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Chapter 20: Problem 36

In addition to \(\mathrm{CO}_{2}\), two additional stable oxides of carbon form. The space-filling models for \(\mathrm{CO}_{2}\) and the other two stable oxides are: What are the formulas for the two additional stable oxides of carbon? Explain the bonding in each of these two forms using the localized electron model.

Short Answer

Expert verified
The two additional stable oxides of carbon are carbon monoxide (CO) and carbon suboxide (C3O2). In carbon monoxide, a triple bond (1蟽 and 2蟺 bonds) between carbon and oxygen satisfies the octet rule for both atoms. Carbon suboxide has a linear structure O=C=C=C=O, in which the central carbon atom is doubly bonded (1蟽 and 1蟺 bond) to the two adjacent carbon atoms, which are also doubly bonded to oxygen atoms.

Step by step solution

01

Identify the two additional stable oxides of carbon

The two additional stable oxides of carbon are carbon monoxide (CO) and carbon suboxide (C3O2).
02

Describe the bonding in CO using the localized electron model

First, let's analyze carbon monoxide (CO). Carbon has 4 valence electrons, and oxygen has 6 valence electrons. Using the localized electron model, we can represent the bonding as follows: 1. Place a single bond (2 electrons shared) between the carbon and oxygen atoms. 2. Assign the remaining 6 electrons to the oxygen atom in three lone pairs. 3. Assign the remaining 2 electrons to the carbon atom in a lone pair. 4. Since carbon needs to have 8 electrons in its outer shell to achieve an octet, it forms a triple bond with the oxygen atom. The Lewis structure would look like this: C鈮 In the localized electron model, carbon monoxide consists of a triple bond between carbon and oxygen, with one sigma (蟽) bond and two pi (蟺) bonds. It has no formal charge, as the octet rule is satisfied for both atoms.
03

Describe the bonding in C3O2 using the localized electron model

Now, let's analyze carbon suboxide (C3O2). The molecule consists of 3 carbon atoms and 2 oxygen atoms. Carbon has 4 valence electrons, and oxygen has 6 valence electrons. Using the localized electron model, we can represent the bonding as follows: 1. Place a single bond (2 electrons shared) between each pair of carbon atoms. 2. Place a single bond (2 electrons shared) between each carbon at the ends and an oxygen atom. 3. Assign the remaining electrons in such a way that each atom achieves an octet. 4. The resulting Lewis structure would look like this: O=C=C=C=O In the localized electron model, carbon suboxide consists of a central carbon atom doubly bonded to two adjacent carbon atoms, which are in turn doubly bonded to oxygen atoms. There is a 蟽 bond between adjacent carbon atoms and 蟺 bonds between the adjacent double-bonded atoms (C=C and C=O). Each atom has no formal charge, as the octet rule is satisfied.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Localized Electron Model
The localized electron model helps us understand how atoms bond together. In this model, we represent molecules by showing the sharing of electrons between atoms.
Each shared pair of electrons forms a bond, represented in diagrams as lines or sticks. - It is most effective for explaining how covalent bonds are formed. - The model highlights which electrons participate in bonding and which form lone pairs. Unlike more complex models, the localized electron model focuses on fixed positions, making it accessible and simpler for beginners learning molecular structures.
Lewis Structure
A Lewis structure is a diagrammatic way to represent the atoms and bonding in a molecule. It shows the arrangement of valence electrons around atoms. Structure is key in learning how atoms share electrons to reach a stable configuration. - Lewis structures use dots to represent valence electrons. - Lines between atoms symbolize shared pairs, i.e., bonds. For example, in carbon monoxide (CO), a Lewis structure clearly illustrates the triple bond between carbon and oxygen. By providing a visual guide, Lewis structures make complex bonding concepts more accessible for students.
Triple Bond
A triple bond involves the sharing of three pairs of electrons between two atoms. This is a strong type of chemical bond and is essential for the stability of molecules like carbon monoxide. - Consists of one sigma (蟽) bond and two pi (蟺) bonds.
- Provides high bond strength and short bond length. In the case of CO, a triple bond forms due to carbon's need for more electrons to complete its octet. This triple bond ensures both carbon and oxygen achieve their stable electron configurations, making CO a remarkably stable molecule.
Carbon Monoxide
Carbon monoxide is a well-known compound with the formula CO. It features a triple bond between carbon and oxygen, which provides its remarkable stability. - It has one carbon atom and one oxygen atom. - Formed by sharing 8 valence electrons between the two atoms. Due to its stable triple bond, CO remains a prevalent molecule in various industrial processes. However, despite its uses, it is also a toxic gas, such as when burning fuels without adequate ventilation.
Carbon Suboxide
Carbon suboxide, with the formula C3O2, is an intriguing molecule with unique bonding. - It includes three carbon atoms and two oxygen atoms.
- Each carbon atom connects through double bonds, creating a linear chain. Bonded as O=C=C=C=O, this structure is fascinating in its arrangement but less common as it can rapidly polymerize under certain conditions. Understanding carbon suboxide enriches the study of organic chemistry, showcasing different ways carbon-based compounds can exhibit stability.

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Most popular questions from this chapter

Although nitrogen trifluoride \(\left(\mathrm{NF}_{3}\right)\) is a thermally stable compound, nitrogen triiodide \(\left(\mathrm{NI}_{3}\right)\) is known to be a highly explosive material. \(\mathrm{NI}_{3}\) can be synthesized according to the equation $$ \mathrm{BN}(s)+3 \mathrm{IF}(g) \longrightarrow \mathrm{BF}_{3}(g)+\mathrm{NI}_{3}(g) $$ a. What is the enthalpy of formation for \(\mathrm{NI}_{3}(s)\) given the enthalpy of reaction \((-307 \mathrm{~kJ})\) and the enthalpies of formation for \(\mathrm{BN}(s)(-254 \mathrm{~kJ} / \mathrm{mol}), \mathrm{IF}(g)(-96 \mathrm{~kJ} / \mathrm{mol})\), and \(\mathrm{BF}_{3}(g)(-1136 \mathrm{~kJ} / \mathrm{mol})\) ? b. It is reported that when the synthesis of \(\mathrm{NI}_{3}\) is conducted using 4 moles of IF for every 1 mole of \(\mathrm{BN}\), one of the by-products isolated is \(\left[\mathrm{IF}_{2}\right]^{+}\left[\mathrm{BF}_{4}\right]_{-}^{-} .\) What are the molecular geometries of the species in this by-product? What are the hybridizations of the central atoms in each species in the by-product?

Captain Kirk has set a trap for the Klingons who are threatening an innocent planet. He has sent small groups of fighter rockets to sites that are invisible to Klingon radar and put a decoy in the open. He calls this the "fishhook" strategy. Mr. Spock has sent a coded message to the chemists on the fighters to tell the ships what to do next. The outline of the message is Fill in the blanks of the message using the following clues. (1) Symbol of the halogen whose hydride has the second highest boiling point in the series of HX compounds that are hydrogen halides. (2) Symbol of the halogen that is the only hydrogen halide, \(\mathrm{HX}\), that is a weak acid in aqueous solution. (3) Symbol of the element whose existence on the sun was known before its existence on earth was discovered. (4) The Group \(5 \mathrm{~A}\) element in Table \(20.13\) that should have the most metallic character. (5) Symbol of the Group 6 A element that, like selenium, is a semiconductor. (6) Symbol for the element known in rhombic and monoclinic forms. (7) Symbol for the element that exists as diatomic molecules in a yellow-green gas when not combined with another element. (8) Symbol for the most abundant element in and near the earth's crust. (9) Symbol for the element that seems to give some protection against cancer when a diet rich in this element is consumed. (10) Symbol for the smallest noble gas that forms compounds with fluorine having the general formula \(\mathrm{AF}_{2}\) and \(\mathrm{AF}_{4}\) (reverse the symbol and split the letters as shown). (11) Symbol for the toxic element that, like phosphorus and antimony, forms tetrameric molecules when uncombined with other elements (split the letters of the symbol as shown). (12) Symbol for the element that occurs as an inert component of air but is a very prominent part of fertilizers and explosives.

Write balanced equations describing the reaction of \(\mathrm{Sr}\) with each of the following: \(\mathrm{O}_{2}, \mathrm{~S}, \mathrm{Cl}_{2}, \mathrm{P}_{4}, \mathrm{H}_{2}, \mathrm{H}_{2} \mathrm{O}\), and \(\mathrm{HCl}\).

Explain why \(\mathrm{HF}\) is a weak acid, whereas \(\mathrm{HCl}, \mathrm{HBr}\), and \(\mathrm{HI}\) are all strong acids.

What is a disproportionation reaction? Use the following reduction potentials \(\mathrm{ClO}_{3}^{-}+3 \mathrm{H}^{+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{HClO}_{2}+\mathrm{H}_{2} \mathrm{O} \quad \mathscr{E}^{\circ}=1.21 \mathrm{~V}\) \(\mathrm{HClO}_{2}+2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{HClO}+\mathrm{H}_{2} \mathrm{O} \quad \mathscr{E}^{\circ}=1.65 \mathrm{~V}\) to predict whether \(\mathrm{HClO}_{2}\) will disproportionate.
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