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Chapter 9: Problem 11

In the hybrid orbital model, compare and contrast \(\sigma\) bonds with \(\pi\) bonds. What orbitals form the \(\sigma\) bonds and what orbitals form the \(\pi\) bonds? Assume the \(z\) -axis is the internuclear axis.

Short Answer

Expert verified
In the hybrid orbital model, \(\sigma\) bonds are formed by end-to-end overlap of atomic orbitals, such as s-s, s-p, and p-p (specifically p\(_z\)) combinations, while \(\pi\) bonds are formed by side-by-side overlap of p\(_x\) or p\(_y\) orbitals. The \(\sigma\) bond has electron density concentrated along the internuclear axis, while the \(\pi\) bond has two regions of electron density above and below the internuclear axis.

Step by step solution

01

Definition of Sigma and Pi bonds.

A \(\sigma\) bond is a covalent bond formed between two atoms by the end-to-end (head-on) overlapping of their atomic orbitals. This overlapping forms a region of high electron density along the internuclear axis, which is the z-axis in this case. A \(\pi\) bond, on the other hand, is another type of covalent bond formed between two atoms when their atomic orbitals overlap side-by-side. This overlapping creates two regions of high electron density above and below the internuclear axis.
02

Orbitals that form Sigma bonds

In the hybrid orbital model, the atomic orbitals involved in the formation of \(\sigma\) bonds are: 1. s orbitals: An s orbital can overlap with another s orbital head-on, forming an \(\text{s-s}\) \(\sigma\) bond. 2. s and p orbitals: An s orbital can also overlap with a p orbital along the internuclear axis, resulting in an \(\text{s-p}\) \(\sigma\) bond. 3. p orbitals: A p orbital, specifically the one aligned along the z-axis (p\(_z\)), can overlap with another p\(_z\) orbital head-on, creating a \(\text{p-p}\) \(\sigma\) bond.
03

Orbitals that form Pi bonds

In the hybrid orbital model, the atomic orbitals involved in the formation of \(\pi\) bonds are: 1. p orbitals: The p orbitals responsible for forming \(\pi\) bonds are the ones not aligned along the internuclear axis (z-axis). These are the p\(_x\) and p\(_y\) orbitals. They can interact with each other in a side-by-side fashion to form a \(\pi\) bond. In summary, the hybrid orbital model shows that \(\sigma\) and \(\pi\) bonds are formed by different types of atomic orbital overlap. \(\sigma\) bonds are created through end-to-end overlap of orbitals (like s, s-p, p-p), while \(\pi\) bonds are established by side-by-side overlap of p orbitals (p\(_x\) or p\(_y\)).

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Most popular questions from this chapter

What are the relationships among bond order, bond energy, and bond length? Which of these quantities can be measured?

The atoms in a single bond can rotate about the internuclear axis without breaking the bond. The atoms in a double and triple bond cannot rotate about the internuclear axis unless the bond is broken. Why?

Why are \(d\) orbitals sometimes used to form hybrid orbitals? Which period of elements does not use \(d\) orbitals for hybridization? If necessary, which \(d\) orbitals \((3 d, 4 d, 5 d, \text { or } 6 d)\) would sulfur use to form hybrid orbitals requiring \(d\) atomic orbitals? Answer the same question for arsenic and for iodine.

In which of the following diatomic molecules would the bond strength be expected to weaken as an electron is removed? $$ \begin{array}{ll}{\text { a. } H_{2}} & {\text { c. } C_{2}^{2-}} \\ {\text { b. } B_{2}} & {\text { d. OF }}\end{array} $$

The molecules \(\mathrm{N}_{2}\) and \(\mathrm{CO}\) are isoelectronic but their properties are quite different. Although as a first approximation we often use the same MO diagram for both, suggest how the \(\mathrm{MOs}\) in \(\mathrm{N}_{2}\) and \(\mathrm{CO}\) might be different.
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