Electrochemistry. Dr. A. R. Ramesh Assistant Professor of Chemistry Govt. Engineering College, Kozhikode
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1 Electrochemistry Dr. A. R. Ramesh Assistant Professor of Chemistry Govt. Engineering College, Kozhikode 1
2 Electro Chemistry : Chemistry of flow of electrons Redox Reaction The electrons flow through the external circuit Electrochemical Cell Chemical to Electrical Anode = -ve Electrolytic Cell Electrical to Chemical (Reverse of Electrochemical Cell) Anode = +ve 2
3 A voltaic cell or Galvanic Cell Two-half cells separated by a porous boundary with solid electrodes connected by an external circuit Electrons always travel in the external circuit from anode to cathode Internally, cationsmove toward the cathode, anions move toward the anode, keeping the solution neutral (ionic movement through electrolyte) 3
4 Single Electrode Potential (Origin of electrode potential) When a metal rod is dipped in a solution of its own ions Either oxidation or Reduction takes place Oxidation Layer of negative charge (e - ) at the electrode surface -vely charged electrode surface attract a layer of +vely charged ions at the interface Develop an electrical double layer (EDL) at the metal-solution interface The potential difference between the metal and solution at the interface (EDL) is the single electrode potential 4
5 Oxidation Reduction (Reverse case) Standard Electrode Potential (E 0 ) is the electrode potential when the electrodeisin contactwitha solutionof unitconcentrationat298k. It measures the tendency of the metallic electrode to lose (oxidation potential) or gain (reduction potential) electrons, when it is in contact withitsown saltsolutionof1m concentrationat25 0 C 5
6 The electrode potential depends upon The nature of the metal and its ions Concentration of the ions in the solution and Temperature Helmholtz Double Layer A Helmholtz double layer (HDL) is an electrical double layer (EDL) of positive and negative charges one molecule thick. This occurs at the surface of a metal immersed in a solution. Potential difference ε = dielectric constant of the medium σ = charge density a = distance between the layers Layer of aligned ions, which is one particle thick and then immediately 6 next to that, free solution.
7 Gouy-Chapman Model There is not a simple layer of ions but, an ionic distribution that extends some distance from the surface - called diffused layer Stern Model Rigid Helmholtz layer Dispersed outside the Helmholtz plane 7
8 Terms Used for Voltaic Cells Half Cell & Cell 8
9 Electromotive Force (emf) Water only spontaneously flows one way in a waterfall. Likewise, electrons only spontaneously flow one way in a redox reaction from higher to lower potential energy. 9
10 Electromotive Force (emf) The potential difference between the cathode and anode in a cell is called the electromotive force (emf). It is also called the cell potential, and is designated E cell. 10
11 Daniell Cell Anode (oxidation) ive Zn metal salt bridge Cathode (reduction) +ive Cu metal ZnSO 4 (aq) CuSO 4 (aq) Zn(s) Zn 2+ (aq) + 2e Cu 2+ (aq) + 2e Cu(s) 11
12 Galvanic Cells (cont.) In turns out that we still will not get electron flow in the example cell. This is because charge buildup results in truncation of the electron flow. We need to complete the circuit by allowing positive ions to flow as well. We do this using a salt bridge which will allow charge neutrality in each cell to be maintained. 12
13 Salt Bridge Salt bridge is an inverted U tube filled with a concentrated solution of KCl or KNO 3 or NH 4 NO 3 in agar-agar or gelatin Functions Complete the inner electrical circuit, Daniell Cell without salt bridge maintain electrical neutrality Salt bridge makes cell construction and operation easier. Liquid-liquid interface Carefully merge two solutions. Make CuSO 4 more dense than ZnSO 4. Sheath Cu electrode in glass. Pack tube with a viscous, aqueous solution of KCl or KNO 3. The viscosity prevents mixing with the electrolytes. The ions permit exchange of charge. The chosen ions have similar mobility to minimize junction potentials. 13
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15 Electrochemical Series Is an arrangement of elements in the increasing order of their reduction potential 15
16 Electrochemical Series - Applications To know relative ease of oxidation and reduction E 0 = +ve (Reduction) = -ve (Oxidation) To predict whether metal react with acid to give hydrogen E 0 = -ve only react with H 2 To calculate standard EMF of cell E 0 Cell= E 0 R E 0 L = E 0 Cathode E 0 Anode 16
17 Nernst Equation E =E o n E =E o n lnq logq Note the difference between using natural logarithms and base10 logarithms. Be aware of the significance of n the number of moles of electrons transferred in the process according to the stoichiometry chosen. 17
18 EMF measurements (Poggendorf s compensation method) Cannot measured using voltmeter-cause current flow- change in concentration Principle: The EMFofthe cell isopposed by an external source ofemf. When there is no net flow of current in the circuit the imposed potentialwillbeequalto the EMFof the cell. 18
19 E = Battery (whose EMF is greater than Cell) Ex = Cell (unknown EMF) Es = Standard Cell (known EMF, Weston Cell) AB = Potentiometer wire length D = null deflection for Ex (AD) D = null deflection for Es (AD ) 19
20 Standard Hydrogen Electrode The convention is to select a particular electrode and assign its standard reduction potential the value of V. This electrode is the Standard Hydrogen Electrode. 2H + (aq) + 2e H 2 (g) H 2 Pt The standard aspect to this cell is that the activity of H 2 (g) and that of H + (aq) are both 1. This means that the pressure of H 2 is 1 atm and the concentration of H + is 1M, given that these are our standard reference states. H + 20
21 Standard Hydrogen Electrode E = 0 V (by definition; arbitrarily selected) 2H + + 2e - H 2 21
22 Calculating Cell Potential Because we tabulate reduction potentials, the cell potential is calculated (from those tabulated numbers) as E cell = E cathode - E anode The minus sign is present only because we are using reduction potential tables and, by definition, an anode is where oxidation occurs. 22
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26 Concentration Cells...a cell in which both compartments have the same components but at different concentrations 26
27 Concentration Cells Notice that the Nernst equation implies that a cell could be created that has the same substance at both electrodes. For such a cell, E cell would be 0, but Q would not. Therefore, as long as the concentrations are different, E will not be 0. 27
28 e e Ag e Porous disk e Ag 0.1 M Ag M NO 3 1 M Ag + 1 M NO 3 Anode Cathode 28
29 29
30 Batteries A battery is a galvanic cell or, more commonly, a group of galvanic cells connected in series. 30
31 How Does a Battery Work Assume a generalized battery Seal/cap Electrolyte Paste cathode (+) anode (-) 31
32 Battery Placing the battery into a flashlight, etc., and turning the power on completes the circuit and allows electron flow to occur Electrolyte paste: ion migration occurs here e - flow cathode (+): Reduction occurs here anode (-): oxidation occurs here 32
33 How Does a Battery Work Battery reaction when producing electricity (spontaneous): Cathode: O 1 + e - R 1 Anode: R 2 O 2 + e - Overall: O 1 + R 2 R 1 + O 2 Recharging a secondary cell Redox reaction must be reversed, i.e., current is reversed (nonspontaneous) Recharge: O 2 + R 1 R 2 + O 1 Performed using electrical energy from an external power source 33
34 34
35 Alkaline Dry Cell 35
36 Alkaline Dry Cell Plated steel (+) Cathode: Mixture of MnO 2 and C (graphite) Brass rod Anode: Mixture of Zn and KOH(aq) Insulators Paper or fabric Separator Plated steel (-) 36
37 Alkaline Dry Cell Half-reactions anode: Zn(s) + 2OH - (aq) --> ZnO(s) + H 2 O(l) + 2e - cathode: 2MnO 2 (s) + H 2 O(l) + 2e - --> Mn 2 O 3 (s) + 2OH - (aq) overall: Zn(s) + 2MnO 2 (s) --> Mn 2 O 3 (s) + ZnO(s) E cell = 1.54 V 37
38 Batteries are Galvanic Cells Car batteries are lead storage batteries. Pb +PbO 2 +H 2 SO 4 PbSO 4 (s) +H 2 O 38
39 Lead Storage Battery (anode) (cathode) 6 x 2V = 12 V 39
40 Half-reactions Lead Storage Battery anode: Pb(s) + HSO 4 2- (aq) --> PbSO 4 (s) + H + + 2e - cathode: PbO 2 (s) + 3H + (aq) + HSO 4 2- (aq) + 2e - --> PbSO 4 (s) + 2H 2 O(l) overall: Pb(s) + PbO 2 (s) + 2H + + 2HSO 4- (aq) --> 2PbSO 4 (s) + 2H 2 O(l) Cell reaction reversed during recharging. 40
41 Lead Storage Battery Half-reactions during recharging (nonspontaneous) cathode: PbSO 4 (s) + H + + 2e - --> Pb(s) + HSO 2-4 (aq) anode: PbSO 4 (s) + 2H 2 O(l) --> PbO 2 (s) + 3H + (aq) + HSO 2-4 (aq) + 2e - overall: 2PbSO 4 (s) + 2H 2 O(l) --> PbO 2 (s) + Pb(s) + 2H + + 2HSO 4- (aq) Cell converted into electrolytic cell via application of external electrical energy. 41
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43 Ni-Cad Battery Anode (-) Cd + 2 OH - ---> Cd(OH) 2 + 2e- Cathode (+) NiO(OH) + H 2 O + e- ---> Ni(OH) 2 + OH - 43
44 Fuel Cells Voltaic-like cell that operates with continuous supply of energetic reactants (fuel) to the electrodes utilize combustion reactions do not store chemical energy Not self-contained since reactants must be supplied to the electrodes Example: Hydrogen-Oxygen fuel cell 44
45 Hydrogen-Oxygen Fuel Cell 45
46 Hydrogen-Oxygen Fuel Cell Half-reactions anode: 2H 2 (g) + 4OH - (aq) --> 4H 2 O(l) + 4e - cathode: O 2 (g) + 2H 2 O(l) + 4e - --> 4OH - (aq) overall: 2H 2 (g) + O 2 (g) --> 2H 2 O(l) 46
47 Fuel Cells Galvanic cells Reactants are continuously supplied. 2H 2(g) + O 2(g) 2H 2 O (l) anode: 2H 2 + 4OH 4H 2 O + 4e cathode: 4e + O 2 + 2H 2 O 4OH Dr. A. R. Ramesh-GEC CLT - 47
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