O3 is the chemical formula for ozone. Chemically, it is a trioxygen gas molecule having a molar mass of 48 g/mol. The ozone gas is pale and bluish in color.
So, if you are curious whether a homoatomic ozone molecule is polar or non-polar, then this article is for you.
Is Ozone (O3) polar or non-polar?
Ozone (O3) is a polar molecule. O3 consists of three identical oxygen (O) atoms. One oxygen atom is present at the center of the molecule, while the other two oxygen (O) atoms occupy terminal positions, one on each side. The central O-atom has a lone pair of electrons; thus, O3 occupies a bent, angular, or V-shape.
Each oxygen atom in O3 has the same electronegativity value, i.e., 3.44 units. Therefore, no electronegativity difference exists between the bonded atoms in the O=O and O-O bonds present in the O3 molecule. So O3 should technically be non-polar, right? But that is not the case.
The polarity of O3 is accredited to its asymmetrical bent shape. The central O-atom carries a +1 formal charge, while the single-bonded O-atom carries a -1 formal charge in O3. The O-atom carrying +1 charge is slightly electron deficient, while the O-atom carrying -1 formal charge is slightly electron rich.
The electron-deficient O-atom attracts O-O bonded electrons to a greater extent. Consequently, it is due to the presence of these formal charges and bent shape that O3 contains a non-uniform charge distribution overall. It is thus slightly polar (net µ = 0.53 D).
A molecule is polar if there is a non-uniform charge distribution present in it. If the charge distribution is equally balanced in different parts of the molecule, then it is considered non-polar.
The following three factors mainly influence the polarity of a molecule:
The electronegativity difference between two or more covalently bonded atoms
Dipole moment
Molecular geometry or shape
Now let’s see to what extent the above three factors affect the polarity of an ozone (O3) molecule.
Factors affecting the polarity of ozone (O3)
Electronegativity
It is defined as the ability of an atom to attract a shared pair of electrons from a covalent chemical bond.
Electronegativity decreases down the group in the Periodic Table of elements while it increases across a period.
Greater the electronegativity difference between bonded atoms in a molecule, the higher the bond polarity.
Oxygen (O) is present in Group VI A (or 16) of the Periodic Table. The electronic configuration of an oxygen atom is 1s2 2s2 2p4. Valence electrons in oxygen are a total of 6 in number, so each O-atom has a deficiency of 2 more valence electrons for it to complete its octet.
In Ozone (O3), two O-atoms are covalently bonded to the central O-atom, one on each side of the bent, angular or V-shape.
One terminal O-atom is bonded via a double-covalent bond with the central O-atom, while the other terminal O-atom is single-covalently bonded to the central O-atom in O3.
The octet of one terminal O-atom is complete with O-O single bond and three lone pairs of electrons present on it. Similarly, the octet of the other terminal O-atom is complete with an O=O double bond and two lone pairs of electrons.
The central O-atom contains 1 lone pair of electrons. This lone pair distorts the geometry of the molecule; it thus adopts a bent shape.
Atom
Electronic configuration
Valence electrons
Oxygen (8O)
1s2 2s2 2p4
6
Both the O-atoms have the same electronegativity value, i.e., 3.44 units. Thus zero or no electronegativity difference exists between the bonded O-atoms in either of the O-O single bond or O=O double covalent bond present in O3. Thus the covalent bonds present in the O3 molecule are individually non-polar.
There are no partial positive and/or partial negative centers present in O3. As per ‘bond polarity’’ the O-O the electron cloud is equally shared between the bonded oxygen atoms in each of the O-O and O=O bonds in the ozone molecule (as shown below).
Dipole Moment
The dipole moment is the product of electrical charge (Q) and bond length (r) between two bonded atoms. It is a vector quantity expressed in Debye (D) units.
It is represented by a Greek symbol µ and measures the polarity of a bond.
The dipole moment of a polar covalent bond conventionally points from the positive center to the center of the negative charge.
But as O3 consists of non-polar covalent bonds and there are no partial positive or partial negative centers present in it thus, O-O and O=O does not contain a dipole moment value individually.
Molecular geometry
According to the valence shell electron pair repulsion (VSEPR) theory of chemical bonding, O3 is an AX2E1-type molecule. To one O-atom at the center (A), two other O-atoms are covalently bonded (X), and there is one lone pair of electrons (E) on the central oxygen atom.
So, the molecular geometry or shape of O3 is bent, also known as angular or V-shaped. The O=O-O bond angle is 116.8° in O3.
The presence of a lone pair of electrons on the central oxygen atom in O3 leads to strong lone pair-bond pair electronic repulsions in addition to a bond pair-bond pair repulsive effect.
This strong repulsive effect pushes the bonded atoms away from the lone pair at the center and decreases the bond angle to 116.8° from an ideal bond angle of 120° in a symmetrical AX3-type trigonal planar molecule.
It is due to the distortion present in O3 molecular geometry or shape that not only does the symmetry of the molecule gets disturbed, but it also leads to an asymmetric charge distribution over the molecule.
The presence of a +1 and -1 formal charge on the covalently bonded O-atoms further disturbs the electron cloud distribution. Thus, O3 possesses a net dipole moment of 0.53 D, and it is overall polar.
Difference between polar and nonpolar?
Polar molecule
Non-polar molecule
Atoms must have a difference in
electronegativity
Atoms may have the same or different electronegativity values
Unequal charge distribution overall
Equal charge distribution overall
Net dipole moment greater than zero
Net dipole moment equals to zero
Examples include water (H2O), ethanol (CH3CH2OH), ammonia (NH3), sulfur dioxide (SO2), bromine trifluoride (BrF3), ozone (O3), etc.
Examples include oxygen (O2), nitrogen (N2), methane (CH4), sulfur trioxide (SO3), etc.
Both the O-O single bond and O=O double bond present in O3 is individually non-polar as no electronegativity difference exists between two identical oxygen atoms.
However, the presence of formal charges and an asymmetric bent shape of O3 due to a lone pair of electrons present on it leads to a net dipole moment µ > 0.
The electron cloud stays non-uniformly distributed between the bonded atoms; thus, O3 is overall polar.
What is the difference between bond polarity and molecular polarity?
Bond polarity only depends on the electronegativity difference between two bonded atoms in a covalent chemical bond.
Molecular polarity depends on molecular geometry or shape in addition to the electronegativity difference between bonded atoms and dipole moment.
The net dipole moment of a polar molecule is greater than zero. The net dipole moment of a non-polar molecule is equal to zero.
Can a molecule be overall non-polar, having polar bonds in it?
Yes. If the dipole moments of individually polar bonds get canceled equally due to the symmetrical shape of a molecule, then the molecule will be overall non-polar.
For example, in carbon dioxide (CO2), a high electronegativity difference of 0.89 units exists between a carbon and an oxygen atom in each C=O bond. So both C=O bonds are polar.
However, it is due to the symmetrical linear shape of CO2 that C=O dipole moments get canceled equally, so CO2 is overall non-polar (net µ = 0).
Can a molecule be overall polar even though having non-polar bonds?
Yes. It’s a rare case, but examples are there, for instance, the ozone (O3) molecule. O3 has non-polar O-O and O=O bonds due to zero electronegativity difference between identical oxygen atoms.
However, the asymmetrical bent shape of O3 leads to a net µ > 0. Thus, it is overall polar.
What are the formal charges present on bonded atoms in O3?
Formal charge of an atom = [ valence electrons – non-bonding electrons- ½ (bonding electrons)]
+1 formal charge is present at the central O-atom, -1 formal charge is present at the single-bonded O-atom, and zero formal charges are present at the double-bonded O-atom in O3.
Thus, the net formal charge present on the ozone molecule is +1 + (-1) + (0) = 0.
Why O3 molecule is polar, but the O2 molecule is non-polar?
In O3, there is a lone pair of electrons present on the central O-atom. Electronic repulsions distort the shape and geometry of the molecule. An asymmetric bent shape ensures the polarity of O3 (net µ = 0.53 D).
Contrarily, O2 is a diatomic molecule with a symmetrical linear molecular geometry and an ideal bond angle of 180°. Both O-atoms have identical electronegativity, both share equal charges and there are no partial charges present on any atom. Therefore, the O2 molecule has zero dipole moment (net µ =0) and is non-polar.
Summary
Ozone O3 is a polar molecule.
It consists of one O-O single bond and one O=O double bond having no electronegativity difference between identical bonded atoms.
Ozone O3 has a bent, angular, or V-shape with a 116.8° bond angle.
The asymmetric bent shape ensures that the electron cloud is not uniformly spread over the 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/
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