Yeungnam University Basic Concepts in Electrochemistry Mosab Kaseem Plasticity Control and Mechanical Modeling Laboratory School of Materials Science and Engineering, Yeungnam University, Korea Contents redox reactions Balancing redox reactions electrochemical cells • electrode processes • construction • notation cell potential and Go standard reduction potentials (Eo) non-equilibrium conditions (Q) Yeungnam University School of Materials Science and Engineering Introduction 1) Redox Reaction involves transfer of electrons from one species to another. Redox reactions: electron transfer processes Oxidation: loss e- ………causes reduction “reducing agent” Reduction: gain e- ……… causes oxidation “ oxidizing agent” Zn metal With time, Cu plates out onto Zn metal strip, and Zn strip “disappears.” Cu+2 ions Yeungnam University School of Materials Science and Engineering Redox Reaction • Zn is oxidized reducing agent Zn(s) Zn2+(aq) + 2e• Cu2+ is reduced oxidizing agent Cu2+(aq) + 2e- Cu(s) wire elect rons Electrons travel thrugh external wire. Salt bridge allows anions and cations to move between electrode compartments. This maintains electrical neutrality. Zn Zn2+ ions Yeungnam University salt bridge Cu Cu2+ ions School of Materials Science and Engineering The first two reactions are known as “1/2 cell reactions” Include electrons in their equation 3.) The net reaction is known as the total cell reaction No free electrons in its equation 2.) ½ cell reactions: Net Reaction: 4.) In order for a redox reaction to occur, both reduction of one compound and oxidation of another must take place simultaneously Total number of electrons is constant Yeungnam University School of Materials Science and Engineering Balancing redox reactions in acidic solution Cr2O72-(aq) + I-(aq) → Cr3+(aq) + I2(s) Step 1: Divide the reaction into half-reactions. Cr2O72- → Cr3+ I→ I2 Step 2: Balance the atoms and charges in each half-reaction. For the Cr2O72-/Cr3+ half-reaction: Balance atoms other than O and H: Cr2O72- → 2Cr3+ Balance O atoms by adding H2O molecules: Cr2O72- → 2Cr3+ + 7H2O Yeungnam University School of Materials Science and Engineering Balancing redox reactions in acidic solution Balance H atoms by adding H+ ions: 14H+ + Cr2O72- → 2Cr3+ + 7H2O Balance charges by adding electrons: 6e- + 14H+ + Cr2O72- → 2Cr3+ + 7H2O For the I-/I2 half-reaction: Balance atoms other than O and H: 2I- → I2 There are no O or H atoms, so we balance charges by adding electrons: 2I- → I2 + 2eStep 3: try to make the number of e- lost in the oxidation equals the number of egained in the reduction. 6e- + 14H+ + Cr2O72- → 2Cr3+ + 7H2O 6I- → 3I2 + 6e6I-(aq) + 14H+(aq) + Cr2O72-(aq) → 3I2(s) + 7H2O(l) + 2Cr3+(aq) Yeungnam University School of Materials Science and Engineering Balancing redox reactions in acidic solution Example: MnO4-(aq) + C2O4-2 (aq) → Mn+2 (aq) + CO2 (g) ClO3-(aq) + SO2 (g) → SO4-2 (aq) + Cl- (aq) Yeungnam University School of Materials Science and Engineering Balancing redox reactions in basic solution MnO4-(aq) + Br- (aq) → MnO2 (s) + BrO3- (aq) Se + Cr(OH)3 → Cr + SeO3-2 Yeungnam University School of Materials Science and Engineering Electrical parameters 1.) Electric Charge (q) Measured in coulombs (C) Charge of a single electron is 1.602x10-19C Faraday constant (F) : 9.649x104C is the charge of a mole of electrons Relation between charge and moles Coulombs 2.) moles Electric current Quantity of charge flowing each second through a circuit -Ampere: unit of current (C/sec) Yeungnam University School of Materials Science and Engineering 3.) Electric Potential (E) Measured in volts (V) Work (energy) needed when moving an electric charge from one point to another -Measure of force pushing on electrons Relation between free energy, work and voltage: Joules Volts Higher potential difference Coulombs Higher potential difference requires more work to lift water (electrons) to higher trough Yeungnam University School of Materials Science and Engineering 4.) Electric Potential (E) Combining definition of electrical charge and potential Describes the voltage that can be generated by a chemical reaction 5.) Ohm’s Law Current (I) is directly proportional to the potential difference (voltage) across a circuit and inversely proportional to the resistance (R) -Ohms (W) - units of resistance Yeungnam University School of Materials Science and Engineering 6.) Power (P) Work done per unit time -Units: joules per second J/sec or watts (W) Yeungnam University School of Materials Science and Engineering Electrochemical Cells There are two types of electrochemical cells Galvanic cell: ones that spontaneous produce electrical energy. Electrolytic cell: ones that consume electrical energy. Yeungnam University School of Materials Science and Engineering Electrochemical cells 1.) Galvanic or Voltaic cell Spontaneous chemical reaction to generate electricity -One reagent oxidized the other reduced -two reagents cannot be in contact Electrons flow from reducing agent to oxidizing agent -Flow through external circuit to go from one reagent to the other Reduction: Oxidation: Net Reaction: AgCl(s) is reduced to Ag(s) Ag deposited on electrode and Clgoes into solution Cd(s) is oxidized to Cd2+ Cd2+ goes into solution Yeungnam University School of Materials Science and Engineering Galvanic or Voltaic cell Daniell cell Salt Bridge Connects & separates two half-cell reactions Prevents charge build-up and allows counter-ion migration -0.76 V Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) A potential difference between two electrodes represents a tendency for the reaction to o ccur. • Chemical energy electrical energy • Spontaneous: (so Ecell is positive) Example: Primary (non-rechargeable) • Le Clanche (dry cell) Secondary (rechargeable) • Lead storage battery Hydrogen-Oxygen Fuel Cell Yeungnam University School of Materials Science and Engineering Exercise1 Consider a simple galvanic cell consisting of two beakers connected by a salt bridge. One beaker contains a solution of MnO4− in dilute sulfuric acid and has a Pt electrode. The other beaker contains a solution of Sn+2 in dilute sulfuric acid, also with a Pt electrode. When the two electrodes are connected by a wire, current flows and a spontaneous reaction occurs that is described by the following balanced chemical equation: 2MnO4−(aq) + 5Sn+2 (aq) + 16H+(aq) → 2Mn+2 (aq) + 5Sn+4 (aq) + 8H2O(l) For this galvanic cell, a. write the half-reaction that occurs at each electrode. b. indicate which electrode is the cathode and which is the anode. c. indicate which electrode is positive and which is negative. Yeungnam University School of Materials Science and Engineering Applications of Galvanic Cells- Lead-Acid Battery Reactions PbO2 + 4H+ + 2e- + SO4-2 PbSO4 + 2H2O Pb(s) + SO4-2 PbSO4(s) + 2e- E° = 1.685V E° = 0.356 V Total reaction PbO2(s) + Pb(s) + 2H2SO4 2 PbSO4(s) + 2 H2O E° = 2 V Lead-acid batteries Yeungnam University School of Materials Science and Engineering Electrolytic cell Electrolytic cell: is a cell in which current flows, power is consumed, and the cell reaction being driven is the reverse of the spontaneous cell reaction. Cathode (-) Anode (+) 2 H2O O2(g) + 4 H+ + 4 H2O + 4e- 4e- 2H2 + 4 OH- • Electrical energy chemical energy • Non-spontaneous: (Ecell is negative) Example: • Electrolysis of water Yeungnam University School of Materials Science and Engineering Comparison between electrolytic and galvanic cells Energy Energy • Energy is released from spontaneous redox reaction • Anode is negative • Energy is absorbed to drive nonspontaneous redox reaction. • Anode is positive In both voltaic and electrolytic cells, electrons flow from the anode to the cathode in the external circuit Yeungnam University School of Materials Science and Engineering Schematic Representation of Electrochemical Cells The anode is always on the left : boundaries : salt bridge For copper-zinc cell: Other conditions like concentration are listed just after each species The cell reaction H2 (g) + Cu2+ (aq) 2 H+ (aq) + Cu (s) Pt H2 (g) H+ (aq) Cu2+ (aq) Cu (s) Pt Yeungnam University School of Materials Science and Engineering Cell potential, Eo For Zn/Cu, voltage is 1.10 V at 25°C and when [Zn2+] and [Cu2+] = 1.0 M. This is the Standard cell potential, Eo Eo is a quantitative measure of the tendency of reactants to proceed to products when all are in their standard states at 25 °C. Eo and Go Eo is related to Go, the free energy change for the reaction. Go = - n F Eo F = Faraday constant = 9.6485 x 104 J/V•mol n = the number of moles of electrons transferred. Yeungnam University School of Materials Science and Engineering Standard cell potential, Eo • Can’t measure half- reaction Eo directly. Therefore, measure it relative to a standard half cell: the Standard Hydrogen Electrode (SHE) 2 H+(aq, 1 M) + 2e- H2(g, 1 atm) Eo = 0.0 V Yeungnam University School of Materials Science and Engineering Standard Potentials When combining two ½ cell reaction together to get a complete net reaction, the total cell potential (Ecell) is given by: E0cell = E0cathode - E0anode Predict voltage observed when two half-cells are connected • Reactions always written as reduction Potentials measured at standard conditions • All concentrations (or activities) = 1M • 25oC, 1 atm pressure Best Oxidizing agent? Best Reducing agent ? As Eo increases, the more favorable the reaction and the more easily the compound is reduced (better oxidizing agent). Yeungnam University School of Materials Science and Engineering E0cell > 0 G0 < 0 Spontaneous reaction proceeds to right (products) E0cell < 0 G0 > 0 Non-spontaneous reaction proceeds to left (reactants) E0cell = 0 G0 = 0 Equilibrium reaction Exercise? Br2(aq)+2V3+ +2H2O(l) 2VO2+(aq)+ 4H+(aq)+ 2Br-(aq) Given: E0cell = +1.39 V E0Br2 = +1.07 V What is E0V3+ and is the reaction spontaneous? Yeungnam University School of Materials Science and Engineering • Eo only applies to [ ] = 1 M for all aqueous species • at other concentrations, the cell potential differs • Ecell can be predicted by Nernst equation E = Eo - RT nF ln (Q) n = # e- transferred F = Faraday’s constant = 9.6485 x 104 J/V•mol Zn / Zn2+ (0.5 M) // Cu2+ (2.0 M) / Cu Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s) [Zn2+] Q= [Cu2+] E = 1.10 - (0.0257) ln ( [Zn2+]/[Cu2+] ) 2 E = 1.10 - (-0.018) = 1.118 V Yeungnam University School of Materials Science and Engineering Electrodes 1) Working electrode In corrosion testing, the working electrode is a sample of the corroding metal. The working electrode can be bare or coated metal. 2) Auxiliary (counter) electrode The auxiliary electrode is a conductor that completes the cell circuit. Such as an inert conductor like platinum or graphite The current that flows into the solution via the working electrode leaves the solution via the auxiliary electrode. Yeungnam University School of Materials Science and Engineering 3) Reference electrode A reference electrode is used in measuring the working electrode potential. A reference electrode should have a constant electrochemical potential as long as no current flows through it. Standard hydrogen electrode (SHE) The standard reference electrode which is often referred to is the standard hydrogen electrode (SHE). Al other reference electrodes can be expressed in terms of some constant deviation from SHE. The SHE reaction is based upon the reaction: 2H 2e H 2 g Eo =0.000VSHE Where Eo is the reference state at standard temperature and pressure. It should be noted that this potential is independent of temperature. Yeungnam University School of Materials Science and Engineering Reference electrode Saturated calomel electrode(SCE) The saturated calomel electrode consists of pure mercury covering a platinum wire which passes through a sealed glass tube. The mercury is covered with mercurous chloride and imme rse in saturated potassium chloride. The temperature coefficient is -6.6x 10 -4 V/oC Hg 2Cl2 2e 2 Hg 2Cl Eo =0.268VSHE Silver-silver chloride reference electrode: Ag-AgCl electrode consits of a silver, the surface which has been chloridized (transformed to silve r chloride), typically in dilute hydrochloric acid,. The temperature coefficient is –4.3x 10 -4 V/oC . Eo = 0.222 VSHE Yeungnam University School of Materials Science and Engineering Potentials of common reference electrodes Common Name Electrode V vs NHE Saturated Calomel Electrode (SCE) Hg/Hg2Cl2/sat. KCl +0.241 Calomel Hg/Hg2Cl2/1M KCl +0.280 Mercurous sulphate Hg/Hg2SO4/sat. K2SO4 +0.640 Mercurous oxide Hg/HgO/1M NaOH +0.098 Silver chloride Ag/AgCl/sat. KCl +0.197 Copper sulphate Cu/sat. CuSO 4 +0.316 Zinc in seawater Zn/seawater ~ -0.8 Yeungnam University School of Materials Science and Engineering References (1) Modern electrochemistry- 2nd edition-John O’M. Bocker and Amulya K.N. Reddy (2) Electrochemical methods ; fundamental and applications- Allen J. Bard and Larry R. Faulkner Yeungnam University School of Materials Science and Engineering