Oxidation and Redox
Redox Chemistry
Redox chemistry is about the transfer of electrons. The word comes from reduction and oxidation, which always happen together: when one atom loses electrons, another gains them. Oxidation is the loss of electrons, and reduction is the gain. This language makes sense when you remember that oxidation state is a kind of charge bookkeeping: adding an electron reduces the oxidation state because electrons carry charge \(-1\).
Rusting is the most familiar case. Metallic iron has loosely held electrons that it can give up. Oxygen, by contrast, has six valence electrons and would prefer to have eight. When iron is exposed to air and moisture, atoms at the surface lose electrons and become Fe²⁺ or Fe³⁺ (there’s the sign change: by losing electrons the atom becomes a positively charged ion, Fe²⁺ or Fe³⁺). Those electrons flow to oxygen molecules dissolved in water, reducing each oxygen atom from oxidation state 0 in O₂ to −2 in hydroxide ions (OH⁻). The iron ions then combine with these oxygen species and water to form hydrated oxides, which we see as rust—see callout. The whole process is a redox reaction: iron is oxidized, oxygen is reduced.
Electrons released from iron flow to dissolved O₂ at cathodic spots on the wet surface. In water, O₂ is reduced as
\[ \mathrm{O_2 + 4e^- + 2H_2O \;\to\; 4OH^-}. \]
In this step, the two oxygen atoms that were in O₂ each change oxidation state 0 → −2 and end up as the oxygen in two of the OH⁻ ions; the other two OH⁻ oxygens come from the two H₂O molecules and already had oxidation state −2. The resulting OH⁻ then reacts with Fe²⁺/Fe³⁺ and additional O₂/H₂O to give Fe(OH)₂, Fe(OH)₃/FeOOH, and ultimately hydrated iron(III) oxides (rust).
Why is oxygen thought of as “−2”? A neutral oxygen atom has eight protons and eight electrons. It has six in its outer shell and tends to gain two more to complete the octet. When it does, it carries two extra negative charges compared to neutral, so in oxidation state bookkeeping we assign oxygen the value −2. In iron oxide, three O²⁻ ions (\(−6\) in total) are balanced by two Fe³⁺ ions (\(+6\)), giving an electrically neutral compound.
The reverse process is smelting, where iron ore such as Fe₂O₃ is converted back to metallic iron. In the ore, iron is in the +3 oxidation state, and to become neutral metal it must gain three electrons. That is reduction. The oxide ions, which started as O²⁻, are forced to give up their electrons and leave as O₂ gas. That is oxidation. So the same element, oxygen, is oxidized in one context and reduced in another.
Water plays a special role in rusting. It does not provide the oxygen atoms, which come from dissolved O₂ gas, but it enables the whole process. Water dissolves both oxygen and iron ions, allows them to move, and directly participates in the reduction reaction. Rusting is best imagined as a microscopic galvanic cell (see callout), with tiny anode and cathode spots forming on the metal surface. Without water, the cycle stalls, which is why dry iron does not rust.
In this way the iron cycle alternates between oxidation in the environment and reduction in the furnace. Oxidizing means rusting, reducing means smelting, and together they tell the story of how iron interacts with air and water.
Rusting works like a tiny, uncontrolled battery. Different spots on the iron surface act as the two poles of a galvanic cell:
Anode (oxidation): Fe → Fe²⁺ + 2e⁻ Iron atoms lose electrons and go into solution.
Cathode (reduction): O₂ + 4e⁻ + 2H₂O → 4OH⁻ Dissolved oxygen molecules gain the electrons.
The electrons released at the anodic spots travel through the iron itself to cathodic spots, where oxygen is reduced. Meanwhile, ions move through the thin water layer acting as an electrolyte. In this way the corroding iron surface becomes a patchwork of microscopic galvanic cells, each driving the redox reactions that create rust.
The formula Fe₂O₃ does not describe a molecule but a lattice. In hematite, iron and oxygen atoms are arranged in an extended 3D solid, with Fe³⁺ and O²⁻ ions repeating in a regular pattern. There are no discrete “Fe₂O₃ units” floating around, only the overall ratio of two irons to three oxygens.
The bonding is mostly ionic, since iron readily loses electrons to oxygen, but it is not purely so. Overlap between Fe 3d and O 2p orbitals gives the bonds some covalent character, which helps explain hematite’s properties such as its color and magnetism. Like many metal oxides, Fe₂O₃ sits on the spectrum between ionic and covalent bonding, but the lattice structure is the clear sign that it is not a molecular compound.
Other Examples
Other examples of redox are found throughout metallurgy and everyday life. Copper, for instance, was one of the first metals smelted by humans. The ore copper(II) oxide, CuO, contains copper in the +2 oxidation state. In a furnace it is reduced by carbon monoxide, which itself is oxidized to carbon dioxide: \[ \mathrm{CuO + CO \;\to\; Cu + CO_2}. \] Here copper(II) gains two electrons to become metallic copper, while carbon goes from +2 in CO to +4 in CO₂.
Aluminium is much harder to extract, because its oxide Al₂O₃ is extremely stable. The reduction cannot be done by simple heating with carbon. Instead, aluminium is produced by electrolysis in the Hall–Héroult process. Al³⁺ ions in molten cryolite are reduced at the cathode: \[ \mathrm{Al^{3+} + 3e^- \;\to\; Al}, \] while oxide ions are oxidized at the anode: \[ \mathrm{2O^{2-} \;\to\; O_2 + 4e^-}. \] This is a redox process driven by electrical energy rather than chemical fuel.
Zinc extraction also relies on reduction. In zinc oxide, Zn is in the +2 state. Heating ZnO with carbon produces metallic zinc and carbon monoxide: \[ \mathrm{ZnO + C \;\to\; Zn + CO}. \] Again the metal is reduced to oxidation state 0, while carbon is oxidized from 0 in graphite to +2 in CO.
Redox is not limited to reactions with oxygen. A simple case is the displacement of copper by zinc in solution. If a strip of zinc is placed in copper(II) sulfate, the zinc is oxidized from 0 to +2, while copper is reduced from +2 to 0: \[ \mathrm{Zn(s) + Cu^{2+}(aq) \;\to\; Zn^{2+}(aq) + Cu(s)}. \] This is the basis of the Daniell cell, one of the earliest batteries.
Halogen chemistry also provides clean oxygen-free examples. If bromine water is added to potassium iodide, bromine is reduced from 0 to −1, while iodide is oxidized from −1 to 0, liberating iodine: \[ \mathrm{Br_2(aq) + 2I^-(aq) \;\to\; 2Br^-(aq) + I_2(aq)}. \]
Another important oxygen-free redox system is the Fe–Ce titration used in analysis. In acidic solution, iron(II) is oxidized to iron(III), while cerium(IV) is reduced to cerium(III): \[ \mathrm{Fe^{2+}(aq) + Ce^{4+}(aq) \;\to\; Fe^{3+}(aq) + Ce^{3+}(aq)}. \]
These examples show that redox is broader than rusting or combustion: it is any process where electrons are transferred, regardless of whether oxygen is involved.
Beyond metallurgy, redox processes dominate our world. Combustion is a classic case: hydrocarbons are oxidized to carbon dioxide and water, with oxygen reduced in the process, releasing energy. In respiration the same chemistry occurs more gently inside cells, with glucose oxidized and oxygen reduced, capturing the energy in ATP. Photosynthesis is the reverse: carbon dioxide is reduced to carbohydrates while water is oxidized to oxygen, powered by sunlight. And in every battery the same principle appears in controlled form, as one electrode undergoes oxidation, the other reduction, and the flow of electrons provides usable electrical power.
Oxidation States
Oxidation states are the bookkeeping system that keeps track of electron transfer in all these reactions. Each element has certain states that are especially stable, often linked to its electron configuration. Iron, for example, is commonly found as Fe²⁺ and Fe³⁺, and it flips between these two in rusting. Copper also has two familiar states, Cu⁺ and Cu²⁺, with the latter usually more stable. Zinc, by contrast, shows almost exclusively the +2 state, reflecting the filled d-shell that resists further change.
Some elements are versatile, with a whole range of possible oxidation states. Manganese has the most, from −3 up to +7, with +2, +4, and +7 being common; this is why it appears in so many colored minerals and reagents. Vanadium and chromium show similar variety. At the other extreme, the noble gases were once thought to be chemically inert with only the 0 state, but heavier members like xenon can be forced into compounds with positive states such as +2, +4, +6, and +8.
The halogens (from Greek hals (ἅλς, salt) + -gen (producer) -> “salt-formers”) are another instructive case. Fluorine is the most electronegative element and is found at −1 in every compound, without exception. Chlorine, by contrast, is usually −1 but can also appear at +1, +3, +5, and +7 in compounds with oxygen or fluorine, such as chlorates and perchlorates. These higher states are less stable, but they are important as strong oxidizers and in industrial chemistry.
The “preferred” states reflect a balance between gaining the stability of filled or half-filled electron shells and the energy required to move electrons around. Oxygen nearly always sits at −2 because that completes its octet. Alkali metals are almost always +1, alkaline earths +2, both because they readily shed their few valence electrons and then stop. Transition metals sit in the middle, where several states are accessible, and their rich redox chemistry underlies much of inorganic and bioinorganic chemistry.
Carbon spans a wide range of oxidation states, from −4 in methane (CH₄) through 0 in elemental graphite and diamond, up to +4 in carbon dioxide (CO₂). Intermediate values are found in countless compounds, such as −3 in propane, +2 in carbon monoxide, and +4 in carbonate. This flexibility underlies carbon’s unique role in chemistry: it can act as both electron donor and acceptor, it forms stable bonds across the entire range, and it drives the central redox cycles of life itself. In respiration, carbon is oxidized step by step from glucose (average oxidation state near 0) to CO₂ at +4, while photosynthesis reverses the path.
Sulfur is another element with a wide range of oxidation states, from −2 in sulfides like FeS, through 0 in elemental S₈, up to +4 in sulfites and +6 in sulfates. The −2 state is common in minerals and in biological systems such as amino acids, while the higher states appear in atmospheric chemistry and industrial processes. This breadth reflects sulfur’s ability to access multiple bonding arrangements beyond a simple octet, giving it an unusually rich redox chemistry.
Thinking in terms of oxidation states allows us to see redox processes as a shifting landscape. Some elements are locked into one role, others are flexible actors able to move between several states. It is this flexibility that powers the variety of redox reactions from smelting ores to burning fuels to the delicate electron transfers in living cells.
In high school chemistry, carbon is often introduced as “valence 4,” meaning it has four valence electrons and can complete an octet by forming four bonds. Pushed to the extremes, you can picture carbon either gaining four electrons to make C⁴⁻ (as in methane, −4) or losing four to make C⁴⁺ (as in carbon dioxide, +4).
That picture is correct, but it is not the whole story. In covalent bonds the electrons are shared, and oxidation state bookkeeping assigns them to the more electronegative atom. This allows carbon to sit comfortably at almost any integer state between −4 and +4, depending on what it is bonded to. That flexibility explains why carbon chemistry is so rich: it can be oxidized step by step in respiration, reduced step by step in photosynthesis, and form stable compounds at every stage in between.
| Element | Common oxidation states | Examples |
|---|---|---|
| Carbon | −4 to +4 | CH₄ (−4), CO ( +2), CO₂ (+4), CO₃²⁻ (+4) |
| Nitrogen | −3 to +5 | NH₃ (−3), N₂ (0), NO₂ ( +4), NO₃⁻ (+5) |
| Sulfur | −2 to +6 | H₂S (−2), S₈ (0), SO₂ (+4), SO₄²⁻ (+6) |
Glossary of Chemicals
- Fe₂O₃ – iron(III) oxide (hematite, common iron ore)
- Fe(OH)₂ – iron(II) hydroxide (intermediate in rusting)
- Fe(OH)₃ / FeOOH – iron(III) hydroxide / iron oxyhydroxide (components of rust)
- CuO – copper(II) oxide (black oxide of copper)
- CuSO₄ – copper(II) sulfate (blue vitriol, used in agriculture and chemistry)
- Al₂O₃ – aluminium oxide (corundum, main ore is bauxite; used as an abrasive)
- NaCl – sodium chloride (common salt, table salt)
- ZnO – zinc oxide (white pigment, sunscreen ingredient)
- CO – carbon monoxide (toxic gas, important industrial reducing agent in metallurgy)
- CO₂ – carbon dioxide (greenhouse gas, dry ice in solid form)
- CO₃²⁻ – carbonate ion (in limestones, chalk, shells)
- CH₄ – methane (natural gas, fuel)
- C₃H₈ – propane (bottled gas for heating/cooking)
- C₆H₁₂O₆ – glucose (blood sugar, central to metabolism)
- O₂ – oxygen molecule (respiration, combustion)
- OH⁻ – hydroxide ion (in alkalis such as NaOH)
- SO₂ – sulfur dioxide (acid rain precursor, food preservative)
- SO₃²⁻ – sulfite ion (preservatives, wine-making)
- SO₄²⁻ – sulfate ion (in gypsum, Epsom salts)
- H₂S – hydrogen sulfide (toxic gas with “rotten egg” smell)
- S₈ – elemental sulfur (yellow solid, used in vulcanizing rubber)
- NH₃ – ammonia (fertilizer, cleaning agent)
- N₂ – nitrogen molecule (main component of air, inert atmosphere)
- NO₂ – nitrogen dioxide (brown gas, air pollutant)
- NO₃⁻ – nitrate ion (fertilizers, explosives)
- Br₂ – bromine molecule (red-brown liquid, disinfectants, flame retardants)
- Br⁻ – bromide ion (in salts, sedatives historically)
- I₂ – iodine molecule (antiseptic, nutrition supplement)
- I⁻ – iodide ion (in iodized salt, thyroid function)
- Ce³⁺ / Ce⁴⁺ – cerium(III) and cerium(IV) ions (oxidizing agents in analysis, catalysts)
- Fe²⁺ / Fe³⁺ – iron(II) and iron(III) ions (found in minerals, rust chemistry, biological systems)
- Roman numerals in names indicate the oxidation state of the metal. For example, iron(III) oxide (Fe₂O₃) has Fe in the +3 state; iron(II) oxide (FeO) has Fe in the +2 state. This avoids ambiguity for metals with multiple common states.
- -ide ending: simple binary compounds with that element in its most reduced common state. Examples: chloride (Cl⁻), sulfide (S²⁻), oxide (O²⁻).
- -ite ending: oxyanions with fewer oxygen atoms. Example: sulfite (SO₃²⁻), nitrite (NO₂⁻).
- -ate ending: oxyanions with more oxygen atoms. Example: sulfate (SO₄²⁻), nitrate (NO₃⁻).
- per-…-ate: oxyanions with the maximum oxygen number in the series (e.g. perchlorate, ClO₄⁻).
- hypo-…-ite: oxyanions with the minimum oxygen number in the series (e.g. hypochlorite, ClO⁻).