Redox

Rates, mechanisms, and energies

Redox reactions can occur slowly, as in the formation of rust, or rapidly, as in the case of burning fuel. Electron transfer reactions are generally fast, occurring within the time of mixing.

The mechanisms of atom-transfer reactions are highly variable because many kinds of atoms can be transferred, and such reactions can involve several steps. Electron-transfer reactions, by contrast, proceed by two distinct pathways. In inner-sphere transfer, the two reactants share a bridging ligand through which the electron passes; in outer-sphere transfer, the electron moves between reactants whose coordination shells remain intact. Henry Taube received the 1983 Nobel Prize in Chemistry for distinguishing these pathways through experiments on metal complexes.

The rate of an outer-sphere electron transfer is described by Marcus theory, developed by Rudolph A. Marcus. The theory expresses the activation energy in terms of two quantities: the standard free-energy change of the reaction and the reorganization energy, the energy needed to distort the reactants and the surrounding solvent into the configuration of the products before the electron moves. It predicts an "inverted region", in which the rate falls once the driving force exceeds the reorganization energy. Marcus received the 1992 Nobel Prize in Chemistry for the theory.

Analysis of bond energies and ionization energies in water allows calculation of the thermodynamic aspects of redox reactions.

Standard electrode potentials (reduction potentials)

Each half-reaction has a standard electrode potential (E^o
_cell), which is equal to the potential difference or voltage at equilibrium under standard conditions of an electrochemical cell in which the cathode reaction is the half-reaction considered, and the anode is a standard hydrogen electrode where hydrogen is oxidized:

1⁄2 H₂ → H⁺ + e⁻

The electrode potential of each half-reaction is also known as its reduction potential (E^o
_red), or potential when the half-reaction takes place at a cathode. The reduction potential is a measure of the tendency of the oxidizing agent to be reduced. Its value is zero for H⁺ + e⁻ → 1⁄2H₂ by definition, positive for oxidizing agents stronger than H⁺ (e.g., +2.866 V for F₂) and negative for oxidizing agents that are weaker than H⁺ (e.g., −0.763V for Zn²⁺).

For a redox reaction that takes place in a cell, the potential difference is:

E^o
_cell = E^o
_cathode − E^o
_anode

However, the potential of the reaction at the anode is sometimes expressed as an oxidation potential:

E^o
_ox = −E^o
_red

The oxidation potential is a measure of the tendency of the reducing agent to be oxidized but does not represent the physical potential at an electrode. With this notation, the cell voltage equation is written with a plus sign

E^o
_cell = E^o
_red(cathode) + E^o
_ox(anode)

Examples of redox reactions

In the reaction between hydrogen and fluorine, hydrogen is being oxidized and fluorine is being reduced:

H₂ + F₂ → 2 HF

This spontaneous reaction releases a large amount of energy (542 kJ per 2 g of hydrogen) because two H-F bonds are much stronger than one H-H bond and one F-F bond. This reaction can be analyzed as two half-reactions. The oxidation reaction converts hydrogen to protons:

H₂ → 2 H⁺ + 2 e⁻

The reduction reaction converts fluorine to the fluoride anion:

F₂ + 2 e⁻ → 2 F⁻

The half-reactions are combined so that the electrons cancel:

The protons and fluoride combine to form hydrogen fluoride in a non-redox reaction:

2 H⁺ + 2 F⁻ → 2 HF

The overall reaction is:

H₂ + F₂ → 2 HF

Metal displacement

In this type of reaction, a metal atom in a compound or solution is replaced by an atom of another metal. For example, copper is deposited when zinc metal is placed in a copper(II) sulfate solution:

Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

In the above reaction, zinc metal displaces the copper(II) ion from the copper sulfate solution, thus liberating free copper metal. The reaction is spontaneous and releases 213 kJ per 65 g of zinc.

The ionic equation for this reaction is:

Zn + Cu²⁺ → Zn²⁺ + Cu

As two half-reactions, it is seen that the zinc is oxidized:

Zn → Zn²⁺ + 2 e⁻

And the copper is reduced:

Cu²⁺ + 2 e⁻ → Cu

Other examples

• The reduction of nitrate to nitrogen in the presence of an acid (denitrification):

2 NO−3 + 10 e⁻ + 12 H⁺ → N₂ + 6 H₂O

• The combustion of hydrocarbons, such as in an internal combustion engine, produces water, carbon dioxide, some partially oxidized forms such as carbon monoxide, and heat energy. Complete oxidation of materials containing carbon produces carbon dioxide.
• The stepwise oxidation of a hydrocarbon by oxygen, in organic chemistry, produces water and, successively: an alcohol, an aldehyde or a ketone, a carboxylic acid, and then a peroxide.

Corrosion and rusting

• The term corrosion refers to the electrochemical oxidation of metals in reaction with an oxidant such as oxygen. Rusting, the formation of iron oxides, is a well-known example of electrochemical corrosion: it forms as a result of the oxidation of iron metal. Common rust often refers to iron(III) oxide, formed in the following chemical reaction:

4 Fe + 3 O₂ → 2 Fe₂O₃

• The oxidation of iron(II) to iron(III) by hydrogen peroxide in the presence of an acid:

Fe²⁺ → Fe³⁺ + e⁻

H₂O₂ + 2 e⁻ → 2 OH⁻

Here the overall equation involves adding the reduction equation to twice the oxidation equation, so that the electrons cancel:

2 Fe²⁺ + H₂O₂ + 2 H⁺ → 2 Fe³⁺ + 2 H₂O

Disproportionation

A disproportionation reaction is one in which a single substance is both oxidized and reduced. For example, thiosulfate ion with sulfur in oxidation state +2 can react in the presence of acid to form elemental sulfur (oxidation state 0) and sulfur dioxide (oxidation state +4).

S₂O2−3 + 2 H⁺ → S + SO₂ + H₂O

Thus one sulfur atom is reduced from +2 to 0, while the other is oxidized from +2 to +4.

Redox reactions in industry

Cathodic protection is a technique used to control the corrosion of a metal surface by making it the cathode of an electrochemical cell. A simple method of protection connects protected metal to a more easily corroded "sacrificial anode" to act as the anode. The sacrificial metal, instead of the protected metal, then corrodes.

Oxidation is used in many industries, such as in the production of cleaning products and oxidizing ammonia to produce nitric acid, whose principal industrial route is the catalytic oxidation of ammonia in the Ostwald process.

Redox reactions are the foundation of electrochemical cells, which can generate electrical energy or support electrosynthesis. Metal ores often contain metals in oxidized states, such as oxides or sulfides, from which the pure metals are extracted by smelting at high temperatures in the presence of a reducing agent. The process of electroplating uses redox reactions to coat objects with a thin layer of a material, as in chrome-plated automotive parts, silver plating cutlery, galvanization and gold-plated jewelry.

Redox reactions in biology

Many essential biological processes involve redox reactions. Before some of these processes can begin, iron must be assimilated from the environment.

Aerobic cellular respiration, for instance, is the oxidation of substrates (in this case, glucose [C₆H₁₂O₆]) and the reduction of oxygen to water. The summary equation for aerobic respiration is:

C₆H₁₂O₆ + 6 O₂ → 6 CO₂ + 6 H₂O + Energy

The process of cellular respiration also depends heavily on the reduction of NAD⁺ to NADH and the reverse reaction (the oxidation of NADH to NAD⁺). Photosynthesis and cellular respiration are complementary, but photosynthesis is not the reverse of the redox reaction in cellular respiration:

6 CO₂ + 6 H₂O + light energy → C₆H₁₂O₆ + 6 O₂

Biological energy is frequently stored and released using redox reactions. Photosynthesis involves the reduction of carbon dioxide into sugars and the oxidation of water into molecular oxygen. The reverse reaction, respiration, oxidizes sugars to produce carbon dioxide and water. As intermediate steps, the reduced carbon compounds are used to reduce nicotinamide adenine dinucleotide (NAD⁺) to NADH, which then contributes to the creation of a proton gradient, which drives the synthesis of adenosine triphosphate (ATP) and is maintained by the reduction of oxygen. In animal cells, mitochondria perform similar functions.

The term redox state is often used to describe the balance of GSH/GSSG, NAD⁺/NADH and NADP⁺/NADPH in a biological system such as a cell or organ. The redox state is reflected in the balance of several sets of metabolites (e.g., lactate and pyruvate, beta-hydroxybutyrate and acetoacetate), whose interconversion is dependent on these ratios. Redox mechanisms also control some cellular processes. Redox proteins and their genes must be co-located for redox regulation according to the CoRR hypothesis for the function of DNA in mitochondria and chloroplasts.

Redox cycling

Many aromatic compounds are enzymatically reduced to form free radicals that contain one more electron than their parent compounds. In general, the electron donor is a flavoenzyme or one of its coenzymes. Once formed, these anion free radicals reduce molecular oxygen to superoxide and regenerate the unchanged parent compound. The net reaction is the oxidation of the flavoenzyme's coenzymes and the reduction of molecular oxygen to form superoxide. This catalytic behavior has been described as a futile cycle or redox cycling.

Redox reactions in geology

Minerals are generally oxidized derivatives of metals. Iron is mined as ores such as magnetite (Fe₃O₄) and hematite (Fe₂O₃). Titanium is mined as its dioxide, usually in the form of rutile (TiO₂). These oxides must be reduced to obtain the corresponding metals, often achieved by heating these oxides with carbon or carbon monoxide as reducing agents. Blast furnaces are the reactors where iron oxides and coke (a form of carbon) are combined to produce molten iron. The main chemical reaction producing the molten iron is:

Fe₂O₃ + 3 CO → 2 Fe + 3 CO₂

Redox reactions in soils

Electron transfer reactions are central to many processes and properties in soils, and redox potential, quantified as Eh (platinum electrode potential (voltage) relative to the standard hydrogen electrode) or pe (analogous to pH as −log electron activity), is a master variable, along with pH, that controls and is governed by chemical reactions and biological processes. Early theoretical research with applications to flooded soils and paddy rice production shaped subsequent work on thermodynamic aspects of redox and plant root growth in soils. Later work extended it to redox reactions related to heavy metal oxidation state changes, pedogenesis and morphology, organic compound degradation and formation, free radical chemistry, wetland delineation, soil remediation, and methods for characterizing the redox status of soils.

Mnemonics

The key terms involved in redox can be confusing. For example, a reagent that is oxidized loses electrons; however, that reagent is referred to as the reducing agent. Likewise, a reagent that is reduced gains electrons and is referred to as the oxidizing agent. These mnemonics are commonly used by students to help memorise the terminology:

• "OIL RIG"—oxidation is loss of electrons, reduction is gain of electrons
• "LEO the lion says GER [grr]"—loss of electrons is oxidation, gain of electrons is reduction
• "LEORA says GEROA"—the loss of electrons is called oxidation (reducing agent); the gain of electrons is called reduction (oxidizing agent).
• "RED CAT" and "AN OX", or "AnOx RedCat" ("an ox-red cat")—reduction occurs at the cathode and the anode is for oxidation
• "RED CAT gains what AN OX loses" – reduction at the cathode gains (electrons) what anode oxidation loses (electrons)
• "PANIC" – Positive Anode and Negative is Cathode. This applies to electrolytic cells that release stored electricity, and can be recharged with electricity. PANIC does not apply to cells that can be recharged with redox materials. These galvanic or voltaic cells, such as fuel cells, produce electricity from internal redox reactions. Here, the positive electrode is the cathode and the negative is the anode.