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R01 - R02 non-redox reactions - redox reactions
Aim: To show that there are electrical aspects to
certain chemical reactions and hence show the
existence of electron transfer reactions
Chemical reactions can be classified into various types:
synthesis, decomposition, dissociation, replacement,
precipitation, acid-base and redox to name but a few.
Some of these types of reaction involve changes in the
electron structure of the (charged) atom others do not.
e.g. (i) the dissociation in water of an ionic compound
into its component ions; and (ii) the formation of a
crystalline precipitate from ions in solution. However,
they are all chemical reactions.
Non-redox reactions
Illustration R1 depicts simple experiments illustrating two
types of reaction in which there are no changes in the
electron structure of the (charged) atoms.
The left-hand side of R1 depicts the addition of an
aqueous solution of sodium hydroxide (NaOH), which is
colourless, to a solution of a weak acid, the indicator
phenol red, which is yellow coloured. Phenol red is a
weak acid which dissociates in water:
If an acid is added the reaction is displaced to the left
and the yellow-coloured HA is formed. If an alkali, e.g.
sodium hydroxide, is added, the reaction is displaced to
-
the right and the red-coloured A ions are formed.
Since there is neither proton transfer nor formation of a
precipitate this must belong to a different class of
reaction. The type of reaction involved can be
demonstrated by allowing the same reagents to react
with each other in a different situation as shown on the
right-hand side of illustration R2.
Here the two solutions are not mixed as in the
experiment depicted on the left-hand side, but are
brought into contact with each other by means of a
so-called salt bridge. This allows ionic transport by
cations, positively charged ions, as well as anions,
negatively charged ions, without allowing the two
solutions to mix. An example of a salt bridge is a
U-shaped tube filled with agar gel and containing as
electrolyte potassium chloride (KCl).
Inert platinum rods are placed in both solutions and
these rods are in turn linked by wires to a small low
voltage lamp.
An alternative experiment is to replace the lamp by an
ammeter or voltmeter which enable even a small current
to be detected. The current can be increased by
replacing the platinum rods by platinum foil which has a
larger surface area.
The glowing filament in the lamp shows that a current is
flowing through the circuit i.e. that electrons are being
transferred between the chemical reagents.
This type of reaction is classified as an electron transfer
reaction. It can only be demonstrated in a set-up in
which the reagents are not mixed. If electrodes are
placed in the bottom of the beaker shown on the
left-hand side of R2, no current flows between the
electrodes. This does not mean that electron transfer
does not occur when potassium iodide (KI) is directly
reacted with iron(III) chloride (FeCl ). It only means that
3
the electrons are not transferred to the electrodes, but
-
directly from I ions to iron(III) ions (Fe
represented by:
3+
). This can be
This type of acid-base reaction is classified as a proton
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transfer reaction, a proton being abstracted from HA by
-
the OH ions.
The right-hand side of illustration R1 depicts the reaction
between an aqueous solution of nickel chloride (NiCl ),
2
which is green, and an aqueous solution of sodium
sulphide (Na S), which is colourless.
2
When these aqueous solutions are mixed a colour change
is observed which cannot be explained by physical
dilution.
The original green colour of the nickel chloride is
replaced by a less coloured liquid with a black precipitate
distributed in it. This black precipitate is nickel sulphide
(NiS).
When the two salts nickel chloride and sodium sulphide
are dissolved in water, they ionise:
Which itself can be represented by two half-reaction
equations :
-
All the electrons which are donated by the I ions are
3+
accepted by the Fe ions. In the set-up on the left, the
direct electron transfer reaction, no electrons end up in
the water. They are directly exchanged from particle (ion
or molecule) to particle (molecule or ion).
In the case of a set-up such as that on the right of
illustration R2, the electrons are transferred - more
-
2+
2-
The nickel ions (Ni ) react with the sulphide ions (S )
to form crystals of nickel sulphide( NiS ), which
precipitate.
This type of reaction is a precipitation reaction.
Not all reactions in which a coloured solution and a
colourless solution are mixed can be classified into
proton transfer or precipitation reactions.
Redox reactions
The left-hand side of illustration R2 depicts a simple
experiment in which the electron structure of the
(charged) atoms do change. A colourless aqueous
solution of potassium iodide(KI) is mixed with a yellow
aqueous solution of iron(III) chloride, (FeCl ). The
3+
slowly than in the first case - from I to Fe via the
conducting wires, through the bulb wire.
The transport of charge throughout the salt bridge and in
the solutions themselves is carried out by different
charge carriers: ions (cations as well as anions).
More details are given on illustration R3.
Moreover, the salt bridge ensures that the two solutions
do not diffuse.
This short search reveals that such special and extremely
interesting chemical reactions do exist.
However, should you think that all chemical reactions can
be traced to this one type, control tests on other
reactions, such as the ones indicated on illustration R1,
will be absolutely necessary. In set-ups analogous to the
one on illustration R2 (right) no current flow can be
registered for proton transfer reactions (left on R1) or for
precipitation reactions (right on R1).
3
resulting aqueous solution is red-brown in colour.
The two salts ionise in water
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R3 The classic electrochemical cell
Electrodes are represented according to IUPAC as solid
vertical lines I, the salt bridge by double vertical
dashedlines.
In practice sulphate salts are used in this cell although
nitrate, chloride or hydrogen sulphate salts could be
used.
The salt bridge, as also shown in illustration R2, contains
KCl or KNO . It is a sufficiently good conductor of ions to
3
produce good contact between the solutions.
It does not contain possibly interfering species.
Aim: To demonstrate the classic electrochemical
cell and compare it with an unusual but
analogous example
In illustration R2, it was shown that the electron transfer
-
3+
between I and iron(III) ions (Fe ) can be used to
produce an electric current by placing platinum
electrodes in the two solutions, connecting the electrodes
with a wire and connecting the solutions with a salt
bridge. The two half-reactions, which make up the
electron transfer reaction, take place in the solutions in
the two beakers.
This is an electrochemical cell, and the beakers in which
the two half-reactions take place are its half-cells. An
electrochemical cell is therefore a device which allows an
electron transfer reaction to take place in such a way
that a flow of electrons is produced between the
electrodes.
If a piece of zinc metal is placed in a blue aqueous
solution of copper(II) sulphate (CuSO ) then the colour
4
intensity spontaneously decreases and the zinc takes on
a red-brown colour. An electron transfer takes place
between the zinc metal and the copper(II) ions with
copper being deposited on the zinc metal, the solution
becoming colourless as the copper(II) ions are reduced
on the zinc metal to copper and the zinc metal is
oxidised to colourless zinc ions. This reaction can be
represented as:
This electron transfer can, in principle, be observed
indirectly by an increase in temperature as the electron
transfer energy is changed to heat energy.
When the electrodes are connected, current begins to
flow and the reaction starts. As the electrons flow
between the zinc and copper electrodes the
concentration of zinc ions increases in one half-cell while
the concentration of copper(II) ions decreases in the
other half-cell leading to an ionic imbalance.
This results in a decrease in the potential difference
between the two half-cells. The salt bridge between the
two solutions enables an ionic balance to be maintained
without the transport of ions from one half-cell solution
to the other. The positive ions (cations) of the salt bridge
diffuse to the sulphate solution and the negative ions
(anions) to the solution of zinc ions.
The reaction proceeds and the potential difference
between the electrodes decreases eventually reaching
zero when either all the copper(II) ions have or the zinc
electrode has been totally consumed.
On the right-hand side of illustration R3 a lemon is
shown with two electrodes stuck in it. One is zinc and
one is copper. If these electrodes are joined by a wire to
a voltmeter then a voltage (admittedly lower than 1.1 V)
can be measured. Again the zinc is functioning as the
anode and the copper as the cathode.
The lemon must contain ions which function in the same
2+
2+
way as the Cu and Zn ions in the Daniell cell. The
diffusion of the ions is slow and so the voltage is lower
but is measurable, even though both electrodes are in
one solution without a salt bridge.
Cell electrodes
A metal consists of cations of the metal (positively
charged ions) surrounded by electrons. If a strip of
metal is placed in a solution of its ions some of the
cations in the metal may dissolve leaving a build up of
electrons on the metal:
This redox reaction can be broken down into the
following half-reaction equations :
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The metal strip will become negatively charged.
Alternatively metal ions in the solution may take
electrons from the strip of metal and be discharged as
metal atoms:
If these half-reactions are set-up in separate beakers as
in illustration R2 (right-hand side), each containing the
metals in the form of electrodes immersed in an aqueous
solution of the corresponding metal salt and with the
solution connected by a salt bridge and the metal
electrodes linked by wires to a lamp as shown in
illustration R3, we again create a source of electric
current.
This demonstrates the involvement of electron transfer.
If this simple electrochemical cell is kept at a
temperature of 25°C (298K) and the concentrations of
the zinc and copper ions are both 1mol/L, then we obtain
what is known as the Daniell cell.
The tendency of electrons to flow through the external
circuit of a cell is quantified as the e.m.f. , the
electromotive force of a cell, also called the cell potential
E .
cell
The e.m.f of a cell can be measured using a high
resistance voltmeter, which takes negligible current from
the cell.
In the case of the Daniell cell the E
is +1.1 V.
cell
The cell potential is by convention taken to be acting
from left to right through the cell with the (more)
negative electrode on the left.
Thus
The potential difference between the strip of metal and
the solution depends on the nature of the metal and the
concentration of the ions involved in the equilibrium at
the metal’s surface. If we compare zinc (Zn) and copper
(Cu) for example, then, for the same concentration of
ions in each solution, zinc acquires a more negative
potential than copper since it has a greater tendency to
dissolve ions and a smaller tendency to be deposited as
a metal. Hence in such a system as is shown in
illustration R3 the zinc electrode is the anode and the
copper electrode the cathode.
At the anode oxidation takes place : here zinc metal is
converted into zinc ions.
At the cathode reduction takes place : copper(II) ions
are reduced to copper metal.
In galvanic cells, the anode is indicated by the negative
sign s and the cathode by the sign r.
In electrolytic cells (see R11) an anode, being the
electrode where oxidation takes place, gets the positive
sign r. A cathode of an electrolytic cell is indicated by
the negative sign s.
For a galvanic cell, the cell voltage or cell potential, in
terms of anode and cathode is written as :
The cell can be represented as :
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R04 Table of electronegativity
Below is a schematic diagram of the deduction of these
oxidation numbers:
Aim: To highlight the electronegativity of the
elements in the description of redox reactions in
general and of the oxidation number in particular
The above-mentioned cases concern metals and
monoatomic metal ions among which an electron transfer
took place. In these cases, a complete exchange of
electrons is evident.
A great deal of electron transfer reactions, however,
concern uncharged covalent molecules and/or multiatom
ions. In the reaction which takes place in a gas cooker,
methane reacts with dioxygen, both covalent molecules.
The reaction products CO and H O also belong to the
2
2
group of exclusively covalent substances. Such
combustion reactions can essentially be reduced to a
transfer of electrons between atoms as well.
This is easy to prove by letting both elements slowly
react in a galvanic cell, in this case a fuel cell. This cell
actually produces electricity (electrical energy) from
chemical conversions (chemical energy).
After the introduction of the concept "oxidation number
of an atom" it is understandable that one began to use
"oxidation-reduction reactions" or "redox reactions" as
general names for electron transfer reactions.
When in a particle the O.N of an atom increases due to
the release of (valency) electrons, it is said that this
atom is being oxidized and that the particle as a whole
acts as a reductor. The released electrons will be
absorbed by an(other) atom into a(nother) particle. It is
then said that the atom is being reduced and that the
particle as a whole acts as an oxidant.
Important:
All electrons which are released by a reductor in a redox
reaction have to be absorbed by the oxidant. Therefore,
the stoichiometric figures of the oxidant and the reductor
in a redox reaction have to be filled in.
The symbols Red and Ox stand for "reduced form" and
"oxidised form" respectively.
Oxidation number
In order to be able to describe such reactions as electron
transfer reactions as well, we have to follow accounting
rules. Only then can chemists agree on the gain-loss
balances of electrons during a chemical conversion.
The cornerstone of this electron accounting is the
socalled oxidation number, which is assigned to a specific
atom in a specific particle (ion or molecule):
the oxidation number (O.N.) of an individual atom in a
specific particle is the electrical charge in electron charge
units actually carried by this atom, as an ion in an ionic
bond, or assigned to this atom by representing all
covalent bonds as ionic bonds.
1 of 2
In this reaction equation of a redox reaction, A
Red
Red and A
1
Ox
or
or Ox - by analogy with acid base
1
reactions - can be called the conjugated oxidantreductor pairs.
This is also valid for B or Ox and B
or Red
Ox
2
Red
2
Further on, on illustration R5, we will see that this
fundamental balance reaction, whether self-sustaining or
not, mainly progresses from the direction of a relatively
strong reductor and a relatively strong oxidant towards a
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Oxidation numbers are assigned to atoms of a molecular
structural formula showing all valence electrons, on the
basis of the following rules: all binding electrons in a
covalent bond are entirely displaced to the most
electronegative partner (see R4). In the case of equal
electronegative partners, the binding electrons are
equally distributed. The amount of electrical charge (not
in coulomb, but in electron charge units) assigned to the
atom is that obtained by the atom's group number less
the number of valency electrons remaining on this atom.
Since "the oxidation number of an atom" does not
necessarily correspond with a real ionic charge, such an
O.N. is indicated by a Roman figure (instead of an Arabic
figure) preceded by + or - (instead of followed by the
charge symbol). Referring to the O.N. of an atom (in an
ion or in a neutral molecule) one will therefore say, for
instance, +II (plus two) and not 2+ (two plus).
This means that the difference between O.N. and ionic
charge is not only legible, but also audible!
Examples :
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weaker oxidant and a weaker reductor.
Certain techniques which are applied in class to deduce
the stoichiometric values of a redox reaction are often
based on the application of the oxidation number
accounting. Nevertheless, the usual atom balance
method is applicable to most redox reactions as well.
Electronegative values
The ability of an atom to attract the electrons of a
covalent bond is measured by the electronegativity of
the atom. An atom of high electronegativity will attract
electrons away from one of lower electronegativity.
There are several electronegativity scales. In the Pauling
electronegativity scale fluorine is given the value of 4.0
and the other elements are allotted values based on
their electronegativities relative to fluorine.
Put in another way, the electronegativity of an element
expresses on an arbitrary scale the relative attraction of
one of its atoms for the electrons of a single covalent
bond formed with an atom of another element.
Pauling electronegativities are given for selected
elements in illustration R4.
Dealing with redox reactions we saw that the deduction
of oxidation numbers is very important.
To be able to unmistakably deduce the oxidation number
of a covalently bonded atom, the displacement of binding
electrons (rule of play) has to be applied to the
molecule's or the ion's Lewis representation.
For I the Lewis notation becomes :
2
after assignment of the binding electrons and the O.N.(I)
will be 0, since the group number of I is seven and the
number of valency electrons in an ionic representation of
the covalent bond is seven as well:
group number 7 - 7 valency electrons = 0.
2 of 2
Therefore, a good ready knowledge of the
electronegative value of certain elements is just as
important as the knowledge of their electron
configuration (number of valency electrons). Besides,
memorizing these values is very simple: For Li, 1.0
applies as Electronegative Value and up to F an
increment equal to 0.5 applies; from K to CI an
increment equal to 0.3 applies, except for the last two
elements S and CI.
In this scale the Electronegative Value of H is 2.1.
Hydrogen is therefore not a metal, but a non-metal!
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R05 Ordering of half-reaction equations for metals
Aim: To arrange several metal redox couples
according to their standard reduction potential
A half-reaction equation represents a substance in
equilibrium with an oxidised or reduced form of the same
substance. According to IUPAC recommendations a
halfreaction equation is always written in the form of a
reduction with a substance on the left of the equation
and a reduced form of the substance on the right of the
equation. For example:
At the same time zinc is oxidised to zinc(II) ions, the
oxidation number of zinc changing from 0 to +II.
Adding these two half-reaction equations (top: right side;
bottom : left side) so that no electrons are left, gives the
overall redox reaction equation:
2+
We can therefore conclude that the Zn
/Zn half-reaction
2+
By reduced form of a substance is meant a species in
which an element is characterised by a lower oxidation
state than in the oxidised containing the same element.
For a strip of copper in equilibrium with a solution of
copper(II) ions the half-reaction equation is:
equation is more strongly reducing than the Pb /Pb
half-reaction equation.
In the second experiment a lead rod reduces copper(II)
ions to copper metal, the oxidation number of copper
changing from +II to 0. The blue colour of the aqueous
solution is seen to disappear. At the same time the lead
is itself oxidised to lead(II) ions, the oxidation number of
lead changing from 0 to +II. The lead rod becomes
brown in colour as copper is deposited on it.
Examples of half-reaction equations for a dissolved gas in
equilibrium with its ions are:
Adding the two half-reaction equations gives the overall
redox reaction equation:
These so-called redox couples have an ability
spontaneously to donate or accept electrons from
1 of 2
This shows that the lead half-reaction equation is more
strongly reducing than the copper half-reaction equation.
In the third experiment the zinc rod reduces the
copper(II) ions to copper metal, the oxidation number of
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another substance depending on their potentiality for so
doing.
The potentiality of a particular redox couple to donate or
accept electrons from another redox couple can be
qualitatively established by observing the effects of
bringing different redox couples together in a redox
reaction.
2+
2+
copper changing from + II to 0, resulting in the original
blue solution becoming colourless and a red-brown
deposit of copper being formed on the zinc rod. At the
same time the zinc is itself oxidised to zinc(II) ions, the
oxidation number of zinc changing from 0 to +II.
2+
Let us consider the Zn /Zn, Pb /Pb and Cu /Cu redox
couples. Illustration R5 shows three simple redox
reactions. In the first a zinc rod in an aqueous solution of
Pb(NO ) (lead(II) nitrate) reduces the lead(II) ions to
3 2
Adding the two half-reaction equations gives the overall
redox reaction equation.
lead, the oxidation number of lead changing from +II to
0. This is observed as a lead “tree” of dendritic lead
crystals.
2+
We can therefore conclude that the Zn
2+
is more reducing than the Pb
/Pb half-reaction, which in
turn is more reducing than the Cu
equation.
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/Zn half-reaction
2+
/Cu halfreaction
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R06 Table of half-reaction equation equations for non-metals
In this reaction Br
2(aq)
has oxidised
itself being reduced to
Aim: To arrange some halogen redox reactions
according to their oxidation capacity
oxidizing agent and
The Br
At the top right-hand side of the Periodic Table we find
the halogen group; fluorine, chlorine, bromine and
iodine.
They are normally found as diatomic molecules, X .
. Br
2(aq)
to I
, while
2(aq)
is therefore the
the reducing agent.
-
/2Br half- reaction equation is the stronger
2(aq)
oxidiser and the two half-reaction equations can be
written as :
2
The halogens (Greek for salt makers) are seldom found
free in nature. This does not mean that the bonding
between the atoms to form diatomic molecules is weak,
but is rather a reflection of the fact that the halogens
easily form ionic halides.
In the reaction depicted on the right-hand side of
illustration R6 chlorine (Cl ), a light green gas,is bubbled
2
This is a redox reaction with a transfer of electrons
taking place. Halogens are well known oxidising agents,
which indicates that the halogen/halide ion redox couple
has a potentiality for accepting electrons from another
reagents i.e. oxidising that agent.
-
The potentiality of a particular X /2X redox couple
2
through a colourless aqueous solution of sodium bromide
(NaBr). A light brown aqueous solution of bromine
results. Again the colour can be intensified by shaking
with a small volume of tetrachloromethane. Since the
bromine dissolves in CCl better than in water, the CCl
4
4
layer at the bottom becomes brown. Bromine, like iodine,
is apolar and therefore dissolves better in the apolar CCl
4
than in water which is highly polar.
The reaction can be represented by two half-reaction
equation equations:
-
relative to other X /2X redox couples can be
2
qualitatively established by observing the effects of
bringing different redox couples together. In illustration
R6 two simple redox experiments are shown.
On the left-hand side of illustration R6 a brown liquid Br
2
is mixed with an aqueous solution of iodide ions for
example from potassium iodide (KI). A light brown
solution is obtained showing the presence of a small
quantity of iodine. To demonstrate the presence of iodine
more convincingly,a small volume of tetrachloromethane
(CCl ) is added and the vessel shaken. After leaving to
4
stand for a few minutes the bottom layer of CCl has a
4
1 of 2
It is not possible by means of such a simple reaction to
show that fluorine reacts as an oxidising agent with
respect to chlorine, because it reacts violently with water
in a redox reaction during which a gas, HF, is produced.
I , Br , and Cl , also react with water but less violently.
2
2
2
We may assume that a redox reaction between F
2(aq)
oxidising agent and
as
as the reducing agent can be
represented by the half-reaction equations:
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purple colour of I , while the aqueous layer has lost most
2
of its brown colour. Another test would be to add a few
drops of starch solution, whereupon the presence of a
blueblack solution would indicate the presence of iodine.
The total reaction which is occurring can be represented
as:
From these experiments it can be concluded that iodine
is the least oxidising of the four halogen elements and
-
that the oxidising power i.e. the potentiality of the X /2X
2
couple to accept electrons increases in the order
iodine < bromine < chlorine < fluorine
Although these simple experiments are not all carried out
under the standard conditions of temperature and
pressure necessary for the measurements of standard
reduction potentials, they achieve their purpose of
demonstrating convincingly the order of their potentiality
to accept electrons in the non-metallic halogen group.
N.B.: When working with any halogen element all
necessary safety precautions should be taken.
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R07 Table showing several standard reduction potentials
because we substractthe reduction potential of the left
electrode from the reduction potential of the right
electrode.
On the basis of the standard half-cell redox potentials of
2+
Zn
/Zn and Cu
2+
/Cu one would expect Zn metal to play
2+
2+
the role of the anode (Zn oxidised to Zn ) and Cu the
role of the cathode (oxidising agent, being reduced).
Then :
This positive value for E
cell
ensures a self-sustaining
2+
redox reaction between Zn and Cu , even when energy
is not (continously) supplied to the system.
Aim: To show how, using a limited table of
standard reduction potentials for redox
half-reactions many experimental results can be
explained and many more predicted
Using illustration R5 it has been shown that the reducing
power of redox couples i.e. the potentiality to donate
electrons increases in the order:
and using illustration R6 it has been shown that the
oxidising power and halogen redox couples i.e. the
potentiality to accept electrons increases in the order:
Important reducing and oxidising agents.
The table given below shows many more redox
halfreaction equations than have been considered in the
preceding series of experiments. They have been written
in accordance with the IUPAC convention, all values are
standard reduction potentials and the letters (g) (l) (s)
and (aq) have been added to show the physical state:
gas, liquid, solid and aqueous respectively. These letters
may be written on the same level as the symbol and
immediately next to it; e.g. Li+(aq) or
this qualitative information on the relative reducing and
oxidising power of redox couples is quantified by
measuring their standard reduction potentials. Standard
reduction potentials of half-reactions are referred to the
standard hydrogen electrode.
Standard hydrogen electrode
The standard hydrogen electrode has, by convention, a
standard reduction potential of zero volts at unit effective
concentration, a temperature of 298 K and a hydrogen
pressure of one atmosphere. The hydrogen electrode
consists of a platinised platinum electrode immersed in a
1 mol/L solution hydrogen ions. Hydrogen gas at a
pressure of 1 atm is bubbled over the platinum
electrode. On the surface of the platinum an equilibrium
is established between hydrogen gas and hydrogen ions.
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A potential develops on the surface of the platinum. It is
assigned a value of zero volts.
Compounds with both oxidising and reducing
properties
Determination of standard reduction potentials
There are situations in which the ions indicated by a
halfreaction equation can function as both an oxidising
and reducing agent depending on the other ions in the
reaction and the type of medium. Hydrogen peroxide
can, for example, function as both an oxidising and as a
reducing agent. In the table (left) the substance H O is
Standard reduction potentials are determined by
connecting the system represented in the half-reaction
equation concerned to a standard hydrogen electrode via
a salt bridge and reading off the potential difference on
either a previously calibrated potentiometer or a high
resistance voltmeter. The standard hydrogen electrode is
by convention always placed on the left.
The configuration of the cells whose standard reduction
potential is being measured depends upon the species
involved. In the case of a metal ion/metal couple the
metal is present as an electrode immersed in a 1mol/L
solution of its ions at 25° C. In the case of an ion/ion
couple a platinum electrode is immersed in a solution
containing 1 mol/L of each of the ions at 25° C. Finally,
in the case of a gas/ion couple the gas at a pressure of 1
atm is bubbled over a platinum electrode immersed in a
solution of 1 mol/L of ions at 25° C.
2
Simultaneous oxidation and reduction of a compound is
known as disproportionation. The fact that it can behave
in this way does not affect neutral aqueous hydrogen
peroxide when being stored . Aqueous solutions only
containing H O remain stable. Decomposition only
2
When the values for
are tabulated it is the convention
to write the most strongly reducing redox couple at the
top of the table and the most strongly oxidising couple at
the bottom.
he table of standard reduction potentials in illustration R7
shows the
values for redox half-reactions which have
been measured with respect to the standard hydrogen
electrode as reference.
It will be noticed that the
value of the redox
half-reaction equations at the top of the table are large
and negative, while those for the half-reaction equations
at the bottom of the table are large and positive. Strong
reducing agents and hence the least oxidising redox
+
2+
couples, such as Na /Na and Zn /Zn have large
negative
values, whereas the least reducing and
-
hence most oxidising redox couples such as Cl /2Cl and
2
occurs when H O comes into contact with a catalyst
2
such as Fe
Table of standard reduction potentials
2
found top right and bottom left :
2+
, Fe
2
3+
, I-, MnO , blood, etc. It is commonly
2
used for cleaning contact lenses, bleaching hair and as a
disinfectant.
Ammonium nitrate (NH NO ) is another compound that
4
3
can disproportionate, but this is not linked to the type of
medium being used and a real catalyst does not need to
be present. This can be explained by using the above
table from which it can be seen that the
ion is at the
top right-hand side of the table and is therefore a strong
reducing agent. The nitrate ion (
) is at the bottom
left-hand side of the table and is therefore a possible
strong oxidising agent. In an aqueous solution or in the
solid, the
and
ions can remain stable.
A jolt however can be sufficient to start a
disproportionation or autoxidoreduction reactions.
Ammonium nitrate was used for many years as a
fertilizer but it is also a powerful explosive, decomposing
on heating:
2
-
F /2F have large positive
2
values and are to be found
at the bottom of the table.
Ammonium dichromate (NH ) Cr O (s) also reacts in a
4 2
Use of standard reduction potentials
This table can be used to predict the outcome of
classroom experiments and also to explain various
industrial processes which are based on redox
half-reaction equations by using their
values.
Bromine is produced industrially by bubbling chlorine gas
through a solution of bromide ions. The same principle
cannot be used to produce chlorine industrially, because
the installation necessary for the production of fluorine is
too expensive. Other oxidising agents such as the
,
2 of 3
2
7
similar way on heating decomposing spectacularly from
orange crystals to a large heap of green flakes.
Thermodynamic capability versus kinetic reality
Above we were clearly confronted with examples of
reductors and oxidants which, even strongly mixed, do
not proceed to an electron transfer or a redox reaction.
Consider H O , that does not disintegrate spontaneously,
2
2
or NH NO , that does not explode spontaneously.
4
3
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which can be used in the laboratory scale production of
chlorine, are also too expensive for large scale use.
Chlorine is produced industrially by the electrolysis of an
aqueous solution of sodium chloride. The process is not
self-sustaining and electrical energy must be constantly
supplied. Electrolysis is the only method of producing
fluorine because there are no oxidising agents which are
more powerful than fluorine.
The table also clarifies why the alkali metals such as
sodium (Na) react so violently with the hydrogen ions in
neutral water. The half-reaction equations of these
metals are above that of hydrogen ions and so they
reduce them. Zinc only reduces water in the presence of
+
additional H ions, i.e. in an acid medium ( see
illustration R9 for the influence of ion concentration on
reduction potentials).
The table also quantifies the qualitative difference in
2+
oxidising power of the Zn
2+
/Zn, Pb
/Pb and Cu
2+
/Cu
and quantifies the redox reactions depicted in illustration
R5. It is the convention for measuring the cellpotential
(E ) or electromotive force (e.m.f.) of electrochemical
cell
cells always to place the more negative electrode system
(the anode) on the left :
This leads to :
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The fact that one reagent is an element to be found at
the top right and the other reagent is an element to be
found at the bottom left of the standard potential table
does not always guarantee a self-sustaining redox
reaction.
Some cases require something more: an impulse, a push
or a poke, in the form of temperature or in the form of a
catalyst. In this way H O , which can safely be stored in
2
2
(dark) bottles, will yet start to spontaneously
disintegrate as soon as a pinch of a catalysing substance
is added to it. Interesting to see here is the possible
impact of iron (II) salts and iron (III) salts: their
half-reaction is somewhere between the two
half-reactions for hydrogen peroxide.
Useful as it may be, the table of standard potentials does
not allow prediction with certainty of which selfsustaining
redox reactions are actually possible.
It is said that the table of standard potentials is
normative for the thermodynamic tendency of reductors
and oxidants to react in a self-sustaining way. No matter
how big the difference between their E°-values is, this
does not say anything about the kinetics of the redox
conversion, i.e. the ease and the speed with which this
can happen in practice.
A lack of visible reaction upon combination of redox
couples from the tables, which thermodynamically should
produce a self-sustaining reaction, may indicate a kinetic
problem or could also indicate the formation of an
insoluble salt on the surface of the metal, especially if we
think of lead in an aqueous solution of zinc ions. The zinc
salt should be choses with care and zinc nitrate is
probably the most suitable for use here.
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R08 Concentration dependence of reduction potentials
The reduction of an acidified solution of
ions to
2+
Mn ions clearly illustrates the influence of ion
concentrations upon the reduction potential observed.
This reaction can be described by the equation:
That this is a 5 electron reduction reaction can be
checked by considering the change in the oxidation
number of manganese from +VII in
to +II in Mn
ions. The reaction quotient for this reaction,
Aim: To show the relationship between the actual
reduction potential and the standard reduction
potential and the concentrations of the reagents
present.
The standard reduction potential of a half-reaction
equation when measured with respect to a standard
hydrogen electrode is expressed in V. For a standard
hydrogen electrode:
2+
The expression, Q , should not be confused with the
c
equilibrium constant of a reaction which is always
measured at equilibrium, the reaction quotient only
relating to the concentrations at a particular moment in
the reaction i.e. not necessarily at equilibrium.
The concentration factor for the solvent, water , can be
omitted from the equation since it can be taken as 1
(see the Chemical equilibria module in DIDAC 2).
Thus at 25°C :
= 0 V at a temperature of 25 °C, a partial pressure of
hydrogen of 1 atmosphere (1013 hPa) and a
+
concentration of H ions of 1 mol/L (i.e. pH = 0).
The temperature, the partial pressure of gases, if the
reagents are gases, and the concentrations of the
reagents involved in the half-reaction equation have an
effect on the electrochemical potential.
The first relationship between these factors was found by
Nernst, a German chemist and physicist, who in 1920
received a Nobel prize for his contributions to
thermodynamics and electrolyte solutions.
If the respective concentrations are known then the
actual value for E at 25°C can be calculated.
It can be seen from the above equation that E is
dependent on the concentration of hydrogen ions i.e. the
acidity of the solution. As the hydrogen ion concentration
increases, i.e. pH decreases, E becomes more positive
i.e. the oxidising strength of the reducing agent
increases.
Nernst stated that:
Note :
• From the Nernst equation for the actual half-reaction
it can be deduced that - at least for a
pH-dependent half-reaction - the oxidising capacity
+
increases (the E value becomes more positive) as [H ]
increases, i.e. as the pH decreases. This explains why
for a pH-dependent oxidant we prefer to represent the
half-reaction as for an acid environment.he
half-reaction as for an acid environment.
• For a pH-dependent reductor the reducing capacity
+
increases (more negative E value) by keeping [H ] low
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-
or [OH ] high.
and log 10, the conversion factor from ln to log, is 2.3.
e
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R09 Electrode-concentration cells
Hydrogen is in its gaseous state therefore its partial
pressure in atmospheres is used in the Nernst equation;
in this case p(H ) = 1 atmosphere and the concentration
2
+
of H = 0.1, hence the reduction potential of the anode,
E , is
A
This value indicates that the left-hand half-cell is a
stronger reducing agent than the right-hand standard
half-cell. The lower the concentration of hydrogen ions in
the left-hand half-cell , i.e. the higher the pH, the more
negative the potential of the left-hand half-cell. The
potential for this cell, E , is the difference between the
Aim: To show that a difference in concentration
gives rise to a potential difference in a cell
consisting of two otherwise identical hydrogen
reference electrodes
cell
reduction potential of the cathode (right-hand half-cell)
E and the reduction potential of the anode (left-hand
C
half-cell), E :
A
Illustration R9 shows an electrochemical cell consisting of
two hydrogen electrodes as half-cells, only differing in
the concentration of hydrochloric acid: 0.1 mol/L in the
half-cell on the left and 1 mol/L in the half-cell on the
right. Hydrogen is passed over both platinum electrodes
at a pressure of 1 atmosphere (1013 hPa) and the
temperature is 25°C. The right-hand electrode is
therefore a standard hydrogen electrode and
=0V.
The pH of the left-hand half-cell is 1 (Concentration of H
= 0.1 mol/L). The half-reaction equation for both
half-cells is:
+
This can be seen to have a positive value numerically
equal to the value for E calculated using the Nernst
A
equation. The cell convention is to assign a positive
value of E
to an electrochemical cell-reaction when
cell
the reaction is written down in the direction of
self-sustaining change. A negative value for
is
required for a self-sustaining (spontaneous) reaction and
since
= -nFE
it is clear that if E
is positive
is
cell
Which is a 2 electron reaction with an expression for Q
c
of :
The voltmeter detects a potential difference between the
two half-cells. This can only result from the difference in
concentration, since all other factors are identical The
actual reduction potential for the left-hand, non-standard
half-cell reaction can be determined using the Nernst
equation:
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cell
negative.
The fact that the potential difference between two
otherwise identical cells varies with pH is used in a
potentiometric pH meter, despite the fact that the
hydrogen electrode is not well suited for routine pH
measurement since it is a source of gaseous hydrogen
and is sensitive to various poisons that inhibit the
activity of the platinised surface of the electrode. This
potential difference can also be used to calculate an
unknown concentration of a solution using the same
electrochemical cell as shown in illustration R9.
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R10 Non-standard reduction potentials
+
-7
for p(H ) = one atmosphere and [H ] = 10
2
mol/L.
Thus :
For the bottom set of E values on illustration R10 i.e. the
half-reaction O /2H O :
2
2
Aim: To compare actual reduction potentials with
standar reduction potentials for two very common
half-reactions
Illustration R10 shows the
and actual reduction
potential values, E, for the two half-reaction equations:
The process represents the reduction of oxygen from
oxidation number 0 to -II. Again the non-standard
reduction potential can be calculated using the Nernstequation.
at various pH's
For each half-reaction equation there are four different
representations depending on the pH of the medium in
-
which they are to be found: two for bases (with OH ) and
+
two with acids (with H ). However only three E values
are shown for each half-reaction equation.
For the top set of E values on illustration R10 i.e. the
+
Note :
half-reaction 2H /H :
2
Most reactions in neutral aqueous solution are carried out
at non-standard reduction potentials. For the half-reaction
+
H /H the reduction potential used is -0.41 V and for the
2
half-reaction O /H O this is +0.81 V.
2
2
Illustration RR10 shows that the value does not change
-
The middle values are not standard reduction potentials
+
because the concentration of H is not 1 mol/L and
therefore the concentration is quoted in the table. Lack
of space prevents either the units or the physical state of
the reactants being quoted. Strictly speaking these
should be quoted as (g) (l) (s) (aq) depending on
whether the reagent is a gas, a liquid, a solid or an
aqueous solution. The values given at the top and
bottom of the table are the standard reduction potentials
1 of 2
+
whether the half-reaction is written with OH or H .
It can be concluded that hydrogen is a stronger reducing
agent and oxygen a stronger oxidising agent in water at
pH 7 than under standard conditions.
In a neutral solution hydrogen ions, which are present in
very small concentration, are less likely to oxidise zinc
metal than in an acid solution in which their concentration
is high.
This model can also be extended to consider oxygen gas in
acid solution in which it is more likely to oxidise iron than
in a neutral solution because the concentration of
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for each of the half-reactions. Fundamentally the process hydrogen ions is greater.
is the same, the reduction of hydrogen, H, from oxidation
number +I to 0. The
values can be found in standard
tables, but the non-standard reduction potentials (middle
values) are not. These can be calculated at 25°C for the
+
half-reaction 2H /H using the Nernst-equation:
2
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R11 Electrolytic cell versus electrochemical cell
A heavy arrow between reactants and products indicates
that the reduction goes to completion, but has no
bearing on the kinetic pathway. Two such reactions are
presented below:
Aim: To illustrate the relationship between a
electrochemical cell and an electrolytic cell or
between a self-sustaining redox reaction and a
sustained redox reaction.
Electron transfer reactions, like most chemical reactions,
can in principle proceed in either direction according to
the equation or even in both directions at once. In an
electrochemical cell the electrode reactions are
selfsustaining proceeding until one reactant is consumed
or until the reduction potentials of the two half-cells
become equal. In an electrolytic cell, on the other hand,
electrical energy is supplied which sustains a redox
reaction that would otherwise not take place : a so-called
sustained redox reaction.
The half -reactions of the Daniell cell (see illustration R3
and R11) proceed to completion without the use of a
catalyst or the addition of external energy. This is also
true for direct mixing of the same reagents (see
illustration R5), these reactions continuing as long as
there are reactants present.
The formation of water from hydrogen and oxygen can
also be written with a heavy arrow because this
transformation only stops when at least one of the
reagents is completely used up. This reaction is,
however, nonspontaneous, initiation either requiring a
catalyst or an increase in temperature of the reactants.
Once the reaction has started, however, it proceeds to
completion without additional energy being supplied.
This is an example of a non-spontaneous self-sustaining
redox reaction going to completion. These reactions both
proceed in the direction from the stronger reducing
agent and stronger oxidising agent to the weaker
oxidising agent and weaker reducing agent (see
illustration R10).
2. Self-sustaining equilibrium reactions
Cells, whether they be electrochemical or electrolytic, are
by convention always shown with the negative electrode
on the left-hand side i.e. in electrochemical cells electron
flow in self-sustaining redox reactions is always depicted
from left to right, whereas the flow of electrons supplied
to a redox reaction is always depicted from right to left.
2+
In illustration R11 two cells are shown with Zn
/Zn and
2+
Cu /Cu redox couples in their half-cells. An
electrochemical cell configuration is depicted on the
left-hand side, a Daniell cell in which electrons flow from
2+
the half-cell with the Zn
/Zn redox couple to that with
A redox reaction does not always proceed to completion
despite proceeding from the stronger oxidising agent
and stronger reducing agents to the weaker ones. If the
redox half-reaction couples have very similar standard
reduction potentials, the actual reduction potentials of
the redox half-reaction may converge as the reaction
proceeds, the reduction stopping when they become
identical. In the half-reactions
2+
the Cu /Cu redox couple in a self-sustaining redox
reaction which proceeds until either the zinc metal or the
copper(II) ions have been consumed.
An electrolytic cell is depicted on the right-hand side of
illustration R11 in which electrons, supplied by a battery,
flows from the half-cell with the Cu
2+
2+
/Cu redox couple to
that with the Zn /Zn redox couple until either the
copper metal or the zinc(II) ions have been consumed in
a sustained redox reaction.
1 of 3
at certain concentrations the difference in
reducing/oxidising capacity is so small, that the Gibbs
free energy change involved gives rise to the possibility
of a reaction in either direction. The system is said to
have reached equilibrium. This is called a reversible
equilibrium redox reaction and it can be represented by
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a double headed arrow.
The redox reactions taking place in these cell are
summarised below:
This reaction is self-sustaining in both directions (no
energy has to be supplied) and it is spontaneous (no
initial energy has to be provided). The system reaches
equilibrium when no further drop in Gibbs free energy is
possible and not when one of the reagents has been
used up.
The constant external supply of electricity forces the
redox reaction to take place and in the example quoted
zinc ions are forced to oxidise the copper metal which in
turn behaves as a reducing agent.
In this electrolytic cell the zinc electrode acts as the
cathode and the copper metal acts as the anode. The
cathode requires a constant supply of electrons which are
available from the electric current.
Cathodes and anodes
A cathode is the electrode at which reduction occurs. In
an electrochemical cell it receives electrons from the
anode and has a positive sign, whereas in an electrolytic
cell it is the electrode at which cations attracted to it are
reduced by the electrons supplied to it. The cathode of
an electrolytic cell has a negative sign due to electrons
supplied to it by the current source.
An anode is the electrode at which oxidation occurs. In
an electrochemical cell it supplies electrons to the
cathode and has a negative sign, whereas in an
electrolytic cell it is the electrode at which anions
attracted to it are oxidised. The anode of an electrolytic
cell has a positive sign due to its being connected to the
positive pole of the current source.
3. Non-self-sustaining reactions
A redox reaction can, like any other reaction, occur in
the reverse direction until one of the reagents is used
up. It is, for example, possible to change copper metal
back to copper ions by reacting it with zinc ions which in
turn become zinc metal. Such a redox reaction does not,
however, occur spontaneously and is never
selfsustaining. An external energy supply is required
during the whole process. This is called a sustained
redox reaction and is accompanied by an increase in
Gibbs free energy. In order to predict whether two
species – possibly after a short preparation or with the
help of a catalyst - can show a self-sustaining redox
reaction, we used the following practical rule: a species
at the TOP RIGHT (in the table) can in principle enter
into such a redox reaction with a species at the BOTTOM
LEFT.
This qualitative rule can also be expressed quantitatively
in terms of the electromotive force (e.m.f.) of the
electrochemical cell made of the substances written in
these half cells: a self-sustaining redox reaction is only
possible if the e.m.f. of this cell turns out to be positive.
For measuring the electromotive force (e.m.f.) or cell
potential (E ) of an electrochemical cell, it is the
cell
convention to substract the reduction potential of the
anode from the reduction potential of the cathode. Not
all redox reactions can be sustained by supplying
external energy. There may be insurmountable kinetic
and/or technical difficulties. A simple way of supplying
energy to a system is to apply an electric current i.e a
stream of electrons. Instead of producing energy as in
an electrochemical cell such as a Daniell cell, energy is
supplied to a cell which is then known as an electrolytic
cell. Supply of external energy, the electric current,
gives rise to redox reactions which can be predicted by
using the table of standard reduction potentials in
illustration R7. That way it is possible to obtain zinc
2+
In electrochemical and electrolytic cells with the same
redox couples in their half-cells, such as depicted in
illustration R11, the redox couples of the cathode and
anode of the electrochemical cell therefore become the
redox couples of the anode and cathode respectively of
the electrolytic cell
2 of 3
metal and Cu by continuously adding electrical energy
to a "cell" of zinc ions and copper metal. In such cases a
weak oxidant (the most oxidised form from a more
reducing redox couple) must react with a weak reductor
(the most reduced form of a more oxidising half cell).a
more oxidising half cell).
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Note :
Self-sustaining and non-self-sustaining redox
reactions
In a mixture A
Red
A
Ox
and B
Ox
+B
Ox
and B
Red
and A
Ox
+B
Red
, where A
Red
and
are both conjugate redox couples,
the following reactions are therefore possible:
1. Self-sustaining reactions proceeding to
completion
The voltage required from the external energy source is
often greater than that predicted by the difference,
,
of the two reduction potentials. Other reagents present
in aqueous solution can also be forced into a redox
reaction by this large voltage. In the example in
illustration R11 sulphate ions
, are attracted towards
the copper rod. The oxygen rather than the sulphur from
the
is oxidised, oxygen gas being formed. This
phenomenon can also be observed in the electrolysis of
water which has been acidified by H SO .
2
4
The main requirement for a self-sustaining reaction is
that there is a decrease in Gibbs free energy due to the
generation of heat a nd an increase in disorder. These
conditions are met for reactions proceeding from the
stronger reducing agent and stronger oxidising agent to
the weaker reducing agent and weaker oxidising agent.
Strong reducing agents are to be found at the top
righthand side of the table of redox couple reduction
potentials in illustration R7 and strong oxidising agents
at the bottom left-hand side of this table.
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R12 Atmospheric corrosion
In the absence of kinetic obstacles the redox reaction
can be predicted to be one which is self-sustaining.
The iron(II) ions which are formed in the water come
into contact with the rapidly diffusing hydroxide ions and
react to form insoluble iron(II) hydroxide. This is in turn
oxidised by the air to iron(III) oxide.
The cathode or cathodic zone is more than just the iron
metal. Iron together with the oxides FeO, Fe O , and
2
3
hydroxides Fe(OH) and Fe(OH) forms the complex
Aim: To show how atmospheric corrosion occurs
and which redox half-reactions are responsible
for it
Corrosion has posed a problem for centuries and is
evident in the home, garden, transport vehicles (from
bicycles to cars), shipping, industry and underground
piping. The consequences of corrosion are all too familiar,
parts have to be replaced, customers become dissatisfied
and there are other adverse financial consequences.
th
1/8 of the annual UK production of steel is needed to
replace iron lost through rusting.
Combating of corrosion often requires expensive surface
treatment e.g. painting, galvanisation, tinning etc., which
often is associated with products which themselves cause
serious contamination of land, water and air.
Rusting of iron.
Corrosion is a widespread problem which can be
explained in terms of redox reactions as represented on
illustration R12. The top right-hand side shows an iron
surface which is in contact with water containing a little
dissolved oxygen.
All the necessary reagents are present to form a
corrosion cell. (This is a simple electrochemical cell in
which corrosion occurs.) The etched out corrosion cell
depicted has an anodic area, a cathodic area, a suitable
transport medium for electrons (the metal itself) and an
aqueous solution through which ions can move. The
following electrochemical half-reactions are important
when considering how the anodic and cathodic areas are
formed:
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2
3
cathodic surface which is called rust.
Constant diffusion takes place between the anodic and
cathodic areas and, because the cathode receives its
oxygen supply from the air, this type of corrosion is
called atmospheric corrosion.
A few characteristics of the rusting process are:
• The presence of salt in the water leading to a greater
degree of corrosion, because it is ionic and when
dissolved in water its ions encourage the transport of
ions already within the system.
• Corrosion only continuing if the rust which is formed
conducts electrons. For iron this is the case but if the
iron is incorporated into a steel alloy e.g. stainless
steel ( 18% Cr, and 8% Ni ) then an insulating layer
of CrO forms on the surface of the iron and if a
2
rusting process starts it stops almost immediately due
to lack of electron transport.
+
• Presence of an acid results in H ions from the acid
reacting with the hydroxide ions produced during
corrosion to form water. This will result in the
formation of more hydroxide ions and more corrosion.
Corrosion at the junction of copper and iron
When two metals, one of which is iron, are joined
together it is important that corrosion of the iron is
prevented. At the bottom left-hand side of R12 a copper
bolt has been used to join two iron plates. Copper metal
has a much lower reducing power than iron. The redox
couples to be considered are:
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From the above
-values it can be deduced that Fe, the
metal surface, is the anode. The top of the water droplet
is in contact with the oxygen in the air and so the
concentration of oxygen is greater here than in the water
droplet itself. This region of higher concentration is the
main cathodic zone. The Nernst equation (R8) shows that
the oxidising potential increases with an increase in
oxygen partial pressure.
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The values of
are very similar. In non-standard
situations the order can even be reversed and the
oxidizing and reducing power lost. Copper corrodes very
little but is an excellent conductor of electrons. It
therefore provides an excellent cathodic surface in the
presence of iron ,thus increasing the corrosion of the
iron. The copper bolt still exists after the two iron plates
have corroded away!
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R13 Combating corrosion: surface treatment
When iron is galvanised with 0.1 to 0.5 mm layer of zinc
it is protected in two ways:
2+
the more negative
value of the Zn /Zn redox couple
ensures that corrosion of zinc to zinc oxide will be
favoured over oxidation of iron.
Furthermore, even when microscopic defects are present
in the zinc layer an electrochemical cell is set up with
2+
the more strongly reducing Zn
/Zn redox couple as an
2+
anode and the Fe /Fe redox couple as the cathode
resulting in the zinc anode being preferentially oxidised,
so protecting the iron underneath and preventing it from
rusting.
Although the layer of zinc is very thin the protection is
good for a long time.
Aim: To illustrate the advantages and
disadvantages of surface treatment in combating
corrosion
There are various ways in which corrosion can be
combated. These involve changing:
1. the anode surface or the anodic half-reaction.
2. the cathode surface
3. the electrolyte
4. the nature of the oxidising agent
5. the ability to transport electrons
The anodic surface can be changed by protecting the
metal with a thin layer of paint (lacquering), zinc
(galvanising) or tin (tinning).
When considering illustration R13, it is necessary to bear
in mind the table of standard reduction potentials for the
half-reactions involved in the process.
In the food industry iron is protected by a very thin layer
of tin in tin cans. The
of Fe
2+
/Fe redox couple is
2+
above that of Sn /Sn therefore if the tin layer in the
can is scratched it will rust since iron oxidises in
preference to tin as shown on the right-hand side of
illustration R13. Considerable localised rusting can occur
and food in dented tin cans should be avoided as these
can be rusty on the inside. The tin ions which may then
enter the food are, however, not very toxic. Despite
these disadvantages tin is used extensively in the food
industry as a good way of protecting the iron and the
chances of it being scratched on the inside are minimal.
Standard reduction potentials for the half-reactions
considered in illustrations R12 and R13.
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R14 Combating corrosion: sacrificial protection and electrolytic protection
As these sacrificial electrodes are oxidised away ions are
released, which in the cases of zinc and aluminium ions
have a polluting effect on the sea or ground water.
Electrolytic protection
Aim: To show that relatively small sacrificial
blocks of metal, or the application of a relatively
high potential difference can protect large surface
areas of metal from corrosion
Sacrificial protection
Examples of the use of sacrificial blocks of metal to
combat corrosion are to be found in illustration R14. The
table of standard reduction potentials in the text
accompanying illustration R13 are useful in clarifying the
processes involved.
Sacrificial protection of metals involves the use of small
blocks of metal which due to their higher
-values will
oxidise before the metal which they are there to protect.
They are replaced at regular intervals as they are
oxidized away.
Magnesium, aluminium or zinc are used in sacrificial
blocks, all having higher
-values than iron. Examples
of applications are :
An alternative to the use of sacrificial blocks in the
protection of the steel hulls of ships is to apply a
negative charge to the iron surface. An electrolytic cell is
thereby set-up with the same half-cells which in the
absence of this electron source make up the
electrochemical cell responsible for the corrosion
process. Whereas in the electrochemical cell the iron
surface acts as the anode at which the iron is oxidised
and oxygen in the water acts as the cathode at which
oxygen is reduced to hydroxide ions, the application of
an electric current results in iron acting as the cathode
making oxidation more difficult.
A decrease in the potential at the iron surface of
100 mV reduces the rate of its oxidation by a factor of
between 5 and 100.
In practice a cathodic current is applied to the metal
being protected and the metal is surrounded by a good
but inert conductor such as graphite which acts as the
anode, such is shown in the top left of illustration R14.
This method has the advantage over sacrificial blocks of
being easier to control and quantify, of requiring less
sacrificial material and of reducing pollution due to
corrosion products. It is used mainly in the
petrochemical industry to protect pipelines and storage
tanks, in concrete structures to protect the iron
reinforcement and in the shipping world to protect ships,
oil platforms and harbour installations. It is however an
expensive form of corrosion protection.
zinc blocks welded to the steel hulls of ships, as
shown in illustration R14;
zinc blocks welded to the steel supports of oil rigs;
and
magnesium, aluminium zinc or even graphite
blocks
electrically connected to pipe lines at regular
intervals, as shown at the top right of illustration
R14.
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R15 A lead-acid accumulator
The battery then functions as an electrochemical cell.
Electric current is produced by the self-sustaining redox
reaction between the top right substances with the
bottom left substances :
Each compartment in the battery corresponds to a cell
voltage
as long as concentrations and temperature are standard.
Aim: To describe the construction, operation and
characteristics of a lead-acid accumulator
An accumulator, or in more general terms a battery, is a
classic example of the application of electrochemical
reactions. It is an electrochemical cell, which
accumulates or stores electric charge. It consists of
chemical components, which function as the anode and
others which function as the cathode, connected together
by an electrolyte solution or an electrolyte paste and a
porous barrier or salt bridge between the anode and
cathode cells.
If a resistance, e.g. an engine, or a lamp is placed
between the anode and cathode, chemical energy will be
converted into electrical energy in a self-sustaining
electron transfer reaction. An accumulator is a chemical
source of electricity.
Batteries and accumulators are very much part of the
modern way of life. They vary in size from minute button
batteries, such as are found in watches, to large heavy
blocks, that function as emergency electricity sources for
hospitals, ships and vehicles. They are usually made for
a specific use and the lead-acid accumulator is widely
used as an electrochemical source of electricity in cars
and lorries. The construction and operation of a lead-acid
battery is shown in illustration R15.
Construction of a lead-acid battery
It consists of a polypropene container, containing six
lead-acid electrochemical cells, each giving an e.m.f. of 2
V (2.05 V), connected in series. While the voltage of a
lead-acid battery is always a multiple of 2 V, its power
can vary considerably. It can be increased by dividing
each cell with microporous polyethene dividers. All the
anode and cathode plates within a single cell are linked
together, the anode of one cell being linked with the
1 of 2
In use, lead-acid batteries gradually become less
efficient. The main redox reaction (but also other
processes) consume sulphuric acid. Since the mass
density of the electrolytic solutions increases with
concentration of H SO , any efficiency loss will result in
2
4
a decrease in the density of the battery acid. By
measuring the mass density (specific gravity) of the
electrolyte solutions, the efficiency of the battery can be
checked.
Charging:
A lead-acid battery has a reasonable lifetime due to a
possible reversion of the redox reactions.
The reversed reactions
can regenerate the substances used during the
discharging process. Charging is only possible by
continuing applications of a potential from an outside
power source, e.g. a car’s alternator or generator, or a
charger plugged into the mains.
2+
As long as current is supplied to the battery, Pb ions
are reduced to lead metal (at the lead electrode) while,
2+
at the PbO electrode, Pb
2
ions are oxidised to PbO . Pb
2
and PbO are deposited on the electrodes.
2
The half-cell reactions during charging and discharging
are exactly the same. The charging process however
involves the weaker oxidising agent (top left) and the
weaker reducing agent (bottom right) :
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anode of the next cell etc. and the cathodes of the cells
being likewise linked together.
The anodes are all made of lead and the cathodes all of
solid lead(IV) oxide or lead dioxide, which is mounted on
a lead base. Both poles of the battery are also made of
lead. The electrolyte is a 80% solution of sulphuric acid
saturated with PbSO .
4
Electrochemistry of a lead-acid battery
Discharging :
During discharge the electrode reactions to be taken into
account are, at the negative pole (now the anode)
at the positive pole (now th ecathode)
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R16 A zinc battery and a button cell
The overall reaction during discharge is:
The standard reduction potentials for the two
half-reactions can be used to calculate the standard
reduction potential for the cell:
In practice the cell potential is 1.55 V, the difference
being due to non-standard concentration of ions in the
cell (see illustration R8).
Aim: To explain the construction and
electrochemical processes of a small zinc battery
and a button cell
This cell is non-rechargeable due to the following
irreversible reactions :
Illustration R16 shows two different batteries which are
in everyday use. Their operation can be compared with
that of the lead-acid accumulator (see illustration R15).
Zinc battery
This battery, which has a rather unusual construction,
still carries the name of its inventor, the French scientist
Leclanché. It is the basis of the familiar 1.5 V batteries
used in torches and other household appliances requiring
batteries. The cathode consists of a central current
collector (which can for example be made of copper)
surrounded by a paste of MnO , KOH, ZnO and small
During the last few years, rechargeable batteries have
been developed using a similar construction by
preventing these irreversible reactions from taking place.
Button cells.
Button cells, which are used in watches, cameras,
calculators etc., work on a similar principle. The half-cellreactions are:
2
graphite particles. The graphite particles, which are
present in high concentration, are there to conduct
electrons to the central collector. The anode, which forms
the shell of the battery, is made of zinc. During discharge
of the battery this zinc layer is slowly oxidised to zinc
oxide, whilst in the cathode compartment Mn(IV) ions in
MnO are reduced at the surface of the graphite
2
particles.
The two electrode-compartments are separated by
waxed paper and/or wax.
The half-cell-reactions for the Leclanché cell are:
The reducing agent (anode) is once again metallic zinc.
The oxidising agent, which is contained in the cathode
compartment, is silver(I) oxide mixed with graphite
particles, mercury(II) oxide, and potassium hydroxide
(the electrolyte). The overall discharge reaction is:
These batteries are small and therefore only have a
small capacity.
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Zinc (the anode, at the top right of the table of the
standard reduction potentials in illustration R7) is the
reducing agent, being oxidised to zinc ions.
The oxidising agent is the MnO (in the cathode
2
compartment) which is found at the bottom left of a
standard table of reduction potentials. Mn(IV) ions are
reduced during discharge, its oxidation number being
changed from +IV to +III.
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R17 An electrolysis chamber
Illustration R17 has been included to show that
electrolysis is not just of theoretical interest, but is used
industrially on a large scale. The illustration shows a
huge electrolysis chamber, in which impure (96-98%)
copper is purified to 99.9 % copper. Copper metal with
this purity has the high conductivity required for a wide
range of electrical applications.
The illustration provides a link with transparencies R18
and R19, which discuss the electrolytic production of
NaOH and Cl .
2
Illustration R20 deals with electrolytic refining and
extraction.
Aim: To give an idea of large scale industrial uses
of electrolysis
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R18 - R19 Producing chlorine in the chlorine-alkali industry: the mercury cell - diaphragm-, and
membrane-cells
Aim: To describe schematically processes used
world-wide for the electrolytic production of
NaOH and Cl2, two of the most important basic
industrial chemicals.
The chlorine-alkali industry is an important branch of the
chemical industry, producing chlorine and sodium
hydroxide by the electrolysis of common salt. The main
raw material is brine, a saturated aqueous solution of
sodium chloride (NaCl) obtained from natural salt
deposits.
50 Million tons of Cl are produced annually. Whilst some
2
of this is marketed as chlorine gas, e.g. to disinfect
drinking water and swimming pools, most of it is used in
the synthesis of (poly)vinylchloride and other
chlorohydrocarbons, such as chlorohydrocarbon solvents,
cooling liquids, insecticides, pesticides, fungicides etc.
Electrolytic production of chlorine also generates two
useful by-products: sodium hydroxide, and hydrogen.
The half-reaction equations give an interesting insight
into the process:
+
with Na , H , or H O molecules; all three are above and
2
to the left of Cl in the list of standard reduction
2
potentials . None of these reactions would be
1 of 3
2
This method only produces a fraction of the chlorine and
sodium hydroxide used by industry as it has certain
disadvantages: mercury is expensive and toxic, and
whilst it is recirculated, some always escapes with the
spent brine with which it reacts to form mercury(II)
chloride.
In the past this effluent was discharged into lakes and
rivers, leading to the accumulation of high levels of
mercury in fish, which absorbed the mercury compound
but could not re-excrete it. Nowadays the spent brine is
treated before discharge, the mercury being precipitated
as mercury(II) sulphide.
In recent years a large share of chlorine and sodium
hydroxide production has been produced in two other
types of cell, which do not use mercury: the membrane
cell and the diaphragm cell. In the UK 1 diaphragm cell is
in use for every 20 mercury cells and in the USA 2
diaphragm cells are in use for every mercury cell.
2
-
In theory chlorine could be produced by reacting Cl ions
+
Industrial installations consist of some 200 mercury cells
in series, each one measuring 15 m x 2 m x 0.3 m. The
process, which takes place at a very high voltage, uses
an enormous quantity of electricity: 3 MWh per ton Cl .
2. The diaphragm cell.
This is illustrated on the left of illustration R19. On the
negatively charged cathode surface NaOH and H are
-
The concentrations of the H and OH are those for
neutral water and the E value is therefore -0.41 V.
+
The brine consumed in the electrolyte cell is continually
replenished. The overall reaction is:
formed directly and Cl is formed at the anode. To
2
separate the chlorine from the sodium hydroxide, the
two half-cells were traditionally separated by a porous
asbestos diaphragm, which needed to be replaced every
two months. This was environmentally detrimental owing
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self-sustaining. For them to proceed, a constant source
of electrical energy would be required: i.e. an electrolytic
process would be needed. An electrolysis process, in
-
+
which Cl ions reduce Na ions, would, in theory, give
rise not only to Cl , but also Na(s). The sodium metal
2
would, however, immediately react with any water and/or
+
H ions present in the cell to produce hydrogen gas and
-
OH ions (i.e. NaOH solution).
Such a process would be very useful, since three
important chemicals would be produced during a single
process. However, there is a problem since the chlorine,
produced in the presence of a basic solution of sodium
-
hydroxide, would combine with it to form ClO ions and
-
Cl ions. This results in the production of sodium
chlorate(I), NaClO, a component of household bleach. To
overcome this problem the chlorine and sodium
hydroxide must be removed from the cell before they can
react.
Three industrial electrolytic processes are currently used
for the production of chlorine, all of which have
overcome the above-mentioned problem.
1. The mercury cell.
This cell, schematically shown in illustration R18, is used
in the electrolytic production of both sodium hydroxide
and chlorine from a saturated solution of brine and
operates at 4.4 V. Electrolysis of brine would normally
generate hydrogen at the cathode, but if mercury is used
as the cathode material this does not occur. At a
negatively charged mercury cathode hydrogen has a high
overvoltage, meaning that a higher negative potential is
required for the discharge of hydrogen ions. At the same
time the discharge potential of Na+ ions is lowered, since
the sodium atoms combine with mercury to form an
amalgam. This also protects the sodium from contact
and therefore reaction with water so the problem of the
alkaline medium is avoided.
to the need of disposing of large quantities of asbestos.
Such frequent replacement is fortunately no longer
necessary, the asbestos having now been replaced in
part by polymers resulting in diaphragms with a much
longer life.
The anode is made of a titanium-steel alloy, and the
cathode of steel. Calcium- and magnesium-ion impurities
must be removed from the brine before it is electrolysed,
otherwise they will precipitate out as insoluble
hydroxides and block the pores of the diaphragm. To
ensure seepage of brine through the diaphragm from the
anode to the cathode, the level is kept higher in the
anode compartment. The diaphragm cell is now
technologically the most advanced of all three cells and
has a high electrochemical performance.
3. The membrane cell.
A membrane cell is illustrated on the right of illustration
R19. This is very similar to the diaphragm cell, and the
same reactions occur. The main difference is that the two
electrodes are here separated by an ion-selective
polymer membrane which only allows cations to pass
through, instead of an asbestos diaphragm. Brine is
pumped in at the top of the anode compartment, and
water is introduced at the top of the cathode
compartment.
At the negatively charged cathode, hydrogen ions in the
water are reduced to hydrogen gas. At the positively
-
charged anode Cl ions from the brine are oxidised to Cl
2
+
gas. The Na ions flow through the membrane to the
cathode compartment, thereby carrying the current and
-
form NaOH with the leftover OH ions. The chloride ions
cannot pass through, so the chlorine does not come into
contact with the sodium hydroxide which is formed in the
cathode compartment. The sodium hydroxide is removed
from the bottom of the cell.
The overall reaction for the diaphragm and membrane
cells is:
Chlorine is produced at the positively charged anode,
which traditionally consisted of a series of suspended
graphite rods but are now being replaced by more
expensive, but more durable ,Ti or Pt-steel alloys. The
chorine which is formed collects at the top of the cell.
The sodium amalgam is run off from the bottom of the
cell into a separate chamber containing graphite balls.
The graphite catalyses the self-sustaining dissociation of
sodium amalgam. The sodium released reacts with water
to form sodium hydroxide and hydrogen, the mercury
being recovered and returned to the electrolysis cell.
2 of 3
The advantage that the two membrane cells have
(relative to the mercury cell) is that the Cl -gas and
2
NaOH do not come into contact with each other. They
also use less electricity and are therefor cheaper to
operate.
Reference :
Buchner, Schliebs, Winter and Büchel, Industrial
Inorganic Chemistry, VCH, Weinheim, Germany.
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R20 Electrolytic extraction and purification of metals
The resulting molten salt electrolyte is electrolysed with
graphite anodes in a cell lined with graphite which acts
as the cathode. The overall electrolytic reaction is:
where Q indicates the consumption of electricity.
Aluminium is discharged at the cathode and oxygen is
evolved at the anode, which oxidises the graphite anode
to carbon dioxide.
The appropriate half-reaction equations are:
Aim: To describe the installations and processes
used to extract one metal (Al) from its ore and to
purify another (Cu).
Electrolysis is frequently used to extract a metal from its
ore, or to purify an impure metal.
1. Electrolytic extraction of metals.
The commonest metal ores contain the metals as oxides
or carbonates. The metal can in theory be obtained by
the electrolytic reduction of the ores. However, in
practice chemical reduction using carbon or sulphur is
more often used because it is cheaper, but this produces
an impure metal and can cause environmental problems.
When product purity is important, electrolysis is the
preferred method. Al, Co, Cr, Mn, Ta, and Mg are all
extracted using electrolytic processes.
Illustration R20 schematically shows an electrolytic
process for the extraction of Al from its ore (bauxite, an
oxide of aluminium containing silicon and other
impurities). This process is used to produce 25 million
tons of aluminium per year. After chlorine, this is the
most important product of the electrochemical industry.
Aluminium is in great demand for the automobile,
shipbuilding, aircraft, electrotechnical and building
industries and, since there are ample reserves of
bauxite, the future for the aluminium industry seems
bright.
The electrolytic extraction process of aluminium from
bauxite was originally developed by Hall (USA) and
Héroult (France) in 1886 and improved in 1887 by Bayer
(Germany).
The first step in this process is the dissolution of bauxite
in sodium hydroxide under pressure as sodium
aluminate. This is a self-sustaining reaction, which does
not involve electron transfer. The SiO and Fe O
2
2
3
impurities are precipitated and removed at this stage.
1 of 2
The second equation contains a multiple of 3 in order to
balance the electrons in the two half-equations.
This process is operated under the following conditions:
Voltage = 4.5 V (cf. theoretical value of 2.2 V).
Current = 150 kA
Cell size = 3 x 8 x 0.7m, containing 8 graphite
3
blocks (200 single cells, each 15 m , in series).
15 MWh of electricity are required to produce 1
ton of aluminium, which is 5 times more electricity
than is required to produce 1 ton of chlorine.
2. Electrolytic refining of metals.
Impure metals can be purified by electrolysis. In an
electrolytic cell the anode is made from the crude metal
needing to be purified, the cathode from the purified
metal. The electrode potential is selected to ensure the
very selective reduction of the metal at the cathode.
During the transfer of the metal from the impure metal
of the positively charged anode to the negatively
charged cathode the impurities remain behind in the
electrolyte solution. The electrolyte is chosen according
to which element is to be purified. For Cu, Ag, Au and Pt,
aqueous solutions are used, whereas for Na, Mg, Ca and
Al molten salts are employed.
Illustration R20 illustrates the electrolytic refining of
copper. Impure copper is the anode, pure copper the
cathode and an aqueous mixture of sulphuric acid and
copper sulphate is the electrolyte. The electrode
potential is carefully selected so that only copper is
reduced at the cathode.
This is the most important method of copper
purification, producing 100,000 tons of purified Cu per
year. It is based on the following considerations:
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The next step is the precipitation of aluminium oxide
from the sodium aluminate solution. This is achieved by
dilution with water, seeding with solid aluminium(III)
hydroxide or treating with carbon dioxide.
The precipitated aluminium hydroxide is then separated
off and heated to obtain aluminium oxide, which is then
dissolved in molten cryolite (Na AlF ).
3
6
Ag, Au and Pt, all precious metals, have a lower
reduction potential than Cu. They are not oxidised
to ions at the anode. As the copper ions are
formed, the anode crumbles away allowing the
precious metal impurities to fall as sludge to the
bottom of the cell.
Sn, Bi and Sb have a larger reduction potential
than Cu. They are therefore oxidised at the anode
but their ions react with the electrolyte to form
insoluble oxides and hydroxides. These are also
deposited in the sludge.
Pb is also oxidised, but forms insoluble PbSO ,
4
which again sinks into the sludge.
Fe, Ni, Co and Sn are oxidised at the anode, but
remain as ions in the electrolyte. This is a
consequence of the careful choice of electrode
potential, which ensures that only the copper ions
are reduced at the cathode, the other metal ions
needing a higher electrode potential to be
reduced.
Electrolytic methods are also used to recover the
precious metals from the anodic sludge.
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R21 Electrolytic methods in electronics
2. Electroless deposition of surface layers
As with electrodeposition, the aim here is to reduce
metal ions to metal, which is then deposited as a
continuous layer on the surface of the substrate. In
electroless deposition the driving force for the
reduction is provided by an additional electro-active
component in the bath: a reducing agent. For this
process to work well, the surface of the substrate
must have a sufficiently catalytic character, that both
the reduction of the metal ions and oxidation of the
reducing agent can occur rapidly on the (catalytic)
surface. To ensure that the surface of the metal
substrate is sufficiently catalytically active, it is
pretreated with an etching bath, the metal thereby
achieving the highest possible surface area and traces
of oil etc.being removed at the same time.
Aim: To show and illustrate that many surface
treatments for metals and non-metals are based
on electrolytic methods
Many everyday objects have been subjected to some
form of surface treatment, take, for example, automobile
parts, kitchen utensils, cans for preserving food, building
materials (for window frames or roofing) etc. Similar
techniques are used in the production of electronic parts,
such as printed circuit boards, electrical contacts and
capacitors. For the most part electrolytic treatments are
used, which can be illustrated by the following examples.
1. Electrolytic deposition of metals and alloys
The aim of electrodeposition is to modify a metal’s
surface to obtain certain surface properties: hardness,
wear- and corrosion-resistance, gloss etc. The adhesion
between the deposited layer and the substrate must be
perfect. The principle of electroplating is simple:
electrolysis, i.e. a sustained redox reactions. The object
to be treated is the cathode, the anode being a
conductive inert material (Pt or Ti alloys) or a pure
sample of the material to be deposited. Illustration R21
shows the plating of a printed circuit board with copper,
the printed circuit boards being the negatively charged
cathode and the copper the positively charged anode.
The electrolyte is the most critical component of the cell.
It contains a suitable salt of the metal to be deposited,
usually complex salts since these tend to have a higher
stability and solubility than common salts such as
chlorides and sulphates. In the example in illustration
R21, the electrolyte is a mixture of copper(II) sulphate
and sulphuric acid. The tricks of the platers’ trade lie in
the formulation of the bath, small quantities of various
(usually organic) additives being present, which are
responsible for the deposited layer acquiring the required
properties of :
1 of 2
Phosphoric acid and formaldehyde are typically used
as reducing agents in electroless deposition
formulations.
Since the power of the reducing agent can be
pH-dependent, the buffering of the bath is also
important. Once again, the deposited layer is
1-100 !m thick.
3. Electrochemical conversion of surfaces
The presence of a passive layer on the surface of a
metal can increase its corrosion resistance, isolate it
electrically or improve its appearance. In this case
oxides, phosphates or chromates are typically
deposited on the surface.
Anodising
This process is used for aluminium, titanium,
copper, steel, tantalum and niobium surfaces.
During anodising, the surface of the metal is
converted into its oxide, the metal being the
anode of an electrolytic cell with a solution of
sulphuric, phosphoric or oxalic acid, as an
electrolyte.
The half-reactions involved in the care of the
anodizing of aluminium are:
The overall electrolyte reaction is:
The current required varies between 1 and
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2
Gloss: the additives control the micro-roughness
of the deposited layer. The mechanism of this
effect is usually unclear.
Wetting: hydrogen gas is usually formed during
the electrodeposition. Gas bubbles can become
trapped under the deposited layer, giving adhesion
problems.
This can be avoided by appropriate additives.
Layer structure: certain additives modify the
crystalline structure of the deposit, giving
improved physical properties (ease of soldering,
corrosion-resistance, hardness).
The thickness of the electrodeposited layer is in the
range of 1-100 !m.
100 A/m .
Depending on the duration of the treatment,
layerthicknesses between 0.5 !m (capacitors)
and 100 !m (building materials) can be
obtained.
Phosphating
Phosphating is mainly used to prepare surfaces
for further coating with paint or organic
coatings. The corrosion protection and adhesive
properties of the metal are noticeably improved
by this priming layer. The most important
applications are the phosphating of steel and
aluminium for use in the automobile and
building industries.
The chemical and electrochemical processes at
work here have not been fully characterised and
the technique is empirical rather than scientific.
Chromating
Chromate solutions, despite their toxicity, are
much used for depositing protective and
decorative layers especially on aluminium and
zinc. The most common applications are in the
food industry, e.g. chromate treatment of
aluminium cans.
In recent years the toxicity of chromates has
triggered a movement away from chromating to
chromophosphating or phosphating.
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Untitled Document
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R22 Electrochemistry in the blast furnace
Higher up the furnace the CO reacts with more coke in
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an endothermic reaction to produce carbon monoxide:
The carbon monoxide reduces the iron oxides
exothermically, the iron is formed dropping to the
bottom of the furnace, where the temperature is high
enough to melt it.
Aim: To illustrate the electrochemical processes
which occur between gases and solids in a blast
furnace
The production of iron in blast furnaces has been of
major economic importance for centuries. Illustration
R22 shows the processes involved. The reduction of iron
ore (oxides of iron(II), iron(III), silicon and other metals)
to iron with coke can be regarded as an example of a
redox reaction. A blast furnace is designed to realise the
following overall reaction:
A pool of molten iron forms on the bottom of the
furnace. Formation of a molten slag results from the
limestone (which is included in the charge together with
the iron ore and coke) dissociating to form calcium oxide
and carbon dioxide
and then combining with the silicon oxide and impurities
from the ore. This slag trickles down to the bottom and,
being less dense, forms a layer on top of the molten
iron.
The iron and slag are tapped off every few hours. A
modern furnace makes 3000 tons of iron a day.
For the sake of investigating the energetics of this
reaction, it can be regarded as taking place in the
presence of water. Under such conditions the reaction
can be presented by the following half-reaction equations
:
Although the reduction of iron(II) and iron(III) oxides
and the oxidation of coke are self-sustaining reactions,
the activation energy threshold for both is high. In reality
the processes are high temperature gas-solid reactions
taking place in the absence of water. In view of the high
activation energy threshold, the temperatures in the
reaction zones of the blast furnace need to be very high.
Such high temperatures are achieved by blowing hot air
over the coke, which is thereby initially oxidised
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exothermically to carbon dioxide:
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