6.2
To investigate the effect of the positions of ions in the electrochemical series on the
selective discharge of ions and the products of electrolysis of aqueous solutions
Problem statement Procedure
How do the positions of ions in the electrochemical 1 Aqueous copper(II) sulphate solution is put into
series determine the types of ions selectively an electrolytic cell with carbon electrodes.
discharged during electrolysis?
2 Two test tubes, filled with copper(II) sulphate
Hypothesis solution are inverted over the carbon anode and
cathode respectively (Figure 6.5).
If an aqueous solution consists of more than one
Experiment 6.2 type of ion, the lower the position of the ion in the
6 electrochemical series, the higher the tendency it is
for the ion to be discharged.
Variables
(a) Manipulated variable : Position of ions in the
electrochemical series
(b) Responding variable : Types of ions discharged at
the anode and the cathode
(c) Constant variable : Concentration of electro
lytes, types of electrodes,
duration of electrolysis
Apparatus Figure 6.5 Electrolysis of copper(II) sulphate
Batteries, electrolytic cell, carbon electrodes, ammeter, solution
switch, connecting wires with crocodile clips and
test tubes. 3 The switch is turned on and electric current is
allowed to flow for 15 minutes.
Materials
Aqueous 0.5 mol dm–3 copper(II) sulphate, CuSO4 4 Steps 1 to 3 of the experiment are repeated by
solution, 0.5 mol dm–3 dilute sulphuric acid, H2SO4 replacing copper(II) sulphate solution with dilute
and 0.5 mol dm–3 sodium nitrate, NaNO3 solution. sulphuric acid and sodium nitrate solution in
turn.
Results
Electrolyte Observation Inference
Copper(II) sulphate At the cathode
solution • Brown deposit is formed Copper metal is deposited
At the anode
• Gas bubbles are formed Oxygen gas is produced
• Gas produced lights up a glowing wooden splint
Colour of electrolyte
• The blue colour of the solution becomes paler Concentration of Cu2+ ion decreases
Dilute sulphuric At the cathode
acid and sodium
nitrate solution • Gas bubbles are formed Hydrogen gas is produced
• When a lighted wooden splint is placed near the
mouth of the test tube, a ‘pop’ sound is produced
At the anode Oxygen gas is produced
• Gas bubbles are formed
• Gas lights up a glowing wooden splint
Electrochemistry 144
Discussion 8 The ions present in the sodium nitrate solution
are Na+ ions, NO3– ions, H+ ions and OH– ions.
1 The ions present in the aqueous copper(II)
sulphate solution are Cu2+ ions, SO42– ions, H+ NaNO3 → Na+ + NO3–
ions and OH– ions.
H2O H+ + OH–
CuSO4 → Cu2+ + SO42– 9 Both types of cations, H+ ions and Na+ ions are
attracted to the cathode. H+ ions are selectively
H2O H+ + OH– discharged at the cathode because the position
of the H+ ion is lower than that of the Na+ ion
2 Both types of cations, Cu2+ ions and H+ ions are in the electrochemical series. Hydrogen ions are 6
attracted to the cathode. Cu2+ ions are selectively discharged to form hydrogen gas.
discharged at the cathode because the position
of the Cu2+ ion is lower than that of the H+ Half-equation at the cathode:
ion in the electrochemical series. A Cu2+ ion is
discharged by accepting 2 electrons to form a 2H+ + 2e– → H2
copper atom at the cathode.
10 Both types of anions, NO3– and OH– ions are
Half-equation at the cathode: attracted to the anode. OH– ions are selectively
discharged at the anode because of the position
Cu2+ + 2e– → Cu of the OH– ion in the electrochemical series.
Hydroxide ions are discharged to form water and
3 Both types otfheanainoonds,eS. OO4H2––ioionnssanadreOdHis–cihoanrsgaerde oxygen.
attracted to
Half-equation at the anode:
at the anode because the position of the OH– ion is
4OH– → 2H2O + O2 + 4e–
lower than that of itohnesSaOre42d–iisocnhainrgtehdebeyledcotrnoactihneg
mical series. 4OH– 11 Electrolysis of dilute sulphuric acid and sodium
nitrate solution is actually electrolysis of water
4 electrons to form water and oxygen. because H+ ions and OH– ions are discharged.
The decrease in the concentrations of H+ ions
Half-equation at the anode: and OH– ions in the solution results in the
increase of sulphuric acid and sodium nitrate
4OH– → 2H2O + O2 + 4e– concentration during electrolysis.
4 The blue colour of the copper sulphate solution 1 2 The ratio of H2 gas to O2 gas produced is 2 : 1.
This is because the release of 4 electrons in the
is due to the presence of copper(II) ion, Cu2+. formation of 1 molecule of O2 results in the
formation of 2 molecules of H2 when these 4
The blue colour of the electrolyte becomes paler electrons are accepted by 4 H+ ions.
during electrolysis because the Cu2+ ion concen
tration decreases when Cu2+ ions are discharged. Conclusion
5 The ions present in the dilute sulphuric acid are
SO42– ions, H+ ions and OH– ions. 1 Electrolysis of aqueous copper(II) sulphate
solution using carbon electrodes produces copper
H2SO4 → 2H+ + SO42– at the cathode and oxygen gas at the anode.
H2O H+ + OH– 2 Cu2+ ions and OH– ions are selectively
discharged because of their lower positions in
6 Hydrogen ions are attracted to the cathode. H+ the electrochemical series.
ion is discharged by receiving one electron to
form a hydrogen atom. Two hydrogen atoms will 3 Electrolysis of dilute sulphuric acid and sodium
combine to form a hydrogen gas molecule. nitrate solution using carbon electrodes produces
hydrogen gas at the cathode and oxygen gas at
Half-equation at the cathode: the anode.
2H+ + 2e– → H2 4 H+ ions and OH– ions are selectively discharged
because their positions are lower in the
7 Both types of anions, SO42– ions and OH– ions are electrochemical series.
attracted to the anode. OH– ions are selectively
discharged at the anode because of the position The hypothesis is accepted.
of the OH– ion in the electrochemical series. Four
OH– ions are discharged to form water and oxygen.
Half-equation at the anode:
4OH– → 2H2O + O2 + 4e–
145 Electrochemistry
6.3 SPM
’09/P2
To investigate the effect of the concentration of ions on the selective discharge of
ions and the products of electrolysis of aqueous solutions
Problem statement
How does the concentration of ions determine the
types of ions discharged during electrolysis?
Hypothesis
Ions of higher concentration will be selectively
discharged during electrolysis.
Experiment 6.3 Variables
6
(a) Manipulated variable: Concentration of ions in Figure 6.6 Electrolysis of copper(II) chloride
the solution solution
(b) Responding variable: Types of ions to be 2 Two test tubes, filled with copper(II) chloride
discharged at the anode and cathode solution are inverted over the carbon anode and
cathode respectively (Figure 6.6).
(c) Constant variables: Types of ions in the electrolyte,
types of electrodes, duration of electrolysis. 3 The switch is turned on and electric current is
allowed to flow for 15 minutes.
Apparatus
4 Any change in colour of the electrolyte and
Batteries, electrolytic cell, carbon electrodes, ammeter, any other changes that occur around the carbon
switch, connecting wires with crocodile clips and electrodes are recorded.
test tubes.
5 Steps 1 to 4 of the experiment are repeated
Materials using the dilute copper(II) chloride solution of
0.001 mol dm–3 to replace the concentrated
Aqueous 2.0 mol dm–3 copper(II) cdhmlo–3ridceo,pCpeurC(I1I)2 copper(II) chloride solution.
solution and aqueous 0.001 mol
chloride solution.
Procedure
1 Concentrated aqueous copper(II) chloride
solution of 2.0 mol dm–3 is put into an electrolytic
cell with carbon electrodes.
Results
Electrolyte Observation Inference
Concentrated Copper metal is produced
copper(II) At the cathode: Chlorine gas is produced
chloride solution Brown deposit is formed
of 2.0 mol dm–3 Concentration of Cu2+ ions in
At the anode: copper(II) chloride solution decreases
Dilute copper(II) Bubbles of pungent greenish-yellow gas are Copper metal is produced
chloride solution produced. The gas turns the damp blue litmus Oxygen gas is produced
of 0.001 mol dm–3 paper to red and then bleaches it Concentration of Cu2+ ions in
copper(II) chloride solution decreases
Colour of electrolyte:
The blue colour of the solution becomes paler
At the cathode:
Brown deposit is formed
At the anode:
Bubbles of colourless gas are produced.
The gas lights up a glowing wooden splint
Colour of electrolyte:
The blue colour of the solution becomes paler
Electrochemistry 146
Discussion Cu2+ ions are discharged because the position
of Cu2+ ions is lower than that of H+ ions
1 Aqueous copper(II) chloride consists of Cu2+ in the electrochemical series. Hence copper
ions, H+ ions, OH– ions and Cl– ions. metal is deposited.
CuCl2 → Cu2+ + 2Cl– Half-equation at the anode:
H2O H+ + OH– Cu2+ + 2e– → Cu
2 Electrolysis of concentrated aqueous copper(II) (b) The blue colour of the solution becomes paler Experiment 6.4 6
because the concentration of the Cu2+ ions
chloride solution (2 mol dm–3): decreases when Cu2+ ions are discharged at
the cathode during electrolysis.
At the anode: 2 types of anions, Cl– ions and
Conclusion
OH– ions are attracted to the anode. Cl– ions are
discharged because the concentration of Cl– 1 In the electrolysis of concentrated aqueous
ions is higher than that of OH– ions. copper(II) chloride solution, copper metal
is produced at the cathode and chlorine gas is
Half-equation at the anode: produced at the anode. At the anode, the Cl– ions
are selectively discharg ed, producing chlorine
2Cl– → Cl2 + 2e– gas because the concentration of Cl– ions is
higher than that of OH– ions.
Hence chlorine gas is produced.
3 Electrolysis of dilute aqueous copper(II) chloride 2 In the electrolysis of dilute aqueous copper(II)
chloride solution, copper metal is produced at
solution (0.001 mol dm–3): the cathode and oxygen gas is produced at the
At the anode: 2 types of anions, Cl– ions and anode. At the anode, OH– ions are selectively
OH– ions are attracted to the anode. OH– ions discharged, producing oxygen gas because the
are discharged because the concentration of Cl– concentration of Cl– ions is low.
ions is lower than that of OH– ions.
3 The type of ions that is selectively discharged at
Half-equation at the anode: the electrode is determined by the concentration
of the ions.
4OH– → 2H2O + O2 + 4e–
The hypothesis is accepted.
Hence oxygen gas is produced.
4 (a) At the cathodes of both concentrated and
dilute aqueous copper(II) chloride solution:
6.4 SPM
’09/P3
To investigate the effect of the types of electrodes on the selective discharge of ions
and the products of electrolysis of aqueous solutions
Problem statement (c) Constant variables : Types of ions in the
electrolyte and the concen
How do the types of electrodes determine the types tration of ions
of ions discharged during electrolysis?
Hypothesis Apparatus
The products of electrolysis of copper(II) sulphate Batteries, electrolytic cell, carbon electrodes, copper
solution with copper electrodes are different from electrodes, ammeter, switch, rheostat, connecting
that with carbon electrodes. wires with crocodile clips and test tubes.
Variables Materials
Aqueous 1.0 mol dm–3 copper(II) sulphate, CuSO4
(a) Manipulated variable : Types of electrodes solution
(b) Responding variable : Products of electrolysis
147 Electrochemistry
Procedure Figure 6.7 Electrolysis of copper(II)
sulphate solution
1 Aqueous 1.0 mol dm–3 copper(II) sulphate solution
is put into an electrolytic cell with carbon electrodes.
2 A test tube filled with copper(II) sulphate solution
is inverted over the carbon anode (Figure 6.7).
3 The switch is turned on and the electric current is
allowed to flow for 15 minutes.
4 Any change in colour of the electrolyte and
any other changes that occur around the carbon
electrodes are recorded.
5 Steps 1 to 4 of the experiment are repeated using
copper electrodes to replace carbon electrodes.
6 Results
Type of electrodes Observation Inference
Carbon
At the cathode: Copper metal is produced
Copper Brown deposit is formed
Oxygen gas is produced
At the anode:
Bubbles of colourless gas are produced Concentration of Cu2+ ions
The gas lights up a glowing wooden splint decreases
Colour of electrolyte: Copper metal is produced
The blue colour of the solution becomes paler Copper anode dissolves to form
Cu2+ ions
At the cathode: Concentration of Cu2+ ions in
Formation of brown deposit makes the cathode thicker copper(II) sulphate remains
constant
At the anode:
Anode corrodes and becomes thinner
Colour of electrolyte:
The blue colour of the solution remains
unchanged
Discussion 4 The colour intensity of the blue solution
decreases because the concentration of Cu2+
1 Aqueous copper(II) sulphate solution consists of ions in the copper(II) sulphate decreases.
Cu2+ ions, H+ ions, SO42– ions and OH– ions. 5 During electrolysis, the concentration of OH–
2 During electrolysis, OH– ions elaencdtroSdOe42i–s ions ions decreases, leaving H+ ions behind. As a
move to the anode. If a carbon used result, the solution becomes acidic.
as the anode, OH– ion is selectively discharged 6 If copper is used as the anode, both SO42– ions
and OH– ions are not discharged. Instead, the
due to its position in the electrochemical series. copper anode dissolves by releasing electrons
to form Cu2+ ions. Hence, the mass of the anode
Oxygen gas is produced. decreases and the anode becomes thinner. The
types of electrodes (copper electrode is an active
Half-equation at the anode: electrode) determine the product formed at the
anode during electrolysis.
4OH– → 2H2O + O2 + 4e–
Half-equation at the anode:
3 Cu2+ ions and H+ ions move to the cathode. Cu2+
ion which is at a lower position than the H+ ion Cu → Cu2+ + 2e–
in the electrochemical series will be discharged.
Copper metal is produced. 7 Cu2+ ions are still discharged at the cathode,
producing copper metal. This causes the mass
Half-equation at the cathode:
Cu2+ + 2e– → Cu
Electrochemistry 148
of the cathode to increase and the cathode Conclusion
becomes thicker.
Half-equation at the cathode: 1 In the electrolysis of aqueous copper(II) sulphate
Cu2+ + 2e– → Cu solution:
8 The colour intensity of the blue solution does (a) If a carbon electrode is used as the anode,
not change because the concentration of Cu2+
ions is constant. The number of moles of Cu2+ OH– ions are discharged and oxygen gas is
ions discharged at the cathode is the same as the
number of moles of Cu2+ ions produced at the produced.
anode.
(b) If a copper electrode is used as the anode, both
9 The mass of the electrodes can be weighed before
and after electrolysis using an electronic balance. OInHst–eaidonthseacnodpSpeOr4a2–nioodnes are not discharged.
dissolves to produce
1 0 The decrease in the mass of the copper anode
is the same as the increase in the mass of the Cu2+ ions.
copper cathode.
(c) Cu2+ ions are discharged at the cathode producing
copper metal whether the cathode used is a
carbon electrode or a copper electrode. 6
2 The types of electrodes used during electrolysis
determine the types of ions discharged and the
products of electrolysis.
The hypothesis is accepted.
How to Predict the Products of Electrolysis of Aqueous Solutions ’08
1
The diagram shows the arrangement of apparatus for the electrolysis of silver nitrate solution.
Write the half-equation representing the reaction at electrode Q.
Comments
The ions present in the electrolyte are Ag+, NO3–, H+ and OH–.
Electrode Q is the anode as it is connected to the positive terminal of the battery. OH– ions are discharged at the
anode producing oxygen gas. (Factor: positions of ions in the electrochemical series)
4OH– → 2H2O + O2 + 4e–
The cations at a higher position in the electrochemical series are very stable. These ions are unlikely to accept electrons
to form neutral atoms. Hence K+ ions and Na+ ions are never discharged in an aqueous solution in electrolysis. The
cations at the lower position of the electrochemical series are less stable. They are more likely to accept electrons to
form neutral atoms.
Similarly with anions. Anions at a higher position in the electrochemical series are never discharged in an aqueous
solution in electrolysis. Hence F–, –S, OI–42o–r aOnHd–NioOn3s– ions are stable compared to the lower anions at a lower position in
the electrochemical series. Cl–, Br will be discharged instead depending on their ionic concentration in
the aqueous solution.
149 Electrochemistry
How to predict the products of electrolysis of aqueous solutions. SPM
’10/P2, ’11/P2
Step 1 Step 2
Identify the cations and anions that are Identify the anode and cathode
present in the aqueous solution. • Anode is the electrode connected to
Generally, an aqueous solution of MaXb
will produce ions as follows: the positive terminal of the battery.
Positive electrode attracts negative ions
Cation Anion (anions).
From MaXb : M b+ , X a– • Cathode is the electrode connected to
From H2O : H+ , OH– the negative terminal of the battery.
Negative electrode attracts positive ions
6 (cations).
Step 4 Step 3
Identify the ions to be discharged at the Identify the movements of ions
cathode • Mb+ ions and H+ ions move to the
H+ ions are discharged at the cathode
(producing H2 gas) except if Mb+ is Cu2+ cathode.
or Ag+ (because these ions are lower than • Xa– ions and OH– ions move to the
H+ ions in the electrochemical series).
• H+ ions discharge by accepting anode.
electrons to form H2 gas. Step 5
2H+ + 2e– → H2 Identify the types of electrode as anode
Inert electrode: carbon or platinum.
• Cu2+ ions or Ag+ ions discharge by Active electrode: Ag or Cu.
accepting electrons to form metal atom.
Cu2+ + 2e– → Cu
Ag+ + e – → Ag
Inert electrode SPM Active electrode
’09/P1 Anions are not discharged. Instead
anode dissolves to form ions.
OH– ions are discharged (producing O2 gas) except Examples:
if Xb– ions are Cl–/Br–/I– of high concentration.
• OH– ions discharge by donating electrons to Ag → Ag+ + e–
Cu → Cu2+ + 2e–
form water and O2 gas.
4OH– → O2 + 2H2O + 4e–
• Cl–/Br –/I– of high concentration discharge by
donating electrons to form halogen.
Example: 2Cl– → Cl2 + 2e–.
Electrochemistry 150
Electrolysis of dilute acids and alkalis Electrolysis of aqueous copper(II) nitrate 6
solution, Cu(NO3)2 with copper electrodes
1 Electrolysis of all dilute acids (HCl/H2SO4/
HNO3) and dilute alkali solutions (NaOH/ 1 At the cathode: Cu2+ ions are discharged.
KOH) is actually the electrolysis of water. Copper metal is deposited. Mass of cathode
will increase. (Factor: positions of ions in
2 At the cathode: H+ ions are discharged the electrochemical series).
producing hydrogen gas.
Cu2+ + 2e– → Cu
2H+ + 2e– → H2
2 At the anode: Copper dissolves from the
3 At the anode: OH– ions are discharged anode forming Cu2+ ions. Mass of the anode
producing oxygen gas. will decrease. (Factor: types of electrodes).
4OH– → 2H2O + O2 + 4e– Cu → Cu2+ + 2e–
4 The removal of H+ ions and OH– ions from Electrolysis of aqueous silver nitrate solution,
the solution during electrolysis causes the AgNO3 with carbon electrodes
concentration of the acid or alkali to increase.
1 At the cathode: Ag+ ions are discharged.
Electrolysis of concentrated aqueous sodium Silver metal is deposited. (Factor: positions
chloride solution, NaCl with carbon electrodes of ions in the electrochemical series).
1 Cations present : Na+ ions, H+ ions SPM Ag+ + e– → Ag
Anions present : Cl– ions, OH– ions
’08/P2 2 At the anode: OH– ions are discharged
producing oxygen gas. (Factor: positions of
2 At the cathode: H+ ions are discharged ions in the electrochemical series).
producing hydrogen gas. (Factor: positions 4OH– → 2H2O + O2 + 4e–
of ions in the electrochemical series). Electrolysis of aqueous silver nitrate solution,
AgNO3 with silver electrodes
2H+ + 2e– → H2
1 At the cathode: Ag+ ions are discharged.
3 At the anode: Cl– ions are discharged producing Silver metal is deposited. Mass of silver
chlorine gas. (Factor: concentration of ions). cathode will increase. (Factor: positions of
ions in the electrochemical series).
2Cl– → Cl2 + 2e–
Ag+ + e– → Ag
Electrolysis of aqueous copper(II) nitrate
solution, Cu(NO3)2 with carbon electrodes 2 At the anode: silver anode dissolves forming
Ag+ ions. Mass of silver anode electrode will
1 Cations present : Cu2+ ions, H+ ions decrease. (Factor: types of electrodes).
Anions present : NO3– ions, OH– ions
Ag → Ag+ + e–
2 At the cathode: Cu2+ ions are discharged.
Copper metal is deposited. (Factor: positions
of ions in the electrochemical series).
Cu2+ + 2e– → Cu
3 At the anode: OH– ions are discharged
producing oxygen gas. (Factor: positions of
ions in the electrochemical series).
4OH– → 2H2O + O2 + 4e–
151 Electrochemistry
6.3 such as sodium, calcium, magnesium
and aluminium are extracted from their
1 Write the formulae of the ions present in the compounds using electrolysis.
aqueous solutions below. Identify the ions that will 2 Electrolysis is used because these reactive
be discharged at the anode and the cathode during metals cannot be extracted from the minerals
electrolysis using carbon electrodes in every case. by reduction using carbon.
3 Examples are:
Aqueous Ions Ions Ions (a) Extraction of aluminium from aluminium
solution present discharged discharged
in the oxide, Al2O3.
solution at the at the (b) Extraction of sodium from sodium
cathode anode
chloride, NaCl.
6 (a) Aqueous
nitric acid Extraction of aluminium metal from the
mineral bauxite (Hall-Heroult process)
(b) Silver
nitrate 1 Bauxite is the major ore of aluminium
consisting of aluminium oxide, Al2O3.
(c) Very dilute
copper(II) 2 oCxriydoelittoelo(Nwear3AitlsFm6)eilstinagddpeodinttofroamlum20in00iu°mC
chloride to about 950°C.
2 3 Molten aluminium oxide is electrolysed using
carbon as electrodes in the electrolytic cell
as shown in Figure 6.8.
The figure above shows the arrangement of Figure 6.8 Extraction of aluminium metal from
apparatus in an electrolysis experiment. aluminium oxide.
(a) Name all the ions present in Cell I and Cell II.
(b) Electrolysis is carried out for 20 minutes. 4 Molten aluminium oxide dissociates into
(i) Predict the observations at electrodes P aluminium and oxide ions as follows:
and Q. Al2O3(l) → 2Al3+(l) + 3O2–(l)
(ii) What will be the change in colour in Cell I?
(a) At the cathode: Al3+ ions discharge to
Explain your answer. form aluminium metal
(iii) Predict the observations at electrodes R
Al3+ + 3e– → Al
and S.
(iv) Write the half-equations at electrodes P and R. (b) At the anode: O2– ions discharge to
form oxygen gas
(c) What is the factor that determines the formation
of the products in 2O2– → O2 + 4e–
(i) electrode Q? (ii) electrode R? Overall equation of electrolysis:
6.4 Electrolysis in Industries 2Al2O3 → 4Al + 3O2
Uses of Electrolysis in Industries SPM 5 The carbon anode is required to be replaced
from time to time because the oxygen gas
’09/P1 generated oxidises the carbon anode to
form carbon dioxide.
1 Electrolysis process is used widely in industries.
2 Some common industrial applications of
electrolysis are
(a) extraction of reactive metals
(b) purification of metals
(c) electroplating of metals
(A) Extraction of Reactive Metals Using Electrolysis
1 Metals that are very reactive (placed at the
top positions of the electrochemical series)
Electrochemistry 152
Extraction of sodium metal from sodium chloride (Downs process)
1 Sodium chloride is the most abundant and 4 Molten sodium chloride dissociates into
cheapest sodium compound. ions as follows:
2 Electrolysis of molten sodium chloride is NaCl(l) → Na+(l) + Cl–(l)
carried out using iron as cathode and carbon
as anode as in Figure 6.9. (a) At the cathode: Na+ ions discharge
to form sodium metal. Sodium metal
chlorine is less dense and floats on top of the
gas electrolyte to be collected.
molten sodium sodium is Na+ + e– → Na Activity 6.2 6
chloride produced here
(b) At the anode: Cl– ions discharge to
cathode (iron) form chlorine gas. Chlorine gas is a
useful by-product.
anode (carbon)
2Cl– → Cl2 + 2e–
Figure 6.9 Extraction of sodium metal from
molten sodium chloride
3 Calcium chloride is added to lower the Overall chemical equation of the
melting point of sodium chloride. electrolysis:
2NaCl → 2Na + Cl2
(B) Purification of Metals Using Electrolysis the anode to the cathode. After electrolysis,
1 Impure metals containing impurities can be the mass of anode is reduced while that of the
purified using electrolysis as outlined below.
(a) The impure metal is used as the anode. cathode is increased.
(b) A piece of pure metal is used as the cathode. 3 For example, in the purification of impure
(c) The electrolyte is a solution containing
the ions of the metal to be purified. copper metal:
(a) Anode: impure copper
2 In the process of purification of a metal (b) Cathode: pure copper
using electrolysis, metal is transferred from (c) Electrolyte: copper(II) ion solutions such
as copper(II) sulphate
To investigate the purification of copper metal using SPM
electrolysis
’10/P1
Apparatus 2 A piece of impure copper plate is connected to
Batteries, electrolytic cell, beaker, connecting wires the positive terminal of the batteries. This plate
with crocodile clips, ammeter and rheostat. acts as the anode.
Materials 3 A piece of pure copper metal is connected to the
1 mol dm–3 copper(II) sulphate solution, pure copper negative terminal of the batteries. This plate acts
plate and impure copper plate. as the cathode.
Procedure 4 The circuit is completed using the connecting
1 About 200 cm3 of 1 mol dm–3 copper(II) sulphate wires, rheostat and ammeter. The two copper
plates are immersed in the copper(II) sulphate
solution is poured into a beaker. solution. The solution is electrolysed for 30
minutes (Figure 6.10).
153 Electrochemistry
Discussion
1 At the anode: Copper dissolves from the anode
by releasing electrons to form Cu2+ ions. The
mass of the anode decreases.
Half-equation: Cu → Cu2+ + 2e–
Figure 6.10 Purification of copper metal 2 At the cathode: Cu2+ ions are discharged by
receiving electrons to form copper atoms.
Results Copper metal is deposited on the surface of the
cathode. As a result, the copper cathode becomes
Observation thicker. The mass of the cathode increases.
6 At the anode:
• The copper anode
Inference Half-equation: Cu2+ + 2e– → Cu
becomes thinner
• Impurities are Copper anode 3 In this process, Cu2+ ions are transferred from the
dissolves to form Cu2+ anode to the cathode and are deposited as pure
deposited below the ions copper metal. Impurities that are collected below
anode the anode are known as anode mud.
Copper metal is
At the cathode: deposited at the 4 The colour intensity of the blue solution does not
• The copper cathode cathode change because the concentration of Cu2+ ions
remains constant throughout electrolysis. The
becomes thicker Concentration of Cu2+ rate of formation of Cu2+ ions at the anode is the
ions in the electrolyte same as the rate of discharge of Cu2+ ions at the
Colour of electrolyte: does not change cathode.
• Colour intensity of
Conclusion
the blue solution
does not change When copper(II) sulphate solution is electrolysed
using pure copper as the cathode and impure copper
as the anode, purification of copper takes place. Pure
copper is deposited at the cathode.
(C) Electroplating of Metals Using Electrolysis (a) The object to be electroplated is used as
the cathode.
1 Electroplating is a process carried out to coat
the surface of metal objects with a thin and (b) The anode is the electroplating metal.
even layer of another metal. (c) The electrolyte is a solution that contains
2 Two main aims of electroplating metals are the electroplating metal ions.
(a) to prevent corrosion. For example, iron 4 An even and lasting layer of metal can be
objects are plated with a thin layer of
chromium or nickel metal to protect the produced if
iron from rusting.
(b) to improve the appearance. For example, (a) the surface of the object to be electroplated
electroplating with gold, platinum and is first polished using sandpaper.
silver makes the surface of the objects
appear shiny and more attractive. (b) a low electric current is used so that
electroplating is carried out slowly during
3 In general, there are three conditions in the
SPM electroplating of metal. the electroplating process.
’08/P1 (c) the object to be plated is rotated steadily
during electrolysis.
Electrochemistry 154
2 SPM Comments
’10/P1 At the anode, the silver foil dissolves by releasing
electrons, thus becoming thinner:
The diagram shows the set-up of the apparatus
used to electroplate an iron key with silver. Ag → Ag+ + e–
At the cathode, silver ions discharge by receiving
electrons to form silver atoms, forming a shiny grey
deposit on the iron key:
Ag+ + e– → Ag
Answer B
What is observed at the anode and cathode after Experiment 6.5 6
30 minutes?
Anode Cathode Presently, plastic electroplating is carried out to coat a
thin layer of metal onto the surface of plastic objects.
A Shiny grey deposits Silver foil becomes The object produced will have the advantanges of
are formed thicker plastic: light, cheap, resistant to corrosion as well as
having a shiny surface like a metal. As plastic is an
B Silver foil becomes Shiny grey deposits electric insulator, a layer of graphite powder is coated
thinner are formed onto the surface of the plastic, so that it can conduct
electricity, before electroplating is carried out.
C Shiny grey deposits Gas bubbles are
are formed released
D Silver foil becomes Gas bubbles are
thinner released
6.5
To investigate the electroplating of an iron spoon with copper using electrolysis
Problem statement
How is electrolysis used to electroplate an iron spoon with copper metal?
Hypothesis
Electroplating of an iron spoon with copper occurs if the iron spoon is used as the cathode, copper metal is
used as the anode and aqueous copper(II) sulphate solution as the electrolyte.
Variables
(a) Manipulated variable : The position of the iron spoon as an electrode
(b) Responding variable : The deposition of copper on the iron spoon
(c) Constant variable : Type of electrolyte and arrangement of apparatus
Apparatus
Batteries, electrolytic cell, beaker, connecting wires with crocodile clips, ammeter and rheostat.
Materials
0.5 mol dm–3 copper(II) sulphate solution, copper plate and iron spoon.
Procedure
1 About 200 cm3 of 0.5 mol dm–3 copper(II) sulphate solution is poured into a beaker.
2 An iron spoon is polished using sandpaper and is connected to the negative terminal of the batteries. The
spoon acts as the cathode.
3 A piece of copper metal, as the anode, is connected to the positive terminal of the batteries.
155 Electrochemistry
4 The circuit is completed using the connecting Figure 6.11 Electroplating of an iron spoon
wires, rheostat and ammeter. The iron spoon and with copper
the copper metal are immersed in the copper(II)
sulphate solution. The solution is electrolysed for
30 minutes using a small current (0.5 A).
5 Steps 1 to 4 of the experiment are repeated by
interchanging the positions of the iron spoon and
copper metal, whereby the iron spoon is made
the anode and the copper metal is made the
cathode.
6 Set Observation Inference
Set 1: Iron spoon as At the cathode: The iron spoon is plated with
the cathode, copper A brown metal is deposited on the copper metal
metal as the anode surface of the iron spoon
At the anode: The copper anode dissolves to form
The copper anode becomes thinner Cu2+ ions
Colour of the electrolyte: Concentration of Cu2+ ions in the
Colour intensity of the blue solution electrolyte remains constant
does not change
Set 2: Copper metal At the cathode: Copper metal is deposited on the copper
as the cathode, iron The copper plate becomes thicker electrode
spoon as the anode
At the anode: Electroplating of copper on the iron
No noticeable change in the appearance spoon does not take place
of the iron spoon
Colour of the electrolyte: Concentration of Cu2+ ions in the
The blue colour of the solution becomes electrolyte decreases
paler
Discussion 3 The colour intensity of the blue solution does not
change because the concentration of Cu2+ ions
1 The brown metal deposited on the iron spoon is remains constant throughout the electroplating
copper metal. process.
(a) At the anode: copper dissolves from the
anode by releasing electrons to form Cu2+ 4 A slow electrolysis process using a small current
ions. will ensure that the layer of copper sticks firmly
to the surface of the iron spoon.
Half-equation: Cu → Cu2+ + 2e–
Conclusion
(b) At the cathode: Cu2+ ions are discharged
by accepting electrons to form copper atoms. 1 In electroplating an iron spoon with copper using
Copper metal is deposited on the surface of electrolysis, the iron spoon is made the cathode
the iron spoon. and a piece of copper metal is made the anode.
Half-equation: Cu2+ + 2e– → Cu 2 Copper metal is transferred from the copper
anode to the iron spoon and is deposited as a thin
2 In this process, Cu2+ ions are transferred from layer of copper metal.
the anode to the cathode (iron spoon) and are
deposited as a thin and even layer of copper metal. 3 Electroplating does not take place if the iron
spoon is made the anode.
The hypothesis is accepted.
Electrochemistry 156
Benefits and Harmful Effects of Electrolysis in Industries
1 Table 6.1 shows the advantages and disadvantages of using electrolysis.
Table 6.1 Advantages and disadvantages of using electrolysis.
Advantages of using electrolysis Disadvantages of using electrolysis
1 Electrolysis is an effective method of extracting 1 Electrolysis is a process that uses a large 6
reactive metals from their compounds. Some quantity of electricity. For example, the recycling
chemical substances such as chlorine and sodium of aluminium requires only 9% of the electrical
can be manufactured in large quantities using energy used to produce the same quantity of
electrolysis. aluminium in electrolysis.
2 In electroplating, the whole surface area of a 2 The problem of environmental pollution
metal such as iron is coated with a thin, even especially in the electroplating process.
and valuable metal (such as gold, platinum (a) In the electroplating of iron by chromium
and silver). This layer of metal protects iron and nickel, the waste chemicals contain
from being exposed to air and water so as chromium ions and nickel ions that can
to prevent corrosion. Besides that, the layer endanger human health as well as pollute
of metal also gives an attractive appearance. water sources.
Electroplating can also be performed using (b) In silver electroplating, potassium silver
polymers as the coating material. This has cyanide, KAg(CN)2 solution is sometimes
been used to coat new cars with paint. The used as an electrolyte. The waste chemical
advantage of this process is that it can be done of the electrolyte contains cyanide ions
in water thereby eliminating the use of volatile which are toxic.
organic solvents used in spray-paints. (c) Metal objects to be electroplated are
cleaned by acids to remove the layer
3 Electrolysis process is used to purify metals of metal oxide on the surface before
such as zinc, silver, nickel, copper, lead and electroplating. The used acid wastes will
aluminium. This process is also known as electro- pollute water in the drains, rivers and
refining. lakes, thus destroying aquatic life.
2 Steps taken to overcome the problems of (b) Waste chemicals in electrolyte from
electroplating are treated to remove the
electrolysis in industries toxic substances before being drained
(a) Recycling such as the recycling of aluminium
out as effluent.
cans is encouraged to reduce the use of
electrolysis in the extraction of aluminium.
6.4 (a) What is a suitable metal that can be used as
metal M?
1 State three main uses of the electrolysis process in
industries. (b) State two observations that will be obtained in
this experiment.
2
(c) Write the half-equations for the reactions that
A student carried out an experiment to electroplate take place at
an iron key with silver using the apparatus as shown
in the above figure. (i) the iron key
(ii) the metal M
(d) How will the concentration of silver nitrate
solution change after electrolysis? Explain your
answer.
(e) Explain how the student can ensure that an even
and lasting layer of silver metal stays on the
surface of the iron key.
157 Electrochemistry
6.5 Voltaic Cells series) will accept electrons and acts as the
positive terminal (cathode).
Simple Voltaic Cells (Chemical Cells) SPM 5 A continuous flow of electrons from the
’09/P2 negative terminal to the positive terminal of
Experiment 6.6 1 A simple voltaic cell can be made by the cell through the external circuit produces
6 immersing two different types of metals in an electric current.
an electrolyte and connecting the two metals 6 The flow of electric current can be detected
by the lighting up of a light bulb or the
by wires in the external circuit. deflection of a galvanometer needle.
2 In a simple voltaic cell, electrons flow from 7 Voltaic cells are also known as galvanic cells
or chemical cells.
one metal to another metal through the 8 The potential difference (voltage) of the cell
is the electromotive force (e.m.f.) that moves
connecting wires in the external circuit. electrons and can be measured by a voltmeter.
3 The more electropositive metal (metal that 9 The further the distance between the positions
is at a higher position in the electrochemical of two metals in the electrochemical series, the
bigger the voltage of the cell. For example, a
series) will release electrons and thus acts as
the negative terminal (anode) of the voltaic magnesium/copper cell will produce a higher
cell. voltage than a zinc/copper cell.
4 The less electropositive metal (metal that
is at a lower position in the electrochemical
6.6
To investigate the production of electricity from chemical reactions in a simple SPM
voltaic cell ’05/P2
Problem statement Procedure
How does a chemical reaction produce electrical 1 A piece of magnesium plate and a piece of
energy in a simple voltaic cell? copper plate are polished with sandpaper.
Hypothesis 2 Both pieces of the magnesium and copper plates
Electric current is produced when two different are immersed in 200 cm3 of aqueous sodium
metals connected by wires are immersed in an chloride solution in a beaker as shown in
electrolyte. Figure 6.12.
Variables 3 Both plates are connected by the connecting wire
(a) Manipulated variable : Pairs of different metals to a voltmeter.
(b) Responding variable : Deflection of a voltmeter
4 The experiment is repeated using two pieces of
needle by the electric copper plates as electrodes.
current produced
(c) Constant variable : Types of electrolyte and
arrangement of apparatus
Apparatus
Voltmeter, beaker, connecting wires with crocodile
clips and sandpaper.
Materials Figure 6.12
1 mol dm–3 sodium chloride solution, copper plates
and magnesium plate.
Electrochemistry 158
Results Observation Inference
Type of metal used as • Voltmeter needle deflects but the • Electric current is produced. The
electrodes deflection decreases after awhile voltage produced is not constant and
decreases rapidly
Magnesium metal and
copper metal • Magnesium dissolves to form Mg2+
ions
• Magnesium metal corrodes
• Hydrogen gas is produced
Two pieces of copper • Bubbles of colourless gas are 6
metal evolved around the copper metal • Electric current is not produced
• Voltmeter needle does not show a • No reaction occurs
deflection
• No noticeable change occurs at the
copper electrode
Discussion Figure 6.13 Movement of electrons and ions in
a simple Mg/Cu cell using sodium
1 The deflection of the voltmeter needle shows that chloride solution as the electrolyte
an electric current is produced. The decreasing
deflection indicates that the electric current 8 If copper(II) solution is used as the electrolyte,
decreases rapidly. Cu2+ ions will receive electrons and are
discharged because its position is lower than H+
2 Magnesium metal is more electropositive ions and Mg2+ ions in the electrochemical series.
than copper (at a higher position in the Copper metal is produced. The overall equation
electrochemical series). Hence, it has a higher of the cell will be
tendency to donate electrons than copper.
Mg + Cu2+ → Mg2+ + Cu
3 Magnesium atoms will donate electrons to form
magnesium ions, Mg2+ in the solution, hence Conclusion
magnesium metal corrodes. 1 An electric current is produced when a chemical
Half-equation: Mg → Mg2+ + 2e– reaction occurs in a simple voltaic cell consisting
of two different metals, connected by wires
4 Electrons accumulate at the surface of the externally and immersed in an electrolyte.
magnesium metal. This makes magnesium act 2 In a simple voltaic cell, chemical energy released
as the negative terminal (also known as the from chemical reactions is converted into
anode) of the cell. electrical energy.
3 No electric current will be produced if both
5 The electrons flow through the external circuit electrodes are of the same material because there
from the magnesium metal (negative terminal) to is no potential difference between them.
the copper metal (positive terminal or cathode The hypothesis is accepted.
of the cell) producing electricity.
6 When sodium chloride solution is used as the
electrolyte, H+ ions (from water), Na+ ions and
Mg2+ ions move towards the copper metal. H+
ions will accept electrons from the copper metal
and be discharged because its position is lower
than Na+ ions and Mg2+ ions in the electro
chemical series. Hydrogen gas is produced.
Half-equation: 2H+ + 2e– → H2
7 The overall chemical equation in the cell is:
Mg + 2H+ → H2 + Mg2+
159 Electrochemistry
3 ’04
The diagram shows the set-up of the Metal P Metal Q
apparatus of a simple chemical cell.
What are metals P and Q? A Iron Aluminium
B Copper Magnesium
C Zinc Iron
D Lead Magnesium
6 Comments
The diagram shows that electrons flow from metal P to
metal Q. Metal P must be more electropositive (higher
in position in the electrochemical series) than metal Q.
This is because the more electropositive metal will release
electrons to become the negative terminal of the chemical
cell.
Answer C
Different Types of Voltaic Cells SPM
Voltaic cells can be divided into two categories as ’10/P2
shown in Table 6.2.
Figure 6.14 Daniell cell using a salt bridge
Table 6.2 Two categories voltaic cells (b) A porous pot as shown in Figure 6.15.
Primary cells Secondary cells Figure 6.15 Daniell cell using a porous pot
4 A salt bridge contains inert ions or salt that
• Non-rechargeable • Rechargeable cells
cells (cells that can be does not react with the electrolyte. Examples
(cells that cannot be charged again) are sodium chloride, potassium chloride,
charged again) potassium nitrate, ammonium chloride and
dilute sulphuric acid.
• Examples • Examples 5 A simple salt bridge can be made by immersing
(a) Daniell cell (a) Lead-acid a piece of filter paper in sulphuric acid or in a
(b) Dry cell accumulator salt solution.
(c) Mercury cell (b) Nickel/cadmium 6 A porous pot has fine pores that allow ions
(d) Alkaline cell cell to flow through but can prevent the two
different aqueous solutions from mixing.
Daniell Cell
1 A Daniell cell has copper metal as the positive
terminal and zinc metal as the negative
terminal.
2 The zinc metal is immersed in zinc sulphate
solution and the copper metal is immersed in
copper(II) sulphate solution.
3 The two solutions of the Daniell cell are
connected using either of the following:
(a) A salt bridge as shown in Figure 6.14.
Electrochemistry 160
7 The functions of salt bridges and porous pots are 2 The electrolyte is ammonium chloride in the 6
SPM (a) to allow the flow of ions so that the form of a paste.
’11/P1 circuit is completed.
3 The cross-section of a dry cell is shown in
(b) to prevent the two aqueous solutions Figure 6.16.
from mixing. This will prevent displace
ment reaction between a more electro Figure 6.16 Dry cell
positive metal and the salt solution of the 4 When the dry cell is in use, the zinc metal
less electropositive metal from taking place.
releases electrons and dissolves to form Zn2+
8 In a simple voltaic cell made by immersing ions.
both the zinc metal and copper metal in At the negative terminal:
copper(II) sulphate solution,
(a) zinc metal reacts directly with copper(II) Zn → Zn2+ + 2e–
sulphate solution in a displacement 5 Electrons flow from the zinc metal casing
reaction.
through the external circuit to the carbon rod,
Zn + CuSO4 → ZnSO4 + Cu where NH4+ ions receive electrons to produce
ammonia gas and hydrogen gas.
As a result, the zinc metal will be coated At the positive terminal:
by a layer of copper metal.
(b) the electric current decreases rapidly. 2NH4+ + 2e– → 2NH3 + H2
9 The salt bridge or porous pot prevents the 6 When the cell produces an electric current,
zinc from reacting directly with the copper(II)
sulphate solution. zinc metal dissolves. When the zinc metal
10 When the negative terminal (zinc) is connec casing is perforated and the electrolyte starts
ted to the positive terminal (copper), the to leak out, the dry cell can no longer be used.
highest voltage produced is 1.10 V if both zinc 7 Usually a dry cell produces quite a stable
sulphate solution and copper(II) sulphate voltage of about 1.5 V.
solution have a concentration of 1.0 mol dm–3.
11 The reactions that take place in a Daniell cell Dry cells of different sizes
are as follows:
At the negative terminal: Zn → Zn2+ + 2e–
At the positive terminal: Cu2+ + 2e– → Cu
Overall chemical equation of cell:
Zn + Cu2+ → Zn2+ + Cu
12 When the Daniell cell is in use,
(a) the copper metal becomes thicker (mass
increases),
(b) the zinc metal becomes thinner (mass
decreases),
(c) the concentration of copper(II) sulphate
solution decreases, hence the blue colour
of the solution becomes paler,
(d) the concentration of zinc sulphate solution
increases.
Dry Cell Dry cells that are out of electricity (old) need to be
removed from the electrical appliance. This is because
1 A dry cell consists of a carbon rod (positive when the container is corroded, the electrolyte will
terminal) and a metal casing made of zinc leak out to damage the electrical appliance.
(negative terminal).
161 Electrochemistry
Lead-acid Accumulator 6 In the recharging process,
(a) reverse reactions occur at both electrodes.
1 A lead-acid accumulator is made of pieces (b) lead(II) sulphate is converted back into
of lead plates immersed in moderately lead(IV) oxide and hence lead(II) sulphate
concentrated sulphuric acid as shown in dissolves.
Figure 6.17. (c) sulphuric acid is formed.
7 An accumulator normally consists of 6 pairs
of plates and produces a voltage of 12 V.
Other Types of Voltaic Cells
6 Mercury cell
1 A mercury cell consists of zinc (negative
Figure 6.17 Lead-acid accumulator is used as car
batteries. terminal), mercury(II) oxide, HgO (positive
terminal) and a mixture of potassium
2 When the accumulator is used to produce hydroxide, KOH and zinc oxide, ZnO as
current, the following changes occur. electrolyte.
(a) At the negative terminal, a lead atom Mercury cells are small
donates 2 electrons to form a Pb2+ ion. 2 Mercury cells are small and long-lasting,
Pb → Pb2+ + 2e– producing a constant voltage of 1.3 V.
3 Mercury cells are used in hearing aids,
(b) At the positive terminal, lead(IV) oxide
accepts electrons and reacts with H+ ions digital watches and heart pacemakers.
in dilute sulphuric acid to form Pb2+ ions.
Alkaline cell
PbO2 + 4H+ + 2e– → Pb2+ + 2H2O
1 Alkaline cells are non-rechargeable cells.
(c) Lead(II) ions from both electrodes 2 An alkaline cell consists of zinc (negative
ctoomprboidnuecwe iltehaSdO(I4I2)– ions in sulphuric acid terminal), carbon rod (positive terminal)
sulphate, PbSO4. surrounded by manganese(IV) oxide, MnO2
and alkali (potassium hydroxide and sodium
Pb2+ + SO42– → PbSO4 hydroxide) as the electrolyte.
The overall chemical reaction is represented
by the equation below:
Pb + PbO2 + 4H+ + 2SO42– → 2PbSO4 + 2H2O
3 In this reaction, sulphuric acid is used up and
water is produced. Hence, sulphuric acid
becomes more dilute and its density decreases.
4 Lead(II) sulphate is insoluble and exists as a
white precipitate. When the precipitate covers
the surface of both electrodes, further reaction
is prevented and no electric current will be
produced.
5 The accumulator can be recharged by passing
an electric current in the opposite direction
to renew the cell.
Electrochemistry 162
Nickel-cadmium cell
Nickel-cadmium cells are A nickel-cadmium cell Other new types of cells include
rechargeable cells consists of cadmium lithium ion, nickel hydride and
(negative terminal), polymeric cells. These cells are
nickel(IV) oxide, NiO2
(positive terminal) rechargeable cells. Unlike the
and alkali, potassium
lithium ion and nickel hydride cells
hydroxide, KOH as the
electrolyte. which require battery casings, the
polymeric cell is flexible and can
be specifically shaped to fit the
device it will power.
Advantages and Disadvantages of Various Types of Voltaic Cells 6
Table 6.3 Advantages and disadvantages of various types of voltaic cells
Type of cell Advantages Disadvantages
1 Daniell cell • Can be prepared easily in the • A type of wet cell, the electrolyte
laboratory spills easily
• Voltage is not constant
2 Dry cell • Cheap • Does not last long
• No spillage as it is a dry cell • Cannot be recharged
• Produces a moderately constant • Zinc metal casing dissolves and the
current and voltage electrolyte that leaks out may corrode
• Portable, can be carried around easily electrical instruments
• Available in different sizes • Current and voltage produced is low
3 Alkaline cell • Lasts longer than a dry cell • Cannot be recharged
• Produces a higher and more stable • Cost more than a dry cell
• If leakage occurs, electrolyte is
voltage
• Portable corrosive
4 Mercury cell • Can be made into very small sizes • Expensive
• Produce a constant voltage for a long • Cannot be recharged
• Mercury which is produced is toxic
period
• Can last for a long period of time
5 Lead-acid accumulator • Can be recharged repeatedly • Acid may spill
• Produces a high current (175 A), • Heavy and is difficult to be carried
suitable for heavy work such as around
starting a car engine • Loss of charge occurs if not used for a
• Produce a high voltage (12 V) for a
long period long time
• Lead plates are easily corroded after a
long period of usage
6 Nickel-cadmium cell • Can be charged repeatedly • Expensive
• No spillage occurs because it is a dry • Requires a transformer for the
cell recharging process
• The size is smaller that an accumulator
163 Electrochemistry
Comparison of Voltaic Cells and Electrolytic Cells SPM
’10/P2
1 Table 6.4 below shows several similarities and differences between an electrolytic cell and a voltaic cell.
Table 6.4
Electrolytic cell Voltaic cell
6 Figure 6.18 Electrolytic cell Figure 6.19 Voltaic cell
Similarities
• Contains an electrolyte • Positive ions and negative ions move in the
• Consists of an anode and a cathode electrolyte
• Electrons move from the anode to the cathode in
• Chemical reactions involve the donation (at the
the external circuit (connecting wires) anode) or acceptance (at the cathode) of electrons
Electrolytic cell Differences Voltaic cell
• A battery is required to supply Basic structure • A battery is not required to
electrical energy supply electrical energy
Energy conversion
• Graphite (carbon) is usually • Graphite (carbon) is not used
used as electrodes Transfer of electrons at as electrodes
the positive electrode
• Electrodes are not made up of • Electrodes are made up of two
different metals Transfer of electrons different metals
at the negative
• Electrical energy is converted electrode • Chemical energy is converted
into chemical energy into electrical energy
Transfer of electrons
• Anode (positive electrode): in the external circuit • Cathode (positive electrode):
Anions (negative ions) lose Oxidising agent accepts
electrons at the anode electrons from the cathode
2X – → X2 + 2e– ... oxidation Cu+(aq) + 2e– → Cu(s)
… reduction
• Cathode (negative
electrode): Cations (positive • Anode (negative electrode):
ions) accept electrons from the Reducing agent releases
cathode electrons
Y 2+ + 2e– → Y ... reduction Zn(s) → Zn2+(aq) + 2e–
… oxidation
• Electrons flow from the anode
(positive electrode) to the • Electrons flow from the anode
cathode (negative electrode) (negative electrode) to the
cathode (positive electrode)
Electrochemistry 164
The table shows the differences between the terminals in voltaic and electrolytic cells as a result of the transfer of electrons.
Transfer of electrons Electrons are donated Electrons are accepted
Type of cells
At the negative terminal (anode) At the positive terminal (cathode)
Voltaic cells At the positive terminal (anode) At the negative terminal (cathode)
Electrolytic cells
6.5 (d) Predict the observations obtained after the 6
voltaic cell is used for some time.
1 (a) Draw the circuit diagram for a simple voltaic
cell consisting of iron metal, copper metal and 2 Compare the advantages and disadvantages of dry
copper(II) sulphate solution. Show the direction cells and alkaline cells.
of the flow of electrons in the circuit diagram.
3 Give an example of a voltaic cell that can be recharged.
(b) Which metal serves as the negative terminal? Explain how the reactions occur at the positive terminal
(c) Write the half-equations for the reactions that and the negative terminal to produce an electric current.
occur at both electrodes.
6.6 The Electrochemical 4 The electrochemical series can be constructed
Series
by two methods.
1 The electrochemical series is an arrangement (a) The potential difference (voltage difference)
of elements according to their tendencies to
form ions. between pairs of metals
(b) The ability of a metal to displace another
2 In the electrochemical series, a metal that has
a higher tendency to ionise and form positive metal from its salt solution.
ions (by releasing electrons) is placed at a higher
position in the series. Hence, metal ions at (A) To Construct the Electrochemical
Series Based on the Potential SPM
the upper positions of the electrochemical Difference (Voltage Difference) ’08/P1
’09/P2
series are less likely to receive electrons to
1 Metals are arranged in the electrochemical
form metal atoms. series according to their tendencies to donate
3 Part of the electrochemical series (for metal electrons to form cations.
elements) is shown below. 2 The electrochemical series can be constructed
based on the measurement of the potential
Metals Positive ions (cations) difference between two metals in voltaic cells.
K ⎯⎯⎯→ K+ + e– 3 When two different metals (immersed in their
Na ⎯⎯⎯→ Na+ + e–
Ca ⎯⎯⎯→ Ca2+ + 2e– respective salt solutions) are connected in the
Tendency Mg ⎯⎯⎯→ Mg2+ + 2e– Tendency
of metal Al ⎯⎯⎯→ Al3+ + 3e– of cations external circuit through a voltmeter and a salt
atoms to Zn ⎯⎯⎯→ Zn2+ + 2e– to accept
donate Fe ⎯⎯⎯→ Fe2+ + 2e– electrons bridge:
electrons Sn ⎯⎯⎯→ Sn2+ + 2e– to form (a) The metal that serves as the negative terminal
to form Pb ⎯⎯⎯→ Pb2+ + 2e– metals
ions H ⎯⎯⎯→ H+ + e– increases of the voltaic cell has a higher tendency to
increases release electrons. Hence, that metal is placed
Cu ⎯⎯⎯→ Cu2+ + 2e– at a higher position in the electrochemical
Ag ⎯⎯⎯→ Ag+ + e–
series. Conversely, the metal that serves as
the positive terminal is placed at a lower
position in the electrochemical series.
(b) The further apart the positions of two
metals in the electrochemical series, the
greater the potential difference (voltage).
165 Electrochemistry
6.7
To construct the electrochemical series through the potential difference (voltage) SPM
of pairs of metals ’06/07
P3
Problem statement Procedure
How to construct the electrochemical series based
on the measurement of the potential differences
between pairs of metals in simple voltaic cells?
Experiment 6.7 Hypothesis Figure 6.20
6
Two principles are used in the construction of the 1 Pieces of zinc, magnesium, iron, aluminium and
electrochemical series: silver metals are polished with sandpaper.
(a) A metal that serves as the negative terminal
2 A piece of zinc metal and a piece of copper metal
of a cell is placed at a higher position in the are connected to a voltmeter by the connecting
electrochemical series. wires with crocodile clips.
(b) The bigger the voltage differences of the voltaic
cells, the further apart the positions of the two 3 The two metals are then dipped in the sodium
metals in the electrochemical series. chloride solution in a beaker as shown in
Figure 6.20.
Variables
4 The highest cell voltage obtained is recorded.
(a) Manipulated variable : Pairs of metals as elec- 5 The direction of the flow of electrons is also noted
trodes
to determine the terminals of the voltaic cell.
(b) Responding variable : Voltage values of voltaic Electrons flow from the negative terminal to the
cells positive terminal. If the voltmeter reading shows
a negative value, the metal pairs connected to the
(c) Constant variables : Type and concentration terminals of the voltmeter should be reversed.
of electrolytes 6 Zinc metal is then replaced by other metals in
turn: magnesium, iron, aluminium and silver.
Apparatus The highest cell voltage of every pair of metals
is recorded.
Voltmeter, beaker, connecting wires with crocodile
clips and sandpaper.
Materials
Sodium chloride solution of 1.0 mol dm–3, pieces of
copper, zinc, magnesium, iron, aluminium and silver
metals.
Results
Pairs of metals Positive terminal Negative terminal Potential difference (V)
Zn/Cu Copper Zinc 1.1
Mg/Cu Copper Magnesium 2.7
Fe/Cu Copper Iron 0.8
Al/Cu Copper Aluminium 2.0
Ag/Cu Silver 1.1
Copper
Conclusion 2 Silver serves as the positive terminal when it is
connected to copper. Hence, silver is placed at a
1 Copper metal serves as the positive terminal of lower position than copper in the electrochemical
the voltaic cells when paired with zinc, magne series.
sium, iron and aluminium metal. Hence, copper
is at a lower position than zinc, magnesium, iron 3 The further apart the distance between the metals
and aluminium in the electrochemical series. in the electrochemical series, the greater the
potential difference (voltage).
Electrochemistry 166
4 The arrangement of the metals in the electrochemical series based on the voltage (potential difference)
of the cell is as follows:
Higher tendency Magnesium 2.0 V 2.7 V The bigger the voltage
to release Aluminium reading, the further the
electrons Zinc 1.1 V distance between the
Iron metals
Copper 0.8 V
Silver 1.1 V
6
4 SPM
’10/P1
The table shows information about three simple cells. Comments
• In the metal pair of P and Q, P is the positive
Metal pairs Potential Positive
difference (V) terminal terminal. Hence P is placed below Q in the
P and Q electrochemical series. Similarly, S is placed below
Q and S 1.7 P Q in the electrochemical series.
R and S • The potential difference between Q and S is bigger
2.1 S than that between Q and P. Thus S is placed below P.
• In the metal pair R and S, R is the positive terminal.
0.6 R Hence R is placed below S.
• The arrangement of the metals according to their
What is the potential difference of the metal pair P increasing tendencies to form metal ions is as
and R? follows:
Q 1.7 V 2.1 V
P
Higher tendency to release S The bigger the voltage reading, the
electrons R further the distance between the
metals.
0.6 V
Answer The potential difference between P and R = (2.1 – 1.7) + 0.6 = 1.0 V
(B) To Construct the Electrochemical (a) metal M is more likely to release electrons
Series from the Displacement than metal N.
Reactions of Metals
(b) metal M is more electropositive than
1 The electrochemical series can also be metal N.
constructed based on the ability of a metal
to displace another metal from its salt (c) metal M is placed at a higher position than
solution. metal N in the electrochemical series.
2 If metal M can displace metal N from an 3 Alternatively, if metal P is immersed in an
aqueous N salt solution, then aqueous Q2+ ion solution and no reaction
takes place, then metal P is at a lower position
than metal Q in the electrochem ical series.
167 Electrochemistry
6.8
SPM
’08/P2,
To construct the electrochemical series from displacement reactions ’07/P1,
’04/P2
Problem statement 3 A piece of magnesium metal is placed in the
How to construct the electrochemical series based on solution of every test tube except that of its salt
the ability of a metal to displace another metal from solution (Figure 6.21).
its salt solution?
Hypothesis
Experiment 6.8 A metal that can displace another metal from its
6 salt solution is placed at a higher position in the
electrochemical series. The greater the number of metals
that can be displaced by a metal from their solutions, the
higher its position in the electrochemical series.
Variables
(a) Manipulated variable : Different types of metal Figure 6.21
and their salt solutions
(b) Responding variable : Deposition of metals or
colour change in the salt
solutions
(c) Constant variable : Concentration of nitrate
salt solutions
Apparatus
Test tubes, test-tube rack and sandpaper. 4 Observations are made after awhile to check if
(a) there is any colour change in the solution,
Materials (b) there are any solid deposits on the
Pieces of magnesium, zinc, iron, tin, lead and copper magnesium metal,
metals, solutions of copper(II) nitrate, lead(II) (c) magnesium metal dissolves.
nitrate, tin(II) nitrate, iron(II) nitrate, zinc nitrate and 5 If any of the above occurrences (a), (b) or (c)
magnesium nitrate (concentration and volume of all salt
solutions are 0.5 mol dm–3 and 10 cm3 respectively). is observed, displacement reaction has taken
Procedure place: a tick symbol, (✓) is marked in the table
1 Pieces of magnesium, zinc, copper, tin, lead and of results.
iron metals are polished with sandpaper. 6 If there is no noticeable observation, a cross
2 10 cm3 of 0.5 mol dm–3 solutions of copper(II) symbol, (✗) is marked at the table to indicate that
nitrate, lead(II) nitrate, tin(II) nitrate, iron(II)
nitrate, zinc nitrate and magnesium nitrate are displacement reaction did not take place.
placed into separate test tubes. 7 The experiment is repeated using different
metals and fresh solutions of ions. The results of
the experiment are shown in the table below.
Results
Metal Solution Pb(NO3)2 Sn(NO3)2 Fe(NO3)2 Zn(NO3)2 Mg(NO3)2
Cu(NO3)2
Magnesium, Mg ✓ ✓ ✓ ✓ ✓ –
Zinc, Zn ✓✓✓✓ – ✗
Iron, Fe ✓✓✓ – ✗ ✗
Tin, Sn ✓✓ – ✗ ✗ ✗
Lead, Pb ✓– ✗ ✗ ✗ ✗
Copper, Cu –✗✗✗✗✗
Electrochemistry 168
Conclusion 4 The result of the experiment shows that the
order of the positions of the metals in the
1 Metals can be arranged according to the number electrochemical series is:
of tick symbols (3) recorded (or the number of
metals displaced in reactions). The more (3) Mg Zn Fe Sn Pb Cu
symbols, the more reactive the metal is and
the position of the metal is placed higher in the Electropositivity of metal decreases
electrochemical series.
5 The electrochemical series can be constructed
2 Magnesium is placed at the highest position in from displacement reactions.
the electrochemical series because it can displace The hypothesis is accepted.
all the other metals from their solutions.
6
3 Copper is placed at the lowest position in the
electrochemical series because copper cannot
displace any other metals in this experiment.
The Uses of the Electrochemical Series To predict the ability of a metal to displace
another metal from its salt solution
To determine the terminals of voltaic cells
1 A metal that is at a higher position in the
1 When two different metals are connected by electrochemical series can displace another
wires and then immersed in an electrolyte, metal that is lower than itself in the electro
a simple voltaic cell is formed. The metal chemical series from its salt solution.
that is placed at a higher position in the
electrochemical series will become the 2 For example, aluminium is above iron in the
negative terminal of the cell. The metal that electrochemical series, hence aluminium can
is placed lower in the electrochemical series displace iron from an iron(II) salt solution
will become the positive terminal of the cell. (such as iron(II) sulphate solution).
2 The metal that is placed higher in the To predict whether a metal can displace
electrochemical series is more electropositive hydrogen from an acid
and has a higher tendency to release electrons.
Electrons will flow from the negative terminal 1 Hydrogen ion is placed between lead(II) ion
to the positive terminal. and copper(II) ion in the electrochemical
series.
3 For example, in the zinc/copper simple
voltaic cell, zinc metal will become the 2 All other metals that are placed at higher
negative terminal of the cell because zinc is positions than hydrogen ion in the
above copper in the electrochemical series. electrochemical series can displace hydrogen
Copper metal will become the positive from acids.
terminal of the cell. For example, zinc is above hydrogen in
the electrochemical series; hence zinc can
To compare the standard voltage of SPM displace hydrogen from hydrochloric acid.
the voltaic cell ’09/P1 Zn + 2HCl → ZnCl2 + H2
1 The further the distance between two 3 Metals below hydrogen ion in the electro
metals in the electrochemical series, the chemical series cannot displace hydrogen
greater the cell voltage will be. from acids. Examples are copper, mercury
and silver.
2 For example, the distance between magnesium
and copper is further than that between
zinc and copper in the electrochemical
series. Hence the cell voltage produced by
a magnesium/copper voltaic cell is greater
than that from a zinc/copper voltaic cell.
169 Electrochemistry
6.9
To confirm the predictions of displacement reactions
Problem statement Materials
Experiment 6.9 How can the prediction of the displacement reaction Pieces of magnesium, iron and copper, solutions of
6 of a metal from its salt solution by another metal be copper(II) sulphate, iron(II) sulphate and magnesium
confirmed? sulphate (concentration and volume of all salt
solutions are 0.5 mol dm–3 and 10 cm3 respectively).
Hypothesis
If metal X is at a higher position than metal Y in Procedure
the electrochemical series, then metal X can displace
metal Y from a salt solution of metal Y. 1 Pieces of magnesium, copper and iron metals are
polished with sandpaper.
Variables
(a) Manipulated variable : Different types of metals 2 10 cm3 of 0.5 mol dm–3 solutions of copper(II)
sulphate, iron(II) sulphate and magnesium
and salt solutions sulphate are put into different test tubes.
(b) Responding variable : Deposits of metal or
3 Magnesium, copper and iron metals are placed in
colour change in the salt different salt solutions in the test tubes.
solution
(c) Constant variable : Concentration of salt 4 Observations are made after 20 minutes to check
solutions if there is
(a) any change in colour of the solution,
Apparatus (b) any deposits of metal,
Test tubes, test-tube rack and sandpaper. (c) any corrosion of metal.
5 The result of the experiment is tabulated in the
table below.
Results
Metal + salt solution Prediction Observation
Magnesium + iron(II) sulphate solution Displacement occurs because • Grey deposit is formed
magnesium is more • The colour of the green solution
electropositive than iron
becomes paler
• Magnesium dissolves
Magnesium + copper(II) sulphate Displacement occurs because • Brown deposit is formed
solution magnesium is more • The colour of the blue solution
electropositive than copper
becomes paler
• Magnesium dissolves
Iron + magnesium sulphate solution Displacement does not occur No noticeable change
because iron is less
electropositive than magnesium
Iron + copper(II) sulphate solution Displacement occurs because • Brown deposit is formed
iron is more electropositive • The blue coloured solution
than copper
changes to green
• Iron dissolves
Copper + magnesium sulphate solution Displacement does not No noticeable change
occur because copper is less
electropositive than magnesium
Copper + iron(II) sulphate solution Displacement does not No noticeable change
occur because copper is less
electropositive than iron
Electrochemistry 170
Discussion Fe + Cu2+ → Cu + Fe2+ 6
1 A more electropositive metal can displace a less
4 Copper cannot displace either magnesium or iron
electropositive metal from its salt solution. from their salt solutions because copper is below
2 Magnesium is at a higher position than iron magnesium and iron in the electrochemical
series.
and copper in the electrochemical series. Hence
magnesium can displace iron from iron(II) Conclusion
sulphate solution and copper from copper(II) The prediction that a metal at a higher position in
sulphate solution. the electrochemical series can displace a metal
which is at a lower position from its salt solution
Mg + Fe2+ → Mg2+ + Fe is confirmed.
Mg + Cu2+ → Mg2+ + Cu The hypothesis is accepted.
3 Iron is at a higher position than copper in the
electrochemical series. Hence iron can displace
copper from copper(II) sulphate solution.
A series of metals based on their Electrochemical series: This series is used to:
tendencies to form metal ions (a) predict the tendency of a metal to
Potassium, K
A metal placed higher in the Sodium, Na form ions.
series has a higher tendency to Calcium, Ca (b) predict the ability of a metal to
form positive ions (it is more Magnesium, Mg
electropositive). For example: Aluminium, Al displace another metal from its ionic
Zn is more electropositive than Fe. Zinc, Zn solution.
Iron, Fe (c) determine the terminals and voltage
A metal placed at a lower position Tin, Sn of a voltaic cell.
in the series can be displaced by a Lead, Pb
metal above it. Hydrogen, H In a voltaic cell, the metal that becomes
Copper, Cu the negative terminal is the metal that is
Silver, Ag higher in position in the electrochemical
Gold, Au series as it has a higher tendency to form
ions.
The further the distance between the
metal pairs in a voltaic cell, the greater the
cell voltage will be.
6.6 (i) Which metal will become the negative
terminal of the cell?
1 The table below shows the voltages obtained from
three voltaic cells using different pairs of metals. (ii) Predict the voltage of the cell.
(c) Predict what will happen if
Voltaic cell Metal pairs Voltage (V) Positive
electrode (i) metal Y is immersed in a solution of Z salt.
(ii) metal X is immersed in a solution of Y salt.
1 X and Y 1.2 X (iii) metal W is immersed in a solution of X salt.
2 X and Z 0.9 X 2 Silver is placed at a position lower than copper and
magnesium in the electrochemical series. Predict
3 Y and W 0.4 Y the observation and reaction that will occur in the
following experiment:
(a) Based on the observation above, arrange the (a) Silver in copper(II) sulphate solution,
metals W, X, Y and Z in an ascending order (b) Copper in silver nitrate solution,
according to their electropositivity. (c) Magnesium in silver nitrate solution.
(b) A voltaic cell is made from metal Z and metal W.
171 Electrochemistry
6 6.7 Developing Awareness 2 However, a safe and systematic method of
and Responsible disposal of used batteries and industrial
Practices when Handling by-products in electrochemical industries is
Chemicals used in the important to prevent environmental pollution.
Electrochemical Industries (a) Used batteries should be separated from
other household disposal. They are
1 Electrochemical industries play an important required to be disposed off separately
role in our daily life by improving our quality to prevent the chemicals of the batteries
of life. from leaking and polluting water sources.
(a) For example, useful metals such as Parts of batteries that are useful should be
aluminium, sodium and magnesium are recycled.
extracted from their minerals or compounds (b) Chemical wastes from electrolytic
using electrolysis. industries should be treated to remove
(b) Useful chemical substances such as chlorine poisonous chemicals before being disposed
and sodium hydroxide are manufactured as industrial waste. For example,
on a large scale using electrolysis.
(c) Electroplating of iron with chromium (i) acids that are used to clean metals
protects the iron components of machinery before electroplating should be
from corrosion. Silver-plating is commonly diluted and neutralised before
used in the making of fine cutleries. draining off as waste water.
(d) Various voltaic cells are used in different
devices such as radio, torchlight, quartz (ii) metal ions that are toxic and hazardous
watch, handphone and others. to human health such as cadmium
ion, chromium ion and nickel ion
need to be treated and removed from
industrial effluent.
1 An electrolyte is a chemical compound which conducts 8 The factors that determine the types of ions to be
electricity in the molten state or in an aqueous
solution and undergoes chemical changes. discharged at the electrodes are
(a) positions of ions in the electrochemical
2 A non-electrolyte is a chemical compound which
does not conduct electricity in any state. series: The lower positioned ion will be
3 Electrolysis is the decomposition of an electrolyte discharged
(b) concentration of ions in the solution: The
(molten or in aqueous solution) by the passage of
an electric current. more concentrated ion will be discharged
4 Graphite or platinum is usually used as electrodes (c) types of electrodes used
because they are inert. 9 Uses of electrolysis in industries
5 The anode is the electrode connected to the (a) Extraction of reactive metals. For example:
positive terminal of the batteries.
6 The cathode is the electrode connected to the Extraction of aluminium from molten bauxite
negative terminal of the batteries.
7 Two steps occur during electrolysis. (b) R(Aelf2iOn3in) gusoinf gmceatrablos.nFoerleecxtraomdepsle. : Purification of
(a) Movement of ions to the electrodes:
copper.
Cations (positive ions) move towards the (c) Electroplating of metals. For example: Copper
cathode (negative electrode) whereas anions
(negative ions) move towards the anode plating or silver plating.
10 There are two types of voltaic cells:
(positive electrode).
(b) Discharge of ions at the electrodes: (a) Primary cells: Non-rechargeable cells (cells
Cations discharge by receiving electrons. that cannot be charged again).
(b) Secondary cells: Rechargeable cells (cells that
Generally: An+ + ne– → A
Anions discharge by releasing electrons. can be charged again).
11 The electrochemical series is an arrangement of
Generally: Bn– → B + ne–
elements based on their tendencies to form ions.
Electrochemistry 172
6
Multiple-choice Questions
6.1 Electrolytes and Non- 4 Which of the following A Zinc metal is formed at 6
electrolytes statements are true about an
electrode X.
1 Which of the following can ’06 electrolyte? B Chlorine gas is formed at
conduct electricity? I It has ions that conduct
A Ethanol electrode Y.
B Solid lead(II) nitrate electricity in the solid state. C Zinc ions are attracted to the
C Magnesium chloride solution II It can conduct electricity
D Liquid tetrachloromethane anode.
in the molten state or in D Chloride ions are discharged
2 Calcium carbonate powder does aqueous solution.
not conduct electricity because III It is a compound with ionic at the positive electrode.
A it does not contain ions. bonds only.
B it contains covalent IV It can be decomposed by 7 When molten lead(II) iodide
molecules. electric current.
C it contains calcium ions and A I and II only solution is electrolysed using
carbonate ions that are not B III and IV only
free to move. C II and IV only carbon electrodes, which of
D all the atoms in calcium D I, II, III and IV
carbonate are bonded by the following represents the
strong covalent bonds. 5 Which of the following statements
are true about electrolysis? half-equation that occurs at the
3 The diagram shows the set-up of
apparatus to test the conductivity I The cathode is the positive cathode?
electrode.
’06 of the chemical in the beaker. A Pb2+ + 2e– → Pb
It was found that there is no II Molten covalent compounds B Pb → Pb2+ + 2e–
deflection on the ammeter can be used as electrolytes. CD 2I2I–→+ 2l– + 2e–
needle. 2e– → I2
III Platinum can be used as
inert electrodes. 8 Which of the following
IV A compound is decomposed substances will produce
by electric current.
aluminium metal when
A I and II only
B III and IV only electrolysis is carried out using
C II, III and IV only
D I, III and IV only carbon electrodes?
A Aqueous aluminium sulphate
solution
B Aqueous aluminium chloride
solution
C Solid aluminium oxide
D Molten aluminium oxide
6.2 Electrolysis of Molten 6.3 Electrolysis of Aqueous
6 Compounds Solutions
A 9 Which of the following compounds
produces oxygen gas and
Which of the following action ?@ hydrogen gas during electrolysis?
A Aqueous potassium
will cause a deflection of the heating hydroxide solution
B Saturated sodium chloride
ammeter’s needle? Which of the following occurs solution
A Add more dry cells in the when molten zinc chloride is C Aqueous copper(II) nitrate
electrolysed in t6h/9e apparatus as solution
circuit shown in the diagram? D Concentrated hydrochloric acid
B Add water to glacial ethanoic
10 The diagram below shows
acid the apparatus set-up for the
C Add ethanol to glacial
’09 electrolysis of potassium nitrate
ethanoic acid solution, KNO3.
D Substitute the platinum
Electrochemistry
electrodes with carbon
electrodes
173
carbon B Chlorine gas is evolved at the 17 In an experiment, dilute
electrode Y cathode. aqueous potassium iodide
solution is electrolysed using
potassium C Copper metal deposits at the carbon electrodes. Which of the
nitrate anode. following statements are true
solution about this experiment?
carbon D The intensity of the blue
electrode X colour of the solution remains I A gas that produces a small
constant. ‘pop’ sound when tested with
What are the products formed at a lighted wooden splint is
electrodTeCs 5X4and Y ? 14 The diagram shows the set-up produced at the cathode.
of apparatus for the electrolysis
II A gas that rekindles a glowing
’03 of iron(II) nitrate solution. wooden splint is produced at
the anode.
6 X Y
A Nitrogen Hydrogen III The solution around the
gas What is formed at electrode X ? anode changes to brown
gas Potassium A Iron colour.
B Nitrogen B Oxygen
Hydrogen C Hydrogen gas IV The concentration of
dioxide gas gas D Nitrogen dioxide gas potassium iodide solution
C Oxygen gas Oxygen gas increases.
D Hydrogen A I and II only
gas B II and IV only
C I and III only
11 The products formed at the 15 When aqueous magnesium D I, II and IV only
electrodes during the electrolysis of
aqueous sodium sulphate solution sulphate solution is electrolysed 18 Metal Y is placed at a high
using carbon electrodes are position in the electrochemical
using graphite electrodes, series. When a dilute Y chloride
Cathode Anode A the mass of cathode solution is electrolysed using
carbon electrodes, the product
A Sodium Sulphur increases. formed at the cathode is
B the mass of anode A hydrogen
B Hydrogen Sulphur dioxide B oxygen
decreases. C chlorine
C Hydrogen Oxygen C magnesium metal deposits at D metal Y
D Sodium Oxygen the cathode. 6.4 Electrolysis in Industries
D the concentration of
12 Electrolysis of dilute sodium 19 Electrolysis is used to extract
chloride solution using carbon magnesium sulphate solution aluminium metal from molten
electrodes produces oxygen aluminium oxide. Which
and hydrogen at the anode increases. chemical is used to lower the
and cathode respectively. melting point of aluminium
The products formed at the 16 Electrolysis of aqueous sodium oxide to 900°C?
electrodes will change if A Cryolite
A platinum is used as the iodide solution is carried out B Bauxite
cathode. C Silicon dioxide
B a bigger current flows ’09 using carbon electrodes. Which D Calcium carbonate
through the circuit.
C a concentrated sodium half-equation shows the reaction
chloride solution is used.
D the distance between the at the cathode?
electrodes is reduced.
A 2I– → I2 + 2e– + O2 + 4e–
13 Which of the following is true B 4OH– → 2H2O
about the electrolysis of aqueous C Na+ + e– → Na
copper(II) chloride solution using
copper electrodes? D 2H+ + 2e– → H2
A The mass of cathode
decreases. 20 What are the suitable chemicals and apparatus used to electroplate a
spoon with silver metal by electrolysis?
Anode Cathode Electrolyte
A Silver Spoon Silver chloride solution
B Spoon Silver Silver nitrate solution
C Carbon Spoon Silver nitrate solution
D Silver Spoon Silver nitrate solution
Electrochemistry 174
21 The presence of foreign metals 6.5 Voltaic Cells 26 Voltaic cells that are used in
in copper metal can reduce watches and calculators are
the conductivity of copper 24 The diagram below shows A dry cells
wire. Which of the following a simple chemical cell. Two B alkaline cells
is suitable to be used as the C mercury cells
cathode in the purification of ’11 different metals are used as D nickel-cadmium cells
copper by electrolysis? electrodes.
A Pure copper 27 The diagram shows the set-up
B Impure copper of apparatus of a chemical cell.
C Carbon
D Platinum 1 2 3 4 ’05
0
22 The diagram shows the 5 Which of the following occurs in
arrangement of apparatus to the chemical cell?
copper zinc A The magnesium rod becomes
’06 electroplate a metal key with plate plate
chromium. It is found that thicker.
electroplating does not occur. sodium B The iron rod becomes thinner.
How would you change the chloride C Electrons flow from iron to
arrangement of the apparatus solution
in order to plate a layer of magnesium.
chromium on the surface of the D The green colour of iron(II) 6
key?
Which of the following metals sulphate becomes paler.
A Replace chromium nitrate 28 The diagram shows the set-up of
solution with chromium can be used to replace the zinc
chloride solution apparatus for an electrochemical
plate to obtain the brightest light cell.
B Change the supply of direct
current to alternating current in the light bulb and the highest
C Reverse the terminals of the voltage reading?
batteries A Magnesium
B Iron
D Replace the chromium C Aluminum
electrode with carbon D Lead
25 When magnesium metal and
copper metal are connected
by wire and then immersed in
copper(II) sulphate solution,
which of the following does not
happen?
A Electron flows from copper
metal to magnesium metal.
B Mass of copper increases.
C Mass of magnesium
decreases.
D The colour intensity of blue
copper(II) sulphate solution
decreases.
23 The diagram shows the set-up of the apparatus used for the purification of a
metal through electrolysis.
Which of the following combinations is suitably used for the purification of Which of the following
copper metal? observations are true for this
experiment?
Electrode X Electrode Y Electrode Z I Zinc electrode becomes
A Pure copper Impure copper Copper(II) sulphate thinner.
B Impure copper Pure copper Copper(II) nitrate II Brown colour is formed
C Pure copper Impure copper Sulphuric acid
D Impure copper Pure copper Copper(II) carbonate around electrode X.
III Gray deposit is formed at
175
electrode Y.
IV Intensity of blue colour in
beaker M becomes paler.
A I and III only
B II and III only
C II and IV only
D I, II and IV only
Electrochemistry
29 The diagram shows the set-up Metal Negative Potential B Concentration of magnesium
of apparatus for a simple cell. electrode terminal difference
ion in the solution decreases.
pairs (V) C A brown deposit is formed.
D Zinc ions are formed.
Q-P Q 2.7
R-P R 1.1 36 The table below shows
S-T S 1.3 information about three voltaic
S-P S 2.1
’10 cells. Metals P, Q, R and S are
used as electrodes in the cells.
Which of the following pairs What is the potential difference Voltaic Negative Positive Voltage
of metals gives the highest cell terminal terminal (V)
voltmeter reading? of a voltaic cell made of metal IP
Q 0.9
6 Metal X Metal Y electrode pair Q-T ? II R
A 0.8 V Q 1.3
A Magnesium Iron B 1.4 V III R
B Zinc Copper C 1.9 V S 2.1
C Aluminium Silver D 3.5 V
D Silver Copper
33 If a piece of metal X is What is the order of the metals
30 A voltaic cell is made using
metal X and Y as the electrodes. immersed in copper(II) sulphate from the most electropositive to
If electrons flow from metal X
to metal Y, metal X and metal Y solution, a brown deposit is the least electropositive?
may be A P, Q, R, S
formed. Metal X may be B P, R, Q, S
A copper C R, P, Q, S
B platinum D S, Q, P, R
C aluminium
D silver 37 Which of the following pairs can
Metal X Metal Y 34 Two voltaic cells are constructed undergo a displacement reaction?
as shown in the diagram. The A Magnesium and potassium
A Iron Silver voltmeter reading of cell I is 1.1
V while that of cell II is 2.5 V. chloride solution.
B Silver Copper B Calcium and zinc sulphate
C Iron Magnesium
solution.
D Copper Zinc C Iron and calcium nitrate
31 Which of the following solution.
statements is not true about D Copper and magnesium
lead-acid accumulator? nitrate solution.
A The electrolyte used is
Which of the following is true of 38 When an iron nail is immersed
sulphuric acid. a voltaic cell constructed using in X solution, Fe2+ ions are
B Lead plate is the negative metal Q and metal R? produced. Solution X may be
A Metal Q will be the negative A magnesium sulphate
terminal. B zinc nitrate
C Carbon is the positive terminal. C copper(II) nitrate
B Electrons will flow from metal D sodium chloride
terminal.
D Lead(II) sulphate is formed R to metal Q. 39 Excess metal X powder is added
C The cell will produce a to copper(II) sulphate solution
when the cell is being and is stirred. After half an hour,
reading of 3.6 V. the solution becomes colourless
used. D R ions and Q ions are and brown deposit is formed.
Metal X may be
6.6 The Electrochemical formed.
Series I calcium
35 A piece of zinc metal is II aluminium
32 An experiment is carried out immersed in a beaker containing III magnesium
to measure the potential a mixture of copper(II) sulphate IV silver
and magnesium nitrate solution.
’06 differences produced in voltaic Which of the following does not A I and II only
cells made from metal electrode happen? B III and IV only
pairs Q-P, R-P, S-T or S-P metals. A Zinc metal dissolves. C I, II and III only
The results of the experiment is D II, III and IV only
recorded in the table below.
Electrochemistry 176
40 The diagram shows four simple chemical cells. In each cell, copper is one In which cell does copper act as
of the electrodes.
the negative terminal?
’05 A Cell I
B Cell II
C Cell III
D Cell IV
6
Structured Questions 2 In an experiment, electrolysis of 0.001 mol dm–3
hydrochloric acid is carried out using a electrolytic cell
1 In an experiment, different chemical substances are as shown in Diagram 2. Gases are collected at both
tested using the set-up of apparatus as shown in the electrodes.
Diagram 1.
Diagram 1 Diagram 2
(a) When naphthalene is used as the chemical in (a) Write the formulae of all the ions present in
the experiment, the light bulb does not light up. hydrochloric acid. [1 mark]
Explain this observation. [1 mark]
(b) Name a suitable material that can be used as the
(b) When lead(II) bromide solid is used as the electrodes in this experiment. [1 mark]
chemical in the experiment, the light bulb does
not light up but lights up when lead(II) bromide (c) (i) Name gas X and gas Y. [2 marks]
is heated to the molten form. Explain this
observation. [2 marks] (ii) Write the half-equation for the reaction that
(c) Predict the observation that will take place occurs at the anode. [1 mark]
at the anode and the cathode when molten
lead(II) bromide is used as the chemical in this (d) After the electrolysis is carried out for 50 minutes,
experiment. [2 marks]
the concentration of hydrochloric acid increases
(d) Write the half-equations for the reactions that
occur at the anode and the cathode in (c). and a different gas is collected at the anode.
[2 marks]
(i) Explain why the concentration of
(e) Predict the products that will be formed if molten
zinc chloride is used instead of lead(II) bromide hydrochloric acid increases. [2 marks]
in this experiment. [2 marks]
(ii) Name the new gas collected at the anode
and explain why this gas is produced.
[2 marks]
(iii) Write the half-equation for the reaction that
occurs at the anode in (ii). [1 mark]
177 Electrochemistry
3 Diagram 3 shows the arrangement of apparatus in an (a) Write the formula of all the cations present in the
electrochemistry experiment.
copper(II) sulphate solution. [1 mark]
(b) State the direction of the flow of electrons in Cell
Q. [1 mark]
(c) (i) State the observation at the cathode of Cell
P. [1 mark]
(ii) Write a half-equation for the reaction that
occurred at the cathode of Cell P. [1 mark]
(iii) Predict the change of colour intensity of the
copper(II) sulphate solution of cell P.
Diagram 3 [1 mark]
(a) What is the difference between the energy (iv) Name the product formed at the anode if
change in Cell A and Cell B? [2 marks] copper electrodes in Cell P are replaced by
6 carbon electrodes. [1 mark]
(b) Write half-equations for the reactions that occur (d) Based on cell Q:
at the (i) State the observation on the zinc plate.
(i) magnesium electrode in Cell A. [1 mark] [1 mark]
(ii) copper electrode in Cell A. [1 mark] (ii) Write the half-equations for the reaction that
(c) (i) Name the electrode that serves as the occurs at the zinc plate. [1 mark]
negative terminal in Cell B. [1 mark]
(iii) Write an overall ionic equation for the
(ii) State the reason for your answer in (i).
reaction that has taken place. [1 mark]
[1 mark]
(iii) State the direction of the flow of electrons (iv) What happens to the cell voltage if the
in Diagram 3. [1 mark] copper plate is replaced with a silver plate?
[1 mark]
(d) In Cell B, 5 Diagram 5 shows a voltaic cell that is formed from
copper metal and lead metal.
(i) name the product formed at the carbon
electrode Q and write an equation for the
reaction that occurs. [2 marks]
(ii) name the product formed at the carbon
electrode P and write an equation for the
reaction that occurs. [2 marks]
(e) What would happen if the magnesium electrode
in Cell A is replaced with a silver electrode?
[1 mark]
(f) What would happen if carbon electrodes P and Q Diagram 5
are replaced with copper electrodes? [1 mark]
4 Diagram 4 shows two types of cell. (a) State the positive terminal and the negative
terminal of the voltaic cell. [2 marks]
(b) Write ionic equations showing the reactions that
occur at
(i) the negative terminal of the cell. [1 mark]
(ii) the positive terminal of the cell. [1 mark]
(c) Write the overall ionic equation of the cell.
[1 mark]
(d) What is the function of the salt bridge? [1 mark]
(e) The voltage of the above cell is 0.57 V. If
magnesium is above lead in the electrochemical
series, what would be the expected voltage
produced from a magnesium/copper voltaic cell?
[1 mark]
6 An electrolysis process is carried out using the
arrangement of apparatus as shown in Diagram 6.
Diagram 4
Electrochemistry 178
(c) Write the ionic equation that occurs at
(i) electrode L
(ii) electrode M [2 marks]
(d) What is the product of electrolysis formed at
(i) electrode R?
(ii) electrode S? [2 marks]
Diagram 6 (e) Predict any colour change of the solution that
may occur in beakers I and II after electrolysis
(a) Name the electrodes that serve as the anode. has been carried out for an hour. [2 marks]
[2 marks]
(f) (i) Name instrument Q in the diagram.
(b) Write the formulae of all the ions present in (ii) What is the function of instrument Q?
beaker I. [2 marks] [2 marks]
6
Essay Questions
1 (a) What is meant by the term electrolysis? [2 marks] (b) Using a labelled diagram, describe an experiment
(b) Discuss in terms of ionic theory, the reasons why to show how you can electroplate an iron spoon
solid magnesium chloride (crystals) does not with another metal. In your description, give the
conduct electricity whereas molten magnesium observation and equations for the reactions that
chloride does. [4 marks] occur. [8 marks]
(c) You are supplied with magnesium chloride 3 (a) What is the difference between an electrolytic cell
crystals and all the necessary apparatus.
Describe an experiment to extract magnesium metal and a voltaic cell? [4 marks]
from magnesium chloride crystals using electrolysis.
What would you observe in this experiment? (b) You are supplied with metal P, metal Q, their
Using ionic theory, explain how the products are
formed at the cathode and the anode. nitrate salt solutions and all the necessary
[14 marks] apparatus. Metal P is higher than metal Q in the
electrochemical series and both metals have
a valency of 2. Describe an experiment to show
how you can produce an electric current from
2 (a) The products of electrolysis may be different even chemical reactions. Include a circuit diagram and
though the same type of electrolyte is used. Using show how you can detect the flow of electric
a suitable electrolyte, explain how current in your description. [12 marks]
(i) the types of electrodes, (c) Predict what will happen when a piece of metal
(ii) the concentration of ions can determine P is placed in Q nitrate solution.
the products of electrolysis of an aqueous Explain your answer. [4 marks]
solution. [12 marks]
Experiments
1 A group of students carried out three experiments to determine the products of electrolysis of sodium
hydroxide solution, potassium iodide solution and aqueous X solution using carbon electrodes.
The results of the experiment obtained is tabulated in Table 1.
Experiment Chemical substance Observation at the cathode Observation at the anode
I 0.1 mol dm–3 Colourless gas is evolved Colourless gas is evolved
which lights up a glowing
sodium hydroxide which produces a ‘pop’ sound wooden splint.
solution when a lighted wooden splint
is placed near the mouth of
the test tube.
179 Electrochemistry
Experiment Chemical substance Observation at the cathode Observation at the anode
II 0.5 mol dm–3 Colourless gas is evolved A brown solution is formed.
aqueous potassium which produces a ‘pop’ sound
iodide solution when a lighted wooden splint
is placed near the mouth of
the test tube.
III 0.5 mol dm–3 Brown deposit is formed. A brown gas is evolved which
changes blue litmus paper
aqueous X solution to red and decolourises the
litmus paper subsequently.
6 Table 1
(a) (i) In experiment I, name the products formed at the anode and cathode. [3 marks]
(ii) What factor determines the type of ions discharged at the cathode? [3 marks]
[3 marks]
(b) (i) What is the product formed at the anode in experiment II?
(ii) Suggest a test to identify the product of (i). [3 marks]
(iii) What factor determines the type of ions discharged at the anode in experiment II?
(c) Write ionic equations for the formation of the product(s)
(i) at the anode and the cathode in experiment II.
(ii) at the anode in experiment I.
(d) In experiment III, aqueous X solution is blue in colour.
(i) What is the brown deposit formed at the cathode?
(ii) Name the brown gas produced at the anode.
(iii) Suggest a chemical substance that may be X.
2 Diagram 1 and Diagram 2 show the set-ups of two electrolytic cells using copper(II) chloride solutions
of different concentrations. Plan an experiment to investigate the factor that affects the products of
electrolysis of aqueous solutions as shown in Diagrams 1 and 2.
Diagram 1 Diagram 2
Your planning should include the following aspects:
(a) Aim of experiment
(b) Statement of hypothesis
(c) All the variables
(d) List of substances and apparatus
(e) Procedure of the experiment
(f) Tabulation of data [17 marks]
Electrochemistry 180
7CHAPTER FORM 4
THEME: Interaction between Chemicals
Acids and Bases
SPM Topical Analysis
Year 2008 2009 2010 2011
Paper 1 2 31 2 31 2 31 2 3
Section ABC ABC ABC ABC
Number of questions 5 – – —21 – 5 – – —23 1 3 1 – – 1 4 1 – —21 –
ONCEPT MAP
pH scale: measurement of the H+ ion ACIDS AND BASES Concentration: units in
concentration • g dm–3
• Acids: pH < 7 • mol dm–3
• Alkalis: pH > 7 Relationship between pH values
• pH value changes with the concentration and concentration
and strength of acids/bases
Acids: compounds that produce H+ ions in water Bases: compounds that react with acids to form salts
• Strong acids: complete ionisation to form H+ ions in
and water
water • Alkalis: soluble bases that produce OH– ions in water
• Weak acids: partial ionisation to form H+ ions in • Strong alkalis: complete ionisation to form OH– ions
water in water
• Weak alkalis: partial ionisation to form OH– ions in
water
Properties of Acids: Neutralisation: Properties of Bases:
• Colour change with indicators • Reactions between acids and • Colour change with indicators
• React with bases • React with acids
• React with reactive metals bases • React with ammonium salt on
• React with metal carbonates • Uses of acids/bases and
heating
neutralisation • React with metal ions to form
• Determination of end point in
metal hydroxide
titration using acid/base indicators
or a computer interface
7.1 Characteristics and 5 Without the presence of hydrogen ions, a
Properties of Acids and substance does not show any acidic property.
Bases Dry hydrogen chloride gas, HCl(g) dissolved
in an organic solvent (such as methylbenzene),
The Meaning of Acids glacial ethanoic acid and solid ethanedioic
acid do not show any acidic property.
1 The definition of acids according to Arrhenius
6 Acids can be divided into two types: mineral
Theory: an acid is a chemical compound that acids and organic acids. Mineral acids are
obtained from minerals and most do not
produces hydrogen ions, H+ or hydroxonium contain the element carbon. Organic acids are
extracted from living things and contain the
2 Aionssu, bHst3aOn+cewhheans it dissolves in water. element carbon.
acidic properties because
7 of the formation of hydrogen ions or Table 7.1 Examples of mineral acids and organic
acids
hydroxonium ions in water.
3 Dissociation of acids in water produces Type of acid Examples
SPM hydrogen ions and anions. Examples:
’09/P1
(a) HCl(g) ⎯H⎯2O→ H+(aq) + Cl–(aq) Mineral acid Hydrochloric acid, HCl, sulphuric
acid, H2SO4 and nitric acid, HNO3
hydrogen hydrogen chloride
chloride ion ion Organic acid Ethanoic acid (CH3COOH),
methanoic acid (HCOOH),
(b) HNO3(l) ⎯H⎯2O→ H+(aq) + NO3–(aq) ethanedioic acid (H2C2O4), citric
nitric acid hydrogen nitrate acid, tartaric acid, malic acid and
ascorbic acid.
ion ion
(c) H2SO4(l) ⎯H⎯2O→ 2H+(aq) + SO42–(aq)
sulphuric hydrogen sulphate
Malic acid is found in apples.
acid ion ion Citric acid is found in citrus fruits
such as oranges.
(d) CH3COOH(l) H2O H+(aq) + CH3COO–(aq) Tartaric acid is found in grapes.
ethanoic hydrogen ethanoate Ascorbic acid is vitamin C.
Ethanoic acid is found in
acid ion ion vinegar.
Lactic acid is found in sour milk.
4 In actual fact, the hydrogen ion, H+ does not Tannic acid is found in tea
exist individually but is combined with a water leaves.
molecule (hydrated) to form a hydroxonium
ion. An acid is a substance that produces hydrogen ions
in the presence of water.
H+ + H2O → H3O+
However, H3O+ is usually written as H+(aq) in
the simplified way.
Figure 7.1 The dissociation (ionisation) of a hydrogen Hydrochloric acid, nitric acid and sulphuric acid are
chloride molecule to produce hydroxonium acids that are usually used in the school laboratories.
ion in water. Our stomachs contain hydrochloric acid that is required
for digestion of food. Aspirin, which is used as an
analgesic (a type of medicine for reducing pain), is
also a type of acid.
Acids and Bases 182
The Meaning of Bases and Alkalis 6 A chemical substance has alkaline properties
because of the formation of freely moving
1 A base is defined as a chemical substance that
can neutralise an acid to produce salt and hydroxide ions, OH– in water.
water only. For example, 7 In the presence of water, an alkali dissociates
HCl + NaOH → NaCl + H2O to hydroxide ions and cations.
acid base salt water Examples:
2 Examples of bases are metal oxides and metal (a) NH3(g) + H2O(l) → NH4+(aq) + OH–(aq)
hydroxides that contain oxide ions, O2– and ammonia ammonium hydroxide
hydroxide ions, OH– respectively. Examples:
copper(II) oxide, magnesium hydroxide. ion ion
3 The reaction between an acid and a base is ⎯H2→O K+(aq) + OH–(aq)
known as neutralisation. In neutralisation the (b) KOH(s)
O2– ions or the OH– ions of a base react with
the H+ ions of an acid to form water. potassium potassium hydroxide
O2– + 2H+ → H2O hydroxide ion ion 7
OH– + H+ → H2O
⎯H2→O Na+(aq) + OH–(aq)
4 Most bases are not soluble in water. Bases that (c) NaOH(s)
are soluble in water are known as alkalis.
sodium sodium hydroxide
5 An alkali is defined as a chemical compound
that dissolves in water to produce freely hydroxide ion ion
moving hydroxide ions, OH–.
(d) Ca(OH)2(s) ⎯H2→O Ca2+(aq) + 2OH–(aq)
calcium calcium hydroxide
hydroxide ion ion
8 A compound does not show any alkaline property
in the absence of freely moving hydroxide ions.
Examples: dry ammonia gas, ammonia gas
dissolved in organic solvent (such as propanone),
solid sodium hydroxide and solid potassium
hydroxide do not show alkaline properties.
Figure 7.2 Venn diagram for bases and alkalis Figure 7.4 The association (ionisation) of an ammonia
Bases molecule to produce a hydroxide ion
Bases that are Bases that are soluble • An alkali is a compound that produces hydroxide
insoluble in water in water (alkalis) ions in the presence of water.
examples examples • A base is a compound that neutralises an acid and
produces salt and water only.
Zinc oxide, zinc Sodium oxide, Theories on acids and alkalis:
hydroxide, copper(II) sodium hydroxide, (a) Arrhenius theory: An acid is a compound that
oxide, copper(II) potassium oxide,
hydroxide potassium hydroxide, produces hydrogen ions when it dissolves in water.
calcium hydroxide,
ammonia An alkali is a compound that produces hydroxide
Figure 7.3 Flowchart showing types and examples ions when it dissolves in water.
of bases (b) Brnsted–Lowry theory: An acid is a proton
(hydrogen ion) donor. An alkali is a proton acceptor.
183 Acids and Bases
1 ’07
Which of the following statements is true about all Comments
All bases react with acids to form salts. Only soluble
bases? bases (alkalis) dissolve in water to produce hydroxide
A React with acids ions that change red litmus paper to blue.
B Dissolve in water
C Produces hydroxide ions Answer A
D Change red litmus paper to blue
Uses of Acids, Bases and Alkalis in SPM 3 Examples of bases and their uses are given in
Our Daily Life ’08/P2, Table 7.3.
’09/P1
7
1 Acids and bases are widely used in our everyday Table 7.3 Uses of bases
life in agriculture, medicine, industry and in
the preparation of food. Base Uses
2 Examples of acids and their uses are given in Sodium hydroxide To make soaps, detergents,
bleaching agents and
Table 7.2.
fertilisers
Table 7.2 Uses of acids Ammonia To make fertilisers, nitric
acid, grease remover and
Acid Uses to maintain latex in liquid
form
Sulphuric acid To make paints,
detergents, polymers,
fertilisers, as an Calcium hydroxide To make cement, limewater
and to neutralise the acidity
electrolyte in lead-acid of soil
accumulators
Hydrochloric acid To clean metals before Magnesium To make toothpaste, gastric
electroplating hydroxide medicine (antacid)
Nitric acid To make fertilisers, Aluminium To make gastric medicine
explosive substances hydroxide (antacid)
(such as T.N.T.), dyes and
plastics
Benzoic acid To preserve food
Carbonic acid To make gassy
(carbonated) drinks
Ethanoic acid A component of vinegar
Tartaric acid To make baking powder
Cleaning agent contains ammonia
Methanoic acid is used
in the coagulation of
rubber latex
Fertilisers are made Soaps and detergents are made from sodium
from acids and alkalis hydroxide
Acids and Bases 184
7.1 SPM
’09/P2, ’10/P3, ’11/P2
To investigate the role of water in showing the properties of acids
Problem statement Conclusion Experiment 7.1 7
Is water needed for an acid to show its acidic 1 Aqueous ethanoic acid turns blue litmus paper to
properties? red, indicating its acidic property.
Hypothesis 2 Ethanoic acid in a dry condition or dissolved
in organic solvents does not show any acidic
An acid will only show its acidic properties when property.
dissolved in water.
3 Ionisation of acids will only occur in the presence
Variables of water to produce hydrogen ions which are
responsible for the acidic properties.
(a) Manipulated variable : Types of solvents-water
and propanone 4 Water is essential for the formation of hydrogen
ions which gives the acidic properties in an acid.
(b) Responding variable : Change in the colour of The hypothesis is accepted.
blue litmus
Discussion
(c) Constant variable : Type of acid and blue
litmus paper 1 In the presence of water, an acid dissociates into
hydrogen ions that cause acidity in an acid.
Apparatus
2 Dry acids do not show any acidic properties
Test tube and droppers. in the absence of water because dry acids exist
as covalent molecules. Hydrogen ions are not
Materials produced.
Glacial (dry) ethanoic acid, aqueous ethanoic acid, 3 Solvents such as methylbenzene, propanone
ethanoic acid dissolved in dry propanone and blue and trichloromethane cannot replace water for
litmus paper. an acid to show its acidic properties. This is
because an acid exists as covalent molecules in
Procedure these organic solvents; H+ ions are not produced
in these solutions.
1 A piece of dry blue litmus paper is placed in a
test tube. 4 Glacial ethanoic acid (CH3COOH) consists of
acid molecules only. CH3COOH molecule is a
2 A few drops of glacial ethanoic acid are placed covalent compound.
onto the blue litmus paper using a dropper.
5 Figure 7.5 shows the types of particles that are
3 The effect of the glacial ethanoic acid on the present in ethanoic acid dissolved in propanone
blue litmus paper is recorded. and in water.
4 Steps 1 to 3 of the experiment are repeated using
aqueous ethanoic acid and ethanoic acid dissolved
in propanone to replace glacial ethanoic acid.
5 The observations are then tabulated.
Results
Condition of Observation Inference
ethanoic acid
Glacial (dry) No noticeable Does not show
colour change any acidic
Aqueous in the litmus properties
(dissolved in paper
water) Shows acidic
Dissolved in Blue litmus properties
propanone paper has
changed to red Does not show
any acidic Figure 7.5 Particles in ethanoic acid dissolved in
No noticeable properties (a) propanone (b) water
colour change
in the litmus
paper
185 Acids and Bases
7.2
To investigate the role of water in showing the alkaline properties of alkali
Problem statement 2 The test tube must be stoppered immediately
after the red litmus paper is put in.
Is water essential for an alkali to show its alkaline
properties?
Hypothesis Results
Experiment 7.2 An alkali will only show its alkaline properties when Condition of Observation Inference
7 dissolved in water. ammonia
No colour Does not
Variables Dry change in the red show alkaline
litmus paper property
(a) Manipulated variable : Types of solvents–water Aqueous Red litmus has
and propanone (dissolved in changed to blue Shows
water) alkaline
(b) Responding variable : Change in the colour of Dissolved in No colour properties
red litmus paper propanone change in the red
litmus paper Does not
(c) Constant variable : Type of alkali and red show alkaline
litmus paper property
Apparatus Test tubes and droppers.
Materials Dry ammonia gas stoppered in a Conclusion
test tube, ammonia gas dissolved
in propanone, aqueous ammonia 1 Aqueous ammonia solution turns the red litmus
solution and red litmus paper. paper to blue, indicating its alkaline property.
Procedure 2 Dry ammonia gas or ammonia gas dissolved in
organic solvents does not show any alkaline
1 A piece of dry red litmus paper is put into a stoppered property.
test tube of dry ammonia gas and the test tube is then
stoppered back immediately (Figure 7.6). 3 An alkali shows its alkaline properties only in
the presence of water. When water is present,
2 The effect of the dry ammonia gas on the red ammonia ionises to produce OH– ions that are
litmus paper is recorded. responsible for its alkaline properties.
3 Another piece of dry red litmus paper is put in 4 Water is essential for the formation of
5 cm3 of aqueous ammonia solution in a separate hydroxide ions that cause alkalinity in an alkali.
test tube. The hypothesis is accepted.
4 Step 3 of the experiment is repeated using ammonia Discussion
dissolved in propanone to replace aqueous ammonia
solution. 1 In the presence of water, an alkali ionises to form
hydroxide ions, OH– that change red litmus
paper to blue.
2 Aqueous ammonia solution (ammonia dissolved
in water) consists of NH4+ ions, OH– ions and
NH3 molecules. An aqueous ammonia solution
is alkaline due to the presence of hydroxide ions.
Figure 7.6 Testing for the alkaline properties NH3 + H2O NH4+ + OH–
of ammonia gas
3 Dry alkalis, solid alkalis (such as solid calcium
Safety precautions
hydroxide and barium hydroxide) and alkalis
1 Ammonia gas is poisonous. This experiment
involving dry ammonia gas should be carried out dissolved in organic solvents (such as propanone)
in a fume cupboard.
do not show any alkaline properties. This
is because the alkalis do not dissociate into
hydroxide ions.
Acids and Bases 186
Glacial ethanoic acid is the pure and dry form of ethanoic acid. It is named ‘glacial’ because it appears as ice when it
solidifies below its melting point.
Chemical Properties of Acids
1 If the electrical conductivity of ethanoic acid in pro 1 Acids are sour in taste. Activity 7.1 7
panone and aqueous ethanoic acid is tested in 2 Acid solutions have pH values of less than 7.
turn, only the aqueous solution of acid conducts 3 Acids change colours of indicators as shown
electricity (light bulb is lighted up or ammeter
needle is deflected). in Table 7.4.
4 Acids can react with
2 This shows the presence of freely moving ions in
an aqueous solution of acid. (a) bases to produce salts and water,
(b) metals to produce salts and hydrogen gas,
CH3COOH(l) CH3COO–(aq) + H+(aq) (c) carbonates to produce salts, carbon
ethanoic acid ethanoate ion hydrogen ion
dioxide gas and water.
cHo+niodnusctaenldecCtrHic3iCtyOO– ions H+ ions change blue Table 7.4 Effects of acids on indicators
litmus to red
Colour of indicator
3 Dry acids do not conduct electricity. This is because Indicator in acidic solution
there are no freely moving ions. Dry acid exists as Blue litmus paper Red
covalent molecules. Universal indicator Orange and red
4 Similarly, ammonia dissolved in propanone does not
Methyl orange Red
conduct electricity. It exists as covalent molecules.
5 An aqueous ammonia solution can conduct electricity,
showing the presence of freely moving ions.
To investigate the chemical properties of acids SPM
’10/P2
Apparatus Test tube, test tube holder, spatula,
Bunsen burner, delivery tubes with
stopper and wooden splint.
Materials 1.0 mol dm–3 sulphuric acid, copper(II)
oxide, zinc powder, sodium carbonate
powder and limewater.
Procedure Figure 7.7 An acid with Figure 7.8 An acid with
a base a metal
1 A little copper(II) oxide is added to 5 cm3 of
sulphuric acid in a test tube. The mixture is Figure 7.9 An acid with a metal carbonate
heated slowly (Figure 7.7) and any changes that
occur are recorded.
2 A little zinc powder is added to 5 cm3 of dilute
sulphuric acid in a test tube. The gas evolved is
tested by placing a lighted wooden splint near
the mouth of the test tube (Figure 7.8).
3 A little sodium carbonate powder is added to 5 cm3
of dilute sulphuric acid in a test tube. The gas
evolved is tested with limewater (Figure 7.9).
187 Acids and Bases
Results Observation Inference
Test on acid • Black powder dissolved • Copper(II) salt solution is
Heating with copper(II) • Blue solution is formed formed
oxide
• Hydrogen gas is produced
Test with zinc powder • Effervescence occurred • A salt solution is formed
Test with sodium carbonate • Gas produced a ‘pop’ sound when it is
• Carbon dioxide gas is produced
7 tested with a lighted wooden splint • A salt solution is formed
• Zinc powder dissolved
• Effervescence occurred
• Gas evolved turned limewater milky
• White solid of sodium carbonate
dissolved
Discussion (b) Magnesium dissolves in ethanoic acid
1 A dilute acid reacts with a base to produce salt to form a salt, magnesium ethanoate and
hydrogen gas.
and water only.
Mg(s) + 2CH3COOH(aq) →
Acid + base → salt + water Mg(CH3COO)2(aq) + H2(g)
If the salt solution is evaporated until saturated, 3 A dilute acid will react with a metal carbonate to
salt crystals will form upon cooling. produce a salt, carbon dioxide gas and water.
Examples:
(a) Black copper(II) oxide powder (a base) acid + metal carbonate →
salt + carbon dioxide + water
dissolves in dilute sulphuric acid to produce a
salt, copper(II) sulphate (blue colour) and water. Examples:
(a) Sodium carbonate reacts with dilute sulphuric
CuO(s) + H2SO4(aq) → CuSO4(aq) + H2O(l)
acid to produce a salt, sodium sulphate, carbon
(b) Copper(II) oxide dissolves in ethanoic acid to dioxide gas and water.
form a salt, copper(II) ethanoate and water.
CuO(s) + 2CH3COOH(aq) → Na2CO3(s) + H2SO4(aq) →
Cu(CH3COO)2(aq) + H2O(l) Na2SO4(aq) + CO2(g) + H2O(l)
(c) Nitric acid reacts with sodium hydroxide (an (b) Calcium carbonate reacts with dilute
alkali) to produce a salt, sodium nitrate and hydrochloric acid to produce a salt, calcium
water. chloride, carbon dioxide gas and water.
HNO3(aq) + NaOH(aq) → NaNO3(aq) + H2O(l) CaCO3(s) + 2HCl(aq) →
CaCl2(aq) + CO2(g) + H2O(l)
2 A dilute acid will react with a reactive metal to
produce a salt and hydrogen gas. Conclusion
1 Sulphuric acid reacts with a base (copper(II)
acid + reactive metal → salt + hydrogen
oxide) to produce salt and water.
Examples: 2 Sulphuric acid reacts with a reactive metal (zinc)
(a) Zinc dissolves in sulphuric acid to form a
to produce a salt and hydrogen gas.
salt, zinc sulphate and hydrogen gas. 3 Sulphuric acid reacts with a metal carbonate
Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g) (sodium carbonate) to produce a salt, water and
carbon dioxide gas.
Acids and Bases 188
Chemical Properties of Alkalis
Non-reactive metals such as copper and silver do 1 Alkalis are bitter in taste and feel soapy.
not react with dilute acid. Very reactive metals such 2 Alkaline solutions have pH values of more
as sodium and potassium will react with dilute acid
vigorously and may produce an explosion. than 7.
3 Alkalis change the colours of indicators as
2 ’01
shown in Table 7.5 below.
Which of the following compounds reacts with
limestone powder to produce a gas that turns lime- Table 7.5 Effects of alkalis on indicators
water milky?
Indicator Colour of indicator in
A Nitrogen dioxide gas alkaline solution
B Hydrogen chloride gas dissolved in tetra
Red litmus paper Blue
chloromethane
C Sulphur dioxide gas dissolved in propanone Universal indicator Blue or purple Activity 7.2 7
D Sulphur dioxide gas dissolved in water
Methyl orange Yellow
Comments
An acidic gas must first dissolve in water before 4 An alkali reacts with an acid to produce salt
reacting with calcium carbonate (limestone) and water. For example:
powder to produce carbon dioxide gas which turns
limewater milky. KOH(aq) + HCl(aq) → KCl(aq) + H2O(l)
Answer D 5 When an alkali is heated with an ammonium
salt, ammonia gas is produced. For example:
Generally,
(a) metal oxides and metal hydroxides are basic. For NH4+(aq) + OH–(aq) → NH3(g) + H2O(l)
example: from ammonium salt from alkali
MgO + H2O → Mg(OH)2
6 An aqueous alkali forms metal hydroxide as
magnesium oxide magnesium hydroxide precipitate when added to an aqueous salt
(b) non-metal oxides are acidic. For example, SO2, solution. For example:
NO2 or CO2. Cu2+(aq) + 2OH–(aq) → Cu(OH)2(s)
SO2 + H2O → H2SO3
from copper(II) from alkali copper(II) hydroxide
sulphur dioxide sulphurous acid salt solution as blue precipitate
To investigate the chemical properties of alkalis
Apparatus changes that occur are recorded.
2 A little ammonium chloride powder is added to 5
Test tubes, test tube holder, spatula, Bunsen burner,
delivery tubes with stopper and red litmus paper. cm3 of sodium hydroxide solution in a test tube.
The mixture is heated gently. The gas evolved is
Materials tested with a piece of damp red litmus paper.
3 5 cm3 of sodium hydroxide solution is added to
2.0 mol dm–3 sodium hydroxide solution, benzoic 5 cm3 of iron(III) sulphate solution in a test tube.
acid powder, ammonium chloride powder, 1.0 mol Any changes that occur are recorded.
dm–3 iron(III) sulphate solution.
Procedure
1 A little benzoic acid powder is added to 5 cm3
of sodium hydroxide solution in a test tube. Any
189 Acids and Bases
Results NaOH(aq) + C6H5COOH(s) →
C6H5COONa(aq) + H2O(l)
Test on sodium Observation Inference
hydroxide 2 If the salt solution is evaporated in an evaporating
dish until a saturated solution is formed, white
With benzoic White powder A salt solution crystals of sodium benzoate will be crystallised
acid powder dissolves and is formed upon cooling.
added a colourless
solution is 3 When sodium hydroxide is heated with ammonium
formed chloride (an ammonium salt), ammonia gas is
produced.
Heating with A pungent gas Ammonia gas
ammonium that turns damp is produced NH4Cl(s) + NaOH(aq) →
chloride red litmus paper NaCl(aq) + H2O(l) + NH3(g)
7 powder blue is evolved
In this reaction, ammonium ions react with
With iron(III) A brown Iron(III) hydroxide ions to produce ammonia gas. The ionic
hydroxide is equation for this reaction is
sulphate precipitate is formed
NH4+(aq) + OH–(aq) → NH3(g) + H2O(l)
solution added formed
4 Sodium hydroxide solution dissociates to
Conclusion hydroxide ions in water.
1 Sodium hydroxide reacts with benzoic acid to NaOH → Na+ + OH–
produce salt and water.
Hydroxide ions combine with iron(III) ions from
2 When sodium hydroxide is heated with iron(III) sulphate solution to form insoluble
ammonium chloride, ammonia gas which turns iron(III) hydroxide as a brown precipitate.
red litmus to blue is produced.
Fe3+(aq) + 3OH–(aq) → Fe(OH)3(s)
3 Sodium hydroxide solution reacts with an
aqueous iron(III) solution to produce a brown from iron(III) from sodium iron(III) hydroxide as
precipitate, iron(III) hydroxide. sulphate hydroxide brown precipitate
Discussion
1 Sodium hydroxide as an alkali reacts with benzoic
acid, C6H5COOH to produce a salt, sodium
benzoate and water in a neutralisation reaction.
Basicity of Acids CH3COOH(aq) CH3COO–(aq) + H+(aq)
1 Basicity of an acid is the number of moles of Three H atoms bonded to Only one H atom
OH– ions that are required to react with one carbon do not dissociate dissociates to form H+ ion
mole of the acid.
4 A diprotic acid (or dibasic acid) is an acid
2 Since one mole of OH– ions reacts with one that will produce two moles of H+ ions from
mole of H+ ion, the basicity of an acid is also one mole of the acid in water.
the number of moles of H+ ion that can be
produced by one mole of the acid when it For example:
dissolves in water.
H2SO4(aq) → 2H+ (aq) + SO42–(aq)
3 A monoprotic acid (or monobasic acid) is
an acid that will produce one mole of H+ 1 mol sulphuric 2 mol sulphate ion
ion when one mole of the acid dissolves in acid hydrogen ions
water.
For example, although ethanoic acid has four 5 A triprotic acid (or tribasic acid) is an acid
hydrogen atoms in the molecule, only one of
the hydrogen dissociates to form H+ ion in that will produce three moles of H+ ions
water.
from one mole of the acid in water.
For example:
H3PO4(aq) 3H+(aq) + PO43–(aq)
1 mol 3 mol phosphate ion
phosphoric acid hydrogen ions
Acids and Bases 190
Table 7.6 Examples of monoprotic acid and diprotic acid
Examples of Hydrochloric acid (HCl), nitric acid (HNO3), Basicity of an acid is not the same
monoprotic acid ethanoic acid (CH3COOH) and methanoic acid as the number of H atoms in the
(HCOOH) formula of the acid.
Examples of Sulphuric acid (H2SO4), ethanedioic acid (H2C2O4), Basicity is the number of moles of
diprotic acid carbonic acid (H2CO3) and chromic acid (H2CrO4) H+ ions produced by one mole of
acid in water.
7.1 3 Identify the chemicals Q, R, X, Y and gas Z in the 7
following reactions:
1 (a) Explain what you understand by the term (a) H2SO4 + Q → MgSO4 + H2O + CO2
(i) an acid (b) Ca(OH)2 + 2R → Ca(NO3)2 + 2H2O
(ii) a base (c) 2Al + 6X → 2AlCl3 + 3H2
(iii) an alkali ( d) Y + NH4NO3 ⎯he⎯at→ KNO3 + H2O + Z
(b) What is the effect of an acid and an alkali on 4 Write equations to show the reactions between
moist litmus paper? (a) sulphuric acid and magnesium oxide
(b) nitric acid and aluminium metal
2 Identify the correct uses of the following acids and (c) hydrochloric acid and calcium carbonate
bases. (d) ethanoic acid and sodium hydroxide
(e) potassium hydroxide and ammonium chloride
Acids or bases: MHg2(SOOH4,)2,HNNHO3,3,NaCOaHO, Ca(OH)2, when heated
Uses Acids or bases 5 Effervescence occurs when magnesium powder is
To make antacid added to aqueous hydrochloric acid. However, no
To make fertiliser noticeable change takes place when magnesium
To make soap powder is added to hydrogen chloride dissolved in
To neutralise acidity in soil methylbenzene. Explain why.
Acid Alkali
dissolves in water dissolves in water
produces H+ ions produces OH– ions
reacts with reacts with
carbonate metal base acid ammonium salt metal ions
heat metal
salt salt + hydrogen salt + water
+ ammonia gas hydroxides
carbon dioxide
+ Acids and Bases
water
191
7.2 The Strength of Acids • pH > 7 ⇒ alkaline solution
and Alkalis 3 pH is actually a measurement of the concen
The pH Scale tration of hydrogen, H+ ions in a solution.
4 The higher the concentration of the H+ ions,
1 The pH scale is a set of numbers used to
indicate the degree of acidity or alkalinity of the lower the pH value and the more acidic
a solution.
the solution.
2 The values of the pH scale range from 0 to 14. 5 The higher the concentration of the OH–
• pH < 7 ⇒ acidic solution
• pH = 7 ⇒ neutral solution ions, the higher the pH value and the more
alkaline the solution.
6 The relationship between the pH scale, acidity
or alkalinity and concentration of H+ ions is
shown below.
Activity 7.3 • All acids have pH < 7. • All alkalis have pH > 7.
7 • The lower the pH value, the higher the H+ ion • The higher the pH value, the higher the OH– ion
concentration. concentration.
Measurement of pH Value of a Solution 4 A pH meter is an
1 The pH value of a solution can be measured electric meter that is
by using
(a) universal indicator or pH paper used to measure the
(b) a pH meter (with or without a computer
interface) pH value of a solution
2 Universal indicator is a mixture of indicators accurately. A pH meter
that gives different colours corresponding to
different pH values as shown in Table 7.7. will show the pH
3 Universal indicator is used in the form of value when its probe is
(a) solution, or
(b) paper strips (also known as pH paper). immersed in a solution pH meter
to be tested.
5 With a computer interface,
the exact pH value can be displayed on the
computer screen when the pH meter is placed
in the solution.
Table 7.7 Colours of universal indicator
pH value 0, 1, 2 3 4 5 6 7 8 9 10 11 12, 13, 14
Colour purple
red orange orange orange yellow green greenish- blue blue bluish-
red yellow blue purple
To measure the pH values of some solutions used in daily life
Apparatus Materials
Soap solution, carbonated drink, tap water, orange
Beakers, universal indicator solution, dropper, fruit juice, distilled water, milk, tea, dilute sodium
standard colour chart of universal indicator. hydroxide and hydrochloric acid.
Acids and Bases 192
Procedure
1 About 10 cm3 of soap solution is placed in a small beaker.
2 Two drops of universal indicator solution are added to the soap solution. The solution is then stirred.
3 The colour of the solution produced is matched against the standard colour chart of universal indicator. The
corresponding pH value of the colour is noted and recorded.
4 The experiment is repeated using carbonated drink, tap water, orange fruit juice, distilled water, milk, tea,
dilute sodium hydroxide and hydrochloric acid in place of the soap solution.
Results
Solution Soap Carbonated Tap Orange Distilled Milk Tea Dilute Dilute
solution drink water juice water sodium hydrochloric
hydroxide acid
pH value 10 5 64 7 6 5 13 1 7
Conclusion
1 Different solutions have different pH values.
2 The pH value of a solution can be measured using the universal indicator solution.
Degree of Dissociation HCl → H+ + Cl–
HNO3 → H+ + NO3–
1 The strength of an acid or an alkali depends H2SO4 → 2H+ + SO42–
on the degree of dissociation (also known as
the degree of ionisation). (The one–way arrow → indicates complete
dissociation)
2 The degree of dissociation measures the 3 Complete dissociation (100%) in water by a
percentage or fraction of molecules that strong acid produces a high concentration of
dissociates into ions when dissolved in water. H+ ions and hence a low pH.
4 Weak acids are chemicals that dissociate
3 For example, the degree of dissociation of partially (incomplete dissociation) into
hydrochloric acid is 100% or 1. This means hydrogen ions H+ in water.
5 Most of the organic acids such as ethanoic
that all the hydrogen chloride molecules in acid, ethanedioic acid, methanoic acid, citric
acid and tartaric acid are weak acids.
hydrochloric acid will ionise to form H+ ions
CH3COOH CH3COO– + H+
and Cl– ions when dissolved in water. H2C2O4 2H+ + C2O42–
4 In a 1.0 mol dm–3 aqueous ethanoic solution,
only 4 out of 1000 molecules of ethanoic acid
dissociate to form ions. Degree of dissociation
of a 1.0 mol dm–3 aqueous ethanoic solution is
—1—04—0—0 = 0.004 or 0.4%.
5 Acids can be divided into 2 categories: strong (The two–way arrow indicates reversible
acids and weak acids, depending on their reaction)
degree of dissociation. Examples of weak acids
6 Alkalis can be divided into 2 categories: strong Ethanoic acid, CH3COOH
Methanoic acid, HCOOH
alkalis and weak alkalis, depending on their Ethanedioic acid, H2C2O4
Carbonic acid, H2CO3
degree of dissociation. Phosphoric acid, H3PO4
Chromic acid, H2CrO4
Strong and Weak Acids SPM Nitrous acid, HNO2
Sulphurous acid, H2SO3
’08/P2,
’09/P1,
’10/P3
1 A strong acid is a chemical substance that
dissociates completely (degree of dissociation
is 100%) into hydrogen ions, H+ in water.
2 Mineral acids such as hydrochloric acid, nitric
acid and sulphuric acid are strong acids.
193 Acids and Bases
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