Oxygen is a chemical element with an atomic number 8 and symbol O. It lies in group 16 (chalcogen group) of the periodic table and is a highly reactive non-metallic substance. It is an oxidizing agent that forms oxides with multiple compounds readily. There are many known allotropes of oxygen, out of which O2 is the most stable one. It is also known as diatomic oxygen or molecular oxygen and has a significant presence in the atmosphere.
So, is O2 Polar or Non-Polar? The oxygen (O2) molecule is nonpolar because the molecule is diatomic and both atoms have equal electronegativity. As a result, both atoms share equal charges and there are no partial charges on any atom. Consequently, O2 comes out to be a nonpolar molecule with a zero dipole moment.
The basis of Polarity for any molecule
A polar bond is the one in which the centers of negative and positive charges do not coincide.
If two atoms are involved in the formation of a polar bond, then one of them is bound to be more electronegative than the other, to be able to draw some partial negative charge on itself.
As a result, some positive charge is left on the less electronegative atom, contributing to the existence of a dipole moment.
So, Polar molecules have:
- A dipole moment greater than 0D
- An electronegative difference greater than 0.4
- Accumulation of positive and negative charges at the poles of the molecule
- High melting and boiling points
- Lack of Symmetricity
- Good solubility in a polar solvent and insoluble/sparingly soluble in non-polar solvents
- Check out the article for the reason of the polarity of PCl3.
Non-polar molecules differ in properties from polar molecules in the following manner:
- They have a net-zero dipole moment
- The electronegative difference of atoms lies between 0-0.4 (negligible)
- No charge build-up at the poles
- Lower melting and boiling points
- Symmetric molecular structure
- Check out the article for the reason for non-polarity of CS2.
So, while concluding the polarity of any molecule or element, the above properties can be compared for a concrete inference.
Why is O2 nonpolar?
The O2 molecule considered a nonpolar molecule due to the below parameters. Let us study the below important points to check the polarity of a compound.
Electronegativity Difference: The electronegativity of a single oxygen atom (O) is 3.44.
When two of these atoms organize themselves with a double bond between them, it forms the structure of molecular O2 (diatomic oxygen). Since both the atoms are the same, the difference in their electronegativity is 0.
Dipole Moment of O2: Both the atoms forming O2 are the same and hence have an equal and opposite effect on each other.
The magnitude of the pull exerted on the shared electrons is equal from both sides, thereby leading to a net-zero force.
So, no charge build-up occurs at any pole and the net dipole moment of the molecular oxygen remains 0 Debye.
Symmetrical shape: The O2 molecule is linear in shape due to the diatomic molecule. Two atoms form double bonds to complete their octet and form a linear geometrical structure.
Below is the image of the geometrical structure of the Oxygen gas molecule
Nature of Bond in O2: When two oxygen atoms come together to form a bond, they end up forming a double bond that is covalent in nature i.e. there is equal sharing of valence electrons.
The electron configuration of Oxygen is [He] 2s² 2p⁴. It is clear that there are six valence electrons in the second shell of the atom, which is available for bonding.
When two oxygen atoms form a bond, they share one pair of electron each (i.e. two electrons are shared by each atom), thereby forming a covalent double bond.
Since there is equal sharing of electrons with no partial charges occurring at any of the atoms, therefore the net charge on the atoms in the molecule is zero each.
Hence, by virtue of covalent bond and zero permanent dipole moment, the diatomic oxygen molecule (O2) is non-polar in nature.
Lewis structure of O2
While determining the Lewis structure of diatomic oxygen molecule, two possibilities arise- one with a single bond between the two oxygen atoms and another with a double bond. However, only one out of these two structures is stable enough to exist.
- Although both structures (a) and (b) have complete octets for both oxygen atoms, still the two arrangements differ majorly when it comes to stability.
- In structure (a), there is an unpaired electron, called radical on both the oxygen atoms. This radical is extremely chemically reactive and triggers instability for the entire structure.
- Conversely, in structure (b), there is no presence of any radical. The formation of a double bond keeps the octet complete and also stabilizes the structure by including the individual radicals in structure (a) into the bond.
For any element, the structure should not only focus on the completion of the octet, but also on stability.
Hence, the preferred structure in the case of elemental oxygen O2 is the structure (b), with a double bond between the two oxygen atoms.
The Most Stable Allotrope of Oxygen – O2
O2 is the most stable form of the existence of oxygen. It is colorless in its gaseous form and imparts a pale blue color in liquid and solid forms of existence.
The formation of O2 is mainly attributed to the process of photosynthesis which can be described by the following equation:
6CO2(Carbon-dioxide) + 6H2O(Water) ——photons(Sunlight)—-> C6H12O6 (Glucose) + 6O2 (Dioxygen)
By mass, oxygen is the most abundant chemical element on the Earth’s surface, and the third most abundant element in the entire universe (after Hydrogen and Helium).
It makes up about 49.2% of the Earth’s crust (by mass) and 0.9% of the Sun’s mass.
The major presence of oxygen is in the form of oxide compounds such as silicon dioxide (SiO2).
It is also present in dissolved form in water and as a gas in the atmosphere. 88.8% of the mass of the oceans across the world is comprised of oxygen, while it takes up 20.8% of the volume of the atmosphere.
Properties of O2
- Molecular weight = 16g
- Density (at STP/gaseous form) = 1.43 g/L
- Density (in liquid form) = 1.14 g/cm3
- Melting point = -218.8 °C
- Boiling point = -183 °C
- At STP, Oxygen exists as a colorless gas, while in a liquid state it acquires a pale blue color.
- It is a highly reactive non-metal with the heat of fusion equal to 0.44 kJ/mol and heat of vaporization equal to 6.82 kJ/mol.
- Pure oxygen is 1.1 times heavier than air.
- Oxygen is the most prominent oxidizing agent with oxidation states -1, -2, and +2 (only in compounds with Fluorine).
- Oxygen dissolves readily in water (more readily in freshwater than seawater).
- Oxygen supports combustion.
Uses of O2
- Existence/Breathing: Living animals on land, in air and in water breathe O2 for survival. It is present in water in dissolved form for fish and other water animals.
- Medicine: Oxygen supplementation is used in medicine too. Many diseases such as pneumonia, heart disorders, emphysema, etc. are cured by Oxygen therapy.
- Life Support: As a low pressure breathing gas, it is deployed in space suits for astronauts. It is also consumed artificially by divers (Scuba and another underwater) via cylinders.
- Industrial Application: The process of smelting of iron ore for its conversion to steel consumes about half of the commercially produced oxygen. It is also used in welding and water treatment plants.
A number of reactions are carried out in the chemical industry for the formation of oxides, polymers, etc. with the help of oxygen.
- Formation of Ozone (O3): Ozone is a gaseous compound that occurs in the earth’s atmosphere and protects it from harmful UV rays of the sun. An endothermic reaction on three moles of O2 can lead to the formation of two moles of O3.3O2 (Oxygen) ——–> 2O3 (Ozone)
Oxygen forms a diatomic molecule as O2. Both atoms form. The electronegativity of both atoms is equal due to which both atoms have equal influence on the charges. The dipole moment of the O2 molecule turns out to be zero depicting it as a nonpolar molecule.
So friends, if you have any questions regarding polarity or non-polarity. feel free to reach out to us via comments.