In the experiment, the \(\text{Cu}^{2+}\) ions from the \(\color{blue}{\textbf{blue
copper(II) sulfate}}\) solution were reduced (gained electrons) to copper metal, which
was then deposited as a layer on the solid zinc. The zinc atoms were oxidised (lost
electrons) to form \(\text{Zn}^{2+}\) ions in the solution.
\(\text{Zn}^{2+}(\text{aq})\) is colourless, therefore the blue solution lost colour. As
discussed in Grade 11, the half-reactions are as follows:
Remember that there was an increase in the temperature of the reaction when you carried out
this experiment (it was exothermic). An exothermic reaction releases
energy. This raises a few questions:
Is it possible that this heat energy could be converted into electrical
energy?
Can we use a chemical reaction with an exchange of electrons, to produce
electricity?
If we supplied an electrical current could we cause some type of
chemical reaction to take place?
The answers to these questions are the focus of this chapter:
The energy of a chemical reaction can be converted to electrical potential
energy, which forms an electric current.
The transfer of electrons in a chemical reaction can cause electrical current
to flow.
If you supply an electric current it can cause a chemical reaction to take
place, by supplying the electrons (and potential energy) necessary for
the reactions taking place within the cell.
These types of reactions are called electrochemical reactions. An
electrochemical reaction is a reaction where:
a chemical reaction creates an electrical potential difference, and
therefore an electric current in the external conducting wires
or
an electric current provides electrical potential energy and electrons, and
therefore a chemical reaction takes place
Electrochemical reaction
An electrochemical reaction involves a transfer of electrons. There is a
conversion of chemical potential energy to electrical potential energy,
or electrical potential energy to chemical potential
energy.
Electrochemistry is the branch of chemistry that studies these
electrochemical reactions. An electrochemical cell is a device in which
electrochemical reactions take place.
Electrochemical cell
A device where electrochemical reactions take place.
temp text
Electrochemical reactions
Textbook Exercise 13.3
In each of the following equations, say which elements in the
reactants are oxidised and which are reduced.
The oxidation number of the ion \(\text{SO}_{4}\) is
\(-\text{2}\) as both a reactant and a product.
The oxidation number of \(\text{O}\) is
\(-\text{2}\). \(\text{O}\) is neither oxidised
nor reduced.
The oxidation number of \(\text{O}_{4}\) in
\(\text{SO}_{4}^{2-}\) is \(-\text{8}\).
Therefore the oxidation number of \(\text{S}\)
in \(\text{SO}_{4}^{2-}\) is \(\text{+6}\).
\(\text{S}\) is neither oxidised nor reduced.
The oxidation number of the ion \(\text{NO}_{3}\) is
\(-\text{1}\) as both a reactant and a product.
The oxidation number of \(\text{O}\) is
\(-\text{2}\). \(\text{O}\) is neither oxidised
nor reduced.
The oxidation number of \(\text{O}_{3}\) in
\(\text{NO}_{3}\) is \(-\text{6}\). Therefore
the oxidation number of \(\text{N}\) in
\(\text{NO}_{3}^{-}\) is \(\text{+5}\).
\(\text{N}\) is neither oxidised nor reduced.
The oxidation number of \(\text{Zn}\) in:
\(\text{Zn}(\text{s})\) is
\(\text{0}\)
\(\text{Zn}(\text{NO}_{3})_{2}(\text{aq})\)
is \(\text{+2}\)
The oxidising agent causes another reactant to be oxidised
and is itself reduced.
\(\text{H}^{+}\) remains as \(\text{H}^{+}\) in
\(\text{H}_{2}\text{O}\)
\(\text{Cr}^{3+}\) is a product.
\(\text{SO}_{2}\) contains \(\text{S}^{4+}\), while
\(\text{SO}_{4}^{2-}\) contains \(\text{S}^{6+}\).
Therefore \(\text{S}^{4+}\) loses two electrons and is
oxidised.
\(\text{Cr}_{2}\text{O}_{7}^{2-}\) contains
\(\text{Cr}^{6+}\). Therefore \(\text{Cr}^{6+}\) gains
three electrons to become \(\text{Cr}^{3+}\) and is
reduced. It is the oxidising agent.
d) \(\text{Cr}_{2}\text{O}_{7}^{2-}\)
There are two types of electrochemical cells we will be looking more closely in this chapter:
galvanic and electrolytic cells. Before we go into
detail on galvanic and electrolytic cells you'll need to know a few definitions:
Electrode
An electrode is an electrical conductor that connects the electrochemical
species from its solution to the external electrical circuit of the
cell.
There are two types of electrodes in an electrochemical cell, the
\(\color{blue}{\textbf{anode}}\) and the \(\color{red}{\textbf{cathode}}\).
\(\color{blue}{\text{Oxidation}}\) always occurs at the \(\color{blue}{\textbf{anode}}\)
while \(\color{red}{\text{reduction}}\) always occurs at the
\(\color{red}{\textbf{cathode}}\). So when trying to determine which electrode you are
looking at first determine whether oxidation or reduction is occurring there. An easy
way to remember this is:
Oxidation is the loss of electrons.
Reduction is the gain of electrons.
\(\color{blue}{\textbf{O}}\)xidation
\(\color{blue}{\textbf{i}}\)s
\(\color{blue}{\textbf{l}}\)oss of electrons
\(\color{blue}{\textbf{OIL}}\)
\(\color{red}{\textbf{R}}\)eduction
\(\color{red}{\textbf{i}}\)s
\(\color{red}{\textbf{g}}\)ain of electrons
\(\color{red}{\textbf{RIG}}\)
\(\color{blue}{\textbf{O}}\)xidation is
\(\color{blue}{\textbf{l}}\)oss of electrons at the
\(\color{blue}{\textbf{a}}\)node
\(\color{blue}{\textbf{An Ox}}\)
\(\color{red}{\textbf{Red}}\)uction is gain of electrons at
the \(\color{red}{\textbf{cat}}\)hode
\(\color{red}{\textbf{Red Cat}}\)
Table 13.1: A summary of phrases useful to help you remember the oxidation
and reduction rules.
The electrode is placed in an electrolyte solution within the cell. If the
cell is made up of two compartments, those compartments will be connected by a
salt bridge.
Electrolyte
An electrolyte is a solution that contains free ions, and which therefore
behaves as a conductor of charges (electrical conductor) in solution.
Salt bridge
A salt bridge is a material which contains electrolytic solution and acts as
a connection between two half-cells (completes the circuit). It
maintains electrical neutrality in and between the electrolytes in the
half-cell compartments.
Galvanic cells (ESCR5)
A galvanic cell (which is also sometimes referred to as a
voltaic or wet cell) consists of two half-cells, which
convert chemical potential energy into electrical potential energy.
Galvanic cell
A galvanic cell is an electrochemical cell which converts chemical potential
energy to electrical potential energy through a spontaneous chemical
reaction.
In a galvanic cell there are two half-cells. Each half-cell contains an electrode in an
electrolyte. The separation is necessary to prevent direct chemical contact of the
oxidation and reduction reactions, creating a potential difference. The electrons
released in the oxidation reaction travel through an external circuit (and do work)
before being used by the reduction reaction.
Remember that cells are not two-dimensional, although when asked to sketch a cell you
should draw it as shown in Figure 13.5 or Figure
13.6.
In a galvanic cell (for example the cell shown in Figure 13.5):
The electrons released in the oxidation of \(\text{X}\) remain on the anode,
\(\text{X}^{+}\) moves into solution.
The \(\text{Y}^{+}\) ions are being reduced at the cathode and forming solid
\(\text{Y}\).
The metal at the anode is \(\color{blue}{\textbf{X}}\).
\(\color{blue}{\textbf{Ox}}\)idation is loss of electrons at the
\(\color{blue}{\textbf{An}}\)ode.
The \(\color{blue}{\textbf{anode half-reaction}}\) is
\(\color{blue}{{\textbf{X(s)}} \to
{\textbf{X}}^{+}{\textbf{(aq) + e}}^{-}}\)
This half-reaction occurs in the half-cell containing the
X(s) anode and the \(\text{X}^{+}(\text{aq})\)
electrolyte solution.
The electrons released in the
\(\color{blue}{\textbf{oxidation}}\) of the metal remain
on the \(\color{blue}{\textbf{anode}}\), while the metal
cations formed move into solution.
The metal at the cathode is Y. \(\color{red}{\textbf{Red}}\)uction is gain of
electrons at the \(\color{red}{\textbf{Cat}}\)hode.
The \(\color{red}{\textbf{cathode half-reaction}}\) is
\(\color{red}{{\textbf{Y}}^{+}{\textbf{(aq) + e}}^{-}
\to {\textbf{Y(s)}}}\)
This half-reaction occurs in the half-cell containing the
Y(s) cathode and the \(\text{Y}^{+}(\text{aq})\)
electrolyte solution.
At the \(\color{red}{\textbf{cathode}}\) metal ions in the
solution are being \(\color{red}{\textbf{reduced}}\)
(accepting electrons) and deposited on the electrode.
There are more electrons at the anode than at the cathode.
Electrons will flow from areas of high concentration to areas of low
concentration, therefore the electrons move \(\color{blue}{\textbf{from
the anode}}\), through the external circuit, \(\color{red}{\textbf{to
the}}\) \(\color{red}{\textbf{cathode}}\)
Conventional current is measured as a flow of positive charge and so is in
the opposite direction (from the cathode to the anode)
The overall reaction is: \(\text{X}(\text{s}) +
\text{Y}^{+}(\text{aq})\) \(\to\)
\(\text{X}^{+}(\text{aq}) + \text{Y}(\text{s})\).
To represent this reaction using standard cell notation we
write the following:
The \(\color{blue}{\textbf{anode}}\) is always written on the
\(\color{blue}{\textbf{left}}\).
The \(\color{red}{\textbf{cathode}}\) is always written on
the \(\color{red}{\textbf{right}}\).
The anode and cathode half-cells are divided by \(||\)
representing the salt bridge.
The different phases within each half-cell (solid (s) and
aqueous (aq) here) are separated by \(|\).
The electrodes in each half-cell are connected through a wire in the external
circuit. There is also a salt bridge between the individual half-cells.
A galvanic cell uses the reactions that take place at at the two electrodes to produce
electrical energy, i.e. the reaction occurs without the need to add energy.
A spontaneous reaction is one that will occur without the need for external energy.
Refer to the subsection on spontaneity for more information.
The zinc-copper reaction you performed in Grade 11 can be modified to make a galvanic cell.
Bars of zinc and copper are used as electrodes, with zinc(II) sulfate
and copper(II) sulfate solutions as the electrolytes.
In the galvanic cell experiment, make sure that the sodium chloride paste is highly
concentrated and fills the U-tube for the best results.
A galvanic cell
Aim
To investigate the reactions that take place in a galvanic cell.
Weigh the copper and zinc plates and record their mass.
Pour \(\text{200}\) \(\text{ml}\) of the zinc sulfate
solution into a beaker and place the zinc plate in the
beaker.
Pour \(\text{200}\) \(\text{ml}\) of the copper(II) sulfate
solution into the second beaker and place the copper
plate in the beaker.
Fill the U-tube with the \(\text{NaCl}\) paste and seal the
ends of the tubes with the cotton wool (making a
salt-bridge). The cotton will help stop the paste from
dissolving in the electrolyte.
Connect the zinc and copper plates to the zero-centered
ammeter and observe the ammeter.
Place the U-tube so that one end is in the copper(II) sulfate
solution and the other end is in the zinc sulfate
solution. Observe the ammeter.
Take the ammeter away and connect the copper and zinc plates
to each other directly using copper wire. Leave to stand
for about one day.
After a day, remove the two plates and rinse them: first with
distilled water, then with alcohol, and finally with
ether (if available). Dry the plates using a hair dryer.
Weigh the zinc and copper plates and record their mass.
Note
A voltmeter can also be used in place of the zero-centered ammeter. A
zero-centered voltmeter will measure the potential difference across the
cell (not the flow of electrons), while an ammeter will measure the
current.
Discussion
Did the ammeter record a reading before the salt-bridge was
placed in the solutions?
Did the ammeter record a reading after the salt-bridge was
placed in the solutions? If yes, in what direction does
the current flow?
Fill in the table below:
Plate
Initial mass
Final mass
Zinc
Copper
How did the mass of the zinc and copper plates change?
Based on what you know of oxidation and reduction, why did
those mass changes take place?
Which electrode is the anode and which is the cathode?
Results
During the experiment, you should have noticed the following:
When the salt bridge was absent, there was no reading on the
ammeter.
When the salt bridge was connected, a reading was recorded on
the ammeter.
The direction of electron flow is from the zinc plate towards
the copper plate, meaning that conventional current flow
is from the copper plate towards the zinc plate.
After the plates had been connected directly to each other
and left for a day, there was a change in their mass.
The mass of the zinc plate decreased, while the mass of
the copper plate increased.
\(\color{blue}{\textbf{O}}\)xidation
\(\color{blue}{\textbf{i}}\)s
\(\color{blue}{\textbf{l}}\)oss of electrons,
\(\color{red}{\textbf{R}}\)eduction
\(\color{red}{\textbf{i}}\)s
\(\color{red}{\textbf{G}}\)ain of electrons.
The zinc electrode lost mass. This implies that solid Zn
metal atoms become ions and move into the electrolyte
solution: \(\text{Zn}(\text{s})\) \(\to\)
\(\text{Zn}^{2+}(\text{aq}) + 2\text{e}^{-}\). Oxidation
occurs at the zinc electrode.
The copper electrode gained mass. This implies that the Cu
metal ions in the electrolyte solution become metal
atoms and deposit on the electrode:
\(\text{Cu}^{2+}(\text{aq}) + 2\text{e}^{-}\) \(\to\)
\(\text{Cu}(\text{s})\). Reduction occurs at the copper
electrode.
\(\color{blue}{\textbf{Ox}}\)idation is loss of electrons at
the \(\color{blue}{\textbf{an}}\)ode. Oxidation occurs
at the zinc electrode, therefore the zinc plate is the
anode.
\(\color{red}{\textbf{Red}}\)uction is gain of electrons at
the \(\color{red}{\textbf{cat}}\)hode. Reduction occurs
at the copper electrode, therefore the copper plate is
the cathode.
When \(\text{Zn}(\text{s})\) \(\to\)
\(\text{Zn}^{2+}(\text{aq}) + 2\text{e}^{-}\) the
electrons are deposited on the anode, which becomes
negatively charged.
When \(\text{Cu}^{2+}(\text{aq}) + 2\text{e}^{-}\) \(\to\)
\(\text{Cu}(\text{s})\) the electrons are taken from the
cathode, which becomes positively charged.
Conclusions
When a zinc(II) sulfate solution containing a zinc plate is connected by a
salt bridge to a copper(II) sulfate solution containing a copper plate,
reactions occur in both solutions. The decrease in mass of the zinc
plate suggests that the zinc metal electrode has been oxidised to form
\(\text{Zn}^{2+}\) ions in solution. The increase in mass of the copper
plate suggests that reduction of \(\text{Cu}^{2+}\) ions has occurred
here to produce more copper metal.
The important thing to notice in this experiment is that:
the chemical reactions that take place at the two electrodes cause an
electric current to flow through the external circuit
the overall reaction must be a spontaneous redox reaction
chemical energy is converted to electrical
energy
the zinc-copper cell is one example of a galvanic cell
It was the Italian physician and anatomist Luigi Galvani who marked the birth of
electrochemistry by making a link between chemical reactions and electricity. In
1780, Galvani discovered that when two different metals (copper and zinc for
example) were connected to each other and then both touched to different parts
of a nerve of a frog leg at the same time, they made the leg contract. He called
this 'animal electricity'.
In the zinc-copper cell, the copper and zinc plates are the electrodes. The
salt bridge plays a very important role in a galvanic cell:
An electrolyte solution consists of metal cations and spectator anions.
\(\text{NaCl}(\text{aq})\) is \(\text{Na}^{+}(\text{aq})\) and
\(\text{Cl}^{-}(\text{aq})\) in the paste.
There is a build up of positive charge in the \(\color{blue}{\textbf{anode
half-cell compartment}}\) as solid metal is
\(\color{blue}{\textbf{oxidised}}\) and the positive ions move into
solution. So there are more postive metal ions in the
electrolyte than negative ions.
\(\color{blue}{\textbf{Zn(s)} \to \textbf{Zn}^{2+}\textbf{(aq)}}\), while the
number of \(\text{SO}_{4}^{2-}\) ions remains the same.
To balance the charge, negative ions from the salt bridge move into the anode
half-cell compartment.
\(\text{Cl}^{-}\) ions move from the salt bridge to the anode half-cell
compartment.
There is a decrease in positive charge in the \(\color{red}{\textbf{cathode
half-cell compartment}}\) as metal ions are
\(\color{red}{\textbf{reduced}}\) and form solid metal. So there are
more negative ions in the electrolyte than positive
metal ions.
\(\color{red}{\textbf{Cu}^{2+}\text{(aq)} \to \textbf{Cu(s)}}\), while the
number of \(\text{SO}_{4}^{2-}\) ions remains the same.
To balance the charge, positive ions from the salt bridge move into the
cathode half-cell compartment.
\(\text{Na}^{+}\) ions move from the salt bridge to the cathode half-cell
compartment.
The salt bridge acts as a transfer medium that allows ions to flow through without allowing
the different solutions to mix and react directly. It allows a balancing of the charges
in the electrolyte solutions, and allows the reactions in the cell to continue.
Without the salt bridge, the flow of electrons in the outer circuit stops completely. This is
because the salt bridge is needed to complete the circuit.
Electrolytic cells (ESCR6)
In an electrolytic cell electrical potential energy is converted to chemical
potential energy. An electrolytic cell uses an electric current to force a particular
chemical reaction to occur, which would otherwise not take place.
Electrolytic cell
An electrolytic cell is an electrochemical cell that converts electrical
potential energy to chemical potential energy by using electricity to
drive a non-spontaneous chemical reaction.
Sometimes galvanic cells are just called electrochemical cells. While they are
electrochemical cells, electrolytic cells are also electrochemical cells.
Electrolytic and galvanic cells are not the same however.
An electrolytic cell is activated by applying an electrical potential across the electrodes
to force an internal chemical reaction between the electrodes and the ions that are in
the electrolyte solution. This process is called electrolysis.
Electrolysis
Electrolysis is a method of driving chemical reactions by passing an electric
current through an electrolyte.
In an electrolytic cell (for example the cell shown in Figure 13.6):
The electrolyte solution consists of the metal cations and spectator anions.
The oxidation and reduction reactions occur in the same container but are
non-spontaneous.They require the electrodes to be connected to an
external power source to proceed.
The electrodes in an electrolytic cell can be the same metal or different
metals. The prinicple is the same. Let there be only one metal, and let
that metal be Z.
An electrode is connected to the \(\color{blue}{\textbf{positive terminal}}\)
of the battery.
To balance the charge at the \(\color{blue}{\textbf{positive
electrode}}\) metal atoms are
\(\color{blue}{\textbf{oxidised}}\) to form metal ions.
The ions move into solution, leaving their electrons on
the electrode.
The following reaction takes place:
\(\color{blue}{\textbf{Z(s)} \to
{\textbf{Z}}^{+}{\textbf{(aq)}} + {\textbf{e}}^{-}}\)
\(\color{blue}{\textbf{Ox}}\)idation is loss at the
\(\color{blue}{\textbf{an}}\)ode, therefore this
electrode is the \(\color{blue}{\textbf{anode}}\).
An electrode is connected to the \(\color{red}{\textbf{negative terminal}}\)
of the battery.
When positive ions come in contact with the
\(\color{red}{\textbf{negative electrode}}\) the ions
gain electrons and are
\(\color{red}{\textbf{reduced}}\).
The following reaction takes place:
\(\color{red}{\textbf{Z}^{+}{\textbf{(aq)}} +
{\textbf{e}}^{-} \to {\textbf{Z(s)}}}\)
\(\color{red}{\textbf{Red}}\)uction is gain at the
\(\color{red}{\textbf{cat}}\)hode, therefore this
electrode is the \(\color{red}{\textbf{cathode}}\).
This means that the overall reaction is: \(\text{Z}(\text{s}) +
\text{Z}^{+}(\text{aq})\) \(\to\) \(\text{Z}^{+}(\text{aq}) +
\text{Z}(\text{s})\). While this might seem trivial this is an important
technique to purify metals (see
Section 13.7).
In the movement of coloured ions experiment, to make the ammonia and ammonium
chloride buffer solution you need to combine equimolar amounts of ammonia and
ammonium chloride. The volume doesn't matter, so long as there are the same
number of moles of each compound in the solution.
Concentrated, strong bases can cause serious burns. Please remind the learners to be
careful and wear the appropriate safety equipment when handling all chemicals,
especially strong, concentrated bases. The safety equipment includes gloves,
safety glasses, and protective clothing.
The movement of coloured ions under the effect of electrical charge
Aim
To demonstrate how ions migrate in solution towards oppositely charged
electrodes.
Apparatus
Filter paper, glass slide, a 9V battery, two crocodile clips
connected to wires, tape
Connect the wire from one crocodile clip to one end of the
battery and secure with tape. Repeat with the other
crocodile clip wire and the other end of the battery.
Soak a piece of filter paper in the ammonia and ammonium
chloride buffer solution and place it on the glass
slide.
Connect the filter paper to the battery using one of the
crocodile clips, keep the other one nearby.
Place a line of copper(II) chromate solution at the centre of
the filter paper. The colour of this solution is
initially green-brown.
Attach the other crocodile clip opposite the first one (as
shown in the diagram) and leave the experiment to run
for about \(\text{20}\) \(\text{minutes}\).
Results
After \(\text{20}\) \(\text{minutes}\) you should see that
the central coloured band disappears and is replaced by
two bands, one yellow and the other blue, which seem to
have separated out from the first band of copper(II)
chromate.
The cell that is used to supply an electric current sets up a
potential difference across the circuit, so that one of
the electrodes is positive and the other is negative.
The chromate (\(\text{CrO}_{4}^{2-}\)) ions in the copper(II)
chromate solution are attracted to the positive
electrode, this creates a yellow band. The
\(\text{Cu}^{2+}\) ions are attracted to the negative
electrode, this creates a blue band.
Conclusion
The movement of ions occurs because the electric current in the external
circuit provides a potential difference between the two electrodes.
In the electrolytic cell experiment it is important that the learners weigh the
copper electrodes carefully before the experiment, and wash and dry the
electrodes before weighing after the experiment.
An electrolytic cell
Aim
To investigate the reactions that take place in an electrolytic cell.
Apparatus
Two copper plates (of equal size and mass), copper(II)
sulfate (\(\text{CuSO}_{4}\)) solution (\(\text{1}\)
\(\text{mol.dm$^{-3}$}\))
A \(\text{9}\) \(\text{V}\) battery, two connecting wires, a
beaker.
Method
Half fill the beaker with copper(II) sulfate solution. What
colour is the solution?
Weigh each copper electrode carefully and record the weight.
Place the two copper electrodes (of known mass) in the
solution and make sure they are not touching each other.
Connect the electrodes to the battery as shown below and
leave the experiment for a day. What colour is the
solution after a day?
Discussion
What colour was the copper(II) sulfate solution before the
experiment?
What colour was the copper(II) sulfate solution after the
experiment?
Examine the two electrodes, what do you observe?
What is the charge on each electrode?
Which electrode is the anode and which is the cathode?
Observations
The initial blue colour of the solution remains unchanged
throughout the experiment.
It appears that copper has been deposited on one of
the electrodes (it increased in mass) but
dissolved from the other (it decreased in
mass).
The electrode connected to the negative terminal of the
battery will have a negative charge. The electrode
connected to the positive terminal of the battery will
have a positive charge.
When positively charged \(\text{Cu}^{2+}\) ions encounter the
negatively charged electrode they gain electrons and are
reduced to form copper metal. This metal is deposited on
the electrode. The half-reaction that takes place is as
follows:
Reduction occurs at the
cathode. Therefore, the electrode which
increased in mass is the cathode.
At the positive electrode, copper metal is oxidised to form
\(\text{Cu}^{2+}\) ions, leaving electrons on the
electrode. The half-reaction that takes place is as
follows:
Oxidation occurs at the
anode. Therefore, the electrode which
decreased in mass is the anode.
The amount of copper that is deposited at one
electrode is approximately the same as the amount of
copper that is dissolved from the other. The
number of \(\text{Cu}^{2+}\) ions in the solution
therefore remains almost the same, and the blue colour
of the solution is unchanged.
Conclusion
In this demonstration, the container held aqueous \(\text{CuSO}_{4}\)
(\(\text{Cu}^{2+}(\text{aq})\) and \(\text{SO}_{4}^{2-}(\text{aq})\)).
The copper atoms of the electrode connected to the positive terminal
(the anode) were oxidised and formed \(\text{Cu}^{2+}(\text{aq})\) ions,
causing a decrease in mass. The copper atoms of the electrode connected
to the negative terminal (the cathode) were reduced to form solid
copper, causing an increase in mass. This process is called
electrolysis, and is very useful in the purification of metals.
Note that the cathode is negative and the anode is positive. Reduction still
occurs at the cathode (Red Cat), and oxidation still occurs at the anode
(An Ox).
The electrolysis of water
August Wilhelm von Hofmann was a German chemist who invented the Hofmann cell, which
uses a current to form \(\text{H}_{2}(\text{g})\) and \(\text{O}_{2}(\text{g})\)
from water through electrolysis.
Water can undergo electrolysis to form hydrogen gas and oxygen gas according to the following
reaction:
This reaction is very important because hydrogen gas has the potential to be used as an
energy source. The electrolytic cell for this reaction consists of two electrodes,
submerged in an electrolyte and connected to a source of electric current (Figure 13.7).
The \(\color{blue}{\textbf{oxidation half-reaction}}\) is as follows: The
\(\color{red}{\textbf{reduction half-reaction}}\) is as follows:
The informal experiment is the electrolysis of sodium iodide and water. It would be
advisable to check that current flows through the pencils, as the pencil lead
may be broken. If you have access to graphite rods they can be used instead. In
the second part of the experiment, there should be an obvious colour change when
the phenolphthalein and \(\text{NaOH}\) mix at the cathode.
Learners are required to work with a concentrated, strong acid. Concentrated, strong
acids can cause serious burns. Please remind the learners to be careful and wear
the appropriate safety equipment when handling all chemicals, especially
concentrated acids. The safety equipment includes gloves, safety glasses and
protective clothing.
temp text
The electrolysis of sodium iodide and water
Aim
To study the electrolysis of water and sodium iodide.
Apparatus
\(\text{4}\) pencils, copper wire attached to crocodile
clips, \(\text{9}\) \(\text{V}\) battery, \(\text{2}\)
beakers, pencil sharpener, spatula, glass rod
Label one beaker 1 and half fill it with
distilled \(\text{H}_{2}\text{O}\).
Sharpen both ends of two of the pencils. Strip away some of
the wood to expose more of the graphite. (Graphite rods
can be used instead of pencils if they are available).
Attach one end of the crocodile clips to the pencils and the
other end to the battery.
Place the pencils in the beaker, making sure they are not
touching each other. Observe what happens.
Pour approximately \(\text{5}\) \(\text{cm$^{3}$}\) of the
\(\text{H}_{2}\text{SO}_{4}\) solution into the beaker.
Observe what happens.
Label a second beaker 2 and add \(\text{5}\)
spatula tips of \(\text{NaI}\) to the beaker.
Half fill beaker \(\text{2}\) with distilled
\(\text{H}_{2}\text{O}\) and stir with the glass rod
until all the sodium iodide has dissolved.
Repeat steps \(\text{2}\) - \(\text{4}\) with the second set
of pencils.
After a few minutes add \(\text{3}\) - \(\text{4}\) drops of
the phenolphthalein to the beaker. Observe what happens.
Questions
In beaker 1:
What happened when the pencils were first put
in the water?
What happened when the sulfuric acid was
added to the beaker?
Why is the sulfuric acid necessary in this
reaction?
What is happening at the negative electrode
(the pencil attached to the negative
terminal of the battery)?
What is happening at the positive electrode
(the pencil attached to the positive
terminal of the battery)?
Which electrode is the anode and which is the
cathode?
In beaker 2:
What happened when the pencils were first put
in the water?
What is happening at the negative electrode?
What is happening at the positive eletrode?
Which electrode is the anode and which is the
cathode?
What happened when you added the
phenolphthalein? Why did the change take
place?
Results
In the electrolysis of water two \(\text{H}^{+}\) ions each
gain an electron (are reduced) and combine to form
hydrogen gas (\(\text{H}_{2}(\text{g})\)):
When the positive \(\text{H}^{+}\) ions encounter the
negative electrode they are reduced. Reduction is a gain
of electrons at the cathode, therefore the negative
electrode is the cathode.
When the negative \(\text{O}^{2-}\) ions encounter the
positive electrode they are oxidised. Oxidation is a
loss of electrons at the anode, therefore the positive
electrode is the anode.
Remember that pure water does not conduct electricity. So, an
electrolyte (such as sulfuric acid) is necessary for the
reaction to take place.
A salt dissolved in the water is also an electrolyte, so
sulfuric acid is not necessary in the second beaker.
The electrolysis of water occurs in beaker
2, but there are other reactions taking
place as well because of the presence of the
\(\text{Na}^{+}\) and \(\text{I}^{-}\) ions.
When the negative \(\text{I}^{-}\) ions encounter the
positive electrode they are oxidised:
Remember that phenolphthalein turns pink in the presence of a
base. So when the phenolphthalein is added, the water
around the negative electrode should become pink due to
the \(\text{NaOH}\).
Conclusion
The application of an electric current to water splits the water molecules
and causes hydrogen and oxygen gas to form. This can only happen in the
presence of an electrolyte (for example sulfuric acid). Sodium iodide
dissolved in the water is also an electrolyte and enables the
electrolysis of water. However, the cation and anion of the salt will
also undergo a reaction. \(\text{NaOH}\) will form at the cathode, while
solid iodine will form at the anode.
Galvanic and electrolytic cells
Textbook Exercise 13.4
An electrolytic cell consists of two electrodes in a silver
chloride (\(\text{AgCl}\)) solution, connected to a
source of current. A current is passed through the
solution and \(\text{Ag}^{+}\) ions are reduced to a
silver metal deposit on one of the electrodes.
What is the name of this process?
Electrolysis
Does reduction occur at the electrode where the
deposit formed?
Yes. Reduction is a gain of electrons. In an
electrolytic cell the electrode connected to the
negative terminal of the battery is negative.
When the metal cations encounter this electrode
they gain electrons and form metal atoms that
deposit on the electrode
(\(\text{Ag}^{+}(\text{aq}) +\text{e}^{-}\)
\(\to\) \(\text{Ag}(\text{s})\)).
Give the equation for the reduction half-reaction.
A galvanic cell consists of two half-cells: a copper anode in
a copper nitrate
(\(\text{Cu}(\text{NO}_{3})_{2}(\text{aq})\)) solution,
and a silver cathode in a silver nitrate
(\(\text{AgNO}_{3}(\text{aq})\)) solution.
Give equations for the half-reactions that take place
at the anode and cathode.
You are told that the copper electrode is the anode,
oxidation is a loss of electrons at the anode: