Question

Use covalent Lewis structures to explain why the compound that forms between nitrogen and hydrogen has the formula \(\mathrm{NH}_{3}\). Show why \(\mathrm{NH}_{2}\) and \(\mathrm{NH}_{4}\) are not stable.

Short answer

NH3 is stable because nitrogen has 5 valence electrons and forms three covalent bonds with hydrogen atoms, reaching an octet. NH2 lacks an octet and NH4 goes beyond the octet, both leading to instability.

Step-by-step solution

Recognizing the Valence Electrons

Nitrogen has 5 valence electrons and hydrogen has 1 valence electron. For a stable molecule, nitrogen wants to reach an octet (8 electrons in its outer shell) and hydrogen aims for a duet (2 electrons in its outer shell).

Drawing the Lewis Structure for NH3

Pair each of the 3 hydrogen atoms with 1 of nitrogen's valence electrons to form 3 covalent bonds. The remaining 2 electrons on nitrogen form a lone pair. This satisfies the octet rule for nitrogen (8 total electrons around it) and the duet rule for each hydrogen (2 electrons around each).

Assessing the Stability of NH2

In NH2, nitrogen would have 6 electrons in its valence shell (5 from itself, 1 from each hydrogen), leaving it short of an octet. This would make it reactive and inclined to seek two more electrons.

Assessing the Stability of NH4

In NH4, nitrogen would have 10 electrons surrounding it (5 from itself and 1 from each of the 4 hydrogens), which violates the octet rule for nitrogen. This would result in a structure with too much energy, making it unstable.

Key concepts

Valence Electrons

Understanding valence electrons is critical when it comes to predicting the behavior of atoms during chemical bonding. Valence electrons are the electrons located in the outermost energy level of an atom. They are the 'social butterflies' of atoms, as they participate in bonding by being lost, gained, or shared with neighboring atoms.

For example, in the case of nitrogen (\textbf{N}), it has five valence electrons. Hydrogen (\textbf{H}), on the other hand, has one valence electron. By knowing the number of valence electrons, one can predict the bonding patterns and the resultant molecular structures, such as why ammonia (\textbf{NH}\(_3\)) is stable, while other hypothetical structures like \textbf{NH}\(_2\) and \textbf{NH}\(_4\) are not.

Octet Rule

Moving on to the octet rule, this rule is one of the cornerstones of chemical bonding, and it explains why atoms bond the way they do. It states that atoms are most stable when they have eight electrons in their valence shell, mimicking the electron configuration of noble gases. The octet rule is akin to the social rule of 'better together,' as atoms often share or exchange electrons to fill up their valence shells and reach a stable configuration.

In our exercise, nitrogen achieves stability by forming covalent bonds with three hydrogen atoms, resulting in a total of eight electrons in its valence shell, thus adhering to the octet rule. On the contrary, \textbf{NH}\(_2\) and \textbf{NH}\(_4\) do not adhere to this rule, making them less stable or hypothetical.

Covalent Bonding

Finally, let's delve into covalent bonding. This type of bond involves the sharing of valence electrons between atoms. Picture this as a friendship where two parties share something valuable equally—their electrons. Covalent bonds are strong and directional, meaning they hold molecules together in a specific shape.

In ammonia (\textbf{NH}\(_3\)), each hydrogen atom shares its single electron with nitrogen, forming a strong covalent bond. As a result, the molecule's geometry is such that it maximizes the distance between the pairs of shared electrons, leading to a trigonal pyramidal shape. These shared pairs of electrons allow nitrogen to comply with the octet rule, while each hydrogen follows the 'duet rule,' settling for two electrons in their valence shell, which is the equivalent stability principle for hydrogen.