Question

The "Chemistry in Focus" segment Hiding Carbon Dioxide discusses attempts at sequestering (storing) underground \(\mathrm{CO}_{2}\) produced at power plants so as to diminish the greenhouse effect. Draw all resonance structures of the \(\mathrm{CO}_{2}\) molecule.

Short answer

In the case of the \(\mathrm{CO}_{2}\) molecule, there are no resonance structures. The original Lewis structure of O=C=O properly represents the linear molecule with all valence electrons involved in bonding and no lone pairs on the central carbon atom.

Step-by-step solution

Recall the Lewis structure of carbon dioxide

First, let's recall the Lewis structure of the carbon dioxide (\(\mathrm{CO}_{2}\)) molecule. Carbon has 4 valence electrons, while oxygen has 6 valence electrons. In \(\mathrm{CO}_{2}\), there is one carbon atom and two oxygen atoms. So, the total number of valence electrons is: Total valence electrons = 4 (for Carbon) + 2 x 6 (for each Oxygen) = 4 + 12 = 16 The Lewis structure of \(\mathrm{CO}_{2}\) is linear, with the carbon atom double-bonded to each oxygen atom. Thus, the representation is O=C=O.

Identify the possibility of resonance structures

In order to generate resonance structures, we look for any possible way to redistribute the electrons within the molecule without changing its overall charge. In the case of \(\mathrm{CO}_{2}\), we can observe that all valence electrons are already involved in bonding and there are no lone pairs on the central atom (carbon).

Attempt to create resonance structures

Since there are no lone pairs on the carbon atom and both oxygen atoms are already double-bonded to the carbon atom, we cannot create any additional resonance structures for this molecule. Given the electronic configuration and bonding nature of \(\mathrm{CO}_{2}\), there is no opportunity to redistribute electrons, and it wouldn't make sense to reduce the bond order between carbon and either of the oxygens. Therefore, there are no resonance structures for \(\mathrm{CO}_{2}\). In conclusion, the only proper representation for \(\mathrm{CO}_{2}\) is its original Lewis structure: O=C=O. There are no resonance structures for this molecule.

Key concepts

Chemical Bonding

Chemical bonding is the fundamental process by which atoms combine to form molecules. It entails the sharing or transferring of valence electrons between atoms to satisfy the octet rule, which states that atoms tend to have eight electrons in their valence shell. The Lewis structure of a molecule like carbon dioxide (CO2) visually represents the valence electrons and bonds between atoms.

In CO2, each oxygen atom forms a double bond with the central carbon atom, utilizing four of the carbon's valence electrons, and leaving the remaining four from the two oxygens to form another double bond. This arrangement achieves full octets around all the atoms, resulting in a stable molecule.

Understanding the connection between Lewis structures and bond formation equips students with a clearer perspective on molecular geometry and stability. Notably, the electron configuration of the atoms—and how they share or transfer electrons—influences the strength and type of bonds formed, playing a critical role in the molecule's properties and behavior.

Resonance Structures

Resonance structures come into play when a single Lewis structure is insufficient to describe the electron distribution within a molecule. In essence, resonance is a tool used to depict molecules where electrons are delocalized or shared by more than two atoms, which cannot be adequately represented by only one fixed structure.

In some molecules, electrons can be distributed in various ways, leading to multiple valid Lewis structures known as resonance structures. These structures are not individual entities but rather a blended hybrid that represents the real electronic structure of the molecule. The actual arrangement of electrons is a weighted average of all the resonance forms.

However, not all molecules exhibit resonance. As seen in the CO2 example, if there are no alternative ways to arrange the electrons without breaking the octet rule or altering the molecule's charge, then the molecule does not have resonance structures. CO2, with a stable double bond on each side of the carbon, has no need for resonance since its electron distribution is already optimal.

Valence Electrons

Valence electrons are the outermost electrons of an atom and are crucial in determining how an atom can bond with others. They are involved in chemical reactions and are represented in the Lewis structure as either shared pairs in bonds or lone pairs.

In the CO2 molecule, carbon has four valence electrons, while each oxygen has six. When drawing the Lewis structure, we calculate the total number of valence electrons (4 from carbon plus 6 from each oxygen), resulting in 16 electrons that need to be distributed in the molecule.

The allocation of these electrons is governed by the principles of chemical bonding and the octet rule, leading to a structure where each atom utilizes its valence electrons to form stable electron configurations. In the case of CO2, this results in double bonds between carbon and each oxygen atom, fulfilling the octet rule without any remaining valence electrons. By visualizing valence electrons, students can better understand the formation of molecules and how electron distribution affects molecular stability.