Ozone (O3) Lewis dot structure, molecular geometry or shape, electron geometry, bond angle, formal charge, hybridization
Ozone depletion is a growing environmental concern in the modern scientific world. It is a hot topic everywhere right? But what is ozone and what is its chemical nature?
Well, ozone is chemically trioxygen, a pale blue gas, naturally present in the stratosphere that protects the Earth from harmful UV-B radiation of the sun. It is represented by a homoatomic chemical formula i.e., O3. In this article, we will introduce you to the chemistry behind O3.
We have talked about the Lewis structure of O3, its molecular geometry or shape, electron geometry, bond angle, formal charge, hybridization, etc. In short, everything you need to know about ozone (O3) is summed up in this article. So, let’s begin reading!
|Name of Molecule||Ozone|
|Molecular geometry of O3||Bent, angular, or V-shaped|
|Electron geometry of O3||Trigonal planar|
|Bond angle (O=O-O)||116.8º|
|Total Valence electron in O3||18|
|Overall Formal charge in O3||Zero|
How to draw lewis structure of O3?
The Lewis structure of O3 consists of three identical oxygen (O) atoms. One O-atom acts as the central atom while the other two O-atoms are bonded to it as outer atoms. There are a total of 3 electron density regions around the central O atom. Out of these 3 regions, two are the bonded electron pairs while a lone pair is also present on the central oxygen.
Drawing the Lewis dot structure of O3 is a bit technical but quite an interesting task. To make things easier for you, we have subdivided this task using the following simple steps.
Steps for drawing the Lewis dot structure of O3
1. Count the total valence electrons in O3
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 find the total valence electrons present in it.
The valence electrons can be determined by identifying the concerned elemental atoms in the Periodic Table. In the O3 molecule, there is only 1 type of elemental atom involved i.e., oxygen (O), so we just need to look for oxygen in the Periodic Table of elements.
Oxygen is present in Group VI A of the Periodic Table which means it has a total of 6 valence electrons.
- Total number of valence electrons in oxygen = 6
∴ There are 3 O atoms in the ozone molecule thus total valence electrons available for drawing the Lewis structure of O3 are 3(6) = 18 valence electrons.
2. Choose the central atom
In this second step, usually the least electronegative atom out of all the concerned atoms is chosen and placed as the central atom. This is because the least electronegative atom is the one that is most likely to share its electrons with the atoms spread around it.
But in the case of O3, the situation seems uncomplicated because it involves three identical O atoms. Hence any one O atom can be chosen as the central atom while the other two atoms are placed in its surroundings, as shown in the figure below.
3. Connect outer atoms with the central atom
In this step, the outer atoms are joined to the central atom using single straight lines.
So, the outer O atoms are joined to the central O atom in the Lewis structure of O3 using straight lines.
Each straight line represents a single covalent bond i.e., a bonded electron pair containing 2 electrons. There are a total of 2 single bonds in the diagram shown above thus the total valence electrons used so far out of the 18 valence electrons available are 2(2) = 4 valence electrons.
- Total valence electrons available – electrons used till step 3 = 18-4 = 14 valence electrons.
- This means 14 valence electrons are still available to be accommodated in the Lewis dot structure of O3.
4. Complete the octet of outer atoms
In this step, we need to complete the octet of the two O-atoms bonded to the central oxygen atom. Each O-atom requires a total of 8 valence electrons to achieve a stable octet electronic configuration.
Both the outer O-atoms are bonded to the central oxygen atom using single bonds. That means each outer O-atom already has 2 valence electrons. It is short of 6 electrons that are required to complete its octet.
Therefore, 6 valence electrons are placed around both the outer O atoms as 3 lone pairs respectively. Refer to 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 outer O-atom, shown as dots) = 2(2) + 2(6) = 16 valence electrons.
- Total valence electrons – electrons used till step 4 = 18 – 16 = 2 valence electrons.
There are 2 valence electrons still available thus these 2 electrons are placed as a lone pair on the central O atom.
All 18 valence electrons are now used but the problem that stays is that there are a total of 6 valence electrons only around the central O atom in this structure which denotes an incomplete octet.
An easy solution to this problem is that we convert one electron pair from an outer O atom into a covalent bond. This results in the formation of a double covalent bond between two O atoms i.e., the central oxygen atom and an outer oxygen atom as shown in the figure below.
Now there is 1 single bond + 1 double bond around the central O-atom and a lone pair is present on it thus it has a complete octet.
The octet of the first oxygen atom is complete with 1 double bond and 2 lone pairs. Similarly, the octet of the second oxygen atom is also complete with 1 single bond and 3 lone pairs respectively. This means a favorable situation for all three bonded atoms.
The final step is to check the stability of the above Lewis structure which can be done with the help of 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 three oxygen atoms in the O3 Lewis structure can be determined using the formula given below.
- Formal charge = [ valence electrons – nonbonding electrons- ½ (bonding electrons)]
Now let’s count the formal charge present on O3 atoms using this formula and the Lewis structure obtained in the previous step.
For central oxygen atom
- Valence electrons of oxygen = 6
- Bonding electrons = 1 single bond + 1 double bond = 2 + 2(2) = 6 electrons
- Non-bonding electrons = 1 lone pair = 2 electrons
- Formal charge = 6 -2 – 6/2 = 6 – 2 – 3 = 6 – 5 = +1
For 1st oxygen atom
- Valence electrons of oxygen = 6
- Bonding electrons = 1 double bond = 2(2) = 4 electrons
- Non-bonding electrons = 2 lone pairs = 2(2) = 4 electrons
- Formal charge = 6 – 4 – 4/2 = 6 -4 – 2 = 6 – 6 =0
For the 2nd oxygen atom
- Valence electrons of oxygen = 6
- 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
A +1 formal charge is present on the central oxygen atom while a -1 formal charge is present on one of the other two oxygen atoms.
If we convert a lone pair present on the 2nd oxygen atom into a covalent bond between this oxygen and the central oxygen atom, then the +1 and the -1 charges can be neutralized.
But, this cannot be done because oxygen does not have an expanded octet. The central O atom in the O3 molecule does not have a d subshell so it can accommodate up to 8 valence electrons only.
Therefore, the Lewis structure obtained in step 5 is considered the correct Lewis structure of O3. The +1 and -1 formal charges cancel out each other marking the stability of this Lewis structure overall.
Net formal charge on the molecule = +1 + 0 -1 = 1-1 = 0
In reality, the true Lewis structure of ozone is a resonance hybrid of the following two resonance structures. Both resonance structures contribute equally to the resonance hybrid.
Each resonance structure is a way of representing the Lewis structure of a molecule. Nevertheless, the resonance hybrid is the best possible and the most stable Lewis representation of the molecule.
A lone pair from any one outer atom can transform into a bond pair at a moment. The electrons keep moving from one position to another on the molecule. This is called electronic delocalization.
Similarly, the charges present on O3 atoms are not stationary, rather they keep revolving from one point to another on the molecule, canceling each other’s effect overall.
Also check –
What are the electron and molecular geometry of O3?
The ideal electron geometry of ozone (O3) is trigonal planar. However, its molecular geometry or shape is bent or angular. The molecule adopts a different shape from its ideal electronic geometry due to the presence of a lone pair on the central oxygen atom.
Lone pair-bond pair repulsions distort the geometry of the molecule and make it occupy an asymmetric bent shape.
Molecular geometry of O3
The ozone (O3) molecule has a bent, angular, or V-shape. The central O atom is situated at the tip of the inverted V while the other two O atoms occupy terminal positions. Due to the presence of a lone pair on the central O atom, lone pair-bond pair repulsions exist in the molecule in addition to bond pair-bond pair electronic repulsions.
The terminal O atoms are maximally pushed away from the lone pair at the center and the molecule obtains an asymmetric bent shape.
An important point to consider is that the molecular geometry or shape of a molecule depends on the distinction between lone pairs and bond pairs around the central atom. Contrarily, the ideal electron geometry depends on the total electron density regions around the central atom in the molecule, whether it’s a bond pair or a lone pair.
Therefore, we seriously take into account the repulsive effect of a lone pair on the central O atom while determining the shape of the ozone molecule. But this repulsive effect is ignored while considering the electron geometry of O3. Let’s see how.
Electron geometry of O3
According to the valence shell electron pair repulsion (VSEPR) theory of chemical bonding, a molecule containing 3 regions of electron density around the central atom has a trigonal planar electron geometry.
The double bond between two bonded O atoms in O3 is considered 1 region of electron density. In this way, there are a total of 3 electron density regions (1 double bond + 1 single bond + 1 lone pair) around the central O atom in O3 thus its ideal electron geometry is trigonal planar.
You can also find the electron and molecular geometry or shape of a molecule such as O3 using the AXN method.
AXN is a simple formula that represents the number of bonded atoms and lone pairs around the central atom of a molecule. It is used to predict the shape and geometry of the molecule using the VSEPR concept.
AXN notation for the O3 molecule:
- A in the AXN formula represents the central atom. In the O3 molecule, the oxygen (O) atom is present at the center so A=O.
- X denotes the number of atoms bonded to the central atom. In O3, two O atoms are bonded to the central oxygen atom thus X=2.
- N represents the number of lone pairs present on the central atom in the molecule. In O3, 1 lone pair is present on the central oxygen atom so N=1.
In this way, the AXN generic formula for the O3 molecule is AX2N1.
Now you can use this formula against the VSEPR chart given below to identify the electron and molecular geometry assigned next to it.
According to the VSEPR chart given above, a molecule with an AX2N1 generic formula has a bent molecular geometry or shape while its electron geometry is trigonal planar, as we already identified for the O3 molecule.
Hybridization of O3
The ozone (O3) molecule has sp2 hybridization.
The electronic configuration of oxygen (O) is 1s2 2s2 2p4.
During chemical bond formation, the 2s atomic orbital of the central oxygen atom hybridizes with two half-filled 2p atomic orbitals to yield three sp2 hybrid orbitals. Each hybrid orbital has a 33.3% s character and a 66.7% p character.
One of these three sp2 hybrid orbitals contains paired electrons which are situated as a lone pair on the central O atom in the O3 molecule. The other two sp2 hybrid orbitals contain 1 electron each.
Each sp2 hybrid orbital of the central oxygen atom overlap with the sp2 hybrid orbitals of terminal O atoms to form sigma (σ) O-O bonds on each side of O3. The unhybridized p orbital of the central O atom forms a pi (π) bond with the unhybridized p orbital of a terminal O atom.
A shortcut for memorizing the hybridization present in a molecule is to use the steric number of the molecule against the table given below. The steric number of the central O atom in the O3 molecule is 3 as it has a total of three electron density regions around it. The table shows it has sp2 hybridization.
The O3 bond angles
The ideal bond angle in a trigonal planar molecule is 120°. But it is due to a lone pair present on the central oxygen (O) atom that lone pair-bond pair repulsions decrease the bond angle by tilting the O=O and O-O bonds towards each other.
Thus, the O=O-O bond angle becomes 116.8° in the asymmetrical bent shape of the ozone (O3) molecule. The average O=O bond length is 128 pm.
Also check:- How to determine bond angle?
Is O3 polar or nonpolar?
The polarity of the ozone (O3) molecule is a tricky concept. The ozone molecule is made up of three identical O atoms. Each O atom has the same electronegativity value i.e., 3.44 units. Therefore, no electronegativity difference exists between the bonded O atoms in the O3 molecule.
According to Pauling’s electronegativity scale, the bonded atoms must have an electronegativity difference between 0.5 and 1.6 units for a bond to be considered polar. So, technically O3 should be a non-polar molecule, right? But that is not the case.
This is because of the asymmetrical bent shape of the O3 molecule. Formal charges are present on the bonded O-atoms in the O3 molecule. The central O atom is slightly positively charged while a terminal O atom is slightly negatively charged.
Due to this unequal charge distribution, the electron cloud stays non-uniformly distributed in the molecule overall. Thus, ozone possesses a net dipole moment value (μ=0.53 D) and it is a polar molecule.
Read in details –
How many bond pairs and lone pairs are present in the O3 Lewis structure?
What is the electron pair geometry of O3?
|The ideal electron pair geometry of the O3 molecule is trigonal planar as there are 3 regions of electron density around the central O atom in O3 i.e., 1 double bond + 1 single bond + 1 lone pair.|
Why does ozone have a bent shape?
|Lone pair-bond pair repulsions > bond pair-bond pair repulsions. The presence of a lone pair on the central oxygen atom in the ozone molecule pushes the bonded O atoms away from the central position therefore the molecule Ozone (O3) adopts a bent shape.|
Why is the bond angle in O3 less than 120°?
A lone pair of electrons present on the central oxygen atom in the O3 molecule distorts the shape and geometry of the molecule.
Lone pair-bond pair repulsions decrease the bond angle. The bond angle decreases from an ideal 120° to 116.8° in the asymmetrical bent shape of O3.
Why is the ozone angle greater than the water angle?
The O=O-O bond angle in ozone is 116.8° while the H-O-H bond angle in water is 104.5°. Both the molecules are triatomic and have a bent shape but ozone angle > water angle.
This is because in H2O the central oxygen atom has 2 lone pairs of electrons. Lone pair -lone pair repulsions exist in the molecule in addition to lone pair-bond pair repulsions which push the O-H bonds further away from the central position and they tilt inwards to a greater extent thus the bond angle decreases.
The central oxygen atom in O3 has only 1 lone pair therefore the repulsive effect is comparatively weaker than that in H2O and the bond angle is consequently higher.
Let’s read in detail about – the H2O lewis structure, molecular geometry or shape, and bond angle
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- The total valence electrons available for drawing the O3 Lewis structure are 18.
- Ozone (O3) molecule has a bent or V-shape and molecular geometry.
- The ideal electron geometry of O3 is trigonal planar.
- The Lewis dot structure of O3 has two equivalent resonance forms.
- There are +1 and -1 formal charges present on two of the three bonded oxygen atoms in the O3 Lewis structure.
- The +1 formal charge gets canceled with the -1 charge to mark the stability of the overall Lewis structure of O3.
- O3 has sp2
- The ideal bond angle in a trigonal planar molecule is 120° but it is due to a lone pair present on the central oxygen atom in the O3 molecule that the O=O-O bond angle decreases and becomes 116.8°.
- The presence of formal charges leads to an unequal charge distribution in the molecule overall and consequently a non-uniformly distributed electron cloud. Thus, the O3 molecule is polar in nature (net μ= 0.53 D).