General General Chemistry:Chemistry:Oxidation-Oxidation-Reduction Reduction ReactionsReactions

CE 541CE 541

Oxidation - ReductionOxidation - Reduction

IntroductionIntroduction The rusting of metals, the process involved in The rusting of metals, the process involved in

photography, the way living systems produce and utilize photography, the way living systems produce and utilize energy, and the operation of a car battery, are few energy, and the operation of a car battery, are few examples of a very common and important type of examples of a very common and important type of chemical reaction. These chemical changes are all chemical reaction. These chemical changes are all classified as "electron-transfer" or oxidation-reduction classified as "electron-transfer" or oxidation-reduction reactions. reactions.

The term, oxidation , was derived from the observation The term, oxidation , was derived from the observation that almost all elements reacted with oxygen to form that almost all elements reacted with oxygen to form compounds called, oxides. A typical example is the compounds called, oxides. A typical example is the corrosion or rusting of iron as described by the chemical corrosion or rusting of iron as described by the chemical equation: equation:

4 Fe + 3 O4 Fe + 3 O22 -----> 2 Fe -----> 2 Fe22OO33 Reduction, was the term originally used to describe the Reduction, was the term originally used to describe the

removal of oxygen from metal ores, which "reduced" the removal of oxygen from metal ores, which "reduced" the metal ore to pure metal as shown below: metal ore to pure metal as shown below:

2 Fe2 Fe22OO33 + 3 C -----> 3 CO + 3 C -----> 3 CO22 + 4 Fe + 4 Fe Based on the two examples above, oxidation can be Based on the two examples above, oxidation can be

defined very simply as, the "addition" of oxygen; and defined very simply as, the "addition" of oxygen; and reduction, as the "removal" of oxygen. reduction, as the "removal" of oxygen.

Oxidation States/Oxidation Oxidation States/Oxidation NumbersNumbers

All atoms are electrically neutral even though All atoms are electrically neutral even though they are comprised of charged, subatomic they are comprised of charged, subatomic particles. The terms, oxidation state or oxidation particles. The terms, oxidation state or oxidation number, have been developed to describe this number, have been developed to describe this "electrical state" of the atom. "electrical state" of the atom. The oxidation The oxidation state or oxidation number of an atom is state or oxidation number of an atom is simply defined as the sum of the negative and simply defined as the sum of the negative and positive charges in an atom.positive charges in an atom. Since every atom Since every atom contains equal numbers of positive and negative contains equal numbers of positive and negative charges, the oxidation state or oxidation number charges, the oxidation state or oxidation number of any atom is always zero. This is illustrated by of any atom is always zero. This is illustrated by simply totaling the opposite charges of the atoms simply totaling the opposite charges of the atoms

as shown by the following examples.as shown by the following examples.

Oxidation-Reduction Oxidation-Reduction ReactionsReactions Assigning all atoms an oxidation state of zero serves as an important Assigning all atoms an oxidation state of zero serves as an important

reference point, as oxidation-reduction reactions always involve a reference point, as oxidation-reduction reactions always involve a change in the oxidation state of the atoms or ions involved. This 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. change in oxidation state is due to the "loss" or "gain" of electrons. The loss of electrons from an atom produces a positive oxidation The loss of electrons from an atom produces a positive oxidation state, while the gain of electrons results in negative oxidation states. state, while the gain of electrons results in negative oxidation states.

The changes that occur in the oxidation state of certain elements can The changes that occur in the oxidation state of certain elements can be predicted quickly and accurately by the use of simple guidelines. be predicted quickly and accurately by the use of simple guidelines. These guidelines can be divided into two classes; the metals and These guidelines can be divided into two classes; the metals and nonmetals. nonmetals.

All metal atoms are characterized by their tendency to be oxidized, All metal atoms are characterized by their tendency to be oxidized, losing one or more electrons, forming a positively charged ion, losing one or more electrons, forming a positively charged ion, called a cation. During this oxidation reaction , the oxidation state of called a cation. During this oxidation reaction , the oxidation state of the metal always increases from zero to a positive number, such as the metal always increases from zero to a positive number, such as "+1, +2, +3...." , depending on the number of electrons lost. The "+1, +2, +3...." , depending on the number of electrons lost. The number of electrons lost by these metals and the charge of the cation number of electrons lost by these metals and the charge of the cation formed are always equal to the Group number of the metal as formed are always equal to the Group number of the metal as summarized belowsummarized below. .

The group The group numbers also numbers also correspond to correspond to the electrons the electrons that are found that are found in the in the outermost outermost energy levels of energy levels of these atoms. these atoms. These electrons These electrons are often called are often called valence electrovalence electronsns. .

Group Group NumberNumber

No of No of ElectroElectrons Lostns Lost

Charge Charge of of

Cation Cation FormedFormed

II 11 +1+1

IIII 22 +2+2

IIIIII 33 +3+3

IVIV 44 +4+4

By convention oxidation reactions are By convention oxidation reactions are written in the following form using written in the following form using the element, Calcium, as an example the element, Calcium, as an example

Note that the oxidation state Note that the oxidation state increases from zero to a positive increases from zero to a positive number (from "0" to "+2" in the above number (from "0" to "+2" in the above example) and is always numerically example) and is always numerically equal to the number of electrons lost. equal to the number of electrons lost.

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

Group Group NumbeNumbe

rr

No of No of ElectroElectro

ns ns GainedGained

Charge Charge of of

Anion Anion FormeForme

dd

IVIV 44 -4-4

VV 33 -3-3

VIVI 22 -2-2

VIIVII 11 -1-1

VIIIVIII 00 No No tendency tendency to form to form anionanion

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

By convention reduction reactions are written in the By convention reduction reactions are written in the following way: following way:

Note that the charge of anion formed is always Note that the charge of anion formed is always numerically equal to the number of electrons gained. numerically equal to the number of electrons gained.

One important fact to remember in studying oxidation-One important fact to remember in studying oxidation-reduction reactions is that the process of oxidation reduction reactions is that the process of oxidation cannot occur without a corresponding reduction cannot occur without a corresponding reduction reaction. Oxidation must always be "coupled" with reaction. Oxidation must always be "coupled" with reduction, and the electrons that are "lost" by one reduction, and the electrons that are "lost" by one substance must always be "gained" by another, as substance must always be "gained" by another, as matter (such as electrons) cannot be destroyed or matter (such as electrons) cannot be destroyed or created. Hence, the terms "lost or gained", simply mean created. Hence, the terms "lost or gained", simply mean that the electrons are being transferred from one that the electrons are being transferred from one particle to another. particle to another.

Ionic CompoundsIonic Compounds

The simplest type of oxidation-reduction The simplest type of oxidation-reduction (coupled) reactions is that which occurs (coupled) reactions is that which occurs between metals and nonmetals. The transfer of between metals and nonmetals. The transfer of electrons between the atoms of these elements electrons between the atoms of these elements result in drastic changes to the elements result in drastic changes to the elements involved. This is due to the formation of ionic involved. This is due to the formation of ionic compounds. The reaction between sodium and compounds. The reaction between sodium and chlorine serves as a typical example. The chlorine serves as a typical example. The element sodium is a rather "soft" metal solid, element sodium is a rather "soft" metal solid, with a silver-grey color. Chlorine is greenish with a silver-grey color. Chlorine is greenish colored gas. When a single electron is colored gas. When a single electron is transferred between these elements, their transferred between these elements, their atoms are transformed via a violent reaction atoms are transformed via a violent reaction into a totally different substance called, sodium into a totally different substance called, sodium chloride, commonly called table salt -- a white, chloride, commonly called table salt -- a white, crystalline, and brittle solid. crystalline, and brittle solid.

Sodium chloride exhibits properties quite distinct Sodium chloride exhibits properties quite distinct and different from sodium and chlorine. The and different from sodium and chlorine. The changes in physical as well as chemical changes in physical as well as chemical properties are due to the formation of cations properties are due to the formation of cations and anions via the oxidation-reduction process, and anions via the oxidation-reduction process, and the resultant, powerful attractive force that and the resultant, powerful attractive force that develops between these oppositely charge ions. develops between these oppositely charge ions. This force of attraction is called the ionic or This force of attraction is called the ionic or electrostatic bond, and serves to keep the sodium electrostatic bond, and serves to keep the sodium and chloride ions tightly bound in a highly and chloride ions tightly bound in a highly organized network or lattice of alternating organized network or lattice of alternating positive and negative charges. This entire positive and negative charges. This entire complex of ions is called an ionic compound, and complex of ions is called an ionic compound, and is illustrated below in two dimensions. Note how is illustrated below in two dimensions. Note how the oppositely charged ions are arranged. the oppositely charged ions are arranged.

Ionic FormulaIonic Formula All ionic compounds are comprised of a definite All ionic compounds are comprised of a definite

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

Ionic Compounds Involving Ionic Compounds Involving Transition Metals Transition Metals

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

Concept of Concept of ElectronegativityElectronegativity

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

Note the following trends: Note the following trends: Metals generally have low electronegativity values, Metals generally have low electronegativity values,

while nonmetals have relatively high while nonmetals have relatively high electronegativity values. electronegativity values.

Electronegativity values generally increase from Electronegativity values generally increase from left to right within the Periodic Table of the left to right within the Periodic Table of the elements. elements.

Electronegativity values generally decrease from Electronegativity values generally decrease from top to bottom within each family of elements top to bottom within each family of elements within the Periodic Table. within the Periodic Table.

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

Electronegativity of Metals Electronegativity of Metals and Non-Metalsand Non-Metals

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

Redox Reactions Involving Redox Reactions Involving Nonmetals OnlyNonmetals Only

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

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

Guidelines - Oxidation Guidelines - Oxidation States of NonmetalsStates of Nonmetals

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

assigned an oxidation state of "-1" when combined with any assigned an oxidation state of "-1" when combined with any other element. other element.

Hydrogen, whenever it is combined in a molecule, is Hydrogen, whenever it is combined in a molecule, is assigned an oxidation state of "+1". assigned an oxidation state of "+1".

When hydrogen combines with metals in forming When hydrogen combines with metals in forming compounds called, metal hydrides, it is assigned an oxidation compounds called, metal hydrides, it is assigned an oxidation state of (-1) state of (-1)

2.2. Oxygen, in most compounds, is usually assigned an oxidation Oxygen, in most compounds, is usually assigned an oxidation state of "-2". state of "-2". However, when it is found in peroxides (" O - O bonds ") it However, when it is found in peroxides (" O - O bonds ") it

is assigned a value of "-1"; or when combined with fluorine, is assigned a value of "-1"; or when combined with fluorine, it is assigned a value of "+1". it is assigned a value of "+1".

Guidelines - Oxidation Guidelines - Oxidation States of NonmetalsStates of Nonmetals

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

any monatomic ion has an oxidation state equal any monatomic ion has an oxidation state equal to its charge to its charge

the sum of the oxidation states of all atoms in a the sum of the oxidation states of all atoms in a compound must equal zero. compound must equal zero.

the sum of the oxidation states of all atoms in a the sum of the oxidation states of all atoms in a polyatomic ion must equal the charge of the ion. polyatomic ion must equal the charge of the ion.

Types of Redox Reactions Types of Redox Reactions

Combination ReactionsCombination Reactions Decomposition ReactionsDecomposition Reactions Single Displacement Reactions Single Displacement Reactions

Combination Reactions Combination Reactions One of the simplest types of redox reactions One of the simplest types of redox reactions

is the is the combination reactioncombination reaction. In these . In these reactions, which involve the "combining" of reactions, which involve the "combining" of two elements to form a chemical compound, two elements to form a chemical compound, one element is always oxidized, while the one element is always oxidized, while the other is always reduced as illustrated other is always reduced as illustrated below. below.

Example - Formation of water from Example - Formation of water from hydrogen and oxygen gas.hydrogen and oxygen gas.

Note: Hydrogen is oxidized and oxygen is Note: Hydrogen is oxidized and oxygen is reduced. reduced.

Example - Formation of sulfur trioxide from Example - Formation of sulfur trioxide from oxygen and sulfur.oxygen and sulfur.

Note: Sulfur is oxidized; oxygen is reduced.Note: Sulfur is oxidized; oxygen is reduced.

Decomposition ReactionsDecomposition Reactions

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

Note: Chlorine is reduced, while oxygen is Note: Chlorine is reduced, while oxygen is oxidized.oxidized.

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

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

Single Displacement Single Displacement Reactions Reactions

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

The oxidation of iron is represented by: The oxidation of iron is represented by:

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

The reduction of hydrogen is represented by: The reduction of hydrogen is represented by:

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

Another example is the replacement of silver by Another example is the replacement of silver by copper. copper.

Note: Copper is oxidized; silver is reduced. Note: Copper is oxidized; silver is reduced.

Balancing Redox Reactions Balancing Redox Reactions Using the Half Reaction Using the Half Reaction

MethodMethodMany redox reactions occur in aqueous solutions Many redox reactions occur in aqueous solutions or suspensions. In this medium most of the or suspensions. In this medium most of the reactants and products exist as charged species reactants and products exist as charged species (ions) and their reaction is often affected by the (ions) and their reaction is often affected by the pH of the medium. The following provides pH of the medium. The following provides examples of how these equations may be examples of how these equations may be balanced systematically. The method that is used balanced systematically. The method that is used is called the ion-electron or "half-reaction" is called the ion-electron or "half-reaction" method.method.

Example 1 -- Balancing Example 1 -- Balancing Redox Reactions Which Redox Reactions Which Occur in Acidic SolutionOccur in Acidic Solution

Organic compounds, called alcohols, are Organic compounds, called alcohols, are readily oxidized by acidic solutions of readily oxidized by acidic solutions of dichromate ions. The following reaction, dichromate ions. The following reaction, written in net ionic form, records this written in net ionic form, records this change. The oxidation states of change. The oxidation states of each each atom in each compound atom in each compound is listed in is listed in order to identify the species that are order to identify the species that are oxidized and reduced, respectively.oxidized and reduced, respectively.

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

First, divide the equation into two First, divide the equation into two halves; an oxidation half-reaction halves; an oxidation half-reaction and reduction half- reaction by and reduction half- reaction by grouping appropriate species.grouping appropriate species. (red.) (red.) (Cr2O7)(Cr2O7)-2-2 Cr Cr+3+3     (ox.) (ox.) CC22HH66O O C C22HH44OO

Second, if necessary, balance both equations Second, if necessary, balance both equations by inspection. In doing this ignore any oxygen by inspection. In doing this ignore any oxygen and hydrogen atoms in the formula units. In and hydrogen atoms in the formula units. In other words, balance the non-hydrogen and other words, balance the non-hydrogen and non-oxygen atoms only. By following this non-oxygen atoms only. By following this guideline, only the reduction half-reaction guideline, only the reduction half-reaction needs to be balanced by placing the coefficient, needs to be balanced by placing the coefficient, 2 , in front of Cr2 , in front of Cr+3+3 as shown below. as shown below.

(red.) (red.) (Cr(Cr22OO77))-2-2 2 Cr 2 Cr+3+3     (ox.) (ox.) CC22HH66O O C C22HH44OO

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

The third step involves balancing oxygen The third step involves balancing oxygen atoms. To do this, one must use water (Hatoms. To do this, one must use water (H22O) O) molecules. Use 1 molecule of water for each molecules. Use 1 molecule of water for each oxygen atom that needs to be balanced. Add oxygen atom that needs to be balanced. Add the appropriate number of water molecules to the appropriate number of water molecules to that side of the equation required to balance that side of the equation required to balance the oxygen atoms as shown below. the oxygen atoms as shown below.

(red.) (red.) (Cr(Cr22OO77))-2-2 2 Cr 2 Cr+3+3 + 7 H + 7 H22O   O   (ox.) (ox.) CC22HH66O O C C22HH44O O

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

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

(red.) (red.) 14 H14 H++ + (Cr + (Cr22OO77))-2-2 2 Cr 2 Cr+3+3 + 7 H + 7 H22O O (as there are 14 hydrogen atoms in 7 water molecules, 14 (as there are 14 hydrogen atoms in 7 water molecules, 14

H+ ions must be added to the opposite side of the equation)H+ ions must be added to the opposite side of the equation)

(ox.) (ox.) CC22HH66O O C C22HH44O + 2 HO + 2 H++ (2 hydrogen ions must be added to the "product" side of the (2 hydrogen ions must be added to the "product" side of the

equation to obtain a balance)equation to obtain a balance)

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

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

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

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

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

At this point the two half-reactions appear as:At this point the two half-reactions appear as:

(red) (red) 6e6e-- + 14 H + 14 H++ + (Cr + (Cr22OO77))-2-2 2 Cr 2 Cr+3+3 + 7 + 7 HH22O   O  

(ox) (ox) CC22HH66O O C C22HH44O + 2 HO + 2 H++ + 2e + 2e--

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

The sixth step involves multiplying each The sixth step involves multiplying each half-reaction by the smallest whole half-reaction by the smallest whole number that is required to equalize the number that is required to equalize the number of electrons gained by reduction number of electrons gained by reduction with the number of electrons produced with the number of electrons produced by oxidation. Using this guideline, the by oxidation. Using this guideline, the oxidation half reaction must be oxidation half reaction must be multiplied by "3" to give the 6 electrons multiplied by "3" to give the 6 electrons required by the reduction half-reaction. required by the reduction half-reaction. (ox.) (ox.) 3 C3 C22HH66O O 3 C 3 C22HH44O + 6 HO + 6 H++ + 6e + 6e--

The seventh and last step involves adding the The seventh and last step involves adding the two half reactions and reducing to the smallest two half reactions and reducing to the smallest whole number by cancelling species which on whole number by cancelling species which on both sides of the arrow. both sides of the arrow. 6e6e-- + 14 H + 14 H++ + (Cr + (Cr22OO77))-2-2 2 Cr 2 Cr+3+3 + 7 H + 7 H22O O

3 C3 C22HH66O O 3 C 3 C22HH44O + 6 HO + 6 H++ + 6e + 6e--

adding the two half-reactions above gives adding the two half-reactions above gives the following:the following:

6e6e-- + 14H + 14H++ + (Cr + (Cr22OO77))-2-2 + 3C + 3C22HH66O O 2Cr2Cr+3+3 + 7H + 7H22O + 3CO + 3C22HH44O + 6HO + 6H++ + 6e + 6e--

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

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

Example 2 - Balancing Example 2 - Balancing Redox Reactions in Basic Redox Reactions in Basic

Solutions Solutions The active ingredient in bleach is the The active ingredient in bleach is the hypochlorite (OClhypochlorite (OCl--) ion. This ion is a powerful ) ion. This ion is a powerful oxidizing agent which oxidizes many substances oxidizing agent which oxidizes many substances under basic conditions. A typical reaction is its under basic conditions. A typical reaction is its behavior with iodide (Ibehavior with iodide (I--) ions as shown below in ) ions as shown below in net ionic form.net ionic form.

II-- (aq) + OCl (aq) + OCl--(aq) (aq) I I22 + Cl + Cl-- + H + H22OO

Balancing redox equations in basic solutions is Balancing redox equations in basic solutions is identical to that of acidic solutions except for the identical to that of acidic solutions except for the last few steps. last few steps.

First, divide the equation into two halves; First, divide the equation into two halves; an oxidation half-reaction and reduction-an oxidation half-reaction and reduction-reaction by grouping appropriate species. reaction by grouping appropriate species. (ox) I(ox) I-- I I22

(red) OCl(red) OCl-- Cl Cl-- + H + H22OO Second, if needed, balance both Second, if needed, balance both

equations, by inspection ignoring any equations, by inspection ignoring any oxygen and hydrogen atoms. (The non-oxygen and hydrogen atoms. (The non-hydrogen and non-oxygen atoms are hydrogen and non-oxygen atoms are already balanced, hence skip this step) already balanced, hence skip this step)

Third, balance the oxygen atoms using Third, balance the oxygen atoms using water molecules . (The hydrogen and water molecules . (The hydrogen and oxygen atoms are already balanced; oxygen atoms are already balanced; hence, skip this step also. hence, skip this step also.

Fourth, balance any hydrogen atoms by Fourth, balance any hydrogen atoms by using an (Husing an (H++) for each hydrogen atom ) for each hydrogen atom (ox) 2 I(ox) 2 I-- I I22 (as no hydrogen is present, skip (as no hydrogen is present, skip

this step for this half-reaction)this step for this half-reaction) (red) 2 H(red) 2 H++ + OCl + OCl-- Cl Cl-- + H + H22OO (two (two

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

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

Sixth, equalize the number of Sixth, equalize the number of electrons lost with the number of electrons lost with the number of electrons gained by multiplying by electrons gained by multiplying by an appropriate small whole number. an appropriate small whole number. (ox) 2 I(ox) 2 I-- I I22 + 2e + 2e--

(red) 2e(red) 2e-- + 2 H + 2 H++ + OCl + OCl-- Cl Cl-- + H + H22OO (as the number of electrons lost equals (as the number of electrons lost equals the number of electrons gained, skip the number of electrons gained, skip this step)this step)

Add the two equations, as shown Add the two equations, as shown below. below. 2 e2 e-- + 2 I + 2 I-- + 2 H + 2 H++ + OCl + OCl-- I I22 + Cl + Cl-- + +

HH22O + 2eO + 2e--

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

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

2 OH2 OH-- + 2 I + 2 I-- + 2 H + 2 H++ + OCl + OCl-- I I22 + Cl + Cl-- + H + H22O O + 2 OH+ 2 OH--

Then, on that side of the equation Then, on that side of the equation which contains both (OHwhich contains both (OH--) and (H) and (H++) ) ions, combine them to form Hions, combine them to form H22O. O. Note, combining the 2 OHNote, combining the 2 OH-- with the 2 with the 2 HH++ ions above gives 2 HOH or 2 H ions above gives 2 HOH or 2 H22O O molecules as written below. molecules as written below.

2 H2 H22O + 2 IO + 2 I-- + OCl + OCl-- I I22 + Cl + Cl-- + H + H22O O + 2 OH+ 2 OH--

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

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