Redox II Flashcards
(38 cards)
What is oxidation and reduction?
OXIDATION : loss of electrons
Addition of oxygen
Loss of hydrogen
REDUCTION: gain of electrons
Loss of oxygen
Addition of hydrogen
What are standard conditions for electrode potentials?
Ion conc 1.00 Mol/dm
298K temp
100kPa pressure
- use high resistance voltmeter
What is standard electrode potential?
Potential difference produced when standard half cell is connected to a standard hydrogen cell under standard conditions
What is the standard hydrogen electrode ?
Half cell in which a 1.00 mol/dm solution of hydrogen ions is in equilibrium with hydrogen gas at 100kPa pressure on the surface of a platinum electrode coated with platinum black at 298K
What is a reference electrode? Why are they needed?
Used to measure electrode potentials in place of the standard hydrogen electrode
- equilibrium position can change electrode potential so need to be able to compare electrode potentials of different species by measuring against common reference/standard
Where is reduction and oxidation more likely to occur?
Reduction - at the more + electrode
Oxidation - at the more - electrode
When is a platinum electrode used to find SHE?
When the half cells are
NON METALS/ non metal ions or IONS (with different oxidation states)
- platinum is inert/doesn’t take part in reaction /redox equilibrium established on platinum surface
What is the salt bridge for? What chemicals are used and why?
- has mobile ions that complete the circuit
Potassium chloride /potassium nitrate used too make salt bridge as CHLORIDES/NITRATES are soluble
Should ensure no precipitates form which could affect equilibrium position of half cells
How to calculate emf/potential difference of electrochemical cell?
E right - E left = Ecell
How do the E values infdicate how easily the species can be oxidised or reduced? When is reaction feasible?
More + the value - easier it is to REDUCE the species on the left of half equation
E.g Cl2 + 2e- —> 2Cl- E= +1.36V
- reaction will proceed in forward reaction
Less + the value , easier it is to OXIDISE the species on right of half equation
E.g Cu2+ + 2e- —> Cu. E = +0.34V
- reaction will proceed in backward reaction
Reaction is FEASIBLE when Ecell is +
How is cell potential liked to entropy,S/equilibrium constant,K?
Larger cell potential = BIGGER CHANGE IN ENTROPY
Cell potential directly proportional to entropy change
Limitations of standard electrode potential predictions?
- thermodynamic feasibility of reaction can be deduced but no info about the rate of reaction
- have to consider the le Chatelier’s effect on conc using non-standard conditions , when measuring electrode potentials
- rate of reaction may be too slow- looks like no reaction is occurring when it is
- reaction has high Ea which inhibits reaction
- not all redox reactions are aqueous
What is oxidised and reduced in s,p,d block?
S block elements - oxidised to form +1/+2 ions
P block elements
metal ions: oxidation to form + ions that are consistent/not with group of periodic table
E.g Sn —> SN2+ + 2e-
Al —> Al3+ + 3e-
Non metal ions: reduced to form - ions
Charge : group no. MINUS 8
D block elements - form various ions due to variable oxidation states
- usually oxidised to form + ions (cu2+ , Cr3+, V5+)
What is a fuel cell?
Electrochemical cell in which a fuel donates electrons to 1 electrode and oxygen gains electrons at the other electrode
- methanol/hydrogen rich fuels can be used in fuel cells as well as different electrolytes
Half equations for alkaline hydrogen-oxygen fuel cells?
NEGATIVE ELECTRODE:
H2 (g) + 2OH– (aq) → 2H2O (l) + 2e– Eθ=-0.83V
POSITIVE ELECTRODE:
O2 (g) + 2H2O + 4e– → 4OH– (aq) Eθ =+0.40V
Full equation :
2H2 (g) + O2 (g) → 2H2O (l) Eθ = +1.23 V
Half equations for acidic hydrogen-oxygen fuel cell?
NEGATIVE ELECTRODE:
H2 (g) → 2H+ (aq) + 2e– Eθ = 0.00 V
POSITIVE ELECTRODE:
O2 (g) + 4H+ (aq) + 4e– → 2H2O (l) Eθ = +1.23 V
Full :
2H2 (g) + O2 (g) → 2H2O (l) Eθ = +1.23 V
Benefits of fuel cells?
- water is only by product - good for environment
- Greater efficiency
- no harmful oxides of nitrogen produced
Limitations of fuel cells?
- hydrogen is flammable - storage/production of this is safety hazard
- thick walled cylinders/pipes need to store hydrogen (economic impacts)
- hydrogen is by-product of crude oil industry so relies on non renewable/finite resource
- expensive to produce
Potassium manganate titrations - what is used? When is end point reached?
Manganate (VII) = OXIDISING AGENT /reduced to Mn2+ (burette)
Iron = REDUCING AGENT /oxidised to Fe3+
Dilute sulfuric Acid added to iron before reaction starts (doesn’t oxidise under the conditions/doesn’t react with Mn(VII) IONS )
- purple —> pale pink colour - remains as you swirl
- pale pink due to excess of manganate (VII)ions after all analyte ions reacted with manganate (VII) ions
Iodine thiosulfate titrations ?
2S2O32– (aq) + I2 (aq) → 2I–(aq) + S4O62– (aq)
pale yellow colour of the iodine turns paler as it is converted to colourless iodide ions
When the solution is a straw colour, starch is added to clarify the end point -
Blue/black —> colourless
- titration can be used to determine conc of oxidising agent, which oxidises iodine ions to iodine molecules
Why should starch not be added until nearly all iodine has reacted in iodine thiosulfate titration?
Blue complex formed with high concs of iodine is INSOLUBLE
What are lead acid batteries ?
Consists of 6 cells joined tgt in series - powers starter motor in car engines
- use lead metal at negative electrode + lead oxide at positive electrode
- sulfuric acid electrolyte
Pb (s) + SO42- (aq) → PbSO4 (s) + 2e- Eθ = -0.36 V
PbO2 (s) + 4H+ (aq) + SO42- (aq) + 2e- → PbSO4 (s) + 2H2O (l) Eθ = +1.70 V
overall reaction is
PbO2 (s) + 4H+ (aq) + 2SO42- (aq) + Pb (s) → 2PbSO4 (s) + 2H2O (l) Eθcell = +2.06V
How do lead acid batteries regenerate lead+lead oxide?
Car is in motion - generator provides a push of electrons that REVERSES reaction/regenerate lead and lead oxide
Disadvantage of lead acid batteries?
Heavy
Contain toxin materials
Sulfuric acid electrolyte is corrosive
