Phyisical Chemistry- Electrode Potentials, electrochemical cells

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between half cells + balance out charges -Zinc strip In other half cell same number of electrons taken form external circuit reducing Cu 2+ to copper atoms Voltmeter in external circuit shows voltage between 2 half cells This is cell potential or emf Solution of Zn² (aq) (1 mol dm³) Anode Zn (-) 20 Zn Zn2+ Zn²+ Voltmeter Salt bridge Oxidation half-reaction Zn(s) Zn²+ (aq) +2e 50 2- Cathode (+) Cu Cu2+ Electrode potential Reaction at electrodes are reversible Direction each reaction goes depends on reactivity of metals This measured with electrode potentials A metal that is easily oxidised has a negative electrode potential Cu²+ 20 Reduction half-reaction 2e + Cu²+ (aq) Cu(s) Overall (cell) reaction Zn(s) + Cu²+ (aq) →→→ Zn²+ (aq) + Cu(s) Aqueous ions of same elements Can have half cells involving solutions of 2 aqueous ions of same elements e.g. Fe²+ (aq)/Fe³+ (aq) Conversion form Fe²+ to Fe³+ (+ other way round) happens in surface of platinum electrode Platinum is inert mental so not react with ions Direction of conversion depends on other half cell If it is less reactive than iron, iron will oxidise from Fe²+ to Fe³+ at electrode If other cell more reactive, Fe will reduce from Fe³+ to Fe²+ Whatever is more reactive will be oxidised A metal that is easily reduced has a positive electrode potential Zn²+/Zn -0.76 = reducing agent, more reactive (is oxidised) Cu²+/Cu +0.34 = oxidising agent, less reactive (is reduced) Reducing agent = Negative electrode potential Oxidising agent = Positive electrode potential Drawing electrochemical cells reduced midised micised reduced version version version version salt bridge RIOTTOIR more negative/ More reactive on left hand side Increase agent (a) Tendency for oxidation to occur (b) Power as reducing Elements Li K Ba Ca Na Mg Al Zn Cr Fe Cd Pb Co Ni Sn Pb H₂ Cu Fe Hg Ag Hg N₂ Br₂ 0₂ Cr Cl₂ Au Mn F₂ Electrode Reaction Oxidised Form + ne- Li*(aq) + K'(aq) + e- Ba(aq) + 2e Ca²+ (aq) + 2e Na (aq) + e- Mg2 (aq) + 26 Al³(aq) + 3e Zn²(aq) + 2e Cr³(aq) + 3e- Fe²(aq) + 2e- H₂O(1) + +H₂(g) + OH(aq) Cd²+ (aq) + 2e Cd(s) PbSO4(s) +20- →Pb(s) + SO² (aq) Co (aq) + 2e- Co(s) Ni2+ (aq) + 2e- Ni(s) Sn²(aq) + 2e Sn(s) Pb (aq) + 2e →→→ Pb(s) 2H¹ +2e" Cu(aq) + 2e L₂(s) + 2e Fe³+ (aq) + e Hg₂+ (aq) + 2e- Ag*(aq) + e g2(aq) + 2e NO₂ + 4H+3c- Br.(aq) + 2e- O₂(g) + 2H,O'(aq) + 2e- Cr₂O, +14H+ Cl₂(g) + 2e Au (aq) + 3e MnO,- (aq) + 8H,0¹(aq) + Se Reduced Form Li(s) K(s) Ba(s) Ca(s) Na(s) Mg(s) Al(s) Zn(s) → Cr(s) →Fe(s) H₂(g) (standard electrode) →→→ Cu(s) F₂(g)+2e- 21 (aq) →Fe²(aq) → 2Hg(1) → Ag(s) Hg(1) → NO(g) + 2H₂O →2Br (aq) → 3H₂O 2Cr +7H₂O →→→→→2C1 (ag) →→→ Au(s) Mn²(aq) + 12H₂O(0) 2F (aq) E.g. Lithium and Magnesium Lithium more negative, oxidised more easily so on left side E(volts) <<-3.05 -2.93 -2.90 -2.87 -2.71 -2.37 -1.66 -0.76 -0.74 -0.44 -0.41 <<-0.40 -0.31 -0.28 -0.25 -0.14 -0.13 0.00 +0.34 +0.54 +0.77 +0.79 +0.80 +0.85 +0.97 +1.08 +1.23 +1.33 +1.36 +1.42 +1.51 +2.87 Tendency for reduction to occur (b) Power as oxidising agent (a) Increase Cell notation is RIO||OIR 2+ Lics) Licona Moona Macs) Standard hydrogen electrode (SHE) Potential of all electrodes measured by comparing their potential to standard hydrogen electrode So called primary standard as its standard to which all other potentials compared H₂ at 100 kPa- temperature - 298 K Li(s) → Licon) + C² ре 2+ Mg (aq) + 2e → Mg (S) 2H₁ =26= H₂₂) Pt(s) | H₂(g) | Haas 1.0 MH() V high resistance voltmeter salt bridge Cu 1.0 M Cu² E = +0.DV left hand half cell right hand half cell Pt(s) | H₂(g) | H*(aq) || Cu²+ (aq) | Cu(s) ↑ ↑ ↑ phase boundaries phase boundary salt bridge Hydrogen more easily oxidised (has more electrode potential) so goes on left with platinum electrode Eº = 0.00 V hydrogen gas hydrogen ion Pt Fe PE (5) oxidation anode Fe2+ 1 2H(aq) + 2e H₂(g) Cr₂O7²(aq) + 14H* + 6e¯ ⇒ 2Cr³+ (aq) + 7H₂O e- Anode (oxidation) 1.33V Hº (0₂) reduction cathode ||₂0| Volt meter Salt Bridge ex: NaCl + 2+ Fe²(aq) + 2e MnO4 (aq) + 8H* + 5e¯ ⇒ EⓇ = 1.33 V Cr₂O72- dichromate ion, chromium(III) ion Cr³+ Pt -hydrogen ion e- и U2+ CU 2 | P (0g) (ag) (S) MnO4 Pt 2+ H* Mn²+ Cathode (reduction) Fe(s) 2+ Mn²*(aq) + 4H₂O Nickel Electrode (Anode) Ni²+ (1 M) SO4²- fe felt (5) fels || Mno H² M₁Ⓡ 10-1₁ n* Mn | | | | (aq) (09) (09) Ni(s) = Ni²+ (aq) + 2 e- V Salt Bridge Nat Ni2+ (aq) + 2e Ni(s) 0.76 V(SHE) Ni (5) (48) Ni SO4²- $|| H*(1 M) 2 H+ (aq) + 2 e = H₂(g) M² Platinum Electrode M₂ (99) Pt (S) 2H*(aq) + 2e H₂(g) Pt LO H₂(g) (1 atm) Standard Hydrogen Electrode (Cathode)