Question

Write Lewis structures that obey the octet rule (duet rule for H) for each of the following molecules. a. \(\mathrm{H}_{2} \mathrm{CO} \quad\) c. \(\mathrm{HCN}\) b. \(\mathrm{CO}_{2}\)

Short answer

The Lewis structures for the given molecules are as follows: a. H2CO: H-C=O-H b. CO2: O=C=O c. HCN: H-C≡N

Step-by-step solution

Counting valence electrons

For each molecule, we first count the total number of valence electrons available for bonding. We do this by adding the valence electrons of each individual atom in the molecule. a. H2CO: - Hydrogen: 1 valence electron (each hydrogen atom) - Carbon: 4 valence electrons - Oxygen: 6 valence electrons Total: 1(2) + 4 + 6 = 12 valence electrons b. CO2: - Carbon: 4 valence electrons - Oxygen: 6 valence electrons (each oxygen atom) Total: 4 + 6(2) = 16 valence electrons c. HCN: - Hydrogen: 1 valence electron - Carbon: 4 valence electrons - Nitrogen: 5 valence electrons Total: 1 + 4 + 5 = 10 valence electrons **Step 2: Determine the Central Atom**

Determining the central atom

In each molecule, the atom with the highest bonding capacity, which is usually the least electronegative element, will be the central atom. a. H2CO: Carbon is the central atom. b. CO2: Carbon is the central atom. c. HCN: Carbon is the central atom. **Step 3: Arrange Atoms and Distribute Electrons**

Arranging atoms and electron distribution

Next, we arrange the atoms around the central atom and distribute the electrons to form bonds and fulfill the octet rule for each atom (excluding hydrogen). a. H2CO: Place the carbon atom in the center, with the two hydrogen atoms and the oxygen atom around it. Two electrons form a bond between each hydrogen atom and the carbon atom. Then, we add a double bond between carbon and oxygen, since oxygen needs two more electrons to follow the octet rule, and we have 4 valence electrons left in our pool. This leaves us with the Lewis structure: H-C=O-H. b. CO2: Place the carbon atom in the center with both oxygen atoms on either side. Form two double bonds, one between each oxygen atom and the carbon atom. This uses the 16 valence electrons available, and the Lewis structure is: O=C=O. c. HCN: Place the carbon atom in the center with hydrogen atom and nitrogen atom on either side. One electron pair forms a bond between hydrogen and carbon. A triple bond between carbon and nitrogen uses the remaining six valence electrons, fulfilling the octet rule for both carbon and nitrogen, while hydrogen follows the duet rule. The Lewis structure is: H-C≡N. Now we have successfully drawn the Lewis structures for all three molecules that obey the octet and duet rules.

Key concepts

Octet Rule

The octet rule is a fundamental principle in chemistry that helps explain how atoms form compounds. Simply put, atoms prefer to have eight electrons in their outer shell, similar to the electron configuration of noble gases. This is because a full outer shell provides stability to the atom. However, there is a notable exception: hydrogen only needs two electrons to become stable, a concept known as the duet rule.
In the context of creating Lewis structures, the octet rule directs us to ensure each atom in the molecule (besides hydrogen) has a complete set of eight electrons. This often involves forming bonds where electrons are shared between atoms. For example, in the molecule CO₂ from the exercise, each oxygen atom is connected to carbon through double bonds, satisfying the octet rule.

Valence Electrons

Valence electrons are the electrons found in the outermost shell of an atom. These electrons play a key role in chemical bonding because they can be gained, lost, or shared to stabilize the atom by achieving a full outer shell. Counting valence electrons is the first crucial step in drawing Lewis structures.
For instance, in the case of HCN, hydrogen contributes 1 valence electron, carbon has 4, and nitrogen has 5. This sums up to a total of 10 valence electrons for the HCN molecule, which are then strategically used for bonding between the atoms to satisfy the octet and duet rules.

Chemical Bonding

Chemical bonding refers to the process where atoms connect or "bond" together to form molecules. These bonds result from atoms sharing or transferring valence electrons to achieve stable electron configurations. In the context of Lewis structures, we often deal with covalent bonds, where electron pairs are shared between atoms.
In CO₂, double bonds form between the carbon and each oxygen atom, highlighting a classic example of covalent bonding. These shared electrons help both carbon and oxygen achieve their ideal electron configuration. The formation of triple bonds, as seen in HCN between carbon and nitrogen, illustrates how multiple shared electron pairs can increase bond strength and stability.

Molecular Geometry

Molecular geometry is the three-dimensional arrangement of atoms within a molecule. It is crucial because it influences the molecule's properties, including reactivity, polarity, and color. Understanding molecular geometry involves considering both the types of chemical bonds and how they spatially orient.
While the exercise mainly focuses on drawing Lewis structures in a planar form, it also sets the foundation for understanding geometry. For example, the linear formation of HCN, where both bonds are aligned in a line, affects how this molecule interacts with others. Similarly, the linear shape of CO₂ is due to the double bonds being symmetrically placed opposite to each other, leading to a balanced structure without a net dipole moment.