In the discipline of chemistry, carbonates are one of the most regularly seen and discussed ionic substances. Carbonates, a salt of carbonic acid, are widely employed in a wide range of industrial and home applications. Glass and pottery manufacture, food preservation, and iron extraction are only a few of them.
With a few exceptions, the CO32- ion is the simplest oxocarbon anion that decomposes on heating and is usually water insoluble.
Let’s take a closer look at the CO32- ion’s chemical bonding.
Lewis Structure CO32
You’ve probably already heard of this term if you’re reading this article, right?
It goes without saying that while learning about the nature of chemical bonding between atoms and molecules, Lewis Structure is a topic that cannot be ignored.
Simply sketch a 2D diagrammatic representation of the provided molecule to gain a quick and clear summary of the atomic bonding among elements. Lewis Structure is the name given to a skeletal diagram in which the atoms are represented by symbols and the valence shell electrons are represented by dots.
As a result, Lewis Structure is also known as Electron Dot Structure.
Let’s start by drawing the CO32- ion’s most acceptable LS diagram.
Count the total number of Valence Electrons in Step 1.
One carbon atom, three oxygen atoms, and two negatively charged electrons carry the charge in the CO32- ion.
The number of electrons in an atom’s outermost shell around the nucleus that assist determine its valency is referred to as valence electrons. The value can be simply deduced from the atomic number listed in the periodic chart.
When it comes to the carbonate ion,
Carbon has a valency of 4 and an atomic number of 6. The atomic number of oxygen is 8 and its valency is 6.
CO32- has a total valence electron number of
4 + 63 + 2 = 4 + 63 + 2
Step 2: Identify the Molecule’s Central Atom.
To draw the Lewis Structure, we must first establish which atom in a multiatomic heterogeneous molecule, such as an ion, is the centre atom.
Carbon has an electronegativity value of 2.55 in the carbonate ion, while Oxygen has a high value of 3.44.
The atom with the lowest electronegativity value will serve as the centre atom, as per standard protocol.
The central atom in this picture is carbon.
Step 3: Draw the Molecule’s Skeletal Diagram.
We’ll be able to construct the principal sketch of the carbonate ion using dots for valence electrons and atomic symbols for the elements.
We’ll need to start by incorporating the octet rule.
The Rule of the Octet
The elements in the core group usually adhere to the octet fulfilment principle. This suggests that, like noble gas configurations of the same time, these atomic elements will tend to have eight valence electrons.
Let’s make a skeleton diagram for the CO32- ion:
Step 4: Forming Bonds
We’ve drawn the sketch here, as you can see.
The total number of valence electrons in the system has been set at 24.
Carbon and each of the three oxygen atoms have two electrons in common, indicating the presence of single bonds.
For all of the oxygen atoms, the octet rule has been met.
However, because the Carbon atom has only six electrons, the octet configuration is not met.
We can consider a double bond between any two oxygen atoms and carbon, resulting in carbon having eight valence electrons surrounding it.
Consider the Lewis Structure.
Each oxygen molecule has an octet arrangement.
We have a carbon molecule with an octet configuration.
Between carbon and each of the two oxygen atoms, there are two single bonds.
One double bond exists between C and one O atom.
The Lewis structure looks like this after drawing bonds:
It can also be seen in the image below.
Step 5: Charge Formally
We’ll need to check the formal charge values to see if the preceding design represents the best feasible Lewis Structure of the Carbonate (CO32-) ion.
We sometimes assign a charge to a bonded atom based on the assumption that the charge is distributed evenly among all bonded atoms. The formal charge is what it’s called.
The formal charge formula is as follows:
Let’s see what we can find out for CO32-:
Formal charge for carbon is 4 – 0.5*8 – 0 = 4 – 4 = 0.
Formal charge for each O in a single bond with carbon
-1 = 6 – 0.5*2 – 6 = 6 – 1 – 6
Formal charge for the O atom in a double bond with carbon
6 – 0.5*4 – 4 = 6 – 2 – 4 = zero
Every atom’s formal charge values are kept at their lowest attainable levels, as can be shown. As a result, our Lewis Structure is now complete.
CO32 Molecular Geometry
Is a 2D structure sufficient for a thorough comprehension of the bonding that occurs within a molecule?
A well-drawn Lewis Structure, on the other hand, explains the basic depiction of constituent atoms inside any molecule or ion, as well as the types of bonds that are formed.
However, this is insufficient.
The VSEPR theory (Valence Shell Electron Pair Repulsion Theory) is a model for establishing the three-dimensional character of any chemical composition.
This is referred to as molecular geometry, which includes not only the molecule’s structure but also the bond lengths and angles.
This allows us to get a better and clearer picture of the molecule.
The AXn notation can be used to determine the exact molecular geometry of the CO32- ion.
VSEPR notation is what it’s called. We discuss the minimum repulsion that occurs between negatively charged electron clouds in order to achieve a balanced molecular composition in this theory.
In the notation AXn,
The centre atom is denoted by the letter A.
The number of atoms surrounding the core atom is denoted by the letter X.
The number of bonds linked to the centre atom in the molecule is denoted by the letter n.
The number of lone pairs (non-bonded electron pairs) in the centre atom is denoted by Ex.
When it comes to the carbonate ion,
A= atom of carbon, B= atom of carbon, C= atom of
X stands for the atom of oxygen.
n = 3
Ex = 0
As a result, AX3 is the needed notation.
Let’s have a look at the graph above. We can forecast our molecular shape using the VSEPR chart with AXnEx notations.
The AX3 designation, like the CO32- ion, has a trigonal planar structure with a bond angle of roughly 120 degrees.
If you’re a chemical student, you’re probably familiar with the distinction between an orbit and an orbital.
While orbit refers to an electron’s precise journey around an atomic nucleus, orbital refers to the likelihood of electrons being present in any given region.
Atomic orbitals come in a variety of shapes, including spherical and dumb-bell shapes. As a result, they are referred to as s,p,d, and f.
When chemical bonding takes place, these AOs join together to produce hybridised orbitals, which are involved in the creation of bonds inside a molecule.
Hybridization is the term for this procedure.
Looking at the formal charge notion in Lewis Structure once again, we can see that each of the singly bound O atoms in C-O has a negative charge of -1. The doubly bound O atom in C=O, on the other hand, has no charge value.
The presence of sigma bonds is shown by single bonds, whereas the presence of both sigma and pi bonds is indicated by double bonds.
As a result, we have three sigmas and one pi in the carbonate ion’s core C atom. The formula for quickly determining the H (Hybridization value) of an atom within a molecule is given below.
V = 4, M = 0, C = 0, A = 2 in this case.
3 = H = 0.5 (4 + 0 – 0 + 0 + 0 + 0 + 0 + 0 + 0 + 0 + 0 +
We have sp2 hybridization for the value of three electron pairs.
Diagram of the Molecular Orbital MO of CO32
What exactly is MO theory?
The chemical bonding nature inside different molecular structures is deciphered using Molecular Orbital Theory, a quantum mechanics concept.
This is a complicated but important tool that aids in the creation of MO diagrams for better comprehension. In this theory, electrons are seen as both particles and waves.
We can assume valence electrons to be shared across all atoms, unlike in valence bond theory, where AOs from the same atom can only fuse to generate hybridised orbitals, resulting in hybridization.
As a result, AOs from various atoms can come together to form Molecular Orbitals ( MOs ).
Delocalized pi bonding is the best explanation for molecular orbital theory in the carbonate ion. The electrons in a delocalized pi bond are allowed to migrate across many nuclei, allowing the pi () to appear in numerous conformations.
We discovered that there are 24 valence electrons when designing the Lewis Structure. There are four non-bonding electrons on each oxygen atom in the ion.
We have three sigma bonds, so we’ll utilise six.
6 is the atomic number of carbon.
8 is the atomic number of oxygen.
C has the following electronic configuration: 1s2 2s2 2p2.
O has the following electronic configuration: 1s2 2s2 2p4.
For delocalized pi bonding, the 2pz orbitals of carbon and three O atoms are available. Two electrons fill bonding molecular orbitals, while four electrons fill non-bonding MOs.
As a result, the six pi electrons available are allocated to the lowest-energy MOs — the bonding MOs. This is an illustration of a four-center MO therapy.
We’ve gone over the chemical bonding nature of the well-known carbonate anion in great depth in this comprehensive post.
We’ve gone over how Lewis Structure is formed, as well as the exact molecular geometry and bond angles of 3D CO32-. In addition, we’ve looked into orbital hybridization and quantum MO theory.
Read more: Is CH2Cl2 a polar or a nonpolar compound?