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Chapter 8: Problem 48

Write Lewis structures for the following: (a) \(\mathrm{H}_{2} \mathrm{CO}\) (both \(\mathrm{H}\) atoms are bonded to \(\mathrm{C} ),(\mathbf{b}) \mathrm{H}_{2} \mathrm{O}_{2},(\mathbf{c}) \mathrm{C}_{2} \mathrm{F}_{6}\) (contains \(\mathrm{a} \mathrm{C}-\mathrm{C}\) bond), \((\mathbf{d}) \mathrm{AsO}_{3}^{3-},\) (e) \(\mathrm{H}_{2} \mathrm{SO}_{3}(\mathrm{H}\) is bonded to \(\mathrm{O})\) \((\mathbf{f}) \mathrm{NH}_{2} \mathrm{Cl}\)

Short Answer

Expert verified
The Lewis structures for the requested molecules are as follows: (a) H2CO: \(\mathrm{H-C=O}\) (b) H2O2: \(\mathrm{H-O-O-H}\) (c) C2F6: \(\mathrm{F^-\equiv F^-\equiv C-F^-\equiv F^-\equiv C-F^-\equiv F^-}\) (d) AsO3^3-: \(\mathrm{:O=As=O:}\) (e) H2SO3: \(\mathrm{O=C-OH}\) (f) NH2Cl: \(\mathrm{H-N-Cl}\)

Step by step solution

01

(Step 1: Total valence electrons)

For each molecule, find the total number of valence electrons. (a) H2CO: H has 1 valence electron; C has 4 valence electrons; O has 6 valence electrons. Total = (2x1) + 4 + 6 = 12 valence electrons. (b) H2O2: H has 1 valence electron; O has 6 valence electrons. Total = (2x1) + (2x6) = 14 valence electrons. (c) C2F6: C has 4 valence electrons; F has 7 valence electrons. Total = (2x4) + (6x7) = 46 valence electrons. (d) AsO3^3-: As has 5 valence electrons; O has 6 valence electrons; the 3- charge adds 3 more electrons. Total = 5 + (3x6) + 3 = 26 valence electrons. (e) H2SO3: H has 1 valence electron; S has 6 valence electrons; O has 6 valence electrons. Total = (2x1) + 6 + (3x6) = 26 valence electrons. (f) NH2Cl: N has 5 valence electrons; H has 1 valence electron; Cl has 7 valence electrons. Total = 5 + (2x1) + 7 = 14 valence electrons.
02

(Step 2: Draw skeleton structures)

Draw the skeleton structure for each molecule with a single bond between the atoms. (a) H2CO: `[H-C-H]` and place the O above the C atom. (b) H2O2: `[H-O-O-H]`. (c) C2F6: `[C-C]` with three F atoms around each C atom. (d) AsO3^3-: `[As]` with three O atoms around it. (e) H2SO3: `[S]` with three O atoms around it and one H atom attached to two of the O atoms. (f) NH2Cl: `[N]` with one H atom and one Cl atom placed around the N atom.
03

(Step 3: Distribute electrons)

Distribute the remaining valence electrons as lone pairs (LP) and complete the octets for each atom. (a) H2CO: `[H-C-H]` and place two LP on the O atom, and double bond between C and O. (b) H2O2: `[H-O-O-H]` and place two LP on each O atom. (c) C2F6: `[C-C]` with three F atoms around each C atom and add three LP to each F atom. (d) AsO3^3-: `[As]` with three O atoms around it, place three LP on each O atom and place a double bond between As and each O atom. (e) H2SO3: `[S]` with three O atoms around it, place three LP on each O atom, place two LP on the S atom, and place a double bond between S and one O atom. (f) NH2Cl: `[N]` with one H atom and one Cl atom placed around the N atom, add one LP to the N atom, and three LP to the Cl atom. These are the Lewis structures for the requested molecules.

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Key Concepts

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

Valence Electrons
Valence electrons are the outermost electrons of an atom and play a key role in chemical bonding. They determine how an atom interacts with other atoms. Understanding the number of valence electrons helps us predict how elements will bond.

For instance, carbon, with its 4 valence electrons, typically forms four bonds to achieve a stable configuration, while oxygen, with 6 valence electrons, tends to form two bonds. Recognizing these patterns is essential when drawing Lewis structures.
  • Hydrogen is unique with only 1 valence electron and forms single bonds.
  • Halogens like fluorine have 7 valence electrons, allowing them to form one bond.
  • Some ions, like \( ext{AsO}_3^{3-}\), gain or lose electrons, indicated by charges which must be considered in valence electron counts.
In summary, mastering valence electrons is the first step in building accurate molecular models.
Chemical Bonding
Chemical bonding involves the interaction between valence electrons of different atoms that allow them to achieve stability. Covalent bonds form when atoms share electrons, a crucial concept when creating Lewis structures.

Lewis structures visually depict how atoms share valence electrons. Each bond represents a pair of shared electrons. The type of bond鈥攕ingle, double, or triple鈥攊s determined by the number of shared electron pairs.
  • Single bonds involve one pair of shared electrons.
  • Double bonds share two pairs, as seen in \( ext{H}_2 ext{CO}\) between carbon and oxygen.
  • Triple bonds share three pairs, though not present in the exercise examples.
Bonding explains molecule stability and reactivity, where complete octets or duets (for hydrogen) indicate a stable configuration.
Molecular Geometry
Molecular geometry describes the 3D arrangement of atoms in a molecule, impacting properties like polarity and reactivity. It depends on the number of bonds and lone pairs present around the central atom.

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shape by repelling electron pairs to maximize distance between them. For example:
  • Linear (e.g., \( ext{C}_2 ext{F}_6\) when considering C-C bond linearly, each C surrounded by F atoms as tetrahedral).
  • Tetrahedral (when a central atom forms 4 bonds, like in ammonia derivatives).
  • Bent (as seen in water, \( ext{H}_2 ext{O}\), due to two bonding pairs and two lone pairs).
Each molecule's geometry influences its molecular interactions and characteristics, making this concept vital for predicting behavior and reactions.

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Most popular questions from this chapter

Draw the Lewis structures for each of the following ions or molecules. Identify those in which the octet rule is not obeyed; state which atom in each compound does not follow the octet rule; and state, for those atoms, how many electrons surround these atoms: (a) \(\mathrm{PH}_{3},\) (b) AlH_ \(_{3},(\mathbf{c}) \mathrm{N}_{3}^{-}\) (d) \(\mathrm{CH}_{2} \mathrm{Cl}_{2},(\mathbf{e}) \mathrm{SnF}_{6}^{2-}\)

The substances \(\mathrm{NaF}\) and \(\mathrm{CaO}\) are isoelectronic (have the same number of valence electrons). (a) What are the charges on each of the cations in each compound? (b) What are the charges of each of the anions in each compound? (c) Without looking up lattice energies, which compound is predicted to have the larger lattice energy? (d) Using the lattice energies in Table \(8.1,\) predict the lattice energy of ScN.

Write the Lewis symbol for atoms of each of the following elements: (a) Al, (b) Br, (c) Ar, (d) Sr.

Draw the Lewis structure for \(\mathrm{NO}^{+} .\) Is the nitrogen-oxygen bond in \(\mathrm{NO}^{+}\) longer, shorter, or the same length as the nitrogen- oxygen bond in NO? Explain.

(a) Construct a Lewis structure for \(\mathrm{O}_{2}\) in which each atom achieves an octet of electrons. (b) How many bonding electrons are in the structure? (c) Would you expect the \(\mathrm{O}-\mathrm{O}\) bond in \(\mathrm{O}_{2}\) to be shorter or longer than the \(\mathrm{O}-\mathrm{O}\) bond in compounds that contain an \(\mathrm{O}-\mathrm{O}\) single bond? Explain.
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