Ozone (O₃) is one of the strongest oxidizers found in nature, with a reduction potential of 2.07 volts in acidic solution. In chemical reactions, it breaks apart carbon-carbon double and triple bonds, destroys pathogens in water, degrades rubber and plastics, and plays central roles in both atmospheric protection and ground-level pollution. What ozone does depends entirely on where it reacts and what it meets.
Ozone as an Oxidizer
Ozone’s three oxygen atoms make it inherently unstable. It constantly wants to shed that extra oxygen, and in doing so it rips electrons from whatever molecule is nearby. This is what makes it such a powerful oxidizing agent. For comparison, fluorine is the only common substance with a higher reduction potential. That reactivity means ozone can break bonds that milder oxidizers leave untouched, which is why it shows up in everything from organic synthesis to municipal water treatment.
Breaking Double Bonds: Ozonolysis
The most well-known reaction involving ozone in organic chemistry is ozonolysis, where O₃ cleaves carbon-carbon double bonds. The mechanism follows three distinct steps, first described by Rudolf Criegee.
In the first step, ozone adds across the double bond in a cycloaddition, forming a five-membered ring called the primary ozonide. This intermediate is highly unstable and immediately falls apart into two fragments: a regular carbonyl compound (an aldehyde or ketone) and a reactive species called a carbonyl oxide, sometimes referred to as a Criegee intermediate. These two fragments then recombine in a reversed orientation to form a more stable secondary ozonide.
What happens next depends on how the chemist treats the secondary ozonide. A reductive workup (often using dimethyl sulfide or zinc) yields aldehydes and ketones. An oxidative workup pushes the products further to carboxylic acids. This makes ozonolysis extremely useful for figuring out where double bonds sit in an unknown molecule: you break them, identify the fragments, and work backward.
Ozone also reacts with triple bonds, though much more slowly. The rate constants for alkynes are roughly in the hundreds per molar per second, compared to much faster rates for alkenes, and the activation energies are higher (37 to 48 kJ/mol). When ozone attacks a triple bond, the products typically include carboxylic acids. In one well-studied example, ozone attacking an ethynyl group produced a carboxyl product at 54% yield along with an aldehyde product and hydrogen peroxide released at about 21% yield.
Ozone in the Stratosphere
About 15 to 35 kilometers above the Earth’s surface, ozone absorbs ultraviolet radiation and shields life below. But certain pollutants catalytically destroy it. A single chlorine atom, released from refrigerants and industrial chemicals, reacts with ozone to form chlorine monoxide and ordinary oxygen. That chlorine monoxide then reacts with a free oxygen atom, regenerating the original chlorine atom and producing another molecule of O₂. The net result: one ozone molecule and one oxygen atom become two O₂ molecules, and the chlorine is free to repeat the cycle thousands of times.
Bromine compounds participate in similar cycles, sometimes teaming up with chlorine. When chlorine monoxide reacts with bromine monoxide, the net effect destroys two ozone molecules and produces three O₂ molecules. These catalytic cycles are the chemical engine behind the ozone hole over Antarctica.
Ozone at Ground Level
Near the surface, ozone forms through a completely different process and acts as a pollutant rather than a protector. Volatile organic compounds (VOCs) from vehicle exhaust, industrial emissions, and even vegetation react with nitrogen oxides (NOₓ) in the presence of sunlight. The chemistry is nonlinear, involving two interlocking radical cycles. Hydroxyl, hydroperoxyl, and organic peroxyl radicals (collectively called ROₓ) oxidize nitric oxide to nitrogen dioxide. Sunlight then splits the nitrogen dioxide, releasing an oxygen atom that combines with O₂ to form ozone.
This is why ozone levels spike on hot, sunny afternoons in cities with heavy traffic. The relationship between NOₓ and VOCs is not straightforward, though. Reducing nitrogen oxide emissions alone can sometimes temporarily increase ozone if VOC levels remain high, because NOₓ also scavenges the very radicals that drive ozone production. Effective ozone reduction requires cutting both precursors.
Water Disinfection
Ozone is widely used to purify drinking water because it kills bacteria, viruses, and parasites faster than chlorine and leaves no lasting chemical taste. Its effectiveness is measured using CT values, which multiply the ozone concentration (in mg/L) by the contact time (in minutes). The EPA considers a CT value of about 0.4 mg·min/L at 20°C sufficient for a 99.9% (3-log) reduction in viruses. Giardia cysts, which are tougher, require a CT of roughly 0.72 at the same temperature, rising to 1.43 mg·min/L at colder water temperatures around 10°C.
Because ozone decomposes back into oxygen within minutes, it doesn’t persist in the water supply the way chlorine does. This is both an advantage (no chemical residue) and a limitation (no ongoing protection in the distribution pipes).
Degradation of Rubber and Polymers
Ozone attacks carbon-carbon double bonds in polymers the same way it attacks them in small organic molecules. Natural rubber is especially vulnerable because its backbone is rich in these double bonds. Even trace atmospheric ozone, well below levels you can smell, causes “ozone cracking”: small surface cracks that propagate through the rubber under mechanical stress.
Research on elastomer blends shows that ozone cracks travel through the ozone-reactive rubber phase and skip across ozone-resistant particles (like ethylene-propylene rubber) without cutting through them. This is why tire and hose manufacturers blend ozone-resistant polymers into their rubber compounds and add chemical protectants called antiozonants that intercept ozone before it reaches the polymer chains.
Ozone on Metal Surfaces
When ozone contacts transition metals, it can oxidize them to unusually high oxidation states, creating surface metal oxides that serve as active sites for further reactions. This property is exploited in catalytic systems designed to break down air pollutants. In one approach, platinum particles supported on manganese and cerium oxides decompose ozone at the catalyst surface, generating reactive oxygen species that then oxidize volatile organic compounds like toluene. At just 30°C, optimized catalysts achieved roughly 97% toluene conversion and nearly 100% ozone decomposition. Without the catalyst present, ozone barely reacted with toluene at those temperatures, highlighting how metal surfaces unlock ozone’s oxidizing potential in controlled ways.
Workplace Exposure Limits
Because ozone is so reactive, even small concentrations irritate the lungs. OSHA sets the permissible exposure limit at 0.1 ppm averaged over an eight-hour work shift. You can typically smell ozone at concentrations around 0.01 to 0.02 ppm, which provides some warning, but chronic low-level exposure below the smell threshold can still cause respiratory inflammation over time. Workers near industrial ozone generators, UV sterilization equipment, and high-voltage electrical systems face the highest risk.