Many structures of phosphorus-containing compounds are drawn with some \(\mathrm{P}=0\) bonds. These bonds are not the typical \(\pi\) bonds we've considered, which involve the overlap of two \(p\) orbitals. Instead, they result from the overlap of a \(d\) orbital on the phosphorus atom with a \(p\) orbital on oxygen. This type of \(\pi\) bonding is sometimes used as an explanation for why \(\mathrm{H}_{3} \mathrm{PO}_{3}\) has the first structure below rather than the second: Draw a picture showing how a \(d\) orbital and a \(p\) orbital overlap to form a \(\pi\) bond.

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Phosphorus-containing compounds often contain P=O bonds due to a unique type of π bonding that comes from the overlap of a phosphorus atom's d orbital and an oxygen atom's p orbital. For instance, in H₃PO₃, we see this bonding in its first structure: P(=O)(OH)₂. To visualize this overlap, we can consider a d orbital as a four-lobed shape (taking d(z^2) as an example), with two lobes on the z-axis and a torus around the xy plane. A p orbital can be thought of as a two-lobed shape (taking p(z) as an example), with lobes extending above and below the z-axis. When overlapping, the positive lobe of the p orbital aligns with one of the positive lobes of the d orbital, and the negative lobe of the p orbital aligns with the negative lobe of the d orbital. Through this in-phase overlap, a π bond forms between the phosphorus and oxygen atoms, resulting in P=O bonds in compounds like H₃PO₃.

Step by step solution

01

Understanding P=O bonds in phosphorus-containing compounds

Phosphorus-containing compounds often have P=O bonds that involve a type of π bonding not commonly seen in other compounds. This atypical π bonding is the result of the overlap between a d orbital on the phosphorus atom and a p orbital on the oxygen atom.
02

Visualizing the first structure of H₃PO₃

We are given that H₃PO₃ has the following first structure: P(=O)(OH)₂. In this structure, there is a P=O bond, where a d orbital on phosphorus and a p orbital on oxygen are overlapping to form a π bond.
03

Drawing the overlapping orbitals

To draw a picture of how a d and p orbital overlap to form a π bond, consider the following: - A d orbital is represented as a four-lobed shape oriented along the coordinate axes (let's take the d(z^2) orbital as an example). It has two lobes on the z-axis and a torus around the xy plane. - A p orbital is represented as a two-lobed shape aligned along one of the coordinate axes (let's take p(z) being the oxygen-contributing orbital). In this case, the p orbital has lobes extending above and below the z-axis. When overlapping, the positive lobe of the oxygen's p orbital aligns with one of the positive lobes (say the upper) of the phosphorus's d orbital, while the negative lobe of the p orbital aligns with the negative lobe (the lower) of the d orbital. Similarly, the oxygen's p orbital negative side aligns with the negative side of the phosphorus's d orbital. As a result, the d and p orbitals' in-phase overlap forms a π bond between the phosphorus and oxygen atoms, allowing for the formation of P=O bonds in compounds like H₃PO₃.

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Most popular questions from this chapter

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