Question

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

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

(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.

(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.

(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.

Key concepts

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—single, double, or triple—is 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.