How to calculate formal charges of Sulfate (SO42-) ion 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.
The sulfate [SO_{4}]^{2-} ion present in salts can create a lathering effect to remove dirt, which is the main reason it is present in household cleaners, detergents, and even in your shampoos.
In this article, we will calculate the formal charges present on the bonded atoms in [SO_{4}]^{2- }and also the overall charge present on the molecular ion.
So, continue reading!
Name of the molecular ion | Sulfate |
Chemical formula | [SO_{4}]^{2-} |
The formal charge on the S-atom | 0 |
The formal charges on double-bonded O-atoms | 0 |
The formal charge on single-bonded O-atoms | -1 |
The overall formal charge on [SO_{4}]^{2-} | -2 |
How to calculate the formal charges on SO_{4}^{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 the most preferred Lewis structure of Sulfate ion [SO_{4}]^{2-}.
The most preferred Lewis representation of [SO_{4}]^{2-} is as shown below.
It consists of a total of 32 valence electrons. A Sulfur (S) atom is present at the center, which is bonded to four oxygen (O) atoms. Two atoms of oxygen (O) are bonded to the central S-atom via double covalent bonds, while the remaining two oxygen atoms are bonded via single covalent bonds, respectively.
No lone pair of electrons is present at the central S-atom. Each of the two double-bonded O-atoms contains 2 lone pairs while each of the two single-bonded O-atoms contains 3 lone pairs of electrons.
It is the best possible Lewis structure of Sulfate ion [SO_{4}]^{2-} because the formal charges are minimized in it, and thus, it is the most stable.
Let’s find out how we can determine the formal charges present on each atom in [SO_{4}]^{2- }Lewis structure.
For the central Sulfur atom
- Valence electrons of Sulfur = It is present in Group VI A = 6 valence electrons
- Bonding electrons = 2 single bonds + 2 double bonds = 2(2) + 2(4) = 12 electrons
- Non-bonding electrons = no lone pair = 0 electrons
- Formal charge = 6 – 0 – 12/2 = 6 – 0 – 6 = 6 – 6 = 0
∴ The formal charge on the Sulfur (S) atom in [SO_{4}]^{2-} is 0.
For single-bonded oxygen atoms
- Valence electrons of oxygen = It is present in Group VI A = 6 valence electrons
- Bonding electrons around Oxygen = 1 single bond = 2 electrons
- Non-bonding electrons on Oxygen = 3 lone pairs = 3(2) = 6 electrons
- Formal charge on the single bonded Oxygen atom = 6 – 6 – 2/2 = 6 – 6 – 1 = 6 – 7 = -1
∴ The formal charge on each single-bonded oxygen (O) atom in [SO_{4}]^{2-} is -1.
For double-bonded oxygen atoms
- Valence electrons of oxygen = It is present in Group VI A = 6 valence electrons
- Bonding electrons around Oxygen = 1 double bond = 4 electrons
- Non-bonding electrons on Oxygen = 2 lone pairs = 2(2) = 4 electrons
- Formal charge on the double bonded Oxygen = 6 – 4 – 4/2 = 6 – 4 – 2 = 6 – 6 = 0
∴ The formal charge on each double-bonded oxygen (O) atom in [SO_{4}]^{2-} is 0.
This calculation shows zero formal charge present on the central S atom and two double-bonded O-atoms, while each of the two single-bonded O-atoms carries -1 formal charge, when added up (-1 + -1 = -2) becomes equal to -2, which accounts for an overall -2 formal charge present on the Sulfate ion [SO_{4}]^{2-}, as shown below.
Consequently, the SO_{4}^{2-} Lewis structure is enclosed in square brackets, and a -2 charge is placed at the top right corner.
Also, check –
- How to draw SO_{4}^{2-} lewis structure?
- Formal charge calculator
- SO_{3} formal charge
- CO_{2} formal charge
- HCN formal charge
- ClO_{3}^{–} formal charge
- PO_{4}^{3-} formal charge
- SO_{3}^{2-} formal charge
- CN^{–} formal charge
- SO_{2} 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_{4}]^{2-} formal charges? |
The formal charges present on the bonded atoms in [SO_{4}]^{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 |
What is the formal charge on S-atom in [SO_{4}]^{2-}? |
The central Sulfur (S) atom carries zero formal charges in [SO_{4}]^{2-}. |
What is the formal charge on O-atoms in [SO_{4}]^{2-}? |
-1 formal charge is present on single (S-O) bonded oxygen atoms in [SO_{4}]^{2- }while zero or no formal charge is present on each double (S=O) bonded oxygen atom in [SO_{4}]^{2-}. |
What is the overall formal charge on [SO_{4}]^{2-}? |
The overall formal charge on [SO_{4}]^{2-} is –2. |
Do all four O-atoms carry the same formal charge in [SO_{4}]^{2-}? |
No, the single bonded O-atoms carry a -1 formal charge, while no formal charge is present on each double bonded O-atom in [SO_{4}]^{2-} Lewis structure. |
Summary
- The best possible Lewis structure of a molecule or molecular ion 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].
- In [SO_{4}]^{2-}, zero formal charges are present on the central S-atom.
- The double-bonded O-atoms also have zero formal charges in [SO_{4}]^{2- }.
- The single-bonded O-atoms carry -1 formal charge each in [SO_{4}]^{2- }Lewis structure.
- The overall formal charge on [SO_{4}]^{2-} ion is -2.
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