How to calculate formal charges of NO3- with lewis structure?
In covalently bonded molecules, formal charge is the charge assigned to an atom based on the assumption that the bonded electrons are equally shared between concerning atoms, regardless of their electronegativity.
The overall formal charge present on a molecule is a measure of its stability.
The fewer the formal charges present on the bonded atoms in a molecule (close to zero), the greater the stability of its Lewis structure.
In this article, we will calculate the formal charges present on the nitrate [NO3]– ion using its best possible Lewis representation. In addition, we will also discuss the formal charges present in the different resonance structures of NO3–.
So, what are you waiting for? Let’s start reading!
|Name of the molecule||Nitrate|
|The formal charge on N-atom||+1|
|The formal charge on each single bonded O-atom||-1|
|The formal charge on the double-bonded O-atom||0|
|Overall formal charge on [NO3]–||-1|
How to calculate the formal charges on NO3– atoms?
The formal charges can be calculated using the formula given below:
The formal charge of an atom = [valence electrons of an atom – non-bonding electrons – ½ (bonding electrons)]
- The valence electrons (V.E) of an atom are the total number of electrons present in its valence shell. Valence electrons can be calculated by locating the position of the elemental atom in the Periodic Table.
- Non-bonding electrons (N.E) are the number of lone pairs present on the atom. (1 lone pair means 2 nonbonding electrons).
- Bonding electrons (B.E) are the total electrons shared with the atom via covalent chemical bonds. (1 single bond means 2 bonding electrons).
The most preferred Lewis representation of NO3– is as shown below.
In the above Lewis structure, there are a total of 24 valence electrons. The central nitrogen (N) atom is single-bonded to two oxygen (O) atoms via N-O bonds and double-bonded to the third oxygen (O) atom via the N=O bond.
Now let’s see how we can use this Lewis structure and apply the formal charge formula given above to determine NO3– formal charges.
For nitrogen atom
- Valence electrons of nitrogen = It is present in Group V A = 5 valence electrons
- Bonding electrons =1 double bond + 2 single bonds = 4 + 2(2) = 8 electrons
- Non-bonding electrons = no lone pair = 0 electrons
- Formal charge = 5 – 0 – 8/2 =5 – 0 – 4 = 5 – 4 = +1
∴ The formal charge on the N-atom in NO3– is +1.
For each single-bonded oxygen atom
- Valence electrons of oxygen = It is present in Group VI A = 6 valence electrons
- Bonding electrons = 1 single bond = 2 electrons
- Non-bonding electrons = 3 lone pairs = 3(2) = 6 electrons
- Formal charge = 6 – 6 – 2/2 = 6 – 6 – 1 = 6 – 7 = -1
∴ The formal charge on each single-bonded O-atom in NO3– is -1.
For double-bonded oxygen atom
- Valence electrons of oxygen = It is present in Group VI A = 6 valence electrons
- Bonding electrons = 1 double bond = 4 electrons
- Non-bonding electrons = 2 lone pairs = 2(2) = 4 electrons
- Formal charge = 6 – 4 – 4/2 = 6 – 4 – 2 = 6 – 6 = 0
∴ The formal charge on the double-bonded O-atom in NO3– is 0.
The above calculation shows that zero formal charges are present on the double-bonded O-atom in NO3– Lewis structure. However, a +1 formal charge is present on the central N-atom, while each single-bonded O-atom carries a -1 formal charge. 1+ (-1) + (-1) = -1, which is the charge present on the nitrate ion (NO3–) overall in this most preferred Lewis representation.
Therefore, the NO3– Lewis structure is enclosed in square brackets, and a -1 formal charge is placed at the top right corner, as shown below.
This is the most stable Lewis structure of NO3– because the formal charges are as minimized in it as possible.
However, two other resonance structures are possible for representing the nitrate ion, as shown below.
Each resonance structure is a way of representing the Lewis structure of a molecule or molecular ion. The pi-bonded electrons in the N=O double bond, as well as the lone pairs, present on each single-bonded O-atom in NO3–, keep revolving from one position to another. The formal charges present on bonded atoms simultaneously change positions.
In each case, however, the central N-atom carries a +1 formal charge, while each single-bonded O-atom carries a -1 formal charge. +1 cancels with -1 of one O-atom to yield an overall negative one formal charge in each resonance form.
Therefore, each of the above resonance structures is equivalent and thus equally stable. As a result, the actual Lewis structure of NO3– is a hybrid of the above resonance forms.
Also, check –
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How can you calculate NO3– formal charges?
The formal charges present on the bonded atoms in NO3– can be calculated using the formula given below:
V.E – N.E – B.E/2
⇒ V.E = valence electrons of an atom
⇒ N.E = non-bonding electrons, i.e., lone pairs
⇒ B.E = bonding electrons
What is the formal charge present on the central N-atom in NO3–?
|The central N-atom carries a +1 formal charge in NO3– Lewis structure.|
What is the formal charge present on the O-atoms in NO3–?
|The two single-bonded O-atoms carry -1 formal charges, while the one double-bonded O-atom carries zero or no formal charges in NO3– Lewis structure.|
Do all three O-atoms carry the same formal charges in NO3–?
|No, the single bonded O-atoms carry -1 formal charge each while the double bonded O-atom carries zero formal charges.|
What is the overall charge present on NO3–?
|+1 formal charge on the central N-atom cancels with -1 formal charge of one of the two single-bonded O-atoms. In this way, the overall formal charge on NO3– is -1.|
How many resonance structures are possible for representing NO3–?
Three distinct resonance structures are possible for representing NO3–, as shown below.
Do all three resonance structures of NO3– carry the same formal charges?
Yes. In each NO3– resonance structure, the central N-atom carries a +1 formal charge. The single-bonded O-atoms carry -1 formal charge each, while the double-bonded O-atom carries zero formal charges.
The overall charge present on each structure is thus -1. Therefore, all the resonance structures are equivalent. The actual NO3– structure is a hybrid of the three resonance structures with an equal contribution of each resonance form.
- The best possible Lewis structure of a molecular ion is the one in which the bonded atoms carry formal charges as close to zero as possible.
- The formal charge formula is [ V.E – N.E – B.E/2].
- The nitrate [NO3]– ion consists of three distinct resonance structures.
- In each resonance form, the central N-atom carries a +1 formal charge, while each of the two single-bonded O-atoms carries a -1 formal charge. However, no formal charge is present on the double-bonded O-atom in NO3–.
- +1 formal charge of the central N-atom cancels with -1 charge of one of the two single-bonded O-atoms.
- The overall charge present on NO3– in each of the three resonance forms is -1. Thus, it is a monovalent anion.
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