Sulfur dioxide (SO2) Lewis dot structure, molecular geometry or
shape, electron geometry, bond angle, formal charge, hybridization
SO2 is the chemical formula for sulfur dioxide, colorless gas that is extremely useful in the chemical industry. The pungent, suffocating odor associated with a burning matchstick is because of SO2. Before we can talk any further about why is SO2 important, we need to know the basic structure and chemical nature of SO2.
To make things easier, we have compiled for you in this article important information about sulfur dioxide (SO2) including its Lewis structure, molecular geometry or shape, electron geometry, bond angle, hybridization, formal charge, etc. So, continue reading!
The Lewis structure of SO2 consists of a sulfur (S) atom present at the center of the molecule. It is bonded with the help of two double bonds to two atoms of oxygen (O) at the sides. There are a total of 5 lone pairs in the SO2 lewis structure(one on the sulfur atom and 2 lone pairs on each oxygen atom).
Let’s draw the Lewis structure of SO2 using the following simple guidelines.
Steps for drawing the Lewis dot structure of SO2
1. Count the total valence electrons in SO2
The Lewis dot structure of a molecule is a simplified representation of all the valence electrons present in the molecule. Therefore, the first step whenever we want to draw the Lewis structure of a molecule is to calculate the total valence electrons present in it.
Valence electrons present in an atomic element can be easily determined by using the Periodic Table. Having a quick look at the Periodic Table, we can readily identify that both Sulfur (S) and oxygen (O) are situated in Group VI A. So, both atoms have 6 valence electrons each.
∴ The SO2 molecule is made up of 1 S atom and 2 O atoms. Therefore, the total valence electrons available for drawing the Lewis structure of SO2 are 6 + 2(6) = 18 valence electrons.
2. Find the least electronegative atom and place it at the center
Electronegativity refers to the tendency of an atom to attract a shared pair of electrons from a covalent chemical bond. In contrast to that, the central position in the Lewis structure of a molecule demands an atom that is most likely to share its electrons with other atoms.
To fulfill this requirement, the least electronegative atom is placed at the center of the Lewis structure.
So, in the Lewis structure of SO2, the S atom is placed at the center while the two O atoms are placed at the outer ends.
3. Connect outer atoms with the central atom
In this step, we join the two outer atoms i.e., O atoms with the central S atom using single bonds, as shown in the figure below.
Now, if we count the total valence electrons used in the above structure, there are two single bonds. Each bond represents an electron pair i.e., 2 electrons. Thus, 2(2) = 4 valence electrons are used till step 3.
Total valence electrons available – electrons used till step 3 = 18-4 = 14 valence electrons.
This shows we still have 14 valence electrons to be placed in the Lewis structure of SO2.
4. Complete the octet of outer atoms
The two O atoms are the outer atoms in the Lewis structure of SO2. Each O atom needs to have a total of 8 valence electrons in order to achieve a stable octet electronic configuration.
Each O already contains 2 electrons, so we need to place 6 more electrons (3 electron pairs) around an oxygen atom to complete its octet, as done in the figure below.
5. Complete the octet of the central atom and make a covalent bond if necessary
Total valence electrons used till step 4 = 2 single bonds + 2 ( electrons placed around each O atom as dots) = 2 (2) + 2 (6) = 16 valence electrons.
Total valence electrons available- electrons used till step 4 = 18 – 16 = 2 valence electrons.
Thus, these 2 electrons are placed as a lone pair on the central S atom.
All 18 valence electrons are now used but still, the problem is that there are a total of 6 valence electrons only around the central sulfur atom which means its octet is not yet complete.
An easy solution to this problem is that we convert one of the three lone pairs present on an O atom into a covalent bond, as shown in the figure below.
In this way, the central S atom achieves a complete octet ( 1 single bond + 1 double bond+ 1 lone pair), also the octet of each O atom on the sides is complete. This gives us the correct Lewis structure for the SO2 molecule.
But the thing is we surely don’t know whether the above Lewis structure is stable or not. The stability of this structure can be checked using the formal charge concept.
6. Check the stability of Lewis’s structure with the help of 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 present on two different atoms in the SO2 molecule can be determined using the formula given below.
The above calculation shows that the formal charges on the sulfur (S) atom is +1 and that on the second oxygen (O) atom is -1 which indicates that the Lewis structure obtained in step 5 is not stable.
This means we need to minimize the charges present on the SO2 Lewis structure in order to get a more stable Lewis representation of sulfur dioxide.
7. Minimize the formal charges present on SO2 atoms to increase the stability of its Lewis structure
The formal charges can be minimized by shifting a lone pair present on the negatively charged O atom onto the positively charged S atom and making another covalent bond.
The Lewis structure obtained above shows one lone pair and two double bonds around the central S atom. Each O atom has a complete octet with two lone pairs and a double bond.
On the other hand, the central S atom has a total of 10 valence electrons which is quite possible because sulfur follows the expanded octet rule. Due to the availability of its 3d orbitals to the incoming electrons, sulfur can accommodate more than 8 valence electrons during chemical bonding.
8. Again, check the stability of the SO2 Lewis structure using the formal charge concept
Zero formal charges on all the atoms present in the SO2 Lewis structure ensures the stability of this structure.
In conclusion, the structure shown above is the best possible Lewis structure of SO2. It is a resonance hybrid of the following two resonance structures.
What are the electron and molecular geometry of SO2?
The sulfur dioxide (SO2) molecule has an asymmetric bent or V-shape molecular geometry. However, the ideal electron geometry of the SO2 molecule is trigonal planar. The molecule adopts a different shape from the ideal electron geometry owing to a lone pair of electrons present on the central S-atom.
Molecular geometry of SO2
The SO2 molecule has an asymmetric bent or V-shape. The sulfur (S) atom lies at the center of the inverted V shape while the oxygen (O) atoms occupy the terminals. The presence of a lone pair of electrons on the central sulfur atom makes sure that the O atoms are maximally pushed away from the center.
The lone pair-bond pair repulsions are significantly stronger than the bond pair-bond pair repulsions so the O-atoms tilt slightly closer to each other but far away from the central S-atom in an angular shape.
It should be noted that the presence of a lone pair has a pronounced effect on the molecular geometry or shape of a molecule. On the other hand, the electron geometry depends on the total number of electron pairs (bond pairs + lone pairs) present on the central atom in the molecule.
Therefore, we seriously take into consideration the repulsive effect of the lone pair on the central S atom in the SO2 molecule when determining its molecular geometry. But this repulsive effect can be ignored while considering the ideal electronic geometry of SO2.
Electron geometry of SO2
According to the valence shell electron pair repulsion (VSEPR) theory of chemical bonding, the ideal electronic geometry of a molecule containing a total of 3 electron pairs (i.e., 3 regions of electron density) around the central atom is trigonalplanar.
A region of electron density refers to the group of bonding and nonbonding electrons present around the central atom in a molecule.
Each double bond in the Lewis structure of the SO2 molecule is considered 1 region of electron density. 2 double bonds and a lone pair make a total of 3 electron density regions around the S atom in the SO2 molecule. So, the electron geometry of SO2 is trigonal planar.
A simpler way of finding the electron and molecular geometry of SO2 is the AXN method.
AXN is a simple formula that is used to represent the number of bonded atoms and lone pairs present on the central atom to predict the shape or geometry of a molecule using the VSEPR concept.
AXN notation for SO2 molecule is:
A in the AXN formula represents the central atom. In the SO2 molecule, sulfur is present at the center so A=S.
X denotes the atoms bonded to the central atom. As 2 oxygen (O) atoms are bonded to the central S atom in the SO2 molecule so X=2.
N stands for the lone pair of electrons situated on the central atom in the molecule. As per the Lewis structure of SO2, only 1 lone pair is present on S so N=1.
In this way, the AXN formula for the SO2 molecule is AX2N1.
Now go through the VSEPR chart given below and identify the electron and molecular geometry associated with AX2N1.
According to the VSEPR chart given above, a molecule with an AX2N1 generic formula has a bent or V-shape while its ideal electron geometry is trigonal planar, as we noted down for the SO2 molecule.
Hybridization of SO2
Each of the O atoms and the central S atom in the SO2 molecule are sp2 hybridized.
The electronic configuration of sulfur (S) is 1s22s22p63s23p4. During chemical bonding, the paired 3p electron gets unpaired and shifts to the 3d orbital of the sulfur atom.
The 3s orbital mix with two 3p orbitals to form three sp2 hybrid orbitals. Two of these hybrid orbitals contain 1 electron each while the third sp2 hybrid orbital contains a pair of electrons. This pair lies as a lone pair on S.
The electronic configuration of oxygen (O) is 1s2 2s2 2p4. Hybridization occurs as the 2s orbital mix with two of the three 2p orbitals of oxygen to form three sp2 hybrid orbitals. Two of the three sp2 hybrid orbitals of oxygen contain an electron pair while the third sp2 has a single electron only.
The sp2 hybrid orbitals containing electron pairs are situated as lone pairs on the oxygen atom. Contrarily, the sp2 hybrid orbitals containing one electron only form a sigma (σ) bond with the sp2 hybrid orbitals of sulfur on each side.
The unhybridized p-orbitals of oxygen form pi (π) bonds with the unhybridized 3p and 3d orbitals of sulfur on either side.
A shortcut for memorizing the hybridization present in a molecule is using its steric number against the table given below. The steric number of central S in the SO2 molecule is 3 as it has three regions of electron density around it. So, the table below shows that it has sp2 hybridization.
Steric number
Hybridization
2
sp
3
sp2
4
sp3
5
sp3d
6
sp3d2
The SO2 bond angle
The ideal bond angle in a trigonal planar molecule is 120°. However, there is a lone pair present on the central S atom. Lone pair-bond pair repulsions push the O atoms away from the central S atom.
The molecule (SO2) attains an asymmetric bent shape. Consequently, the O=S=O bond angle decreases from the ideal 120° to 119°. The S=O bond length is 143.1 pm.
A specific electronegativity difference of 0.86 units exists between the bonded sulfur (E.N= 2.58) and oxygen (E.N=3.44) atoms in a S=O bond.
Pauling’s electronegativity scale states that a polar bond has an electronegativity difference greater than 0.5 units between its bonded atoms.
As 0.86 > 0.5 thus each S=O bond in the SO2 molecule is polar and has a specific dipole moment (symbol μ) value.
The dipole moments of individual S=O bonds add up in the overall asymmetric shape of the SO2 molecule. The electron cloud stays non-uniformly distributed. So, the sulfur dioxide (SO2) molecule is polar with net μ = 1.62 Debye.
The best possible Lewis structure of SO2 is shown below.
According to this structure, one sulfur (S) atom at the center of the molecule is bonded to two atoms of oxygen (O), one on each side.
There are 3 regions of electron density around the central S atom.
2 electron density regions are represented as double bonds on each side while the lone pair present on S denotes the third electron density region.
This SO2 Lewis dot structure is a hybrid of two resonance structures.
Why does sulfur in SO2 have 10 electrons in the Lewis structure?
The sulfur atom has an expanded octet. During chemical bonding, sulfur can accommodate more than 8 valence electrons.
This is due to the availability of the 3d orbital to the incoming valence electrons at a relatively low energy level.
After completely filling the 3p orbitals, the electrons start occupying the 3d subshell.
Why is the Lewis structure of SO2 not similar to O3?
Both S and O atoms belong to Group VI A of the Periodic Table which means they have 6 valence electrons each.
In both SO2 and O3 molecules, the central atom is bonded to two O atoms, so each molecule has a total of 18 valence electrons. However, they have different Lewis structures.
This is because oxygen (O) unlike sulfur (S) does not follow the expanded octet rule. It can accommodate a total of 8 valence electrons only.
The Lewis structure of O3 is represented by the following two resonance forms. The central O atom in both forms has a total of 8 valence electrons.
The Lewis structure of SO2 is a resonance hybrid of the following two resonance forms. The central S atom in the resonance hybrid has a total of 10 valence electrons.
How many covalent bonds are present in the Lewis structure of SO2?
The Lewis structure of SO2 has a total of 4 covalent bonds.
One double bond (i.e., a sigma and a pi covalent bond) is present between S and O atoms on each side of the molecule.
What are the similarities and differences in SO2 and SO3 Lewis structures?
Similarities:
Both the Lewis structures contain a sulfur (S) atom at the center which is bonded to oxygen (O) atoms at the terminals.
The total electron density regions in both the molecules are 3.
Differences:
The Lewis structure of SO2 has a total of 18 valence electrons while the Lewis structure of SO3 displays a total of 24 valence electrons.
There are 2 double bonds and a lone pair on the central S atom in the SO2 Lewis structure which makes it occupy a bent shape. Contrarily, there are 3 double bonds and no lone pair on central S in the Lewis structure of SO3, so it has an identical electron and molecular geometry or shape i.e., trigonal planar.
The total valence electrons available for drawing sulfur dioxide (SO2) Lewis structure are 18.
SO2 molecule has a bent or V-shape and molecular geometry.
The electron geometry of SO2 is trigonal planar.
The Lewis dot structure of SO2 has two different resonance forms.
The actual Lewis structure of SO2 is a hybrid of both resonance forms. It consists of two double bonds and a lone pair on the central S atom.
The central S and both the O atoms are sp2 hybridized in SO2.
The ideal O=S=O bond angle in a trigonal planar molecule is 120° but due to lone pair-bond pair repulsions, the O=S=O bond angle decreases in the SO2 molecule and becomes 119°.
The S=O bond length in SO2 is 143.1 pm.
SO2 is a polar molecule overall with net μ=1.62 D.
The overall formal charge on SO2 Lewis structure is zero so it is a stable molecule.
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.
