Electrochemistry experimen Essay

Introduction: Redox reactions are reactions where the oxidation states of the atoms change. The atoms are either oxidized or reduced, depending if they lose or gain electrons. Electrochemical cells are devices that cause a current from redox reactions. It is set up so that electrons lost from one of the reagents can travel to another reagent. This creates a voltage, which is also known as the electric potential difference. This voltage can be read if a high-resistance voltmeter or multimeter is connected to the circuit. Salt bridge is used to allow migration of ions between two electric cells to maintain neutrality of solutions. It is usually made up of a filter paper moistened with an inert solution or an inert solution/gelatine salt bridge to prevent oxidation of certain ions. This experiment is divided into 2 parts: part A and part B. Part A Objective: To investigate the effect of change in lead(II) ion concentration on the potential of the Pb2+(aq) |Pb(s) electrode Introduction: This experiment investigates the e. m. f. of the cell: Cu(s) |Cu2+(aq) |Pb2+(aq)|Pb(s). Keeping the ion concentration in the copper electrode system constant(1M) and varying the ion concentration in the lead electrode system, the effect of change in lead(II) ion concentration on the potential of electrode as well as the Kc of the above reaction can be found. Chemicals: Copper foil x1, lead foil x1, 1M Cu2+ solution, 0. 1M Pb2+ solution, saturated potassium nitrate solution Apparatus: 250 cm3 beakers, 50cm3 beakers multimeter, distilled water bottle, filter papers, electrical wires with electrode holders, forceps, 100ml volumetric flask,10ml pipette x2, dropper Procedure: 1. The 0. 1M, 0. 01M, 0. 001M, 0. 0001M Pb2+ solutions were prepared from 0. 1M Pb2+ solution by dilution(1 portion solution plus 10 portions water). 2. The copper and lead electrodes were cleaned with a sand paper. 3. The circuit was connected as the diagram below. 4. The e. m. f. was recorded when the data shown was stable. 5. The above steps(2-4) were repeated with 0. 01M, 0. 001M and 0. 0001M Pb2+ solutions. Safety precaution: Potassium nitrate solution: Contact with combustible material may lead to fire 1M Copper(II) sulphate solution: harmful and irritating to eyes and skin. –>Safety goggles should be worn. Experimental set-up Results: [Pb2+]/M log[Pb2+] E/V 0. 1 -1 0. 482 0. 01 -2 0. 502 0. 001 -3 0. 521 0. 0001 -4 0. 545 Graph of E against log [Pb2+]: Trend shown: Given the ion concentration in the copper electrode system constant, it is found that the potential of the cell drops as the lead(II) ion concentration increase. The electrode potential is inversely proportional to the ten folds of ion concentrations. Calculation: When the reaction Cu2+(aq) + Pb(s) –> Pb2+(aq) + Cu(s) achieves equilibrium, the net e. m. f. of the cell=0 volt. The equilibrium expression of this reaction is: Kc= [ Pb2+(aq)][ Cu(s)] /[Cu2+(aq)][ Pb(s)] The effective concentration of Pb(s)/ Cu(s) are independent of its amount present and can be considered as constant. This reduces the expression to Kc= [ Pb2+(aq)] /[Cu2+(aq)] = 1Ãâ€"1021 Conclusion The potential of the cell decreases ad the ion concentration of Pb2+ increases. Further Analysis: Using the Nernst equation: E=E? -0. 059/n log[ox]/[red], Take [Pb2+]=0. 1 M as an example, E=0. 47-0. 059/2 log(0. 1/1)=0. 4405(V) [Pb2+]/M log[Pb2+] E/V(calculated) E/V(measured) %difference 0. 1 -1 0. 4405 0. 482 9. 42% 0. 01 -2 0. 529 0. 502 5. 10%.0. 001 -3 0. 5585 0. 521 6. 71% 0. 0001 -4 0. 588 0. 545 7. 31% It is shown that difference is present between the calculated value and measured value. This may be because of different conditions, resistance of the multimeter or errors in preparation of various concentrations of solutions. Part B Objective: To find out the equilibrium constant by e. m. f. measurement Introduction: The equilibrium constant for the below reaction is found out: Ag+(aq) + Fe2+(aq) Fe3+ (aq) + Ag(s) By e. m. f. measurement on the cell Pt |Fe2+(aq), Fe3+(aq)|Ag+(aq)|Ag(s) Chemicals: 0. 1 M Fe3+ solution, 0. 2 M iron(II)sulphate, 0. 2M barium nitrate, 0. 4M silver nitrate, platinum electrode, silver electrode Apparatus: gelatine salt bridge, 250 cm3 beakers, 50cm3 beakers ,multimeter, distilled water bottle, electrical wires with electrode holders, forceps, 10ml pipette x2 Procedure: 1. Equal volumes of 0. 2M FeSO4 and 0. 2M Ba(NO3)2 were mixed and the precipitate was allowed to settle without disturbance. 2. Equal volumes of 0. 1M iron(II) nitrate solution obtained and the iron(III)nitrate solution were mixed. This was the Fe2+(aq)/Fe2+(aq) half-cell. 3. The 0. 4M, 0. 2M, 0. 1M, 0. 05M, 0.025M silver nitrate solutions were prepared from 0. 4M silver nitrate solution by dilution. 4. The silver electrode was cleaned with a sand paper. 5. The circuit was connected as the diagram below. 6. The e. m. f. was recorded when the data shown was stable. 7. The above steps (4-6) were repeated with 0. 2M, 0. 1M, 0. 05M, 0. 025M silver nitrate solutions. Safety Precaution: Silver nitrate: harmful and oxidizing; 1) Poisonous if swallowed or inhaled 2) Skin contact with silver nitrate solid or solutions is likely to leave silver stains on the skin. Barium nitrate is poisonous and very harmful if swallowed. It is also a strong oxidizer, so may be hazardous if mixed with flammable materials. Experimental Set-upResults: [Ag+]/M log[Ag+] E/V 0. 4 -0. 398 0. 023 0. 2 -0. 699 0. 005 0. 1 -1 -0. 030 0. 05 -1. 301 -0. 050 0. 025 -0. 025 -0. 053 Graph of E against log[Ag+(aq)]: Calculation: When the reaction reaches equilibrium, both forward and backward reactions proceed to the same extent. This means that both the half cell reactions would have the same potential to proceed, so that the net e. m. f of the cell =0 volt at equilibrium. From the graph, the x-intercept is log[Ag+(aq)]=-0.72, hence,[Ag+(aq)]eqm=0. 1905 Ag+(aq) + Fe2+(aq) Fe3+ (aq) + Ag(s) KC= [Fe3+ (aq)]/ [Ag+(aq)][ Fe2+(aq)] =0. 1/ (0. 1905X0. 1) =5. 2493(mol2dm-6) ~5. 25(mol2dm-6) Conclusion: The equilibrium constant for the reaction between Fe2+/Fe3+ and Ag+ is 5. 25 mol2dm-6. Discussion(for both parts): 1. Possible Errors: -The electrodes (Cu(s),Pb(s),Ag(s)) were not cleaned very well with a sand paper so that they are not conducting electricity in all parts. The e. m. f. measured may hence be underestimated. -The same ammonium nitrate/gelatine salt bridge was used several times in part II of experiment. Ions of previous measurement may remain in the salt bridge and change the concentration of ions in the next measurement. -The apparatus including pipettes, beakers and volumetric flask may not be washed to be very clean. The error in concentration may be enormous when handling very dilute solutions (e. g. 0. 001M, 0. 0001M, 0. 025M) -The electrode touched the salt bridge once so that the accuracy of measurement of e. m. f. was affected. The solution is not exactly passed. -Air gap may be present in the ammonium nitrate/ gelatine salt bridge, increasing the resistance of ion migration. 2. Difference in effect of ion concentration on electrical potential of cell: -In part A, the cell e. m. f. increases as [Pb2+(aq)] decreases. However, in part B, the e. m. f. drops as [Ag+(aq)]. This suggests that when the species is a stronger oxidizing agent in the reaction and undergoes reduction, the presence of its ions tends to increase the e. m. f. whereas the one which undergoes oxidation tends to reduce the e. m. f. The species with more positive standard reduction potential has a positive effect but the one with less positive potential has a negative effect. Reference: Physical Chemistry II by TM Leung and CC Lee( p. 295-298 &299-301).