Oxygen (O)

In aerobic organisms, O₂ is the terminal electron acceptor in the electron transport chain. Electrons from NADH/FADH₂ reduce O₂ to water, driving ATP synthesis.

\(\mathrm{C_6H_{12}O_6 + 6\,O_2 \rightarrow 6\,CO_2 + 6\,H_2O + \text{energy (ATP)}}\)

This high redox potential of O₂ makes energy capture efficient.

No. Combustion is a redox reaction where a fuel is oxidized by an oxidizer (often O₂). Oxygen usually acts as the oxidizer, not the fuel. However, oxygen can react with even more powerful oxidizers (e.g., fluorine), but under normal conditions we say O₂ does not "burn"—it enables burning.

Most common is −2 (e.g., H₂O, CO₂). Special cases include:

• −1 in peroxides (\(\mathrm{H_2O_2}\), \(\mathrm{Na_2O_2}\))
• −\tfrac{1}{2} in superoxides (\(\mathrm{KO_2}\))
• 0 in elemental forms (O₂, O₃)
• +1 in \(\mathrm{O_2F_2}\) and +2 in \(\mathrm{OF_2}\) because fluorine is more electronegative

Industrial: Fractional distillation of liquefied air after removing CO₂ and H₂O. Liquid O₂ (b.p. ~90 K) is separated from N₂ (b.p. ~77 K).

Laboratory: Decomposition of hydrogen peroxide using a catalyst (MnO₂):

\(\mathrm{2\,H_2O_2(aq) \xrightarrow{MnO_2} 2\,H_2O(l) + O_2(g)}\)

O₂ is a diatomic molecule essential for respiration; it is relatively stable. O₃ is a triatomic allotrope (bent structure) and a much stronger oxidant. In the stratosphere, ozone absorbs harmful UV radiation; at ground level, it can be a pollutant.

\(\mathrm{O_2 + O\cdot \rightarrow O_3}\)

Photosynthetic organisms release O₂ by splitting water:

\(\mathrm{6\,CO_2 + 6\,H_2O \xrightarrow{light,\;chlorophyll} C_6H_{12}O_6 + 6\,O_2}\)

Over geological time, this process oxygenated Earth’s atmosphere, enabling aerobic life.

Glowing splint test: A glowing wooden splint inserted into a sample re-lights in O₂ due to enhanced combustion.

Paramagnetism: Liquid oxygen is pale blue and strongly paramagnetic; O₂ is attracted into a magnetic field because it has two unpaired electrons in antibonding \(\pi^*\) orbitals.

Hydrogen burns in oxygen to form water:

\(\mathrm{2\,H_2(g) + O_2(g) \rightarrow 2\,H_2O(l)}\)

The reaction is highly exothermic (\(\Delta H^\circ < 0\)), releasing about \(\mathrm{\approx -286\,kJ\,mol^{-1}}\) per mole of water formed (value depends on phase/conditions).

As oxygen forms bonds, it typically gains electrons (is reduced), causing other species to lose electrons (be oxidized). For example, in metal oxidation:

\(\mathrm{4\,Fe + 3\,O_2 \rightarrow 2\,Fe_2O_3}\)

Oxygen’s high electronegativity and ability to stabilize the \(\mathrm{O^{2-}}\) state drive these processes.

O₂ itself is not flammable, but it greatly accelerates combustion and lowers ignition temperatures of fuels. Safety tips:

• Keep oils/greases away from high-pressure O₂ systems (risk of spontaneous ignition).
• Avoid sources of sparks/flames; use compatible materials and regulators.
• Store cylinders secured and upright; never use in confined, poorly ventilated spaces without controls.