Electrochemistry study of how electricity produces chemical reactions and chemical reactions produces electricity

1. Electrochemistrystudy of how electricity produces chemical reactions and chemical reactions produces electricity Involves redox reactions Electrochemical cell: any device which converts chemical energy into electrical energy or vs.

2. Electrochemical cell functions galvanically if voltage applied is less than equilibrium voltage Chemical  electrical energy spontaneous Electrochemical cell functions as an equilibrium cell if applied voltage equals equilibrium voltage No electric current flow Electrochemical cell functions electrolytically if voltage applied is greater than equilibrium voltage electrical  chemical energy Nonspontaneous

3. Copper placed into solution of AgNO3- Spontaneous reaction Grayish white silver deposit formed on copper Solution turns blue because of copper (II) nitrate 2Ag+ + Cu  Cu2+ + 2Ag No usable energy harnessed/dissipated as heat Same reaction can occur and produce electricity if placed in galvanic cell

4. Voltaicor Galvanic Cell • Redox reactions (chemical reactions in two half-cells) make electrons move from one substance to another • Spontaneous reactions provide electric energy or current • Electrical energy produced can do work (external electric current flow) • Current created by using two different metals that differ in tendency to lose electrons • First device was voltaic cell/use batteries now

5. Half cells • Electrode • Conductor of electrons in circuit • Wire • Carries electrons in external circuit • Electrolyte solution • Contains atoms, molecules, or ions (Red cat) Gives up electrons to ionic conductor. (An ox) Electrons generated.

6. Wire • Electrons flow from anode (-)  cathode (+) • Creates electricity • Voltages produced affected by • Different metals used for electrodes • Different solution concentrations • Solution (electrolyte) • Contains salts (ions) of electrode in solution • 2 containers store half-reactions • Works because solutions in half-cells remain electrically neutral

7. Salt Bridge • Conducting solution of soluble salt (KCl, NaCl, KNO3) • Electrically neutral • Electron flow always anode  cathode • Negative ions flow toward anode • Positive ions flow toward cathode

8. For voltaic cell to continue to produce external electric Anode builds up positive ions, while cathode adds more electrons Electrons attracted to anode and don’t move away Cathode’s negative charge repels electrons Build-up of charge doesn’t allow for flow of electrons and cell eventually fails Salt bridge contains ions that do not interfere with reaction taking place Movement of ions completes circuit

10. Electronegativity helps determine strongest or weakest oxidizer • Higher the electronegativity • More element will gain electrons • Is reduced • Is very good oxidizing agent • Lower the electronegativity • More element will give away electrons • Is oxidized • Is very good reducing agent • Why lithium is popular metal used in batteries

11. Electrochemical cell created by placing metallic electrodes into an electrolyte where a chemical reaction either uses or generates an electric current • Zn/Cu electrodes placed in solutions of their salts • Electrons flow thru external wire from Zn  Cu • Type of electrode determines if reaction takes place • Zn more readily loses electrons than Cu • Becomes positive ion • Goes into aqueous solution • Decreases mass of Zn electrode • Cu receives 2 electrons • Cu ion from solution  uncharged Cu atom • Deposits on Cu electrode • Increases mass of Cu electrode

12. Two half-reactions are typically written Zn(s)  Zn2+(aq) + 2e- Oxidation (loses electrons) Occurs at anode (- terminal of battery) Cu2+(aq) + 2e-  Cu(s) Reduction (gains electrons) Occurs at cathode (+ terminal of battery) Cu2+(aq) + 2e-  Cu(s) Zn(s)  Zn2+(aq) + 2e- Zn(s) + Cu2+(aq)  Zn+2(aq) + Cu(s)

13. Zn (s) l ZnSO4(aq)llCuSO4(aq) l Cu(s) Half-cell undergoing oxidation (anode) on left Double vertical line represents salt bridge Half-cell undergoing reduction (cathode) goes on right Single vertical lines indicates boundaries of phases in contact Zn rod/zinc sulfate solution are separate phases in physical contact Copper(II) sulfate solution/copper rod are separate phases in physical contact

14. http://college.hmco.com/chemistry/shared/media/animations/anodereaction.htmlhttp://college.hmco.com/chemistry/shared/media/animations/anodereaction.html

15. http://college.hmco.com/chemistry/shared/media/animations/cathodereaction.htmlhttp://college.hmco.com/chemistry/shared/media/animations/cathodereaction.html • Electricalconductor • K+ • Cl- SO4 Reducing agent Oxidizing agent Oxidation half-reaction Zn  Zn2+ + 2e- Reduction half-reaction Cu2+ + 2d-  Cu http://college.hmco.com/chemistry/shared/media/animations/electrochemicalhalf.html

16. Electric flow

17. Write anode/cathode reactions for a galvanic cell that utilizes reaction: Ni(s) + 2 Fe3+ Ni2+ + 2 Fe2+ • Oxidation occurs at anode: Ni | Ni2+ • Ni(s)  Ni2+(aq) + 2 e- • Reduction occurs at cathode: Fe3+ l Fe2+ • 2 Fe3+ + 1 e-  2 Fe2+ • For every Ni atom oxidized, two Fe3+ ions are reduced • Ni(s) | Ni2+(aq) || Fe3+(aq), Fe2+(aq) | Pt(s) • Inert Platinum electrode • Placed in solution • Provides e’s for reduction

18. http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/animations/ZnCutransfer.htmlhttp://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/animations/ZnCutransfer.html • http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/animations/PbAgtransfer.html • http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/animations/CuZncell.html

19. Comparison: AnodeCathode • Oxidation Reduction • Electrons produced Electrons consumed • Anions migrate Cations migrate toward - sign toward + sign

20. Electrode Potential • Lemon, apple, tomato, potato, pineapple, orange • Strips of magnesium, aluminum, zinc, iron, nickel, tin, copper • Multimeter • Make two slits few cm apart and place copper (cathode) and iron (anode) in each. • Connect leads to get a positive voltage reading. • Remove iron and replace with others. • Try with different foods.

21. Rank the cell voltage of the anodes tested when using copper as the cathode. • Why was copper a good choice for the cathode in this activity? • Compare the voltages of the following combinations of metals: magnesium-copper, zinc-copper, tin-copper. What might this suggest about the ability of magnesium, zinc, and tin to give up electrons? Consult the table of standard reduction potentials to check your answers. • Does the type of fruit affect the rank-order of electrode potential?

22. Standard Reduction Potential (E°cell) • Reduction potential • How easily substance accepts electrons & is reduced • Determines if proposed reaction under standard conditions will be spontaneous • Electrons always flow from half-cell w/lower (negative) standard reduction potential (oxidation side/anode) to half-cell w/higher (positive) reduction potential (reduction side/ cathode) • Cell potential • Measure of difference between two electrode potentials (between two points in V, volts) • Calculated as potential for reduction at electrode

23. Standard Hydrogen Electrode25oC/1 atm • Measure reduction potential of all electrodes against one standard electrode, (standard hydrogen electrode, SHE) • Voltage of single electrode cannot be measured because voltage is difference in electron pressure • Compare other electrode voltage to this to get standard electrode potential of any electrode compared to SHE • Sometimes called standard reduction potential because electrode reactions are usually written as reduction reactions

24. positive negative • cathode • anode • cathode • anode Electrode acts as oxidation or reduction half-reaction depending upon half-cell to which it is connected http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/animations/SHEZnV7.html http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/animations/SHECu.html

25. Simulation for voltaic cell • http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/electroChem/voltaicCell20.html • http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/electroChem/voltaicCellEMF.html • Gives reduction potentials/different molarities

26. Equilibrium Cell Voltages • Equilibrium cell • Neither producing nor using electrical energy • Neither electrolytic nor galvanic cell • Voltage called equilibrium voltage • No current flows when applied voltage exactly matches equilibrium voltage • If applied voltage differs from equilibrium voltage even a little bit, current flows in one direction

28. Values are ability to be reduced Positive cell potential Spontaneous reaction Occurs at reduction & contains cathode More positive E0 values, greater tendency for substance to be reduced Good oxidizing agents (in general, nonmetals) Negative cell potential Nonspontaneous reaction (reverse is spontaneous) Occurs at oxidation & contains anode Good reducing agents (in general, metals) Any electrode can be oxidized by any other electrode whose SRP is greater than its own value, and can be reduced by any electrode with SRP less than its own value

29. Which of the following elements would be most easily oxidized: Ca, Cu, Fe, Li, or Au? • Use reduction potential chart: • Nonmetals at top and most easily reduced • Metals at bottom and most easily oxidized • Li-bottom, most easily oxidized • Order from most easily to least easily oxidized: Au, Fe, Cu, Ca, Li • Which one of the following would be the best oxidizing agent: Ba, Na, Cl, F, or Br? • Oxidizing agents elements most easily reduced (F)

30. To calculate standard cell potential for reaction • Electrical potential: ability of voltaic cell to produce electrical current • Write oxidation/reduction half-reactions for cell • Look up reduction potential, Eoreduction,  for reduction half-reaction • Look up reduction potential for oxidation half-reaction/reverse sign to obtain oxidation potential Eooxidation = - Eoreduction. • Add potentials of half-cells to get overall standard cell potential Eocell = Eoreduction + Eooxidation

31. Find standard cell potential for following electrochemical cell Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) Write half-reactions for each process Zn(s)  Zn2+(aq) + 2 e- Cu2+(aq) + 2 e-   Cu(s) Look up standard potentials for redcution half-reaction Eoreduction of Cu2+ = + 0.339 V Look up standard reduction potential for reverse of oxidation reaction and change the sign Eoreduction of Zn2+ = - 0.762 V Eooxidation of Zn = - ( - 0.762 V) = + 0.762 V Add cell potentials to get overall standard cell potential oxidation:  Zn(s)  Zn2+(aq) + 2 e- Eoox= - Eored = - (- 0.762 V)=+ 0.762V  reduction: Cu2+(aq) + 2 e-  Cu(s)  Eored. = + 0.339 V________ overall: Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) Eocell = + 1.101 V

32. Example: • Zn2+(aq) + Pb(s) Zn(s) + Pb2+(aq) • Oxidation half-reaction: Pb2+(aq) + 2e- Pb(s) • E° = –0.1262 V • Reduction half-reaction: Zn2+(aq) + 2e- Zn(s) • E° = –0.762 V • Calculate standard cell potential • E°cell = E°red + E°ox • = –0.762V + 0.1262V • = –0.636 V • Overall reaction is not spontaneous, because standard cell potential is negative

33. Calculate cell potential for following electrochemical cell • Before you add half-reactions, must take care that electrons cancel • Mg l Mg2+ ll Cr3+ lCr • Oxidation half-reaction: 3Mg  3Mg2+ + 6e- • E° = -2.372 V • Reduction half-reaction: 2Cr3+ + 6e- 2Cr • E° = -0.744 V • E°cell = E°red + E°ox • = -0.744 V + 2.372 V • = 1.628 V • Overall reaction is spontaneous, because standard cell potential is positive

34. (lose electrons best) Most active nonmetals (gain electrons best) F2 Cl2 Br2 I2 Least active nonmetals

35. Homework: Read 21.1, pp. 662-672 Q pg. 672, #9, 10, 11 Q pg. 692, #30, 43-46

37. Batteries Common batteries are one or more voltaic cells Electrolytes used as energy storage medium When two electrolytes, each containing ions with different oxidation states, are allowed to exchange charges, electric current is formed (redox)

38. Dry Cell (Leclanche cell) are voltaic cells Inactive Starch thickens it to paste-like consistency to prevent leaks

39. Battery: one/more galvanic cells connected in series Produce/store electric energy chemically Primary battery Not easily reversed Deliver current until reactants gone and battery is discarded Secondary battery (storage battery) Rechargeable (reversible redox reactions) Alkaline battery KOH/NaOH replaces NH4Cl http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/animations/flashlightV8.html

42. Mercury Battery • Used in medicine/ electronics (hearing aids/pacemakers/ light meters/electric watches) • No change in electrolyte composition during operation • Overall reaction involves only solid substance (Strong alkaline electrolyte)

43. Lead Storage Battery • Series of 6 identical cells, 2V each (12V) • Anode/cathode in aqueous solution of sulfuric acid which acts as electrolyte • Each electrode consists of several grids w/large surface area to deliver high currents required to start engine (Pb goes from 0  +2) (Pb goes from +4  +2) www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/animations/PbbatteryV9web.html

44. Could go on forever but • Lead sulfate formed in reduction falls to bottom of battery • Lead electrodes get thinner/more brittle over time • Eventually break, short out cell, battery needs replacing • “Gone dead” in cold climates • EMF of many electrochemical cells decrease w/decreasing temperature • Electrolyte must be fully conducting to function properly • If cold, ions move more slowly through viscous medium, decreasing power output of battery • If warmed, recovers normal power

45. Solid-state Lithium Battery • Solid electrolyte (not aqueous solution or water-based paste) • Lithium used as anode because most negative E0 value

47. Fuel Cells • Do not store chemical energy as batteries do • Reactants constantly resupplied • Cell doesn’t “die” when oxidized substance used up • Products constantly removed • Up to 70% efficient KOH (aq)

49. Corrosion…loss of metal resulting from an oxidation-reduction reaction of the metal with substances in the environment …and preventing it