Redox and galvanic cells

Question 1

Oxidation and reduction

Answer 1

Interdependent, simultaneous, electron transfer process

Question 2

Oxidation

Answer 2

Involves losing electrons (Lose Electron Oxidation, LEO)

Question 3

Reduction

Answer 3

Involves gaining electrons (Gain Electron Reduction, GER)

Question 4

Oxidising agent

Answer 4

Substance that has the ability to remove electrons from a substrate, substance that is reduced

Question 5

Reducing agent

Answer 5

Substance that has the ability to donate electrons to a substrate, substance that is oxidised

Question 6

Anode

Answer 6

Involves oxidation, loses mass

Question 7

Cathode

Answer 7

Involves reduction, gains mass

Question 8

Galvanic cells

Answer 8

Redox reactions generate electricity if the two half-reactions are physically separated so that the transferred electrons can be directed through an external circuit (produces power)

Question 9

Purpose of galvanic cells and electrolytic cell

Answer 9

Galvanic- Turn chemical potential energy into electrical energy
Electrolytic- Turn electrical energy into chemical potential energy

Question 10

What does the salt bridge do in a galvanic cell

Answer 10

Maintains electrical neutrality, Completes circuit, Facilitates ion movement

Question 11

Why may the voltage reading not be accurate to calculations

Answer 11

Not being at standard conditions, Not accurate/correct concentration of solutions

Question 12

Primary batteries- dry cell

Answer 12

Single use, Oxidation: Zn —> Zn2+ + 2e, Reduction: 2MnO2 + 2NH4+ + 2e- —> Mn2O3 + H20 + 2NH3, Porous separator between cathode and anode which acts as a site for electron transfer, Moist paste which allows flow of charge (cations and anions), Moist past contains NH4Cl which provides needed acidic conditions, Anode is the negative terminal

Question 13

Secondary batteries- lead accumulator

Answer 13

Rechargeable, Cathode: PbO2 + 4H+ + SO4 2- <—> 2H2O + PbSO4, Anode: Pb + SO4 2- <—> PbSO4 + 2e-, Discharge: Pb + 2SO4- + 4H+ + Pbo2 —> 2PbSO4 + 2H2O, causes decrease in pH, H2SO4 acts as the electrolyte, forms a precipitate

Question 14

Fuel cells

Answer 14

Anode: 2H2 + 2OH- —> 2H20 + 2e-, Cathode: O2 + 2H2O + 4e- —> 4OH-, Continuous supply of oxygen and hydrogen (from fossil fuels), Porous electrode of nickel or platinum which acts as a catalyst, High temperature makes it less efficient, Acidic or alkaline conditions, Constant rate or electricity- depends on reaction rates

Question 15

Half reactions in galvanic cells vs electrolytic cells

Answer 15

Galvanic- spontaneous (positive E value), separate half-reactions
Electrolytic- Non-spontaneous (negative E value), Half reactions not separate (all in one container) but sometimes products need to be separated

Question 16

Anode and cathode in galvanic cells vs electrolytic cells

Answer 16

Galvanic- Anode- oxidation, negative, Cathode- reduction, positive
Electrolytic- Anode- oxidation, positive (plugged into positive terminal), Cathode- reduction, negative (plugged into negative terminal of DC power supply)

Question 17

Movement of ions in galvanic cells vs electrolytic cells

Answer 17

Galvanic- Anions move towards the anode, Cations move towards the cathode, Electrons go from anode to cathode
Electrolytic- Anions move towards the anode, Cations move towards the cathode, Electrons travel from negative terminal to anode to cathode and positive terminal

Question 18

Use of a salt bridge in galvanic cells vs electrolytic cells

Answer 18

Galvanic- maintains electrical neutrality
Electrolytic- not needed

Question 19

Commercial cells

Answer 19

Storage cells and fuel cells

Question 20

Storage cells

Answer 20

Standard batteries, ‘batch’ process, one lot of reactants

Question 21

Fuel cells

Answer 21

Continuous process, reactants continually delivered, products constantly removed, Types of storage cells: primary (one life), secondary (rechargeable)

Question 22

Galvanic cells criteria

Answer 22

Spontaneous, Separate half reactions, External circuit- carry electrons, Internal circuit- salt bridge, Electrolyte solution

Question 23

Wet corrosion

Answer 23

Iron rusts due to the porous and non-adherent oxide coating which flakes easily and comes off whereas other metals have an non-porous adherent coating, Anode reaction- Fe (s) —> Fe2+ (aq) + 2e-, Cathode reaction (always the same)- O2 (g) + 2H20 (l) + 4e- —> 4OH- (aq), Carbon atom- provides surface for electron transfer (anode), Iron undergoes oxidation (most exposed to oxygen- differential aeration) and Fe2+ ions move towards the cathode, OH- moves towards the anode and collides with Fe2+ ions to form green Fe(OH)2, Fe(OH)2 is further oxidised to brown Fe(OH)3 by reacting with more dissolved oxygen- 4Fe(OH)2 (s) + O2 (aq) + 2H20 (l) —> 4 Fe(OH)3, Fe(OH)3 dehydrates to form rust - 2 Fe(OH)3 (s) —> Fe2O3.nH2O (s) (dry and flaky) + 3-nH20 (l)

Question 24

How corrosion is prevented (6)

Answer 24

Modifying material, modifying environment, galvanizing with zinc (protective metal coating), tin coating, cathodic protection (sacrificial anode), impressed current

Question 25

Modifying material

Answer 25

Alloy, mixing/adding metal with a non-porous, adherent oxide coating such as stainless steel which has chromium

Question 26

Modifying environment

Answer 26

Stopping water from getting on material, Eg factories which use reverse cycle air conditioning that removes water and clean rooms which have water and oxygen removed

Question 27

Protective non-metal coatings

Answer 27

Must have an unbroken surface as needs to stop water and oxygen coming in contact with iron (barrier), Oil, Grease- non polar so won’t mix with non-polar water, Paint- however if there is a chip, rust can occur can continue to occur under the paint, Enamels, Glass, Plastic

Question 28

Protective metal coating: galvanizing with zinc

Answer 28

Zinc produces a protective oxide coating (acts as barrier) as zinc has a non-porous, adherent oxide coating, If zinc coating is broken, zinc acts as a sacrificial anode (Zn —> Zn2+ + 2e-) as it preferentially oxidises as it is a stronger reducing agent (show with half equations and E naught value), Iron becomes cathodic- O2 (g) + 2H20 (l) + 4e- —> 4OH- (aq), Zn2+ reacts with OH- to precipitate Zn(OH)2, Zn(OH)2 reacts with carbon dioxide to form 2nCO3.2nOH which is non-porous and adherent (Zn(OH)2 + CO2 —> ZnCO3.2nOH + H20)

Question 29

Tin coatings

Answer 29

“Tin can” has a layer of tin on the inside and outside which acts a barrier (excludes oxygen and water), Iron preferentially oxidises because it is a stronger reducing agent compared to tin (use half equations to show), When tin coating is broken the rate of steel corrosion increases (as iron is exposed) which is helpful when tin cans end up in landfill and need to rot away, Cathode surface (cathodic)- Sn —> Sn2+ + 2e-, Preferentially oxidises- Fe (s) —> Fe2+ + 2e-

Question 30

Cathodic protection- sacrificial anode

Answer 30

Goal- make iron the cathode, Makes steel cathodic by electrically connecting it to a more reactive metal which preferentially oxidises as it is a stronger reducing agent (Eg Mg, Zn or Al), Once the anode rusts away the bottom or a water tank starts to rust as it has little exposure to oxygen, Used on outboard motors, pipelines and water heaters

Question 31

Impressed current

Answer 31

Steelwork is put at a negative potential by connecting it at intervals to a low-voltage DC source, A DC attaches to the jetty making it cathodic due to the electrolysis, The DC is connected to scrap iron or a sacrificial anode which rusts away instead, Needs a DC converter as if the current constantly changed the jetty would rust, Used on jetties and pipelines