# How to calculate formal charges of Phosphate (PO43-) ion with lewis structure?

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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.

Phosphate is one of the necessary ingredients in fertilizer production and is a major part of DNA as well.

In this article, we will calculate the formal charges present on the bonded atoms in [PO4]3- and also the overall charge present on the molecular ion.

So, without any further delay, let’s start reading! Happy learning.

 Name of the molecular ion Phosphate Chemical formula [PO4]3- The formal charge on the P-atom 0 The formal charges on double-bonded O-atoms 0 The formal charge on each single-bonded O-atoms -1 The overall formal charge on [PO4]3- -3

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## How to calculate the formal charges on [PO4]3- 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 determined 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).

Now let us use this formula to calculate the formal charges in the most preferred Lewis structure of the phosphate [PO4]3- ion.

The most preferred Lewis representation of [PO4]3- is as shown below.

It consists of a total of 32 valence electrons. A Phosphorus (P) atom is present at the center, which is bonded to four Oxygen atoms, one atom of oxygen (O) via a double covalent bond, and three oxygen atoms via single covalent bonds.

No lone pair of electrons is present at the central P-atom. Double-bonded O-atom contains 2 lone pairs of electrons, while each of the three single-bonded O-atoms contains 3 lone pairs of electrons.

It is the best possible Lewis structure of Phosphate ion [PO4]3- because the formal charges are minimized in it, and thus, it is the most stable.

Let’s find out how we can determine the formal charges present on each atom in [PO4]3- Lewis structure.

For the central Phosphorus atom

• Valence electrons of Phosphorus = It is present in Group V-A = 5 valence electrons
• Bonding electrons around Phosphorus = 3 single bonds + 1 double bond = 3(2) + 1(4) = 10 electrons
• Non-bonding electrons on Phosphorus = no lone pair = 0 electrons
• Formal charge on the Phosphorus atom = 5 – 0 – 10/2 = 5 – 0 – 5 = 5 – 5 = 0

The formal charge on the Phosphorus (P) atom in [PO4]3- is 0.

For single-bonded oxygen atoms

• Valence electrons of Oxygen = It is present in Group VI-A = 6 valence electrons
• Bonding electrons around Oxygen = 1 single bond = 2 electrons
• Non-bonding electrons on Oxygen = 3 lone pairs = 3(2) = 6 electrons
• Formal charge on the single bonded Oxygen atom = 6 – 6 – 2/2 = 6 – 6 – 1 = 6 – 7 = -1

The formal charge on each single bonded oxygen (O) atom in [PO4]3- is -1.

For double-bonded oxygen atom

• Valence electrons of oxygen = It is present in Group VI A = 6 valence electrons
• Bonding electrons around Oxygen = 1 double bond = 4 electrons
• Non-bonding electrons on Oxygen = 2 lone pairs = 2(2) = 4 electrons
• Formal charge on the double bonded Oxygen atom = 6 – 4 – 4/2 = 6 – 4 – 2 = 6 – 6 = 0

The formal charge on the double-bonded oxygen (O) atom in [PO4]3- is 0.

This calculation shows zero formal charges are present on the central P atom and on the double-bonded O-atom, while each of the three single-bonded O-atoms carries -1 formal charge, that when added (-1 + (-1) + (-1) = -3) makes up -3 that accounts for the formal charge present on the phosphate ion [PO4]3- overall, as shown below.

As a final step, the PO43- Lewis structure is enclosed in square brackets, and a -3 charge is placed at the top right corner.

Also, check –

## FAQ

### How can you calculate [PO4]3-  formal charges?

The formal charges present on the bonded atoms in [PO4]3- can be calculated using the formula given below:

V.E – N.E – B.E/2

Where –

⇒ 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 on P-atom in [PO4]3-?

The central Phosphorus (P) atom carries zero formal charges in [PO4]3-.

### What is the formal charge on O-atoms in [PO4]3-?

Zero or no formal charge is present on double (P=O) bonded oxygen atoms in [PO4]3- while -1 formal charge is present on each of the three single (P-O) bonded oxygen atoms in [PO4]3-.

### What is the overall formal charge on [PO4]3-?

The overall formal charge on [PO4]3- is –3.

### Do all three O-atoms carry the same formal charge in [PO4]3-?

No, the single bonded O-atoms carry -1 formal charge each, while no formal charge is present on the double bonded O-atom in [PO4]3- Lewis structure.

## Summary

• The best possible Lewis structure of a molecule 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].
• In [PO4]3-, zero formal charges is present on the central P-atom.
• The double-bonded O-atom also has zero formal charges in [PO4]3-.
• The single-bonded O-atoms carry a -1 formal charge in [PO4]3-.
• The overall formal charge on [PO4]3- is -3.
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