Ozone: An Allotrope of Oxygen

Ozone (O₃) is a reactive form of oxygen found mainly in the atmosphere around 20 km above sea level. It forms from oxygen when exposed to sunlight and is crucial in protecting Earth by absorbing harmful ultraviolet (UV) radiation.

1.0Structure of Ozone

Ozone (O₃) is a polar molecule with a bent geometry. It consists of two resonating structures, where the central oxygen atom has a formal charge of +1, and the terminal oxygen atoms each have a formal charge of -1. The bent geometry of ozone, confirmed by microwave spectroscopy, results in a bond angle of 116.7° and an O–O the bond length of 127.2 pm.

The central oxygen atom in ozone is sp² hybridised and contains one lone pair of electrons.

Due to its geometry and charge distribution, ozone has a dipole moment of 0.53 D and exhibits C₂v symmetry.

2.0Preparation of Ozone

Ozone can be produced by passing dry oxygen through a silent electrical discharge, creating ozonised oxygen (10%O₃).

• 3O₂ → 2O₃        (ΔH (298 K) = +142 kJ mol⁻¹)

Laboratory Preparation of Ozone 

Ozone (O₃) can be prepared in the laboratory by passing electrical discharges through dry oxygen using an apparatus known as an ozoniser. The most commonly known ozonisers are:

• Siemens Ozoniser: This device passes a silent electrical discharge through dry oxygen. Electric discharge splits oxygen molecules (O₂) into individual oxygen atoms (O), which then react with other oxygen molecules to form ozone (O₃).
• Brodie's Ozoniser: This ozoniser operates similarly, utilising a high-voltage electrical discharge to convert oxygen into ozone.

The general reaction for the formation of ozone in an ozoniser can be represented as 

                      3O₂   ⟶   2O₃ 

In both Siemens and Brodie's ozonisers, the process involves the following steps:

• Dry Oxygen Supply: Dry oxygen gas is fed into the ozoniser.
• Electric Discharge: A high-voltage electric discharge is applied to the oxygen gas, causing the oxygen molecules to dissociate into oxygen atoms.
• Ozone Formation: The free oxygen atoms combine with oxygen molecules to form ozone.

This method efficiently generates ozone, which can be collected and used for various applications such as disinfection, water treatment, and chemical synthesis.

3.0General Properties of Ozone

• Pale blue gas, dark blue liquid, violet-black solid.
• Characteristic smell: harmless in small amounts but harmful above 100 ppm, causing headaches and nausea.
• It is thermodynamically unstable and decomposes to oxygen (O₂) with heat release (negative ΔH) and increased entropy (positive ΔS), making it potentially explosive at high concentrations.
• Releases nascent oxygen (O₃ → O₂ + O), oxidising substances like lead sulfide to lead sulfate and iodide ions to iodine.
• • PbS + 4O₃ → PbSO₄  + 4O₂
  • 2I⁻ + H₂O + O₃→ 2OH⁻ + I₂ + O₂

4.0Chemical Properties of Ozone

Decomposition:

When heated, ozone decomposes into oxygen. This decomposition is completed at 300°C with catalysts such as black manganese dioxide.

•   2O₃→3O₂ Oxidizing Properties

Ozone is a strong oxidising agent. As it decomposes, it generates nascent oxygen, making it much stronger than dioxygen.

•   O₃→O₂+O

Oxidation Reactions:

Oxidation of Metals Ozone reacts directly with metals like silver (Ag), mercury (Hg), and copper (Cu) to form their respective oxides.

•   2Ag+O₃→Ag₂O+O₂
•   2Cu+O₃→Cu₂O+O₂

Oxidation of Metalloids Metalloids such as arsenic (As) and antimony (Sb) react with ozone to produce their respective oxyacids.

• Antimony: 2Sb + 5O₃ + 3H₂O → 2H₃SbO₄  +5O₂
• Arsenic: 2As+5O₃+3H₂O→2H₃AsO₄+5O₂

Oxidation of Non-Metals Non-metals like iodine (I), sulfur (S), and phosphorus (P) react with ozone in the presence of water to form their respective oxyacids.

• Iodine: I₂ + 5O₃ + H₂O → 2HIO₃ + 5O₂
• Sulfur: S + 5O₃ + H₂O → 2H₂SO₄+ 3O₂
• Phosphorus: P+5O₃+3H₂O→2H₃PO₄ + 5O₂

Oxidation of Compounds:

• Lead Sulfide: PbS + 4O₃  → PbSO₄+4O₂
• Iodides: 2KI + H₂O + O₃ → 2KOH+I₂+O₂
• Nitrites: KNO₂ + O₃ → KNO₃+O₂
• Potassium Manganate: 2K₂MnO₄(green) + O₃ + H₂O → 2KMnO,(pink)+2KOH+O₂

5.0Uses of Ozone

• Germicide and Disinfectant: Ozone is widely used as an antiseptic and disinfectant for sterilising water.
• UV Protection: It absorbs strong UV rays, protecting us from harmful UV radiation from the sun.
• Water Purification: Ozone is used for purification in water treatment plants, eliminating the need for filtration systems.
• Common Equipment: Ozone can be formed using equipment like photocopiers, laser printers, and other electrical devices.
• Antimicrobial Effects: Ozone can minimise the effects of bacteria, viruses, fungi, yeast, and protozoa.
• Ozone Therapy: Used to disinfect and treat diseases and hair-loss treatments.

6.0Reactivity and Environmental Impact

Ozone (O₃) is a molecule consisting of three oxygen atoms. It forms when heat and sunlight trigger chemical reactions between nitrogen oxides (NOₓ) and volatile organic compounds (VOCs), also known as hydrocarbons.

These reactions can occur near the ground, contributing to smog and high in the atmosphere, where ozone forms the protective ozone layer.

Stratospheric ozone, known as "good" ozone, forms about 10-30 miles above the Earth's surface and creates the ozone layer. This protective layer shields us from excessive ultraviolet (UV) radiation from the sun.

Conversely, ground-level ozone, or "bad" ozone, harms human health and the environment. It forms close to the ground primarily due to:

• NOₓ and VOCs: Emissions from mobile sources and industrial processes
• UV Radiation: Sunlight

Nitrogen Oxides: Rapidly react with ozone, potentially depleting the ozone layer.

• NO + O₃ → NO₂ + O₂

Freons: Used in aerosol sprays and refrigerants, posing a threat to the ozone layer.

Individuals with respiratory conditions such as asthma or those active outdoors on days with elevated ozone levels may experience shortness of breath, wheezing, and coughing. To safeguard our health and mitigate the harmful effects of this pollutant, it is crucial to take steps to limit both exposure and emissions.

7.0Ozone Layer Depletion

• Chlorofluorocarbons (CFCs): The primary source of ozone layer depletion. UV radiation in the stratosphere breaks down CFC molecules, releasing chlorine atoms that react with and degrade the ozone layer.
• Rocket Emissions: Unregulated rocket launches deplete the ozone layer more than CFCs. If unaddressed, significant depletion may occur by 2050.
• Nitrogen Oxides: Emissions from supersonic jet aeroplanes (NO₂, NO, N₂O) might slowly deplete the ozone concentration in the upper atmosphere.
• Freons: Used in aerosol sprays and as refrigerants, freons pose a threat to the ozone layer.
• Natural Phenomena: Stratospheric winds and sunspots contribute to ozone layer degradation, but they only cause a 1-2 percent reduction.

8.0Effects of Ozone Layer Depletion

• Increased exposure to UV radiation
• Health problems: skin disorders, cancer, weakened immune system, sunburns, cataracts, and accelerated ageing
• Epidermal damage in whales
• Increased solar damage in aquatic species
• Reduced growth, blooming, and photosynthesis in plants
• Forests affected by harmful UV radiation
• Significant impact on plankton life
• Disruption of the food chain if plankton are harmed.