The oxygen people breathe almost always appears as a molecule composed of two atoms bonded together, a form known as dioxygen (O2). This common molecular structure leads many to believe that oxygen cannot exist as an individual, unbound atom. However, oxygen can exist as a single atom. This solitary form, known as atomic oxygen (O), is chemically different from the common gas and is highly unstable in the dense atmosphere near the Earth’s surface, requiring specialized, high-energy conditions to be maintained.
The Stability of Molecular Oxygen
The reason the molecular form (O2) is so prevalent is rooted in the rules of chemical stability. An oxygen atom has six valence electrons in its outermost shell. To achieve a stable configuration, it needs a total of eight valence electrons.
Two oxygen atoms satisfy this requirement by sharing two pairs of electrons between them, which creates a double covalent bond. This sharing allows both atoms to effectively count eight electrons in their outer shell, fulfilling the stability rule. The resulting O2 molecule represents the lowest possible energy state for two oxygen atoms under normal atmospheric conditions.
This low-energy, double-bonded structure makes molecular oxygen relatively unreactive at room temperature, ensuring its abundance in the lower atmosphere. Breaking this strong double bond requires a significant input of energy, measured at approximately 117.2 kilocalories per mole. This large energy barrier protects the molecular form from spontaneously breaking down into its single-atom components.
Defining Atomic Oxygen
Atomic oxygen (O) is defined as a single, unbound atom of the element and is chemically distinct from the molecular gas. With only six valence electrons, a single oxygen atom is in a high-energy, highly reactive state. This configuration makes the atom an extremely powerful oxidizing agent.
A single oxygen atom rapidly reacts with nearly any other available element or compound it encounters to achieve a stable bond. This extreme reactivity means the free atom exists for only a fraction of a second in the dense air near the ground before bonding with something else. For instance, atomic oxygen readily bonds with a molecule of O2 to form ozone (O3).
Atomic oxygen is a short-lived transient species because of its high energy and tendency to quickly react. The energy required to create it, by breaking the double bond of O2, is retained by the single atom. This makes it an energetic particle that aggressively seeks a reaction partner. This difference in energy and reactivity separates the common O2 molecule from the chemically aggressive O atom.
Environments Where Single-Atom Oxygen Occurs
Since atomic oxygen is so reactive, its sustained presence requires environments that continuously create it and where reaction partners are sparse. The most significant natural habitat for O is the Earth’s upper atmosphere, specifically the thermosphere and the exosphere, starting at altitudes above 80 kilometers.
In the upper atmosphere, intense ultraviolet (UV) radiation from the sun constantly bombards molecular oxygen. This high-energy light causes a process called photodissociation. Photodissociation occurs when UV photons provide enough energy to break the double covalent bond of O2, splitting it into two separate O atoms. This process is the primary natural source of atomic oxygen.
The upper atmosphere has an extremely low density, meaning the particles are very far apart. The single oxygen atoms do not collide frequently enough with each other or with O2 molecules to rapidly reform the stable O2 or O3 molecules. This low collision rate allows the highly reactive atomic oxygen to persist for long periods, making it the most abundant atmospheric constituent above an altitude of about 200 kilometers.
Atomic oxygen is also intentionally created in specialized high-vacuum laboratory settings or industrial processes, such as plasma etching used in microchip manufacturing. These artificial environments use high energy inputs to generate the single atoms and then maintain a vacuum to prevent them from immediately reacting. This mimics the low-density conditions of the upper atmosphere and allows scientists and engineers to harness the atom’s high reactivity for specific material modification tasks.