The materials to prepare students for practical lessons of inorganic chemistry

The materials to prepare students for practical lessons of inorganic chemistry

LESSON № 11.

Theme. Reactions with transfer of electrons.

Plan

Reactions with transfer of electrons.

Redox-reactions. Oxidation-reduction properties. Electrolysis in fusions and solutions of electrolytes. Bases of electrochemistry.

Oxidation-reduction reactions. Oxidation-reduction properties of elements and their compounds. Oxidatioumber of element atom in compounds. Change of oxidatioumber in redox-reactions. Influence of environment (рН) on oxidation-reduction reactions (for example of КМnО₄). Role oxidation-reduction reactions on metabolism compound in human body.

 

REACTIONS WITH TRANSFER OF ELECTRONS.

Electron activity in chemical reactions

So far in our discussions on electricity and electric circuits, we have not discussed in any detail how batteries function. Rather, we have simply assumed that they produce constant voltage through some sort of mysterious process. Here, we will explore that process to some degree and cover some of the practical considerations involved with real batteries and their use in power systems.

In the first chapter of this book, the concept of an atom was discussed, as being the basic building-block of all material objects. Atoms, in turn, are composed of even smaller pieces of matter called particles. Electrons, protons, and neutrons are the basic types of particles found in atoms. Each of these particle types plays a distinct role in the behavior of an atom. While electrical activity involves the motion of electrons, the chemical identity of an atom (which largely determines how conductive the material will be) is determined by the number of protons in the nucleus (center).

The protons in an atom’s nucleus are extremely difficult to dislodge, and so the chemical identity of any atom is very stable. One of the goals of the ancient alchemists (to turn lead into gold) was foiled by this sub-atomic stability. All efforts to alter this property of an atom by means of heat, light, or friction were met with failure. The electrons of an atom, however, are much more easily dislodged. As we have already seen, friction is one way in which electrons can be transferred from one atom to another (glass and silk, wax and wool), and so is heat (generating voltage by heating a junction of dissimilar metals, as in the case of thermocouples).

Electrons can do much more than just move around and between atoms: they can also serve to link different atoms together. This linking of atoms by electrons is called a chemical bond. A crude (and simplified) representation of such a bond between two atoms might look like this:

There are several types of chemical bonds, the one shown above being representative of a covalent bond, where electrons are shared between atoms. Because chemical bonds are based on links formed by electrons, these bonds are only as strong as the immobility of the electrons forming them. That is to say, chemical bonds can be created or broken by the same forces that force electrons to move: heat, light, friction, etc.

When atoms are joined by chemical bonds, they form materials with unique properties known as molecules. The dual-atom picture shown above is an example of a simple molecule formed by two atoms of the same type. Most molecules are unions of different types of atoms. Even molecules formed by atoms of the same type can have radically different physical properties. Take the element carbon, for instance: in one form, graphite, carbon atoms link together to form flat “plates” which slide against one another very easily, giving graphite its natural lubricating properties. In another form, diamond, the same carbon atoms link together in a different configuration, this time in the shapes of interlocking pyramids, forming a material of exceeding hardness. In yet another form, Fullerene, dozens of carbon atoms form each molecule, which looks something like a soccer ball. Fullerene molecules are very fragile and lightweight. The airy soot formed by excessively rich combustion of acetylene gas (as in the initial ignition of an oxy-acetylene welding/cutting torch) contains many Fullerene molecules.

When alchemists succeeded in changing the properties of a substance by heat, light, friction, or mixture with other substances, they were really observing changes in the types of molecules formed by atoms breaking and forming bonds with other atoms. Chemistry is the modern counterpart to alchemy, and concerns itself primarily with the properties of these chemical bonds and the reactions associated with them.

A type of chemical bond of particular interest to our study of batteries is the so-called ionic bond, and it differs from the covalent bond in that one atom of the molecule possesses an excess of electrons while another atom lacks electrons, the bonds between them being a result of the electrostatic attraction between the two unlike charges. When ionic bonds are formed from neutral atoms, there is a transfer of electrons between the positively and negatively charged atoms. An atom that gains an excess of electrons is said to be reduced; an atom with a deficiency of electrons is said to be oxidized. A mnemonic to help remember the definitions is OIL RIG (oxidized is less; reduced is gained). It is important to note that molecules will often contain both ionic and covalent bonds. Sodium hydroxide (lye, NaOH) has an ionic bond between the sodium atom (positive) and the hydroxyl ion (negative). The hydroxyl ion has a covalent bond (shown as a bar) between the hydrogen and oxygen atoms:

Na⁺ O—H^–

Sodium only loses one electron, so its charge is +1 in the above example. If an atom loses more than one electron, the resulting charge can be indicated as +2, +3, +4, etc. or by a Romaumeral in parentheses showing the oxidation state, such as (I), (II), (IV), etc. Some atoms can have multiple oxidation states, and it is sometimes important to include the oxidation state in the molecular formula to avoid ambiguity.

The formation of ions and ionic bonds from neutral atoms or molecules (or vice versa) involves the transfer of electrons. That transfer of electrons can be harnessed to generate an electric current. A device constructed to do just this is called a voltaic cell, or cell for short, usually consisting of two metal electrodes immersed in a chemical mixture (called an electrolyte) designed to facilitate such an electrochemical (oxidation/reduction) reaction:

In the common “lead-acid” cell (the kind commonly used in automobiles), the negative electrode is made of lead (Pb) and the positive is made of lead (IV) dioxide (Pb0₂), both metallic substances. It is important to note that lead dioxide is metallic and is an electrical conductor, unlike other metal oxides that are usually insulators. The electrolyte solution is a dilute sulfuric acid (H₂SO₄ + H₂O). If the electrodes of the cell are connected to an external circuit, such that electrons have a place to flow from one to the other, lead(IV) atoms in the positive electrode (PbO₂) will gain two electrons each to produce Pb(II)O. The oxygen atoms which are “left over” combine with positively charged hydrogen ions (H)⁺to form water (H₂O). This flow of electrons into into the lead dioxide (PbO₂) electrode, gives it a positive electrical charge. Consequently, lead atoms in the negative electrode give up two electrons each to produce lead Pb(II), which combines with sulfate ions (SO₄^-2) produced from the disassociation of the hydrogen ions (H⁺) from the sulfuric acid (H₂SO₄) to form lead sulfate (PbSO₄). The flow of electrons out of the lead electrode gives it a negative electrical charge. These reactions are shown diagrammitically below:

Note on lead oxide nomenclature

The nomenclature for lead oxides can be confusing. The term, lead oxide can refer to either Pb(II)O or Pb(IV)O₂, and the correct compound can be determined usually from context. Other synonyms for Pb(IV)O₂ are: lead dioxide, lead peroxide, plumbic oxide, lead oxide brown, and lead superoxide. The term, lead peroxide is particularly confusing, as it implies a compound of lead (II) with two oxygen atoms, Pb(II)O₂, which apparently does not exist. Unfortunately, the term lead peroxide has persisted in industrial literature. In this section, lead dioxide will be used to refer to Pb(IV)O₂, and lead oxide will refer to Pb(II)O. The oxidation states will not be shown usually.

This process of the cell providing electrical energy to supply a load is called discharging, since it is depleting its internal chemical reserves. Theoretically, after all of the sulfuric acid has been exhausted, the result will be two electrodes of lead sulfate (PbSO₄) and an electrolyte solution of pure water (H₂O), leaving no more capacity for additional ionic bonding. In this state, the cell is said to be fully discharged. In a lead-acid cell, the state of charge can be determined by an analysis of acid strength. This is easily accomplished with a device called a hydrometer, which measures the specific gravity (density) of the electrolyte. Sulfuric acid is denser than water, so the greater the charge of a cell, the greater the acid concentration, and thus a denser electrolyte solution.

There is no single chemical reaction representative of all voltaic cells, so any detailed discussion of chemistry is bound to have limited application. The important thing to understand is that electrons are motivated to and/or from the cell’s electrodes via ionic reactions between the electrode molecules and the electrolyte molecules. The reaction is enabled when there is an external path for electric current, and ceases when that path is broken.

Being that the motivation for electrons to move through a cell is chemical iature, the amount of voltage (electromotive force) generated by any cell will be specific to the particular chemical reaction for that cell type. For instance, the lead-acid cell just described has a nominal voltage of 2.04 volts per cell, based on a fully “charged” cell (acid concentration strong) in good physical condition. There are other types of cells with different specific voltage outputs. The Edison cell, for example, with a positive electrode made of nickel oxide, a negative electrode made of iron, and an electrolyte solution of potassium hydroxide (a caustic, not acid, substance) generates a nominal voltage of only 1.2 volts, due to the specific differences in chemical reaction with those electrode and electrolyte substances.

The chemical reactions of some types of cells can be reversed by forcing electric current backwards through the cell (in the negative electrode and out the positive electrode). This process is called charging. Any such (rechargeable) cell is called a secondary cell. A cell whose chemistry cannot be reversed by a reverse current is called a primary cell.

When a lead-acid cell is charged by an external current source, the chemical reactions experienced during discharge are reversed:

REVIEW:

·                     Atoms bound together by electrons are called molecules.

·                     Ionic bonds are molecular unions formed when an electron-deficient atom (a positive ion) joins with an electron-excessive atom (a negative ion).

·                     Electrochemical reactions involve the transfer of electrons between atoms. This transfer can be harnessed to form an electric current.

·                     A cell is a device constructed to harness such chemical reactions to generate electric current.

·                     A cell is said to be discharged when its internal chemical reserves have been depleted through use.

·                     A secondary cell’s chemistry can be reversed (recharged) by forcing current backwards through it.

·                     A primary cell cannot be practically recharged.

·                     Lead-acid cell charge can be assessed with an instrument called a hydrometer, which measures the density of the electrolyte liquid. The denser the electrolyte, the stronger the acid concentration, and the greater charge state of the cell.

Electron transfer (ET) occurs when an electron moves from an atom or a chemical species (e.g. a molecule) to another atom or chemical species. ET is a mechanistic description of the thermodynamic concept of redox, wherein the oxidation states of both reaction partners change.

Numerous biological processes involve ET reactions. These processes include oxygen binding, photosynthesis, respiration, and detoxification. Additionally, the process of energy transfer can be formalized as a two-electron exchange (two concurrent ET events in opposite directions) in case of small distances between the transferring molecules. ET reactions commonly involve transition metal complexes, but there are now many examples of ET in organic chemistry.

Outer sphere refers to an electron transfer (ET) event that occurs between chemical species that remain separate intact before, during, and after the ET event. In contrast, for inner sphere electron transfer the participating redox sites undergoing ET become connected by a chemical bridge. Because the ET in outer sphere electron transfer occurs between two non-connected species, the electron is forced to move through space from one redox center to the other.

Marcus Theory

The main theory that describes the rates of outer-sphere electron transfer was developed by Rudolph A. Marcus in the 1950s. A major aspect of Marcus theory is the dependence of the electron transfer rate on the thermodynamic driving force (difference in the redox potentials of the electron-exchanging sites). For most reactions, the rates increase with increased driving force. A second aspect is that the rate of outer-sphere electron-transfer depends inversely on the “reorganizational energy.” Reorganization energy describes the changes in bond lengths and angles that are required for the oxidant and reductant to switch their oxidation states. This energy is assessed by measurements of the self-exchange rates (see below). Outer sphere electron transfer is the most common type of electron transfer, especially in biochemistry, where redox centers are separated by several (up to about 11) angstroms by intervening protein. In biochemistry, there are two main types of outer sphere ET: ET between two biological molecules or fixed distance electron transfer, in which the electron transfers within a single biomolecule (e.g., intraprotein).

Examples

Self-exchange

Outer sphere electron transfer can occur between chemical species that are identical save for their oxidation state. This process is termed self-exchange. An example is the degenerate reaction between the tetrahedral ions permanganate and manganate:

[MnO₄]^– + [Mn*O₄]^2- → [MnO₄]^2- + [Mn*O₄]^–

For octahedral metal complexes, the rate constant for self-exchange reactions correlates with changes the population of the e_g orbitals, the population of which most strongly affects the length of metal-ligand bonds:

·                     For the [Co(bipy)₃]⁺/[Co(bipy)₃]²⁺ pair, self exchange proceeds at 10⁹ M^-1s^-1. In this case, the electron configuration changes from Co(I): (t_2g)⁶(e_g)² to Co(II): (t_2g)⁵(e_g)².

·                     For the [Co[bipy)₃]²⁺/[Co(bipy)₃]³⁺ pair, self exchange proceeds at 18 M^-1s^-1. In this case, the electron configuration changes from Co(II): (t_2g)⁵(e_g)² to Co(III): (t_2g)⁶(e_g)⁰.

Iron-sulfur proteins

Outer sphere ET is the basis of the biological function of the iron-sulfur proteins. The Fe centers are typically further coordinated by cysteinyl ligands. The [Fe₄S₄] electron-transfer proteins ([Fe₄S₄] ferredoxins) may be further subdivided into low-potential (bacterial-type) and high-potential (HiPIP) ferredoxins. Low- and high-potential ferredoxins are related by the following redox scheme:

Because of the small structural differences between the individual redox states, ET is rapid between these clusters.

Inner sphere or bonded electron transfer is a redox chemical reaction that proceeds via a covalent linkage—a strong electronic interaction—between the oxidant and the reductant reactants. In Inner Sphere (IS) electron transfer (ET), a ligand bridges the two metal redox centers during the electron transfer event. Inner sphere reactions are inhibited by large ligands, which prevent the formation of the crucial bridged intermediate. Thus, IS ET is rare in biological systems, where redox sites are often shielded by bulky proteins. Inner sphere ET is usually used to describe reactions involving transition metal complexes and most of this article is written from this perspective. However, redox centers can consist of organic groups rather than metal centers.

The bridging ligand could be virtually any entity that can convey electrons. Typically, such a ligand has more than one lone electron pair, such that it can serve as an electron donor to both the reductant and the oxidant. Common bridging ligands include the halides and the pseudohalides such as hydroxide and thiocyanate. More complex bridging ligands are also well known including oxalate, malonate, and pyrazine. Prior to ET, the bridged complex must form, and such processes are often highly reversible. Electron transfer occurs through the bridge once it is established. In some cases, the stable bridged structure may exist in the ground state; in other cases, the bridged structure may be a transiently-formed intermediate, or else as a transition state during the reaction.

The alternative to inner sphere electron transfer is outer sphere electron transfer. In any transition metal redox process, the mechanism can be assumed to be outer sphere unless the conditions of the inner sphere are met. Inner sphere electron transfer is generally enthalpically more favorable than outer sphere electron transfer due to a larger degree of interaction between the metal centers involved, however, inner sphere electron transfer is usually entropically less favorable since the two sites involved must become more ordered (come together via a bridge) than in outer sphere electron transfer.

Oxidation, any chemical reaction in which an atom of an element loses one or more of its electrons to an atom of a different element. Originally, the term was applied only to a reaction in which oxygen combines with another element or group of elements to form a compound called an oxide.

Familiar examples of oxidation include the burning of coal, which is rapid oxidation; and the rusting of iron, which is slow oxidation. Oxidation also occurs in animals and plants in the process of respiration. The carbon dioxide exhaled from the lungs is formed during this process.

The element whose atoms lose electrons during oxidation is said to be oxidized. The other element, whose atoms gain electrons, is called the oxidizing agent. The oxidizing agent is said to be reduced, and the process of gaining electrons is called reduction. Oxidation and reduction always occur simultaneously, and therefore chemists often use the term oxidation-reduction (or redox) when referring to this type of reaction.

The oxygen in the air serves as the oxidizing agent in most oxidation-reduction reactions. When iron rusts, for example, oxygen in the air combines chemically with the iron to form a coating of iron oxide on the surface of the iron. In this reaction, the iron atoms give up electrons to the oxygen atoms. The iron is oxidized and the oxygen is reduced.

The halogen elements (fluorine, chlorine, bromine, iodine, and astatine) also serve as oxidizing agents. When chlorine combines with sodium to form sodium chloride, or common salt, the sodium atoms give up electrons to the chlorine atoms. In this reaction, the sodium is oxidized and the chlorine is reduced.

“Redox” is a portmanteau of “reduction” and “oxidation”.

The word oxidation originally implied reaction with oxygen to form an oxide, since (di)oxygen was historically the first recognized oxidizing agent. Later, the term was expanded to encompass oxygen-like substances that accomplished parallel chemical reactions. Ultimately, the meaning was generalized to include all processes involving loss of electrons.

The word reduction originally referred to the loss in weight upon heating a metallic ore such as a metal oxide to extract the metal. In other words, ore was “reduced” to metal. Antoine Lavoisier (1743-1794) showed that this loss of weight was due to the loss of oxygen as a gas. Later, scientists realized that the metal atom gains electrons in this process. The meaning of reduction then became generalized to include all processes involving gain of electrons. Even though “reduction” seems counter-intuitive when speaking of the gain of electrons, it might help to think of reduction as the loss of oxygen, which was its historical meaning.

The electrochemist John Bockris has used the words electronation and deelectronation to describe reduction and oxidation processes respectively when they occur at electrodes. These words are analogous to protonation and deprotonation, but they have not been widely adopted by chemists.

The term “hydrogenation” could be used instead of reduction, since hydrogen is the reducing agent in a large number of reactions, especially in organic chemistry and biochemistry. But unlike oxidation, which has been generalized beyond its root element, hydrogenation has maintained its specific connection to reactions that add hydrogen to another substance (e.g., the hydrogenation of unsaturated fats in saturated fats, R-CH=CH-R + H₂ → R-CH₂-CH₂-R).

 

REDUCTION-OXIDATION REACTIONS

The Process of Discovery: Oxidation and Reduction

The first step toward a theory of chemical reactions was taken by Georg Ernst Stahl in 1697 when he proposed the phlogiston theory, which was based on the following observations.

• Metals have many properties in common.

• Metals often produce a “calx” when heated. (The term calx is defined as the crumbly residue left after a mineral or metal is roasted.)

• These calxes are not as dense as the metals from which they are produced.

• Some of these calxes form metals when heated with charcoal.

• With only a few exceptions, the calx is found iature, not the metal.

These observations led Stahl to the following conclusions.

• Phlogiston (from the Greek phlogistos, “to burn”) is given off whenever something burns.

• Wood and charcoal are particularly rich in phlogiston because they leave very little ash when they burn. (Candles must be almost pure phlogiston because they leave no ash.)

• Because they are found iature, calxes must be simpler than metals.

• Metals form a calx by giving off phlogiston.

Metal = calx + phlogiston

• Metals can be made by adding phlogiston to the calx.

Calx + phlogiston = metal

• Because charcoal is rich in phlogiston, heating calxes in the presence of charcoal sometimes produces metals.

This model was remarkably successful. It explained why metals have similar properties they all contained phlogiston. It explained the relationship between metals and their calxes they were related by the gain or loss of phlogiston. It even explained why a candle goes out when placed in a bell jar  the air eventually becomes saturated with phlogiston.

There was only one problem with the phlogiston theory. As early as 1630, Jean Rey noted that tin gains weight when it forms a calx. (The calx is about 25% heavier than the metal.) From our point of view, this seems to be a fatal flaw: If phlogiston is given off when a metal forms a calx, why does the calx weigh more than the metal? This observation didn’t bother proponents of the phlogiston theory. Stahl explained it by suggesting that the weight increased because air entered the metal to fill the vacuum left after the phlogiston escaped.

The phlogiston theory was the basis for research in chemistry for most of the 18th century. It was not until 1772 that Antoine Lavoisier noted that nonmetals gain large amounts of weight when burned in air. (The weight of phosphorus, for example, increases by a factor of about 2.3.) The magnitude of this change led Lavoisier to conclude that phosphorus must combine with something in air when it burns. This conclusion was reinforced by the observation that the volume of air decreases by a factor of 1/5th when phosphorus burns in a limited amount of air.

Lavoisier proposed the name oxygene (literally, “acid-former”) for the substance absorbed from air when a compound burns. He chose this name because the products of the combustion of nonmetals such as phosphorus are acids when they dissolve in water.

Lavoisier’s oxygen theory of combustion was eventually accepted and chemists began to describe any reaction between an element or compound and oxygen as oxidation. The reaction between magnesium metal and oxygen, for example, involves the oxidation of magnesium.

2 Mg(s) + O₂(g) = 2 MgO(s)

By the turn of the 20th century, it seemed that all oxidation reactions had one thing in common oxidation always seemed to involve the loss of electrons. Chemists therefore developed a model for these reactions that focused on the transfer of electrons. Magnesium metal, for example, was thought to lose electrons to form Mg²⁺ ions when it reacted with oxygen. By convention, the element or compound that gained these electrons was said to undergo reduction. In this case, O₂ molecules were said to be reduced to form O^2- ions.

A classic demonstration of oxidation-reduction reactions involves placing a piece of copper wire into an aqueous solution of the Ag⁺ion. The reaction involves the net transfer of electrons from copper metal to Ag⁺ ions to produce whiskers of silver metal that grow out from the copper wire and Cu²⁺ ions.

Cu(s) + 2 Ag⁺(aq) = Cu²⁺(aq) + 2 Ag(s)

The Cu²⁺ ions formed in this reaction are responsible for the light-blue color of the solution. Their presence can be confirmed by adding ammonia to this solution to form the deep-blue Cu(NH₃)₄²⁺ complex ion.

Chemists eventually recognized that oxidation-reduction reactions don’t always involve the transfer of electrons. There is no change in the number of valence electrons on any of the atoms when CO₂ reacts with H₂, for example,

CO₂(g) + H₂(g) = CO(g) + H₂O(g)

as shown by the following Lewis structures:

Chemists therefore developed the concept of oxidatioumber to extend the idea of oxidation and reduction to reactions in which electrons are not really gained or lost. The most powerful model of oxidation-reduction reactions is based on the following definitions.

Oxidation involves an increase in the oxidatioumber of an atom.

Reduction occurs when the oxidation number of an atom decreases.

According to this model, CO₂ is reduced when it reacts with hydrogen because the oxidatioumber of the carbon decreases from +4 to +2. Hydrogen is oxidized in this reaction because its oxidatioumber increases from 0 to +1.

Oxidation-Reduction Reactions

We find examples of oxidation-reduction or redox reactions almost every time we analyze the reactions used as sources of either heat or work. Wheatural gas burns, for example, an oxidation-reduction reaction occurs that releases more than 800 kJ/mol of energy.

CH₄(g) + 2 O₂(g) = CO₂(g) + 2 H₂O(g)

Within our bodies, a sequence of oxidation-reduction reactions are used to burn sugars, such as glucose (C₆H₁₂O₆) and the fatty acids in the fats we eat.

We don’t have to restrict ourselves to reactions that can be used as a source of energy, however, to find examples of oxidation-reduction reactions. Silver metal, for example, is oxidized when it comes in contact with trace quantities of H₂S or SO₂ in the atmosphere or foods, such as eggs, that are rich in sulfur compounds.

4 Ag(s) + 2 H₂S(g) + O₂(g) = 2 Ag₂S(s) + 2 H₂O(g)

Fortunately, the film of Ag₂S that collects on the metal surface forms a protective coating that slows down further oxidation of the silver metal.

The tarnishing of silver is just one example of a broad class of oxidation-reduction reactions that fall under the general heading ofcorrosion. Another example is the series of reactions that occur when iron or steel rusts. When heated, iron reacts with oxygen to form a mixture of iron(II) and iron(III) oxides.

Molten iron even reacts with water to form an aqueous solution of Fe²⁺ ions and H₂ gas.

Fe(l) + 2 H₂O(l) = Fe²⁺(aq) + 2 OH^–(aq) + H₂(g)

At room temperature, however, all three of these reactions are so slow they can be ignored.

Iron only corrodes at room temperature in the presence of both oxygen and water. In the course of this reaction, the iron is oxidized to give a hydrated form of iron(II) oxide.

2 Fe(s) + O₂(aq) + 2 H₂O(l) = 2 FeO H₂O(s)

Because this compound has the same empirical formula as Fe(OH)₂, it is often mistakenly called iron(II), or ferrous, hydroxide. The FeO H₂O formed in this reaction is further oxidized by O₂ dissolved in water to give a hydrated form of iron(III), or ferric, oxide.

4 FeO H₂O(s) + O₂(aq) + 2 H₂O(l) = 2 Fe₂O₃ 3 H₂O(s)

To further complicate matters, FeO H₂O formed at the metal surface combines with Fe₂O₃ 3 H₂O to give a hydrated form of magnetic iron oxide (Fe₃O₄).

FeO H₂O(s) + Fe₂O₃ = 3 H₂O(s) Fe₃O₄ n H₂O(s)

Because these reactions only occur in the presence of both water and oxygen, cars tend to rust where water collects. Furthermore, because the simplest way of preventing iron from rusting is to coat the metal so that it doesn’t come in contact with water, cars were originally painted for only one reason   to slow down the formation of rust.

Redox (reduction-oxidation) reactions include all chemical reactions in which atoms have their oxidation state changed. This can be either a simple redox process, such as the oxidation of carbon to yield carbon dioxide (CO₂) or the reduction of carbon by hydrogen to yield methane (CH₄), or a complex process such as the oxidation of glucose (C₆H₁₂O₆) in the human body through a series of complex electron transfer processes.

Fundamentally, redox reactions are a family of reactions that are concerned with the transfer of electrons between species. The term comes from the two concepts of reduction and oxidation. It can be explained in simple terms:

·                     Oxidation is the loss of electrons or an increase in oxidation state by a molecule, atom, or ion.

·                     Reduction is the gain of electrons or a decrease in oxidation state by a molecule, atom, or ion.

Although oxidation reactions are commonly associated with the formation of oxides from oxygen molecules, these are only specific examples of a more general concept of reactions involving electron transfer.

The bridging ligand could be virtually any entity that can convey electrons. Typically, such a ligand has more than one lone electron pair, such that it can serve as an electron donor to both the reductant and the oxidant. Common bridging ligands include the halides and the pseudohalides such as hydroxide and thiocyanate. More complex bridging ligands are also well known including oxalate, malonate, and pyrazine. Prior to ET, the bridged complex must form, and such processes are often highly reversible. Electron transfer occurs through the bridge once it is established. In some cases, the stable bridged structure may exist in the ground state; in other cases, the bridged structure may be a transiently-formed intermediate, or else as a transition state during the reaction.

The alternative to inner sphere electron transfer is outer sphere electron transfer. In any transition metal redox process, the mechanism can be assumed to be outer sphere unless the conditions of the inner sphere are met. Inner sphere electron transfer is generally enthalpically more favorable than outer sphere electron transfer due to a larger degree of interaction between the metal centers involved, however, inner sphere electron transfer is usually entropically less favorable since the two sites involved must become more ordered (come together via a bridge) than in outer sphere electron transfer.

Oxidation, any chemical reaction in which an atom of an element loses one or more of its electrons to an atom of a different element. Originally, the term was applied only to a reaction in which oxygen combines with another element or group of elements to form a compound called an oxide.Redox reactions, or oxidation-reduction reactions, have a number of similarities to acid–base reactions. Like acid–base reactions, redox reactions are a matched set, that is, there cannot be an oxidation reaction without a reduction reaction happening simultaneously. The oxidation alone and the reduction alone are each called a half-reaction, because two half-reactions always occur together to form a whole reaction. When writing half-reactions, the gained or lost electrons are typically included explicitly in order that the half-reaction be balanced with respect to electric charge.

Though sufficient for many purposes, these descriptions are not precisely correct. Oxidation and reduction properly refer to a change in oxidation state — the actual transfer of electrons may never occur. Thus, oxidation is better defined as an increase in oxidation state, and reduction as a decrease in oxidation state. In practice, the transfer of electrons will always cause a change in oxidation state, but there are many reactions that are classed as “redox” even though no electron transfer occurs (such as those involving covalent bonds).

Redox reactions — reactions in which there’s a simultaneous transfer of electrons from one chemical species to another — are really composed of two different reactions: oxidation (a loss of electrons) and reduction (a gain of electrons).

The electrons that are lost in the oxidation reaction are the same electrons that are gained in the reduction reaction. These two reactions are commonly called half-reactions; the overall reaction is called a redox (reduction/oxidation) reaction.

Oxidation

There are three definitions you can use for oxidation:

·           The loss of electrons

·           The gain of oxygen

·           The loss of hydrogen

Loss of electrons

One way to define oxidation is with the reaction in which a chemical substance loses electrons in going from reactant to product. For example, when sodium metal reacts with chlorine gas to form sodium chloride (NaCl), the sodium metal loses an electron, which is then gained by chlorine.

The following equation shows sodium losing the electron:

Na (s) = Na⁺ + e

When it loses the electron, chemists say that the sodium metal has been oxidized to the sodium cation. (A cation is an ion with a positive charge due to the loss of electrons.)

Reactions of this type are quite common in electrochemical reactions, reactions that produce or use electricity.

Gain of oxygen

Sometimes, in certain oxidation reactions, it’s obvious that oxygen has been gained in going from reactant to product. Reactions where the gain of oxygen is more obvious than the gain of electrons include combustion reactions (burning) and the rusting of iron. Here are two examples.

Burning of coal:

C (s) + O₂ (g) = 2 Fe₂O₃

Rusting of iron:

2Fe(s) + 3 O₂ (g) = 2 Fe₂O₃ (s)

In these cases, chemists say that the carbon and the iron metal have been oxidized to carbon dioxide and rust, respectively.

Loss of hydrogen

In other reactions, oxidation can best be seen as the loss of hydrogen. Methyl alcohol (wood alcohol) can be oxidized to formaldehyde:

CH₃OH (l) = CH₂O (l) + H₂ (g)

In going from methanol to formaldehyde, the compound went from having four hydrogen atoms to having two hydrogen atoms.

Reduction

Like oxidation, there are three definitions you can use to describe reduction:

·  The gain of electrons

·  The loss of oxygen

·  The gain of hydrogen

Gain of electrons

Reduction is often seen as the gain of electrons. In the process of electroplating silver onto a teapot, for example, the silver cation is reduced to silver metal by the gain of an electron. The following equation shows the silver cation gaining the electron:

Ag+ + e = Ag

When it gains the electron, chemists say that the silver cation has been reduced to silver metal.

Loss of oxygen

In other reactions, it’s easier to see reduction as the loss of oxygen in going from reactant to product. For example, iron ore (primarily rust) is reduced to iron metal in a blast furnace by a reaction with carbon monoxide:

2Fe(s) + 3 CO (g) = 2 Fe  (s) + 3 CO₂ (g)

The iron has lost oxygen, so chemists say that the iron ion has been reduced to iron metal.

Gain of hydrogen

In certain cases, a reduction can also be described as the gain of hydrogen atoms in going from reactant to product. For example, carbon monoxide and hydrogen gas can be reduced to methyl alcohol:

CO(g) + 2 H₂ (g) = CH₃OH (l)

In this reduction process, the CO has gained the hydrogen atoms.

One’s loss is the other’s gain

Neither oxidation nor reduction can take place without the other. When those electrons are lost, something has to gain them.

Consider, for example, the net-ionic equation (the equation showing just the chemical substances that are changed during a reaction) for a reaction with zinc metal and an aqueous copper(II) sulfate solution:

Zn(s) + Cu²⁺  = Zn ²⁺+ Cu

This overall reaction is really composed of two half-reactions, shown below.

Oxidation half-reaction — the loss of electrons:

Zn(s) = Zn ²⁺+ 2e

Reduction half-reaction — the gain of electrons:

Cu2+ 2 e = Cu (s)

Zinc loses two electrons; the copper(II) cation gains those same two electrons. Zn is being oxidized. But without that copper cation (the oxidizing agent) present, nothing will happen. It’s a necessary agent for the oxidation process to proceed. The oxidizing agent accepts the electrons from the chemical species that is being oxidized.

The copper(II) cation is reduced as it gains electrons. The species that furnishes the electrons is called the reducing agent. In this case, the reducing agent is zinc metal.

The oxidizing agent is the species that’s being reduced, and the reducing agent is the species that’s being oxidized. Both the oxidizing and reducing agents are on the left (reactant) side of the redox equation.

 

Taube’s experiment

The discoverer of the inner sphere mechanism was Henry Taube, who was awarded the Nobel Prize in Chemistry in 1983 for his pioneering studies. A particularly historic finding is summarized in the abstract of the seminal publication. “When Co(NH₃)₅Cl⁺⁺ is reduced by Cr⁺⁺ in M {meaning 1M} HClO₄, 1 Cl⁻ appears attached to Cr for each Cr(III) which is formed or Co(III) reduced. When the reaction is carried on in a medium containing radioactive Cl, the mixing of the Cl⁻ attached to Cr(III) with that in solution is less than 0.5%. This experiment shows that transfer of Cl to the reducing agent from the oxidizing agent is direct…” The paper and the excerpt above can be described with the following equation:

[CoCl(NH₃)₅]²⁺ + [Cr(H₂O)₆]²⁺ → [Co(NH₃)₅(H₂O)]²⁺ + [CrCl(H₂O)₅]²⁺

The point of interest is that the chloride that was originally bonded to the cobalt, the oxidant, becomes bonded to chromium, which in its +3 oxidation state, forms kinetically inert bonds to its ligands. This observation implies the intermediacy of the bimetallic complex [Co(NH₃)₅(μ-Cl)Cr(H₂O)₅]⁴⁺, wherein “μ-Cl” indicates that the chloride bridges between the Cr and Co atoms, serving as a ligand for both. This chloride serves as a conduit for electron flow from Cr(II) to Co(III), forming Cr(III) and Co(II).

The Creutz-Taube ion

In the preceding example, the occurrence of the chloride bridge is inferred from the product analysis, but it was not observed. One complex that serves as a model for the bridged intermediate is the “Creutz Taube complex,” [(NH₃)₅RuNC₄H₄NRu(NH₃)₅]⁵⁺. This species is named after Carol Creutz, who prepared the ion during her PhD studies with Henry Taube. The bridging ligand is the heterocycle pyrazine, 1,4-C₄H₄N₂. In the Creutz-Taube Ion, the average oxidation state of Ru is 2.5+. Spectroscopic studies, however, show that the two Ru centers are equivalent, which indicates the ease with which the electron hole communicates between the two metals. The significance of the Creutz-Taube ion is its simplicity, which facilitates theoretical analysis, and its high symmetry, which ensures a high degree of delocalization. Many more complex mixed valence species are known both as molecules and polymeric materials.

Mixed valence compounds

Mixed valence compounds contain an element which is present in more than one oxidation state. Well-known mixed valence compounds include the Creutz-Taube complex, Prussian blue and Molybdenum blue. Many solids are mixed-valency including indium chalcogenides. Mixed valency is required for organic metals to exhibit electrical conductivity.

As the extinction coefficient decreases, the coupling constant decreases, influencing the angle to increase.

Mixed-valence compounds are subdivided into three groups, according to the Robin-Day Classification:

·            Class I, where the valences are “trapped,” or localized on a single site, such as Pb₃O₄ and antimony tetroxide. There are distinct sites with different specific valences in the complex that cannot easily interconvert.

·            Class II, which are intermediate in character. There is some localization of distinct valences, but there is a low activation energy for their interconversion. Some thermal activation is required to induce electron transfer from one site to another via the bridge. These species exhibit an intense Intervalence charge transfer (IT or IVCT) band, a broad intense absorption in the IR- or visible part of the spectrum, and also exhibit magnetic exchange coupling at low temperatures. The degree of interaction between the metal sites can be estimated from the absorption profile of the IVCT band and the spacing between the sites. This type of complex is common when metals are in different ligand fields. For example, Prussian blue is an iron(II,III)–cyanide complex in which there is an iron(II) atom surrounded by six carbon atoms of six cyanide ligands bridged to an iron(III) atom by their nitrogen ends. In the Turnbull’s blue preparation, an iron(II) solution is mixed with an iron(III) cyanide (c-linked) complex. An electron-transfer reaction occurs via the cyanide ligands to give iron(III) associated with an iron(II)-cyanide complex.

·            Class III, wherein mixed valence is not distinguishable by spectroscopic methods as the valence is completely delocalized. The Creutz-Taube Ion is an example of this class of complexes. These species also exhibit an IT band. Each site exhibits an intermediate oxidation state, which can be half-integer in value. This class is possible when the ligand environment is similar or identical for each of the two metal sites in the complex. The bridging ligand needs to be very good at electron transfer, be highly conjugated, and be easily reduced.

Etymology

“Redox” is a portmanteau of “reduction” and “oxidation”, more commonly known as oxidation-reduction.

The word oxidation originally implied reaction with oxygen to form an oxide, since (di)oxygen was historically the first recognized oxidizing agent. Later, the term was expanded to encompass oxygen-like substances that accomplished parallel chemical reactions. Ultimately, the meaning was generalized to include all processes involving loss of electrons.

The word reduction originally referred to the loss in weight upon heating a metallic ore such as a metal oxide to extract the metal. In other words, ore was “reduced” to metal. Antoine Lavoisier (1743-1794) showed that this loss of weight was due to the loss of oxygen as a gas. Later, scientists realized that the metal atom gains electrons in this process. The meaning of reduction then became generalized to include all processes involving gain of electrons. Even though “reduction” seems counter-intuitive when speaking of the gain of electrons, it might help to think of reduction as the loss of oxygen, which was its historical meaning.

The electrochemist John Bockris has used the words electronation and deelectronation to describe reduction and oxidation processes respectively when they occur at electrodes. These words are analogous to protonation and deprotonation, but they have not been widely adopted by chemists.

The term “hydrogenation” could be used instead of reduction, since hydrogen is the reducing agent in a large number of reactions, especially in organic chemistry and biochemistry. But unlike oxidation, which has been generalized beyond its root element, hydrogenation has maintained its specific connection to reactions which add hydrogen to another substance (e.g., the hydrogenation of unsaturated fats in saturated fats, R-CH=CH-R + H₂ → R-CH₂-CH₂-R).

Oxidizing and reducing agents

In redox processes, the reductant transfers electrons to the oxidant. Thus, in the reaction, the reductant or reducing agent loses electrons and is oxidized, and the oxidant or oxidizing agent gains electrons and is reduced. The pair of an oxidizing and reducing agent that are involved in a particular reaction is called a redox pair. A redox couple is a reducing species and its corresponding oxidized form, e.g., Fe²⁺/Fe³⁺.

Assigning Oxidation Numbers

The key to identifying oxidation-reduction reactions is recognizing when a chemical reaction leads to a change in the oxidatioumber of one or more atoms. It is therefore a good idea to take another look at the rules for assigning oxidation numbers. By definition, the oxidatioumber of an atom is equal to the charge that would be present on the atom if the compound was composed of ions. If we assume that CH₄contains C^4- and H⁺ ions, for example, the oxidatioumbers of the carbon and hydrogen atoms would be -4 and +1.

Note that it doesn’t matter whether the compound actually contains ions. The oxidatioumber is the charge an atom would have if the compound was ionic. The concept of oxidatioumber is nothing more than a bookkeeping system used to keep track of electrons in chemical reactions. This system is based on a series of rules, summarized in the table below.

 

Oxidizers

Substances that have the ability to oxidize other substances are said to be oxidative or oxidizing and are known as oxidizing agents, oxidants, or oxidizers. That is, the oxidant (oxidizing agent) removes electrons from another substance; i.e., it oxidizes other substances, and is thus itself reduced. And, because it “accepts” electrons, it is also called an electron acceptor.

Oxidants are usually chemical substances with elements in high oxidation states (e.g., H₂O₂, MnO⁻⁴, CrO₃, Cr₂O₂⁻⁷, OsO₄), or else highly electronegative elements (O₂, F₂, Cl₂, Br₂) that can gain extra electrons by oxidizing another substance.

Reducers

Substances that have the ability to reduce other substances are said to be reductive or reducing and are known as reducing agents, reductants, or reducers. The reductant (reducing agent) transfers electrons to another substance; i.e., it reduces others, and is thus itself oxidized. And, because it “donates” electrons, it is also called an electron donor. Electron donors can also form charge transfer complexes with electron acceptors.

Reductants in chemistry are very diverse. Electropositive elemental metals, such as lithium, sodium, magnesium, iron, zinc, and aluminium, are good reducing agents. These metals donate or give away electrons readily. Hydride transfer reagents, such as NaBH4 and LiAlH4, are widely used in organic chemistry, primarily in the reduction of carbonyl compounds to alcohols. Another method of reduction involves the use of hydrogen gas (H₂) with a palladium, platinum, or nickel catalyst. These catalytic reductions are used primarily in the reduction of carbon-carbon double or triple bonds.

Standard electrode potentials (reduction potentials)

Each half-reaction has a standard electrode potential (E⁰_cell), which is equal to the potential difference (or voltage) (E⁰_cell) 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: ½ H₂ → H⁺ + e^–.

The electrode potential of each half-reaction is also known as its reduction potential E⁰_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⁻ → ½ H₂ by definition, positive for oxidizing agents stronger than H⁺ (e.g., +2.866 V for F₂) and negative for oxidizing agents which are weaker than H⁺ (e.g. –0.763 V for Zn²⁺).

 For a redox reaction which takes place in a cell, the potential difference E⁰_cell = E⁰_cathode – E⁰_anode

Historically, however, the potential of the reaction at the anode was sometimes expressed as an oxidation potential, E⁰_ox = – E⁰. 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 E⁰_cell = E⁰_cathode + E⁰_{ox (anode)}

Examples of redox reactions

Illustration of a redox reaction

A good example is the reaction between hydrogen and fluorine in which hydrogen is being oxidized and fluorine is being reduced:

H₂ + F₂ → 2 HF

We can write this overall reaction as two half-reactions:

the oxidation reaction:

H₂ → 2 H⁺ + 2 e⁻

and the reduction reaction:

F₂ + 2 e⁻ → 2 F⁻

Analyzing each half-reaction in isolation can often make the overall chemical process clearer. Because there is no net change in charge during a redox reaction, the number of electrons in excess in the oxidation reaction must equal the number consumed by the reduction reaction (as shown above).

Elements, even in molecular form, always have an oxidation state of zero. In the first half-reaction, hydrogen is oxidized from an oxidation state of zero to an oxidation state of +1. In the second half-reaction, fluorine is reduced from an oxidation state of zero to an oxidation state of −1.

When adding the reactions together the electrons are canceled:

And the ions combine to form hydrogen fluoride:

2 H⁺ + 2 F⁻ → 2 HF

The overall reaction is:

H₂ + F₂ → 2 HF

Displacement reactions

Redox occurs in single displacement reactions or substitution reactions. The redox component of these types of reactions is the change of oxidation state (charge) on certain atoms, not the actual exchange of atoms in the compounds.

For example, in the reaction between iron and copper(II) sulfate solution:

Fe + CuSO₄ → FeSO₄ + Cu

The ionic equation for this reaction is:

Fe + Cu²⁺ → Fe²⁺ + Cu

As two half-equations, it is seen that the iron is oxidized:

Fe → Fe²⁺ + 2 e−

And the copper is reduced:

Cu²⁺ + 2 e− → Cu

Other examples

·                     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⁻

Overall equation:

2 Fe²⁺ + H₂O₂ + 2 H⁺ → 2 Fe³⁺ + 2 H₂O

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

2 NO₃⁻ + 10 e⁻ + 12 H⁺ → N₂ + 6 H₂O

Iron rusting in pyrite cubes

·                     Oxidation of elemental iron to iron(III) oxide by oxygen (commonly known as rusting, which is similar to tarnishing):

4 Fe + 3 O₂ → 2 Fe₂O₃

·       The combustion of hydrocarbons, such as in an internal combustion engine, which produces water, carbon dioxide, some partially oxidized forms such as carbon monoxide, and heat energy. Complete oxidation of materials containing carbon produces carbon dioxide.

·       In organic chemistry, the stepwise oxidation of a hydrocarbon by oxygen produces water and, successively, an alcohol, an aldehyde or a ketone, a carboxylic acid, and then a peroxide.

Redox reactions in industry

The primary process of reducing ore at high temperature to produce metals is known as smelting.

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

Redox reactions are the foundation of electrochemical cells.

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, and gold-plated jewelry.

The production of compact discs depends on a redox reaction, which coats the disc with a thin layer of metal film

 

Redox reactions in biology

Top: ascorbic acid (reduced form of Vitamin C)

Bottom: dehydroascorbic acid (oxidized form of Vitamin C)

Many important biological processes involve redox reactions.

Cellular respiration, for instance, is the oxidation of glucose (C₆H₁₂O₆) to CO₂ 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 + light energy → C₆H₁₂O₆ + 6 O₂

Biological energy is frequently stored and released by means of 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⁺), 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 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 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. An abnormal redox state can develop in a variety of deleterious situations, such as hypoxia, shock, and sepsis. Redox signaling involves the control of cellular processes by redox 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

A wide variety of aromatic compounds are enzymatically 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 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 futile cycle or redox cycling.

Examples of redox cycling-inducing molecules are the herbicide paraquat and other viologens and quinones such as menadione.

Redox reactions in geology

A uranium mine, near Moab, Utah. Note alternating red and white/green sandstone. This corresponds to oxidized and reduced conditions in groundwater redox chemistry. The rock forms in oxidizing conditions, and is then “bleached” to the white/green state when a reducing fluid passes through the rock. The reduced fluid can also carry uranium-bearing minerals.

In geology, redox is important to both the formation of minerals, mobilization of minerals, and in some depositional environments. In general, the redox state of most rocks can be seen in the color of the rock. Red is associated with oxidizing conditions of formation, and green is typically associated with reducing conditions. White can also be associated with reducing conditions. Famous examples of redox conditions affecting geological processes include uranium deposits and Moqui marbles.

Balancing redox reactions

Describing the overall electrochemical reaction for a redox process requires a balancing of the component half-reactions for oxidation and reduction. In general, for reactions in aqueous solution, this involves adding H⁺, OH⁻, H₂O, and electrons to compensate for the oxidation changes.

Acidic media

In acidic media, H⁺ ions and water are added to half reactions to balance the overall reaction.

For instance, when manganese(II) reacts with sodium bismuthate:

The reaction is balanced by scaling the two half-cell reactions to involve the same number of electrons (multiplying the oxidation reaction by the number of electrons in the reduction step and vice versa):

8 H₂O(l) + 2 Mn²⁺(aq) → 2 MnO−

4(aq) + 16 H⁺(aq) + 10 e−

10 e− + 30 H⁺ + 5 BiO−

3(s) → 5 Bi³⁺(aq) + 15 H₂O(l)

Adding these two reactions eliminates the electrons terms and yields the balanced reaction:

14 H⁺(aq) + 2 Mn²⁺(aq) + 5 NaBiO₃(s) → 7 H₂O(l) + 2 MnO−

4(aq) + 5 Bi³⁺(aq) + 5 Na⁺(aq)

Basic media

In basic media, OH⁻ ions and water are added to half reactions to balance the overall reaction.

For example, in the reaction between potassium permanganate and sodium sulfite:

Balancing the number of electrons in the two half-cell reactions gives:

6 e− + 4 H₂O + 2 MnO₄⁻ → 2 MnO₂ + 8 OH⁻

6 OH⁻ + 3 SO₃²⁻ → 3 SO₄²⁻ + 3 H₂O + 6 e−

Adding these two half-cell reactions together gives the balanced equation:

2 KMnO₄ + 3 Na₂SO₃ + H₂O → 2 MnO₂ + 3 Na₂SO₄ + 2 KOH

 

Redox reactions, or oxidation-reduction reactions have a number of similarities to acid-base reactions. Fundamentally, redox reactions are a family of reactions that are concerned with the transfer of electrons between species. Like acid-base reactions, redox reactions are a matched set — you don’t have an oxidation reaction without a reduction reaction happening at the same time. Oxidation refers to the loss of electrons, while reduction refers to the gain of electrons. Each reaction by itself is called a “half-reaction”, simply because we need two (2) half-reactions to form a whole reaction. Iotating redox reactions, chemists typically write out the electrons explicitly:

Cu (s) —-> Cu²⁺ + 2 e^–

This half-reaction says that we have solid copper (with no charge) being oxidized (losing electrons) to form a copper ion with a plus 2 charge. Notice that, like the stoichiometry notation, we have a “balance” between both sides of the reaction. We have one (1) copper atom on both sides, and the charges balance as well. The symbol “e^–” represents a free electron with a negative charge that caow go out and reduce some other species, such as in the half-reaction:

2 Ag⁺ (aq) + 2 e^– ——> 2 Ag (s)

Here, two silver ions (silver with a positive charge) are being reduced through the addition of two (2) electrons to form solid silver. The abbreviations “aq” and “s” mean aqueous and solid, respectively. We caow combine the two (2) half-reactions to form a redox equation:

We can also discuss the individual components of these reactions as follows. If a chemical causes another substance to be oxidized, we call it the oxidizing agent. In the equation above, Ag⁺ is the oxidizing agent, because it causes Cu(s) to lose electrons. Oxidants get reduced in the process by a reducing agent. Cu(s) is, naturally, the reducing agent in this case, as it causes Ag⁺ to gain electrons.

As a summary, here are the steps to follow to balance a redox equation in acidic medium (add the starred step in a basic medium):

1.                Divide the equation into an oxidation half-reaction and a reduction half-reaction

2.                Balance these

o         Balance the elements other than H and O

o         Balance the O by adding H₂O

o         Balance the H by adding H⁺

o         Balance the charge by adding e^–

3.                Multiply each half-reaction by an integer such that the number of e^– lost in one equals the number gained in the other

4.                Combine the half-reactions and cancel

5.                **Add OH^– to each side until all H⁺ is gone and then cancel again**

In considering redox reactions, you must have some sense of the oxidatioumber (ON) of the compound. The oxidatioumber is defined as the effective charge on an atom in a compound, calculated according to a prescribed set of rules. An increase in oxidation number corresponds to oxidation, and a decrease to reduction. The oxidation number of a compound has some analogy to the pH and pK measurements found in acids and bases — the oxidatioumber suggests the strength or tendency of the compound to be oxidized or reduced, to serve as an oxidizing agent or reducing agent. The rules are shown below. Go through them in the order given until you have an oxidatioumber assigned.

1.                For atoms in their elemental form, the oxidatioumber is 0

2.                For ions, the oxidatioumber is equal to their charge

3.                For single hydrogen, the number is usually +1 but in some cases it is -1

4.                For oxygen, the number is usually -2

5.                The sum of the oxidatioumber (ONs) of all the atoms in the molecule or ion is equal to its total charge.

As a side note, the term “oxidation”, with its obvious root from the word “oxygen”, assumes that oxygen has an oxidatioumber of -2. Using this as a benchmark, oxidatioumbers were assigned to all other elements. For example, if we look at H₂O, and assign the value of -2 to the oxygen atom, the hydrogens must each have an oxidatioumber of +1 by default, since water is a neutral molecule. As an example, what is the oxidatioumber of sulfur in sulfur dioxide (SO₂)? Given that each oxygen atom has a -2 charge, and knowing that the molecule is neutral, the oxidatioumber for sulfur must be +4. What about for a sulfate ion (SO₄ with a total charge of -2)? Again, the charge of all the oxygen atoms is 4 x -2 = -8. Sulfur must then have an oxidatioumber of +6, since +6 + (-8) = -2, the total charge on the ion. Since the sulfur in sulfate has a higher oxidatioumber than in sulfur dioxide, it is said to be more highly oxidized.

Working with redox reactions is fundamentally a bookkeeping issue. You need to be able to account for all of the electrons as they transfer from one species to another. There are a number of rules and tricks for balancing redox reactions, but basically they all boil down to dealing with each of the two half-reactions individually. Consider for example the reaction of aluminum metal to form alumina (Al₂O₃). The unbalanced reaction is as follows:

Looking at each half reaction separately:

This reaction shows aluminum metal being oxidized to form an aluminum ion with a +3 charge. The half-reaction below shows oxygen being reduced to form two (2) oxygen ions, each with a charge of -2.

If we combine those two (2) half-reactions, we must make the number of electrons equal on both sides. The number 12 is a common multiple of three (3) and four (4), so we multiply the aluminum reaction by four (4) and the oxygen reaction by three (3) to get 12 electrons on both sides. Now, simply combine the reactions. Notice that we have 12 electrons on both sides, which cancel out. The final step is to combine the aluminum and oxygen ions on the right side using a cross multiply technique:

Taking care of the number of atoms, you should end up with:

One of the more useful calculations in redox reactions is the Nernst Equation. This equation allows us to calculate the electric potential of a redox reaction in “non-standard” situations. There exist tables of how much voltage, or potential, a reaction is capable of producing or consuming. These tables, known as standard potential tables, are created by measuring potential at “standard” conditions, with a pressure of 1 bar (≅1 atm), a temperature of 298° K (or 25° C, or room temperature) and with a concentration of 1.0 M for each of the products. This standard potential, or E°, can be corrected by a factor that includes the actual temperature of the reaction, the number of moles of electrons being transferred, and the concentrations of the redox reactants and products. The equation is:

Perhaps the best way of understanding this equation is through an example. Suppose we have this reaction:

Fe(s) + Cd²⁺(aq) ——> Fe²⁺(aq) + Cd(s)

In this reaction iron (Fe) is being oxidized to iron(II) ion, while the cadmium ion (Cd²⁺) in aqueous solution is being reduced to cadmium solid. The question is: how does this reaction behave in “non-standard” conditions?

The first thing to answer is how does it behave in standard conditions? We need to look at the standard potential for each half-reaction, then combine them to get a net potential for the reaction. The two (2) half-reactions are:

Fe²⁺ (aq) + 2 e^– ——> Fe (s), E° = -0.44 V

Cd²⁺ (aq) +2 e^– ——> Cd (s), E° = -0.40 V

Notice that both half-reactions are shown as reductions — the species gains electrons, and is changed to a new form. But in the complete reaction above, Fe is oxidized, so the half-reaction needs to be reversed. Quite simply, the potential for the half-reaction of iron is now 0.44 V. To get the potential for the entire reaction, we add up the two (2) half-reactions to get 0.04 V for the standard potential.

The questioow is: what is the total potential (in volts) for a nonstandard reaction? Suppose again that we have the same reaction, except now we have 0.0100 M Fe²⁺ instead of the standard 1.0 M. We need to use the Nernst equation to help us calculate that value. If you go to the Redox Half-Reaction Calculator, you should notice that the reaction is selected and the appropriate values are entered into the boxes. Since we don’t have any species “B” or “D”, we have entered zero for their concentrations. The concentration of the solid Fe is 1.0 M (actually, concentrations of solids and solvents (liquids) don’t enter into the Nernst equation, but we set them to 1.0 so that the mathematics works out). If you click on the “Evaluate” button, you should learn that the standard potential is -0.44 V, while the nonstandard potential is -0.5 V. If you scroll down on the calculator, you can enter 0.5 as the first half-reaction. We again change the sign since we’re actually reversing the Fe reaction

Using the calculator again, we calculate the nonstandard potential of the Cd reaction. Suppose we now have a concentration of Cd²⁺ of 0.005 M, what is its potential? The calculator should return a standard potential of -0.4 V and a nonstandard potential of -0.47 V. Place this value in the box for the second half-reaction, then click on “Evaluate”. You should learn that the net nonstandard potential is 0.03 V, slightly less than the value of the net standard potential. Since this value is less than the net standard potential of 0.04 V, there is less of a tendency for this reaction to transfer electrons from reactants to products. In other words, less iron will be oxidized and cadmium will be reduced than at standard conditions.

Test your use of the redox calculator by calculating the net standard potential for this reaction:

2 Ag⁺ (aq, 0.80 M) + Hg (l)——> 2 Ag (s) + Hg²⁺ (aq, 0.0010M)

Answer: 0.025 V. Since the value is positive, the reaction will work to form the products indicated. Negative values of the potential indicate that the reaction tends to stay as reactants and not form the products. The net standard potential for this reaction is 0.01 V — since the nonstandard potential is higher, this reaction will form products than the standard reaction.

Free energy and the standard potential can also be related through the following equation:

Where:

ΔG = change in free energy n = number of moles

If a reaction is spontaneous, it will have a positive E^o, and negative ΔG, and a large K value (where K is the equilibrium constant-this is discussed more in the kinetics section).

The energy released in any spontaneous redox reaction can be used to perform electrical work using an electrochemical cell (a device where electron transfer is forced to take an external pathway instead of going directly between the reactants. Think of the reaction between zinc and copper. Instead of placing a piece of zinc directly into a solution containing copper, we can form a cell where solid pieces of zinc and copper are placed in two different solutions such as sodium nitrate. The two solids are called electrodes. The anode is the electrode where oxidation occurs and mass is lost where as the cathode is the electrode where reduction occurs and mass is gained. The two electrodes are connected by a circuit and the two (2) solutions are connected by a “salt bridge” which allows ions to pass through. The anions are the negative ions and they move towards the anode. The cations are the positive ions and they move towards the cathode.

The following is a diagram of an electrochemical cell with zinc and copper acting as the electrodes.

An external electric current hooked up to an electrochemical cell will make the electrons go backwards. This process is called electrolysis. This is used, for example, to make something gold plated. You would put the copper in a solution with gold and add a current which causes the gold ions to bond to the copper and therefore coating the copper. The time, current, and electrons needed determine how much “coating” occurs. The key to solving electolysis problems is learning how to convert between the units. Useful information: 1 A=1 C/sec; 96,500 coulombs can produce one (1) mole of e^–; the electrons needed is determined by the charge of the ion involved

Example Problem: If you are trying to coat a strip with aluminum and you have a current of 10.0 A (amperes) running for one hour, what mass of Al is formed?

The solution of this problem involves a lengthly unit conversion process:

 

The oxidatioumber corresponds to the number of electrons, e^–, that an atom loses, gains, or appears to use when joining with other atoms in compounds. When determining the Oxidation State of an atom there are seven guidelines to follow:

1.                The Oxidation State of an individual atom is 0.

2.                The total Oxidation State of all atoms in: a neutral species is 0 and in an ion is equal to the ion charge.

3.                Group 1 metals have an Oxidation State of +1 and group 2 an Oxidation State of +2

4.                The Oxidation State of fluorine is -1, when in compounds

5.                Hydrogen generally has an Oxidation State of +1 in compounds

6.                Oxygen generally has an Oxidation State of -2 in compounds

7.                In binary metal compounds, group 17 elements have an Oxidation State of -1, group 16 of -2, and group 15 of -3.

(Note: The sum of the oxidation states is equal to zero for neutral compounds and equal to the charge for polyatomic ion species.)

Sample Problems: Determine the oxidation states:

1.                Fe(s) + O₂(g) → Fe₂O₃(g)

2.                Fe²⁺

3.                Ag(s) + H₂S → Ag₂S(g) + H₂(g)

Solutions

1.                Fe and O₂ are free elements, therefore they have an O.S. of “0” according to Rule #1. The product has a total O.S. equal to “0” and following Rule #6, O₃ has an O.S. of -2, which means Fe₂ has an O.S. of +2.

2.                The O.S. of Fe corresponds to its charge, therefore the O.S. is +2.

3.                Ag has an O.S. of 0, H₂ has an O.S. of +1 according to Rule #5 and S has an O.S. of -2 according to Rule #7.

Example 1: Determine the oxidation state of the bold element in each of the following:

1.                Na₃PO₃

2.                H₂PO₄^–

Answers to Example 1:

1.                The oxidatioumbers of Na and O are +1 and -2. Since sodium phosphite is neutral, the sum of the oxidatioumbers must be zero.. Letting x be the oxidatioumber of phosphorus then, 0= 3(+1) + x + 3(-2). x=oxidation number of P= +3. 

2.                Hydrogen and oxygen have oxidatioumbers of +1 and -2. The ion has a charge of -1, so the sum of the oxidatioumbers must be -1. Letting y be the oxidatioumber of phosphorus, -1= y + 2(+1) +4(-2), y= oxidatioumber of P= +5.

Example 2: Determine which element is oxidized and which element is reduced in the following reactions (be sure to include the oxidation state of each): 

1.                Zn + 2H⁺ → Zn²⁺ + H₂

2.                2Al + 3Cu²⁺→2Al³⁺ +3Cu

3.                CO₃^2- + 2H⁺→ CO₂ + H₂O

Answers to Example 2:

1.                Zn is oxidized (Oxidatioumber: 0 → +2); H⁺ is reduced (Oxidatioumber: +1 → 0)

2.                Al is oxidized (Oxidatioumber: 0 → +3); Cu²⁺ is reduced (+2 → 0)

3.                This is not a redox type because each element has the same oxidatioumber in both reactants and products: O= -2, H= +1, C= +4.

 

Oxidizing and Reducing Agents

An atom is oxidized when it oxidation number increases, the reducing agent, and an atom is reduced when its oxidation number decreases, the oxidizing agent. In other words, what is oxidized is the reducing agent and what is reduced is the oxidizing agent. (Note: the oxidizing and reducing agents can be the same element or compound).

Oxidation-Reduction Reactions

Redox reactions are comprised of two parts, a reduced half and an oxidized half, that always occur together. The reduced half gains electrons and the oxidatioumber decreases, while the oxidized half losses electrons and the oxidatioumber increases. Simple ways to remember this are the mnemonic devices OIL RIG meaning “oxidation is loss” and “reduction is gain” or LEO says GER meaning “loss of e^– = oxidation” and “gain of e^– = reduced.” There is no net change in the number of electrons in a redox reaction. Those given off in the oxidation half reaction are taken on by another species in the reduction half reaction.

The two species that exchange electrons in a redox reaction are given special names. The ion or molecule that accepts electrons is called the oxidizing agent; by accepting electrons it brings about the oxidation of another species. Conversely, the species that donates electrons is called the reducing agent; when reaction occurs it reduces the other species. In other words, what is oxidized is the reducing agent and what is reduced is the oxidizing agent. (Note: the oxidizing and reducing agents can be the same element or compound This will be further discussed under Types of Redox Reactions: Disproportionation).

A good example of a redox reaction is the thermite reaction in which iron atoms of ferric oxide lose (or give up) O atoms to Al atoms, producing Al₂O₃.

Fe₂O₃(s) + 2Al(s) → Al₂O₃(s) + 2Fe(l)

Another example of the redox reaction is the reaction between Zinc and Copper sulfate.

Example 3. Using the equations from the previous examples determine what is oxidized?

Zn + 2H⁺ → Zn²⁺ + H₂

Answer:

The O.S. of H goes from +1 to 0 and the O.S. of Zn goes from 0 to 2+. Hence, Zn is oxidized and acts as the reducing agent.

Example 4. What is reduced?

Zn + 2H⁺ → Zn²⁺ + H₂

Answer:

The O.S. of H goes from +1 to 0 and the O.S. of Zn goes from 0 to 2+. Hence, H⁺ ion is reduced and acts as the oxidizing agent.

Half Reactions

Before one can balance an overall redox equation, one has to be able to balance two half-equations, one for oxidation (electron loss) and one for reduction (electron gain). Collectively, oxidation and reduction are known as redox, or an electron transfer reaction. After balancing the two half-equations one can determine the total net reaction.

Each equation is balanced by adjusting coefficients and adding H₂O, H⁺, and e^– in this order:

1) Balance the number of atoms of each element. 

2) Balance the number of electrons transferred.

3) Balance the total charge on reactants and products

(Note: If #1 and #2 are done correctly, #3 will follow. Thus, it serves as a means of checking your work).

 

 BALANCING OF OXIDATION-REDUCTION REACTIONS

Predicting the directions of oxidation-reduction reactions

To solve redox reactions accurately, you must first understand how to balance chemical equations. Though this process is more difficult thaormal balancing it is a required step in the process of redox reactions. One of the most accepted methods of balancing a redox reaction is known as the half-equation method, however it can become more complex when involving basic or acidic solutions. In this module, a brief introduction to this different method will be explored.

In oxidation-reduction (“redox”) reactions, electrons are transferred from a donor (reducing agent) to an acceptor (oxidizing agent.) But how can you predict whether, or in which direction, such a reaction will actually go? We present here a very simple way of understanding how different redox reactions are related. It is essentially the same as we did for acid-base reactions, in which we saw how protons fall from higher-energy sources (acids) to lower-energy sinks (bases). Similarly, electrons-transfer reactions spontaneously proceed in the direction in which electrons fall (in free energy) from sources (reducing agents) to sinks (oxidizing agents.)

Here is a simple example: will the reaction

Zn + Cu²⁺= Zn²⁺ + Cu

go in the forward or reverse direction? (Assume equal effective concentrations of the two ions to avoid favoring one or the other direction.) We make use of an electron free energy diagram, the relevant portion of which is shown here:

 

Because electrons have a higher energy on Zn than they do on Cu, copper ions will serve as an electron sink to Zn, and the reaction will go to the right: the Zn gets oxidized to Zn²⁺, and the Cu²⁺ is reduced to metallic copper.

After you have digested this example, take a look at the following two pages:

Falling through the activity series of the elements

Falling through the respiratory ladder

If you already know some electrochemistry, you probably know how to use the Nernst Equation to carry out quantitative calculations. Nevertheless, this is still a very helpful picture when you have to deal with multiple redox systems which occur very commonly in environmental chemistsry, analytical chemistry, and biochemistry.

In molecular substances, we use these rules to give the approximate charges on the atoms. Consider the molecule SO . Oxygen atoms tend to attract electrons, pulling them from other atoms (sulfur in the case of SO 2 ). As a result, an oxygen atom in SO₂ takes on a negative charge relative to the sulfur atom. The magnitude of the charge on an oxygen atom in a molecule is not a full -2 charge as in the O ion. However, it is convenient to assign an oxidatioumber of -2 to oxygen in SO (and in most other compounds of oxygen) to help us express the approximate charge distribution in the molecule.

Rule 3 in Table 4.5 says that an oxygen atom has an oxidatioumber of -2 in most of its compounds. Rules 4 and 5 are similar in that they tell you what to expect for the oxidatioumber of certain elements in their compounds.

Rule 4, for instance, says that hydrogen has an oxidatioumber of -1 in most of its compounds.

Rule 6 states that the sum of the oxidatioumbers of the atoms in a compound is zero. This rule follows from the interpretation of oxidatioumbers as (hypothetical) charges on the atoms. Because any compound is electrically neutral, the sum of the charges on its atoms must be zero. This rule is easily extended to ions: the sum of the oxidatioumbers (hypothetical charges) of the atoms in a polyatomic ion equals the charge on the ion. You can use Rule 6 to obtain the oxidatioumber of one atom in a compound or ion, if you know the oxidatioumbers of the other atoms in the compound or ion. Consider the SO 2

Oxidation is the half-reaction in which there is a loss of electrons by a species (or an increase of oxidatioumber of an atom). Reduction is the half-reaction in which there is a gain of electrons by a species (or a decrease in the oxidatioumber of an atom). Thus, the equation represents the oxidation half-reaction, and the equation represents the reduction half-reaction. Recall that a species that is oxidized loses electrons (or contains an atom that increases in oxidatioumber) and a species that is reduced gains electrons (or contains an atom that decreases in oxidatioumber). An oxidizing agent is a species that oxidizes another species; it is itself reduced. Similarly, a reducing agent is a species that reduces another species; it is itself oxidized. In our example reaction, the copper(II) ion is the oxidizing agent, whereas iron metal is the reducing agent. The relationships among these terms are shown in the following diagram for the reaction of iron with copper(II) ion.

Some Common Oxidation–Reduction Reactions

Many oxidation–reduction reactions can be described as one of the following:

1. Combination reaction

2. Decomposition reaction

3. Displacement reaction

4. Combustion reaction

 

Half-Equation Method

The half-equation method (for neutral reactions) involves three basic steps which are as follows:

·                     Write and balance the half reactions.

·                     Adjust coefficients in both equations so that the same number of electrons appears in each half.

·                     Add together both halves, canceling out electrons, to obtain the overall equation.

Balancing in Basic and Acidic Solution

Balancing in acidic solution is similar to balancing ieutral solutions however, instead of three steps to follow, there are six. These rules are:

·                     Write and balance the half reactions.

·                     Balance oxygen, O, by adding with H₂O

·                     Balance hydrogen, H, by adding H⁺ (acidic)

·                     Balance charge by adding electrons (you should be adding the same number of electrons as H⁺ ions)

·                     Multiply both half reactions by some integer to cancel out electrons

·                     Add the half reactions together and cancel out what appears on both sides

Example:

Balance the redox reaction in acidic solution: MnO₄^– + I^– –> Mn²⁺ + I_2(s)

Write and balance the half reactions:

       MnO₄^– + I^– –> Mn²⁺ + I_2(s)

_{O.S:  +7 -2       -1          +2          0}        (Mn is reduced and I- is oxidized)

Oxidation Rx:2I^–_(aq) –> I_2(s) + 2e^–

Reduction Rx: MnO₄^– + 5e^– –> Mn²⁺

Balance oxygen, O, by adding H₂O

Oxidation Rx: 2I^–_(aq) –> I_2(s) + 2e^–

Reduction Rx: MnO₄^– + 5e^– –> Mn²⁺ + 4H₂O

Balance hydrogen, H, by adding H⁺

Oxidation Rx: 2I^–_(aq) –> I_2(s) + 2e^–

Reduction Rx: MnO₄^– + 5e^– + 8H⁺ –> Mn²⁺ + 4H₂O

Balance charge by adding electrons

Oxidation Rx: 2I^–_(aq) –> I_2(s) + 2e^–

Reduction Rx: MnO₄^– + 5e^– + 8H⁺ –> Mn²⁺ + 4H₂O

Multiply both half reactions by some integer to cancel out electrons

(Oxidation Rx: 2I^–_(aq) –> I_2(s) + 2e^–) * 5

(Reduction Rx: MnO₄^– + 5e^– + 8H⁺ –> Mn²⁺ + 4H₂O) *2

Oxidation Rx: 10I^–_(aq) –> 5I_2(s) + 10e^–

Reduction Rx: 2MnO₄^– + 10e^– + 16H⁺ –> 2Mn²⁺ + 8H₂O

Add the half reactions together and cancel out what appears on both sides:

10I^–_(aq) + 2MnO₄^–_(aq) + 16H⁺_(aq) –> 2Mn²⁺_(aq) + 5I_2(s) + 8H₂O_(l)

(Note: Don’t forget the states of matter! Generally, anything with a charge is _(aq) and H2O is _(l))

Balancing in basic solution follows balancing in acidic solutions in three steps:

·                     Balance the reaction in acidic solution

·                     Add the same amount of OH^– ions as H⁺ ions to both sides of the equation. On one side, the OH^– and H⁺ will react to form water (H2O) in a 1:1 ratio.

·                     Cancel out water molecules appearing on both sides

Example:

Balance the above redox reaction in basic solution:

Balance the reaction in acidic solution

10I^–_(aq) + 2MnO₄^–_(aq) + 16H⁺_(aq) –> 2Mn²⁺_(aq) + 5I_2(s) + 8H₂O_(l)

Add the same amount of OH^– ions as H⁺ ions to both sides of the equation.

10I^–_(aq) + 2MnO₄^–_(aq) + 16H⁺_(aq) + 16OH^– –> 2Mn²⁺_(aq) + 5I_2(s) + 8H₂O_(l) + 16OH^–

On one side, the OH^– and H⁺ will react to form water (H2O) in a 1:1 ratio.

10I^–_(aq) + 2MnO₄^–_(aq) + 16H₂O –> 2Mn²⁺_(aq) + 5I_2(s) + 8H₂O_(l) + 16OH^–

Cancel out the water molecules appearing on both sides

10I^–_(aq) + 2MnO₄^–_(aq) + 8H₂O_(l) –> 2Mn²⁺_(aq) + 5I_2(s) + 16OH^–_(aq)

Example:

Balance the following half-equation:

(1) MnO₄^– → Mn²⁺

Answers:

(1) a. Because there is one atom of Mn on both sides, no adjustment is required.

b. Because manganese is reduced from an oxidatioumber of +7 to +2, five electrons must be added to the left (MnO₄^– + 5e^– → Mn²⁺)

c. There is a total charge of -6 on the left versus +2 on the right. To balance, add eight H⁺ to the left to give a charge of +2 on both sides. (MnO₄^– + 8H⁺ + 5e^– → Mn²⁺)

d. To balance the eight H⁺ ions on the left, add four H₂O molecules to the right. MnO₄^– + 8H⁺ + 5e^– → Mn²⁺ + 4H₂O

e. Note that there are the same number of oxygen atoms, four, on both sides, as there should be. The equation shown in green is the correctly balanced reduction half-equation.

Types of Redox Reactions:

Combination

Combination reactions are some of the simplest redox reactions and as the name suggests involves the “combining” of elements to form a chemical compound. As usual, oxidation and reduction occur together. General Equation:

A + B → AB

Sample 1.

Equation: H₂ + O₂ → H₂O

Calculation: 0 + 0 → (2)(+1) + (-2) = 0

Explanation: In this equation both H₂ and O₂ are free elements and following Rule 1, their oxidation state is “0.” The product is H₂O, which has a total oxidation state of “0.” According to Rule #6, the O.S. of oxygen is usually -2. So, the O.S. of H₂ must be +1.

Decomposition

General Equation: AB → A + B

Decomposition reactions are the reverse of combination reactions, meaning they are the breakdown of a chemical compound into the individual elements.

Sample 2.

Equation: H₂O → H₂ + O₂

Calculation: (2)(+1) + (-2) = 0 → 0 + 0

Explanation: In this equation the water is “decomposed” into a Hydrogen and Oxygen. Similar to the previous sample the H₂O has a total oxidation state of “0,” thus according to Rule#6 the O.S. of oxygen is usually -2 so the O.S. of H₂ must be +1.

Displacement Reactions

Displacement reactions, also known as replacement reactions, involve compounds and the “replacing” of elements. They occur as single replacement and double replacement reactions.

Single Replacement

General Equation: A + BC → AB + C

A single replacement reaction involves the “replacing” of an element in the reactants with another element in the products.

Sample 3.

Equation: Cl₂ + NaBr → NaCl + Br₂

Calculation: (0) + ((+1) + (-1) = 0) -> ((+1) + (-1) = 0) + 0

Explanation: In this equation Br is replaced with Cl and Cl is reduced, while Br is oxidized.

Double Replacement

General Equation: AB + CD → AD + CB

A double replacement reaction is similar to a double replacement reaction, but involves “replacing” two elements in the reactants, with two in the products.

Sample 4.

Equation: Fe₂O₃ + HCl → FeCl₃ + H₂O

Explanation: In this equation Fe and H trade places and oxygen and chlorine trade places.

Combustion

Combustion reactions always involve oxygen, in the form of O₂ and are almost always exothermic, meaning they produce heat.

General Equation: C_xH_y + O₂ → CO₂ + H₂O

Disproportionation

General Equation: 2A → A’ + A”

In some redox reactions substances can be both oxidized and reduced. These are known as disproportionation reactions, which have some practical significance in everyday life including the reaction of hydrogen peroxide, H₂O₂ poured over a cut. This a decomposition reaction of hydrogen peroxide, which produces oxygen and water. Oxygen is present in all parts of the chemical equation and as a result it is both oxidized and reduced.

Reaction: 2H₂O₂(aq) → 2H₂O(l) + O₂(g)

Explanation: In the reactants H has an O has an O.S. of -1, which changes to -2 for the product, H₂O (reduced) and 0 for the product, O₂ (oxidized).

 

Example: 

Balance the following equations in both acidic and basic environments:

1) H₂(g) + O₂(g)  → H₂O(l)

Acidic Answer: 2H₂(g) + O₂(g) → 2H₂O(l)

Basic Answer: 2H₂(g) + O₂(g) → 2H₂O(l)

2) Cr₂O₇^2-(aq) + C₂H₅OH(l) → Cr³⁺(aq) + CO₂(g)

Acidic Answer: 2Cr₂O₇^2-(aq) + C₂H₅OH(l) + 16H⁺(aq) → 4Cr³⁺(aq) + 2CO₂(g) + 11H₂O(l)

Basic Answer: 2Cr₂O₇^2-(aq) + C₂H₅OH(l) + 5 H₂O(l) → 4Cr³⁺(aq) + 2CO₂(g) + 16OH^–(aq)

3) Fe²⁺(aq) + MnO₄^–(aq) → Fe³⁺(aq) + Mn²⁺(aq)

Acidic Answer: MnO₄^–(aq) + 5Fe²⁺(aq) + 8H⁺(aq) → Mn²⁺(aq) + 5Fe³⁺(aq) + 4H₂O(l)

Basic Answer: MnO₄^–(aq) + 5Fe²⁺(aq) + 4H₂O(l) → Mn²⁺(aq) + 5Fe³⁺(aq) + 8OH^–(aq)

4) Zn(s) + NO^3-(aq) → Zn²⁺(aq) + NO(g)

Acidic Answer: 3Zn(s) + 2NO₃^–(aq) + 8H⁺(aq) → 3Zn²⁺(aq) + 2NO(g) + 4H₂O(l)

Basic Answer: 3Zn(s) + 2NO₃^–(aq) + 4H₂O(l) → 3Zn²⁺(aq) + 2NO(g) + 8OH^–(aq)

5) Al(s) + H₂O(l) + O₂(g) → [Al(OH)₄]^–(aq)

Acidic Answer: 4Al(s) + 3O₂(g) + 10 H₂O(l) → 4[Al(OH)₄]^–(aq) + 4H⁺(aq)

Basic Answer: 4Al(s) + 3O₂(g) + 6H₂O(l) + 4OH^–(aq) → 4[Al(OH)₄]^–(aq)

 

Example: 

Copper metal added to concentrated nitric acid

Cu_(s)  +  HNO_3(aq) —–>  Cu(NO₃)_2(aq)  +  NO_2(g)

a. first divide the equation into half-reactions (notice that the copper is going from a 0 oxidation state to 2+ which is oxidation, and some of the nitrogen is being reduced from 5+ in the nitrate ion to 4+ in the nitrogen dioxide).

Oxidation:  Cu_(s) —–>  Cu²⁺_(aq)  +  2e^–

Reduction:  NO₃^1-_(aq)  +  e^– —–>  NO_2(g)

b. balance each half-reaction with respect to atoms first, then with respect to electrons.

Oxidation:  Cu_(s) —–>  Cu²⁺_(aq)  +  2e^–

Reduction:  2NO₃^1-_(aq)  +  2 e^– —–>  2NO_2(g))

Notice that this time there are oxygen atoms in the reduction step that are not balanced. When this happens, add as many water molecules on the right as are needed to balance the total oxygens on the left. Then add hydrogen ions on the left side of the arrow to balance the number of hydrogen atoms that were introduced with the water molecules.

Oxidation:  Cu_(s) —–>   Cu²⁺_(aq)  +  2e^–

Reduction:  2NO₃^1-_(aq)  +  2 e^–  +  4H¹⁺_(aq) —–>  2NO_2(g))  +  2H₂O_(l)

c. add the two half-reactions together cancelling the electrons which are now equal on each side of the arrow.

Cu_(s)  +  2NO₃^1-_(aq)  +  4H¹⁺_(aq) —–>  Cu²⁺_(aq)  +  2NO_2(g))  +   2H₂O_(l)

Note that this is a Net Ionic Equation for the reaction. The other two nitrate ions that would be with the four hydrogen ions in the nitric acid would remain unchanged and provide the nitrate ions that would form copper(II) nitrate.

 

Kinetics of Electron Transfer

In this section we will develop a quantitative model for the influence of the electrode voltage on the rate of electron transfer. For simplicity we will consider a single electron transfer reaction between two species (O) and (R)

The current flowing in either the reductive or oxidative steps can be predicted using the following expressions

For the reduction reaction the current (i_c) is related to the electrode area (A), the surface concentration of the reactant [O]_o, the rate constant for the electron transfer (k_Red or k_Ox) and Faraday’s constant (F). A similar expression is valid for the oxidation, now the current is labelled (i_a), with the surface concentration that of the species R. Similarly the rate constant for electron transfer corresponds to that of the oxidation process. Note that by definition the reductive current is negative and the oxidative positive, the difference in sign simply tells us that current flows in opposite directions across the interface depending upon whether we are studying an oxidation or reduction.

To establish how the rate constants k_Ox and k_Red are influenced by the applied voltage we will use transition state theory from chemical kinetics. You will recall that in this theory the reaction is considered to proceed via an energy barrier. The summit of this barrier is referred to as the transition state.

The rate of reaction for a chemical process (eg)

is predicted by an equation of the form

where the term in the exponential is the free energy change in taking the reactant from its initial value to the transition state divided by the temperature and gas constant. This free energy plot is also qualitatively valid for electrode reactions

where the free energy plot below corresponds to the thermodynamic response at a single fixed voltage.

Using this picture the activation free energy for the reduction and oxidation reactions are

and so the corresponding reaction rates are given by

So for a single applied voltage the free energy profiles appear qualitatively to be the same as corresponding chemical processes. However if we now plot a series of these free energy profiles as a function of voltage it is apparent that the plots alter as a function of the voltage. It is important to note that the left handside of the figure corresponding to the free energy of R is invariant with voltage, whereas the right handside ( O + e) shows a strong dependence.

At voltage V₁ the formation of the species O is thermodynamically favoured. However as we move through the voltages to V₆ the formation of R becomes the thermodynamically favoured product. This can be explained in terms of the Fermi level diagrams noted earlier, as the voltage is altered the Fermi level is raised (or lowered) changing the energy state of the electrons. However it is not just the thermodynamic aspects of the reaction that can be influenced by this voltage change as the overall barrier height (ie activation energy) can also be seen to alter as a function of the applied voltage. We might therefore predict that the rate constants for the forward and reverse reactions will be altered by the applied voltage. In order to formulate a model we will assume that the effect of voltage on the free energy change will follow a linear relationship (this is undoubtedly an over simplification). Using this linear relationship the activation free energies for reduction and oxidation will vary as a function of the applied voltage (V) as follows

 

 

The parameter is called the transfer coefficient and typically is found to have a value of 0.5. Physically it provides an insight into the way the transition state is influenced by the voltage. A value of one half means that the transition state behaves mid way between the reactants and products response to applied voltage. The free energy on the right hand side of both of the above equations can be considered as the chemical component of the activation free energy change, ie it is only dependent upon the chemical species and not the applied voltage. We caow substitute the activation free energy terms above into the expressions for the oxidation and reduction rate constants, which gives

These results show us the that rate constants for the electron transfer steps are proportional to the exponential of the applied voltage. So the rate of electrolysis can be changed simply by varying the applied voltage. This result provides the fundamental basis of the experimental technique called voltammetry which we will look at more closely later.

In conclusion we have seen that the rate of electron transfer can be influenced by the applied voltage and it is found experimentally that this behaviour can be quantified well using the simple model presented above. However the kinetics of the electron transfer is not the only process which can control the electrolysis reaction. In many circumstances it is the rate of transport to the electrode which controls the overall reaction.

 

ROLE OXIDATION-REDUCTION REACTIONS ON METABOLISM COMPOUND IN HUMAN BODY

 

Oxidation-reduction reactions are of central importance in organic chemistry and biochemistry. The burning of fuels that provides the energy to maintain our civilization and the metabolism of foods that furnish the energy that keeps us alive both involve redox reactions.

The burning of natural gas is not only a combustion reaction but also a redox reaction. Similar reactions include the burning of gasoline and coal. These are also redox reactions.

All combustion reactions are also redox reactions. A typical combustion reaction is the burning of methane, the principal component of natural gas

CH₄+ 2O₂→ CO₂+ 2H₂O

In respiration, the biochemical process by which the oxygen we inhale in air oxidizes foodstuffs to carbon dioxide and water, redox reactions provide energy to living cells. A typical respiratory reaction is the oxidation of glucose (C₆H₁₂O₆), the simple sugar we encountered in the chapter-opening essay that makes up the diet of yeast:

C₆H₁₂O₆+ 6O₂→ 6CO₂+ 6H₂O

Organic chemists use a variety of redox reactions. For example, potassium dichromate (K₂Cr₂O₇) is a common oxidizing agent that can be used to oxidize alcohols (symbolized by the general formula ROH). The product of the reaction depends on the location of the OH functional group in the alcohol molecule, the relative proportions of alcohol and the dichromate ion, and reaction conditions such as temperature. If the OH group is attached to a terminal carbon atom and the product is distilled off as it forms, the product is an aldehyde, which has a terminal carbonyl group(C=O) and is often written as RCHO. One example is the reaction used by the Breathalyzer to detect ethyl alcohol (C₂H₅OH) in a person’s breath:

3C₂H₅OH + Cr₂O₇²⁻+ 8H⁺ → 3CH₃CHO + 2Cr³⁺+ 7H₂O

If the product acetaldehyde (CH₃CHO) is not removed as it forms, it is further oxidized to acetic acid (CH₃COOH). In this case, the overall reaction is as follows:

3C₂H₅OH + 2Cr₂O₇²⁻+ 16H⁺ → 3CH₃COOH + 4Cr³⁺+ 11H₂O

In this reaction, the chromium atom is reduced from Cr₂O₇²⁻ to Cr³⁺, and the ethanol is oxidized to acetic acid.

When the OH group of the alcohol is bonded to an interior carbon atom, the oxidation of an alcohol will produce a ketone. (The formulas of ketones are often written as RCOR, and the carbon–oxygen bond is a double bond.) The simplest ketone is derived from 2-propanol (CH₃CHOHCH₃). It is the common solvent acetone [(CH₃)₂CO], which is used in varnishes, lacquers, rubber cement, and nail polish remover. Acetone can be formed by the following redox reaction:

3CH₃CHOHCH₃+ Cr₂O₇²⁻+ 8H⁺→ 3(CH₃)₂CO + 2Cr³⁺+ 7H₂O

As we have just seen, aldehydes and ketones can be formed by the oxidation of alcohols. Conversely, aldehydes and ketones can be reduced to alcohols. Reduction of the carbonyl group is important in living organisms. For example, in anaerobic metabolism, in which biochemical processes take place in the absence of oxygen, pyruvic acid (CH₃COCOOH) is reduced to lactic acid (CH₃CHOHCOOH) in the muscles.

CH₃COCOOH → CH₃CHOHCOOH

(Pyruvic acid is both a carboxylic acid and a ketone; only the ketone group is reduced.) The buildup of lactic acid during vigorous exercise is responsible in large part for the fatigue that we experience.

In food chemistry, the substances known as antioxidants are reducing agents. Ascorbic acid (vitamin C; C₆H₈O₆) is thought to retard potentially damaging oxidation of living cells. In the process, it is oxidized to dehydroascorbic acid (C₆H₆O₆). In the stomach, ascorbic acid reduces the nitrite ion (NO₂⁻) to nitric oxide (NO):

C₆H₈O₆+ 2H⁺+ 2NO₂⁻→ C₆H₆O₆+ 2H₂O + 2NO

If this reaction did not occur, nitrite ions from foods would oxidize the iron in hemoglobin, destroying its ability to carry oxygen.

Tocopherol (vitamin E) is also an antioxidant. In the body, vitamin E is thought to act by scavenging harmful by-products of metabolism, such as the highly reactive molecular fragments called free radicals. In foods, vitamin E acts to prevent fats from being oxidized and thus becoming rancid. Vitamin C is also a good antioxidant

Citrus fruits, such as oranges, lemons, and limes, are good sources of vitamin C, which is an antioxidant.

Finally, and of greatest importance, green plants carry out the redox reaction that makes possible almost all life on Earth. They do this through a process called photosynthesis, in which carbon dioxide and water are converted to glucose (C₆H₁₂O₆). The synthesis of glucose requires a variety of proteins called enzymes and a green pigment called chlorophyll that converts sunlight into chemical energy. The overall change that occurs is as follows:

6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂

Photosynthesis is the fundamental process by which plants use sunlight to convert carbon dioxide and water into glucose and oxygen. Then plants make more complex carbohydrates. It is the ultimate source of all food on Earth, and it is a redox reaction.

Practice Problem 1:

Determine the oxidatioumber of each element in the following compounds:

(a) BaO₂     (b) (NH₄)₂MoO₄      (c)Na₃Co(NO₂)₆      (d) CS₂

Answer

(a)  Ba is +2 and O is -1.

(b) H is +1, N is -3, Mo is -6, and O is -2.

(c) N is +3, O is -2, and Co is +3.

(d) S is -2 and C is +4.

Solution

(a) If the oxidatioumber of the oxygen in BaO₂ were -2, the oxidatioumber of the barium would have to be +4.  But elements in Group IIA can’t form +4 ions.  This compound must be barium peroxide, [Ba²⁺][O₂^2-].   Barium therefore is +2 and oxygen is -1.

(b) (NH₄)₂MoO₄ contains the NH⁴⁺ ion, in which hydrogen is +1 and nitrogen is -3.  Because there are two NH⁴⁺ ions, the other half of the compound must be an MoO₄^2- ion, in which molybdenum is -6 and oxygen is -2.

(c) Sodium is in the +1 oxidation state in all of its compounds.  This compound therefore contains the Co(NO₂)₆^3- complex ion.  This complex ion contains six NO^2- ions in which the oxidatioumber of nitrogen is +3 and oxygen is -2.  The oxidation state of the cobalt atom is therefore +3.

(d) The most elecronegative element in a compound always has a negative oxidation number.  Since sulfur tends to form -2 ions, the oxidatioumber of the sulfur in CS₂, is -2 and the carbon is +4.

Practice Problem 2:

Classify each of the following as either a metathesis or an oxidation-reduction reaction.  Note that mercury usually exists in one of three oxidation states: mercury metal, Hg₂²⁺ ions, or Hg²⁺ ions.

(a) Hg₂²⁺(aq) + 2 OH^–(aq)   Hg₂O(s) + H₂O(l)

(b) Hg₂²⁺(aq)  + Sn²⁺(aq)     2 Hg(l) + Sn⁴⁺(aq)

(c) Hg₂²⁺(aq) + H₂S(aq)     Hg(l) + HgS(s) + 2 H⁺(aq)

(d) Hg₂CrO₄(s) + 2 OH^–(aq)   Hg₂O(s) + CrO₄^2-(aq) + H₂O(l)

Answer

(a) Metathesis.  Mercury is in the +1 oxidation state in both the Hg₂²⁺ ion and in [Hg₂²⁺][O^2-].

(b) Oxidation-reduction.  Mercury is reduced from the +1 to the 0 oxidation state, while tin is oxidized from +2 to +4.

(c) Oxidation-reduction.  This is a disproportionation reaction (see chapter 10) in which mercury is simultaneously reduced from +1 to 0 and oxidized from +1 to +2.

(d) Metathesis.  Mercury is in the +1 oxidation state in both [Hg₂²⁺][CrO4^2-] and [Hg₂²⁺][O^2-].

 

 

 

References:

1. The abstract of the lecture.

2. intranet.tdmu.edu.ua/auth.php

3. Atkins P. W. Physical chemistry / P.W. Atkins. – New York, 1994. – P.299‑307.

4. Cotton F. A. Chemical Applications of Group Theory / F. A. Cotton. ‑ John Wiley & Sons : New York, 1990.

5. Girolami G. S. Synthesis and Technique in Inorganic Chemistry / G. S. Girolami, T. B. Rauchfuss, R. J. Angelici. ‑ University Science Books : Mill Valley, CA, 1999.

6. Russell J. B. General chemistry / J B. Russell. New York.1992. – P. 550‑599.

7. Lawrence D. D. Analytical chemistry / D. D. Lawrence. –New York, 1992. – P. 218–224.

8. http://www.lsbu.ac.uk/water/ionish.html

 

Prepared by PhD Falfushynska H.