Question

Which has the greater \(\mathrm{O}-\mathrm{N}-\mathrm{O}\) bond angle, \(\mathrm{NO}_{2}^{-}\) or \(\mathrm{NO}_{2}^{+} ?\) Explain briefly.

Short answer

The NO2+ ion has a greater bond angle (180°) than NO2-, due to its linear structure and absence of lone pairs.

Step-by-step solution

Understanding the Basics

The bond angle in a molecule is influenced by the electron pair geometry around the central atom. In both \(\mathrm{NO}_2^-\) and \(\mathrm{NO}_2^+\), nitrogen is the central atom, and the differing electron arrangement affects the bond angles.

Analyzing NO2-

For the \(\mathrm{NO}_2^-\) ion, nitrogen is surrounded by three pairs of electrons: two bonding pairs and one lone pair. Lone pairs repel more strongly than bonding pairs, causing the bond angle to be less than \(120^\circ\), which is typical of a trigonal planar arrangement.

Analyzing NO2+

For the \(\mathrm{NO}_2^+\) ion, nitrogen is surrounded by two bonding pairs and no lone pairs. This leads to a linear geometry with a bond angle of exactly \(180^\circ\).

Comparing the Bond Angles

Since lone pairs repel more strongly, \(\mathrm{NO}_2^-\) has a smaller bond angle due to the presence of the lone pair, which reduces the angle from the ideal \(120^\circ\). In contrast, \(\mathrm{NO}_2^+\) has a bond angle of \(180^\circ\) due to its linear structure and lack of lone pairs.

Conclusion

The \(\mathrm{NO}_2^+\) ion has the greater \(\mathrm{O}-\mathrm{N}-\mathrm{O}\) bond angle compared to \(\mathrm{NO}_2^-\), because it lacks lone pairs and has a linear shape.

Key concepts

Electron Pair Geometry

Electron pair geometry describes the spatial arrangement of all electron pairs around a central atom. This encompasses both bonding pairs, which connect atoms, and lone pairs, which are non-bonding. Understanding this geometry is crucial to predict molecule shapes and bond angles.
In the context of \( \text{NO}_2^- \) and \( \text{NO}_2^+ \):

• For \( \text{NO}_2^- \), the nitrogen atom is surrounded by two bonding pairs and one lone pair of electrons.
• For \( \text{NO}_2^+ \), the nitrogen atom has two bonding pairs and zero lone pairs.

The presence or absence of lone pairs has a significant impact on the shape and bond angle, as lone pairs repel more strongly than bonding pairs.

Trigonal Planar Arrangement

Trigonal planar arrangement occurs when three regions of electron density surround a central atom, aiming to minimize repulsion by being equidistant. Typically, this results in bond angles of \( 120^\circ \).
The \( \text{NO}_2^- \) ion tries to adopt this geometry, but its lone pair causes deviations.
The lone pair's repulsive force is stronger, thus reducing the bond angle below \( 120^\circ \).
As electrons try to balance repulsive forces, the geometric arrangement shifts away from being purely trigonal planar.

Linear Geometry

Linear geometry is noted when two regions of electron density surround a central atom. These regions align themselves on opposite sides to minimize repulsion, creating a straight line.
In \( \text{NO}_2^+ \), two bonding pairs with no lone pairs result in a perfect linear shape.
As a result, the bond angle is \( 180^\circ \). This is a perfect configuration in terms of electron pair repulsion as there are no lone pairs to alter this linear shape.

Lone Pair Repulsion

Lone pair repulsion has a significant influence on molecular shape. This concept states that lone pairs occupy more space than bonding pairs due to their non-shared nature, thereby exerting greater repulsive forces.
With \( \text{NO}_2^- \), the presence of a lone pair causes a reduction in bond angle, as these pairs push the oxygen atoms closer together than they would be in an ideal configuration.

• Lone pairs increase electron cloud density, creating more repulsion.
• This drives adjacent bonding pairs closer, reducing the overall bond angle.

Understanding this principle is essential to interpreting how molecules deviate from ideal angles based on their geometry and electron pair arrangement.