chapter 20

Chapter 20Oxidation-Reduction Reactions

Some Common Reactions

The combustion of gasoline in an automobile engine requires oxygen

Burning of wood in a fireplace requires oxygen

The reactions that break down food in your body and release energy use oxygen

The oxide of hydrogen is water

Charcoal oxidizes when it burn forming CO2

Bleaching stains in fabric is still oxidation even though it does not burn.

Oxidation

When methane burns in air, it oxidizes and forms oxides of carbon and hydrogen.

CH4(g) + 2O2(g) CO2(g) + 2H2O(g)

When elemental iron turns to rust, it slowly oxidizes to compounds such as iron (III) oxide.

4Fe(s) + 3O2(g) 2Fe2O3(s)

Reduction

Originally reduction meant a loss of oxygen from a compound

2Fe2O3(s) + 3C(s) 4Fe(s) + 3CO2(g) iron oxide carbon iron carbon dioxide

Reduction of iron ore to metallic iron involves the removal of oxygen from iron (III) oxide.

Involves heating the ore with carbon.

Question

What happens to magnesium and oxygen when they react to form magnesium oxide?

+2 -2

2Mg + O2 2MgO magnesium oxygen magnesium oxide

Magnesium loses electrons to form Mg2+

Oxygen gains electrons to form O2-

Electron Shift in Redox Reactions

The modern concept of oxidation and reduction have been extended to include many reactions that do not even involve oxygen.

Oxygen is the most electronegative element (besides fluorine)

When oxygen bonds with an atom of a different element (except fluorine), electrons from that atom shift toward oxygen.

Redox Reactions

Redox reactions are currently understood to involve any shift of electron between reactants.

OxidationOxidation – a process that involves a complete or partial loss of electrons or a gain of oxygen.

• Results in an increase in the oxidation number of an atom

ReductionReduction – a process that involves a complete or partial gain of electrons or the loss of oxygen.

• Results in a decrease in the oxidation number of an atom

Redox Reactions

Oxidation and reduction always occur simultaneously.

The substance gaining oxygen or losing electrons is oxidized

The substance losing oxygen or gaining electrons is reduced.

Redox Reactions

During a reaction between a metal and a nonmetal, electrons are transferred from atoms of the metal to atoms of the nonmetal.

Mg + S Mg2+S2- magnesium sulfur magnesium sulfide

2 electrons are transferred from a magnesium atom to a sulfur atom.

Magnesium atoms become more stable by the loss of electrons. Sulfur atoms become more stable by the gain of electrons

Redox Reactions

Mg + S Mg2+S2- magnesium sulfur magnesium sulfide

Oxidation: Mg Mg2+ + 2e- (loss of electrons)

Reduction: S + 2e- S2- (gain of electrons)

Magnesium atom is said to be oxidized to a magnesium ion

Sulfur atom is said to be reduced to a sulfide ion.

Redox Reactions

When a metal combines with oxygen, it loses electrons

When oxygen is removed from the oxide of a metal, the metal gains electrons.

This knowledge is what led to the broader definition of oxidation and reduction as an exchange of electrons.

Redox Reactions

Reducing agentReducing agent – the substance that loses the electrons

Mg + S MgS reducing agent

oxidized

Oxidizing agentOxidizing agent – the substance that accepts electrons

Mg + S MgS oxidizing agent reduced

Sample Problem

Silver nitrate reacts with copper to form copper nitrate and silver. From the equation below, determine what is oxidized and what is reduced. Identify the oxidizing agent and the reducing agent.

2AgNO3 + Cu Cu(NO3)2 + 2Ag

Rewrite the equation in ionic form

2Ag+ + 2NO3- + Cu Cu2+ + 2NO3

- + 2Ag

Sample Problem

2Ag+ + 2NO3- + Cu Cu2+ + 2NO3

- + 2Ag

Oxidation: 2 e- are lost from copper when it becomes a Cu2+

Reduction: 2e- are gained by two silver ions which become neutral silver atoms.

2Ag+ + 2NO3- + Cu Cu2+ + 2NO3

- + 2Ag

oxidizing reducingAgent agentreduced oxidized

Sample Problem

2Na + S 2Na+ + S2- oxidized reduced reducing agent oxidizing agent

4Al + 3O2 4Al3+ + 3O2-

Oxidized reduced Reducing agent oxidizing agent

2I- I2 + 2e-

oxidation

Zn2+ + 2e- Zn

reduction

Redox with Covalent CompoundsSome reactions involve covalent compounds. In

these compounds complete electron transfer does not occur.

2H2H22 (g) + O2 (g) 2H2O (l)

In each reactant hydrogen molecule, the bonding electrons are shared equally between the hydrogen atoms.

In water, however, the bonding electrons are pulled toward oxygen because it is much more electronegative than hydrogen.

Redox with Covalent Compounds

2H2H22 (g) + O2 (g) 2H2O (l)

There is a shift of bonding electrons away from hydrogen, even though there is not a complete transfer.

Hydrogen is oxidized because it undergoes a partial loss of electrons.

Redox with Covalent Compounds2H2 (g) + OO2 2 (g) 2H2O (l)

In oxygen, the bonding electrons are share equally between oxygen atoms in the reactant oxygen molecule.

When oxygen bonds to hydrogen in the water molecule, there is a shift of electrons toward oxygen.

Oxygen is thus reduced because it undergoes a partial gain of electrons.


H H O O H O shift of bonding

e- shared e- shared e- away from H

equally equally H and toward O

H is reducing O is oxidizing agent agent


Some reactions involving covalent reactants or products, the partial electron shifts are less obvious.

General guideline for covalent reactants or products:

• for carbon compounds, the addition of oxygen or the removal of hydrogen is always oxidation

Processes Leading to Oxidation & Reduction

Processes Leading to Oxidation & Reduction

Oxidation Reduction

Complete loss of electrons (ionic reactions)

Complete gain of electrons (ionic reactions)

Shift of electrons away from an atom in a covalent bond

Shift of electrons toward an atom in a covalent bond

Gain of oxygen Loss of oxygen

Loss of hydrogen by a covalent compound

Gain of hydrogen by a covalent compound

Increase in oxidation Increase in oxidation numbernumber

Decrease in oxidation Decrease in oxidation numbernumber

Corrosion

Iron, often used in the form of the alloy steel, corrodes by being oxidized to ions of iron by oxygen.

2Fe(s) + O2 (g) + 2H2O (l) 2Fe(OH)2 (s)

4Fe(OH)2(s) + O2 (g) + 2H2O (l) 4Fe(OH)3 (s)

• Equations describe the corrosion of iron to iron hydroxides in moist conditions.

Corrosion

Water in the environment accelerates the rate of corrosion.

Corrosion occurs more rapidly in the presence of salts and acids.

Salts and acids produce electrically conducting solutions that make electron transfer easier.

Resistance to Corrosion

Not all metals corrode easily.

Gold and platinum are called noble metals because they are very resistant to losing their e- by corrosion.

Other metals lose electrons easily but are protected from extensive corrosion by the oxide coating formed on their surface.

• Aluminum oxidizes quickly in air to form a coating of very tightly packed aluminum oxide particles.

Resistance to Corrosion

Iron also forms a coating when it corrodes

But the coating of iron oxide that forms is not tightly packed.

Water and air can penetrate the coating and attack the iron metal below it.

Corrosion continues until the iron object becomes only a pile of rust.

Controlling Corrosion

To prevent corrosion, the metal surface can be coated with oil, paint, plastic or another metal.

Coatings exclude air and water from the surface, preventing corrosion.

If coating is scratched or worn away, however, the exposed metal will begin to corrode.


Another method of corrosion control

One metal is “sacrificed” or allowed to corrode, in order to save a second metal.

To protect an iron object, a piece of magnesium (or other active metal) may be placed in electrical contact with the iron.

When oxygen and water attack the iron object, the iron atoms lose electrons as the iron being to be oxidized.


Because magnesium is a better reducing agent than iron and is more easily oxidized

the magnesium immediately transfers electrons to the iron, preventing their oxidation to iron ions.

Magnesium is “sacrificed” by oxidation and protects the iron in the process.


Sacrificial zinc and magnesium blocks are sometimes attached to piers and ship hulls to prevent corrosion damage in areas submerged in water.

Underground pipelines and storage tanks may be connected to magnesium block for protection

It is easier and cheaper to replace a block of magnesium or zinc than to replace a bride or a pipeline.

Question

Can you identify the common chemical characteristic of all metal corrosion?

The transfer of electrons from metals to oxidizing agents.

Questions

Define oxidation and reduction in terms of the gain or loss of oxygen.

Oxidation is the gain of oxygenReduction is the loss of oxygen

Define oxidation and reduction in terms of the gain or loss of electrons.

LEO the lion goes GER

Loss of Electrons is Oxidation Gain of Electrons is Reduction

Questions

What happens to the atoms in an iron nail that corrodes?

Iron atoms are oxidized when iron corrodes

How do you identify the oxidizing agent and the reducing agent in a redox reaction?

The species reduced is the oxidizing agent. The species oxidized is the reducing agent.

Questions

Use electron transfer or electron shift to identify what is oxidized and what is reduced in each reaction. (use electronegativity values for molecular compounds)

2Na(s) + Br2(l) 2NaBr(s)

Na oxidized, Br2 reduced

H2(g) + Cl2(g) 2HCl(g)

H2 oxidized, Cl2 reduced

Questions


2Li(s) + F2(g) 2LiF(s)

Li oxidized, F2 reduced

S(s) + Cl2(g) SCl2(g)

S oxidized, Cl2 reduced

Questions


N2(g) + 2O2(g) 2NO2(s)

N2 oxidized, O2 reduced

Mg(s) + Cu(NO3)2(aq) Mg(NO3)2(aq) + Cu(s)

Mg oxidized, Cu reduced

Questions

Identify the reducing agent and the oxidizing agent for each reaction.

2Na(s) + Br2(l) 2NaBr(s)

Na reducing agent, Br2 oxidizing agent

H2(g) + Cl2(g) 2HCl(g)

H2 reducing agent, Cl2 oxidizing agent

Questions


2Li(s) + F2(g) 2LiF(s)

Li reducing agent, F2 oxidizing agent

S(s) + Cl2(g) SCl2(g)

S reducing agent, Cl2 oxidizing agent

Questions


N2(g) + 2O2(g) 2NO2(s)

N2 reducing agent, O2 oxidizing agent

Mg(s) + Cu(NO3)2(aq) Mg(NO3)2(aq) + Cu(s)

Mg reducing agent, Cu oxidizing agent

End of Section 20.1

Oxidation Numbers

Oxidation numberOxidation number is a + or – number assigned to an atom to indicate its degree of oxidation or reduction.

General Rule

A bonded atom’s oxidation # is the charge that it would have if the e- in the bond were assigned to the atom of the more electronegative element.

Rules for Assigning Oxidation Numbers

1.The oxidation number of a monatomic ion is = in magnitude and sign to its ionic charge.

Bromide (Br1-) is -1 Iron III (Fe3+) is +3

2.The oxidation number of hydrogen in a compound is +1, except in metal hydrides, such as NaH, where it is -1

3.The oxidation number of oxygen in a compound is -2 except in peroxides, such as H2O2, where it is -1 and in compounds with the more electronegative fluorine, where it is positive.

Rules for Assigning Oxidation Numbers

4.The oxidation number of an atom in uncombined (elemental) form is 0.

Potassium metal (K) is 0 Nitrogen Gas (N2) is 0

5.For any neutral compound, the sum of the oxidation numbers of the atoms in the compound must equal 0

6.For polyatomic ion, the sum of the oxidation numbers must equal the ionic charge of the ion.

Some Thought

Determining oxidation numbers of elements in compounds is a way for chemists to keep track of electron transfer during redox reactions

What are other examples where items are numbered to keep track of movement?

The numbers on a sports player’s jersey

The area codes assigned to telephone numbers in different regions

Binary Ionic Compounds

In binary ionic compounds, such as NaCl and CaCl2, the oxidation numbers of the atoms equal their ionic charges.

Na1+ + Cl-1 NaCl oxidation # +1 -1 neutral

Ca2+ + Cl-1 CaCl2 oxidation # +2 -1 neutral

Note the sign I put before the oxidation number

Molecular Compounds

No ionic charges are associated with atoms of molecular compounds.

However, oxygen is reduced in the formation of water for example.

In water the two shared e- in the H – O bond are

shifted toward the O and away from the H.

Imagine the e- contributed by the two H atoms are completely transferred to the O.

Molecular Compounds

The charge that would result from the transfer are the oxidation numbers of the bonded elements.

The oxidation number of O is -2 and the oxidation number of each hydrogen is +1

Oxidation numbers are often written above the chemical symbols in a formula.

+1 -2

H2O

Multiple Oxidation Numbers

Many elements can have several different oxidation numbers.

+1 +6 -2

K2CrO4 – Potassium Chromate

+1 +12 -2

K2CrO7 – Potassium Dichromate

Sample Problems

What is the oxidation number of each kind of atom in the following ions and compounds?

+4 -2

SO2

+4 -2

CO32-

+1 +6 -2

Na2SO4

-3 +1 -2

(NH4)2S

Sample Problems

Determine the oxidation number of each element in the following.

+3 -2

S2O3

+1 -1

Na2O2

+5 -2

P2O5

+5 -2

NO3-

Sample Problems

Determine the oxidation number of chlorine in each of the following substances.

+1 +5 -2

KClO3

0

Cl2 +2 +7 -2

Ca(ClO4)2

+1 -2

Cl2O

Sample Problems

What are the oxidation numbers of iodine in the following?

+1 +7 -2

HIO4

+1 +5 -2

HIO3

+1 +1 -2

HIO 0

I2

+1 -1

HI

Oxidation Number Changes

When copper wire is placed in a solution of silver nitrate the following reaction occurs

+1 +5 -3 0 +2 +5 -2 0

2AgNO3(aq) + Cu(s) Cu(NO3)2(aq) + 2Ag(s)

Ag is reduced from +1 to 0

Cu is oxidized from 0 to +2

Oxidation Number Changes

An increase in the oxidation number of an atom or ion indicates oxidation.

A decrease in the oxidation number of an atom or ion indicates reduction.

+2 +6 -2 0 0 +2 +6 -2

2CuSO4(aq) + Fe2(s) 2Cu(s) + 2FeSO4

Cu is reduced from +2 to 0

Fe is oxidized from 0 to +2

Sample Problem

Identify which atoms are oxidized and which are reduced in the following reaction. Also identify the oxidizing agent and the reducing agent.

+1 -1 0 +1 -1 0

2HBr(aq) + Cl2(g) 2HCl(aq) + Br2(l)

Cl is reduced from 0 to -1, so Cl2 is the oxidizing agent

Br is oxidized from -1 to 0, so Br1- is the reducing agent

Sample Problem

Identify which atoms are oxidized and which are reduced in each reaction

0 0 +1 -2

O2(g) + 2H2(g) 2H2O(l)

O2 is reduced from 0 to -2

H2 is oxidized from 0 to +1

Sample Problem

Identify which atoms are oxidized and which are reduced in each reaction

+1 +5 -2 +1 +3 -2 0

2KNO3(s) 2KNO2(s) + O2(g)

N is reduced from +5 to +3

O is oxidized from -2 to 0

Sample Problem

Identify the oxidizing agent and the reducing agent in each equation.

0 0 +1 -2

O2(g) + 2H2(g) 2H2O(l)

O2 is reduced, thus O2 is the oxidizing agent

H2 is oxidized, thus H2 is the reducing agent

Sample Problem

Identify the oxidizing agent and the reducing agent in each equation.

+1 +5 -2 +1 +3 -2 0

2KNO3(s) 2KNO2(s) + O2(g)

N is reduced, thus N is the oxidizing agent

O is oxidized, thus O is the reducing agent

Questions

What is the general rule for assigning oxidation numbers?

The oxidation number is the charge a bonded atom would have if the electrons in the bond were

assigned to the more electronegative element

How is a change I oxidation number related to the process of oxidation and reduction?

An increase in oxidation number indicates oxidation; a decrease in oxidation number indicates reduction.

Sample Problem

Identify which atoms are oxidized and which are reduced in each reaction. Also identify the oxidizing agent and the reducing agent.

0 0 +1 -2

2Na(s) + Cl2(g) 2NaCl(s)

Cl2 is reduced, thus Cl2 is the oxidizing agent

Na is oxidized, thus Na is the reducing agent

Sample Problem

Identify which atoms are oxidized and which are reduced in each reaction. Also identify the oxidizing agent and the reducing agent.

0 0 +1 -2

2HNO3(aq) + 6HI(aq) 2NO(g) + 3I2(s) + 4H2O(l)

O2 is reduced, thus O2 is the oxidizing agent

H2 is oxidized, thus H2 is the reducing agent

End of section 20.2

Identifying Redox Reactions

In general, all chemical reaction can be assigned to one of two classes

1. Redox reactions in which electrons are transferred from one reacting species to another. a. Many single-replacement reactions,

combination reactions, decomposition reactions and combustion reactions are redox reactions.

2. All other reactions in which no electron transfer occurs. a. Double-replacement reactions and acid-base

reactions are not redox reactions

What Kind of Reaction Is It?

2Mg(s) + O2(g) 2MgO(s)

Combination reaction – two or more substances react to form a single new substance

2HgO (s) 2Hg (l) + O2 (g)

Decomposition reaction – a single compound breaks down into two or more simpler products.

Many are redox reactions


Zn(s) + Cu(NO3)2(aq) Cu(s) + Zn(NO3)2 (aq)

Single Replacement Reaction – one element replaces a second element in a compound.

2C6H18(l) + 25O2 (g) 16CO2(g) + 18H2O(l)

Combustion Reactions – an element or a compound reacts with oxygen often producing energy in the form of heat and light..

Many are redox reactions


Na2S(aq) + Cd(NO3)2(aq) CdS(s) + 2NaNO3(aq)

Double Replacement Reaction – involving an exchange of positive ions between two compounds.

2C6H18(l) + 25O2 (g) 16CO2(g) + 18H2O(l)

Combustion Reactions – an element or a compound reacts with oxygen often producing energy in the form of heat and light..

Many are not redox reactions

Identifying Redox Reactions

If the oxidation number of an element in a reacting species change, then that element has undergone either oxidation or reduction.

Many reactions in which color changes occur are redox reactions.

Balancing Redox Equations

Many redox reaction are too complex to be balanced by trial and error.

Two systematic methods are available to balance redox reactions

The two methods are based on the fact that the total number of electrons gained in reduction must equal the total number of electrons lost in oxidation.

One method used oxidation number changes, and the other used half reactions.

Using Oxidation Number Changes to Balance Redox Equations

Oxidation Number Method

You balance by comparing the increases and the decreases in oxidation numbers.

+3 -2 +2 -2 0 +4 -2

Fe2O3(s) + CO(g) Fe(S) + CO2(g)

Step 1 – assign oxidation numbers to all the atoms in the equation.



C oxidized (+2)

+3 -2 +2 -2 0 +4 -2


Fe reduced (-3)

Step 2 – Identify which atoms are oxidized and which are reduced


Oxidation Number Method (+2) oxidized

+3 -2 +2 -2 0 +4 -2


(-3) reduced

Step 3 – Use one bracketing line to connect the atoms that undergo oxidation and another such line to connect those that undergo reduction.


In a balanced redox equation, the total increase in oxidation number of the species oxidized must be balance by the total decrease in the oxidation number of the species reduced.


Oxidation Number Method 3 x (+2) = +6

+3 -2 +2 -2 0 +4 -2

Fe2O3(s) + 33CO(g) 22Fe(S) + 33CO2(g)

2 x (-3) = -6

Step 4 – Make the total increase in oxidation number equal to the total decrease in oxidation number by using appropriate coefficients.



Fe2O3(s) + 3CO(g) 2Fe(S) + 3CO2(g)

Step 5 – Make sure that the equation is balanced for both atoms and charge.

Sample Problem


S oxidized (+4) x 3= +12

+1 +6 -2 +1 -2 0 +1 -2 +1 +3 -2 +4 -2

22K2Cr22O7(aq) + H2O(l) + 33S(s) KOH(aq) + 22Cr22O3(s) + 33SO2 (g)

Cr reduced (-3) x 4 = -12

4 Cr atoms must be reduced for each 3 S atom that are oxidized

Sample Problem


S oxidized (+4) x 3= +12

+1 +6 -2 +1 -2 0 +1 -2 +1 +3 -2 +4 -2

2K2Cr2O7) + 22H2O + 3S 44KOH + 2Cr2O3 + 3SO2

Cr reduced (-3) x 4 = -12

Check the equation and balance by inspection

4 in front of KOH balances potassium2 in front of H2O balances hydrogen and oxygen.

Sample Problems

Balance each redox equation using the oxidation number change method

KClO3(s) KCl (s) + O2 (g)

2KClO3(s) 2KCl (s) + 3O2 (g)

HNO2 (aq) + HI (aq) NO (g) + I2 (s) + H2O (l)

2HNO2 (aq) + 2HI (aq) 2NO (g) + I2 (s) + 2H2O (l)

Sample Problems

Balance each redox equation using the oxidation number change method

Bi2S3 + HNO3 Bi(NO3)3 + NO + S + H2O

Bi2S3 + 8 HNO3 2Bi(NO3)3 + 2NO + 3S + 4H2O

SbCl5 + KI SbCl3 + KCl + I2

SbCl5 + 2KI SbCl3 + 2KCl + I2

Using Half Reactions to Balance Redox Equations

Half Reaction Method

Half reaction is an equation showing just the oxidation or just the reduction that takes place

You write and balance the oxidation and reduction half reaction separately before combining them into a balanced redox equation

Then you balance the electrons gained in the reduction with the electrons lost in the oxidation.



S + HNO3 SO2 + NO + H2O (unbalanced)

S + H+ + NO3- SO2 + NO + H2O

In this case only HNO3 is ionized. The products are covalent compounds

Step 1 – write the unbalanced equation in ionic form



Oxidation Half Reaction 0 +2

S SO2

Reduction Half Reaction +5 +2

NO3- NO

Step 2 – Write separate half reactions for the oxidation and reduction processes.


Half Reaction Method0 +2

S SO2

Sulfur is already balanced, but oxygen is not.

The reaction takes place in acid solution, so H2O and H+ are present and can be used to balance oxygen and hydrogen as needed.

2H2H22OO + S SO2 + 4H4H++

Step 3 – Balance the atoms in the half reactions.


Half Reaction Method +5 +2

NO3- NO

Nitrogen is already balanced, but oxygen is not.

4H4H++ + NO3- NO + 2H2H22OO



0 +4

2H2O + S SO2 + 4H+ + 4e4e - -

S is oxidized going from 0 to +4, a loss of 4 e-

+5 +2 3e3e-- + 4H+ + NO3

- NO + 2H2O

N is reduced going from +5 to +2, a gain of 3 e-

Step 4 – Add enough electrons to one side of each half reaction to balance the charges.


Half Reaction Method 0 +4

(x 3)(x 3) 6H2O + 3S 3SO2 + 12H+ + 12e -

S is oxidized going from 0 to +4, a loss of 4 e-

+5 +2 (x 4) 12(x 4) 12e-- + 16H+ + 4NO3

- 4NO + 8H2O

N is reduced going from +5 to +2, a gain of 3 e-

Step 5 – Multiply each half reaction by an appropriate number to make the numbers of electrons equal in both.



6H2O + 3S 3SO2 + 12H+ + 12e -

12e- + 16H+ + 4NO3- 4NO + 8H2O

6H2O + 3S + 16H+ + 4NO3- + 12e - 3SO2 + 12H+ +

4NO + 8H2O + 12e –

3S + 4H+ + 4NO3- 3SO2 + 4NO + 2H2O

Step 6 – Add the balanced half reactions to show an overall equation and then subtract the terms that appear in both sides of the equation



3S + 4HNO3- 3SO2 + 4NO + 2H2O

Spectator ionSpectator ion – are present during a reaction, but do not participate in or change during a reaction.

Because none of the ions in the reactants appear in the products, there are no spectator ions in this particular example.

Step 7 – Add the spectator ions and balance the equation.

Sample Problem


KMnO4 + HCl MnCl2 + Cl2 + H2O + KCl

K+ + MnO4- + H+ + Cl- Mn2+ + 2Cl- + Cl2 + H2O + K+ + Cl-

Step 1 – write the unbalanced equation in ionic form

Sample Problem



MnO4- Mn2+

Oxidation Half Reaction -1 0

2Cl- Cl2

Step 2 – Write separate half reactions for the oxidation and reduction processes.

Sample Problem


8H+ + MnO4- Mn2+ + 4H4H22OO

Oxidation Half Reaction -1 0

2Cl- Cl2

Solution is acidic, so H2O and H+ ions are used to balance equation. (If solution is basic, H2O and OH- are used)

Step 3 – Balance the atoms in the half reactions.

Sample Problem


+7 +2

5e5e-- + 8H+ + MnO4- Mn2+ + 4H2O

Mn is reduced going from +7 to +2, a gain of 5 e-

-1 0

2Cl- Cl2 + 2e-2e-

Cl is oxidized going from -1 to 0, a loss of 2 e-

Step 4 – Add enough electrons to one side of each half reaction to balance the charges.

Sample Problem

Half Reaction Method +7 +2

(x 2)(x 2) 10e- + 16H+ + 2MnO4- 2Mn2+ + 8H2O

Mn is reduced going from +7 to +2, a gain of 5 e-

-1 0

(x5)(x5) 10Cl- 5Cl2 + 10e-

Cl is oxidized going from -1 to 0, a loss of 2 e-

Step 5 – Multiply each half reaction by an appropriate number to make the numbers of electrons equal in both.

Sample Problem


10e- + 16H+ + 2MnO4- 2Mn2+ + 8H2O

10Cl- 5Cl2 + 10e-

10e- + 16H+ + 2MnO4- + 10Cl- 2Mn2+ + 8H2O + 5Cl2

+ 10e-

16H+ + 2MnO4- + 10Cl- 2Mn2+ + 8H2O + 5Cl2

Step 6 – Add the balanced half reactions to show an overall equation and then subtract the terms that appear in both sides of the equation

Sample Problem


16H+ + 6Cl- + 2MnO4- + 2K+ + 10Cl-

5Cl2 + 2Mn2+ + 4Cl- + 8H2O + 2K+ + 2Cl-

16H+ = 6Cl- + 10Cl- (from the HCl in original equation)

2MnO4- = 2K+ (from the KMnO4 in original equation)

2Mn2+ = 4Cl- (from the MnCl2 in original equation)

2K+ + 2Cl- (from the KCl in original equation)


Sample Problem


16H+ + 16Cl- + 2MnO4- + 2K+

5Cl2 + 2Mn2+ + 6Cl- + 8H2O + 2K+

Adding spectator and non-spectator Cl- on each side gives

16HCl + 2KMnO4 5Cl2 + 2MnCl2 + 8H2O + 2KCl


Sample Problem 2

The following takes place in basic solution. Balance the equation using the half reaction method

Zn + NO3

- NH3 + Zn(OH4)2-

4Zn + NO3- + 6H2O + 7OH-

4Zn(OH4)2- + NH3

Sample Problem 3

The following takes place in basic solution. Balance the equation using the half reaction method

Zn + As2O3 AsH3 + Zn2+

6Zn + As2O3 + 9H2O 6Zn2+ + 2AsH3 + 12OH-

Choosing a Balancing Method

In some redox reactions, the same element is both oxidized and reduced. (called a disproportional reaction)

3Br2 + 6KOH 5KBr + KBrO3 + 3H2O

Br is reduced from 0 to -1Br is oxidized from 0 to + 5

Equations like above and equation for reactions that take place in acidic or alkaline solution are best balanced by the half reaction method.

End of Chapter 20

chapter 20

Education

oxygen gains electrons

oxygen atoms

shift of electrons

gain of oxygen

removal of oxygen

electronswhen oxygen

exchange of electrons

partial loss of electrons