Question

Consider the molecule \(\mathrm{N}_{2} \mathrm{O}\) (connected \(\mathrm{N}-\mathrm{N}-\mathrm{O})\) (a) Draw the dot diagram. (b) Draw the molecule's three-dimensional shape, and label the numeric value of all bond angles. (c) What is the shape of this molecule? (d) Draw in the individual bond dipole moments. (e) Is the molecule polar? If yes, draw the molecular dipole moment vector.

Short answer

(a) The dot diagram for N2O is: N = N - O with two lone pairs next to each N and four lone pairs next to O. (b) The 3D structure of N2O is linear, with a bond angle of 180° between the N-N-O atoms. (c) The shape of the N2O molecule is linear. (d) Individual bond dipole moments are pointing from the less electronegative N atoms to the more electronegative N and O atoms. (e) N2O is a nonpolar molecule, as the bond dipole moments cancel each other out and there is no net molecular dipole moment.

Step-by-step solution

(a) Drawing the Dot Diagram

To draw the dot diagram, we need to first determine the total number of valence electrons in the molecule. Nitrogen (N) has 5 valence electrons, and Oxygen (O) has 6 valence electrons. Since there are two Nitrogen atoms and one Oxygen atom, the total number of valence electrons is (5 x 2) + 6 = 16. We can arrange these 16 electrons around the atoms as follows: 1. Place one pair of electrons (a "bonding pair") between the two Nitrogen atoms and another pair between one Nitrogen and the Oxygen atom. This forms single bonds between the atoms. 2. Fill in the remaining valence electrons as "lone pairs" around the Nitrogen and Oxygen atoms until each atom has 8 electrons (or a full octet) in its valence shell. The dot diagram for N2O looks like this: N = N - O where two dots next to each N and 4 dots next to O represent the respective lone pair electrons.

(b) Drawing the Three-Dimensional Shape and Labeling Bond Angles

To draw the 3D structure and label the bond angles, we can use the VSEPR (Valence-Shell Electron-Pair Repulsion) theory. The central Nitrogen atom has three electron groups (one from the single bond to the second N atom, one from the single bond to O, and one lone pair). This gives Nitrogen a trigonal pyramidal electron-group geometry but due to the presence of one lone pair, N2O molecule has bent molecular geometry for the central Nitrogen atom. For a linear molecule like N2O, the bond angle at the Oxygen atom between the two Nitrogen atoms is 180°. So the 3D structure can be drawn with a straight line connecting the N-N-O atoms, with the bond angle between the atoms labeled as 180°.

(c) Identifying the Shape of the Molecule

The shape of the N2O molecule is linear, as the bond angle between the Nitrogen and Oxygen atoms is 180° and there is no distortion from the electron groups on the central Nitrogen atom.

(d) Drawing the Individual Bond Dipole Moments

A bond dipole moment is a measure of the polarity in a chemical bond, which is determined by the electronegativity difference between the two atoms involved in the bond. Since Nitrogen and Oxygen are more electronegative than Nitrogen, there are bond dipole moments pointing from the less electronegative N atoms to the more electronegative N and O atoms in each of the two N-N and N-O bonds, respectively. To draw the individual bond dipole moments, draw an arrow from the Nitrogen atom bonded to the central Nitrogen atom towards the central N atom. Similarly, draw an arrow from the central Nitrogen atom towards the Oxygen atom.

(e) Determining Polarity and Drawing the Molecular Dipole Moment Vector

A molecule is considered polar if it has an uneven distribution of electron density that results in a net molecular dipole moment. In the case of N2O, since the molecule is linear and two bond dipole moments are in opposite directions, their effects cancel out, resulting in a net molecular dipole moment of zero. Therefore, N2O is a nonpolar molecule, and there is no need to draw a molecular dipole moment vector.

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom. They are important because they participate in chemical bonding. In the molecule \(_{2}o\), both nitrogen (N) atoms have 5 valence electrons each, while oxygen (O) has 6 valence electrons.

To determine the total number of valence electrons in \(_{2}o\), simply add them up: \(5 \times 2 + 6 = 16\).

These valence electrons can form different types of bonds and arrangements around the atoms, influencing the molecule's structure. Each nitrogen atom in \(_{2}o\) is involved in bonding, contributing to forming single bonds with each other and with oxygen, while maintaining some electrons as lone pairs distributed to respect the octet rule.

VSEPR Theory

VSEPR theory stands for Valence-Shell Electron-Pair Repulsion theory. It is a model used to predict the three-dimensional shape of molecules. The idea is that electron pairs around a central atom will arrange themselves to minimize repulsion, leading to specific molecular geometries.

In \(_{2}o\), the central nitrogen atom has three groups around it: two bonds (one with another nitrogen and one with oxygen) and one lone pair.

Although the electron geometry is technically trigonal pyramidal due to the lone pair, the presence of a lone pair influences the arrangement, giving the molecule a bent (angular) shape for its microscopic structure. Since the surrounding atoms are in a linear arrangement with the oxygen atom at one end, the bond angle remains close to 180°, making the macroscopic shape linear.

Bond Dipole Moment

A bond dipole moment arises from differences in electronegativity between bonded atoms. It is a vector quantity that has both magnitude and direction, indicating the polarity of a bond.

In \(_{2}o\), there are two bonds to consider: the bond between the two nitrogen atoms and the bond between a nitrogen and the oxygen atom.

Since oxygen is more electronegative than nitrogen, there is a dipole moment pointing from the center nitrogen to the oxygen atom. Additionally, a smaller dipole moment points from one nitrogen towards the central nitrogen, but since nitrogen-nitrogen bonds are between identical atoms, the dipole is weaker compared to nitrogen-oxygen.

These dipole moments are depicted as arrows: from less electronegative nitrogen towards both central nitrogen and the most electronegative oxygen, showing the direction of electron density.

Molecular Polarity

Molecular polarity depends on the shape of the molecule and the distribution of bond dipoles within the molecule. When bond dipoles do not cancel out, the molecule is polar and has a net dipole moment.

\(_{2}o\) is a linear molecule with two bond dipole moments that are opposite in direction: one from nitrogen to central nitrogen, and one from central nitrogen to oxygen. However, because it is linear, these two dipole moments cancel each other out.

The result is no net dipole moment, making \(_{2}o\) a nonpolar molecule, even though the individual bonds within the molecule are polar. This distinction is important in predicting interactions with other substances since polarity affects properties like solubility and boiling/melting points.