H2PO4- Lewis structure

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.

Alternative method: Lewis structure of H₂PO₄^–

Rough sketch

• 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

Lone pair

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:

Formal charge

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.

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

Final structure

The final structure of H₂PO₄^– comprises a central phosphorus atom connected to four oxygen atoms, two of which are bonded to individual hydrogen atoms. Within this layout, the phosphorus atom utilizes an expanded valence shell to form one double bond with a lone oxygen atom and single bonds with the three remaining oxygen atoms. One of these single-bonded oxygen atoms carries a formal charge of -1, while the other atoms, including the phosphorus and the two hydroxyl (-OH) groups, maintain a formal charge of zero. The double-bonded oxygen satisfies its octet with two lone pairs, the hydroxyl oxygens each retain two lone pairs, and the negatively charged single-bonded oxygen holds three lone pairs. Consequently, this specific electronic distribution serves as the definitive and most accurate Lewis representation of the dihydrogen phosphate ion.

To complete the representation, draw square brackets around the entire Lewis structure and place a “-” or “-1” sign as a superscript outside the upper right bracket. This notation signifies that the negative charge is a property of the whole ion.

Next: HPO₄^2- Lewis structure

External video

• How to Draw the Lewis Dot Structure for H2PO4-: Dihydrogen phosphate ion – YouTube • Wayne Breslyn