Part I - Basic Concepts
Introduction
The rusting of metals, the process involved in photography, the way living systems produce and utilize energy, and the operation of a car battery, are but a few examples of a very common and important type of chemical reaction. These chemical changes are all classified as "electron-transfer" or oxidation-reduction reactions.
The term, oxidation , was derived from the observation that almost all elements reacted with oxygen to form compounds called, oxides. A typical example is the corrosion or rusting of iron as described by the chemical equation:
4 Fe + 3 O₂ -----> 2 Fe₂O₃
Reduction, was the term originally used to describe the removal of oxygen from metal ores, which "reduced" the metal ore to pure metal as shown below:
2 Fe₂O₃ + 3 C -----> 3 CO₂ + 4 Fe
Based on the two examples above, oxidation can be defined very simply as, the "addition" of oxygen; and reduction, as the "removal" of oxygen. But there is a lot more to "oxidation-reduction", as described in the following sections.

Oxidation States / Oxidation Numbers

The logical starting point in the discussion of oxidation-reduction reactions is the atom, and the terms and conventions used by chemists in describing this phenomenon.
All atoms are electrically neutral even though they are comprised of charged, subatomic particles. The terms, oxidation state or oxidation number, have been developed to describe this "electrical state" of the atom. The oxidation state or oxidation number of an atom is simply defined as the sum of the negative and positive charges in an atom. Since every atom contains equal numbers of positive and negative charges, the oxidation state or oxidation number of any atom is always zero. This is illustrated by simply totaling the opposite charges of the atoms as shown by the following examples.

Note, in every instance the sum of the positive and negative charges is zero; hence the oxidation state of any atom is always zero.
Oxidation- Reducation Reactions:
A Basic Modal

Assigning all atoms an oxidation state of zero serves as an important reference point, as oxidation-reduction reactions always involve a change in the oxidation state of the atoms or ions involved. This change in oxidation state is due to the "loss" or "gain" of electrons. The loss of electrons from an atom produces a positive oxidation state, while the gain of electrons results in negative oxidation states.
The changes that occur in the oxidation state of certain elements can be predicted quickly and accurately by the use of simple guidelines. These guidelines are based on the behavior of the Representative Elements, which can be divided into two classes; the metals and nonmetals.
All metal atoms are characterized by their tendency to be oxidized, losing one or more electrons, forming a positively charged ion, called a cation. During this oxidation reaction , the oxidation state of the metal always increases from zero to a positive number, such as "+1, +2, +3...." , depending on the number of electrons lost. The number of electrons lost by these Representative metals and the charge of the cation formed are always equal to the Group number of the metal as summarized below.

Group Number	Number of Electrons Lost	Charge of Cation Formed
I	1	+1
II	2	+2
III	3	+3
IV	4	+4

The group numbers also correspond to the electrons that are found in the outermost energy levels of these atoms. These electrons are often called valence electrons.
By convention oxidation reactions are written in the following form using the element, Calcium, as an example

Note that the oxidation state increases from zero to a positive number (from "0" to "+2" in the above example) and is always numerically equal to the number of electrons lost. See Exercise 1
The electrons lost by the metal are not destroyed but gained by the nonmetal, which is said to be reduced. As the nonmetal gains the electrons lost by the metal, it forms a negatively charged ion, called an anion. During this reduction reaction, the oxidation state of the nonmetal always decreases from zero to a negative value (-1, -2, -3 ...) depending on the number of electrons gained. The number of electrons gained by any Representative nonmetal and the charge of the anion formed, can be predicted by use of the following guidelines.

Group Number	Number of Electrons Gained	Charge of Anion Formed
IV	4	-4
V	3	-3
VI	2	-2
VII	1	-1
VIII	0	no tendency to form anions

Note, the GROUP VIII nonmetals have no tendency to gain additional electrons, hence they are unreactive in terms of oxidation-reduction. This is one the reasons why this family of elements was originally called the Inert Gases.
By convention reduction reactions are written in the following way:

Note that the charge of anion formed is always numerically equal to the number of electrons gained. See exercise 2
One important fact to remember in studying oxidation-reduction reactions is that the process of oxidation cannot occur without a corresponding reduction reaction. Oxidation must always be "coupled" with reduction, and the electrons that are "lost" by one substance must always be "gained" by another as matter (such as electrons) cannot be destroyed or created. Hence, the terms "lost or gained", simply mean that the electrons are being transferred from one particle to another.
Ionic Compounds
The simplest type of oxidation-reduction (coupled) reactions is that which occurs between metals and nonmetals of the Representative Elements. The transfer of electrons between the atoms of these elements result in drastic changes to the elements involved. This is due to the formation of ionic compounds. The reaction between sodium and chlorine serves as a typical example. The element sodium is a rather "soft" metal solid, with a silver-grey color. Chlorine is greenish colored gas. When a single electron is transferred between these elements, their atoms are transformed via a violent reaction into a totally different substance called, sodium chloride, commonly called table salt -- a white, crystalline, and brittle solid.
Sodium chloride exhibits properties quite distinct and different from sodium and chlorine. The changes in physical as well as chemical properties are due to the formation of cations and anions via the oxidation-reduction process, and the resultant, powerful attractive force that develops between these oppositely charge ions. This force of attraction is called the ionic or electrostatic bond, and serves to keep the sodium and chloride ions tightly bound in a highly organized network or lattice of alternating positive and negative charges. This entire complex of ions is called an ionic compound, and is illustrated below in two dimensions. Note how the oppositely charged ions are arranged.

Ionic Formulas
As indicated in the previous section, all ionic compounds are comprised of a definite ratio of cations and anions. This ratio of ions within the ionic compound is determined by the oxidation state of the cation and anion. In every ionic compound, the total positive charge of the cations must always equal the total negative charge of the anions, so that the net charge of the complex is always zero. Every ionic compound can be described by an ionic formula unit which lists the simplest whole number ratio of the ions in the ionic crystal lattice formed. The simplest whole number ratio of the sodium and chloride ions the network of ions shown above is:

Hence, the chemical formula for the ionic compound sodium chloride, " NaCl ". See exercise 3

Ionic Compounds-Transition Metals
The behavior of the Transition metals is similar to that of the Representative metals. They are also oxidized by nonmetals, losing their electrons to the nonmetal and forming ionic compounds. However, many Transition metals exhibit multiple oxidation states, forming cations with different positive charges. This is due to the fact that many Transition Metals are characterized by a partially filled inner electron level, inside the valence shell. Electrons within this inner shell may sometimes behave as valence electrons and are lost along with the outermost electrons during oxidation. The number of electrons lost depends on the conditions under which the chemical reactions occur. Hence, many of these metals can exhibit "multiple oxidation" states, forming cations of different charges. A typical example is iron. Depending on the conditions of the reaction, iron may form a cation with a "+2" or "+3" charge, by losing two or three electrons, respectively. Manganese, another Transition metal and an extreme example, may exist in the following oxidation states: "+2, +3, +4, +6, and +7, by losing 2, 3, 4, 6, or 7 electrons, respectively. Because the number of electrons lost by the metal depends on so many variables (temperature, amount and nature of nonmetal, etc.) the exact chemical formula of ionic compounds formed by the Transition Metals must be determined experimentally. The simple whole number ratio of the atoms in the derived formula can then be used to determine the oxidation state of the Transition Metal. See exercise 4

Part II - Extended Concepts

Concept of Electro negativity

Over the years the definition of oxidation-reduction has been broadened to include processes which involve combinations of atoms in which there is no clearcut transfer of electrons between them. An understanding of this behavior is provided by the concept of electronegativity. According to this concept, each kind of atom has a certain attraction for the electrons involved in a chemical bond. This "electron-attracting" power of each atom can be listed numerically on an electronegativity scale. Fluorine, which has the greatest attraction for electrons in bond-forming situations, is assigned the highest value on this scale. All other atoms are assigned values less than that of fluorine as shown.

Note the following trends:

Metals generally have low electronegativity values, while nonmetals have relatively high electronegativity values.
Electronegativity values generally increase from left to right within the Periodic Table of the elements.
Electronegativity values generally decrease from top to bottom within each family of elements within the Periodic Table.

When atoms react with each other, they "compete" for the electrons involved in a chemical bond. The atom with the higher electronegativity value, will always "pull" the electrons away from the atom that has the lower electronegativity value. The degree of "movement or shift" of these electrons toward the more electronegative atom is dependent on the difference in electronegativities between the atoms involved.

Electronegativity & Metal/Nonmetal Compounds

As indicated by the table shown in the previous section and below, metals generally have low electronegativity values compared to nonmetals. Hence when metals react with nonmetals, the difference in their electronegativity values is sufficient to justify the generalization that metal atoms will "lose" their valence electrons and that nonmetal atoms will gain these electrons. This generalization is the basis for following guideline. Reactions between metals and nonmetals will usually result in the formation of ionic compounds.

Redox Reactions Involving Nonmetals Only

The situation is a bit more complex when nonmetals atoms are involved. As all nonmetals have similarly high electronegativity values, it is unreasonable to assume that there will be a transfer of electrons between them in an oxidation-reduction reaction. In these instances the valence electrons involved can no longer be thought of as being "lost or gained" between the atoms, but instead, are only partially transferred, moving closer to that atom which has the higher electronegativity (and away from the atom of lower electronegativity). This "shift" of electrons results in an unequal distribution of charge, as the more electronegative atom becomes more "negative" and the atom of lower electronegativity becomes more "positive".

The accurate determination of the distribution of charge resulting from these "electron shifts" is very difficult, but guidelines have been devised to simplify the process. In general, these guidelines assign the more electronegative atom a negative oxidation state, and the atom with the lower electronegativity, a positive oxidation state. One should be aware that these guidelines are at best, arbitrary approximations, and in some instances may have to be supplemented by additional methods.

Guidelines - Oxidation States of Nonmetals

When two, nonmetals react with each other, the more electronegative element is assigned the negative oxidation state.

Fluorine, the most electronegative element, is always assigned an oxidation state of "-1" when combined with any other element.
Hydrogen, whenever it is combined in a molecule, is assigned an oxidation state of "+1".
When hydrogen combines with metals in forming compounds called, metal hydrides, it is assigned an oxidation state of (-1)

Oxygen, in most compounds, is usually assigned an oxidation state of "-2".

However, when it is found in peroxides (" O - O bonds ") it is assigned a value of "-1"; or when combined with fluorine, it is assigned a value of "+1".

The sum of the oxidation states of every element in a substance or species (it may be an ion or a molecule) must always equal the electrical charge indicated for that substance or species.

any monatomic ion has an oxidation state equal to its charge
the sum of the oxidation states of all atoms in a compound must equal zero.
the sum of the oxidation states of all atoms in a polyatomic ion must equal the charge of the ion.

Types of Red ox Reactions

Combination Reactions

One of the simplest types of redox reactions is the combination reaction. In these reactions, which involve the "combining" of two elements to form a chemical compound, one element is always oxidized, while the other is always reduced as illustrated below.
Example - Formation of water from hydrogen and oxygen gas.

Note: Hydrogen is oxidized and oxygen is reduced.

Example - Formation of sulfur trioxide from oxygen and sulfur.

Note: Sulfur is oxidized; oxygen is reduced.

Decomposition Reactions

The result of a combination reaction can be reversed; in other words, a compound can be decomposed into the components from which it was formed. This type of reaction is called a decomposition reaction. Many decomposition reactions occur via oxidation-reduction as illustrated below.

Note: Chlorine is reduced, while oxygen is oxidized.

But many other decomposition reactions do not involve a corresponding oxidation and reduction of the substances as shown below.

Note, that in this example of chemical decomposition, the oxidation states of the elements involved remain constant.

Single Displacement Reactions

Another type of redox reaction is one in which an element replaces or displaces another from a compound. In these reactions, known as single replacement reactions, the element which replaces that which is in a compound is always oxidized. The element being displaced, is always reduced. This is illustrated by the displacement of hydrogen gas by metallic iron in the example below:

The oxidation of iron is represented by:

Note that the net charge on both sides of the arrow must always be equal to each other.

The reduction of hydrogen is represented by:

Note: In both oxidation and reduction, the net charge of both sides of the arrow must always be equal.

Another example is the replacement of silver by copper.

Note: Copper is oxidized; silver is reduced. Balancing Redox Reactions Using the Half Reaction Method

Many redox reactions occur in aqueous solutions or suspensions. In this medium most of the reactants and products exist as charged species (ions) and their reaction is often affected by the pH of the medium. The following provides examples of how these equations may be balanced systematically. The method that is used is called the ion-electron or "half-reaction" method.
Example 1 -- Balancing Redox Reactions Which Occur in Acidic Solution
Organic compounds, called alcohols, are readily oxidized by acidic solutions of dichromate ions. The following reaction, written in net ionic form, records this change. The oxidation states of each atom in each compound is listed in order to identify the species that are oxidized and reduced, respectively.

An examination of the oxidation states, indicates that carbon is being oxidized, and chromium, is being reduced. To balance the equation, use the following steps:

First, divide the equation into two halves; an oxidation half-reaction and reduction half- reaction by grouping appropriate species.

2.           (red.)   (Cr₂O₇)^-2   ---->   Cr⁺³

3.

(ox.)    C₂H₆O     ---->   C₂H₄O

Second, if necessary, balance both equations by inspection. In doing this ignore any oxygen and hydrogen atoms in the formula units. In other words, balance the non-hydrogen and non-oxygen atoms only. By following this guideline in the example above, only the reduction half-reaction needs to be balanced by placing the coefficient, 2 , in front of Cr⁺³ as shown below.

5.           (red.)    (Cr₂O₇)^-2  ---->  2 Cr⁺³

6.

(ox.)    C₂H₆O     ---->  C₂H₄O

(as there are equal numbers of carbon atoms on both sides of this equation, skip this step for this half-reaction. Remember, in this step, one concentrates on balancing only non-hydrogen and non-oxygen atoms)

The third step involves balancing oxygen atoms. To do this, one must use water (H₂O) molecules. Use 1 molecule of water for each oxygen atom that needs to be balanced. Add the appropriate number of water molecules to that side of the equation required to balance the oxygen atoms as shown below.

8.           (red.)   (Cr₂O₇)^-2   ---->  2 Cr⁺³  +  7 H₂O

9.

(ox.)    C₂H₆O      ---->  C₂H₄O

(as there are equal numbers of oxygen atoms, skip this step for this half-reaction)

The fourth step involves balancing the hydrogen atoms. To do this one must use hydrogen ions (H⁺). Use one (1) H⁺ ion for every hydrogen atom that needs to balanced. Add the appropriate number of hydrogen ions to that side of the equation required to balance the hydrogen atoms as shown below

(red.)   14 H⁺ + (Cr₂O₇)^-2   --->   2 Cr⁺³ + 7 H₂O

(as there are 14 hydrogen atoms in 7 water molecules, 14 H⁺ ions must be added to the opposite side of the equation)

(ox.)    C₂H₆O ---> C₂H₄O + 2 H⁺

(2 hydrogen ions must be added to the "product" side ofthe equation to obtain a balance)

The fifth step involves the balancing of positive and negative charges. This is done by adding electrons (e-). Each electron has a charge equal to (-1). To determine the number of electrons required, find the net charge of each side the equation.

The electrons must always be added to that side which has the greater positive charge as shown below.

note: the net charge on each side of the equation does not have to equal zero.

The same step is repeated for the oxidation half-reaction.

As there is a net chargae of +2 on the product side, two electrons must be added to that side of the equation as shown below.

At this point the two half-reactions appear as:

(red)   6e^-  +  14 H⁺  +  (Cr₂O₇)^-2  ------->  2 Cr⁺³ + 7 H₂O

(ox)    C₂H₆O  ------>  C₂H₄O  +  2 H⁺  +  2e^-

The reduction half-reaction requires 6 e-, while the oxidation half-reaction produces 2 e-.

The sixth step involves multiplying each half-reaction by the smallest whole number that is required to equalize the number of electrons gained by reduction with the number of electrons produced by oxidation. Using this guideline, the oxidation half reaction must be multiplied by "3" to give the 6 electrons required by the reduction half-reaction.

(ox.)   3 C₂H₆O ---> 3 C₂H₄O + 6 H⁺ + 6e^-

The seventh and last step involves adding the two half reactions and reducing to the smallest whole number by cancelling species which on both sides of the arrow.

14.       6e^-  +  14 H⁺  +   (Cr₂O₇)^-2  ----->   2 Cr⁺³ + 7 H₂O

                    3 C₂H₆O   ----->   3 C₂H₄O + 6 H⁺ + 6e^-

adding the two half-reactions above gives the following:

6e^-+ 14H⁺ + (Cr₂O₇)^-2 + 3C₂H₆O ---> 2Cr⁺³ + 7H₂O + 3C₂H₄O + 6H⁺ + 6e^-

Note that the above equation can be further simplified by subtracting out 6 e- and 6 H⁺ ions from both sides of the equation to give the final equation.

Note: the equation above is completely balanced in terms of having an equal number of atoms as well as charges.

Example 2 - Balancing Redox Reactions in Basic Solutions
The active ingredient in bleach is the hypochlorite (OCl^-) ion. This ion is a powerful oxidizing agent which oxidizes many substances under basic conditions. A typical reaction is its behavior with iodide (I^-) ions as shown below in net ionic form.
I^- (aq) + OCl^-(aq) ------> I₂ + Cl^- + H₂O
Balancing redox equations in basic solutions is identical to that of acidic solutions except for the last few steps as shown below.

First, divide the equation into two halves; an oxidation half-reaction and reduction-reaction by grouping appropriate species.

2.           (ox)    I^-   ----> I₂

3.

(red)   OCl^- ----> Cl^- + H₂O

Second, if needed, balance both equations, by inspection ignoring any oxygen and hydrogen atoms. (The non-hydrogen and non-oxygen atoms are already balanced,hence skip this step)
Third, balance the oxygen atoms using water molecules . (The hydrogen and oxygen atoms are already balanced; hence, skip this step also.
Fourth, balance any hydrogen atoms by using an (H+) for each hydrogen atom

(ox)    2 I^- ----> I₂

(as no hydrogens are present, skip this step for this half-reaction)

(red)   2 H⁺ + OCl^- -----> Cl^- + H₂O

(two hydrogen ions must be added to balance the hydrogens in the water molecule).

Fifth, use electrons (e-) to equalize the net charge on both sides of the equation. Note; each electron (e-) represents a charge of (-1).

Sixth, equalize the number of electrons lost with the number of electrons gained by multiplying by an appropriate small whole number.

9.           (ox)   2 I^-  ---->  I₂  +  2e^-

(red)  2e^-  +  2 H⁺  +  OCl^-  ---->  Cl^-  +  H₂O

(as the number of electrons lost equals the number of electrons gained, skip this step)

Add the two equations, as shown below.

2 e^-  +  2 I^-  +  2 H⁺  +  OCl^-  ----> I₂  +  Cl^-  +  H₂O  +  2e^-

and subtract "like" terms from both sides of the equation. Subtracting "2e-" from both sides of the equation gives the net equation:

To indicate the fact that the reaction takes place in a basic solution, one must now add one (OH-) unit for every (H+) present in the equation. The OH- ions must be added to both sides of the equation as shown below.

2 OH^- + 2 I^- + 2 H⁺ + OCl^-  ----->  I₂ + Cl^- + H₂O + 2 OH^-

Then, on that side of the equation which contains both (OH^-) and (H⁺) ions, combine them to form H₂O. Note, combining the 2 OH^- with the 2 H⁺ ions above gives 2 HOH or 2 H₂O molecules as written below.

2 H₂O + 2 I^- + OCl^-  ---->  I₂ + Cl^- + H₂O + 2 OH^-

Simplify the equation by subtracting out water molecules, to obtain the final, balanced equation.

Note that both the atoms and charges are equal on both sides of the equation, and the presence of hydroxide ions (OH^-) indicates that the reaction occurs in basic solution.

Bleaching Agents

Bleaching agents are compounds which are used to remove color from substances such as textiles. In earlier times textiles were bleached by exposure to the sun and air. Today most commercial bleaches are oxidizing agents, such as sodium hypochlorite (NaOCl) or hydrogen peroxide (H₂O₂) which are quite effective in "decolorizing" substances via oxidation. The action of these bleaches can be illustrated in the following simplified way:

Recall that an oxidizing agent is any substance which causes another substance to lose one or more electrons. The decolorizing action of bleaches is due in part to their ability to remove these electrons which are activated by visible light to produce the various colors. The hypochlorite ion (OCl^-), found in many commercial preparations, is reduced to chloride ions and hydroxide ions forming a basic solution as it accepts electrons from the colored material as shown below.

   OCl^- + 2e^- + HOH --------> Cl^- + 2 OH^-

Bleaches are often combined with "optical brighteners". These compounds are quite different from bleaches. They are capable of absorbing wavelengths of ultraviolet light invisible to the human eye, and converting these wavelengths to blue or blue-green light. The blue or blue-green light is then reflected by the substance making the fabric appear much "whiter and brighter" as more visible light is seen by the eye.

Photosynthesis

An example of naturally-occuring biological oxidation-reduction reactions is the process of photosynthesis. It is a very complex process carried out by green plants, blue-green algae, and certain bacteria. These organisms are able to harness the energy contained in sunlight, and via a series of oxidation-reduction reactions, produce oxygen and sugar, as well as other compounds which may be utilized for energy as well as the synthesis of other compounds. The overall equation for the photosynthetic process may be expressed as:

     6 CO₂   +   6H₂O   -------->   C₆H₁₂O₆   +   6 O₂

                                   (glucose)

The equation is the net result of two processes. One process involves the splitting of water. This process is really an oxidative process that requires light, and is often referred to as the "light reaction". This reaction may be written as:

     12 H₂O  ----------------------->  6 O₂  +  24 H⁺  +  24e^-

             light or radiant energy

The oxidation of water is accompanied by a reduction reaction resulting in the formation of a compound, called nicotinamide adenine dinucleotide phosphate (NADPH). This reaction is illustrated below:

     NADP⁺         +  H₂0  ------->     NADPH      +  H⁺     +  O

  (oxidized form)                  (reduced form)           (oxygen)

This reaction is linked or coupled to yet another reaction resulting in the formation of a highly energetic compound, called adenosine triphosphate, (ATP). As this reaction involves the addition of a phosphate group (labeled, as P_i) to a compound called, adenosine diphosphate (ADP) during the light reaction, it is called photophosphorylation.

     ADP + P_i ------------> ATP

Think of the light reaction, as a process by which organisms "capture and store" radiant energy as they produce oxygen gas. This energy is stored in the form of chemical bonds of compounds such as NADPH and ATP.
The energy contained in both NADPH and ATP is then used to reduce carbon dioxide to glucose, a type of sugar (C₆H₁₂O₆). This reaction, shown below, does not require light, and it is often referred to as the "dark reaction".

     6 CO₂  +  24 H⁺  +  24 e^-  ------>  C₆H₁₂O₆  +  6 H₂O

The chemical bonds present in glucose also contain a considerable amount of potential energy. This stored energy is released whenever glucose is catabolized (broken down) to drive cellular processes. The carbon skeleton in glucose also serves as a source of carbon for the synthesis of other important biochemical compounds such as, lipids, amino acids, and nucleic acids.
In simplest terms, the process of photosynthessis can be viewed as one-half of the carbon cycle. In this half, energy from the sun is captured and transformed into nutrients which can be utilized by higher organisms in the food chain. The release of this energy during the metabolic re-conversion of glucose to water and carbon dioxide represents the second half of the carbon cycle and it may be referred to as catabolism or "oxidative processes".

Metabolism

Metabolism is a general term used to refer to all of the chemical reactions which occur in a living system. Metabolism can be divided into two parts; anabolism, or reactions involving the synthesis of compounds; and catabolism, or reactions involving the breakdown of compounds. In terms of oxidation-reduction principles, anabolic reactions are primarily characterized by reduction reactions, such as the dark reaction in photosynthesis where carbon dioxide is reduced to form glucose. Catabolic reactions are primarily oxidation reactions. Although catabolism involves many separate reactions, an example of such as process can be described by the oxidation of glucose as shown below. Note that this equation is the reverse of the photosynthetic equation.

  C₆H₁₂O₆  +  6 O₂  ---------->  6 CO₂  +  6 H₂O  + Energy

Note also, that in this reaction, the carbon atoms in glucose are oxidized, undergoing an increase in oxidation state (each carbon loses 2 electrons) as they are converted to carbon dioxide. At the same time, each oxygen atom is reduced by gaining 2 electrons when it is converted to water. Part of the energy is released as heat and the remainder is stored in the chemical bonds of "energetic" compounds such as adenosine triphosphate (ATP) and nicotinamide adenine dinucleotide (NADH).
Catabolic reactions can be divided into many different groups of reactions called, catabolic pathways. In these pathways (referred to as Glycolysis, the Citric Acid Cycle, and Electron Transport) the carbon atoms are slowly oxidized by a series of reactions which gradually modify the carbon skeleton of the compound as well as the oxidation state of carbon. Coupled to these reactions are other reversible oxidation-reduction reactions designed to capture the energy released and temporarily store it within the chemical bonds of compounds called adenosine triphosphate (ATP) and nicotamide dinucleotide (NADH) . These compounds are then utilized to provide energy for driving the cellular machinery.

Nitrogen Fixation

Nitrogen is the most abundant element in our atmosphere. It is a vital element as many classes of compounds essential to living systems are nitrogen-containing compounds. Nitrogen is a primary nutrient for all green plants, but it must be modified before it can be readily utilized by most living systems. Nitrogen fixation is one process by which molecular nitrogen is reduced to form ammonia. This complex process is carried out by nitrogen-fixing bacteria present in the soil. Although nitrogen-fixation involves a number of oxidation-reduction reactions that occur sequentially, that reaction which describes its reduction can be written in a simplified way as:

   N₂  +  6 e^-  +  8H⁺   --->  2 NH₄⁺  (ammonium ion)

The ammonium ion (the conjugate acid of ammonia, NH₃) that is produced by this reaction is the form of nitrogen that is used by living systems in the synthesis of many bio-organic compounds.
Another way by which ammonia may be formed is by the process called nitrification. In this process compounds called nitrates and nitrites, released by decaying organic matter are converted to ammonium ions by nitrifying bacteria present in the soil. The process carried out by these bacteria is also a complex series of oxidation-reduction reactions. The reduction reactions involving nitrate and nitrite ions can be simplified as:

      NO₃^-     +  2e^-  +  2H⁺  ----------->   NO₂^-        +  H₂O

 (nitrate ion)                          (nitrite ion)

      NO₂^-     +  6e^-  +  2H⁺  ---------->    NH₄⁺        +  2 H₂O

Another way in which molecular nitrogen is modified is via the discharge of lightning. The tremendous energy released by the electrical discharges in our atmosphere breaks the rather strong bonds between nitrogen atoms, causing them to react with oxygen. Note in this process, nitrogen is oxidized and oxygen is reduced.

                     lightning

      N₂  +  O₂   -------------->   2 NO (nitric oxide)

The nitrous oxide formed combines with oxygen to form nitrogen dioxide.

    2 NO  +  O₂   --------------->  2NO₂

Nitrogen dioxide readily dissolves in water to product nitric and nitrous acids;

    2 NO₂  +  H₂O  ------->  HNO₃  +  HNO₂

These acids readily release the hydrogen forming nitrate and nitrite ions which can be readily utilized by plants and micro-organisms.

    HNO₃   -------->   H⁺  +  NO₃^-  (nitrate ions)

    HNO₂   -------->   H⁺  +  NO₂^-  (nitrite ions)

Denitrifying bacteria, act on ammonia as well as nitrates produced by death and decay, recycling these compounds as free nitrogen (N₂). The nitrogen that is fixed by the processes described above is eventually returned to the atmosphere by this denitrification process, to complete what is commonly referred to as the "nitrogen cycle".

Combustion of Fuels

The combustion or the burning of fuels, is perhaps the most common and obvious example of oxidation and reduction. Combustion is also that process which converts the potential energy of fuels into kinetic energy (heat and light). Most fuels (gasoline, diesel oil, propane, etc.) are compounds comprised primarily of carbon and hydrogen. These hydrocarbons represent an excellent source of potential energy which is released as heat during the combustion process. A common example is the oxidation of propane, the fuel used for gas ranges:

     C₃H₈  +  5 O₂  ----->  3 H₂O  +  CO₂  +  Heat

As propane burns in air, its carbon atoms are oxidized when they combine with oxygen to form carbon dioxide. In turn, molecular oxygen is reduced by the hydrogen atoms, forming water. The heat produced can be used directly such as in the cooking of foods or to cause the expansion of the gaseous products produced to perform mechanical work such as in an internal combustion or steam engine.
Many other substances besides hydrocarbons can be used as fuels. For example, the alcohols, such as methanol (CH₃OH) and ethanol (CH₃CH₂OH) are often used in racing cars. Ethanol mixed with gasoline, called gasohol , is currently being explored as a substitute for gasoline. Among the simplest fuels is molecular hydrogen (H₂) which readily reacts with oxygen forming water as shown:

     2 H₂  +  O₂  ------>  2 H₂O  +  Energy

The simplicity and "nonpolluting" aspect of this oxidation-reduction reaction, the amount of energy produced, and the relative abundance of both hydrogen and oxygen in our environment, makes hydrogen a very attractive alternative fuel source. Research efforts are currently focused on further developing the technology to broaden its use as a source of energy.

The "Dry-Cell" Battery

The most common type of battery used today is the "dry cell" battery. There are many different types of batteries ranging from the relatively large "flashlight" batteries to the minaturized versions used for wristwatches or calculators. Although they vary widely in composition and form, they all work on the sample principle. A "dry-cell" battery is essentially comprised of a metal electrode or graphite rod (elemental carbon) surrounded by a moist electrolyte paste enclosed in a metal cylinder as shown below.

In the most common type of dry cell battery, the cathode is composed of a form of elemental carbon called graphite, which serves as a solid support for the reduction half-reaction. In an acidic dry cell, the reduction reaction occurs within the moist paste comprised of ammonium chloride (NH₄Cl) and manganese dioxide (MnO₂):

   2 NH₄⁺  +  2 MnO₂  +  2e^-  ------>  Mn₂O₃  +  2 NH₃  +  H₂O

A thin zinc cylinder serves as the anode and it undergoes oxidation:

   Zn (s)  --------------->  Zn⁺²  +  2e^-

This dry cell "couple" produces about 1.5 volts. ( These "dry cells" can also be linked in series to boost the voltage produced). In the alkaline version or "alkaline battery", the ammonium chloride is replaced by KOH or NaOH and the half-cell reactions are:

       Zn  +     2 OH^-     ------->  ZnO    +  H₂O    +  2e^-

   2 MnO₂  +  2e^-  +  H₂O  ------->  Mn₂O₃  +  2 OH^-

The alkaline dry cell lasts much longer as the zinc anode corrodes less rapidly under basic conditons than under acidic conditions.
Other types of dry cell batteries are the silver battery in which silver metal serves as an inert cathode to support the reduction of silver oxide (Ag₂O) and the oxidation of zinc (anode) in a basic medium. The type of battery commonly used for calculators is the mercury cell. In this type of battery, HgO serves as the oxidizing agent (cathode) in a basic medium, while zinc metal serves as the anode. Another type of battery is the nickel/cadmium battery, in which cadmium metal serves as the anode and nickel oxide serves as the cathode in an alkaline medium. Unlike the other types of dry cells described above, the nickel/cadmium cell can be recharged like the lead-acid battery.

Electrochemical Cells

Many oxidation-reduction reactions occur spontaneously, giving off energy. An example involves the spontaneous reaction that occurs when zinc metal is placed in a solution of copper ions as described by the net ionic equation shown below.

     Cu⁺² (aq)  +  Zn (s)  ------->  Cu(s)  +  Zn⁺² (aq)

The zinc metal slowly "dissolves" as its oxidation produces zinc ions which enter into solution. At the same time, the copper ions gain electrons and are converted into copper atoms which coats the zinc metal or sediments to the bottom of the container. The energy produced in this reaction is quickly dissipated as heat, but it can be made to do useful work by a device called, an electrochemical cell. This is done in the following way.
An electrochemical cell is composed to two compartments or half-cells, each composed of an electrode dipped in a solution of electrolyte. These half-cells are designed to contain the oxidation half-reaction and reduction half-reaction separately as shown below.

The half-cell, called the anode, is the site at which the oxidation of zinc occurs as shown below.

     Zn (s)  ---------->  Zn⁺² (aq)  +  2e^-

During the oxidation of zinc, the zinc electrode will slowly dissolve to produce zinc ions (Zn⁺²), which enter into the solution containing Zn⁺² (aq) and SO₄^-2 (aq) ions.
The half-cell, called the cathode, is the site at which reduction of copper occurs as shown below.

     Cu⁺² (aq)  +  2e^-  ------->  Cu (s)

When the reduction of copper ions (Cu⁺²) occurs, copper atoms accumulate on the surface of the solid copper electrode.
The reaction in each half-cell does not occur unless the two half cells are connected to each other.
Recall that in order for oxidation to occur, there must be a corresponding reduction reaction that is linked or "coupled" with it. Moreover, in an isolated oxidation or reduction half-cell, an imbalance of electrical charge would occur, the anode would become more positive as zinc cations are produced, and the cathode would become more negative as copper cations are removed from solution. This problem can be solved by using a "salt bridge" connecting the two cells as shown in the diagram below. A "salt bridge" is a porous barrier which prevents the spontaneous mixing of the aqueous solutions in each compartment, but allows the migration of ions in both directions to maintain electrical neutrality. As the oxidation-reduction reaction occurs, cations ( Zn⁺²) from the anode migrate via the salt bridge to the cathode, while the anion, (SO₄)^-2, migrates in the opposite direction to maintain electrical neutrality.
The two half-cells are also connected externally. In this arrangement, electrons provided by the oxidation reaction are forced to travel via an external circuit to the site of the reduction reaction. The fact that the reaction occurs spontaneously once these half cells are connected indicates that there is a difference in potential energy. This difference in potential energy is called an electomotive force (emf) and is measured in terms of volts. The zinc/copper cell has an emf of about 1.1 volts under standard conditions.

Any electrical device can be "spliced" into the external circuit to utilize this potential energy produced by the cell for useful work. Although the energy available from a single cell is relatively small, electrochemical cells can be linked in series to boost their energy output. A common and useful application of this characteristic is the "battery". An example is the lead-acid battery used in automobiles. In the lead-acid battery, each cell has a lead metal anode and lead (IV) oxide (lead dioxide) cathode both of which are immersed in a solution of sulfuric acid. This single electrochemical cell produces about 2 volts. Linking 6 of these cells in series produces the 12-volt battery found in most cars today. One disadvantage of these "wet cells" such as the lead-acid battery is that it is very heavy and bulky. However, like many other "wet cells", the oxidation-reduction reaction which occurs can be readily reversed via an external current such as that provided by an automobile's alternator. This prolongs the lifetime and usefulness of such devices as an energy source.

Photo-oxidation - Photochromic Glass

Many people who wear eye glasses prefer those made with photochromic lenses or glass lenses which darken when exposed to bright light. These eyeglasses eliminate the need for sunglasses as they can reduce up to 80% of transmitted light. The basis of this change in color in response to light can be explained in terms of oxidation-reduction reactions. Glass consists of a complex matrix of silicates which is ordinarily transparent to visible light. In photochromic lenses, silver chloride (AgCl) and copper (I) chloride (CuCl) crystals are added during the manufacturing of the glass while it is in the molten state and these crystals become uniformly embedded in the glass as it solidifies. One characteristic of silver chloride is its suscepitibility to oxidation and reduction by light as described below.

              Cl^-   ----------->   Cl   +   e^-

                     oxidation

     Ag⁺   +   e^-   ----------->   Ag

                     reduction

The chloride ions are oxidized to produce chlorine atoms and an electron. The electron is then transferred to silver ions to produce silver atoms. These atoms cluster together and block the transmittance of light, causing the lenses to darken. This process occurs almost instantaneously. As the degree of "darkening" is dependent on the intensity of the light, these photochromic lenses are quite convenient and all but eliminate the need for an extra pair of sunglasses.
The photochromic process would not be useful unless it were reversible. The presence of copper (I) chloride reverses the darkening process in the following way. When the lenses are removed from light, the following reactions occur:

       Cl        +       Cu⁺      ------>   Cu⁺²       +       Cl^-

oxidizing agent reducing agent    oxidized species   reduced species

The chlorine atoms formed by the exposure to light are reduced by the copper ions, preventing their escape as gaseous atoms from the matrix. The copper (+1) ion is oxidized to produce copper (+2) ions, which then reacts with the silver atoms as shown.

       Cu⁺²      +        Ag      ------>   Cu⁺¹       +       Ag⁺

oxidizing agent reducing agent    reduced species  oxidized species

The net effect of these reactions is that the lenses become transparent again as the silver and chloride atoms are converted to their original oxidized and reduced states.

Corrosion

Millions of dollars are lost each year because of corrosion. Much of this loss is due to the corrosion of iron and steel, although many other metals may corrode as well. The problem with iron as well as many other metals is that the oxide formed by oxidation does not firmly adhere to the surface of the metal and flakes off easily causing "pitting". Extensive pitting eventually causes structural weakness and disintegration of the metal. (It should be noted, however, that certain metals such as aluminum, form a very tough oxide coating which strongly bonds to the surface of the metal preventing the surface from further exposure to oxygen and corrosion).
Corrosion occurs in the presence of moisture. For example when iron is exposed to moist air, it reacts with oxygen to form rust,

The amount of water complexed with the iron (III) oxide (ferric oxide) varies as indicated by the letter "X". The amount of water present also determines the color of rust, which may vary from black to yellow to orange brown. The formation of rust is a very complex process which is thought to begin with the oxidation of iron to ferrous (iron "+2") ions.

     Fe   ------->   Fe⁺²   +   2 e^-

Both water and oxygen are required for the next sequence of reactions. The iron (+2) ions are further oxidized to form ferric ions (iron "+3") ions.

     Fe⁺²   ------------>   Fe⁺³   +   1 e^-

Tthe electrons provided from both oxidation steps are used to reduce oxygen as shown.

     O₂ (g)   +   2 H₂O   +   4e^-   ------>   4 OH^-

The ferric ions then combine with oxygen to form ferric oxide [iron (III) oxide] which is then hydrated with varying amounts of water. The overall equation for the rust formation may be written as :

The formation of rust can occur at some distance away from the actual pitting or erosion of iron as illustrated below. This is possible because the electrons produced via the initial oxidation of iron can be conducted through the metal and the iron ions can diffuse through the water layer to another point on the metal surface where oxygen is available. This process results in an electrochemical cell in which iron serves as the anode, oxygen gas as the cathode, and the aqueous solution of ions serving as a "salt bridge" as shown below.

The involvement of water accounts for the fact that rusting occurs much more rapidly in moist conditions as compared to a dry environment such as a desert. Many other factors affect the rate of corrosion. For example the presence of salt greatly enhances the rusting of metals. This is due to the fact that the dissolved salt increases the conductivity of the aqueous solution formed at the surface of the metal and enhances the rate of electrochemical corrosion. This is one reason why iron or steel tend to corrode much more quickly when exposed to salt (such as that used to melt snow or ice on roads) or moist salty air near the ocean.

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Thursday 9 June 2011

Basic Inorganic Chemistry