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Chapter 19
Oxidation and Reduction


I: Oxidation and Reduction

transfer of electrons

oxidation = loss of electrons

reduction= gain of electrons

both oxidation and reduction must occur together in the same reaction

use oxidation numbers to tell us about the oxidation states

A: Oxidation States

electron configuration determines the oxidation number i.e. determines the number of electron to be lost of gained to achieve an octet or reach a more stable electron configuration

Table 19-1 page 591 Rules for Assigning Oxidation Numbers

Figure 19-1 page 592 - Colors of chromium solutions in different oxidation states

Homework 19.1

B: Oxidation

oxidation process - definition

Figure 19-2 page 592 - reaction of sodium and chlorine

The reaction is
2 Na(s) + Cl2(g) ---> 2 NaCl(s)

and the oxidation process can be written as
Na0 ---> Na+1 + 1 e-1

sodium change from a zero oxidation state to a +1 oxidation state i.e. lost one electron

oxidized - definition

C: Reduction

reduction is the gain of electrons

Cl2 + 2 e-1 ---> 2 Cl-1

reduced - definition

D: Oxidation and Reduction as a Process

Not only do oxidation and reduction both have to occur simultaneously in a reaction, they have to occur to the same extent - the number of electrons lost has to equal the number of electrons gained.

This follows from the law of conservation of mass.

oxidation reduction reaction - definition

Figure 19-3 page 593 example of redox reaction


If a chemical reaction occurs and no elements gain or lose electrons (do not change oxidation state from reactants to products) then the reaction is not a redox reaction.

1. Redox Reactions and Covalent Bonds

An ion has a charge because it has gained or lost electrons. But in a covalent compound the atoms do not have a charge but they do have a tendency to share a certain number of electrons based on their electron configuration and this tendency is related to their electronegativity as well.

For the reaction

H2 + Cl2 ---> HCl

On the left side of the equation, hydrogen has an oxidation number of zero as does chlorine, since both are diatomic and behave like free elements.
On the right side of the equation, hydrogen has an oxidation number of positive one and chlorine has an oxidation number of negative one. Chlorine is negative because it has the higher electronegativity.

The oxidation half reaction is

H0 ---> H+1 + 1e-1

Since hydrogen is diatomic in our equation, we modify the oxidation half reaction to read

H20 ---> 2 H+1 + 2 e-1

Our reduction half reaction will be

Cl0 + 1 e-1 ---> Cl-1

Since chlorine is diatomic it becomes

Cl20 + 2 e-1 ---> 2 Cl-1

Combining the two we get

H20 ---> 2 H+1 + 2 e-1
Cl20 + 2 e-1 ---> 2 Cl-1

Canceling out the equal number of electrons on both sides we get
H20 + Cl20 --->2 H+1 + 2 Cl-1

Then we transfer the number of atoms to our original equation.

H2 + Cl2 ---> 2 HCl

Since our equation contained only atoms that changed oxidation number, we did not have to balance any other elements.


Another example:

KNO3 ---> KNO2 + O2

nitrogen goes from +5 to +3
oxygen goes from -2 to 0

the half reactions will be
N+5 + 2 e-1 ---> N+3
O-2 ---> O0 +2e-1

Oxygen is diatomic so we adjust its half reaction to be
2O-2 ---> O20 +4e-1
N+5 + 2 e-1 ---> N+3

make the number of electrons equal we get
2O-2 ---> O20 +4e-1
2N+5 + 4 e-1 ---> 2N+3

combine the equations algebraically we get

2O-2 + 2N+5 ---> O20 + 2N+3

Placing the proper number of each atom into the original equation we get

2 KNO3 --->2 KNO2 + O2 then sight balance the equation

2 KNO3 --->2 KNO2 + O2

Another example:

H2S + HNO3 ---> H2SO4 + NO2 + H2O

sulfur goes from a -2 to a +6
S-2 ---> S+6 - 8e-1

nitrogen goes from a +5 to a +4
N+5 + 1e-1 ---> N+4

combining the two equations we get
S-2 ---> S+6 - 8e-1
N+5 + 1e-1 ---> N+4

make the number of electrons equal
S-2 ---> S+6 - 8e-1
8 N+5 + 8e-1 ---> 8 N+4

Algebraically combine the two equations to get
S-2 + 8 N+5 ---> S+6 + 8 N+4

Transfer the correct number of each atom to the original equation:

H2S + 8 HNO3 ---> H2SO4 + 8 NO2 + H2O then sight balance the remaining elements to get

H2S + 8 HNO3 ---> H2SO4 + 8 NO2 + 4 H2O

Homework 19.2

II: Oxidizing and Reducing Agents

Reducing agent - definition

lose electrons and become more positive i.e. it is oxidized

Oxidizing agent - definition

gains electrons and becomes more negative i.e. it is reduced

Table 19-2 page 602 Oxidation-Reduction Terminology

A: Strengths of Oxidizing and Reducing Agents


Table 19-3 Relative Strengths of Oxidizing and Reducing Agents

Fluorine is most electronegative and has the greatest attraction for electrons (wants to gain electrons) therefore it is a the best oxidizing agent on the list. It also makes it the weakest reducing agent on the list.

Positive ion of a strong reducing agent, e.g. Li, is a weak oxidizing agent, e.g. Li+1.

Metals would tend to lose electrons and be good reducing agents. If Zn is above Cu on the reducing agent column of the table, that means that zinc is a better reducing agent than copper or zinc loses electrons more readily than copper.

If you put zinc in contact with copper ions, the zinc will give it electrons to the copper ions; copper will become the copper element and zinc will become the zinc ion.

Zinc in copper solution movie.


Oxidation and reduction of copper metal movie.

http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/redox/home.html

http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/animations/ZnCutransfer.html

Figure 19-5 Zinc reacting with copper ions

As the reaction proceeds the blue color of the copper solution will fade, indicating that the number of copper ions is decreasing.

Every redox reaction has an oxidizing and a reducing agent.

A link to illustrate how the activity series works is at Iowa State University:
http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/redox/home.html

B: Autooxidation

In some reactions the same element is both oxidized and reduced e.g.

2 H2O2 ---> 2 H2O + O2

oxygen in peroxide is -1 and it produces an oxygen in water which is -2 and the diatomic oxygen which is zero.

Autooxidation - definition

Homework 19.3

III: Electrochemistry

Electrochemistry - definition

If we mix the reactants of a redox reaction in the same container, we would produce heat energy. If we keep them separated and join the two containers, we can produce electricity.

A: Electrochemical Cells

A link to some interesting animations for electrochemistry at Iowa State University is
http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/animationsindex.htm

Figure 19-7 heat energy produced by the direct mixing of two reactants in a redox reaction.

Figure 19-8 electrochemical cell

Parts are
anode
cathode
two solutions, each of which is an electrolyte
a porous barrier
a wire connecting the two electrodes

The porous barrier allows ions to pass through but prevents the wholesale mixing of the two solutions.

As electrons move through the wire from one electrode to the other, ions pass from one side of the porous barrier to the other.

Notation for writing electrochemical cell:
Zn|Cu

Two types of electrochemical cells:
a) voltaic (galvanic)
b) electrolytic

1. Voltaic Cells

Voltaic cell - definition

Figure 19-9 page 608 - Voltaic Cell of copper and zinc

Reduction half reaction for figure 19-9 is:
Cu+2 (aq) + 2e-1 ---> Cu(s)

and the oxidation half reaction is:
Zn(s) ---> Zn+2(aq) + 2e-1

According to table 19-3 page 603, zinc is a better reducing agent than copper i.e. it is oxidized more easily than copper. That means that zinc is oxidized and copper is reduced. Way to predict which substance is oxidized and which is reduced.

If zinc is losing electrons it means that electrons are moving from the zinc electrode through the wire to the copper electrode. The copper electrode takes on a negative charge which attracts the copper ions in solution to that electrode. The copper ions take on electrons and are deposited at the copper electrode i.e. the copper electrode is being plated or gaining more copper atoms.

At the zinc electrode, the zinc that makes up the electrode is losing electron which means that the zinc atoms of the electrode are becoming zinc ions which will then go into the solution around the zinc electrode. That means the zinc electrode is being used up or losing zinc atoms.

Since the solution around the zinc electrode is increasing in Zn+2 ions, those ions have to be neutralized. the sulfate ions from the copper side, migrate through the porous barrier and neutralize the zinc ions. The sulfate ions can migrate because the copper ions that brought the sulfate ions are being reduced at the copper electrode.

Link at Iowa State University to illustrate a voltaic cell:
http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/electroChem/volticCell.html

Figure 19-10 page 608 Picture of dry cells which are forms of voltaic cells.

Three common types of dry cells are:
a) zinc-carbon
b) alkaline
c) mercury

Figure 19-11 page 609 - cutaway of zinc-carbon dry cell

a) Zinc-Carbon Dry Cells

Figure 19-11 page 609 Cutaway of zinc-carbon dry cell

zinc container - anode
moist paste of MnO2, graphite, and NH4Cl
zinc atoms are oxidized at the negative electrode or anode: Zn(s) ---> Zn+2 + 2e-1

electrons move through the circuit 0 - battery to bulb back to battery - and reenter the battery through the carbon rod which acts as the cathode or positive electrode. The manganese dioxide is reduced:
2 MnO2(s) + H2O(l) ----> Mn2O3(s) + 2 OH-1(aq)

b) Alkaline Batteries

Figure 19-12 page 609 Cutaway of alkaline battery

Paste of zinc metal with KOH instead of the carbon electrode.

At the anode the reaction is
Zn0(s) + 2 OH-1(aq) ---> Zn(OH)2(s) + 2e-1
the anode is where the oxidation takes place

At the cathode, where reduction occurs,
2 MnO2(s) + H2O(l) ----> Mn2O3(s) + 2 OH-1(aq)
This is the same reaction as in the zinc-carbon dry cell.

c) Mercury Batteries

Figure 19-13 page 610 Cutaway of a mercury battery

small size

anode half reaction is same as in the alkaline battery i.e. Zn0(s) + 2 OH-1(aq) ---> Zn(OH)2(s) + 2e-1

The cathode reaction is
HgO(s) + H2O(l) + 2e-1 ---> Hg(l) + 2 OH-1(aq)

2. Electrolytic Cells

If a redox reaction will not occur spontaneously, we can make it happen by supplying electricity.

Electrolysis - definition

Electrolytic cell - definition

Figure 19-14 A voltaic cell v an electrolytic cell.

The electrode connected to the negative terminal of the battery becomes negative and is the cathode of the cell.

The electrode of the cell connected to the positive terminal of the battery loses electrons and passes them on to the battery which acts like a pump. This electrode is the anode.

In figure 19-14 the copper electrode is losing electrons and the reaction is
Cu0 --- Cu+2 + 2e-1

The zinc electrode is gaining electrons and the reaction is
Zn+2 + 2e-1 ---> Zn0

Voltaic v Electrolytic Cell
a) battery in an electrolytic cell; no battery in a voltaic cell; voltaic cell produces electricity while the electrolytic cell uses electricity.
b) reaction in voltaic cell is spontaneous and produces electricity; reaction in electrolytic cell is not spontaneous and uses electricity.

a: Electroplating

Electroplating - definition

Figure 19-15 page 612 Electroplating cell

To deposit silver on an object you need silver ions in solution and the object must be the electrode connected to the negative terminal of the battery. This will produce electrons on the object which will attract the positive ions in solution to the object where the ions will be reduced and plate the metal on the object e.g.

Ag+1 + 1e-1 ---> Ag0

The other electrode must consist of a pure piece of the metal to be plated. As electrons are removed from this electrode, the metal is converted to ions which go into solution on that side of the cell. The electrons move through the circuit, through the battery and collect on the object to be plated. The pure metal is converted to ions by the reaction


Ag0 ---> Ag+1 + 1e-1

The ions go into solution and migrate through the porous barrier to replace the silver ions that are plating on the other side of the cell.

b: Rechargeable Cells

A rechargeable cell produces electricity until the reaction is exhausted. Then the recharging uses electricity to reverse the reaction and store electricity in the chemicals again.

Such is the case in a car battery.

Figure 19-16 page 613 A 12 v automobile battery.

Six rechargeable cells.

Lead submersed in sulfuric acid serves as the anode and the reaction is

Pb(s) + SO4-2(aq) ---> PbSO4(s) + 2e-1

At the cathode PbO2 is reduced and the reaction is

PbO2(S) + 4 H+(aq) + SO4-2(aq) + 2 e-1 ---> PbSO4(s) + 2 H2O(l)

The net reaction is an algebraic combination of the two half cell reactions.

When the car is started the battery behaves like a voltaic cell by supplying electricity and while you drive the car it is being recharged by the alternator and so is behaving like an electrolytic cell.

As long as all the reactants of both forward and reverse reactions are still present in the proper amounts the battery operates.

3: Electrode Potentials

Look at figure 19-9 on page 608. The zinc and copper electrodes have different tendencies to accept electrons.

Verified by checking table 19-3 on page 603.

Reduction potential - definition

Each electrode is dipping into a solution of its ions: copper in a solution of copper ions and zinc in a solution of zinc ions.

electrode potential - definition

Connect the two half cells and the difference in potential causes electrons to move from one to the other. The difference in potential is measured in volts. For our cell the potential difference is 1.10 v when the solution concentrations are 1 M each.

We measure this difference in potential by using a voltmeter. We cannot measure the potential of an individual cell, only a cell connected to another cell to make up a full voltaic cell. Only when they are connected are electrons able to flow because of the potential difference.

We use a standard half cell as a reference. See figure 19-17 page 614, hydrogen electrode.

The standard hydrogen electrode (SHE) consists of a platinum electrode dipped into a 1.00 M acid solution surrounded by hydrogen gas at 1 atm pressure and 25 degrees Celsius.

Other half cells are connected to the SHE and we measure their ability to reduce hydrogen under these conditions.

The anode reaction (oxidation) for the SHE is

H2(g)0 = 2 H+1(aq) + 2e-1

The cathode reaction for SHE is the reverse of the above.

The value of 0.00v is assigned to each of the reactions of the SHE.

When we attach another half cell to the SHE any voltage on the voltmeter is due to the other half cell, not the hydrogen half cell.

Standard electrode potential has the symbol, Eo.

Standard electrode potential - definition

Electrode potential are expressed as potentials for reduction.

Indicate the tendency of a substance to be reduced.

Table 19-4 page 615, Standard Reduction Potentials

A positive E0 indicates that hydrogen is more willing to give up its electrons than the other electrode. Half reactions with positive reduction potentials are favored.

Effective oxidizing agents (they are reduced) have positive values. e.g. Copper and fluorine.

The opposite is true for those with negative value. These prefer to be oxidized than reduced. It indicates that the metal is more willing to give up electrons than is hydrogen. e.g. Li and Zn.

Since table 19-4 show all reactions as reductions and all redox reaction must involve both oxidation and reduction, we must reverse the half reaction of one of the cells. When we do this we need to reverse the sign of the electrode potential.

e.g.
Zn+2 + 2 e-1 ---> Zn Eo = -0.76v
Zno ---> Zn+2 + 2 e-1 Eo = +0.76v

To measure a half cell’s reduction potential, connect it to the SHE half cell.

In figure 19-18 page 614, we see that the voltage for the zinc half cell connected to the SHE is -0.76v. That tells us that the zinc does not want to be reduced, but oxidized and it loses electrons to the hydrogen which is reduced.

In figure 19-18 page 614, we see that the voltage for the copper half cell connected to the SHE is +0.34v. That tells us that the copper wants to be reduced, and it receives electrons from the hydrogen half cell to do this.

We use standard electrode potentials to predict of a redox reaction will occur spontaneously (naturally). The cell potential will have a positive value if it is spontaneous. Eocell = Eocathode +Eoanode

The half reaction with the more negative standard reduction potential is the anode (where oxidation takes place). We reverse this reaction and also the sign of the standard reduction potential. The other half reaction remains the same as does its sign. We algebraically combine the two half reactions and their standard reduction potential to calculate the cell reaction and the cell potential.

Problem:
Given the two reduction half reactions, calculate the cell potential and the cell reaction:

Fe+3 + 3e-1 ---> Fe Eo = -0.04v
Ag+1 + 1 e-1 ---> Ago Eo = +0.80v

We reverse the reaction that has the lower potential - the iron reaction and we also reverse the sign of the potential. We leave the other reaction as is.

Fe ---> Fe+3 + 3 e-1 Eo = +0.04v
Ag+1 + 1 e-1 ---> Ago Eo = +0.80v

Make the number of electrons lost equal to the number of electrons gained by multiplying the appropriate equation accordingly.

Fe ---> Fe+3 + 3 e-1 Eo = +0.04v
3(Ag+1 + 1 e-1 ---> Ago) Eo = +0.80v

Fe ---> Fe+3 + 3 e-1 Eo = +0.04v
3 Ag+1 + 3 e-1 ---> 3 Ago Eo = +0.80v

Combine the two reactions algebraically:
Fe + 3 Ag+1 ---> Fe+3 + 3 Ago Eo = +0.84v

The above is the cell reaction and the cell potential when you combine a silver and iron half cell.

The reverse of the above reaction would not occur in a voltaic cell but we could use it in an electrolytic cell using electricity.





end of notes

Oxidation processes are reactions in which the atoms or ions of an element experience an increase in oxidation state i.e. the oxidation number increases.back

A species whose oxidation number increases is oxidized.back

Reactions in which the oxidation state of an element decreases are reduction processes.back

A species that undergoes a decrease in oxidation state is reduced.back

Any chemical process in which elements undergo changes in oxidation number is an oxidation reduction reaction or a redox reaction.back

A reducing agent is a substance that has the potential to cause another substance to be reduced (while it is oxidized).back

An oxidizing agent is a substance that has the potential to cause another substance to be oxidized (while it is reduced).back


Autooxidation is a process in which a substance acts as both an oxidizing agent and a reducing agent.back

Electrochemistry is the branch of chemistry that deals with electricity related applications of redox reactions.back

An electrode is a conductor used to establish electrical contact with a nonmetallic part of a circuit, such as and electrolyte.back

A single electrode immersed in a solution of its ions is a half cell.back

An anode is the electrode where oxidation takes place.back

A cathode is the electrode where reduction takes place.back

An electrochemical cell is a system of electrodes and electrolytes in which either chemical reactions produce electrical energy or an electric current produces chemical change.back


Voltaic cell is a redox reaction in an electrochemical cell that occurs spontaneously and produces electrical energy.back

Electrolysis is the process in which an electric current is used to produce a redox reaction.back

Electrolytic cell is where a chemical change occurs in an electrochemical cell in which electrical energy is required to produce the redox reaction.back

Electroplating is an electrolytic process in which a metal ion is reduced and a solid metal is deposited on a surface.back

Reduction potential is the tendency for the half-reaction of either electrode material to occur as a reduction half reaction in an electrochemical cell.back

Electrode potential is the difference in potential between an electrode and its solution.back

Standard electrode potential is the half cell potential measured relative to a potential of zero for the standard hydrogen electrode.back