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Infection and response
Organisation
Cell biology
Biological molecules
Homeostasis and response
Responding to change (a2 only)
Energy transfers (a2 only)
The control of gene expression (a-level only)
Substance exchange
Bioenergetics
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1f industrialisation and the people: britain, c1783-1885
1c the tudors: england, 1485-1603
Britain & the wider world: 1745 -1901
Inter-war germany
World war one
2n revolution and dictatorship: russia, 1917-1953
World war two & the holocaust
1l the quest for political stability: germany, 1871-1991
The cold war
2j america: a nation divided, c1845-1877
Medieval period: 1066 -1509
The fight for female suffrage
2m wars and welfare: britain in transition, 1906-1957
2d religious conflict and the church in england, c1529-c1570
Britain: 1509 -1745
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Electrode potentials 1 Electrochemical Cells Can be made of 2 different metals dipped in salt solutions of their own ions + connected by wire (external circuit) Always 2 reactions occurring, one is oxidation other is reduction (REDOX) When piece of metal dipped into solution of metal ions equilibrium set up There is tendency of metal to form positive ion + go into solution There is tendency of metal ions in solution to gain electrons + form positive ions Setting up electrochemical cell 2 half cells joined together to give complete circuit 1. 2 metals joined with wire (electrons flow through wire) 2. 2 solutions joined with salt bridge (ions flow through it) 3. Voltmeter used in circuit, allows potential difference (emf) to be measured Salt Can be Piece of filter paper soaked with solution of unreactive ions Tube containing unreactive ions in aged gel Compounds KNO3 often used in salt bridge and K+ and NO3- quite unreactive Copper strip- Solution of- Cu² (aq) (1 mol dm ³) High-resistance voltmeter Salt bridge Zn loses electrons more easily than copper (Zn more reactive) So heals cell on right oxidised Zn to Zn ²+ This releases electrons into external circuit Electrons flow through wire form more reactive element to less reactive element Solutions connected by salt bridge (e.g. strip of filter paper soaked in KNO3 (aq) This allows ions to flow...
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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)