Carbonate ion (CO32-) Lewis dot structure, molecular geometry or
shape, electron geometry, bond angle, formal charge,
hybridization
CO32- is the chemical formula for carbonate ion, a polyatomic ion composed of 1 carbon and 3 oxygen atoms. It is present in a carbonate salt i.e., a salt of carbonic acid (H2CO3). It is widely used as a raw material in the industrial sector, especially in glass making, paper, and pulp industry, and for the production of soaps and detergents.
Considering the immense importance and applications of CO32- for a chemist, we have compiled for you in this article interesting information about the carbonate [CO3]2- ion.
We will learn how to draw the Lewis structure of CO32-, what is its molecular geometry or shape, electron geometry, bond angle, hybridization, formal charge, polarity nature, etc. So, what are you waiting for? Let’s start reading.
The Lewis structure of carbonate [CO3]2- ion is made up of a carbon (C) atom and three oxygen (O) atoms. The carbon (C) is present at the center of the molecular ion while oxygen (O) occupies the terminals, one on each side.
There are a total of 3 electron density regions around the central C atom in the Lewis structure of [CO3]2-. All the electron density regions are constituted of bond pairs which denotes there is no lone pair of electrons on the central C-atom in [CO3]2-.
You can draw the Lewis dot structure of [CO3]2- by following the simple steps given below.
Steps for drawing the Lewis dot structure of [CO3]2-
1. Count the total valence electrons in [CO3]2-
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 [CO3]2- is to count the total valence electrons present in the concerned elemental atoms.
There are two different atoms present in [CO3]2- ion i.e., the carbon (C) atom and the oxygen (O) atoms. The total valence electrons present in an atom of carbon and oxygen can be determined by identifying these elements in the Periodic Table.
Carbon (C) is present in Group IV A of the Periodic Table of elements so it has 4 valence electrons. On the other hand, oxygen (O) is present in Group VI A so the total valence electrons present in it are 6.
The [CO3]2- ion consists of 1 C-atom and 3 O-atoms. Thus, the valence electrons in the Lewis dot structure of [CO3]2- = 1(4) + 3(6) = 22 valence electrons.
An important point to remember is that the [CO3]2- ion carries a negative (-2) charge which means 2 extra valence electrons are added in this Lewis structure.
∴ Hence, the total valence electrons available for drawing the Lewis dot structure of [CO3]2- = 22+2 = 24 valence electrons.
2. Choose the central atom
Electronegativity is defined as the ability of an atom to attract a shared pair of electrons from a chemical bond. So, the atom which is least electronegative or most electropositive is placed at the center of a Lewis structure. This is because this atom is most likely to share its electrons with the more electronegative atoms surrounding it.
As carbon (C) is less electronegative than oxygen (O) so, a C-atom is placed at the center of the [CO3]2- Lewis structure while the three O-atoms are placed in its surroundings, as shown below.
3. Connect outer atoms with the central atom
At this step of drawing [CO3]2- Lewis structure, we need to connect the outer atoms with the central atom using single straight lines. As the three O-atoms are the outer atoms in the [CO3]2- Lewis structure while C is the central atom. So all the 3 O-atoms are joined to the central C-atom using straight lines, as shown below.
Each straight line represents a single covalent bond i.e., a bond pair containing 2 electrons. There are a total of 3 straight lines in the above structure which means 3(2) = 6 valence electrons are used so far, out of the 24 initially available.
Total valence electrons available – electrons used till step 3 = 24-6 = 18 valence electrons.
This means we still have 18 valence electrons to be accommodated in the Lewis dot structure of [CO3]2-.
4. Complete the octet of outer atoms
There are three O-atoms present as outer atoms in the Lewis structure of [CO3]2-. Each O-atom needs a total of 8 valence electrons in order to achieve a stable octet electronic configuration.
Each C-O bond already represents 2 electrons therefore all the 3 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 and make a covalent bond if necessary
Total valence electrons used till step 4 = 3 single bonds + 3 (electrons placed around an O-atom, shown as dots) = 3(2) +3(6) = 24 valence electrons.
Total valence electrons available – electrons used till step 4 = 24 – 24 = 0 valence electrons.
As all the valence electrons available are already used so there is no lone pair on the central C-atom in the [CO3]2- Lewis structure.
But, the problem that stays is that there are 3 single bonds around the central C-atom in the Lewis structure drawn so far. 3 single bonds represent a total of 6 valence electrons only. Consequently, the central C-atom still needs 2 more electrons to achieve a complete octet.
An easy solution to this problem is that we can convert one of the lone pairs from any one outer O atom into a covalent bond between C and O atoms.
In this way, the central C-atom has a complete octet with 2 single bonds + 1 double bond. Also, the octet of the outer O-atom is complete with 1 double bond + 2 lone pairs. The octet of the other two O-atoms is also complete with 1 single double +3 lone pairs, each. Refer to the figure given below.
The final step is to check whether the Lewis structure of the carbonate ion obtained above is stable or not. So let’s do it using the formal charge concept.
6. Check the stability of the CO32- 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 shows that a -1 charge is present on each single-bonded O-atom in the Lewis structure of [CO3]2-, while no formal charges on the central C-atom and the double-bonded O-atom, as shown below.
The overall formal charge in the CO32- Lewis structure is -1 + (-1) = -2 which is equal to the charge present on the ion. Thus, it is a correct and stable Lewis structure for the carbonate ion. The carbonate ion Lewis structure is enclosed in square brackets and a negative 2 charge is placed at the top right corner.
An interesting fact is that the actual structure of a carbonate [CO3]2- 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 [CO3]2- atoms are not stationary, rather they keep moving from one position to another. Similarly, a double bond can be formed between the central carbon and any one oxygen atom out of all the three present.
In conclusion, all the above three structures contribute equally to the resonance hybrid.
Now, that we have discussed everything about the Lewis structure of [CO3]2-, let us proceed forward to the next section of the article.
What are the electron and molecular geometry of CO32-?
The carbonate [CO3]2- ion has an identical electron and molecular geometry or shape i.e., trigonal planar. The three O-atoms bonded to the central C-atom lie in the same plane, in a symmetrical arrangement.
No lone pair of electrons is present on the central C-atom in [CO3]2- thus there is no distortion present in the geometry or shape of the molecule.
Molecular geometry of [CO3]2-
The carbonate [CO3]2- ion has a trigonal planar shape and molecular geometry. The O-atoms lie at the vertices of an equilateral triangle while the C-atom is present at the center. There is no lone pair on the central C-atom so no lone pair-bond pair and lone pair-lone pair repulsions exist in the molecular ion.
However, the bond pair-bond pair electronic repulsions between C=O and two C-O bonds keep the terminal O-atoms maximally separated.
Electron geometry of [CO3]2-
The valence shell electron pair repulsion (VSEPR) theory of chemical bonding states that the ideal electronic geometry of a molecule containing 3 regions of electron density around the central atom is trigonal planar.
As per 2 single bonds and 1 double bond, there are a total of 3 electron density regions around the central C-atom in [CO3]2- ion. Thus the ideal electron geometry of carbonate [CO3]2- ion is trigonal planar.
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 [CO3]2- molecular ion
A in the AXN formula represents the central atom. In the [CO3]2- ion, carbon is present at the center so A = Carbon.
X denotes the atoms bonded to the central atom. In [CO3]2-, three Oxygen (O) atoms are bonded to the central C so X = 3.
N stands for the lone pairs present on the central atom. As per the Lewis structure of [CO3]2- there is no lone pair on carbon so N = 0.
So, the AXN generic formula for the [CO3]2- ion is AX3.
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 AX3 generic formula are identical i.e., trigonal planar, as we already noted down for the [CO3]2- ion.
Hybridization of [CO3]2-
The central C atom is sp2 hybridized in the [CO3]2- ion.
The electronic configuration of carbon (C) is 1s2 2s2 2p2.
The electronic configuration of oxygen (O) is 1s2 2s2 2p4.
During chemical bonding, the 2s electrons of carbon get unpaired. One of the two 2s electrons gets excited and shifts to the empty 2p atomic orbital of carbon. This leaves behind half-filled 2s and 2p atomic orbitals.
The 2s orbital mixes with two 2p orbitals to yield three equivalent sp2 hybrid orbitals, each containing 1 electron. Each sp2 hybrid orbital has a 33.3 % s-character and a 66.7% p-character.
Two sp2 hybrid orbitals of carbon overlap with the p-orbitals of oxygen (O) to form two C-O sigma (σ) bonds, one on each side.
Meanwhile, the third sp2 hybrid orbital forms a sigma (σ) bond with the sp2 hybrid orbital of the third oxygen which is involved in C=O bond formation. The unhybridized p-orbital of carbon forms the pi (π) bond with this oxygen atom in C=O via p-p overlap.
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 C in [CO3]2- is 3 so it has sp2 hybridization.
Steric number
Hybridization
2
sp
3
sp2
4
sp3
5
sp3d
6
sp3d2
The steric number of central carbon in CO32- is 3 so it has sp2 hybridization.
The [CO3]2- bond angle
The bonded atoms in [CO3]2- ion form ideal bond angles as expected in a symmetrical trigonal planar molecule. The O-C-O bond angle is 120°. Each C-O bond length in the [CO3]2- ion is equivalent i.e., 128 pm.
Although a C=O double bond is expected to be shorter in length than a C-O single bond. But, it is due to the resonance present in the molecule that each C-O bond length is equal in the [CO3]2- ion.
Each C-O bond present in the [CO3]2- ion is polar due to an electronegativity difference of 0.89 units between the covalently bonded carbon (E.N = 2.55) and oxygen (E.N = 3.44) atoms.
Oxygen more strongly attracts the shared C-O electron cloud as opposed to the carbon atom. Thus each C-O bond has a specific dipole moment value (symbol µ).
However, the symmetrical shape of [CO3]2- ion cancels the dipole moment effect of C-O bonds. The dipole moment of the upwards-pointing C=O bond gets canceled with the net dipole moment of two downwards-pointing C-O bonds. Thus, the [CO3]2- ion is overall non-polar with net µ =0.
There are a total of 24 valence electrons = 24/2 = 12 electron pairs in the Lewis structure of CO32-.
Out of these 12 electron pairs, there are 4 bond pairs and 8 lone pairs.
The central C-atom is bonded to an oxygen (O) atom via a double covalent bond while it is bonded to two other oxygen (O) atoms via single bonds.
There is no lone pair on the central C-atom. However, 3 lone pairs of electrons are present on each of the single-bonded O-atoms while 2 lone pairs are present on the double-bonded O-atom in the lewis structure of CO32-.
How many valence electrons are in CO32- lewis structure?
Total number of valence electrons in carbon = 4
Total number of valence electrons in oxygen = 6
The [CO3]2- ion consists of 1 C-atom, 3 O-atoms, and a (-2) charge. Thus, the valence electrons in the Lewis dot structure of [CO3]2- = 1(4) + 3(6) + 2 = 24 valence electrons.
What are the electron and molecular geometries, respectively, for the carbonate CO32- ion?
The carbonate CO32- ion has an identical electron and molecular geometry or shape i.e., trigonal planar.
What are the molecular geometries of BeCl2, [CO3]2-, CH4, and [H3O] +?
Both beryllium chloride (BeCl2) and methane (CH4) are neutral molecules while carbonate [CO3]2- and hydronium [H3O]+ are molecular ions with negative and positive charges respectively.
According to the AXN generic formula and the VSEPR concept, BeCl2 is an AX2-type molecule and its molecular geometry is linear.
CH4 is an AX4-type molecule and has a tetrahedral geometry.
[CO3]2- is an AX3-type molecular ion and its molecular geometry is trigonal planar.
The AXN generic formula for the [H3O]+ ion is AX3N1 and according to the VSEPR concept, its molecular geometry is triangular pyramidal. The presence of a lone pair on the central O-atom leads to bond pair-lone pair repulsions which makes the molecular ion occupy a different shape or molecular geometry from its ideal electron geometry which is tetrahedral.
How can the bond angle of CO32- ion be explained according to the VSEPR theory?
According to the VSEPR theory, the ideal bond angle of a trigonal planar molecule or molecular ion such as CO32- is 120°.
The absence of any lone pair on the central C-atom in the carbonate ion ensures there is no distortion present in the shape and symmetry of the molecule thus the O-C-O bonds stay at their ideal bond angle.
The total valence electrons available for drawing the carbonate ion[CO3]2- Lewis structure are 24.
The negative 2 charge present on the carbonate ion accounts for 2 extra valence electrons in its Lewis structure.
The [CO3]2- ion has an identical electron geometry and molecular geometry or shape i.e., trigonal planar.
The O-C-O bond angle is 120° while the C-O bond lengths are 128 pm in [CO3]2- ion.
It is due to the resonance present in the carbonate [CO3]2- ion that each C-O bond length is equivalent as opposed to a shorter C=O bond and two longer C-O bonds, as expected.
The central C-atom in [CO3]2- ion is sp2 hybridized.
The [CO3]2- ion is non-polar in nature. The symmetry present in the molecular ion cancels the dipole moment effect of individually polar C-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 [CO3]2- Lewis structure which makes the molecular ion occupy an overall charge of -2. In conclusion, [CO3]2- is a polyatomic dianion.
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