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ELECTROCHEMISTRY

Explore the field of electrochemistry and its applications in industrial processes, including metals production, gas and inorganic compound production, organic compound synthesis, electrodeposition, and electrical energy production. Learn about redox reactions, oxidation states, galvanic cells, and the measurement of standard reduction potentials.

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ELECTROCHEMISTRY

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  1. ELECTROCHEMISTRY Valentim M. B. Nunes Engineering Unit of IPT March 2018

  2. Introduction Electrochemistry  Field of Chemistry that deals with the relationships between electrical and chemical energy. Electrochemical Processes: • Spontaneous reactions in which energy released by a reaction is converted into electricity. • Processes in which electricity is used to force a non spontaneous chemical reaction. Electrochemistry fields: electroanalysis, electrocatalysis, electrosynthesis, etc..

  3. Industrial Processes with basis in Electrochemistry (Examples…) • Metals production: aluminum (Al), sodium (Na), lithium (Li), magnesium (Mg), zinc (Zn) and cooper (Cu); • Production of gases: hydrogen (H2) and chlorine (Cl2); • Production of inorganic compounds: sodium hydroxide (NaOH), sodium hypochlorite (NaOCl); • Production of organic compounds: organophosforide compounds; • Metallic electrodeposition: to modify appearance, hardness, resistance to corrosion, including modern electronic microcircuits; • Electrochemical production of electrical energy: piles, batteries, accumulators.

  4. Redox reactions Redox reactions are those in which there are a transfer of electrons between species. Zn (s) + Cu2+(aq)  Zn2+(aq) + Cu(s) Oxidized specie – looses electrons (reduction agent!) Reduced specie – gains electrons (oxidation agent!) Corrosion! Combustion?

  5. Oxidation state The oxidation state of an element is the charge that can be assumed to him if in a given chemical bond all the electron were attributed to the most electronegative atom. • The oxidation state of any element in any allotropic form is zero (O2, Fe, Al, S8, etc…) • The oxidation state of a monoatomic ion is the correspondent charge (Al3+, Fe2+, O2-, …) • The oxidation sate of hydrogen is always +1 except in hydrides (LiH, CaH2,..) were is -1. • The oxidation state of oxygen is -2 in all compounds except in peroxides (H2O2) were is -1. • All the other oxidation states are calculated considering that the algebraic sum of oxidations state corresponds to the charge of the molecule or polyatomic ion (CO2, MnO4-, NO3-, NH4+, …)

  6. Langmuir Plot Example: Chlorine can assume oxidations state from -1 until +7.

  7. Galvanic cells (production of electrical energy) If the chemical reaction previously described occurs with direct contact of reagents no net work is obtained. But, redox reactions can occur with reagents connected by an electric conductor. Daniell Cell

  8. How it works? Cathode: electrode were the reduction semi-reaction occurs. Cu2+(aq) + 2 e-  Cu(s) Salt bridge: to keep the electro neutrality of solutions. Anode: electrode were the oxidation semi-reaction occurs. Zn(s)  Zn2+(aq) + 2 e- The fact that the electrons are moving indicates that there are a potential difference between the two electrodes; cell potential or electromotive force (emf)

  9. Redox equations balance • Separate the full equation in two semi-equations. • Balance all atoms except O and H, in each semi-equation. • Balance oxygen. For reactions in acidic medium add H2O, and in basic medium add OH-. • Then balance Hydrogen. For reactions in acidic medium add H+ and in basic medium add H2O. • Balance each semi equation for electrical charge adding electrons. • Sum the two semi-equations, eliminating electrons from final equation, using appropriate coefficients.

  10. Example – balance of an equation in acidic medium Balance the equation for the oxidation of cooper by NO3- in acid medium . Cooper is oxidized to Cu2+ and NO3- is reduced to NO. Cu  Cu2+ NO3-  NO Cu  Cu2+ + 2 e- NO3- + 4 H+ + 3 e-NO + 2 H2O ( x3) Cu  Cu2+ + 2 e- ( x2) NO3- + 4 H+ + 3 e-NO + 2 H2O 8 H+ + 2 NO3- + 3 Cu  2 NO + 3 Cu2+ + 4 H2O

  11. Example – balance of an equation in basic medium Balance the equation for the oxidation of Mn2+ by hydrogen peroxide in basic medium. Mn2+ is oxidized to MnO2 and the peroxide reduced to H2O Mn2+ MnO2 H2O2  H2O TRICK: for each OH- you need to add multiply immediately for 2 and add one H2O molecule the other side of equation! Mn2+ + 4 OH-MnO2 + 2 H2O + 2 e- H2O2 + H2O + 2 e-  H2O + 2 OH- Mn2+ + H2O2 + 2 OH- MnO2 + 2 H2O

  12. Spontaneity of reactions Returning back to the Daniell Cell we can rise two important questions: • Why is the cooper ion, Cu2+, that reduces in contact with metallic zinc? • If the concentrations of cooper and zic ions are unitary ( = 1 M), what will be the electromotive force of the cell?

  13. Standard Reduction Potential, Eº Its impossible to measure the absolute value of the potential of an electrode but only potential differences. So its necessary to define a standard electrode that allow us to calculate all the other potentials for redox pairs. The Standard Hydrogen Electrode (SHE) has, by International convention, a standard reduction potential of Eº = 0 V, corresponding to the reversible reaction: 2 H+(aq, 1M) + 2 e-  H2(g, 1 atm)

  14. Measurement of the standard reduction potentials Galvanic cell that permits the calculation of the standard reduction potential of the redox pair Cu2+/Cu, EºCu2+/Cu Eºcell = Eºcathode - Eºanode Eºcell = EºCu2+/Cu - 0 V EºCu2+/Cu = +0.34 V cathode anode

  15. Table of standard reduction potentials at 25 ºC Standard conditions: dissolved species, 1M and gases, p = 1 atm Higher Eº means higher tendency to suffer reduction!

  16. Calculating the electromotive force of a cell Daniell cell: Eºcell = Eºcathode - Eºanode Eºcell = EºCu2+/Cu - EºZn2+/Zn Eºcell = 0.34 V - (-0.76 V) = 1.10 V

  17. Another example EºAg+/Ag > EºCu2+/Cu, then in the silver electrode occurs the reduction (cathode) and in the cooper oxidation (anode) 2 Ag+(aq) + Cu(s)  Cu2+(aq) + 2 Ag(s) Eºcell = Eºcathode - Eºanode Eºcell = 0.80 V - 0.34 V = 0.46 V

  18. Diagonal rule The chemical reactions in the standard potential table are reversible. For instance, cooper can work as the cathode or as the anode in a galvanic cell. DIAGONAL RULE In standard conditions any specie located in the left in a given reaction will spontaneously react with any specie located in the right and below in the table.

  19. Displacement of hydrogen 2 Na(s) + 2 H2O(l)  2 NaOH(aq) + H2(g)  Cd(s) + H2O(l)   Cd(s) + 2 HCl(aq)  CdCl2(aq) + H2(g)  Cd(s) + 2 H+(aq)  Cd2+(aq) + H2(g) Ag(s) + 2 H+(aq)  

  20. Displacement of metals Zn (s) + CuSO4(aq)  ZnSO4(aq) + Cu (s)  Cu(s) + Zn2+(aq)   Cu (s) + 2 Ag+(aq) Cu2+(aq) + Ag(s)

  21. Displacement of halogens Cl2(g) + 2 Br-(aq) 2 Cl-(aq) + Br2(l) Cl2 (aq) + 2 NaI 2 NaCl (aq) + I2 (s)

  22. Electric energy In a galvanic cell chemical energy is converted in electric energy. Electric energy = emf total charge that crosses the circuit Electric energy = volts  coulombs = joules Total charge = nF, n is the number of moles of electrons and F is the Faraday constant. 1 F  96500 C/mol of e- welectric = - nFEcell

  23. Equilibrium constant In standard conditions: welectric = - nFEºcell For a spontaneous process, Eºcell > 0 K Eºcell Reaction > 1 Positive Spontaneous =1 0 Equilibrium < 1 Negative Non spontaneous -nFEºcell = - RT ln K At 25 ºC

  24. Nernst equation Many reactions occurs in non standard conditions! -nFE = -nFEº + RT ln Q in which Q is the reactional quotient.

  25. Examples 1. Writ the Nernst equation for the reaction that occurs in the Daniell cell. 2. Predict if the reaction Cd(s) + Fe2+(aq)  Cd2+(aq) + Fe(s) occurs spontaneously at 25 ºC when [Fe2+] = 0.6 M and [Cd2+] = 0.01 M. Since E > 0, then the reaction is spontaneous

  26. Involving gases If there are gases involved in the reaction their concentrations must be expressed in atmosphere. What is the emf of a cell made with the semi-cell Cd/Cd2+ and the semi-cell Pt/H2/H+ if [Cd2+] = 0.2 M, [H+] = 0.16 M and PH2 = 0.8 atm? 2 H+(aq) + Cd(s)  Cd2+(aq) + H2(g)

  27. Batteries A battery is a galvanic cell, or a set of galvanic cell connected in series, that furnishes continuous current at constant potential.

  28. Dry Cell Battery (Leclanché cell) Utilized in lanterns, toys, portable devices, etc... Anode: Zn(s)  Zn2+(aq) + 2 e- Cathode: 2 NH4+(aq) + MnO2(s) + 2 e-  Mn2O3(s) + 2 NH3(aq) + H2O(l) Ecell 1.5 V

  29. Mercury Battery Utilized in medicine (pacemakers), electronic industry, etc... Anode: Zn(Hg) + 2 OH-(aq)  ZnO(s) +H2O(l) + 2 e- Cathode: HgO(s) + H2O(l) + 2e-  Hg(l) + 2 OH-(aq) Global: Zn(Hg) + HgO(s)  ZnO(s) + Hg(l) Ecell 1.35 V

  30. Lead Storage Battery (automobile battery) Eºcélula = EºPbO2/PbSO4 - EºPbSO4/Pb Eºcélula = 1.74 - (-0.28)  2 V Ebateria = 6  2 V  12 V Anode: Pb(s) + SO42-(aq)  PbSO4(s) + 2 e- Cathode: PbO2(s) + 4 H+(aq) + SO42-(aq) + 2e-  PbSO4(s) + 2 H2O(l) Global: Pb(s) + PbO2(s) + 4 H+(aq) + 2 SO42-(aq) 2 PbSO4(s) + 2 H2O(l) discharge charge

  31. Ni-Cd Batteries (rechargeable)

  32. Solid-State Lithium Batteries Utilize a solid in contact with the electrodes. The solid is a polymeric material that allows the crossing by Li+, but not by electrons. Ecell 3 V Anode: Li  Li+ + e- Cathode:TiS2 + e-  TiS2-

  33. Fuel Cells A fuel cell its a galvanic cell that needs the continuous supply of reagents to work. Anode: H2(g)  2 H+ + 2 e- Cathode O2(g) + 4 H+ + 4 e-  2 H2O Global: 2 H2(g) + O2(g)  2 H2O(l) Eºcell = Eºcathode - Eºanode Eºcell = 1.23 V - 0 Eºcell = 1.23 V

  34. Applications “Fuel Cells” were utilized by the Apollo Missions to supply energy and water for the astronauts Public transport vehicle supplied with hydrogen cells. 

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