How to calculate formal charges of sulfur dioxide (SO2) with lewis structure?
In covalently bonded molecules, formal charge is the charge assigned to an atom based on the assumption that the bonded electrons are equally shared between concerning atoms, regardless of their electronegativity.
The overall formal charge present on a molecule is a measure of its stability.
The fewer the formal charges present on the bonded atoms in a molecule (close to zero), the greater the stability of its Lewis structure.
In this article, we will calculate the formal charges present on bonded atoms in the different resonance structures of sulfur dioxide (SO_{2}).
This will ultimately help us draw the best and most stable Lewis representation of SO_{2} based on the formal charge concept.
So, continue reading!
Name of the molecule | Sulfur Dioxide |
Chemical formula | SO_{2} |
The formal charge on the S-atom | 0 |
The formal charge on each double-bonded O-atom | 0 |
The overall formal charge on SO_{2} | 0 |
The formal charge data in the above table is according to the most preferred Lewis representation of SO_{2}, as shown below.
This structure consists of a total of 18 valence electrons. A sulfur (S) atom is bonded to two atoms of oxygen (O) via double covalent bonds. There is 1 lone pair of electrons on the central S-atom, while 2 lone pairs are present on each outer O-atom.
However, the following two different resonance structures are also possible for representing SO_{2}.
So let us first calculate the formal charges according to each resonance structure so that we can better understand how is the most preferred SO_{2} Lewis structure obtained.
How to calculate the formal charges on SO_{2} atoms?
The formal charges can be calculated using the formula given below:
The formal charge of an atom = [valence electrons of an atom – non-bonding electrons – ½ (bonding electrons)]
- The valence electrons (V.E) of an atom are the total number of electrons present in its valence shell. Valence electrons can be determined by locating the position of the elemental atom in the Periodic Table.
- Non-bonding electrons (N.E) are the number of lone pairs present on the atom. (1 lone pair means 2 nonbonding electrons).
- Bonding electrons (B.E) are the total electrons shared with the atom via covalent chemical bonds. (1 single bond means 2 bonding electrons).
Now let us use this formula to calculate the formal charges in SO_{2} resonance structures.
For the central Sulfur atom
- Valence electrons of Sulfur = It is present in Group VI A = 6 valence electrons
- Bonding electrons = 1 double bond + 1 single bond = 1(4) + 1(2) = 6 electrons
- Non-bonding electrons = One lone pair = 2 electrons
- Formal charge on Sulfur atom = 6 – 2 –6 / 2 = 6 – 2 – 3 = 6 – 5 = +1
∴ The formal charge on the S-atom in SO_{2 }resonance structures is +1.
For double-bonded oxygen atom
- Valence electrons of oxygen = It is present in Group VI-A = 6 valence electrons
- Bonding electrons = 1 double bond = 4 electrons
- Non-bonding electrons = 2 lone pairs = 2(2) = 4 electrons
- Formal charge = 6 – 4 – 4/2 = 6 – 4 – 2 = 6 – 6 = 0
∴ The formal charge on the double-bonded oxygen (O) atom in SO_{2} resonance structures is 0.
For each single-bonded oxygen atom
- Valence electrons of oxygen = It is present in Group VI-A = 6 valence electrons
- Bonding electrons = 1 single bond = 2 electrons
- Non-bonding electrons = 3 lone pairs = 3(2) = 6 electrons
- Formal charge = 6 – 6 – 2/2 = 6 – 6 – 1 = 6 – 7 = -1
∴ The formal charge on the single-bonded oxygen (O) atom in SO_{2} resonance structures is -1.
The above calculation shows that although no formal charges are present on the double bonded O-atom in each SO_{2} resonance structure. However, a +1 and a -1 formal charge is present on the central S-atom and the single-bonded O-atom, respectively.
Therefore, to reduce this formal charge, a lone pair of electrons on the single-bonded O-atom is converted into a covalent chemical bond between the respective O-atom and the central S-atom. This gives us the most preferred SO_{2} Lewis structure.
Now let’s calculate the formal charges in the most preferred Lewis representation.
For the central Sulfur atom
- Valence electrons of Sulfur = It is present in Group VI-A = 6 valence electrons
- Bonding electrons = 2 double bonds = 2(4) = 8 electrons
- Non-bonding electrons = One lone pair = 2 electrons
- Formal charge = 6 – 2 –8/2 = 6 – 2 – 4 = 6 – 6 = 0
∴ The formal charge on the S-atom in SO_{2 }is 0.
For double-bonded oxygen atom
- Valence electrons of oxygen = It is present in Group VI-A = 6 valence electrons
- Bonding electrons = 1 double bond = 4 electrons
- Non-bonding electrons = 2 lone pairs = 2(2) = 4 electrons
- Formal charge = 6 – 4 – 4/2 = 6 – 4 – 2 = 6 – 6 = 0
∴ The formal charge on each double-bonded O-atom in SO_{2} is 0.
Zero formal charges on all the bonded atoms in the most preferred SO_{2} Lewis structure marks the extraordinary stability of this structure.
In this way, the formal charges are minimized, and no overall formal charge is present on the SO_{2} molecule; thus, it is the best possible Lewis representation of SO_{2}. In reality, it is a hybrid of the two SO_{2} resonance structures.
Also, check –
- How to draw SO_{2} lewis structure?
- Formal charge calculator
- SO_{3} formal charge
- CO_{2} formal charge
- HCN formal charge
- SO_{4}^{2-} formal charge
- PO_{4}^{3-} formal charge
- SO_{3}^{2-} formal charge
- CN^{–} formal charge
- ClO_{3}^{–} formal charge
- O_{3} formal charge
- SCN^{–} formal charge
- POCl_{3} formal charge
- NH_{3} formal charge
- CO formal charge
- H_{2}O formal charge
- NH_{4}^{+} formal charge
- H_{3}O^{+} formal charge
- OH^{–} formal charge
- HSO_{4}^{–} formal charge
- ClO^{–} formal charge
- BH_{4}^{–} formal charge
- N_{3}^{–} formal charge
- H_{2}SO_{4} formal charge
- NCO^{–} formal charge
- NO_{3}^{–} formal charge
- NO_{2}^{–} formal charge
- CH_{3} formal charge
FAQ
How can you calculate SO_{2} formal charges? |
The formal charges present on the bonded atoms in SO_{2} can be calculated using the formula given below: V.E – N.E – B.E/2 Where – ⇒ V.E = valence electrons of an atom ⇒ N.E = non-bonding electrons, i.e., lone pairs ⇒ B.E = bonding electrons |
How many resonance structures are possible for SO_{2}? |
The SO_{2} molecule consists of the following two resonance structures. |
In the most stable SO_{2} Lewis structure, what is the formal charge on the S-atom? |
The central Sulfur (S) atom carries zero formal charge in the most stable SO_{2} Lewis structure. |
What is the best possible Lewis representation for SO_{2}? |
The best possible Lewis representation for SO_{2} is the one in which the formal charges are minimized. |
What is the formal charge on O-atoms in SO_{2}? |
Zero or no formal charge is present on both double-bonded (S=O) oxygen atoms in SO_{2}. |
What is the overall formal charge on SO_{2}? |
No overall formal charge is present on the SO_{2} Lewis structure. |
Do both O-atoms carry the same formal charge in SO_{2}? |
Yes, both the O-atoms carry zero formal charges in the best possible SO_{2} Lewis structure. |
Summary
- The best possible Lewis structure of a molecule is the one in which the bonded atoms carry formal charges as close to zero as possible.
- The formal charge formula is [ V.E – N.E – B.E/2].
- Two different resonance structures are possible for drawing the sulfur dioxide (SO_{2}) molecule.
- The most preferred Lewis structure of SO_{2} is a hybrid of these resonance structures.
- In this structure, the central S-atom carries zero formal charges.
- Both the O-atoms also carry zero formal charges.
- The overall formal charge present in SO_{2} is also zero; thus, it is a neutral molecule.
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