1
Q
Define electrolysis
A
Process of passing an electric current though a compound to chemically separate its components
2
Q
What is an electrolytic cell?
A
Device that converts electrical energy into chemical energy
3
Q
What are the parts of an electrolytic cell?
A
1) Battery
2) Electrodes - Anode, cathode
3) Electrolyte
4
Q
Describe the battery
A
- “electron pump”
- electrons enter from the anode and exit to the cathode
- positive charge on anode and negative charge on cathode
5
Q
Describe the electrolyte
A
- electrically conductive substance in molten/aqueous state
- contain mobile ions that act as mobile charge carriers to conduct electricity
- e.g. NaCl(aq), MgBr(aq)
6
Q
Describe the anode
A
- positive electrode connected to positive terminal
- contains delocalised electrons to conduct electricity
- e.g. metal plates and carbon (e.g. graphite) rods
7
Q
Describe the cathode
A
- negative electrode connected to negative terminal
- contains delocalised electrons to conduct electricity
- e.g. metal plates and carbon (e.g. graphite) rods
8
Q
What is the direction of electron flow in the external circuit?
A
Electrons move from the negative terminal to the positive terminal of the power supply
9
Q
What happens at the anode?
A
- Anions are attracted to the positively charged anode
- Anions can lose electrons to the anode and become oxidised
- X^n- → X + ne-
10
Q
What happens at the cathode?
A
- Cations are attracted to the negatively charged cathode
- Cations can gain electrons from the cathode and become reduced
- Y^n+ + ne- → Y
11
Q
What are the steps to writing an overall equation? (e.g. molten NaCl)
A
1) Write half equations:
Cathode: Na⁺(l) + e⁻ → Na(l)
Anode: 2Cl⁻(l) → Cl₂(g) + 2e⁻
2) Balance:
2Na⁺(l) + 2e⁻ → 2Na(l)
3) Combine and cancel:
2Na⁺(l) + 2Cl⁻(l) → 2Na(l) + Cl₂(g)
4) Write the final overall equation:
2NaCl(l) → 2Na(l) + Cl₂(g)
12
Q
Why are electrodes inert?
A
- Electrolysis is carried out to decompose compounds. - If the electrodes take part in the process, contamination of products occurs.
- Hence inert electrodes are used as they are unreactive
13
Q
What are the 2 commonly used materials for electrodes?
A
Graphite and platinum
14
Q
What are the advantages and disadvantages of using graphite as an electrode?
A
Adv:
1) High mp –> Won’t melt during electrolysis of MBIC
Disadv:
1) Reacts to O2 under high temp to produce CO2
2) Needs to be replaced periodically
15
Q
What are the advantages and disadvantages of using platinum as an electrode?
A
Adv:
1) Does not take part in electrolysis
Disadv:
1) Low mp –> Might melt during electrolysis of MBIC
2) Mainly used for aqueous electrolytes
16
Q
What is considered aqueous?
A
Solutions are formed when a solute dissolves in a solvent. Aqueous solutions are formed when water is the solvent
17
Q
What is the equation for water?
A
H2O (l) ⇌ H+ (aq) + OH- (aq)
18
Q
What are the additional cations and anions in aqueous solutions?
A
OH- (hydroxide) and H+ (hydrogen) ions
19
Q
List the metals in order of increasing ease of discharge of cation
A
Potassium
Sodium
Calcium
Magnesium
Aluminium
Carbon
Zinc
Iron
Tin
Lead
Hydrogen
Carbon
Silver
Gold
Platinum
20
Q
Why is it that a more reactive metal is harder to convert back?
A
The more reactive a metal, the more stable its ion - harder to convert ion back to metal
Less reactive a metal, the less stable its ion - easier to convert ion back to metal
21
Q
If an aqueous solution contained Cu2+ and Ag+ ions, which would be selectively discharged?
A
Ag+ ions are selectively discharged at the cathode. Cu2+ only starts getting discharged after all the Ag+ ions are discharged from the solution
22
Q
Which metals can be discharged before hydrogen in an aqueous solution?
A
Copper, silver, gold, platinum. Those above hydrogen will not be discharged first
23
Q
How do anions discharge from dilute aqueous solutions?
A
They discharge based on their position in the electrochemical series
24
Q
The electro chemical series ranks ions in terms of _____
A
electrochemical reactivity
25
State the electrochemical series in order of increasing ease of discharge
F^-, SO4^2-, NO3^-, Cl^-, Br^-, I^-, OH^-
26
If an aqueous dilute solution contains Cl- an OH- ions, which ion has a higher tendency to be selectively discharged?
OH-. It is difficult to discharge any anion other than OH- in dilute aqueous solutions
27
What anions are affected in concentrated aqueous solutions?
A high concentration increases ease of discharge of Cl-, Br- or I- anions, which can override the electrochemical series
28
If a concentrated aqueous solution of NaCl is electrolysed, which anions will be selectively discharged?
Cl- ions will be selectively discharged over OH- ions. Cl2 gas is evolved at the anode instead of O2 gas
29
What is the equation of electrolysis of water?
2H₂O(l) → 2H₂(g) + O₂(g)
30
What is the equation of electrolysis of hydroxide ions?
4OH- (aq) → O2(g) + 2H2O(l) + 4e-
31
What is the equation of electrolysis of hydrogen ions?
2H+(aq) + 2e- → H2(g)
32
What happens to the aqueous solution with dilute NaCl?
H+ and OH- are selectively discharged. Ratio of H2:O2 is 2:1. The solution remains neutral as Na+ and Cl- are left. Water levels drop as electrolysis continues causing the solution to become concentrated enough for Cl- ions to be selectively discharged
33
What is the equation of electrolysis of chlorine?
2Cl⁻(aq) → Cl₂(g) + 2e⁻
34
What happens to the aqueous solution with concentrated NaCl?
Ratio of H2:Cl2 is 1:1. Solution becomes more alkaline/ less acidic as there is net discharge of H+ ions. Remaining Na+ and OH- forms sodium hydroxide (alkaline solution)
35
What gas is produced when OH- ions are discharged?
O2 gas (colourless)
36
What gas is produced when H+ ions are discharged?
H2 gas (colourless)
37
What gas is produced when Cl- ions are discharged?
Cl2 gas (pale green)
38
What is the colour of iodide (I2) gas?
Violet
39
What is the colour of chlorine (Cl2) gas?
Green
40
What is the colour of bromine (Br2) gas?
Brown
41
What is the colour of oxygen (O2) gas?
Colourless
42
Electrolysis is used to purify copper wires, or else the impurities...
Impurities in the copper wire will reduce efficiency of electrical transfer
43
What do reactive electrodes introduce into the electrolyte?
A metal cation
44
What can happen when reactive electrodes are used as the anode?
The electrode is oxidised to form its cation, which is reduced back to the original metal on the cathode
45
If you had a dilute aqueous solution and copper electrodes with an electrolyte containing dissolved Cu^2+ and SO4^2-, what will happen at the anode and at the cathode?
Anode:
SO4^2- and OH^- are attracted, but copper is not an inert electrode so it oxidises, dissolves to form Cu^2+ ions, decreases in mass
Cathode:
Cu^2+ and H^+ are attracted to the cathode, Cu^2+ is selectively discharged, red brown copper is deposited onto the cathode, increases in mass
46
What are the 2 things reactive electrodes allow?
1) Purify metals
2) Plate objects
47
Why does purifying metals and plating objects not have a net reaction?
These processes do not have a net reaction as the metal that is being oxidised at the anode is simultaneously being reduced at the cathode.
48
How is the metal purification set up made?
- The raw, impure copper is
made the anode by connecting it to the positive terminal of the power source
- A small piece of pure copper is made the cathode by connecting it to the negative terminal of the power source.
- The electrodes are then placed in an electrolyte that contain ions of that metal
49
What can you observe at the anode and cathode during metal purification?
The impure copper anode dissolves and decreases in
mass.
Pure copper is deposited onto the cathode.
Impurities present in the impure copper fall to the bottom to form anode slime.
50
What happens if during metal purification, you replace the anode with an inert electrode instead of an impure metal?
- Some pure copper can still be deposited at the cathode. This copper is produced from the Cu²+ ions in the electrolyte.
- The colour intensity of the electrolyte gradually fades from blue to colourless.
- Since the anode is inert, it is not oxidised. Based on the rules we learnt in
Section 13.2, the OH- ions in the electrolyte will be oxidised instead.
51
What is electroplating?
Electroplating allows us to coat a thin layer of metal onto an object.
52
How is the set up for electroplating made?
- The plating metal (metal to be plated on the
object) is made the anode. - The object to be coated with the metal is made
the cathode.
- The two electrodes are then immersed in an aqueous salt solution of the plating metal.
53
What can you observe at the anode and cathode during electroplating?
At the anode:
The copper anode
dissolves to form Cu²* lons
and decreases In mass.
Cu(s)Cu²+(aq) + 2e
At the cathode:
The cathode Is coated with
a layer of copper metal and increases In mass.
Cu2*(aq) +2eCu(s)
Solution:
Since the copper deposited onto the metal spoon comes mainly from the copper anode, the concentration of Cu²+ ions in the electrolyte remains constant throughout. Hence, the
intensity of the blue colour of the solution does not change.
54
What happens when you replace the object to be coated with a non conductive metal object during electroplating?
Non-conductive objects can also be electroplated if they are first coated with a layer of graphite before being immersed in the electrolyte.
55
What are 2 reasons for electroplating?
- Some metals are electroplated to enhance its attractiveness.
- Reactive metals
are also electroplated to protect them from corroding easily
56
What is a simple cell?
A simple cell is a device that converts chemical energy into electrical
energy.
57
How is a simple cell set up made?
1) Two metals of different reactivities connected to an external circuit
2) An electrolyte
Oxidation occurs at the anode and reduction occurs at the cathode, just like in electrolytic cells
58
What happens at the anode and cathode in a simple cell?
The more reactive metal:
- acts as the anode
- is preferentially oxidised,
forming cations that enter
the electrolyte
- and releases electrons that flow through the external circuit.
The less reactive metal:
- acts as the cathode
- and causes cations from the electrolyte to gain electrons and be reduced
59
Describe electron movement in a simple cell
- In electrolytic cells, the flow of electrons is driven by the external circuit.
- In simple cells, the electrons move spontaneously from
where they are produced (anode) to where they are consumed (cathode).
- This movement
of electrons produces electrical energy.
60
Compare the sources of energy in a simple cell and in an electrolytic cell
Simple Cell: Electrical energy is produced through chemical reactions.
Electrolytic Cell: Electrical energy is supplied by an
external source
61
Compare the electron movement of energy in a simple cell and in an electrolytic cell
Simple cell:
Electrons move from the anode to the cathode.
Electrolytic cell:
Electrons move from the battery to the cathode, through the electrolyte,
and into the anode.
62
Compare the polarity of electrodes of energy in a simple cell and in an electrolytic cell
Simple cell:
Anode: negative (-)
Cathode: positive (+)
Electrolytic cell:
Anode: positive (+)
Cathode: negative (-)
63
How do you measure the electrical energy produced?
To measure the electrical energy produced, a voltmeter is used. A voltmeter tells us the potential difference of the cell, which is also commonly referred to as the voltage of the cell.
64
How does distance of metals in reactivity series affect voltage?
Metals are used as electrodes in a simple cell. The further apart two metals are in the reactivity series, the greater the voltage produced.
65
What are hydrogen fuel cells used for?
Hydrogen fuel cells are used to power electric vehicles
66
Describe a hydrogen fuel cell
A fuel cell is similar to a simple cell. It derives its power from a fuel that is
continuously added at the anode and an oxidiser, or an oxidising agent, that
is continuously added at the cathode. The hydrogen fuel cell utilises hydrogen as the fuel and oxygen from air as the oxidiser.
67
Why do hydrogen fuel cells not contribute to climate change?
Hydrogen fuel cells do not directly contribute to climate change because their only product is water, besides electricity.
68
What happens at the anode and cathode of a hydrogen fuel cell?
At the cathode:
O2(g) + 2H2O(2)+4e --> 40H-(aq)
At the anode:
H2(g) + 2OH-(aq) --> 2H,0(l) +2e
69
Describe how hydrogen fuel cells function?
- The ratio of H₂ to O₂ consumed is 2: 1. The overall equation for the reaction that occurs in a hydrogen fuel cell is the same as the combustion of hydrogen.
- The hydrogen fuel does not come into direct contact with oxygen to ensure that the hydrogen is not spontaneously oxidised.
- Excess hydrogen can be pumped back into the anode chamber to
reduce wastage.
- The electrolyte is responsible for "completing" the circuit between the anode
and the cathode.
70
What are the advantages of using hydrogen fuel cells?
1) Hydrogen is a renewable fuel.
2) It can be obtained via the electrolysis of water or from the cracking of hydrocarbons.
3) Hydrogen fuel cells produce only water as a by-product.
4) It is more efficient than fuel-burning electricity sources.
5) A larger percentage of the chemical energy stored in the fuel ends up as useful electricity in a fuel cell.
71
What are the disadvantages of using hydrogen fuel cells?
1) It is difficult to store and transport hydrogen safely as hydrogen is a
highly flammable gas at room temperature and pressure.
2) Hydrogen is
often transported in high-pressure cylinders which can be dangerous to
handle.
3) A large amount of energy is needed to produce hydrogen from
electrolysis.
4) However, this can be mitigated by using renewable sources
like solar and wind energy to power electrolysis.
Pure Chemistry
flashcards
Decks in class (21)
# Cards
-
Chapter 01 - Experimental Chemistry127
-
Chapter 02 - Kinetic Particle Theory64
-
Chapter 03 - Atomic Structure51
-
Chapter 04 - Chemical Bonding24
-
Chapter 05 - Structure and Properties of Materials88
-
Chapter 06 - Chemical Formulae and Equations35
-
Chapter 07 - Mole Concept and Stoichiometry30
-
Chapter 11 - Qualitative Analysis58
-
Chapter 15 - The Reactivity Series82
-
Chapter 12 - Oxidation and Reduction65
-
Chapter 16 - Chemical Energetics20
-
Chapter 17 - Rate of Reactions30
-
Chapter 10 - Ammonia17
-
Chapter 13 - Electrochemistry71
-
Chapter 18 - Fuels and Crude Oil28
-
Chapter 19 - Hydrocarbons54
-
Chapter 20 - Polymers11
-
Chapter 22 - Maintaining Air Quality58
-
Chapter 08 - Acids and Bases49
-
Chapter 09 - Salts25
-
Chapter 14 - Periodic Table42
