The **Lewis structure of NO _{4}^{3-}** contains one double bond and three single bonds, with nitrogen in the center, and four oxygens on either side. The left oxygen atom has two lone pairs. The right oxygen atom, top oxygen atom, and bottom oxygen atom have three lone pairs, and the nitrogen atom does not have any lone pair.

Plus, there is a negative (-1) charge on the right oxygen atom, top oxygen atom, and bottom oxygen atom.

## Steps

By using the following steps, you can easily draw the Lewis structure of NO_{4}^{3-}.

#1 Draw skeleton

#2 Show chemical bond

#3 Mark lone pairs

#4 Calculate formal charge and check stability (if octet is already completed on central atom)

#5 Convert lone pair and calculate formal charge again (if formal charges are not closer to zero)

Let’s one by one discuss each step in detail.

### #1 Draw skeleton

In this step, first calculate the total number of valence electrons. And then, decide the central atom.

- Let’s calculate the total number of valence electrons

We know that… nitrogen is a group 15 element and oxygen is a group 16 element. Hence, nitrogen has **five** valence electrons and oxygen has **six** valence electrons.

Now NO_{4}^{3-} has one nitrogen atom and four oxygen atoms.

So the total number of valence electrons = valence electrons of nitrogen atom + (valence electrons of oxygen atom × 4)

And NO_{4}^{3-} has a negative (-3) charge, so we have to add three more electrons.

Therefore, the **total number of valence electrons** = 5 + 24 + 3 = 32

- Now decide the central atom

The atom with the least electronegative value is placed at the center. By looking at the periodic table, we get the electronegativity values for nitrogen and oxygen as follows:

Electronegativity value of nitrogen = 3.04

Electronegativity value of oxygen = 3.44

Obviously, nitrogen is less electronegative than oxygen. Hence, assume that **nitrogen is the central atom**.

So now, put nitrogen in the center and oxygens on either side. And draw the rough skeleton structure for the Lewis structure of NO_{4}^{3-} something like this:

**Also read:** How to draw Lewis structure of SF_{3}^{+} (4 steps)

### #2 Show chemical bond

Place two electrons between the atoms to show a chemical bond. Since nitrogen is surrounded by four oxygens, use eight electrons to show **four chemical bonds** as follows:

### #3 Mark lone pairs

As calculated earlier, we have a total of 32 valence electrons. And in the above structure, we have already used eight valence electrons. Hence, twenty-four valence electrons are remaining.

Two valence electrons represent one lone pair. So twenty-four valence electrons = **twelve lone pairs**.

Note that both (nitrogen and oxygen) are the period 2 elements, so they can not keep more than 8 electrons in their last shell.

Also, make sure that you start marking these lone pairs on outside atoms first. And then, on the central atom.

The outside atoms are oxygens, so each oxygen will get three lone pairs. And the central atom (nitrogen) will not get any lone pair, because all twelve lone pairs are used.

So the Lewis structure of NO_{4}^{3-} looks something like this:

In the above structure, you can see that the octet is completed on the central atom (nitrogen), and also on the outside atoms. Therefore, the octet rule is satisfied.

Now calculate the formal charge and check the stability of the above structure.

**Also read:** How to draw Lewis structure of C_{2}Cl_{2} (5 steps)

### #4 Calculate formal charge and check stability

The following formula is used to calculate the formal charges on atoms:

Formal charge = valence electrons – nonbonding electrons – ½ bonding electrons

Collect the data from the above structure and then, write it down below as follows:

- For
**nitrogen**atom

Valence electrons = 5

Nonbonding electrons = 0

Bonding electrons = 8

Formal charge = 5 – 0 – ½ (8) = +1

- For
**each oxygen**atom

Valence electrons = 6

Nonbonding electrons = 6

Bonding electrons = 2

Formal charge = 6 – 6 – ½ (2) = -1

Mention the formal charges of atoms on the structure. So the Lewis structure of NO_{4}^{3-} looks something like this:

In the above structure, you can see that the formal charges of atoms are not closer to zero.

Now as mentioned earlier… nitrogen is a period 2 element, so it can not keep more than 8 electrons in its last shell. Hence, the above structure can be the final Lewis structure of NO_{4}^{3-}.

But NO_{4}^{3-} is an ion having a negative (-3) charge, so the overall formal charge in the above Lewis structure should be (-3).

Therefore, let’s convert lone pair and calculate formal charge again.

**Also read:** How to draw Lewis structure of CH_{3}CH_{2}NH_{2} (4 steps)

### #5 Convert lone pair and calculate formal charge again

Note that we have assumed an exception that… in NO_{4}^{3-} nitrogen can keep more than 8 electrons in its last shell.

So convert one lone pair from one oxygen atom to make a new bond with the nitrogen atom. And then, the Lewis structure of NO_{4}^{3-} looks something like this:

Now one last thing we need to do is, calculate the formal charge again and check the stability of the above structure.

Formal charge = valence electrons – nonbonding electrons – ½ bonding electrons

Collect the data from the above structure and then, write it down below as follows:

- For
**nitrogen**atom

Valence electrons = 5

Nonbonding electrons = 0

Bonding electrons = 10

Formal charge = 5 – 0 – ½ (10) = 0

- For
**left oxygen**atom

Valence electrons = 6

Nonbonding electrons = 4

Bonding electrons = 4

Formal charge = 6 – 4 – ½ (4) = 0

- For
**right oxygen**,**top oxygen**, and**bottom oxygen**atom

Valence electrons = 6

Nonbonding electrons = 6

Bonding electrons = 2

Formal charge = 6 – 6 – ½ (2) = -1

Mention the formal charges of atoms on the structure. So the Lewis structure of NO_{4}^{3-} looks something like this:

In the above structure, you can see that the formal charges of atoms are closer to zero. Therefore, this is the **most stable Lewis structure of NO _{4}^{3-}**.

And each horizontal line drawn in the above structure represents a pair of bonding valence electrons.

Now NO_{4}^{3-} is an ion having a negative (-3) charge, so draw brackets around the above Lewis structure and mention that charge on the top right corner. And then, the Lewis structure of NO_{4}^{3-} looks something like this:

## Related

- SF
_{3}^{+}Lewis structure - Lewis structure of C
_{2}Cl_{2} - Lewis structure of CH
_{3}CH_{2}NH_{2} - Lewis structure of NF
_{2}^{–} - Lewis structure of CH
_{3}CH_{2}Cl

## External links

- NO43- Lewis Structure in 5 Steps (With Images) – Pediabay
- What’s the Lewis structure of NO4 -3? – Quora
- NO43- lewis structure, molecular geometry, bond angle, hybridization – Topblogtenz
- Draw the Lewis structure for NO43- – Homework.Study.com
- Draw a Lewis structure for NO4-3 then label the formal charges on each atom and finally propose a structure to minimize the formal charge – Chegg
- Whats the Lewis structure of NO4 -3? – ECHEMI
- Draw a Lewis structure for NO4-3 then label the formal charges on each atom and finally propose a structure to minimize the formal charge – Brainly
- Draw the Lewis Structure for NO4 3- (as well as any resonance structures) and offer an explanation why the NO4 3- anion is unstable – Numerade

Deep

Rootmemory.com was founded by Deep Rana, who is a mechanical engineer by profession and a blogger by passion. He has a good conceptual knowledge on different educational topics and he provides the same on this website. He loves to learn something new everyday and believes that the best utilization of free time is developing a new skill.