Question

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

Short answer

(a) Lewis Structure: O=C-H-H (b) Lewis Structure: H-O-O-H (c) Lewis Structure: F-C-C-F | | F F-F (d) Lewis Structure: [O=As=O]^3- (e) Lewis Structure: O=S-O-H | O-H (f) Lewis Structure: H-C=C-H

Step-by-step solution

Count the valence electrons

In this molecule, Carbon has 4 valence electrons, Oxygen has 6, and each Hydrogen atom has 1. So, the total valence electrons are 4 + 6 + (1x2) = 12.

Sketch the molecule's basic skeleton

Place Carbon in the center forming single bonds between O and H. The structure will look like this: O-H-C-H.

Distribute the remaining electrons

We've already used 4 out of 12 electrons building single bonds between the atoms. Distribute the remaining 8 valence electrons as lone pairs around the Oxygen atom.

Apply the octet rule

Oxygen already has an octet, and Hydrogen atoms have their required 2 electrons each. Carbon atoms however, only have 6 electrons in their valence shell. Move one of the lone pairs on Oxygen and form a double bond with Carbon to fulfill the octet rule for all atoms.

Verify formal charges

The resulting Lewis structure is O=C-H-H and the formal charges of all atoms should be zero. (b) \(\mathrm{H}_{2} \mathrm{O}_{2}\)

Count the valence electrons

In this molecule, each Oxygen atom has 6 valence electrons, and each Hydrogen atom has 1. So, the total valence electrons are (6x2) + (1x2) = 14.

Sketch the molecule's basic skeleton

Place the Oxygen atoms in the center and form single bonds between them. The structure will look like this: H-O-O-H.

Distribute the remaining electrons

We've already used 4 out of 14 electrons building single bonds between the atoms. Distribute the remaining 10 valence electrons as lone pairs around both Oxygen atoms.

Apply the octet rule

Apply the octet rule by placing lone pairs on Oxygen atoms so that both satisfy the octet rule.

Verify formal charges

The resulting Lewis structure is H-O-O-H, with each Oxygen atom containing two lone pairs. The formal charges of all atoms should be zero. The remaining molecules will follow the same steps to determine their Lewis structures: (c) \(\mathrm{C}_{2} \mathrm{F}_{6}\) Lewis Structure: F-C-C-F | | F F-F (d) \(\mathrm{AsO}_{3}^{3-}\) Lewis Structure: [O=As=O]^3- (e) \(\mathrm{H}_{2} \mathrm{SO}_{3}\) Lewis Structure: O=S-O-H | O-H (f) \(\mathrm{C}_{2} \mathrm{H}_{2}\) Lewis Structure: H-C=C-H

Key concepts

Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom that can participate in forming chemical bonds with other atoms. These electrons are significant because they determine how an atom will react chemically and thus play a key role in the structure of molecules.

For instance, in the exercise provided, when constructing the Lewis structure for \(\mathrm{H}_2\mathrm{CO}\), a count of the valence electrons is the first step. Carbon, being in group IV, has 4 valence electrons; Oxygen, in group VI, has 6; and Hydrogen, a member of group I, has 1 valence electron per atom. By adding these up, we find that \(\mathrm{H}_2\mathrm{CO}\) has 12 valence electrons available to form bonds and complete the octets.

Remember that the number of valence electrons more, the more an atom can engage in bonds to achieve stability, typically following the octet rule which we'll discuss next.

Octet Rule

The octet rule is based on the principle that atoms are most stable when they have eight electrons in their valence shell, resembling the electron arrangement of the noble gases. This rule is a guiding principle for many molecules when predicting the bonding behavior and the arrangement of electrons in the Lewis structures.

In the exercise, after sketching the basic skeleton for \(\mathrm{H}_2\mathrm{CO}\), it becomes apparent that Carbon cannot satisfy the octet rule with only single bonds to Hydrogen and Oxygen. This is resolved by forming a double bond between Carbon and Oxygen, thus allowing Carbon to reach the stable configuration of eight electrons around it.

Exceptions to the Octet Rule

It's also important to recognize that there are exceptions to this rule. Atoms in the third period of the periodic table and beyond can utilize d orbitals to expand their valence shell beyond eight electrons, and some molecules are stable with less than an octet.

Formal Charge

The concept of formal charge is crucial for understanding the distribution of electrons in a molecule and helps in determining the most stable Lewis structure. Formal charge is the hypothetical charge you would assign to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms.

To calculate it, use the formula: \[\text{Formal Charge} = (\text{Valence Electrons}) - (0.5 \times \text{Bonding Electrons}) - (\text{Non-bonding Electrons})\] For instance, in the molecule \(\mathrm{H}_2\mathrm{CO}\) once the double bond is introduced, all atoms have a formal charge of zero, which often corresponds to the most stable arrangement of electrons.

Optimizing Formal Charge

When creating a Lewis structure, if the formal charges are not at their lowest, it might indicate that the structure can be improved, either by rearranging the electrons to form multiple bonds or by altering the placement of electrons to better satisfy the octet rule.