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Chapter 8: Problem 49

Consider the carbon-oxygen bond in formaldehyde \(\left(\mathrm{CH}_{2} \mathrm{O}\right)\) and carbon monoxide (CO). In which molecule is the CO bond shorter? In which molecule is the CO bond stronger?

Short Answer

Expert verified
CO has the shorter and stronger C-O bond.

Step by step solution

01

Understand Bond Characteristics

Chemical bonds are characterized by bond length and bond strength. A shorter bond length usually indicates a stronger bond due to the increased overlap of bonding orbitals.
02

Analyze Bond Lengths

The bond length for the carbon-oxygen bond in carbon monoxide (CO) is typically around 112 pm (picometers), whereas in formaldehyde (\( \mathrm{CH}_{2} \mathrm{O}\), it is approximately 123 pm. Thus, CO has the shorter bond.
03

Analyze Bond Strengths

Bond strength is often related to bond order. CO has a bond order of 3, indicating a triple bond, due to its triple shared electron pairs. In formaldehyde, the C=O bond is a double bond with a bond order of 2. Therefore, CO has the stronger bond.

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Key Concepts

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

Bond Length
Bond length is a fundamental concept in understanding chemical bonding. It represents the average distance between the nuclei of two bonded atoms. Different factors, like atomic size and bond order, can influence bond length. A shorter bond length generally suggests a stronger connection between the atoms due to more effective overlap of their orbitals.
For instance, in the context of the carbon-oxygen bond, we can see the difference in bond lengths when comparing formaldehyde
  • Formaldehyde (\( \mathrm{CH}_{2} \mathrm{O}\)): Carbon-oxygen bond length is about 123 pm.
  • Carbon monoxide (CO): Carbon-oxygen bond length is approximately 112 pm.
If we compare these two, the CO bond is shorter than the bond in formaldehyde. This shorter bond length often correlates with other properties like bond strength, which we shall discuss next.
Bond Strength
Bond strength depicts how strongly two atoms are bonded together and can often be estimated by the bond energy required to break the bond. Stronger bonds are generally more stable and require more energy to break.
Bond strength is often connected to bond length and bond order. As already discussed, shorter bonds tend to be stronger. In terms of our examples, carbon monoxide (CO) and formaldehyde (
  • CO has a much stronger bond due to its shorter bond length of 112 pm.
  • Formaldehyde (
    • The carbon-oxygen bond has a bond length of 123 pm, making it weaker than in CO.
By relating this to energy, the triple bond in CO, with a significant bond energy, reflects its superior bond strength compared to the double bond in formaldehyde.
Bond Order
Bond order gives us insight into the number of shared electron pairs between two atoms. A higher bond order implies that the atoms share more pairs of electrons, which usually results in stronger and shorter bonds.
In the case of carbon-oxygen bonds, we see distinct differences:
  • Carbon monoxide (CO) has a bond order of 3, meaning it shares three pairs of electrons, forming a strong triple bond.
  • Formaldehyde (
    • It has a bond order of 2, indicating a double bond with two shared pairs of electrons.
Thus, the bond order is a key factor in determining both the bond's strength and length. Higher bond order (like in CO) results in higher bond strength and shorter bond length, leading to a stable and robust bond.

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Most popular questions from this chapter

What are the orders of the \(\mathrm{N}-\mathrm{O}\) bonds in \(\mathrm{NO}_{2}^{-}\) and \(\mathrm{NO}_{2}^{+} ?\) The nitrogen-oxygen bond length in one of these ions is \(110 \mathrm{pm}\) and \(124 \mathrm{pm}\) in the other. Which bond length corresponds to which ion? Explain briefly.

The cyanate ion, \(\mathrm{OCN}^{-}\), has the least electronegative atom, \(\mathrm{C}\), in the center. The very unstable fulminate ion, CNO \(^{-}\), has the same molecular formula, but the \(\mathrm{N}\) atom is in the center. (a) Draw the three possible resonance structures of \(\mathrm{CNO}^{-}\). (b) On the basis of formal charges, decide on the resonance structure with the most reasonable distribution of charge. (c) Mercury fulminate is so unstable it is used in blasting caps for dynamite. Can you offer an explanation for this instability? (Hint: Are the formal charges in any resonance structure reasonable in view of the relative electronegativities of the atoms?)

In each pair of bonds, predict which is shorter. (a) \(\mathrm{B}-\mathrm{Cl}\) or \(\mathrm{Ga}-\mathrm{Cl}\) (b) \(\mathrm{Sn}-\mathrm{O}\) or \(\mathrm{C}-\mathrm{O}\) (c) \(\mathrm{P}-\mathrm{S}\) or \(\mathrm{P}-\mathrm{O}\) (d) \(\mathrm{C}=\mathrm{O}\) or \(\mathrm{C}=\mathrm{N}\)

Draw a Lewis structure for each of the following molecules: (a) chlorodifluoromethane, CHCIF \(_{2}\) (C is the central atom) (b) acetic acid, \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}\) (basic structure pictured below) (IMAGE NOT COPY) (c) acetonitrile, \(\mathrm{CH}_{3} \mathrm{CN}\) (the framework is \(\mathrm{H}_{3} \mathrm{C}-\mathrm{C}-\mathrm{N}\) ) (d) allene, \(\mathrm{H}_{2} \mathrm{CCCH}_{2}\)

Draw the resonance structures for the formate ion, \(\mathrm{HCO}_{2}^{-},\) and find the formal charge on each atom. If an \(\mathrm{H}^{+}\) ion is attached to \(\mathrm{HCO}_{2}^{-}\) (to form formic acid), does it attach to C or O?
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