Learning Objectives Define oxidation and reduction Determine the valence (charge) of an element. Determine which element is being oxidized and which is being reduced in a redox reaction. Oxidation-Reduction Reactions Earth’s atmosphere contains about 20% molecular oxygen, O 2 , a chemically reactive gas that plays an essential role in the metabolism of aerobic organisms and in many environmental processes that shape the world. The term oxidation was originally used to describe chemical reactions involving O 2 , but its meaning has evolved to refer to a broad and important reaction class known as oxidation-reduction (redox) reactions . A few examples of such reactions will be used to develop a clear picture of this classification. Some redox reactions involve the transfer of electrons between reactant species to yield ionic products, such as the reaction between sodium and chlorine to yield sodium chloride: \[\ce{2Na}(s)+\ce{Cl2}(g)\rightarrow \ce{2NaCl}(s)\] It is helpful to view the process with regard to each individual reactant, that is, to represent the fate of each reactant in the form of an equation called a half-reaction : \[\ce{2Na}(s)\rightarrow \ce{2Na+}(s)+\ce{2e-}\] \[\ce{Cl2}(g)+\ce{2e-}\rightarrow \ce{2Cl-}(s)\] These equations show that Na atoms lose electrons while Cl atoms (in the Cl 2 molecule) gain electrons , the “ s ” subscripts for the resulting ions signifying they are present in the form of a solid ionic compound. For redox reactions of this sort, the loss and gain of electrons define the complementary processes that occur: \(\begin{align} \textbf{oxidation}&=\textrm{loss of electrons}\\ \textbf{reduction}&=\textrm{gain of electrons} \end{align}\) Remembering Oxidation and Reduction It is common to remember the difference between oxidation and reduction using one of two pneumonic devices: "LeO goes GeR" In Greek Mythology, Leo was a lion (like the zodiac sign). And what do lion's do? Well, they roar but that doesn't really fit, so we go with "Ger". L oses e lectrons = O xidation G ains e lectrons = R eduction Figure \(\PageIndex{4}\): Lion image by Tambako, reused under the Creative Commons Attribution 2.0 Generic license. "Oil Rig" O xidation i s l oss R eduction i s g ain Figure \(\PageIndex{5}\): Image by Peter Grima, reused under the Creative Commons Attribution 2.0 Generic license. In this reaction, then, sodium is oxidized and chlorine undergoes reduction . Viewed from a more active perspective, sodium functions as a reducing agent (reductant) , since it provides electrons to (or reduces) chlorine. Likewise, chlorine functions as an oxidizing agent (oxidant) , as it effectively removes electrons from (oxidizes) sodium. \(\begin{align} \textbf{reducing agent}&=\textrm{species that is oxidized}\\ \textbf{oxidizing agent}&=\textrm{species that is reduced} \end{align}\) Some redox processes, however, do not involve the transfer of electrons. Consider, for example, a reaction similar to the one yielding NaCl: \[\ce{H2}(g)+\ce{Cl2}(g)\rightarrow \ce{2HCl}(g)\] The product of this reaction is a covalent compound, so transfer of electrons in the explicit sense is not involved. To clarify the similarity of this reaction to the previous one and permit an unambiguous definition of redox reactions, a property called oxidation number has been defined. The oxidation number (or oxidation state ) of an element in a compound is the charge its atoms would possess if the compound was ionic . Oxidation-reduction (redox) reactions are those in which one or more elements involved undergo a change in oxidation number. While the vast majority of redox reactions involve changes in oxidation number for two or more elements, a few interesting exceptions to this rule do exist as shown below\). Definitions for the complementary processes of this reaction class are correspondingly revised as shown here: \[\begin{align} \textbf{oxidation}&=\textrm{increase in oxidation number}\\ \textbf{reduction}&=\textrm{decrease in oxidation number} \end{align}\] Returning to the reactions used to introduce this topic, they may now both be identified as redox processes. In the reaction between sodium and chlorine to yield sodium chloride, sodium is oxidized (its oxidation number increases from 0 in Na to +1 in NaCl) and chlorine is reduced (its oxidation number decreases from 0 in Cl 2 to −1 in NaCl). In the reaction between molecular hydrogen and chlorine, hydrogen is oxidized (its oxidation number increases from 0 in H 2 to +1 in HCl) and chlorine is reduced (its oxidation number decreases from 0 in Cl 2 to −1 in HCl). Several subclasses of redox reactions are recognized, including combustion reactions in which the reductant (also called a fuel ) and oxidant (often, but not necessarily, molecular oxygen) react vigorously and produce significant amounts of heat, and often light, in the form of a flame. Solid rocket-fuel reactions such as the one depicted below are combustion processes. A typical propellant reaction in which solid aluminum is oxidized by ammonium perchlorate is represented by this equation: \[\ce{10Al}(s)+\ce{6NH4ClO4}(s)\rightarrow \ce{4Al2O3}(s)+\ce{2AlCl3}(s)+\ce{12H2O}(g)+\ce{3N2}(g)\] Watch a brief video showing the test firing of a small-scale, prototype, hybrid rocket engine planned for use in the new Space Launch System being developed by NASA. The first engines firing at 3 s (green flame) use a liquid fuel/oxidant mixture, and the second, more powerful engines firing at 4 s (yellow flame) use a solid mixture. Single-displacement (replacement) reactions are redox reactions in which an ion in solution is displaced (or replaced) via the oxidation of a metallic element. One common example of this type of reaction is the acid oxidation of certain metals: \[\ce{Zn}(s)+\ce{2HCl}(aq)\rightarrow \ce{ZnCl2}(aq)+\ce{H2}(g)\] Metallic elements may also be oxidized by solutions of other metal salts; for example: \[\ce{Cu}(s)+\ce{2AgNO3}(aq)\rightarrow \ce{Cu(NO3)2}(aq)+\ce{2Ag}(s)\] This reaction may be observed by placing copper wire in a solution containing a dissolved silver salt. Silver ions in solution are reduced to elemental silver at the surface of the copper wire, and the resulting Cu 2+ ions dissolve in the solution to yield a characteristic blue color (Figure \(\PageIndex{6}\)). Figure \(\PageIndex{6}\): (a) A copper wire is shown next to a solution containing silver(I) ions. (b) Displacement of dissolved silver ions by copper ions results in (c) accumulation of gray-colored silver metal on the wire and development of a blue color in the solution, due to dissolved copper ions. (credit: modification of work by Mark Ott) Example \(\PageIndex{3}\): Describing Redox Reactions Identify which equations represent redox reactions, providing a name for the reaction if appropriate. For those reactions identified as redox, name the oxidant and reductant. \(\ce{ZnCO3}(s)\rightarrow \ce{ZnO}(s)+\ce{CO2}(g)\) \(\ce{2Ga}(l)+\ce{3Br2}(l)\rightarrow \ce{2GaBr3}(s)\) \(\ce{2H2O2}(aq)\rightarrow \ce{2H2O}(l)+\ce{O2}(g)\) \(\ce{BaCl2}(aq)+\ce{K2SO4}(aq)\rightarrow \ce{BaSO4}(s)+\ce{2KCl}(aq)\) \(\ce{C2H4}(g)+\ce{3O2}(g)\rightarrow \ce{2CO2}(g)+\ce{2H2O}(l)\) S olution Redox reactions are identified per definition if one or more elements undergo a change in oxidation number. This is not a redox reaction, since oxidation numbers remain unchanged for all elements. This is a redox reaction. Gallium is oxidized, its oxidation number increasing from 0 in Ga( l ) to +3 in GaBr 3 ( s ). The reducing agent is Ga( l ). Bromine is reduced, its oxidation number decreasing from 0 in Br 2 ( l ) to −1 in GaBr 3 ( s ). The oxidizing agent is Br 2 ( l ). This is a redox reaction. It is a particularly interesting process, as it involves the same element, oxygen, undergoing both oxidation and reduction (a so-called disproportionation reaction) . Oxygen is oxidized, its oxidation number increasing from −1 in H 2 O 2 ( aq ) to 0 in O 2 ( g ). Oxygen is also reduced, its oxidation number decreasing from −1 in H 2 O 2 ( aq ) to −2 in H 2 O( l ). For disproportionation reactions, the same substance functions as an oxidant and a reductant. This is not a redox reaction, since oxidation numbers remain unchanged for all elements. This is a redox reaction (combustion). Carbon is oxidized, its oxidation number increasing from −2 in C 2 H 4 ( g ) to +4 in CO 2 ( g ). The reducing agent (fuel) is C 2 H 4 ( g ). Oxygen is reduced, its oxidation number decreasing from 0 in O 2 ( g ) to −2 in H 2 O( l ). The oxidizing agent is O 2 ( g ). Exercise \(\PageIndex{3}\) This equation describes the production of tin(II) chloride: \[\ce{Sn}(s)+\ce{2HCl}(g)\rightarrow \ce{SnCl2}(s)+\ce{H2}(g)\] Is this a redox reaction? If so, provide a more specific name for the reaction if appropriate, and identify the oxidant and reductant. Answer Yes, a single-replacement reaction. Sn( s )is the reductant, HCl( g ) is the oxidant. Summary Video \(\PageIndex{1}\): A summary of redox reactions. One of your professors was a script editor for this video. Hint: It was Dr. Clark Redox reactions involve a change in oxidation number for one or more reactant elements. Writing balanced equations for some redox reactions that occur in aqueous solutions is simplified by using a systematic approach called the half-reaction method. Glossary combustion reaction vigorous redox reaction producing significant amounts of energy in the form of heat and, sometimes, light oxidation process in which an element’s oxidation number is increased by loss of electrons oxidation-reduction reaction (also, redox reaction) reaction involving a change in oxidation number for one or more reactant elements oxidation number (also, oxidation state) the charge each atom of an element would have in a compound if the compound were ionic oxidizing agent (also, oxidant) substance that brings about the oxidation of another substance, and in the process becomes reduced reduction process in which an element’s oxidation number is decreased by gain of electrons reducing agent (also, reductant) substance that brings about the reduction of another substance, and in the process becomes oxidized single-displacement reaction (also, replacement) redox reaction involving the oxidation of an elemental substance by an ionic species