules, known as monomers, combine to form a larger molecule called a dimer. In the context of the reaction involving nitrogen dioxide,
the term "dimerization" refers to the transformation of two NO鈧� molecules into dinitrogen tetroxide (N鈧侽鈧�).
  • NO鈧� molecules contain a nitrogen atom that has an unpaired electron, which makes them highly reactive and unstable.
  • By combining two NO鈧� molecules, this unpaired electron can pair up, forming a bond and stabilizing the resulting molecule, N鈧侽鈧�.
  • The newly formed N鈧侽鈧� consists of two nitrogen atoms double-bonded to oxygen atoms, effectively sharing the unpaired electron between the NO鈧� units.
Dimerization is an important reaction because it often increases stability within a molecule,
allowing atoms to achieve more stable electronic configurations compared to their monomer states.
This process is not unique to nitrogen dioxide but can be observed in various chemical systems where monomers seek stability.
Octet Rule
The octet rule is a fundamental principle in chemistry, which states that atoms tend to form compounds in ways that provide them with eight valence electrons,
resulting in electron configurations similar to those of noble gases. This rule guides many chemical bonding and molecular formation processes.
  • In the context of the reactions described, nitrogen dioxide (NO鈧�) dimerizes to form dinitrogen tetroxide (N鈧侽鈧�), allowing the nitrogen atoms to achieve stable octet configurations.
  • For a single NO鈧� molecule, the nitrogen atom is a radical, meaning it does not satisfy the octet rule, leading to reactive behavior.
  • By dimerizing, each nitrogen atom in N鈧侽鈧� attains an octet, which provides greater stability than the monomer state.
Similarly, in the adduct formation with boron trihydride (BH鈧�) and ammonia (NH鈧�),
the octet rule helps explain why NH鈧� donates a pair of electrons to BH鈧�,
satisfying the octet requirement for both the boron and nitrogen atoms.
Understanding the octet rule can give insights into the reactivity and bonding preferences of different atoms and molecules.
Coordinate Bonding
Coordinate bonding, also known as dative covalent bonding, occurs when one atom provides both electrons for a covalent bond.
This is particularly significant in the formation of complex molecules or adducts between electron-rich and electron-deficient atoms.
  • In the case of the reaction between boron trihydride (BH鈧�) and ammonia (NH鈧�), coordinate bonding is demonstrated when the nitrogen atom of NH鈧� donates a lone pair of electrons to the empty orbital of the boron atom in BH鈧�.
  • The nitrogen atom acts as a Lewis base, supplying the electron pair, while the boron atom acts as a Lewis acid, accepting the electron pair.
  • This bonding type allows the formation of BH鈧僋H鈧�, where both the boron and nitrogen atoms achieve a stable electron configuration.
Coordinate bonding is crucial in chemistry for forming stable compounds where traditional covalent bonding is not possible.
It is essential to understand this concept to grasp how molecules with incomplete valence shells can form stable interactions.

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Most popular questions from this chapter

Write Lewis structures that obey the octet rule for each of the following molecules and ions. (In each case the first atom listed is the central atom.) a. \(\mathrm{POCl}_{3}, \mathrm{SO}_{4}^{2-}, \mathrm{XeO}_{4}, \mathrm{PO}_{4}^{3-}, \mathrm{ClO}_{4}^{-}\) b. \(\mathrm{NF}_{3}, \mathrm{SO}_{3}{ }^{2-}, \mathrm{PO}_{3}{ }^{3-}, \mathrm{ClO}_{3}^{-}\) c. \(\mathrm{ClO}_{2}^{-}, \mathrm{SCl}_{2}, \mathrm{PCl}_{2}^{-}\) d. Considering your answers to parts \(\mathrm{a}, \mathrm{b}\), and \(\mathrm{c}\), what conclusions can you draw concerning the structures of species containing the same number of atoms and the same number of valence electrons?

Nitrous oxide \(\left(\mathrm{N}_{2} \mathrm{O}\right)\) has three possible Lewis structures: Given the following bond lengths, $$ \begin{array}{llll} \mathrm{N}-\mathrm{N} & 167 \mathrm{pm} & \mathrm{N}=\mathrm{O} & 115 \mathrm{pm} \\ \mathrm{N}=\mathrm{N} & 120 \mathrm{pm} & \mathrm{N}-\mathrm{O} & 147 \mathrm{pm} \\ \mathrm{N} \equiv \mathrm{N} & 110 \mathrm{pm} & & \end{array} $$ rationalize the observations that the \(\mathrm{N}-\mathrm{N}\) bond length in \(\mathrm{N}_{2} \mathrm{O}\) is \(112 \mathrm{pm}\) and that the \(\mathrm{N}-\mathrm{O}\) bond length is \(119 \mathrm{pm}\). Assign formal charges to the resonance structures for \(\mathrm{N}_{2} \mathrm{O}\). Can you eliminate any of the resonance structures on the basis of formal charges? Is this consistent with observation?

Look up the energies for the bonds in \(\mathrm{CO}\) and \(\mathrm{N}_{2}\). Although the bond in \(\mathrm{CO}\) is stronger, \(\mathrm{CO}\) is considerably more reactive than \(\mathrm{N}_{2}\). Give a possible explanation.

Predict the molecular structure (including bond angles) for each of the following. a. \(\mathrm{SeO}_{3}\) b. \(\mathrm{SeO}_{2}\)

Give three ions that are isoelectronic with neon. Place these ions in order of increasing size.
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