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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-REDUCTION REACTIONS: A BASIC MODEL

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.

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.

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:

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.

EVERYDAY EXAMPLES

Oxidation-reduction reactions have many far-reaching applications in our lives. Some of these applications are so common, that we take them for granted; others are not so obvious. The following are just a few examples of oxidation-reduction reactions.

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