Redox

 ( Soil Chemistry ) 

Redox ( reduction–oxidation, , ) is a type of chemical reaction in which the of substrate change.

• Oxidation is the loss of electrons or an increase in the oxidation state of a chemical or atoms within it.
• Reduction is the gain of electrons or a decrease in the oxidation state of a chemical or atoms within it.

There are two classes of redox reactions:

• electron-transfer where only one (usually) electron flows from the reducing agent to the oxidant. This type of redox reaction is often discussed in term of redox couples and electrode potentials.
• atom transfer, where an atom transfers from one substrate to another. For example, in the rusting of iron, the oxidation state of iron atoms increases as it converts to an oxide and simultaneously the oxidation state of oxygen decreases as it accepts electrons released by iron. Although oxidation reactions are commonly associated with the formation of oxides, other chemical species can serve the same function. In hydrogenation, alkene (and other) bonds are reduced by transfer of hydrogen atoms.

[[Oxygen]] is the quintessential oxidizer.
     

The mechanisms of atom-transfer reactions are highly variable because many kinds of atoms can be transferred. Such reactions can also be quite complex, i.e. involve many steps. The mechanisms of electron-transfer reactions occur by two distinct pathways, inner sphere electron transfer and outer sphere electron transfer.

Analysis of bond energies and ionization energies in water allow calculation of the thermodynamic aspects of redox reactions. For example, the fact that the combustion of methane yields more heat than the combustion of methanol.

 H₂ → H⁺ + e⁻.

The electrode potential of each half-reaction is also known as its reduction potential E, 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⁻ →  H₂ 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., −0.763 V for Zn²⁺).Electrode potential values from:

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

E = E – E

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

E = – E.

E = E + E

+ → 2 HF

This reaction is spontaneous and releases 542 kJ per 2 g of hydrogen because the H-F bond is much stronger than the F-F bond. This reaction can be analyzed as two . The oxidation reaction converts hydrogen to protons:

→ 2 Hydrogen ion + 2 Electron

The reduction reaction converts fluorine to the fluoride anion:

+ 2 e⁻ → 2 Fluoride

The half reactions are summed such that the electrons cancel:

>

|align=right	→	2 H+ + 2 e−
+ 2 e−	→	2 F−
H2 + F2	→	2 H+ + 2 F−

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

2 H⁺ + 2 F⁻ → 2 HF

The overall reaction is:

+ → 2 HF

---

Metal displacement

In this type of reaction, a metal atom in a compound (or in a 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 copper sulfate solution and thus liberates free copper metal. The reaction is spontaneous and releases 213 kJ per 65 g of zinc because relative to zinc, copper metal is lower in energy due to bonding via its partially filled d-orbitals.

The ionic equation for this reaction is:

Zn + Cu²⁺ → Zn²⁺ + Cu

As two , 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):2NO3- + 10e- + 12H+ -> N2 + 6 H2O
• The combustion of , 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. , the formation of , 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:4Fe + 3O2 -> 2Fe2O3
• The oxidation of iron(II) to iron(III) by hydrogen peroxide in the presence of an acid:Fe^{2+} -> {Fe^{3+}} + e-H2O2 + 2e- -> 2OH- Overall equation:{2Fe^{2+}} + {H2O2} + 2H+ -> {2Fe^{3+}} + 2H2O

---

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).

S₂O₃^2- + 2 H⁺ → S + SO₂ + H₂O

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

(2022). 9780130143297, Prentice Hall. ISBN 9780130143297

---

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. A common application of cathodic protection is in galvanization steel, in which a sacrificial coating of zinc on steel parts protects them from rust.

Oxidation is used in a wide variety of industries such as in the production of cleaning products and oxidizing ammonia to produce nitric acid.

Redox reactions are the foundation of electrochemical cells, which can generate electrical energy or support electrosynthesis. Metal often contain metals in oxidized states such as oxides or sulfides, from which the pure metals are extracted by smelting at high temperature 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 plating automotive parts, silver plating cutlery, galvanization and gold-plated jewelry.

---

Redox reactions in biology

Top: ascorbic acid (reducing agent of Vitamin C)
Bottom: dehydroascorbic acid (oxidizing agent of Vitamin C)

Many important biology processes involve redox reactions. Before some of these processes can begin iron must be assimilated from the environment.

(2022). 9783110588903, De Gruyter. ISBN 9783110588903

Cellular respiration, for instance, is the oxidation of glucose (C₆H₁₂O₆) to carbon dioxide and the reduction of oxygen to water. The summary equation for cell respiration is:

C₆H₁₂O₆ + 6 O₂ → 6 CO₂ + 6 H₂O

The process of cell 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 cell respiration:

6 CO₂ + 6 H₂O + photon → C₆H₁₂O₆ + 6 O₂

Biological energy is frequently stored and released by means of redox reactions. Photosynthesis involves the reduction of carbon dioxide into 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. See the Membrane potential article.

Free radical reactions are redox reactions that occur as a part of homeostasis and killing microorganisms, where an electron detaches from a molecule and then reattaches almost instantaneously. Free radicals are a part of redox molecules and can become harmful to the human body if they do not reattach to the redox molecule or an antioxidant. Unsatisfied free radicals can spur the mutation of cells they encounter and are, thus, causes of cancer.

The term redox state is often used to describe the balance of Glutathione, 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., lactic acid and pyruvate, beta-hydroxybutyrate, and acetoacetate), whose interconversion is dependent on these ratios. An abnormal redox state can develop in a variety of deleterious situations, such as hypoxia, shock, and sepsis. Redox mechanism 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

Wide varieties of aromaticity are enzyme reduced to form free radicals that contain one more electron than their parent compounds. In general, the electron donor is any of a wide variety of flavoenzymes and their . 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 its magnetite (Fe₃O₄). Titanium is mined as its dioxide, usually in the form of rutile (TiO₂). To obtain the corresponding metals, these oxides must be reduced, which is often achieved by heating these oxides with carbon or carbon monoxide as reducing agents. 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 + 3CO → 2Fe + 3CO₂

---

Redox reactions in soils

Electron transfer reactions are central to myriad processes and properties in soils, and electron "activity", 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 was seminal for subsequent work on thermodynamic aspects of redox and plant root growth in soils.

F.N. Ponnamperuma (1992). 9780120007240 ISBN 9780120007240

Later work built on this foundation, and expanded it for understanding 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 various methodological approaches for characterizing the redox status of soils.

(2022). 9781439803059, CRC Press. ISBN 9781439803059

---

Mnemonics

The key terms involved in redox can be confusing.

(2022). 9781936137749, National Science Teachers Association. . ISBN 9781936137749

(2022). 9780028282107, Glencoe McGraw-Hill. ISBN 9780028282107

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.

(2022). 9780840068460, Brooks/Cole, Cengage Learning. . ISBN 9780840068460

These mnemonics are commonly used by students to help memorise the terminology:

(2022). 9780547054056, Houghton Mifflin. . ISBN 9780547054056

• "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 which release stored electricity, and can be recharged with electricity. PANIC does not apply to cells that can be recharged with redox materials. These galvanic cell, such as , produce electricity from internal redox reactions. Here, the positive electrode is the cathode and the negative is the anode.

---

See also

• Anaerobic respiration
• Bessemer process
• Bioremediation
• Calvin cycle
• Chemical equation
• Chemical looping combustion
• Citric acid cycle
• Electrochemical series
• Electrochemistry
• Electrolysis
• Electron equivalent
• Electron transport chain
• Electrosynthesis
• Galvanic cell
• Hydrogenation
• Membrane potential
• Microbial fuel cell
• Murburn concept
• Nucleophilic abstraction
• Organic redox reaction
• Oxidative addition and reductive elimination
• Oxidative phosphorylation
• Partial oxidation
• Pro-oxidant
• Redox gradient
• Redox potential
• Reducing agent
• Reducing atmosphere
• Reduction potential
• Exothermic
• Transmetalation
• Sulfur cycle

---

Further reading

• (1999). 9783540665281, Springer-Verlag. ISBN 9783540665281
• (2022). 9780841226524 ISBN 9780841226524

---

External links

• Chemical Equation Balancer – An open-source chemical equation balancer that handles redox reactions.
• Redox reactions calculator
• Redox reactions at Chemguide
• Online redox reaction equation balancer, balances equations of any half-cell and full reactions

Categories: Soil Chemistry, Chemical Reactions, Articles Containing Video Clips, Redox, Reaction Mechanisms