Question

The formal charges on the \(\mathrm{O}\) atoms in the ion \([\mathrm{ONO}]^{+}\) is \((\mathrm{a})-2 ;(\mathrm{b})-1 ;(\mathrm{c}) 0 ;(\mathrm{d})+1\).

Short answer

The formal charge on each Oxygen atom in the [ONO]+ ion is 0.

Step-by-step solution

Determine the number of Valence Electrons

Start by determining the number of valence electrons for the Oxygen atom. Oxygen, being in Group 6 or 16 of the periodic table, has 6 valence electrons.

Build Lewis Structure

Try to build Lewis Structure for the [ONO]+ ion. \(\mathrm{O}\) atom at each end is connected to \(\mathrm{N}\) atom in the middle by a double bond, leaving 2 lone pairs on each Oxygen. The Nitrogen atom is bonded once to a single free electron which results in the ion's positive charge. Therefore, the structure for [ONO]+ ion is:: O=NO+

Calculate the Formal Charge

The formula for calculating formal charge is: Formal charge = Valence electrons - (Number of lone pair electrons + 1/2 * number of bonding electrons). For the Oxygen atoms in [ONO]+, both with a lone pair and double bonds respectively, this amounts to: Formal charge of O = 6 valence electrons - (4 lone pair electrons + 1/2 * 4 bonding electrons) = 6 - (4 + 2) = 0. Therefore the formal charge on the Oxygen atoms is (c) 0.

Key concepts

Formal Charges

Formal charges help us determine the most stable Lewis structure for a molecule. Formal charge is a bookkeeping tool that shows what the charge of atoms would be if every bond was shared equally.
To calculate the formal charge, use this formula:

• Formal charge = Valence electrons - (Number of lone pair electrons + 1/2 * number of bonding electrons).

For instance, let's calculate it for the oxygen atoms in the ion \([\mathrm{ONO}]^{+}\). Each oxygen atom bonds with nitrogen through a double bond and has two lone pairs. So:

• Valence electrons (for Oxygen) = 6 (Group 16 in the periodic table)
• Lone pair electrons = 4 (2 pairs of electrons)
• Bonding electrons = 4 (double bond, shared equally)

Plug into the formula: 6 - (4 + 1/2 * 4) = 0.
This shows us both oxygen atoms have a formal charge of 0, contributing to the neutrality of the structure.

Valence Electrons

Valence electrons are crucial in understanding how atoms bond and form molecules. These are the electrons in the outermost shell of an atom. They are involved in forming bonds with other atoms and determine the chemical properties of the element.
Let's look at oxygen, which plays a central role in molecules like \([\mathrm{ONO}]^{+}\). Oxygen is in Group 16 or Group 6 of the periodic table. This tells us it has 6 valence electrons. During bonding, oxygen often tries to fill its outer shell to a total of 8 electrons, achieving a stable configuration.
Recognizing the number of valence electrons:

• Can help predict how atoms bond together.
• Provides insight into the geometry of molecules and ions.
• Is pivotal in understanding the formal charge calculations.

Remember, molecules form the most stable structures when their atoms achieve a filled valence electron shell.

Oxygen and Nitrogen Bonding

Oxygen and nitrogen are key components of many organic and inorganic compounds. In the context of \([\mathrm{ONO}]^{+}\), oxygen, and nitrogen form a specific structure to stabilize the molecule. The molecule features nitrogen sandwiched between two oxygen atoms.
Both oxygen-nitrogen bonds are double bonds. Double bonds involve sharing two pairs of electrons between atoms. This allows the molecules to become more stable and lower their potential energy. For \([\mathrm{ONO}]^{+}\):

• Nitrogen achieves a filled valence shell by sharing electrons with adjacent oxygen.
• The formation of double bonds helps satisfy the octet rule for oxygen and nitrogen.
• Each oxygen maintains lone pairs, which are not involved in bonding, affecting the electron distribution.

Additionally, the positive charge on the ion is mainly due to nitrogen, as it donates an electron to balance the structure. Recognizing how oxygen and nitrogen bond in this way can help understand more complex ionic and covalent molecules.