Chapter 24: Problem 53
Explain why two \(\mathrm{N}\) atoms can form a double bond or a triple bond, whereas two \(\mathrm{P}\) atoms normally can form only a single bond.
Short Answer
Expert verified
Nitrogen forms multiple bonds due to effective orbital overlap; phosphorus forms single bonds due to less effective overlap.
Step by step solution
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Key Concepts
These are the key concepts you need to understand to accurately answer the question.
Atomic Structure
Understanding the atomic structure of elements like nitrogen and phosphorus can help explain their bonding behavior. Each element has a certain number of electrons distributed in various shells around the nucleus. These electrons are arranged in energy levels or shells, with valence electrons being the outermost. In the periodic table, nitrogen and phosphorus belong to Group 15, showcasing similarities in their electronic configurations.
The nitrogen atom has five valence electrons, needing three more to complete its valence shell, or octet. This element is therefore capable of forming up to three covalent bonds for stability. Conversely, phosphorus also has five valence electrons but due to its larger atomic size, it usually utilizes only a portion of these for bonding, often resulting in single bonds.
The size and structure of an atom influence how well it can overlap with others to share electron pairs. Smaller atomic radii, like that of nitrogen, lead to effective overlap, while larger radii, such as that of phosphorus, result in less efficient orbital interaction. This difference is pivotal in determining the type and strength of bonds formed.
The nitrogen atom has five valence electrons, needing three more to complete its valence shell, or octet. This element is therefore capable of forming up to three covalent bonds for stability. Conversely, phosphorus also has five valence electrons but due to its larger atomic size, it usually utilizes only a portion of these for bonding, often resulting in single bonds.
The size and structure of an atom influence how well it can overlap with others to share electron pairs. Smaller atomic radii, like that of nitrogen, lead to effective overlap, while larger radii, such as that of phosphorus, result in less efficient orbital interaction. This difference is pivotal in determining the type and strength of bonds formed.
Sigma Bond
Sigma bonds (\(\sigma\) bonds) are the strongest type of covalent bond and are formed by the head-on overlap of atomic orbitals. They are characterized by sharing a pair of electrons directly between the bonded nuclei. All single bonds in molecules are sigma bonds, and they are integral for forming stable molecular structures.
In the case of nitrogen, when two atoms form a bond, a sigma bond is the foundation of this interaction. Nitrogen's small atomic size allows this bond to be particularly strong due to efficient overlap of orbitals. Similarly, phosphorus atoms form sigma bonds, but the effectiveness is reduced due to their larger atomic radius.
Sigma bonds allow for free rotation around the bond axis, providing flexibility to single-bonded molecules. However, when multiple bonds involving pi bonds are present (as in nitrogen), this rotation is restricted.
In the case of nitrogen, when two atoms form a bond, a sigma bond is the foundation of this interaction. Nitrogen's small atomic size allows this bond to be particularly strong due to efficient overlap of orbitals. Similarly, phosphorus atoms form sigma bonds, but the effectiveness is reduced due to their larger atomic radius.
Sigma bonds allow for free rotation around the bond axis, providing flexibility to single-bonded molecules. However, when multiple bonds involving pi bonds are present (as in nitrogen), this rotation is restricted.
Pi Bond
Pi bonds (\(\pi\) bonds) arise from the side-to-side overlap of atomic orbitals, and they are found in double and triple bonds alongside sigma bonds. Unlike sigma bonds, pi bonds do not allow for rotation because the orbitals involved are parallel rather than along a single axis.
Nitrogen can form strong pi bonds due to its small atomic size, allowing for effective overlap. This capacity leads to the formation of double (one sigma and one pi bond) or triple bonds (one sigma and two pi bonds). The additional pi bonds give these molecules more rigidity compared to molecules with only sigma bonds.
Phosphorus struggles to form efficient pi bonds. Its larger atomic size prevents close enough overlap to maintain the bonds' stability, making it less likely to support multiple bonds. The weak interactions of pi bonds in larger atoms like phosphorus often make such bonds energetically unfavorable.
Nitrogen can form strong pi bonds due to its small atomic size, allowing for effective overlap. This capacity leads to the formation of double (one sigma and one pi bond) or triple bonds (one sigma and two pi bonds). The additional pi bonds give these molecules more rigidity compared to molecules with only sigma bonds.
Phosphorus struggles to form efficient pi bonds. Its larger atomic size prevents close enough overlap to maintain the bonds' stability, making it less likely to support multiple bonds. The weak interactions of pi bonds in larger atoms like phosphorus often make such bonds energetically unfavorable.
Valence Electrons
Valence electrons are the electrons in the outermost shell of an atom that are involved in chemical bonding. They determine how an element behaves in a chemical reaction, dictating the possible types and numbers of bonds that it can form.
For nitrogen, the five valence electrons make it highly reactive and versatile in forming multiple bonds. It actively seeks to fill its octet, leading to the formation of strong, stable triple bonds in configurations like \( \text{N}_2 \). This ability is due to its small size and effective overlap capabilities.
Phosphorus, though also having five valence electrons, behaves differently. Its larger atomic size hinders the formation of strong overlapping bonds needed for more than a single bond. The availability yet underuse of its valence electrons for multiple bonds reflects its tendency to form primarily single bonds in compounds.
For nitrogen, the five valence electrons make it highly reactive and versatile in forming multiple bonds. It actively seeks to fill its octet, leading to the formation of strong, stable triple bonds in configurations like \( \text{N}_2 \). This ability is due to its small size and effective overlap capabilities.
Phosphorus, though also having five valence electrons, behaves differently. Its larger atomic size hinders the formation of strong overlapping bonds needed for more than a single bond. The availability yet underuse of its valence electrons for multiple bonds reflects its tendency to form primarily single bonds in compounds.
