Question

The three most stable oxides of carbon are carbon monoxide \((\mathrm{CO}),\) carbon dioxide \(\left(\mathrm{CO}_{2}\right),\) and carbon suboxide \(\left(\mathrm{C}_{3} \mathrm{O}_{2}\right) .\) The space-filling models for these three compounds are For each oxide, draw the Lewis structure, predict the molecular structure, and describe the bonding (in terms of the hybrid orbitals for the carbon atoms).

Short answer

The Lewis structures for the three most stable oxides of carbon are: CO: \( \: \: \: \: \: \: \: \: \: C \equiv O : \) CO2: \( \: : O = C = O : \) C3O2: \( \: : O = C = C = C = O : \) The molecular structures of CO, CO2, and C3O2 are all linear. In CO, the carbon atom is sp hybridized, forming a triple bond with an oxygen atom. In CO2, the central carbon atom is sp hybridized, forming double bonds with two oxygen atoms. In C3O2, the terminal carbon atoms are sp2 hybridized, forming double bonds with oxygen atoms and single bonds with the middle carbon atom. The middle carbon atom is sp hybridized, forming double bonds with the two adjacent carbon atoms.

Step-by-step solution

Draw the Lewis Structures

To draw the Lewis structures, we need to follow these steps: 1. Calculate the total number of valence electrons for each compound. 2. Write the symbols for each atom, arranging them to show their connectivity. 3. Distribute the valence electrons as shared (bonding) pairs and lone (non-bonding) pairs. For CO: - Carbon and oxygen have 4 and 6 valence electrons, respectively, so we have a total of 10 valence electrons. - The Lewis structure for CO is: \( \: \: \: \: \: \: \: \: \: \: C \equiv O : \) For CO2: - Carbon and oxygen have 4 and 6 valence electrons, respectively, so for two oxygen atoms and one carbon atom, we have a total of 16 valence electrons. - The Lewis structure for CO2 is: \( \: : O = C = O : \) For C3O2: - Carbon has 4 valence electrons, and oxygen has 6 valence electrons, so for three carbon atoms and two oxygen atoms, we have a total of 22 valence electrons. - The Lewis structure for C3O2 is: \( \: : O = C = C = C = O : \)

Predict the Molecular Structure

To predict the molecular structure, we need to determine the electron domain geometry for each central atom. For CO: - The carbon atom is bonded to an oxygen atom through a triple bond, forming a linear electron domain geometry. - Thus, the molecular structure of CO is linear. For CO2: - The carbon atom is doubly bonded to two oxygen atoms, forming a linear electron domain geometry. - Thus, the molecular structure of CO2 is linear. For C3O2: - In C3O2, we have three carbon atoms and two oxygen atoms connected in a linear chain. - The two terminal carbon atoms are doubly bonded to an oxygen atom, forming a linear electron domain geometry. - The middle carbon atom has two double bonds, one to each adjacent carbon atom, also forming a linear electron domain geometry. - Thus, the molecular structure of C3O2 is linear.

Determine Carbon Hybridization and Bonding

To determine the hybridization of the carbon atoms and describe the bonding in terms of hybrid orbitals, we will examine each compound individually. For CO: - The carbon atom in CO forms a triple bond with an oxygen atom, which consists of one sigma bond and two pi bonds. - The carbon atom is sp hybridized, forming two sp orbitals. - One of the sp orbitals forms a sigma bond with an oxygen atom's 2p orbital, and the other sp orbital has a lone pair of electrons. - The two pi bonds are formed between the unhybridized 2p orbitals of the carbon and oxygen atoms. For CO2: - The central carbon atom in CO2 forms two double bonds with two oxygen atoms, each double bond consisting of one sigma bond and one pi bond. - The central carbon atom is sp hybridized, forming two sp orbitals. - Each oxygen atom forms a sigma bond with one of the sp orbitals of the central carbon atom. - The pi bonds are formed between the unhybridized 2p orbitals of the carbon and oxygen atoms. For C3O2: - The terminal carbon atoms in C3O2 form a double bond with an oxygen atom, which consists of one sigma bond and one pi bond. They also form a sigma bond with the middle carbon atom. - The terminal carbon atoms are sp2 hybridized, forming three sp2 orbitals. - One of the sp2 orbitals forms a sigma bond with an oxygen atom's 2p orbital, another sp2 orbital forms a sigma bond with the middle carbon atom, and the third sp2 orbital has a lone pair of electrons. - The pi bond with the oxygen atom is formed between the unhybridized 2p orbitals of the carbon and oxygen atoms. - The middle carbon atom in C3O2 forms two double bonds with two adjacent carbon atoms, each double bond consisting of one sigma bond and one pi bond. - The middle carbon atom is sp hybridized, forming two sp orbitals. - Each terminal carbon atom forms a sigma bond with one of the sp orbitals of the middle carbon atom. - The pi bonds are formed between the 2p orbitals of the middle and terminal carbon atoms.

Key concepts

Carbon Oxides

In chemistry, carbon oxides are carbon-containing compounds where carbon atoms are bonded to oxygen atoms. The three most stable carbon oxides are carbon monoxide \(\mathrm{CO}\), carbon dioxide \(\mathrm{CO}_{2}\), and carbon suboxide \(\mathrm{C}_{3}\mathrm{O}_{2}\). These compounds are noteworthy because of their occurrence and roles in various natural and industrial processes.
- Carbon monoxide (CO) is a colorless, odorless gas known for its potential toxicity. It is produced during the incomplete combustion of carbon-containing materials.
- Carbon dioxide (CO₂) is a crucial component of Earth's atmosphere and plays a big role in photosynthesis and the global carbon cycle.
- Carbon suboxide (C₃O₂) is less common and can be found in specialized applications, such as in organic synthesis, due to its reactive properties.
Understanding the distinct structures and bonding of these oxides helps in grasping their functional properties and potential applications or impacts in the environment.

Molecular Structure

The molecular structure of a compound determines how atoms are spatially arranged and bonded together. The molecular structure can be predicted using Lewis structures and electron domain geometries. Let's look at each oxide:
- **Carbon Monoxide (CO):** CO consists of a carbon atom triple-bonded to an oxygen atom, creating a linear structure. The triple bond is strong and involves sharing three pairs of electrons between carbon and oxygen.
- **Carbon Dioxide (CO₂):** CO₂ structures as a linear molecule with carbon singly bonded to two oxygen atoms. Each carbon-oxygen bond is a double bond involving an equal sharing of electrons, maintaining a straight 180-degree angle.
- **Carbon Suboxide (C₃O₂):** Interestingly, C₃O₂ forms a chain-like linear structure. It has two terminal carbon atoms each double-bonded to an oxygen atom, with the terminal carbons connected to a central carbon via double bonds. This molecule maintains straight-line geometry due to the nature of the bonding between the atoms.
These different structures affect physical properties such as polarity, reactivity, and interaction with other molecules.

Carbon Hybridization

The concept of hybridization describes how atomic orbitals mix to form new, equivalent hybrid orbitals suitable for molecule bonding. Different types of hybridization lead to different molecular shapes:
- **In CO:** The carbon atom is sp hybridized, which means it's using one s orbital and one p orbital to create two sp orbitals. One sp orbital is used to form a sigma bond with oxygen's 2p orbital, while the other holds a lone pair. The triple bond's remaining pi bonds arise from unhybridized p orbitals. - **In CO₂:** The central carbon in CO₂ is also sp hybridized. Its sp orbitals form sigma bonds with each oxygen. The orientation of the hybrid orbitals explains CO₂'s linear shape and the symmetrical distribution of charge. - **In C₃O₂:** The terminal carbons are sp² hybridized, which allows them to form three sigma bonds: one with oxygen, one with the middle carbon atom, and a lone pair. The middle carbon uses sp hybridization, aligning with the linear structure and enabling double bonding with adjacent carbons.
Understanding hybridization not only explains molecular shape but also provides insights into bond angles and potential interactions with other molecules.

Bonding in Molecules

Bonding in molecules is critical in determining stability, reactivity, and physical properties. Bond types include sigma and pi bonds, each associated with different orbital interactions:
- **Carbon Monoxide (CO):** Features a strong triple bond, comprising one sigma and two pi bonds. The sigma bond arises from the overlap of carbon's sp orbital with oxygen's 2p orbital. The two pi bonds come from parallel overlap of the remaining p orbitals on each atom, reinforcing bond strength and energy. - **Carbon Dioxide (CO₂):** Contains two double bonds between carbon and oxygen atoms. Each bond consists of one sigma bond and one pi bond. The sigma bonds occur via the overlap of each oxygen’s orbital with carbon's sp orbital, while pi bonds arise from side-overlapping of their unhybridized orbitals. - **Carbon Suboxide (C₃O₂):** Showcases diverse bonding arrangements involving single and double bonds. Terminal carbons have sigma bonds with both oxygen and middle carbon atoms, each double bond supported by additional pi bond configurations that contribute to the chain’s stability.
Bonding considerations provide insight into molecular properties, helping predict behavior in various chemical environments.