Environmental Chemistry


Ozone or trioxygen is an inorganic molecule and has the chemical formula O3. Ozone is a pale blue gas. It has a distinctively pungent smell. Ozone is an allotrope of oxygen. It is very less stable than the diatomic allotrope \(O_2\), breaking down in the lower atmosphere to \(O_2\) i.e., dioxygen.

Ozone is made from dioxygen by the action of ultraviolet (UV) light and electrical discharges within the Earth’s atmosphere. Ozone is present in very low concentrations throughout the latter. It has its highest concentration in the ozone layer of the stratosphere that absorbs most of the Sun’s ultraviolet (UV) radiation.



Introduction to Ozone

The odour of ozone is reminiscent of chlorine and detectable by many people at concentrations of as little as 0.1 ppm in air. The structure of ozone \(O_3\) came up in 1865. The molecule of ozone was proven to have a bent structure and to be weakly paramagnetic. Ozone is a pale blue gas that easily condenses at cryogenic temperatures to a dark blue liquid and at the end into a violet-black solid. The instability of ozone with regards to dioxygen such that both concentrated gas and liquid ozone may decompose explosively at elevated temperatures. Therefore, is commercially in use at low concentrations only.

Ozone is a strong oxidant, far more than dioxygen. Ozone has its demand in many industrial and consumer applications related to oxidation. This high oxidizing potential causes ozone to damage mucous and respiratory tissues in animals, and also in plants at concentrations of about 0.1 ppm. Ozone makes a potent respiratory hazard and pollutant near ground level. A higher concentration in the ozone layer (from 2 to 8 ppm) is beneficial, it prevents the damaging of ultraviolet (UV) light from reaching the Earth’s surface. The density of ozone is \(2.14 kg / m^3\). The molecular weight or molar mass of ozone is 48 g/mol.

Structure of Ozone \(O_3\)

Ozone is a bent molecule according to experimental evidence from microwave spectroscopy. It has \(C_{2v}\) symmetry that is similar to the water molecule. The O – O distances in the ozone molecule are 127.2 pm or 1.272 angstroms. The O – O – O angle in the ozone id \(116.78^o\). The central atom of the ozone molecule is \(sp^2\) hybridized and has one lone pair.

It is a polar molecule and has a dipole moment of 0.53 D. The molecule of ozone can be represented as a resonance hybrid with two contributing structures. Each structure with a single bond on one side and a double bond on the other. The arrangement of ozone possesses an overall bond order of 1.5 for both sides. Ozone is isoelectronic with the nitrite anion. Ozone is composed of substituted isotopes like \(^{16}O, ^{17}O, ^{18}O\).

Physical Properties of Ozone \(O_3\)

1)Ozone is a colourless or pale blue gas.

2)Ozone is slightly soluble in water and much more soluble in inert non-polar solvents like sulphuric acid, carbon tetrachloride or fluorocarbons that forms a blue solution.

3)The boiling and the melting point of ozone is \(-112^oC\) and \(-193.2^oC\) respectively.

4)At \(161 K (−112^oC; −170^oF)\), ozone condenses to form a dark blue liquid. It is risky to allow this liquid to warm to its boiling point. It is because both concentrated gaseous ozone and liquid ozone can get detonate. At temperatures below \(80 K (−193.2^oC\); \(−315.7^oF\), it forms a violet-black solid.

5)Most people can detect about \(0.01 \mu mol\)/mol of ozone in the air where it has a very specific sharp odour some what resembling chlorine bleach.

6)Exposure of \(0.1 to 1 \mu mol\)/mol of ozone causes headaches, burning eyes and irritation to the respiratory passages. At low concentrations in the air, ozone is very destructive to organic materials like latex, plastics and animal lung tissue.

Chemical Properties of ozone

1)Ozone easily dissolves in water results in the formation of hydrogen peroxide. The chemical equation for the same is

\(O_3 + 3H_2O \rightarrow  3H_2O_2\)

2)Ozone reacts with lead sulfide that results in the formation of lead sulfate. The chemical equation for the same is

\(3PbS + 4O_3 \rightarrow  3PbSO_4\)

Uses of Ozone

  1. Ozone in use at water treatment plants without filtration systems.
  2. Ozone can also be formed by commonly used equipment such as photocopiers, laser printers, and other electrical devices.
  3. In medicine, by limiting the effects of bacteria, viruses, fungi, yeast, and protozoa, ozone therapy is in use to disinfect and treat diseases.
  4. Many ozone-depleting compounds have properties that make them good refrigerants i.e., they can efficiently transfer heat from one location to another.

FAQs on Ozone

Question 1: Is ozone toxic to breathe?
Answer: Ozone is toxic as ozone can cause damage to the lungs when inhaled. Relatively at low ozone levels, ozone can cause pain in the throat, coughing, shortness of breath and inflammation in the lungs.

Question 2:  Is ozone a resonance structure?
Answer: Ozone has two resonance structures that equally contribute to the molecule’s overall hybrid structure. The two arrangements are similar to a point of equilibrium. Each arrangement is having a positive and negative formal charge on two of the oxygen atoms.

Question 3:  How many lone pairs are in Ozone?
Answer: One single pair is present in the ozone. As ozone has 18 valence electrons out of which there is only one pair of electrons.

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