Redox reactions (reduction-oxidation) describe the transfer of electrons between chemical species. Oxidation is defined as the loss of electrons (LEO), and reduction is the gain of electrons (GER). Since electrons cannot vanish, oxidation and reduction must always occur together, as one substance loses electrons for another to gain them.
The Fundamental Answer: Metals are Oxidized
Elemental metals are overwhelmingly oxidized when they participate in a chemical reaction. This behavior is rooted in the atomic structure of metals, which typically possess a small number of valence electrons in their outermost shell. Metals have a natural tendency to achieve a more stable, lower-energy electron configuration by shedding these valence electrons. This loss of electrons is the definition of oxidation.
A metal atom transitions from a neutral state to a positively charged ion, known as a cation, during this process. For example, when sodium metal reacts, it readily loses its single valence electron to form a sodium ion with a stable +1 charge. Magnesium, a Group 2 metal, loses two valence electrons to form a Mg\(^{2+}\) ion when reacting with oxygen. The low energy required to remove these outer electrons drives the metal to participate in this oxidizing reaction.
The Consequence: Metals Act as Reducing Agents
Because metals are oxidized, they simultaneously play a specific role as a reducing agent in the overall redox reaction. A reducing agent facilitates the reduction of another species by supplying the electrons the other species needs to gain. The metal becomes oxidized itself, but in doing so, it causes the other reactant to be reduced.
The metal undergoes oxidation, but it acts as the reducing agent for the other chemical. For example, in the reaction between sodium and chlorine gas to form table salt, the sodium metal is oxidized, but it is also the agent that reduces the chlorine. Chlorine is reduced by gaining the electrons, and it is therefore classified as the oxidizing agent.
The Exception: Reduction of Metal Ions
While elemental metals are oxidized, the reverse process occurs when a metal is already in its ionic form. Metal ions, which carry a positive charge, are capable of gaining electrons to return to their neutral, elemental metallic state. This process is common in industrial applications where pure metal is extracted or recovered from its compounds.
One commercial example is the smelting of iron, where iron ore (primarily iron oxide, Fe\(^{3+}\) ions) is heated in a blast furnace. Carbon monoxide acts as the reducing agent, supplying electrons to the iron ions so they can transition back to solid, elemental iron. Another example is electroplating, such as the coating of a surface with copper. In this electrical process, copper ions (Cu\(^{2+}\)) dissolved in a solution gain two electrons to deposit as a layer of solid copper metal.
Real-World Impact: Corrosion and Oxidation
The oxidation of metals is a process that affects nearly every metal structure and object in daily life, most notably through corrosion. Corrosion is the uncontrolled oxidation of a metal in the presence of an oxidizing agent, often oxygen, moisture, or a combination of both. The familiar reddish-brown rust on iron is a direct result of iron metal oxidizing to form iron oxide.
Other metals also undergo this reaction, though the results may not be as structurally damaging as rust. Silver tarnishes when it oxidizes by reacting with sulfur compounds in the air, forming a black layer of silver sulfide. Copper develops a green patina, such as that seen on old statues and roofs, which is a layer of copper carbonate that forms as the metal oxidizes. In some cases, like with aluminum, the oxide layer that forms is thin and durable, protecting the underlying metal from further oxidation.