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Electrolysis/Electrochemistry


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electrolysis


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Electrochemistry2

Summary

Introduction

 Electrolysisdecomposition of compound using electricity
 Electrolytean ionic compound which conducts electric current in molten or aqueous solution, being decomposed in the process.
 Electrodea rod or plate where electricity enters or leaves electrolyte during electrolysis.
Reactions occur at electrodes.
 Discharge the removal of electrons from negative ions to form atoms or the gain of electrons of positive ions to become atoms.
 Anode
  • positive electrode connected to positive terminal of d.c. source.
  • Oxidation occurs here.
  • Anode loses negative charge as electrons flow towards the battery, leaving anode positively charged.
  •  This causes anion to discharge its electrons here to replace lost electrons and also, negative charge are attracted to positive charge.
 Cathode
  •  negative electrode connected to negative terminal of d.c. source.
  • Reduction occurs here.
  • Cathode gains negative charge as electrons flow from the battery towards the cathode, making cathode negatively charged.
  • This causes cation to be attracted and gains electrons to be an atom.
 Anion
  •  negative ion
  • attracted to anode.
 Cation
  •  positive ion
  • attracted to cathode.

 Non-electrolytesWeak electrolytes Strong electrolytes 
 Organic liquids or solutions
Weak acids and alkalisStrong acids, alkalis and salt solutions 
ethanol C2H5OH
tetrachloromethane CCl4
trichloromethane CHCl3
pure water H2O
sugar solution C12H22O11
molten sulphur S
limewater Ca(OH)2
ammonia solution NH3
aqueous ethanonoic acid CH3COOH
aqueous sulphurous acid H2SO3
aqueous carbonic acid H2CO3
aqueous sulphuric acid H2SO4
aquous nitric acid HNO3
aquous hydrochloric acid HCl
aqueous potassium hydroxide KOH
aqueous sodium hydroxide NAOH
copper(II) sulphate solution CuSO4

Electrolysis of Molten Compounds

  • Molten/aqueous ionic compounds conduct electricity because ions free to move.
  • In solid state, these ions are held in fixed position within the crystal lattice.
  • Hence solid ionic compounds do not conduct electricity.
  • When molten binary compound is electrolysed, metal is formed on cathode while non-metal is formed on anode.

Example

Electrolysis of molten PbBr2

To make molten lead(II) bromide, PbBr2, we strongly heat the solid until it melts. To electrolyse it, pass current through the molten PbBr2.

What happens:

Ions present: Pb2+ and Br-

Reaction at Anode

Br- loses electrons at anode to become Br atoms. Br atoms created form bond together to make Br2 gas.

2Br-(aq) --> Br2(g)+ 2e-

Reaction at Cathode

Pb2+ gains electrons at cathode to become Pb atoms becoming liquid lead (II).
Pb2+(aq) + 2e- --> Pb(l)

Overall equation

PbBr2(l) --> Pb(l) + Br2(g)

Electrolysis of Aqueous Solution

  • Aqueous solutions contain additional H+ and OH- ions of water, totaling 4 ions in the solution :
    • 2 from electrolyte, 2 from water.
    • Only 2 of these are discharged.
  • Electrolysis of aqueous solutions use the theory of selective discharge.
At cathode
  • In CONCENTRATED solutions of nickel/lead compound, nickel/lead will be discharged instead of hydrogen ions of water which is less reactive than nickel/lead.
  • In VERY DILUTE solutions, hydrogen, copper and silver ions are preferable to be discharged, according to its ease to be discharged.
  • Reactive ions (potassium, sodium, calcium, magnesium, aluminium) will NEVER BE DISCHARGED in either concentrated or dilute condition. Instead, hydrogen ions from water will be discharged at cathode.
At anode
  • In CONCENTRATED solutions, iodine/chlorine/bromine ions are preferable to be discharged, although it’s harder to discharged compared to hydroxide ions.
  • In VERY DILUTE solutions containing iodide/chloride/bromide ions, hydroxide ions of water will be discharged instead of iodide/chloride/bromide, according to ease of discharge.
  • Sulphate and nitrate are NEVER DISCHARGED in concentrated/dilute solutions.

Examples

A. Concentration Solutions

Electrolysis of Concentrated NaCl


What happens:

Ions Present: Na+, H+, OH- and Cl-

Reaction at Anode

  • Cl- loses electrons at anode to become Cl atoms, although OH- is easier to discharge.
  • Cl atoms created form covalent bond together to make Cl2 gas.
  • 2Cl- (aq) --> Cl2 (g) + 2e-
Reaction at Cathode
  • H+ gains electrons at cathode to become H atoms becoming hydrogen gas
  • 2H+ (aq) + 2e- --> H2 (l)
Overall Equation

2HCl(l) --> H2(l) + Cl2(g)

Note: any cation and anion left undischarged in solution forms new bonds between them.
E.g. in above, leftovers Na
+ and OH- combine to form NaOH.

B. Very Dilute Solutions

Electrolysis of Dilute H2SO4

What happens:

Ions Present:
H+, OH- and SO42-

Reaction at Anode
  • OH- loses electrons at anode to become O2 and H2O.
  • 4OH- (aq) --> O2 (g) + 2H2O (l) + 4e-
Reaction at Cathode
  • H+ gains electrons at cathode to become H atoms becoming hydrogen gas.
  • 2H+(aq) + 2e- --> H2 (g)
Overall Equation
  • Both equations must be balanced first.
  • The cathode equation is short 2 electrons. Hence, we should first even them by multiplying cathode equation by 2.
    • (2H+(aq) + 2e- --> H2 (g)) x 2 = 4H +(aq) + 4e- --> 2H2 (g)
  • Now we can combine the equations, forming: 
    • 4H+ (aq) + 4OH+ (aq) --> 2H2 (g) + O2 (g) + 2H2O (l)
  • 4H+ and 4OH+ ions, however, combine to form 4H2O molecules.
    • Hence: 4H2O (l) --> 2H2 (g) + O2 (g)+ 2H2O (l)
  • H2O molecules are formed on both sides.
    • Therefore, they cancel the coefficients: 2H2O (l) --> 2H2 (g) + O2 (g)
  • Since only water is electrolysed, the sulfuric acid now only becomes concentrated.

Electrolysis using different types of electrodes

  • Inert Electrodes are electrodes which do not react with electrolyte or products during electrolysis.
    • Eg. platinum and graphite.
  • Active Electrodes are electrodes which react with products of electrolysis, affecting the course of electrolysis.
    • Eg. copper.
A. Electrolysis of CuSO4 Using Inert Electrodes (e.g. carbon)

What happens:

Ions Present
: Cu2+, H+, OH- and SO42-

Reaction at Anode
  • OH- loses electrons at anode to become O2 and H2O.
  • 4OH- (aq) --> O2 (g) + 2H2O (l) +4e-
Reaction at Cathode
  • Cu2+ gains electrons at cathode to become Cu atoms becoming liquid copper.
  • Hydrogen ions are not discharged because copper is easier to discharge.
  • Cu2+ (aq) + 2e- --> Cu (s)
Overall Equation
  • Both equations must be balanced first.
  • The cathode equation is short 2 electrons. Hence, we should first even them by multiplying cathode equation by 2.
    • (Cu2+ (aq) + 2e- --> Cu (s)) x 2 = 2Cu2+ (aq) + 4e- --> 2Cu (s)
  • Now we can combine the equations, forming:
    • 2Cu(OH)2 (aq) --> 2Cu (s) + O2 (g) + 2H2O (l)
  • Since copper ions in solution are used up, the blue colour fades.
  • Hydrogen and sulphate ions left forms sulphuric acid.
B. Electrolysis of CuSO4 Using Active Electrodes (e.g. copper)

Ions Present
: Cu2+, H+, OH- and SO42-

Reaction at Anode
  • Both SO42- and OH- gets attracted here but not discharged. Instead, the copper anode discharged by losing electrons to form Cu2+. So, the electrode size decreases.
  • Cu (s) --> Cu2+ (aq) + 2e-
Reaction at Cathode
  • Cu2+ produced from anode gains electrons at cathode to become Cu atoms becoming copper. Hence, the copper is deposited here and the electrode grows.
  • Cu2+ (aq) + 2e- --> Cu (s)
Overall Change
  • There is no change in solution contents as for every lost of Cu2+ ions at cathode is replaced by Cu2+ ions released by dissolving anode.
  • Only the cathode increases size by gaining copper and anode decreases size by losing copper.
  • We can use this method to create pure copper on cathode by using pure copper on cathode and impure copper on anode.
  • Impurities of anode falls under it.

Electroplating

  • Electroplating is coating an object with thin layer of metal by electrolysis. This makes the object protected and more attractive.
  • Object to be plated is made to be cathode and the plating metal is made as anode.
  • The electrolyte MUST contain plating metal cation.
Plating Iron object with Nickel

Reaction at Anode
  • Ni2+ discharged from anode into solution. So, the electrode size decreases.
  • Ni (s) --> Ni2+ (aq) + 2e-
Reaction at Cathode
  • Ni2+ produced from anode gains electrons at cathode to become Ni atoms becoming nickel. Hence, the nickel is deposited here and the electrode grows.
  • Ni2+ (aq) + 2e- --> Ni (s)
Overall Change
  • There is no change in solution contents while iron object receives nickel deposit.
  • Uses of Electroplating

Creation of Electric Cells by Electrolysis

  • A Simple cell or an Electric cell is a device that converts chemical energy into electrical energy, and it consists of 2 electrodes made of 2 metals of different reactivity.
  • In a simple cell, the MORE REACTIVE metal/electrode is ALWAYS designated the NEGATIVE electrode
  • The anode (negative electrode) is made of more reactive metal. This is because they have more tendency of losing electrons. 
  • The cathode (positive electrode) is made of less reactive metal.
  • The further apart the metals in the reactivity series, the higher the voltage created.
  • The electrons in a simple cell will ALWAYS flow from the NEGATIVE electrode (made of the MORE reactive metal) to the POSITIVE electrode.
Eg. A simple electric cell using zinc and copper


Observationbubbles of hydrogen gas appear at the copper rod.

Explanation: Zinc is more reactive than copper. Thus, it is more electropositive than copper, meaning that zinc loses electrons more easily than copper. As a result, oxidation occurs at the zinc rod (the anode) and zinc metal loses electrons to become zinc ions, that is, Zn (s) - 2e- --> Zn2+(aq)

The electrons then flow from the zinc rod to the copper rod through the external circuit. At the copper rod, reduction occurs - the hydrogen ions in solution accept these electrons to form hydrogen gas; 
2H+(aq) + 2e- --> H2 (g) 
This explains why bubbles of gas are produced at the copper rod when the two rods are connected by a wire.

The magnitude of the voltage (potential difference) is related to the positions of the two metals in the reactivity series. The further apart the two metals, the larger will be the potential difference (voltage) produced.

Electrolytic cell vs Electrochemical cell

Electrolytic cellElectrochemical cell 
Converts chemical energy into electrical energyConverts electrical energy into chemical energy
Redox reaction is spontaneous and is responsible for the production of electrical energyRedox reaction is not spontaneous and electrical energy has to be supplied to initiate the reaction
The two half-cells are set up in different containers which are connected through a salt bridge or porous partitionBoth the electrodes are placed in the same container in a solution of molten electrolyte.
The anode is the negative electrode while the cathode is the positive electrode.
The reaction at the anode is oxidation and the reaction at the cathode is reduction.
The anode is the positive electrode while the cathode is the negative electrode.
The reaction at the anode is oxidation and the reaction at the cathode is reduction.
Electrons are supplied by the species getting oxidized.
They move from anode to the cathode in the external circuit.
An external battery supplies the electrons, which enter through the cathode and come out through the anode.

Factors affecting electrolysis

  • Concentration
  • Type of electrode
Concentration
  • If the concentration of a particular ion is high, then this can alter the preferential discharge
  • If dilute hydrochloric acid is electrolysed, hydrogen gas is given off at the cathode and oxygen gas at the anode. However, when concentrated hydrochloric acid is electrolysed, hydrogen gas is still given off at the cathode, but chlorine gas is given off at the anode.
  • This is because although the chloride ion is harder to discharge than the hydroxide ion, its high concentration makes it more likely to be discharged.
Type of electrode

Eg. electrolysis of aqueous copper(II) sulphate solution

Use Carbon Electrodes:
  • Carbon electrodes are inert and so do not affect the electrolysis
  • At the anode, we have a choice of sulphate or hydroxide ions, and hydroxide ions are easier to discharge so oxygen gas is given at the anode
  • 4OH- (aq) + O2 (g) ---> O2 (g) + 2H2O (l) + 4e-
  • At the cathode, we have a choice of copper or hydrogen ions. Copper ions are easier to discharge so we will see a pink deposit of copper metal on the carbon electrode
  • Cu2+ (aq) + 2e- ---> Cu (s)
Use Copper Electrodes
  • Copper electrodes are active and so will affect electrolysis
  • At the anode, the copper electrode dissolves into solution
  • Cu (s) ---> Cu2+ (aq) + 2e-
  • At the cathode, the copper ions are deposited as pink copper metal
  • Cu2+ (aq) + 2e- ---> Cu (s)

MCQ Questions

1. Electricity can pass through molten lead(II) bromide because of the presence of
a. free electrons
b. moveable ions
c. moveable atoms
d. lead metal

2. When a dilute salt water is electrolysed, a colorless gas is given off at the anode. The gas is
a. hydrogen
b. steam
c. oxygen
d. chlorine

3. A solution of copper(II) sulphate is electrolysed, using carbon electrodes. The pinkish deposit which forms on one of the electrodes is
a. copper
b. copper(I) oxide
c. copper(II) oxide
d. copper(III) sulphide

4. A solution of copper(II) sulphate is electrolysed, using copper electrodes. Which of the following would happen?
a. the anode loses weight
b. the cathode loses weight
c. the solution darkens in color
d. the solution lightens in color

5. An electrolyte is always
a. an acid or alkali
b. an aqueous solution
c. a liquid
d. a molten solid

6. Anions are formed by
a. metals gaining electrons
b. metals losing electrons
c. non-metals gaining electrons
d. non-metals losing electrons

7. Which of these anions is never discharged at the positive electrode during electrolysis?
a. NO3-
b. OH-
c. I-
d. O2-

8. In the electrolytic manufacture of aluminium, what is the anode made of?
a. copper
b. graphite
c. platinum
d. steel

9. In which electrolyte would a carbon cathode increase in mass during electrolysis?
a. aqueous copper(II) sulphate
b. concentrated hydrochloric acid
c. concentrated aqueous sodium chloride
d. dilute sulphuric acid

10. Chlorine is manufactured commercially by the electrolysis of aqueous sodium chloride (brine). Which other important products are made in the process?
a. hydrochloric acid and hydrogen
b. hydrogen and sodium
c. hydrogen and sodium hydroxide
d. sodium and sodium hydroxide

11. An electric current is passed through aqueous potassium sulphate, K2SO4.
What is formed at the cathode (negative electrode)?
a. hydrogen
b. oxygen
c. potassium
d. sulphur

12. What happens when molten lead(II) chloride is electrolysed?
a. chloride ions gain electrons at the cathode
b chloride ions lose electrons at the anode
c. lead(II) ions lose electrons at the cathode
d. lead(II) ions move towards the anode

13. Which element is liberated at a carbon cathode when aqueous sodium chloride is electrolysed?
a. chlorine
b. hydrogen
c. oxygen
d. sodium

14. Which change always takes place when aqueous copper(II) sulphate is electrolysed?
a. copper is deposited at the negative electrode
b. oxygen is evolved at the positive electrode
c. sulphate ions move towards the negative electrode
d. the color of the solution fades

15. Which element is liberated at the cathode by the electrolysis of an aqueous solution containing its ions?
a. bromine
b. chlorine
c. hydrogen
d. oxygen

16. Aqueous copper(II) sulphate is electrolysed using copper electrodes. Which observations will be made?

  at anode (positive)
 at cathode (negative)
a
 anode dissolves
 pink solid forms
b
  anode dissolves  pink solid forms
c
 color gas forms
 color gas forms
d
 color gas forms  pink solid forms

17. Why is cryolite, Na3AlF6, used in the extraction of aluminium from aluminium oxide?
a. to dissolve aluminium oxide
b. to prevent the anodes from burning away
c. to prevent the oxidation of the aluminium
d. to remove impurities from the aluminium oxide

18. When sodium chloride was electrolysed, sodium was produced at the negative electrode. In which form was the sodium chloride during the electrolysis?
a. concentrated aqueous solution
b. dilute aqueous solution
c. molten
d. solid

19. In which instance is there no change in the concentration of the solution during electrolysis?
a. concentrated sodium chloride solution between carbon electrodes
b. copper(II) sulfate solution between copper electrodes
c. copper(II) sulfate solution between platinum electrodes
d. dilute sodium chloride solution between platinum electrodes

20. An example of a weak electrolyte is
a. alcohol
b. salt solution
c. sugar solution
d. ammonia solution 

21. Electroplating iron with zinc is called galvanising. The reaction at the cathode is shown by the equation
a. Fe (s) ---> Fe2+ (aq) + 2e-
b. Fe2+ (aq) + 2e- ---> Fe (s)
c. Zn (s) ---> Zn2+ (aq) + 2e-
d. Zn2+ (aq) + 2e- ---> Zn (s)

22. The circuit shown below was set up, with brass as the anode.
Which electrode reactions will occur on closing the switch?
            Anode reaction                                  Cathode reaction
a.    Copper dissolves preferentially.                Copper is deposited.
b.    Copper dissolves preferentially.                Hydrogen is evolved.
c.    Zinc dissolves preferentially.                    Hydrogen is evolved.
d.    Zinc and copper both dissolve.                 Copper is deposited.      

23. During the electrolysis of concentrated sodium chloride in a cell, chlorine, hydrogen, and sodium hydroxide are produced. What is the molar ratio of these products?
        Chlorine        Hydrogen        Sodium hydroxide
a.        1                    1                    1
b.        2                    1                    2
c.        2                    1                    1
d.        2                    2                    1                                         

Answers

1. b
2. c
3. a
4. a
5. c
6. c
7. a
8. b
9. a
10. c
11. a (H+ and K+ ions in the electrolyte migrate to the cathode. H+ are preferentially discharged to form hydrogen gas because it is lower down in the electrochemical series than K+ ions)
12. b (the negative chloride ions will migrate to the anode and become oxidised at the anode to form chlorine gas)
13. b (the ions attracted to the cathode are H+ and Na+ ions. H+ is preferentially discharged to form hydrogen gas)
14. b
15. c
16. b
17. a
18. c
19. b
20. d
21. d
22. c
23. b

Structured Question Worked Solutions

1. Dilute sulphuric acid will conduct an electric current.
a. Give the formulae of all of the ions present in dilute sulphuric acid

b. Name the gaseous products which you would expect to be formed during the electrolysis of aqueous potassium sulphate using inert electrodes
at the anode:_____
at the cathode:______

c. Name a metal which is used to electroplate
i. bicycle handlebars
ii. teaspoon

d. Explain why a metal such as aluminium can conduct an electric current but a non-metal such as sulphur cannot conduct a current


Solution

a. H+, OH-, SO42-

b. cathode: Hydrogen
anode: Oxygen

ci. chromium
cii. silver

d. Aluminium consists of positively charged particles in a sea of electrons. The electrons are able to move freely and thus electricity can flow. In sulphur, the atomic arrangement is fixed, so there is no movement of electrons. When electricity is passed through sulphur, electricity will not be able to flow.

2a. When concentrated aqueous sodium chloride is electrolysed using graphite electrodes, hydrogen is collected at the cathode and chlorine at the anode.

When concentrated aqueous sodium chloride is electrolysed using iron electrodes, hydrogen is again collected at the cathode but much less chlorine is collected at the anode.

i. Give the equations for the electrode reactions by which hydrogen and chlorine are formed

ii. Explain why much less chlorine is collected when iron electrodes are used.

iii. Name the product, other than hydrogen and chlorine, which is manufactured by the electrolysis of concentrated aqueous sodium chloride. Give a major use of this product

b. Why is the electrolysis of concentrated hydrochloric acid not used for the manufacture of chlorine?

Answers

ai. 2H+ (aq) + 2e-  ----->  H2 (g)
2Cl- (aq) -----> Cl2 (g) + 2e-

ii. When iron anode is used, some oxygen gas is produced at the same time. Some of the electrical energy is used to liberate oxygen. So less Cl2 is produced.

iii. Sodium hydroxide. It is used to manufacture soap

b. because concentrated HCl is not a cheaply and readily available raw material. It is also a volatile acid. A lot of HCl gas will be emitted.

3.
The diagram shows the electrolytic cell used to produce aluminium. The electrolyte contains aluminium oxide and cryolite (sodium aluminium fluoride) and is molten at about 800oC. The electrodes are made from graphite.

a. Why is a mixture of cryolite and aluminium oxide, and not pure aluminium oxide, used as the electrolyte?

b. Write the equations for the reactions occurring at
i. the positive electrode
ii. the negative electrode

c. Explain why the graphite anodes need to be replaced at regular intervals

d. Calculate the maxiumum mass of aluminium that can be made from 408 tonnes of aluminium oxide.

ei. Aluminium foil is used to make food containers because it does not corrode easily. Explain why aluminium does not corrode easily.

eii. Give a use of aluminium, other than for food containers, together with the physical property that makes aluminium suitable for that use.

eiii. Give a further use of aluminium, that makes aluminium suitable for that use. (excluding that from ei and eii)


Solution

a. Aluminium oxide has a very high melting point. With the addition of cryolite, the melting point is greatly reduced, making it more economical.

bi. 2O2- (l) ---> O2 (g) + 4e-
bii. Al3+ (l) + 3e-  ---> Al (s)

c. The oxygen produced at the graphite anode oxidises the graphite to CO2

d. From the equation 2Al2O3 ---> 4Al + 3O2
1 mole of aluminium oxide gives 2 moles of aluminium. So 102g of Al2O3  gives 54g of aluminium.

Thus 408 tonnes of Al2O3  will give (54/102) x 408 = 216 tonnes of aluminium

ei. Aluminium forms aluminium oxide in the presence of air. This oxide is insoluble and resistant to corrosion so it forms a protective coating for aluminium.

eii. It is used in making cooking utensils since it has very good conductivity, in addition to its good appearance and resistance to corrosion.

eiii. it is a component in several alloys used in aircraft construction. Its favorable use is due to its low density and high tensile strength.

4. Complete the table by naming the products formed when the following liquids are electrolysed using inert electrodes

 liquidproduct formed at cathode
 product formed at anode
 dilute sulphuric acid
  oxygen
 molten calcium bromide
  
 concentrated aqueous sodium chloride
  

Solution


 liquid product formed at cathode  product formed at anode
 dilute sulphuric acid hydrogen oxygen
 molten calcium bromide calcium bromine
 concentrated aqueous sodium chloride hydrogen chlorine

5. Aqueous copper(II) sulphate was electrolysed in two cells using different electrodes as shown below.

a. Write ionic equations, with state symbols, for the reactions which take place at the anode in each cell.
b. Describe one change that you would see happen in both cells.
c. Describe one change that you would see happen in Cell 1 but not in Cell 2.


Solution

a. Anode reaction cell 1: 4OH- (aq) --> O2 (g) + 2H2O (l) + 4e-

Anode reaction cell 2: Cu (s) --> Cu2+ (aq) + 2e-

b. The size of the cathode increases as copper metal is plated onto the cathode in both cells.

Cu2+ (aq) + 2e- --> Cu (s)

c. The blue colour of the electrolyte in cell 1 fades when more and more Cu2+ ions are reduced to copper metal and plated onto the cathode as a pink deposit.

Cu2+ + 2e- --> Cu

6a. Write an ionic equation for the reaction between zinc and aqueous copper(II) sulphate.

This reaction can be used to generate electricity in a cell.


b. Draw an arrow on the diagram to show the direction of the flow of electrons in the wire.

c. The voltage of the cell was measured when the following metals were used as electrode 2.

            copper    iron    lead    zinc


Complete the table by entering the metals in the correct order.


 meter reading/V        
Metal    
 1.10 
 0.78 
 0.21 
 0.00 

d. When metal M was used as electrode 2, it produced a higher voltage than zinc. Suggest a name for metal M.


Solution

a. Zn (s) + Cu2+ (aq) --> Zn2+ (aq) + Cu (s)

Note: zinc, being a more reactive metal, displaces copper ions out of solution as copper metals

b. arrow direction from electrode 2 to electrode 1

Note: Zinc, being a more reactive metal, loses electrons more easily when connected to a metal of lower reactivity, in this case copper.

c.
 meter reading/V
Metal
 1.10 zinc
 0.78 iron
 0.21 lead
 0.00 copper

d. Magnesium

7. Electroplating can be used to coat nickel with a thin coating of silver.

a. Draw a labelled diagram of an apparatus that can be used to electroplate silver onto nickel.

b. Write equations, with state symbols, for the reactions at the anode and cathode.

c. Solutions of two salts, A and B, were electrolysed using carbon electrodes. The following products were collected.

 SaltProducts
 A oxygen and hydrogen
 B chlorine and hydrogen

i. Suggest the names of the two salts, A and B.

ii. Describe tests to confirm the identities of the three gases collected.


Solution

a.
b. reaction at anode: Ag (s) --> Ag+ (aq) + e-

ci. Salt A: sodium sulphate

Salt B: sodium chloride

cii. Collect samples of each gas using test tubes. To test for oxygen, insert a glowing splint into the test tube of gas. The gas that relights the glowing splint is oxygen. To the remaining samples, place a lighted splint at the mouth of each test tube. The gas that extinguishes the lighted splint with a "pop" sound is hydrogen. To identity chlorine, place a piece of moist blue litmus at the mouth of test tube of gas. The litmus turns red and bleaches.

8. One important use of a gas Y is to sterilize swimming pool water. The gas is prepared in the laboratory  by the electrolysis of a solution. A student tried to prepare gas Y by the electrolysis of a very dilute sodium chloride solution as shown below. Contrary to the student's expectation, a colorless gas, instead of gas X, was liberated at the anode.
a. What is the colorless gas liberated?
b. Suggest a chemical test for the colorless gas
c. The experiment was then modified to prepare gas Y.

i. Suggest how the experiment could be modified. Explain your answer.
ii. Suggest the solution left after the electrolysis
iii. Suggest one common use of the solution left.


Solution

8a. oxygen

8b. The gas relights a glowing splint.

8ci. Use concentrated sodium chloride solution instead of a very dilute sodium chloride solution. Concentration of chloride ions in the solution would be much greater than that of hydroxide ions. Therefore, chloride ions would be preferentially discharged to form chlorine gas.

8ciii. manufacture of bleach

9. A dilute copper(II) sulphate solution is electrolysed using carbon electrodes.

a. Describe and explain what would happen at the two carbon electrodes
b. Write half-equations for the reactions at carbon exlectrodes X and Y.
c. What will be the charge in the electrolyte as electrolysis proceeds for some time? Explain your answer.
d. Explain what could happen to the copper(II) sulphate solution if copper electrodes are used in the above experiment.


Solution

9a.
At carbon electrode X:
The sulphate ions and hydroxide ions migrate to electrode X. A hydroxide ion is a stronger reducing agent than a sulphate ion. So hydroxide ions are preferentially discharged.

At carbon electrode Y:
The copper(II) ions and hydrogen ions migrate to electrode Y. A copper(II) ion is a stronger oxidizing agent than a hydrogen ion. So copper(II) ions are preferentially discharged to form a deposit of copper on electrode Y.

9b.
electrode X:
4OH- (aq) --> O2 (g) + 2H2O (l) + 4e-

electrode Y:
Cu2+ (aq) + 2e- --> Cu (s)

9c. The solution becomes sulphuric acid because copper(II) ions and hydroxide ions are consumed in the electrolysis. Hydrogen ions and sulphate ions remain in the solution.

9d. The net effect is the transfer of copper from electrode X to electrode Y. The dilute copper(II) sulphate solution remains the same.

10. Tuning knobs on radios are often made of plastics plated with metal coatings. The plastic knobs are first coated with copper and then electroplated with nickel. The electroplating can be conducted using the following setup.


a. Explain the term 'electroplating'
b. Why is the plastic knob first coated with copper before electroplating?
c. Explain why nickel(II) sulphate solution can conduct electricity
d. Which is the anode, the nickel electrode or the copper-coated knob?
e. Write an ionic half-equation for the reaction at the copper-coated knob
f. Explain why it is better to use a nickel electrode than a carbon electrode in the above process
g. In a nickel-plating factory, the waste water is treated with sodium hydroxide solution to remove nickel(II) ions before discharge. Suggest 2 reasons why it is necessary to remove nickel(II) ions from the waste water before discharge


Solution

10a. It is the coating of an object with a thin layer of a metal by electrolysis
10b. to make the knob conduct electricity
10c. the solution contains mobile ions
10d. the nickel electrode
10e. Ni2+ (aq) + 2e- --> Ni (s)
10f. the concentration of nickel(II) ions in the electrolyte can be maintained
10g.
- to recover the nickel metal
- nickel(II) ions are harmful to marine lives. Humans may get poisoned by eating contaminated seafood

11. The following circuit is set up. Electrodes A and B are made of carbon while electrode C and D are made of copper.

a. What are the functions of the ammeter and rheostat respectively?
b. Explain why no current flows when potassium iodide is in solid state, but a current flows when water is added.
ci. What would be observed at electrode A and B respectively?
cii. Write ionic half-equations for the reactions at electrodes A and B
di. What would be observed at electrodes C and D respectively?
dii. Write ionic half-equations for the reactions at electrodes C and D
diii. Would you expect any color change in the dilute copper(II) sulphate solution during the process? Explain your answer.


Solution

11a. The ammeter is an instrument used to measure the electric current passing through the circuit. The rheostat is used to vary the resistance in the circuit and regulate the current.

11b. In solid state, the ions in potassium iodide are held together by strong attraction. They are not free to move. So solid potassium iodide does not conduct electricity. When water is added to the compound, the compound dissolves in the water and the ions become mobile and a current can then flow through the solution.

11ci. A brown color develops around electrode A. A colorless gas is given off from electrode B.

11cii.
electrode A:
2I- (aq) --> I2 (aq) + 2e-

electrode B:
2H+ (aq) + 2e- --> H2 (g)

11di. Electrode C dissolves. A reddish brown deposit forms on electrode D.

11dii.
electrode C:
Cu (s) --> Cu2+ (aq) + 2e-

electrode D:
Cu2+ (aq) +2e- --> Cu (s)

11diii. The blue color of the dilute copper(II) sulphate solution does not change because the concentration of copper(II) ions in the solution remains the same.

12. A student used the following set-up for passing electricity through some solutions.
The results are shown below:
 ExperimentSolution
Observations
 1 sugar solution
 zero ammeter reading
 2 dilute sulphuric acid
 gas bubbles given off from both electrodes
 3 dilute sodium iodide solution
 ?
 4 dilute silver nitrate solution
 gas bubbles given off from one electrode and silvery solid deposited on the other electrode

a. Why is zero ammeter reading recorded in Experiment 1?
b. For Experiment 2,
i. name the gas liberated at electrodes X and Y respectively
ii. write ionic half-equations for the reactions at the electrodes

c. For Experiment 3,
i. what substance would you expect to form at electrode X. Explain briefly.
ii. what substance would you expect to form at electrode Y. Explain briefly.

d. For Experiment 4,
i. write ionic half-equations for the reactions at the electrodes
ii. state the change in the solution after electricity has been passed through for some time.


Solution

12a. Sugar solution is not a conductor of electricity

12bi. X: oxygen; Y: hydrogen

12bii.
at electrode X:
4OH- (aq) --> O2 (g) + 2H2O (l) + 4e-

at electrode Y:
2H+ (aq) + 2e- --> H2 (g)

12ci. The concentration of iodide ions in the solution is much greater than that of hydroxide ions. Iodide ions are preferentially discharged to form iodine.

12cii. A hydrogen ion is a stronger oxidizing agent than a sodium ion. So hydrogen ions are preferentially discharged to form hydrogen gas.

12ciii. Hydrogen ions and iodide ions are consumed in the electrolysis. Sodium ions and hydroxide ions remain in the solution. Eventually, the solution becomes sodium hydroxide solution.

12di.
electrode X:
4OH- (aq) --> O2 (g) + 2H2O (l) + 4e-

electrode Y:
Ag+ (aq) + e- --> Ag (s)

12dii. Silver ions and hydroxide ions are consumed in the electrolysis. Hydrogen ions and nitrate ions remain in the solution. The solution eventually becomes nitric acid solution.

13. When the circuit in the set-up shown below is closed, the acidified potassium permanganate solution loses its color gradually.
a. Write a half equation for the reaction that occurs in the acidified potassium permanganate solution. Explain whether the permanganate ion is oxidized or reduced.

b. What would be observed in the iron(II) sulphate solution after some time. Write a half equation for the reaction that would occur.

c. Identify the direction of electron flow in the external circuit.

d. Write an ionic equation for the reaction that occurs when acidified potassium permanganate solution and iron(II) sulphate solution are mixed together.

ei. Give the function of the salt bridge set up
eii. Explain whether a sodium sulphite solution can be used instead of a potassium nitrate solution in the salt bridge.

As an alternative to iron(II) sulphate and acidified potassium permanganate solution, potassium iodide solution and iron(III) sulphate solution may be used on the left hand side and the right hand side respectively.

fi. What would be observed in the potassium iodide solution after some time? Write a half equation for the reaction that would occur.
fii. What would be observed in the iron(III) sulphate solution after some time? Write a half equation for the reaction that would occur.


Solution

13a. MnO4- (aq) + 8H+ (aq) + 5e- --> Mn2+ (aq) + 4H2O (l)
MnO4- is reduced because it receives electrons and the oxidation number of Mn changes from +7 to +2.

13b. The solution changes from green to yellow gradually because iron(II) ions are oxidised to iron(III) ions
Fe2+ (aq) --> Fe3+ (aq) + e-

13c. From iron(II) sulphate solution to potassium permanganate solution

13d. 5Fe2+ (aq) + 8H+ (aq) + MnO4- (aq) --> Mn2+ (aq) + 5Fe3+ (aq) + 4H2O (l)

13ei. to allow migration of ions between the two beakers or to complete the circuit

13eii. No. sodium sulphite reacts with potassium permanganate or the sulphite ions can be oxidised by permanganate ions

13fi. the solution turns brown or yellow
2I- (aq) --> I2 (aq) + 2e-

13fii. the solution changes from yellow to green
Fe3+ (aq) + e- --> Fe2+ (aq)

13fiii. 2I- (aq) + 2Fe3+ (aq) --> I2 (aq) + 2Fe2+ (aq)

14. The diagram shows a dry cell

ai. which substance is the positive electrode?
aii. write an ionic half equation for the reaction at the positive electrode

bi. which substance is the negative electrode?
bii. write an ionic half equation for the reaction at the negative electrode

c. explain the function of manganese(IV) oxide

d. the voltage of the cell drops if current is drawn from the cell rapidly. Explain briefly.

e. explain why the zinc case of a used cell is thinner than that of a new cell

f. explain why disposal of 'flat' cells present a pollution problem

g. explain why zinc-carbon cells should be removed from electric appliances when not in use for a long period

Solution
14ai. carbon rod

14aii. 2NH4+ (aq) + 2e- --> 2NH3 (aq) + H2 (g)

14bi. zinc case

14bii. Zn (s) --> Zn2+ (aq) + 2e-

14c. Hydrogen is produced and collected on the surface of the positive electrode. Since hydrogen is a poor conductor of electricity, the accumulation of hydrogen at the positive electrode may hinder further reactions and decrease the current of the cell. Manganese(IV) oxide, an oxidizing agent, is used to remove the hydrogen.

14d. If a current is drawn from the cell rapidly, the gaseous product cannot be removed fast enough. The voltage drops as a result.

14e. The zinc case undergoes oxidation to give zinc ions in the cell reaction.

14f. The materials inside the cells do not decompose even after a long time. These materials may combine with other compounds and form harmful substances which pollute the environment.

14g. There is a slow direct reaction between the zinc electrons and ammonium ions. After some time, the zinc case becomes too thin and the paste leaks out. This may cause damage to electric appliances.


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