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Electrochemical prosses

Electrochemical prosses. The electrochemical processes are called redox reactions which is accompanied by the release or absorption of electricity. The electrodes are the main part of the electrochemical processes.

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Electrochemical prosses

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  1. Electrochemical prosses

  2. The electrochemical processes are called redox reactions which is accompanied by the release or absorption of electricity. • The electrodes are the main part of the electrochemical processes. • The electrode is a conductive material (metal, carbon) in contact with an electrolyte.

  3. Redox processes are the chemical processes associated with the electrons transfer Zn + Cu+2 = Zn+2 + Cu Red1 + Ox2 = Ox1 +Red2 Oxidation: Red1- ne = Ox1 Reduction: Ox2+ne = Red2 symbolic notation: Ox1/Red1; Ox2/Red2 Zn+2/Zn ; Cu+2/Cu

  4. Ox1/Red1Ox2/Red2 Zn+2/Zn Cu+2 / Cu An electric double layer is formed at the interface as due to release of metal ions (Zn +2 / Zn), and (Cu +2 / Cu) by adsorption of metal ions on the surface Zn - 2e = Zn2+Cu - 2e = Cu2+

  5. Equilibrium is established when the electrode potential reaches a value which is capable of stopping the oxidation (or reduction), which occurs the first time after the immersion of the metal in solution

  6.  1  2 The difference between the charges at the interface determines a potential jump  

  7. The magnitude of the potential difference is called an electrode potential ()  depends on: •  the nature of the electrode material • the concentrations of ions in solution • temperature • PH

  8. The absolute value of an electrode potential can not be measured •   The electrode potential measured under standard conditions (T = 298K, P = 1 atm, C = 1M) with respect to another electrode called a standard electrode potential

  9. 2H++2eH2 • °(2H+/H2) = 0V • under standard conditions • P(Н2)= 105 Pа • T = 298K • C = 1 mol/L Standard hydrogen electrode

  10. Measurement of an electrode potential Н2 - 2е = 2Н+

  11. A series of redox potentials Half-reaction0,V Br2 + 2e = 2Br- 1,09 BrO3 + 5H+ + 4e = HBrO + 2H2O 1,45 HBrO + H+ + 2e = Br + H2O 1,33 Cl2 + 2e = 2Cl 1,36 ClO4 + 8H+ + 8e = Cl + 4H2O 1,3 2ClO4 + 16H+ + 14e = Cl2 + 8H2O 1,34 ClO4 + 4H2O + 8e = Cl + 8OH 0,56

  12. A series of redox potentialsof metals Elecrode processо,VElecrode processо,V К - е- = К+ - 2,92 Co - 2e- = Co2+- 0,28 Ba - 2e- = Br2+ - 2,91 Ni - 2e- = Ni2+- 0,25 Ca - 2e- = Ca2+- 2,87 Sn - 2e- = Sn2+- 0,14 Na - e- = Na+ - 2,81 Pb - 2e- = Pb2+- 0,13 Mg - 2e- = Mg2+- 2,36 H2 - 2e- = 2H+0,00 Be - 2e- = Be2+- 1,85 Bi - 3e- = Bi3+ 0,22 Al - 3e- = Al3+- 1,66 Cu - 2e- = Cu2+ 0,34 Mn - 2e- = Mn2+- 1,18 Ag - e- = Ag+0,80 Zn - 2e- = Zn2+- 0,76 Hg - 2e- = Hg2+ 0,85 Fe - 2e- = Fe2+- 0,44 Pt - 2e- = Pt2+1,19 Cd - 2e- = Cd2+- 0,40 Au - 3e- = Au3+ 1,50

  13. A series of redox potentialsof metals Redactive activity metals grow Oxidative activity cations of metals grow

  14. At series of redox potentialsof metals: • Each metal displaces another metal from salt solution which have a larger value of the electrode potential Al + Hg(NO3)2 = Al(NO3)3 + Hg • Metals having (-) potentials displace hydrogen from acid solutions of weak oxidants Zn + H2SO4(dil.) = ZnSO4 + H2

  15. Gibs Energy for electrochemical processes G°= -zF°(stand.cond.) G = -zF (unstand.cond.) z-the least common multiple of the number of electrons in the two half-reactions F – Faraday’s number  = cathod - anode

  16. [Ox]x[Red]y RT nF Nernst’s equation is the calculation of an electrode potential of any conditions =°+ ln [OX] и [Red]equilibrium concentration of the oxidized and reduced forms n- the number of electrons in a half-reaction

  17. For metal electrode F =96487 kL.mol-1, R=8,31 Jmol-1 K-1, T =298K

  18. Chemical power sources • Galvanic Cells • Concentration elements • Fuel cells • Batteries • Due to the direct conversion of chemical energy into electrical Chemical power sources have a high coefficient of performance (about 70-90%)

  19. Galvanic Cell

  20. The electrochemical cell comprises: anode (oxidation), cathode (the recovery process). Symbolic notation: (anode)(-)Zn|Zn2+||Cu2+|Cu(+)(cathode) or (-)Zn|ZnSO4||CuSO4|Cu(+)

  21. The cause of the electric current flow in a galvanic cell is the difference between the electrode potentials (electromotive force - EMF) two Redox systems interconnected EMF = Е=  = cathode - аnode> 0 EMF = Е= (Сu) - (Zn)= 0,34 – (-0,76)= 1,1V G = -zFЕ

  22. Modern Chemical power sources • They have a common electrolyte (alkali) and steel as a construction material buildings and shunts. • Material (-) electrode (reducing agent) usedMg, Zn, Pb, Fe, Cd. • Material (+) electrode (oxidant) was used PbO2, HgO, MnO2, CuCl, etc.

  23. Lead Battery The prosses is based on two reactions: PbO2 + PbPbO Electrolyte H2SO4 (33-37%) H2SO4 HSO4  + H + Charging Ox.RedDischarging

  24. Device Battery(-)Pb│PbSO4│H2SO4 │ PbO2 │ Pb(+) • Cellular alloy plate(Pb + trace of Sb). • A mixture of glycerol with PbO2is filled in the cell • Plates are collected in batteries and dipped in a solution H2SO4 ( = 1,18 g/cm3). • Battery is charged by passing an electric current.

  25. ChargingBattery • (-) cathode: PbSO4 + 2e = Pb + SO42 • (+) anode: PbSO4- 2e + 2H2O = PbO2 + SO42 + 4H+ DischargingBattery (-) anode:Pb 2e + SO42= PbSO4 (+) cathode:PbO2 + 2e + SO42 + 4H+= PbSO4+ 2H2O For a chargedbattery:  = 2,1V When the battery is charging, electrolyte density increases and decreases during discharge. Knowing the density we can judge the state of charge.

  26. Problem • The galvanic cell consists of Mg and Ag electrodes in solutions of their nitrates. The concentration of [Mg2+] = 0,5M; [Ag+] = 2,0M. To write the electrochemical scheme of a galvanic cell, the redox reaction and calculate the EMF of a galvanic cell.

  27. Corrosion of Metals

  28. Corrosion of metals - is the process of spontaneous fracture of metals under the influence of the environment, which is accompanied by the release of energy and material dispersion • Corrosion processes are irreversible, in accordance with the 2nd law of thermodynamics (entropy increase)

  29. Classification of corrosion processes according to mechanism: 1) chemical       a) in solution nonelectrolytes       b) gas   2) electrochemical       a) atmospheric,       b) soil,       c) stray currents

  30. Chemical and electrochemical corrosion is determined by Gibbs free energy change: G°= -RTlnK = - nF° • Chemical corrosion is determined by the equilibrium constant of reversible heterogeneous reactions • Electrochemical corrosion is determined by an electrical work • The kinetics process is very important

  31. Chemical corrosion • occurs in corrosive atmospheres at high temperatures • aggressive organics nonelectrolytes xMe + yO2 = MexOy G° < 0 • The examples: • destruction of combustion engines • destruction of cutting tools • gas turbine blades, nozzles, the exhaust nozzles

  32. Electrochemical corrosion - is the destruction of the metal in the electrolyte medium. • Mechanism of the process: •   anodic oxidation of the metal: Ме - nе = Меn+ cathodic reduction of oxidant: Ox + ne = Red O2 + 2H2O + 4e = 4OH- 2H+ + 2e = H2

  33. Processes of electrochemical corrosion are similar to the processes occurring in the cell. • There are formed microgalvanopairs. • The difference - the absence of the external circuit, i.e. electrons do not come out of the metal but move inside the metal. • Chemical energy of oxidation of the metal is not transmitted in work, but only in heat. • It is possible the secondary chemical reactions: Men+ + nOH- = Me(OH)n

  34. Dangerous areas • point of contact with different metals • areas with different thermal and mechanical treatment • sections stained with oxides and other mineral pigments • metal alloy inhomogeneity Strong corrosive properties are: • seawater • technological solutions. (acid salts, etc.) • groundwater • wastewater • damp air

  35. Corrosion of iron by atmospheric oxygen dissolved in the water air rust Drop of water Ion (Fe)

  36. Soil corrosion of Fe under acidic conditions Mikrogalvanich cellscheme: (-)Fe|Fe2+,2H+,СO32-,SO42-|Fe(+) (- ) Fe - 2e = Fe2 (+ ) 2H++ 4e = H2 Fe2 + SO42- = Fe SO4

  37. Corossion of Fe together with Cu Anode(-) Fe | H2SO4 | Cu (+)Cathode A(-) Fe - 2e = Fe2+°(Fe2+/Fe) = - 0,44V (+)K 2H+ + 2e = H2°(2H+/H2) = 0,0V H+ Fe-anode Cu cathode H+

  38. Methods of Corrosion Protection • Creating a rational constructions • Impact on the environment • Inhibitors • Protective Coating: • lubrication varnishes, paint, polymers, oxidation, phosphating, metal coatings • Protection of the external potential:tread, the current source

  39. Anodic protectionExample: galvanized iron air Drop of water

  40. Cathodic protectionExample: Corrosion "tinplate"

  41. Cathodic protection against corrosion

  42. Electrolysis • Electrolysis is not a spontaneous process (G 0) • This is a redox process called by the electric current passing through a solution or molten electrolyte •   The driving force is the electrolysis voltage applied to the electrodes, which causes the moving of cations and anions towards the cathode and the anode

  43. cathode< equil. The current source а > equil. Cathode (-) Anode(+) Inert electrodes

  44. Factors influencing the electrolysis • electrolyte composition • electrode material • temperature • voltage • the current density etc.

  45. Discharging occurs simultaneously with cations discharging anions; therefore externally imposed voltage is divided into two parts extending at an anodic oxidation and cathodic reduction of the ions or molecules sometimes.

  46. Electrolysis of a fused salt 2Cl-2e=Cl2 Na+ + e=Na 2NaCl = 2Na + Cl2

  47. Sequence ofion discharging If the solution of two or more kinds of cations and anions is most likely the process with a minimum expenditure of energy • At the cathode, reducing cation in order to reduce their potential • Anions are oxidized at the anode in the ascending order of their potential

  48. At series of redox potentialsof metals The order of the oxidation The order of the reduction

  49. Electrolysis of a water solution

  50. Water, as part of the electrolyte solution in the electrode, participates processes • Oxidation of water: 2H2O - 4e = O2 + 4H+o = + 1,23 V • Portentials: F–, SO42-, NO3-, PO43-, CO32- etc. >than+ 1,23 В • Consequently, these anions can not be oxidized in aqueous solution

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