L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
Chpt:20 p.1
REDOX EQUILIBRIA II
II.
ELETROCHEMICAL CELLS電化電池
(A)
Metal/Metal ion system — Qualitative study of Redox System
When a metal is dipped into浸入 a solution containing ions of the same metal, the following
three possible interactions may occur to different extents不同程度, depending on the nature of
the particular metal/metal ion system:
<i>
<ii>
A metal ion in solution may collide with the electrode but undergoes no change.
A metal ion in solution may collide the electrode, gain electrons and can be deposited
on the metal electrode. The metal ion is said to be reduced. The half equation is
Mn+(aq) + ne  M(s)
and can be used to represent this reduction reaction.
<iii>
A metal atom from the electrode may lose electrons and enter the solution as the metal
ion. The metal atom is said to be oxidized. The half equation is
M(s) 
Mn+(aq) + ne
and can be used to represent the oxidation reaction.
An equilibrium is established between the tendencies of <ii> and <iii>.
Mn+(aq) + ne  M(s)
In the forward process, that is the reduction process, the hydrated metal ions in solution consume
electrons from the electrode. This results in a net deficit缺少of electrons in the electrode and
thus a positive charge on the electrode.
On the other hand, in the reverse process, that is the. oxidation process, atoms from the metal
lattice pass into solution to form cations. This leaves a surplus過剩of electrons and results in a
negative charge on the electrode.
The charge on the electrode depends on which process occurs more readily. If
<a> the forward process is more favourable, reduction predominates and thus the electrode
acquires a positive charge.
<b> the reverse process is more favourable, oxidation predominates and thus the electrode
acquires a negative charge.
In either case, there is a separation of charge and thus a potential difference is set up between the
metal and its ions in the solution. The metal/metal ion system is known as a half cell.
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
Chpt:20 p.2
(B) Use of Metal/Metal Ion System as Sources of Electrical Energy
The combination of two metal/metal ion half-cells will give an electrochemical or voltaic cell.
An electrochemical cell is a device which produces an electromotive force (e.m.f.) 電動勢as a
result of chemical reactions taking place at the electrode.
An electrochemical cell is thus a device which convert chemical energy into electrical energy.
Each cell consists of two half cells. In one half-cell an oxidation half-reaction takes place. In the
other a reduction half-reaction takes place.
Example : The Daniell cell丹聶耳電池
Salt bridge
Zn
Cu
Cotton
wool
Daniell cell
In the Daniell cell, zinc and copper electrodes are dipped into a 1M solution of zinc sulphate and
copper(II) sulphate solution. The function of the porous partition of the salt bridge is to enable
electrical contact while separating the two different electrolytes, so that direct chemical redox
reaction
Cu2+(aq) + Zn(s)  Cu(s) + Zn2+(aq)
will not occur. Electrical contact between the solutions is established in the walls of the porous
partition, or through the electrolytes in the salt bridge.
Salt bridge : It can be made by an inverted U-tube containing saturated salt solution such as
potassium chloride, potassium nitrate or ammonium nitrate. I
When these two half cells are connected externally by a conducting wire, owing to the difference
in potentials between the half cells, an electrochemical reaction involving the passage of electrons
can take place.
<1> The zinc goes into the solution as zinc ions. Oxidation takes place.
_____________________________________________________
<2> The electrons produced flow in the external circuit to the copper electrode. In the copper
electrode, copper(II) ions in solution accept these electrons and are deposited.
_____________________________________________________
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
Chpt:20 p.3
The zinc electrode become the negative terminal of the cell whereas the copper electrode
becomes the positive terminal of the call. The Daniell cell therefore provides a means of
converting the chemical potential energy between the two metal/metal ion system into electrical
energy.
The enthalpy of the net reaction
Cu2+(aq) + Zn(s)  Cu(s) + Zn2+(aq)
is liberated 釋放 in the form of electrical energy as electron flow.
(C)
Potential difference of electrochemical cells
When a metal is placed in contact with a solution of one of its salts, an equilibrium is
established between the tendencies.
oxidation
M(s)

Mn+(aq) + ne
reduction
The potentials set up will differ for different metals. These potentials (absolute potentials)
cannot be measured, but the difference in potential between two metal/metal ion system can be
found by incorporating them into an electrochemical or voltaic cell, and measuring the potential
difference between the metal electrodes.
Note
<1> As the resistance in the external circuit is increased, the potential difference between the
two systems increases accordingly. whereas the current flowing through (measured by the
milliammeter) is decreases.
The Maximum potential difference of the whole cell is known as the electromotive force電
動勢 or e.m.f. and it occurs when no current is flowing.
The value of this e.m.f. is a measure of the relative tendencies of the electrode system
involved to liberate electrons by forming ions in solution.
<2> Such set up can be used to measure the e.m.f. values of different metal-metal ion systems,
and their value compared to assess the relative tendencies of metal/metal ion system to
release electrons.
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
(D)
Chpt:20 p.4
Cell Diagrams 電池圖
A cell digram can be used to represent the electrochemical cell and its e.m.f..
Example : The Daniell cell can be represented as follows:
Zn(s)  Zn2+ (aq)

Cu2+ (aq) Cu(s)
E = +1.1V
<1> The solid vertical lines represent boundaries between the electrode and its solution in each
electrode system.
<2> The dotted lines represent the salt bridge or porous pot. Usually,  represents a salt bridge
whereas  represents a porous
<3> E represent the e.m.f. of the cell in volts. The (+) or (-) sign indicates the polarity of the
right hand electrode. In the above case, copper electrode is the positive pole of the cell.
<4> By convention, the half-cell with the cathode陰極 (the positive electrode) is placed on the
right. This is the half-cell in which reduction takes place.
Thus, Cu2+(aq) Cu(s) represents Cu2+(ag) + 2e  Cu(s)
The half-cell with the anode 陽極 (the negative electrode) is placed on the left. This is the
half-cell in which oxidation takes place.
Thus, Zn(s)  Zn2+(aq) represents Zn(s)  Zn2+(aq) +2e
(E)
Types of half-cells
The three most common types are metal-metal ion, non-metal-ion and ion-ion half cells.
1. The Metal-Metal Ion Half-Cell
The zinc-zinc ion and copper-copper(II) ion electrodes.
2. The Non-Metal-Ion Half-Cell
The iodine-iodide ion electrodes (with graphite as the inert electrode). The following
equilibrium is established.
I2(aq) + 2e  2I-(aq)
Cell diagram : <a> As a cathode
<b> As an anode
3.
The Ion-Ion Half-Cell
The iron(III)-iron(II) half cell. In this the following equilibrium is established.
Fe3+ (aq) + e

Cell diagram : <a> As a cathode
Fe2+ (aq)
<b> As an anode
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
III. ELECTRODE
(A)
Chpt:20 p.5
POTENTIALS 電極電勢
Standard Electrode Potential標準電極電勢
If the metal is immersed in a solution of its ions of concentration 1 mol dm3 at 25°C, then the
potential acquired under these standard conditions if the standard electrode potential of the
metal. E0.
Examples : E0Cu = +0.34V
Note:
(B)
E0Zn = -0.76V
Since the electrode potential depend on temperature, concentration and also pressure, it
is necessary to standardize them if they are to be compared.
Finding Electrode Potentials
If a cell is constructed from a standard electrode (i.e., one of known potential) and an electrode
of unknown potential, the e.m.f. of the cell can be used to find the unknown electrode potential.
For a cell with the standard electrode on the left- hand side
E0cell = Eelectrode of unknown potential - Estandard electrode
(C)
Standard Hydrogen Electrode 標準氫電極
The standard hydrogen electrode is the reference electrode with which other electrodes are
compared. It consists of a platinised platinum electrode immersed in a solution of 1 mol dm-3
hydrogen ions. Hydrogen gas at a pressure of 1 atm is bubbled over the platinum electrode. On
the surface of the platinum, equilibrium is established between hydrogen gas and hydrogen ions.
2H3O+(aq) + 2e 
2H+(aq) + 2e 
2H2O(l) + H2(g) or simply
H2(g)
The potential is assigned a value of 0 volts.
In a cell diagram,
a hydrogen electrode is represented as
____________________________
if platinum is the anode
____________________________
if platinum is the cathode
The standard electrode potentials of other system
can be found by combining these with a
standard hydrogen electrode
and measuring the e.m.f. of the cell formed.
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
Example :
Chpt:20 p.6
A voltaic cell which combines a standard zinc electrode and a standard hydrogen
electrode.
The two compartments are connected by a salt bridge The e.m.f. of the cell is -0.76V. The
voltmeter shows that electron flow through the external circuit from zinc to the hydrogen
electrode, showing that zinc has a standard electrode potential of -0.76V
That is Ecell = -0.76V, Estandard electrode = 0V
Eelectrode of unknown potential = Ecell + Estandard electrode
= (-0.76V) + 0
= -0.76V
(D)
The e.m.f. of electrochemical cells
As the electrochemical cell consists of two half-cells. Each half-cell has its own electrode
potential. When these two half-cells are connected into an external circuit, two important
things happen.
1. The cell redox reaction begins to occur. This reaction is the sum of the two half-reaction. As
a result, the equilibrium in the two half-cells is disturbed. As the cell redox reaction proceeds,
the concentration of the oxidized species in the half-cell containing the anode increases. On
the other hand, the concentration of the oxidized species in the half-cell containing the
cathode decreases. Eventually the redox reaction reaches equilibrium, when this occurs the
battery has run down or is “flat”.
2. When the two half-cells are connected into an external circuit, current flows from the
negative electrode, that is the anode, to the positive electrode, that is the cathode. The current
is driven by the potential difference. In the external circuit the current is a flow of electrons.
In the two half- cells the current is carried by the ions. As current is taken from the cell, the
concentration of the oxidized species increases and the concentration of reduced species
decreases. At equilibrium, the potential difference between the two electrodes is zero.
The maximum value of potential difference between the two electrodes is known as the
electromotive force or e.m.f. of the cell.
E0cell = E0anode – E0cathode
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
(E)
Chpt:20 p.7
Measuring Cell e.m.f. and Electrode potentials
Accurate measurement of the e.m.f. of a cell by a voltmeter is impossible since this instrument
draws current from the cell. This causes electrode reactions to occur, thereby changing the
concentration of the solution and altering the potential difference.
The problem can be overcome by the use of a high resistance voltmeter, or just by a digital
voltmeter so that the current taken is small.
By knowing the e.m.f. of the cell, the electrode potential of a metal/metal ion system (a half cell)
can be determined by setting up an electrochemical cell with the hydrogen electrode as the other
half cell. If the hydrogen electrode is the negative electrode, the cell diagram is
Pt  H2(g) H+ (aq)  Mn+ (aq) M(s)
Thus
E0cell = E0 (Mn+/M) – E0( H+/H2)
= E0 (Mn+/M)
Note:
The standard electrode potential values measured may have positive or negative values in volts.
A positive value implies that when this electrode is coupled with the standard hydrogen electrode,
the original electrode will be the cathode and reduction occurs.
A negative value implies that when this electrode is coupled with the standard hydrogen electrode,
the original electrode will be the anode and oxidation occurs.
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
Chpt:20 p.8
Exercise 1
From the following data, determine
(a) the standard cell e.m.f.
(b) the cell reaction
(c) the cell diagram.
Data :
(F)
Zn2+(aq) + 2e  Zn(s)
E0(Zn2+/Zn) = -0.76V
Ni2+(aq) + 2e  Ni(s)
E0(Ni2+/Ni) = -0.26V
Standard Electrode Potentials: Extended Redox Potential Series
Electrode potentials of half-cells are often called redox potentials. By convention, redox
potentials are always quoted as reduction potentials還原電勢. The equilibrium half-reactions
are thus written reduction as the forward reaction:
Oxidized species + ze  Reduced species
A reduction potential is a measure of the tendency for reduction to occur.
Besides involving a metal/metal ion for redox system, there are other redox systems involving
reactions between non-meals and non-metals, and also between ions only.
Examples :
Br2(aq) + 2e  2Br-(aq)
Fe3+ (aq) + e  Fe2+ (aq)
E0 = +1.065V
E0 = +0.77V
Note :
<1> An inert platinum electrode is often inserted into these half-cells system to provide a
pathway for electrical conduction.
<2> When writing the half-cell diagram, the reduced form of the ion is put near the inert
electrode, and separated it from the oxidized form by a comma.
 Br2(aq),,2Br-(aq)Pt or Pt 2Br-(aq),Br2(aq)
 Fe3+(aq) ,Fe2+(aq) Pt
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
Chpt:20 p.9
<3> When the oxidized and reduced forms of an electrode system contain more than one
chemical species (ion or molecule) which participate in the cell reaction, These ions or
molecules should be included in the oxidized and reduced forms of the half cell.
Example:
MnO4-(aq) + 8H+(aq) + 5e  Mn2+(aq) + 4H2O(l)
E0 = +1.51V
The cell diagram should be
 [MnO4- (aq)+8H+(aq)],[Mn2+(aq) + 4H2O(l) ] Pt(s)
Extended Reduction Potentials
The electrode potentials of the various systems, when tabulated together with those of the
metal/metal ion systems, will form the extended reduction potential series.
Interpretation
<1> In such series, electrode system with the greatest negative electrode potential are at the top
of the list, and they have the greatest tendency to exist as cations in solution. That is, they
are strongest reducing agent.
<2> Such electrochemical series shows the order of decreasing reactivity of the system, an also
shows which elements can displace each other from solutions of their salts.
(G)
The electrochemical series – Reduction potential series
When elements are placed in order of their standard electrode potentials , the electrochemical
series is obtained.
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
Chpt:20 p.10
<1> The series can be directly obtained from the redox potential series. That part of the series
which includes metals only is called the electrochemical series of metals. This series
closely approximates to the order of reactivity of metals.
<2> The most electropositive and most reactive metals are at the top of the series. These metals
are readily oxidized and thus ionize easily. As a consequence they are strong reducing
agents.. Metals at the bottom of the series do not oxidize readily. They are thus stable in
their reduced form. Gold and Mercury are examples.
<3> DISPLACEMENT REACTIONS置換反應
Metals high in the series reduce the oxidized forms of metals lower in the series.
In effect they displace the metals from their oxides or from solutions of their salts.
Example 1:
Zinc displaces copper from a solution of copper(II) sulphate.
Zn(s) + Cu2+ (aq)  Zn2+ (aq) + Cu(s)
On the other hand, magnesium, which is higher in the series, displace zinc from its oxides:
Mg(s) +ZnO(s)  MgO(s) + Zn(s)
Example 2:
Metals above hydrogen in the electrochemical series of elements reduce hydrogen ions to
form hydrogen gas.
Mg2+(s) + 2H+ (aq)  Mg (aq) + H2(g)
Example 3:
While the reactivity of metals increases as the electrode potentials become more negative,
the reactivity of non- metals increases as the electrode potentials become more positive.
Chlorine displaces iodine from solution.
Cl2(g) + 2I-(aq)  2Cl-(aq) + I2(s)
(H)
Uses of Standard Reduction potentials
Values of standard reduction potentials can be used to predict possible redox reactions預測氧化
還原反應的可行性 and to calculate the e.m.f. of cells.
A redox reaction can occur spontaneouly自發地進行 (feasible to occur可行) when the e.m.f.
calculated is a positive va1ue.
Note
The value of the e.m.f. is only a state function. It only shows that the reaction is energetically
feasible or not by inspecting the sign of the value. It could not tell the rate of the reaction since
the rate depends on the kinetic factors such as the activation energy of the reaction.
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
(I)
Chpt:20 p.11
Feasibility of redox reaction 氧化還原反應的可行性
Thermodynamics predicts that certain reactions are capable of occurring whereas others are
not. Reactions capable of occurring are said to be feasible. However, although a reaction may be
feasible it may not necessarily occur spontaneously. Most combustion reactions do not occur
spontaneously although they are thermodynamically feasible. This is because the energy barrier
has to be overcome before the reaction can start. So, under given conditions a reaction may be
thermodynamically feasible but not kinetically feasible. The thermodynamic feasibility of a
redox reaction can quite simply be determined by inspecting the electrode potentials of the two
half-reactions.
A reduction half-reaction is feasible if its electrode potential is more positive than the
electrode potential of the other half- reaction. Reduction is not feasible if the electrode potential
is less positive than that of the other half-reaction. The half-reaction with the less positive
electrode potential must be oxidation half-reaction.
Exercise 2
Given that :
Al3+(aq) + 3e  Al(s) E° = -1.66V
Cu2+(aq) + 2e  Cu(s) Eo = -0.34V
(a) Will aluminium metals displace copper(II) ions from the aqueous solutions?
(b) Can we tell the rate of the displacement reaction from the values of electrode potentials?
ANSWER .
Exercise 3
(a) Given that
Br2(aq) + 2e  2BrO2 (g) + 2H2O(l) + 4e 
E0= +1.09V
4OH-(aq) E0= +0.40V
Predict whether Br2 or O2 is a stronger oxidizing agent?
How can such prediction be demonstrated experimentally?
(b) Given that
Ag+(aq) + e  Ag(s)
Eo= +0.779V
AgC1(s) + e  Ag(s) + Cl-(aq) Eo= +0.220V
O2(g) + 2H2O(l) + 4e  4OH-(aq)
Eo= +0.401V
Show that silver metal will not undergo oxidation in the presence of oxygen and water
only, but will do so when chloride ions are present.
ANSWER
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
Chpt:20 p.12
Exercise 4
Given that
Ag (aq) + e  Ag(s)
E = +0.80V
AgC1(s) + e  Ag(s) + Cl (aq)
E°= +0.22V
Devise an electrochemical cell in which a spontaneous reaction can take place. Write the cell
diagram for this cell.
ANSWER
IV.
COMMERCIAL CELLS 商用電池
A battery is portable source of electricity. It consists of one or more electrochemical cells.
A cell which does not regenerate reactants is called primary cell 原電池. (e.g. dry cell)
A cell which can be recharged by regenerating the cell reactants is called secondary cell 次電池
or storage cell.
(e.g. lead acid accumulator鉛酸蓄電池)
In each type of cells, redox reaction should occur when it generates electricity.
(A)
The Dry cell
The dry cell is a primary cell which is the commonest, cheapest and most convenient cell used
at present. All its components are solids or pastes which are tightly sealed from the
environment.
The anode is the zinc container which encases the dry cell.
The cathode is a graphite rod surrounded by a layer of manganese(IV) oxide and carbon.
The electrolyte is a paste consisting of zinc chloride, ammonium chloride and water.
The half-reactions at the electrode are
at the anode
at the cathode
The overall equation is
Zn(s) + 2MnO2(s)+ 2NH4+(aq)  Zn2+(aq)+ Mn2O3(s) + 2NH3(g)+ H2O(l)
Note : <1>
<2>
A dry cell generate between 1.25V and 1.5V.
The formation of ammonia around the cathode would disrupt the electric current. This is
prevented by the complex formation with Zn2+
4NH3 + Zn2+  Zn(NH3)42+
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
(B)
Chpt:20 p.13
Secondary cell- Lead storage cell (Lead-acid accumulator)
Lead accumulator is a secondary cell, the half-reactions at the electrodes are readily reversible.
It consists of a lead anode arid a grid of lead packed with lead(IV) oxide as the cathode. The
electrolyte is sulphuric acid.
The half-reactions at the electrodes are
at the anode:
at the cathode
The overall reaction is
Pb(s)+ PbO2(s)+ 4H+(aq)+ 2SO42-(aq)
discharging

2PbSO4(s)+ 2H20(l)
charging
Note
<1> One cell provides about 2V. The storage battery used in cars normally consists of 6 these
cells arranged in series to provide 12V.
<2> The battery is recharged by applying a current from an external source. This reverses the
electrode reactions. The state of charge can be indicated by the relative density of
sulphuric acid
(i)
When the storage cell is fully charged, the sulphuric acid has a relative density of
1.275.
(ii)
When the cell is used for some time, the concentration of sulphuric acid will
decrease and also its relative density. Upon discharging, lead (II) sulphate forms at
both electrodes, thus reducing the concentration and relative density of the
sulphuric acid. The relative density thus indicates the state of charge in the battery.
Lead acid accumulator
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
(C)
Chpt:20 p.14
Fuel cell 燃料電池– Hydrogen-oxygen fuel cell
A fuel cell is a type of primary cell in which the reactants are continuously replaced as they are
consumed and the products are continuously removed.
In the hydrogen-oxygen fuel cell, hydrogen and oxygen are bubbled through porous carbon
electrodes into a concentrated solution of an alkali. The carbon electrodes contain a platinum
catalyst. The half-reactions at the electrodes are
at the anode
2H2(g) + 4OH-(aq)  4H2O(l) + 4e
at the cathode
O2(g) + 2H2O(l) + 4e 
4OH- (aq)
The overall reaction is
2H2(g) + O2(g)  2H2O(l)
The water is removed and in spacecraft is consumed by the astronauts as the drinking water.
Note:
In some fuel cells an acidic electrolyte is used in which case the electrode reactions are
at the anode :
H2(g)  2H+(aq) +2e
at the cathode :
O2(g) + 4H+(aq) + 4e  2H2O(l)
The overall reaction is
2H2(g) + O2(g)  2H2O(l)
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
V.
Chpt:20 p.15
CORROSION OF IRON AND ITS PREVENTION鐵的腐蝕與防預
(A) Introduction
Corrosion may be regarded as the natural tendency of metals to return to their oxidized species.
Corrosion is a redox process.
reduced species
(metal)

oxidized species + ne
The main sources or causes of corrosion are
<1> the atmosphere - humidity and atmospheric pollution, (most important)
<2> submersing 浸入 in water of aqueous solution,
<3> underground soil attack,
<4> corrosive gases,
<5> immersion in chemicals.
Corrosion of iron is called rusting. Rusting is the electrochemical process by which iron rust away
when exposed to air and water.
(B) The electrochemical process involved in rusting
Rusting is the most common example of corrosion. When iron comes into contact with acid, oxygen an
other species found in the environment, electrochemical reactions takes place.
<1> Consider an iron sheet exposed to open atmosphere. The plate may have other metal impurities
incorporated in it When the plate gets wet due to moisture in the air, the thin layer of water tends
to dissolve oxygen, which diffuses from the edge of the water drop to the interior.
<2>
Around the edge of the drop of water 水滴的邊沿, the concentration of atmospheric oxygen is
higher, and the following reaction may occur on the iron surface immediately in contact with
the edge of the water drop.
1/2 O2(g) + H2O(l) + 2e  2OH-(aq) cathodic reaction
Note
1.
If another metal impurities (which accepts electrons more readily than iron) happens to
be present at this site, then the above reaction may occur even more readily.
2.
The reaction is favored by a low pH, so that the presence of acid , or even carbon
dioxide, will accelerate rusing.
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
Chpt:20 p.16
<3> At the centre of the water drop, the concentration of atmospheric oxygen is lower, so that the
following reaction takes place preferentially.
Fe(s)  Fe2+(aq) + 2e
anodic reaction
Note:
1. Since this process involved release of electrons, it takes more readily on non-uniform surface (e.g.
pointed ends) and mechanically strained areas, where the electrons are of higher energy and can be
released more readily.
2. The presence of a more electropositive metal as impurities also favors this process, as it releases
electrons even more readily than iron.
<4>
The free Fe2+(aq) and OH-(aq) diffuse from their sites of formation and precipitates as Fe(OH)2
precipitate, which is further oxidized by dissolve oxygen to form iron(III) hydroxide. On
standing, this changes to rust (Fe2O3.nH2O), a reddish brown solid. In many cases this coating
either flakes off or is permeable to both water and air, so that corrosion continues unhindered.
NOTE:
(1) The whole process is accelerated by electrolytes like sodium chloride, which gently increases the
conductivity of the solution.
e.g. Iron window frames by the seashores rust more readily as the thin layer of water on iron surface
contains dissolved sodium chloride from salts spray near the sea.
(2) The whole process is also accelerated by high temperature, which increases the rate of a chemical
reaction.
e.g. Car exhaust pipes rust easily as they are subjected to great heat. They have to be replaced every
couple of years.
SUMMARY
Corrosion of iron (rusting) is a local electrochemical process, with the metal at the centre of the water
drop being the anodic site, whereas the metal (or other metal impurties in contact with it) at the
periphery of the drop being the cathodic site.
The corrosion process of iron can be proved by considering their electrode potentials.
Fe2+(aq)  Fe(s)
E0 = -0.44V
anode
0
Pt  O2(g)  OH (aq) E = - 0.4V
cathode
E0cell = (-0.40V) - (-0.44V)
= 0.04V
As the value of E0 is positive, the above cell is spontaneous.
The whole process can be represented by a cell diagram
Fe(s)  Fe2+(aq) OH-(aq) O2(g)  metal impurities(s) or Fe(s)
Such a local voltaic cell is formed by iron together with other metal impurities, as well as water and air.
Dissolved electrolytes, e.g. NaCl and HCO3- from atmosphere, a low pH and a high temperature, will
accelerate the process.
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
Chpt:20 p.17
(C) Prevention of Corrosion
Various methods are used to protect a metal from corrosion.
<1> Coating.
Coating the metal with paint or tin can prevent attack by atmospheric moisture and oxygen.
However, if coating is flawed or broken, corrosion occurs.
e.g. Plating with a thin layer of tin.
When a tin-coated iron can containing lemon juice is scratched. The two metals are exposed to the
acid, an electrochemical reaction would occur.
Fe2+/Fe
= -0.44V
2+
Sn /Sn
= -0.136V
Iron is more reactive than tin so it tends to reduce any tin ions that are present. Thus, the tin ions
aids the oxidation of iron. As a result, the scratched tin-plated object has a stronger tendency to
corrode.
<2> Cathodic protection
The use of zinc to protect iron form corrosion is called cathodic protection.
When iron is coated with zinc, it is called galvanized iron. A galvanized iron can containing acidic
rain water is scratched. The iron surface is exposed. Some metal ions form in solution.
Since
E0 : Fe2+/Fe
-0.44V
Zn2+/Zn
-0.76V
Zinc is more electropositive than iron (more likely to form cation) than iron so that it tends to
reduce any iron ions that form.
Zn(s) + Fe2+(aq)  Zn2+(aq) + Fe(s)
So the presence of zinc prevents the corrosion of iron. Iron will be protected from rusting even the
galvanized iron 鍍鋅鐵 is scratched 刮花.
Note:
<1> Since Zn is higher in the electrochemical series than iron, it acts as the anode and the iron acts as
the cathode. The zinc is thus oxidized in preference to the iron.
Zn(s) 
Zn2+(aq) + 2e
<2> Zinc is not used in food canning since it is higher than tin in the electrochemical series. It is thus
more susceptible to attack by acids in fruit juice. Besides, Zn2+ ion is toxic.
<3> A sacrificial anode is also used as a means of cathodic protection.
A metal (e.g. magnesium ) much higher than iron in the electrochemical series is chosen as the
sacrificial metal. Such metal has a higher tendency to release electron than iron, it always form anode
whenever an electrochemical cell is set up between it and iron.
L.S.T. Leung Chik Wai Memorial School
F.6 Chemistry
Chapter 20: Redox Equilibria II
Chpt:20 p.18
Mg(s)  Mg2+(aq) + 2e (oxidation)
The sacrificial metal forms the anode whereas
the iron surface forms the cathode to accept the
electrons released from anode.
In practice, the sacrificial metal is connected to
the object (e.g. steel pipe) through a conducting wire.
Iron is protected from corrosion as the metal remains.
<4> Aluminium, although high in the electrochemical
series, does not corrode readily since it oxidizes to form
aluminium oxide, A12O3, which does not flake.
It thus forms a strong protective coating.
(C) Socioeconomic implication of corrosion and prevention
As a construction material, iron is suffered from the drawback that it corrodes quite readily. Corrosion
of iron can cost much to the society.
In terms of money, it is estimated in 1980 in the US that 70 billion dollars were lost annually because
of corrosion. Apart from money spent on the replacement of the corroded articles and the prevention
methods, there are indirect costs such as those for the maintenance of machines, and those due to lost
production when machines fail down or when they are shutdown for maintenance.
The wastage of natural resources is one of the social implications of corrosion. It has been estimated
that 1 tonne of steel is converted into rust every 90 seconds in Britain, and that about 40% of the steel
made in the US is used to replace steel lost by rusting. Other than wasting natural resources, corrosion
of iron causes considerable inconvenience to human beings, and even lost of life. This happens because
corrosion results in the formation of cracks and crevices which weakens the strength of metals.
Concrete in buildings may fall off if the steel reinforcing bars inside the concrete corrode. Moreover,
there is evidence that there may be a correlation between the number of serious injuries suffered in road
accidents and the age (i.e. amount of corrosion) of the vehicle.
Despite of the disadvantages, iron is still the most important in present-day society. The disadvantages
of corrosion are outweighed by the relative low cost, abundance and ease of extraction of iron.
Replacement of iron by a more corrosion-resistant metal of a reasonable price will be the goal of many
metallurgists. But, more than they should be done, the society as a whole should assume a greater
responsibility for corrosion prevention in order to have a greater improvement of the situation.