H2PO4- Lewis structure

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H₂PO₄^– (dihydrogen phosphate) has two hydrogen atoms, one phosphorus atom, and four oxygen atoms.

In the H₂PO₄^– Lewis structure, there is one double bond and three single bonds around the phosphorus atom, with four oxygen atoms attached to it. The oxygen atom with a double bond has two lone pairs, the left and right oxygen atom (with which the hydrogen atom is attached) also has two lone pairs, and the bottom oxygen atom with a single bond has three lone pairs.

Also, there is a negative (-1) charge on the bottom oxygen atom.

Steps

Here’s how you can easily draw the H₂PO₄^– Lewis structure step by step:

#1 Draw a rough skeleton structure
#2 Mention lone pairs on the atoms
#3 If needed, mention formal charges on the atoms
#4 Minimize formal charges by converting lone pairs of the atoms, and try to get a stable Lewis structure
#5 Repeat step 4 again if needed, until all charges are minimized

Now, let’s take a closer look at each step mentioned above.

#1 Draw a rough skeleton structure

• First, determine the total number of valence electrons

In the periodic table, hydrogen lies in group 1, phosphorus lies in group 15, and oxygen lies in group 16.

Hence, hydrogen has one valence electron, phosphorus has five valence electrons, and oxygen has six valence electrons.

Since H₂PO₄^– has two hydrogen atoms, one phosphorus atom, and four oxygen atoms, so…

Valence electrons of two hydrogen atoms = 1 × 2 = 2
Valence electrons of one phosphorus atom = 5 × 1 = 5
Valence electrons of four oxygen atoms = 6 × 4 = 24

Now the H₂PO₄^– has a negative (-1) charge, so we have to add one more electron.

So the total valence electrons = 2 + 5 + 24 +1 = 32

Learn how to find: Hydrogen valence electrons, Phosphorus valence electrons, and Oxygen valence electrons

• Second, find the total electron pairs

We have a total of 32 valence electrons. And when we divide this value by two, we get the value of total electron pairs.

Total electron pairs = total valence electrons ÷ 2

So the total electron pairs = 32 ÷ 2 = 16

• Third, determine the central atom

Here hydrogen can not be the central atom. Because the central atom is bonded with at least two other atoms, and hydrogen has only one electron in its last shell, so it can not make more than one bond.

Now we have to choose the central atom from phosphorus and oxygen. Place the least electronegative atom at the center.

Since phosphorus is less electronegative than oxygen, assume that the central atom is phosphorus.

Therefore, place phosphorus in the center and hydrogen and oxygen on either side.

• And finally, draw the rough sketch

#2 Mention lone pairs on the atoms

Here, we have a total of 16 electron pairs. And six bonds are already marked. So we have to only mark the remaining ten electron pairs as lone pairs on the sketch.

Also remember that hydrogen is a period 1 element, so it can not keep more than 2 electrons in its last shell. Phosphorus is a period 3 element, so it can keep more than 8 electrons in its last shell. And oxygen is a period 2 element, so it can not keep more than 8 electrons in its last shell.

Always start to mark the lone pairs from outside atoms. Here, the outside atoms are hydrogens and oxygens. But no need to mark on hydrogen, because each hydrogen has already two electrons.

So for top oxygen and bottom oxygen, there are three lone pairs, for left oxygen and right oxygen, there are two lone pairs, and for phosphorus, there is zero lone pair because all ten electron pairs are over.

Mark the lone pairs on the sketch as follows:

#3 If needed, mention formal charges on the atoms

Use the following formula to calculate the formal charges on atoms:

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

For each hydrogen atom, formal charge = 1 – 0 – ½ (2) = 0

For phosphorus atom, formal charge = 5 – 0 – ½ (8) = +1

For top oxygen and bottom oxygen atom, formal charge = 6 – 6 – ½ (2) = -1

For left oxygen and right oxygen atom, formal charge = 6 – 4 – ½ (4) = 0

Here, both phosphorus and oxygen atoms have charges, so mark them on the sketch as follows:

The above structure is not a stable Lewis structure because both phosphorus and oxygen atoms have charges. Therefore, reduce the charges (as below) by converting lone pairs to bonds.

#4 Minimize formal charges by converting lone pairs of the atoms

Convert a lone pair of the top oxygen atom to make a new P — O bond with the phosphorus atom as follows:

In the above structure, you can see that the central atom (phosphorus) forms an octet. The outside atoms (oxygens) also form an octet, and both hydrogens form a duet. Hence, the octet rule and duet rule are satisfied.

Now there is a negative (-1) charge on the bottom oxygen atom.

This is okay, because the structure with a negative charge on the most electronegative atom is the best Lewis structure. And in this case, the most electronegative element is oxygen.

Also, the above structure is more stable than the previous structures. Therefore, this structure is the most stable Lewis structure of H₂PO₄^–.

And since the H₂SO₄^– has a negative (-1) charge, mention that charge on the Lewis structure by drawing brackets as follows:

Next: HPO₄^2- Lewis structure

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