Question

Oxidation of the cyanide ion produces the stable cyanate ion, OCN \(^{-}\). The fulminate ion, \(\mathrm{CNO}^{-}\), on the other hand, is very unstable. Fulminate salts explode when struck; \(\mathrm{Hg}(\mathrm{CNO})_{2}\) is used in blasting caps. Write the Lewis structures and assign formal charges for the cyanate and fulminate ions. Why is the fulminate ion so unstable? (C is the central atom in OCN - and \(\mathrm{N}\) is the central atom in \(\mathrm{CNO}^{-} .\) )

Short answer

The cyanate ion (OCN-) has a stable Lewis structure with close to zero formal charges on each atom, and the more electronegative nitrogen atom holds the negative charge. In contrast, the fulminate ion (CNO-) has a less stable structure due to the nitrogen atom, which is less electronegative than carbon and oxygen, holding a negative charge. This unequal distribution of charge contributes to its instability and causes it to be highly reactive and explosive.

Step-by-step solution

Drawing Lewis structures

We start by drawing the Lewis structures for both ions. To do this, we will follow these steps: 1. Determine the total number of valence electrons for each ion. 2. Arrange the atoms in the ion, placing the central atom in the middle and surrounding atoms around it. For OCN-, C (carbon) atom is the central atom, and for CNO-, N (nitrogen) atom is the central atom. 3. Connect the atoms with single bonds and distribute the remaining electrons as lone pairs. For OCN- ion: 1. Oxygen has 6 valence electrons, Nitrogen has 5 and Carbon has 4. Adding them we get 6+5+4 = 15 valence electrons. However, since it's a negatively charged ion, one more electron is added, making the total number of valence electrons equal to 16. 2. Place Carbon in the middle with Oxygen on one side and Nitrogen on the other side. 3. Connect the atoms with single bonds and distribute the remaining electrons. For CNO- ion: 1. Carbon has 4 valence electrons, Nitrogen has 5 and Oxygen has 6. Adding them we get 4+5+6 = 15 valence electrons. However, since it's a negatively charged ion, one more electron should be added, so there are a total of 16 valence electrons. 2. Place Nitrogen in the middle with Carbon on one side and Oxygen on the other side 3. Connect the atoms with single bonds and distribute the remaining electrons.

Determine formal charges

To calculate the formal charges for each atom, we use the following formula: Formal charge = Valence electrons - (Number of bonding electrons/2) - Number of non-bonding electrons For OCN- ion: 1. Formal charge of Oxygen: 6 - (2/2) - 4 = 0 2. Formal charge of Carbon: 4 - (6/2) - 0 = 0 3. Formal charge of Nitrogen: 5 - (4/2) - 2 = -1 For CNO- ion: 1. Formal charge of Carbon: 4 - (2/2) - 2 = 0 2. Formal charge of Nitrogen: 5 - (6/2) - 0 = -1 3. Formal charge of Oxygen: 6 - (4/2) - 4 = 0

Analyzing the instability of the fulminate ion

The Lewis structure for the cyanate ion (OCN-) is more stable than the fulminate ion (CNO-) because the formal charges on each atom are close to zero, and the more electronegative element holds the negative charge (nitrogen). However, in the fulminate ion (CNO-), the nitrogen atom holds a negative charge and is placed between carbon and oxygen, which are both more electronegative than nitrogen. This unequal distribution of charge contributes to the instability of the fulminate ion, making it more reactive and explosive.

Key concepts

Formal Charge Calculation

In chemistry, the formal charge is a critical concept that helps describe the distribution of electrons among the atoms within a molecule or ion. To calculate the formal charge, we use the formula:
\[ \text{Formal charge} = \text{Valence electrons} - \left( \text{Number of bonding electrons}/2 \right) - \text{Number of non-bonding electrons} \]
When determining formal charges, it's essential to remember that a formal charge will essentially approximate how electrons are distributed in a molecule or ion. The resulting charge does not necessarily indicate the actual charge distribution within the molecule, as the model assumes equal sharing of electrons in bonds, which is an idealized concept. Nonetheless, molecules and ions with formal charges closer to zero are generally more stable.To ensure clarity in our calculations, let's walk through the calculation of formal charges for the cyanate ion (OCN-) as an example. With 16 valence electrons to consider, upon distribution, we find that oxygen and carbon each have a formal charge of zero, while nitrogen carries a negative formal charge. This calculation helps explain the stability of the cyanate ion, as we look for minimal formal charges on every atom.

Valence Electrons

Valence electrons play a pivotal role in chemical bonding and reactions. They are the electrons in the outermost shell of an atom that are involved in forming bonds. For example, oxygen has 6 valence electrons, nitrogen has 5, and carbon typically has 4.
When chemists construct Lewis structures, they first count the valence electrons to determine how atoms can bond and whether any extra electrons will exist as lone pairs. Calculating the total number of valence electrons for a given molecule or ion sets the stage for placing electrons between atoms as bonds and determines how many electrons will remain to be distributed as lone pairs. In our step-by-step process for the cyanate and fulminate ions, the total count of valence electrons includes the contribution of the negative charge, thus adding one more to the usual count.

Cyanate and Fulminate Ions

The cyanate (OCN-) and fulminate (CNO-) ions showcase the diversity of molecular structures and stability. Cyanate is relatively stable, whereas fulminate is notorious for being highly unstable and explosive.
In the case of the cyanate ion, we see that the more electronegative elements, oxygen and nitrogen, surround carbon. Stability is inferred from the zero formal charges on oxygen and carbon, with the negative charge localized on the nitrogen atom. However, the fulminate ion has a different structure with nitrogen as the central atom, flanked by carbon and oxygen. The negative formal charge on nitrogen, which is less electronegative than both carbon and oxygen, creates an unstable distribution of electrons. This internal tension contributes to the reactive nature of the fulminate ion, making it sensitive to shock and heat, with the ability to explode under such conditions. Understanding this instability is pivotal in safety applications and the handling of fulminate salts like mercury fulminate (Hg(CNO)2), a compound used in blasting caps.