Phosphate [PO4]3- ion Lewis dot structure, molecular geometry or shape, electron geometry, bond angle, hybridization, formal charges, polar vs non-polar
The phosphate [PO4]3- ion is the chemical derivative of phosphoric acid (H3PO4). The phosphate ion is an important constituent for the development of biological structural materials such as bones and teeth. It is also important for plant growth so phosphate salts are widely used in fertilizer manufacturing.
In this article, we have compiled for you some interesting information about the chemistry of a phosphate [PO4]3- ion such as how to draw its Lewis dot structure, what is its molecular geometry or shape, electron geometry, bond angle, hybridization, formal charge, polarity, etc.
So without any further delay, let’s start reading!
The Lewis structure of a phosphate [PO4]3- ion consists of one phosphorus (P) atom and four atoms of oxygen (O). The phosphorus atom is present at the center while the oxygen atoms occupy terminal positions.
There are a total of 4 electron density regions around the central phosphorus atom in [PO4]3-. All the 4 electron density regions are constituted of bond pairs thus there is no lone pair of electrons on the central P-atom in [PO4]3- Lewis structure.
You can easily draw the Lewis dot structure of a phosphate ion with us using the simple guidelines given below.
Steps for drawing the Lewis dot structure of [PO4]3-
1. Count the total valence electrons in [PO4]3-
The Lewis dot structure of a molecule is referred to as a simplified representation of all the valence electrons present in it. Therefore, the very first step while drawing the Lewis structure of [PO4]3- is to count the total valence electrons present in the concerned elemental atoms.
There are two different elemental atoms present in the phosphate [PO4]3- ion i.e., phosphorus (P) and oxygen (O). We can easily determine the valence electrons present in phosphorus and oxygen by identifying these elements in the Periodic Table.
Phosphorus (P) is present in Group V A of the Periodic Table of elements while Oxygen (O) is located in Group VI A. According to this, a total of 5 valence electrons are present in each atom of phosphorus while 6 valence electrons are present in each oxygen atom.
The [PO4]3- ion consists of 1 P-atom and 4 O-atoms. Thus, the valence electrons in the Lewis dot structure of [PO4]3- = 1(5) + 4(6) = 29 valence electrons.
However, the twist here is that the [PO4]3- ion carries a negative (-3) charge which means 3 extra valence electrons are added in this Lewis structure.
∴ Hence, the total valence electrons available for drawing the Lewis dot structure of [PO4]3- = 29+3 = 32 valence electrons.
2. Choose the central atom
In the second step of drawing the Lewis structure of a molecule or a molecular ion, we need to place the least electronegative atom at the center.
As electronegativity refers to the ability of an elemental atom to attract a shared pair of electrons from a covalent chemical bond therefore the least electronegative atom is the one that is most likely to share its electrons with other atoms.
As phosphorus is less electronegative than oxygen so in the Lewis structure of [PO4]3-, we will place the P atom at the center while all the four O-atoms will be spread around it, as shown below.
3. Connect outer atoms with the central atom
Now we need to connect the outer atoms with the central atom of a Lewis structure using single straight lines. As the four oxygen atoms are the outer atoms in [PO4]3- Lewis structure while the phosphorus atom is the central atom, so we will connect all the 4 O-atoms with the central P-atom using straight lines.
Each straight line represents a single covalent bond i.e., a bond pair containing 2 electrons. There are a total of 4 single bonds in the above diagram which means a total of 4(2) = 8 valence electrons are used till this step, out of the 32 initially available.
Total valence electrons available – electrons used till step 3 = 32-8 = 24 valence electrons.
This means we still have 24 valence electrons to be accommodated in the Lewis dot structure of [PO4]3-.
4. Complete the octet of outer atoms
There are four O-atoms present as outer atoms in the Lewis structure of [PO4]3-. Each O-atom needs a total of 8 valence electrons in order to achieve a stable octet electronic configuration.
Each P-O bond already represents 2 electrons therefore all the four O-atoms require 6 more electrons each to complete their octet. Thus, these 6 valence electrons are placed as 3 lone pairs on each O-atom, as shown below.
5. Complete the octet of the central atom
Total valence electrons used till step 4 = 4 single bonds + 4 (electrons placed around an O-atom, shown as dots) = 4(2) + 4(6) = 32 valence electrons.
Total valence electrons available – electrons used till step 4 = 32-32 = 0 valence electrons.
As all the 32 valence electrons are already used up in the Lewis structure drawn till this step so there is no lone pair on the central P atom in the Lewis dot structure of [PO4]3-.
Also, there are a total of 4 single bonds around the central phosphorus atom which means a total of 8 valence electrons are present so the octet of the central P atom is also complete in addition to a complete octet electronic configuration of the terminal O atoms.
But, the thing is, is this Lewis structure stable to be marked as the correct Lewis structure of [PO4]3-? Let us check that using the formal charge concept.
6. Check the stability of the PO43- Lewis structure using the formal charge concept
The less the formal charge on the atoms of a molecule, the better the stability of its Lewis structure.
The formal charge can be calculated using the formula given below.
This calculation shows that a +1 formal charge is present on the central phosphorus atom and a -1 formal charge is present on each of the four oxygen (O) atoms.
But as we already mentioned; the less the formal charge on the bonded atoms of a molecule or molecular ion, the greater the stability of its Lewis structure.
The interesting fact is that phosphorus (P) has the capacity to accommodate more than 8 valence electrons during chemical bonding. Due to the availability of a 3d subshell in its atomic structure, the P-atom follows the expanded octet rule.
Keeping this property of Phosphorus in mind, we can reduce the formal charge present on it by converting a lone pair from any one terminal O-atom into a covalent bond between the central P and that terminal O.
Let us see how that’s done.
7. Minimize the formal charges on atoms by converting lone pairs into covalent bonds
One lone pair from any one terminal oxygen atom is converted into a covalent bond between the central P-atom and the respective O-atom as shown below.
Now there are a total of 3 single bonds and 1 double bond around the central P-atom. We can again check the stability of this Lewis structure using the formal charge formula.
For phosphorus atom
Valence electrons of phosphorus = 5
Bonding electrons = 1 double bond + 3 single bonds = 1(4) + 3(2) = 10 electrons
Non-bonding electrons = no lone pair = 0 electrons
So here, you can see that the formal charges present on the central phosphorus (P) atom and a terminal oxygen (O) atom are reduced to zero. However, there is a -1 formal charge on each of the other three oxygen atoms.
-1 + (-1) + (-1) = -3 which accounts for an overall negative 3 charge on the phosphate [PO4]3- ion. This ensures that it is a correct and stable Lewis representation for the phosphate [PO4]3- ion. The [PO4]3- Lewis structure is enclosed in square brackets and a negative 3 charge is placed at the top right corner, as shown below.
In this way, the phosphorus atom has a total of 10 valence electrons around it in the Lewis structure of phosphate [PO4]3- ion while each terminal O atom has a complete octet.
Another interesting fact to keep in mind is that the actual structure of a phosphate [PO4]3- ion is a hybrid of the following resonance structures. Each resonance structure is a way of representing the Lewis structure of a molecule or an ion.
These resonance structures show that the formal charges present on [PO4]3- atoms are not stationary, rather they keep moving from one position to another. Similarly, a double bond can be formed between the central phosphorus atom and any one terminal oxygen atom out of all the four available.
In conclusion, all the above resonance structures contribute equally to the resonance hybrid which is the best possible Lewis structure of the phosphate [PO4]3- ion.
Now, that we have discussed everything about the Lewis structure of [PO4]3-, we are good to proceed forward to the next section of the article.
What are the electron and molecular geometry of PO43-?
The phosphate [PO4]3- ion has an identical electron and molecular geometry or shape i.e., tetrahedral. The four O-atoms bonded to the central P-atom lie in the same plane, in a symmetrical arrangement, along the four vertices of a tetrahedron.
No lone pair of electrons is present on the central P-atom in [PO4]3- thus there is no distortion present in the geometry or shape of the molecular ion.
Molecular geometry of [PO4]3-
The phosphate [PO4]3- ion has a tetrahedral molecular geometry or shape. The four oxygen atoms lie at the four vertices of a regular tetrahedron while the phosphorus atom is present at the center, refer to the figure below.
P-O and P=O bond pair-bond pair repulsions exist in the molecule which keeps the terminal O-atoms as far apart from one another as possible.
However, there is no lone pair of electrons on the central P-atom therefore no lone pair-bond pair and lone pair-lone pair electronic repulsions are present in the molecule. The shape of the molecule thus stays intact, identical to its ideal electron pair geometry.
Electron geometry of [PO4]3-
According to the valence shell electron pair repulsion (VSEPR) theory of chemical bonding, the ideal electron geometry of a molecule or a molecular ion containing a total of 4 electron density regions around the central atom is tetrahedral.
In the phosphate [PO4]3- ion, there are 3 single bonds and 1 double bond around the central phosphorus atom which makes a total of 4 electron density regions. Thus, its electron geometry is also tetrahedral.
An easy way to find the shape and geometry of the molecule is to use the AXN method.
AXN is a simple formula to represent the number of atoms bonded to the central atom in a molecule and the number of lone pairs present on it.
It is used to predict the geometry or shape of a molecule using the VSEPR concept.
AXN notation for [PO4]3- molecular ion
A in the AXN formula represents the central atom. In the [PO4]3- ion, phosphorus is present at the center so A = Phosphorous.
X denotes the atoms bonded to the central atom. In [PO4]3-, four oxygen (O) atoms are bonded to the central P so X=4.
N stands for the lone pairs present on the central atom. As per the Lewis structure of [PO4]3- there is no lone pair on central phosphorus so N=0.
So, the AXN generic formula for the [PO4]3- ion is AX4.
Now, you may have a look at the VSEPR chart below.
The VSEPR chart reaffirms that the ideal electron geometry and molecular geometry or shape of a molecule with AX4 generic formula are identical i.e., tetrahedral, as we already noted down for the [PO4]3- ion.
Hybridization of [PO4]3-
The phosphate [PO4]3- ion has sp3 hybridization.
The electronic configuration of a phosphorus (P) atom is 1s2 2s2 2p6 3s2 3p3.
The electronic configuration of an oxygen (O) atom is 1s2 2s2 2p4.
During chemical bonding, the paired 3s electrons of phosphorus get unpaired and one of these electrons shifts to an empty 3d atomic orbital. Consequently, the 3s orbital hybridizes with three half-filled 3p orbitals to yield four sp3 hybrid orbitals. Each sp3 hybrid orbital has a 25 % s-character and a 75% p-character and contains a single electron only.
One sp3 hybrid orbital forms the required sigma (σ) bond with the sp2 hybrid orbital of an oxygen atom by an sp3-sp2 overlap in the P=O bond in the [PO4]3- molecular ion. The other three sp3 hybrid orbitals form P-O sigma (σ) bonds with p orbitals of the remaining three oxygen atoms by sp3-p overlap.
However, the unhybridized d-orbital of phosphorus forms the P=O pi (π) bond by overlapping with the unhybridized p-orbital of the concerned oxygen atom.
A shortcut to finding the hybridization present in a molecule or a molecular ion is by using its steric number against the table given below. The steric number of central P in [PO4]3- is 4 so it has sp3 hybridization.
Steric number
Hybridization
2
sp
3
sp2
4
sp3
5
sp3d
6
sp3d2
The steric number of central Phosphorous in PO43- is 4 so it has sp3 hybridization.
The PO43- bond angle
The bonded atoms in [PO4]3- ion form ideal bond angles as expected in a symmetrical tetrahedral molecule. The O-P-O bond angle is 109.5°. Each P-O bond length in the [PO4]3- ion is also equivalent.
Although a P=O double bond is expected to be stronger and shorter in length than a P-O single bond. But, it is due to the resonance present in the molecule that each P-O bond length and O-P-O bond angle is equal in the [PO4]3- ion.
Each P-O bond present in the [PO4]3- ion is highly polar due to an electronegativity difference of 1.25 units between the covalently bonded phosphorus (E.N = 2.19) and oxygen (E.N = 3.44) atoms. Oxygen more strongly attracts the shared P-O electron cloud as opposed to the phosphorus atom. Thus each P-O bond has a specific dipole moment value (symbol µ).
However, the symmetrical tetrahedral shape of [PO4]3- ion cancels the dipole moment effect of individually polar P-O bonds. The dipole moment of the upwards-pointing P=O bond gets cancelled with the net dipole moment of three downwards-pointing P-O bonds. Thus, the [PO4]3- ion is overall non-polar with net µ =0.
How many bond pairs and lone pairs are present in the [PO4]3- Lewis structure?
There are 5 bond pairs and 11 lone pairs in [PO4]3- Lewis structure. Out of the 11 lone pairs of electrons, there are 2 lone pairs on a doubly-bonded oxygen atom while 3 lone pairs are present on each of the three singly bonded oxygen atoms.
However, there is no lone pair on the central P-atom in [PO4]3- Lewis structure, as you can see in the figure shown below.
How many valence electrons are in PO43- lewis structure?
Total number of valence electrons in Phosphorus = 5
Total number of valence electrons in Oxygen = 6
The [PO4]3- ion consists of 1 P-atom, 4 O-atoms, and a (-3) charge. Thus, the valence electrons in the Lewis structure of [PO4]3- = 1(5) + 4(6) + 3 = 32 valence electrons.
What is the formal charge of each element in [PO4]3- Lewis structure?
There are zero formal charges on the central phosphorus (P) atom and one of the four terminal oxygen (O) atoms in [PO4]3-. However, the other three oxygen atoms contain a -1 formal charge which makes an accumulated charge of -3 on the entire phosphate ion.
This charge is represented at the top right corner of the square brackets enclosing the PO43- Lewis structure.
How many oxygen atoms have charges in [PO4]3- Lewis structure?
Three oxygen atoms have a -1 formal charge each in the PO43- Lewis structure.
What is the electronic geometry of PO43-?
According to the VSEPR concept, the AXN generic formula for PO43- ion is AX4 as it has a total of 4 electron density regions around the central P-atom so its ideal electronic geometry is tetrahedral.
What is the molecular geometry of PO43-?
The molecular geometry or shape of PO43- ion is identical to its electronic geometry i.e., tetrahedral. The absence of any lone pair on the central phosphorus atom ensures that there is no distortion present in the shape of this molecular ion.
The total number of valence electrons available for drawing phosphate [PO4]3- ion Lewis structure is 32.
Negative 3 charge present on the phosphate ion accounts for 3 extra valence electrons in its Lewis structure.
The [PO4]3- ion has an identical electron geometry and molecular geometry or shape i.e., tetrahedral.
Each O-P-O bond angle is 109.5° in [PO4]3-.
It is due to the resonance present in the phosphate [PO4]3- ion that each P-O bond length is equivalent as opposed to a shorter P=O bond and three longer P-O bonds, as expected.
The central P-atom in [PO4]3- ion is sp3.
The [PO4]3- ion is non-polar in nature. The symmetry present in the molecular ion cancels the dipole moment effect of individually polar P-O bonds. Thus, the electron cloud stays uniformly distributed overall (net µ=0).
-1 formal charge is present on each singly bonded O-atom in the [PO4]3- Lewis structure which makes the molecular ion occupy an overall charge of -3.
In conclusion, [PO4]3- is a polyatomic anion that carries a triple negative charge.
Vishal Goyal is the founder of Topblogtenz, a comprehensive resource for students seeking guidance and support in their chemistry studies. He holds a degree in B.Tech (Chemical Engineering) and has four years of experience as a chemistry tutor. The team at Topblogtenz includes experts like experienced researchers, professors, and educators, with the goal of making complex subjects like chemistry accessible and understandable for all. A passion for sharing knowledge and a love for chemistry and science drives the team behind the website. Let's connect through LinkedIn: https://www.linkedin.com/in/vishal-goyal-2926a122b/
Topblogtenz is a website dedicated to providing informative and engaging content related to the field of chemistry and science. We aim to make complex subjects, like chemistry, approachable and enjoyable for everyone.
