Question

Write the Lewis structure for each of the following species, and describe its geometry: (a) \(\mathrm{NH}_{4}^{+}\), (b) \(\mathrm{NO}_{2}^{-}\), (c) \(\mathrm{N}_{2} \mathrm{O}\), (d) \(\mathrm{NO}_{2}\).

Short answer

(a) \(\mathrm{NH}_{4}^{+}\): Lewis structure has N with 4 single bonds to H atoms. The geometry is tetrahedral. (b) \(\mathrm{NO}_{2}^{-}\): Lewis structure has N with 2 single bonds to O atoms and one lone pair. The geometry is bent. (c) \(\mathrm{N}_{2} \mathrm{O}\): Lewis structure has a triple bond between two N atoms and a single bond between N and O. The geometry is linear. (d) \(\mathrm{NO}_{2}\): Lewis structure has N with 2 single bonds to O atoms and one lone pair. The geometry is bent.

Step-by-step solution

Count the valence electrons

To begin, calculate the total number of valence electrons for each molecule/ion. For \(\mathrm{NH}_{4}^{+}\), Nitrogen has 5 valence electrons, each Hydrogen has 1, and also consider the positive charge. For \(\mathrm{NO}_{2}^{-}\), Nitrogen has 5 valence electrons, each Oxygen has 6, and also consider the negative charge. For \(\mathrm{N}_{2} \mathrm{O}\), each Nitrogen has 5 valence electrons, and Oxygen has 6. For \(\mathrm{NO}_{2}\), Nitrogen has 5 valence electrons, and each Oxygen has 6.

Drawing Lewis structures for each species

Now, let's draw the Lewis structures for each species, following the octet rule for each atom: (a) \(\mathrm{NH}_{4}^{+}\) Total valence electrons = 5 (from N) + 4x1 (from H) - 1 (due to the +1 charge) = 8 Place 4 single bonds between N and H atoms, so each H has 1 pair of electrons, and the remaining electron pairs are around N. (b) \(\mathrm{NO}_{2}^{-}\) Total valence electrons = 5 (from N) + 2x6 (from O) + 1 (due to the -1 charge) = 18 Place single bonds between N and each O atom. Add lone pairs on O atoms to complete their octet. As a result, 7th electron of N will not be paired, thus creating a free radical structure. (c) \(\mathrm{N}_{2} \mathrm{O}\) Total valence electrons = 2x5 (from N) + 6 (from O) = 16 Place a triple bond between the two N atoms, and a single bond between N and O. Add lone pairs on O atom to complete its octet. (d) \(\mathrm{NO}_{2}\) Total valence electrons = 5 (from N) + 2x6 (from O) = 17 Place single bonds between N and each O atom, and add lone pairs on O atoms to complete their octet. As a result, the extra electron from N will unpair.

Describing the geometry using VSEPR theory

Now that we have the Lewis structures, we can classify the molecular geometries using VSEPR theory. (a) \(\mathrm{NH}_{4}^{+}\) : Tetrahedral Geometry (4 bonding regions, 0 lone pairs) (b) \(\mathrm{NO}_{2}^{-}\) : Bent Geometry (2 bonding regions, 1 lone pair) (c) \(\mathrm{N}_{2} \mathrm{O}\) : Linear Geometry (2 bonding regions, 0 lone pairs) (d) \(\mathrm{NO}_{2}\) : Bent Geometry (2 bonding regions, 1 lone pair)

Key concepts

Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom and play a crucial role in chemical bonding and reactions. Understanding them is vital for creating accurate Lewis structures, which are diagrams that represent the bonding between atoms and the lone pairs of electrons that may exist.

For example, consider Nitrogen (N), which has 5 valence electrons, and Hydrogen (H), with 1 valence electron. In the ammonium ion (H_{4}^{+}), Nitrogen's valence electrons, along with those from Hydrogen atoms, form bonds, adhering to the rule where the total number of valence electrons is affected by the ion's charge. This process reflects the way that elements attain stability by fulfilling the octet rule, which states that atoms are most stable when they have eight valence electrons.

In drawing Lewis structures, first count the valence electrons of each atom, then consider the charges on ions, and proceed to create bonds while complying with the octet rule.

VSEPR Theory

The Valence Shell Electron Pair Repulsion (VSEPR) theory is a model used to predict the shape of individual molecules based on the repulsion between the electron pairs in the valence shell of an atom. The basics of VSEPR theory rest on the idea that electron pairs will arrange themselves as far apart as possible to minimize repulsion.

Using VSEPR theory, we inferred the shapes for various species in our exercise. For instance, the ammonium ion (H_{4}^{+}) has no lone electron pairs around the Nitrogen atom, and it is surrounded symmetrically by four Hydrogen atoms, which places it into a tetrahedral geometry. In contrast, the Nitrite ion (O_{2}^{-}) and the Nitrogen dioxide (O_{2}) molecules both have a 'bent' geometry due to one lone pair of electrons pushing away the bonding pairs.

VSEPR theory is incredibly useful for understanding how molecular geometry is influenced by the distribution of electrons around a central atom.

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. The geometry is determined based on the number of bonding pairs and lone pairs of electrons around the central atom. The shapes of the molecules greatly influence the molecule's properties, including reactivity, polarity, and intermolecular forces.

From our exercise, we can appreciate different molecular geometries: tetrahedral in ammonium (H_{4}^{+}), bent in nitrite (O_{2}^{-}) and nitrogen dioxide (O_{2}), and linear in dinitrogen monoxide (_{2} O). These geometries are derived from the VSEPR theory, which provides a straightforward approach to predicting the shape based on the electron pairs' repulsions.

Understanding molecular geometry is vital as it impacts physical and chemical properties such as boiling and melting points, solubility, and the color of compounds.

Chemical Bonding

Chemical bonding is the force that holds atoms together in molecules. There are several types of chemical bonds, including ionic, covalent, and metallic bonds. In the context of Lewis structures and the molecules discussed from our exercise, covalent bonding is particularly relevant. It involves the sharing of electron pairs between atoms.

For H_{4}^{+}, the nitrogen and hydrogen atoms covalently bond by sharing electrons to satisfy the octet rule. Meanwhile, in the O_{2}^{-} and O_{2} molecules, we observe an odd number of valence electrons, resulting in a free radical—an unpaired electron that can make the molecule highly reactive. Additionally, the concept of formal charge can help to rationalize the most likely arrangement of atoms in molecules with multiple valid Lewis structures.

Chemical bonding not only defines the structure of the molecules but also predicts the reactivity, acidity or basicity, and overall stability of the compounds.