Redox and galvanic cells Flashcards
Oxidation and reduction
Interdependent, simultaneous, electron transfer process
Oxidation
Involves losing electrons (Lose Electron Oxidation, LEO)
Reduction
Involves gaining electrons (Gain Electron Reduction, GER)
Oxidising agent
Substance that has the ability to remove electrons from a substrate, substance that is reduced
Reducing agent
Substance that has the ability to donate electrons to a substrate, substance that is oxidised
Anode
Involves oxidation, loses mass
Cathode
Involves reduction, gains mass
Galvanic cells
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)
Purpose of galvanic cells and electrolytic cell
Galvanic- Turn chemical potential energy into electrical energy
Electrolytic- Turn electrical energy into chemical potential energy
What does the salt bridge do in a galvanic cell
Maintains electrical neutrality, Completes circuit, Facilitates ion movement
Why may the voltage reading not be accurate to calculations
Not being at standard conditions, Not accurate/correct concentration of solutions
Primary batteries- dry cell
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
Secondary batteries- lead accumulator
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
Fuel cells
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
Half reactions in galvanic cells vs electrolytic cells
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
Anode and cathode in galvanic cells vs electrolytic cells
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)
Movement of ions in galvanic cells vs electrolytic cells
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
Use of a salt bridge in galvanic cells vs electrolytic cells
Galvanic- maintains electrical neutrality
Electrolytic- not needed
Commercial cells
Storage cells and fuel cells
Storage cells
Standard batteries, ‘batch’ process, one lot of reactants
Fuel cells
Continuous process, reactants continually delivered, products constantly removed, Types of storage cells: primary (one life), secondary (rechargeable)
Galvanic cells criteria
Spontaneous, Separate half reactions, External circuit- carry electrons, Internal circuit- salt bridge, Electrolyte solution
Wet corrosion
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)
How corrosion is prevented (6)
Modifying material, modifying environment, galvanizing with zinc (protective metal coating), tin coating, cathodic protection (sacrificial anode), impressed current
Modifying material
Alloy, mixing/adding metal with a non-porous, adherent oxide coating such as stainless steel which has chromium
Modifying environment
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
Protective non-metal coatings
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
Protective metal coating: galvanizing with zinc
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)
Tin coatings
“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-
Cathodic protection- sacrificial anode
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
Impressed current
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
