Electrochemical Cell Flashcards
What to do when balancing half-equations?
try use data booklet for relevant half-eqn
Describe what happens when strip of Mg metal is placed in Mg2+ solution
(half-cell config)
- Mg atom on strip surface may lose e- bcome ion OR Mg2+ ion may gain 2 e-, deposit as Mg atom on strip
- When rate Mg2+ ion re-join surface = rate leaving surface, dynamic eqm established
- at eqm, const -ve charge on Mg metal, const no. Mg cations in sol n ard metal
=> separat n of charge across metal-sol n interface, so pd btw Mg metal & Mg2+ sol n. Pd aka electrode potential (E) of metal
What is needed for measuring electrode potential of a half-cell?
- standard condit n (25 deg C/ 298K, Pa = 1 bar, 1 mol dm-3 conc)
- standard electrode as reference electrode
Describe standard hydrogen electrode (SHE)
- reference hydrogen half-cell assigned as zero
- used to find out Eθ value
SHE consists of - platinised platinum electrode (inert)
- immersed in sol n where [H+] = 1mol dm-3
- H2 gas at 1 bar bubled over Pt electrode
- temp kept at 298K
*platinised mean coated w platinum black increase surface area so eqm btw H2 & H+ established rapidly vs polished surface
Define standard electrode potential of a half-cell, Eθ
Eθ is electromotive force, measured at 298K, btw half-cell & standard hydrogen electrode, in which conc of any reacting species in sol n is 1 mol dm-3, any gaseous species at Pa 1 bar
What does Eθ tell?
*can be refer as reduct n potential
- measure tendency reduct n rxn occur relative to SHE
- more +ve, greater tendency reduct n occur
- less +ve, greater tendency oxidat n occur
How to write Eθ fully?
+/- Eθ Mn+/M
where
+/- indicate polarity of electrode wrt SHE,
Mn+/M refer to oxidised & reduced species respectively (subscript)
What are 3 common types of half-cells?
- metal-ion half cell
- gas-ion half cell
- ion-ion half cell
- must draw all types + label + standard condit n present
What must be the conditions to measure standard electrode potential?
- 298K, 1 bar
- all reacting ions hv conc 1 mol dm-3
- any gas must hv Pa 1 bar
- Pt (bcos inert, for gas-ion, ion-ion) as electrode when half cell no include metal
- salt bridge completes circuit (inverted u-shaped tube w high conc non reacting ion in gel) (prevent mixing of 2 sol n but allow ion pass thru, maitain electrical neutrality)
Describe experimental procedure to determine Eθ M n+/M
- use standard condit n 298K, 1 bar, conc of reacting ion (H+, M n+) 1 mol dm-3 each
- SHE used as refer electrode, 2 half cell connected by salt bridge
- pd set up when 2 half cell connected can b measured by high resis voltmeter
- value shown will be standard electrode potential Eθ M n+/M
How does salt bridge maintain electrical neutrality in each half-cell?
- w/o salt bridge, Cl2/Cl- half cell wld slowly bcome more -ve charged as more Cl- produced, while H2/H+ half cell more +ve charged as more H+ produced. Charge imbalance stop cell operat n
- thus, w salt bridge, ions present in salt bridge (eg Na2+, SO4 2-) diffuse out, ensure e- neutrality
(Na+ ion move fr salt bridge to sol n Cl2/Cl- half cell, SO4 2- ion move fr salt bridge to sol n H2/H+ half cell)
What factors affect standard electrode potential? Elaborate concept
By LCP, affected by
- conc ions
- temp
- Pa of gas species (size/surface area of electrode no affect Eθ cell, since M(s) no shift eqm pos n)
Concept:
M n+(aq) + ne- ⇌ M (s)
- value of electrode potential Eθ M n+/M can be affected by above factors, shift eqm pos n
- if shift right, reduct n favoured, so Eθ more +ve
- if shift left, oxidat n favoured, so Eθ less +ve
By convention, electrode potentials of half cells are written as …
reduction processes
What happens to change in Eθ value?
- more +ve Eθ value,
- eqm pos n shift right
- forward rxn favoured
- reduct n more likely occur
- higher tendency gain e-
- stronger OA
eg Cu2+ higher tendency reduced as Cu vs Zn - less +ve/more -ve Eθ value,
- eqm pos n shift left
- backward rxn favoured
- oxidat n more likely occur
- higher tendency lose e-
- stronger RA
eg Zn higher tendency oxidised to Zn2+ vs Cu
How to prevent rusting?
- can b prevented/slowed down by galvanising (eg iron galvanised w zinc)
- iron connected to metal eg Zn w more -ve Eθ (Zn more easily oxidised)
- since Eθ Zn2+/Zn more -ve than Eθ Fe2+/Fe, zinc undergo oxidat n, termed as sacrificial corros n instead of Fe
Describe relative stabilities of metal ions in different mediums, using Fe2+(aq) as example
Fe2+(aq) usually stabilised against oxidat n by add n small amt acid
bcos,
in acidic medium
Fe3+ + e- ⇌ Fe2+ Eθ = +0.77V
- Eθ value more +ve, so Fe2+ less likely oxidised to Fe 3+
- Fe2+ more stable than Fe 3+ in acidic medium
in basic medium
Fe(OH)3 + e- ⇌ Fe(OH)2 + OH-, E= -0.56V
- Eθ value more -ve thus Fe(OH)2 more likely oxidised to Fe(OH)3
- Fe2+ less stable than Fe3+ in basic medium
Define galvanic/voltaic cell
electrochemical cell in which spontaneous rxn occur to convert chem energy to electrical energy
What occurs at anode and cathode of galvanic cell?
anode: oxidation (An Ox)
cathode: reduction (Red Cat)
Describe what happens in a galvanic cell
- no direct transfer e- fr eg Zn metal to Cu2+ ion as separated fr each other (oni travel thru wire)
- e- flow fr anode to cathode thru external ciruit always
- convert chem to electrical energy
- An Ox
- Red Cat
- polarity of anode is -ve, of cathode is +ve (reverse of electrolytic cell)
Define standard cell potential Eθcell/cell e.m.f.
Give formula
p.d. btw cathode & anode of cell
Eθcell = Eθred - Eθoxid
where
Eθred & Eθoxid is standard electrode potential of half-cell undergoing reduct n & oxidat n respectively
* No need multiply by stoichiometric coeff in rxn
Describe steps to apply Eθcell formula
- check Eθ of each half-eqn in data book
- more +ve value will be Eθred
- less +ve value will b Eθoxid - calculate w Eθcell = Eθred - Eθoxid
- write overall eqn (if qn asks)
- if there is specie w multiple OS, calculate (no need show working unless needed) until get Eθcell value < 0 (non-spontaneous rxn)
Describe steps to predict feasibility of redox reaction in cell
- determine if starting rxt oxidised or reduced in eqn based on change in ON (not Eθ value)
- choose relevant half eqn fr data book
- calculate Eθcell = Eθred - Eθoxid
- predict feasibility:
- Eθcell > 0, spontaneous rxn
- Eθcell < 0, not spontaneous rxn
- Eθcell = 0, rxn reached eqm
=> write concluding statement (eg thus Mg2+ (aq) cnt oxidise Zn(s) to Zn2+(aq) )
Explain limitations of standard electrode potentials, Eθ
- energetics vs kinetic feasibility
- Eθ value predict whether rxn energetically feasible but NOT rate rxn
- even if rxn hv Eθ > 0, rxn unlikely occur as slow rate (too kinetically slow) rxn due to vv high Ea - non-standard conditions
- Eθ value relate oni to standard condit n (otherwise, invalid value)
- factors affecting eqm pos n (ie conc, temp, Pa) affect Eθ value
=> if Eθ value change, no more standard condit n
Give Gibbs free energy and electrode potential relationship formulae. What can be deduced?
ΔGθ = -nFEθcell
where
n is no of e- transferred per mole rxn,
F is Faraday const (9.65E4 C mol-1)
*NOTE: units of ΔGθ is J mol-1, NOT kJ mol-1
if combine ΔGθ = -RTlnK w above eqn, we get,
Eθcell = RT(lnK)/(nF)
- since Eθcell depend on temp & conc, non-spontaneous rxn under standard condit n can b made feasible under non-standard condit n (eg increase temp, conc)
Describe signs of ΔGθ and Eθcell for spontaneous and non spontaneous reactions
- spontaneous: ΔGθ < 0, Eθcell > 0
- non-spontaneous: ΔGθ > 0, Eθcell < 0
Describe 2 main battery types
- primary cell
- electrochem cell in which chem rxn generate electrical energy
- redox rxn spontaneous, irreversible (NOT rechargeable)
eg daniel cell, dry cell, alkaline battery - secondary cell
- redox rxn spontaneous, reversible (rechargeable)
eg lead-acid storage cell, nickel-cadmium cell
Describe fuel cell
- consist of fuel eg H oxidised directly by oxidant (usually O). Energy produced is converted to electrical energy
- rxt continuously replaced as they r consumed & pdt continuously removed
What can be deduced from having fuel cell in acidic and alkaline medium?
- Eθcell is const for fuel cell in both acidic, alkaline medium (ie rxn equally spontaneous for both medium)
Name main advantages, disadvantages and extent of application of fuel cells
advantage:
- pollu n free (produce water)
- high pwr to mass ratio (small size, low mass but generate high voltage)
- highly eff (> 70% chem energy converted to electrical energy vs ~ 30% for traditional pwr stat n
disadvantages:
- expensive due to large amt Ni & Pt needed as catalyst for electrode rxn
- catalyst easily poisoned by impurities in fuel & oxygen
- most fuel cell require high temp to react in fuel cell
applicat n:
- limited due to high cost, temp involved (eg store H2 gas difficult)
- now used to supply electrical pwr in space shuttle. Final pdt of water removed, consumed by astronauts
Name main advantages and disadvantages of zinc-carbon dry cell
advantage
- cheap, portable source energy
- supplies reliable voltage
disadvantage
- not rechargeable (gas pdt eg H2 cnt b consumed rapidly enough when current drawn rapidly)
- poor shelf-life when cell is used, zinc casing disintegrate (due to oxidat n of zinc when react w acidic electrolyte NH4Cl give Zn2+ & hydrogen). Paste will leak through outer casing
Name advantage and disadvantage lead-acid storage battery
advantage
- provide initial high voltage needed to start car engine
disadvantage
- large, heavy weight due to lead (provides low pwr compared to mass it hv)
