OME General Chemistry Lecture 10: Electrochemistry Dr. Vladimir Lesnyak Office: Physical Chemistry, Erich Müller-Bau, r. 111 Email: vladimir.lesnyak@chemie.tu-dresden.de Phone: +49 351 463 34907
Outline Oxidation Reduction Reactions 1 Electrolytes 2 Balancing Oxidation Reduction Equations Voltaic Cells 3 Construction of Voltaic Cells 4 Notation for Voltaic Cells 5 Cell Potential 6 Standard Cell Potentials and Standard Electrode Potentials 7 Equilibrium Constants from Cell Potentials 8 Dependence of Cell Potential on Concentration 9 Some Commercial Voltaic Cells Electrolytic Cells 10 Electrolysis of Molten Salts 11 Aqueous Electrolysis 2
Electrolytes and Nonelectrolytes Electrolyte is a substance that dissolves in water to give an electrically conducting solution. In general, ionic solids that dissolve in water are electrolytes. Example: NaCl, table salt Nonelectrolyte is a substance that dissolves in water to give a nonconducting or very poorly conducting solution. Example: sucrose C 12 H 22 O 11, ordinary table sugar 3
Electrical Conductivity of a Solution Pure water Solution of sodium chloride 4
Strong and Weak Electrolytes Strong electrolyte exists in solution almost entirely as ions: most ionic solids that dissolve in water. Weak electrolyte dissolves in water to give a relatively small percentage of ions: most soluble molecular substances. H 2 O NaCl(s) Na + (aq) + Cl (aq) NH 3 (aq) + H 2 O(l) NH 4+ (aq) + OH (aq) 5
Solubility Rules for Ionic Compounds Rule Statement Applies to Exceptions 1 2 3 Group IA and ammonium compounds are soluble Acetates and nitrates are soluble Most chlorides, bromides, and iodides are soluble Li +, Na +, K +, NH + 4 C 2 H 3 O 2, NO 3 Cl, Br, I AgCl, Hg 2Cl 2, PbCl 2, AgBr, HgBr 2, Hg 2 Br 2, PbBr 2, AgI, HgI 2, Hg 2 I 2, PbI 2 4 Most sulfates are soluble SO 2 4 CaSO 4, SrSO 4, BaSO 4, Ag 2 SO 4, Hg 2 SO 4, PbSO 4 5 Most carbonates are insoluble CO 2 3 Group IA carbonates, (NH 4 ) 2 CO 3 6 Most phosphates are insoluble PO 3 4 Group IA phosphates, (NH 4 ) 3 PO 4 7 Most sulfides are insoluble S 2 Group IA sulfides, (NH 4 ) 2 S 8 Most hydroxides are insoluble OH Group IA hydroxides, Ca(OH) 2, Sr(OH) 2, Ba(OH) 2 6
Oxidation Reduction Reactions Reactions involving a transfer of electrons from one species to another: Fe(s) + CuSO 4 (aq) FeSO 4 (aq) + Cu(s) Fe(s) + Cu 2+ (aq) Fe 2+ (aq) + Cu(s) 7
Oxidation Numbers Oxidation number (or oxidation state) of an atom in a substance is the actual charge of the atom if it exists as a monatomic ion, or a hypothetical charge assigned to the atom in the substance by simple rules. In oxidation-reduction reaction one or more atoms change oxidation number transfer electrons. oxidized 0 0 +2 2 2Ca(s) + O 2 (g) 2CaO(s) reduced 8
Oxidation-Number (ON) Rules Rule Applies to Statement 1 Elements ON of an atom in an element is 0 2 Monatomic ions ON of an atom in a monatomic ion = the charge on the ion 3 Oxygen ON of oxygen = 2 in most of its compounds (Exception: O in H 2 O 2 and other peroxides = 1) 4 Hydrogen ON of H = +1 in most of its compounds ( 1 in binary compounds with a metal, such as CaH 2 ) 5 Halogens ON of F = 1 in all of its compounds. Each of the other halogens (Cl, Br, I) has ON = 1 in binary compounds, except when the other element is another halogen above it in the periodic table or the other element is oxygen. 6 Compounds and ions The sum of ONs of the atoms in a compound = 0. The sum of the ONs of the atoms in a polyatomic ion equals the charge on the ion. 9
Oxidation-Reduction Reactions 0 +2 +2 0 Fe(s) + Cu 2+ (aq) Fe 2+ (aq) + Cu(s) reducing agent oxidation oxidizing agent reduction Half-reactions: 0 +2 Fe(s) Fe 2+ (aq) + 2e (electrons lost by Fe) oxidation +2 0 Cu 2+ (aq) + 2e Cu(s) (electrons gained by Cu 2+ ) reduction Oxidation is the half-reaction in which there is a loss of electrons by a species (or an increase of oxidation number of an atom). Reduction is the half-reaction in which there is a gain of electrons by a species (or a decrease in the oxidation number of an atom). 10
Common Oxidation-Reduction Reactions 1. Combination reaction: two substances combine to form a third substance. 2Na(s) + Cl 2 (g) 2NaCl(s) not oxidation-reduction: CaO(s) + SO 2 (g) CaSO 3 (s) 2. Decomposition reaction: a single compound reacts to give two or more substances. 2HgO(s) 2Hg(l) + O 2 (g) not oxidation-reduction: CaCO 3 (s) CaO(s) + CO 2 (g) 3. Displacement reaction: an element reacts with a compound, displacing another element from it. Zn(s) + 2HCl(aq) ZnCl 2 (aq) + H 2 (g) 4. Combustion reaction: a substance reacts with oxygen, usually with the rapid release of heat to produce a flame. 4Fe(s) + 3O 2 (g) 2Fe 2 O 3 (s) 11
Activity Series of the Elements 12
Balancing Oxidation-Reduction Equations 0 +1 +2 0 Zn(s) + Ag + (aq) Zn 2+ (aq) + Ag(s) Zn Zn 2+ Ag + Ag Zn Zn 2+ + 2e Ag + + e Ag (oxidation) (reduction) (oxidation half-reaction) (reduction half-reaction) 1 (Zn Zn 2+ + 2e ) 2 (Ag + + e Ag) Zn(s) + 2Ag + (aq) + 2e Zn 2+ (aq) + 2Ag(s) + 2e Zn(s) + 2Ag + (aq) Zn 2+ (aq) + 2Ag(s) 13
Skeleton Oxidation-Reduction Equations To set up the skeleton equation and then balance it, we need to answer the following questions: 1. What species is being oxidized (or, what is the reducing agent)? What species is being reduced (or, what is the oxidizing agent)? 2. What species result from the oxidation and reduction? 3. Does the reaction occur in acidic or basic solution? +2 +7 +3 +2 Fe 2+ (aq) + MnO 4 (aq) Fe 3+ (aq) + Mn 2+ (aq) (acidic solution) 14
Balancing Oxidation-Reduction Equations in Acidic Solution Step 1: Assign oxidation numbers to each atom. Step 2: Split the skeleton equation into 2 half-reactions: oxidation and reduction. Step 3: Complete and balance each half-reaction. a. Balance all atoms except O and H. b. Balance O atoms by adding H2O s to one side of the equation. c. Balance H atoms by adding H + ions to one side of the equation. d. Balance electric charge by adding electrons (e ) to the more positive side. Step 4: Combine 2 half-reactions to obtain the final balanced oxidation-reduction equation. a. Multiply each half-reaction by a factor such that when the half-reactions are added, the electrons cancel. b. Simplify the balanced equation by canceling species that occur on both sides, and reduce the coefficients to smallest whole numbers. 15
Balancing Oxidation-Reduction Equations in Acidic Solution +2 +7 +3 +2 Fe 2+ (aq) + MnO 4 (aq) Fe 3+ (aq) + Mn 2+ (aq) (acidic solution) +2 +3 Fe 2+ (aq) Fe 3+ (aq) +2 +3 Fe 2+ (aq) Fe 3+ (aq) + e (oxidation half-reaction) +7 +2 MnO 4 (aq) Mn 2+ (aq) MnO 4 (aq) Mn 2+ (aq) + 4H 2 O(l) MnO 4 (aq) + 8H + (aq) Mn 2+ (aq) + 4H 2 O(l) MnO 4 (aq) + 8H + (aq) + 5e Mn 2+ (aq) + 4H 2 O(l) (reduction half-reaction) 5 (Fe 2+ Fe 3+ + e ) 1 (MnO 4 + 8H + + 5e Mn 2+ + 4H 2 O) 5Fe 2+ + MnO 4 + 8H + + 5e 5Fe 3+ + Mn 2+ + 4H 2 O + 5e 5Fe 2+ (aq) + MnO 4 (aq) + 8H + (aq) 5Fe 3+ (aq) + Mn 2+ (aq) + 4H 2 O(l) 16
Additional Steps for Balancing Oxidation-Reduction Equations in Basic Solution Step 5: Note the number of H + ions in the equation. Add this number of OH ions to both sides of the equation. Step 6: Simplify the equation by noting that H + reacts with OH to give H 2 O. Cancel any H 2 O s that occur on both sides of the equation and reduce the equation to simplest terms. MnO 4 (aq) + SO 3 2 (aq) MnO 2 (s) + SO 4 2 (aq) (basic solution) 2MnO 4 + 3SO 3 2 + 2H + 2MnO 2 + 3SO 4 2 + H 2 O 2MnO 4 + 3SO 3 2 + 2H + + 2OH 2MnO 2 + 3SO 4 2 + H 2 O + 2OH 2H 2 O 2MnO 4 (aq) + 3SO 3 2 (aq) + H 2 O(l) 2MnO 2 (s) + 3SO 4 2 (aq) + 2OH (aq) 17
Construction of Voltaic Cells Voltaic cell consists of 2 half-cells that are electrically connected. Electrochemical cell consists of electrodes that dip into an electrolyte and in which a chemical reaction either uses or generates an electric current. Voltaic (or galvanic cell) is an electrochemical cell in which a spontaneous reaction generates an electric current. Electrolytic cell is an electrochemical cell in which an electric current drives an otherwise nonspontaneous reaction. 18
Construction of Voltaic Cells Zn(s) + Cu 2+ (aq) Zn 2+ (aq) + Cu(s) Zn electrode and Cu electrode, without an external circuit no cell reaction 2 electrodes are connected by an external circuit chemical reaction occurs Electrode at which oxidation occurs anode. Electrode at which reduction occurs cathode. 19
Notation for Voltaic Cells Zn(s) + Cu 2+ (aq) Zn 2+ (aq) + Cu(s) always on the left salt bridge Zn(s) Zn 2+ (aq) Cu 2+ (aq) Cu(s) anode cathode always on the right phase boundary Zn(s) Zn 2+ (aq) anode terminal anode electrolyte 20
Notation for Voltaic Cells When the half-reaction involves a gas Hydrogen electrode: hydrogen gas bubbles over a platinum surface, where the half-reaction occurs: 2H + (aq) + 2e H 2 (g) cathode: H + (aq) H 2 (g) Pt anode: Pt H 2 (g) H + (aq) Cathode Cl 2 (g) Cl (aq) Pt Fe 3+ (aq), Fe 2+ (aq) Pt Cd 2+ (aq) Cd(s) Cathode reaction Cl 2 (g) + 2e 2Cl (aq) Fe 3+ (aq) + e Fe 2+ (aq) Cd 2+ (aq) + 2e Cd(s) 21
Cell Potential Potential difference is the difference in electric potential between two points. Volt V is the SI unit of potential difference. The electrical work expended in moving a charge through a conductor: Electrical work = charge potential difference Joules = coulombs volts The Faraday constant, F, is the magnitude of charge on one mole of electrons; F = 9.6485 10 4 C per mole of electrons (96,485 C/mol e ). faraday is a unit of charge = 9.6485 10 4 C. In moving this quantity of charge (1 faraday) from one electrode to another, the work done by a voltaic cell: w = F potential difference 22
Cell Potential The maximum potential difference between the electrodes of a voltaic cell is the cell potential or electromotive force (emf) of the cell E cell. It can be measured by an electronic digital voltmeter. The maximum electrical work of a voltaic cell for molar amounts of reactants (according to the cell equation as written) is: w max = nfe cell n the number of moles of electrons transferred in the overall cell equation, E cell the cell potential, F the Faraday constant, 9.6485 10 4 C/mol e. 23
Standard Cell Potentials and Standard Electrode Potentials Cell potential is a measure of the driving force of the cell reaction. reduced species oxidized species + ne (oxidation/anode) oxidized species + ne reduced species (reduction/cathode) E cell = oxidation potential + reduction potential Oxidation potential for a half-reaction = reduction potential for the reverse half-reaction Reduction potentials are tabulated as electrode potentials, E. 24
Standard Cell Potentials and Standard Electrode Potentials Zn(s) Zn 2+ (aq) Cu 2+ (aq) Cu(s) Zn(s) Zn 2+ (aq) + 2e E Zn Zn 2+ (aq) + 2e Zn(s) E Zn Cu 2+ (aq) + 2e Cu(s) E Cu E cell = E Cu + ( E Zn ) = E Cu E Zn E cell = E cathode E anode The cell potential depends on the concentrations of substances and the temperature of the cell. The standard cell potential, E o cell, is the emf of a voltaic cell operating under standard-state conditions (solute concentrations 1 M, gas pressures 1 atm, temperature usually 25 o C). 25
Tabulating Standard Electrode Potentials The reference chosen for comparing electrode potentials is the standard hydrogen electrode, it is assigned a potential of 0.0 V. Zn(s) Zn 2+ (aq) H + (aq) H 2 (g) Pt Zn(s) Zn 2+ (aq) + 2e ; E o Zn 2H + (aq) + 2e H 2 (g); E o H 2 E cell = E o H 2 + ( E o Zn) E o Zn = 0.76 V 26
Standard Electrode Potentials in Aqueous Solution at 25 o C 27
Strengths of Oxidizing and Reducing Agents oxidized species + ne reduced species The strongest oxidizing agents in a table of standard electrode potentials are the oxidized species corresponding to half-reactions with the largest (most positive) E o values. reduced species oxidized species + ne The strongest reducing agents in a table of standard electrode potentials are the reduced species corresponding to half-reactions with the smallest (most negative) E o values. 28
Calculation of Cell Potentials Using Standard Reduction Potentials Cd 2+ (aq) + 2e Cd(s); E o Cd = 0.4 V Ag + (aq) + e Ag(s); E o Ag = 0.8 V E o cell = E o cathode E o anode E o cell = E o Ag E o Cd = 0.8 ( 0.4) = 1.2 V 29
Equilibrium Constants from Cell Potentials GG = ww mmmmmm GG oo oo = nnnnee cccccccc Free energies of reaction are generally of the order of 10s to 100s of kilojoules. Example: Zn(s) + 2Ag + (aq) Zn 2+ (aq) + 2Ag(s) Zn(s) Zn 2+ (aq) + 2e E o = 0.76 V 2Ag + (aq) + 2e 2Ag(s) E o = 0.80 V Zn(s) + 2Ag + (aq) Zn 2+ (aq) + 2Ag(s) E o cell = 1.56 V GG oo oo = nnnnee cccccccc = 2 mmmmmm ee 96,485 CC mmmmll ee 1.56 VV = 3.01 10 5 JJ 30
Equilibrium Constants from Cell Potentials GG oo oo = nnnnee cccccccc GG oo = RRRRRRRRRR oo nnnnee cccccccc = RRRRRRRRRR oo EE cccccccc = RRRR 2.303RRRR llllll = llooookk nnnn nnnn oo EE cccccccc = 0.0592 nn llooookk (iiii vvvvvvvvvv aaaa 25 oo CC) 31
Relationships among K, G o, and E cell 32
Dependence of Cell Potential on Concentration: Nernst Equation The cell potential of a cell depends on the concentrations of ions and on gas pressures cell potentials provide a way to measure ion concentrations. GG = GG oo + RRRRRRRRRR GG oo oo = nnnnee cccccccc GG = nnnnee cccccccc oo nnnnee cccccccc = nnnnee cccccccc + RRRRRRRRRR oo EE cccccccc = EE cccccccc RRRR nnnn llllqq = EE oo cccccccc 2.303RRRR nnnn lloooooo first derived by the German chemist Walther Nernst (1864 1941) oo EE cccccccc = EE cccccccc 0.0592 nn lloooooo (iiii vvvvvvvvvv aaaa 25 oo CC) Cell potential E cell decreases as the cell reaction proceeds going to 0. 33
Determination of ph Glass electrode pppp = log[hh + ] ph of a solution can be obtained very accurately from cell potential measurements, using the Nernst equation to relate cell potential to ph. The electrode solution is separated from the test solution by a thin glass membrane, which develops a potential across it depending on the H + concentrations on its inner and outer surfaces. Glass electrode is an example of an ion-selective electrode. Many electrodes have been developed that are sensitive to a particular ion, such as K +, NH 4+, Ca 2+, or Mg 2+. They can be used to monitor solutions of that ion. 34
Commercial Voltaic Cells Zinc carbon (Leclanché dry cell) Zn(s) Zn 2+ (aq) + 2e (anode) 2NH 4+ (aq) + 2MnO(s) + 2e Mn 2 O 3 (s) + H 2 O(l) + 2NH 3 (aq) (cathode) The first battery was invented by Alessandro Volta about 1800. He assembled a pile consisting of pairs of zinc and silver disks separated by paper disks soaked in salt water. A battery cell that became popular during the 19 th century was constructed in 1836 by the English chemist John Frederick Daniell. The voltage of this cell is initially about 1.5 V 35
Commercial Voltaic Cells Alkaline dry cell Lithium iodine battery Zn(s) Zn 2+ (aq) + 2e 2MnO 2 (s) + H 2 O(l) + 2e Mn 2 O 3 (s) + 2OH (aq) (anode) (cathode) used to power heart pacemakers 36
Commercial Voltaic Cells Lead storage cell Pb(s) + HSO 4 (aq) PbSO 4 (s) + H + (aq) + 2e (anode) PbO 2 (s) + H + (aq) + HSO 4 (aq) + 2e PbSO 4 (s) + 2H 2 O(l) (cathode) After the lead battery is discharged, it is recharged from an external electric current. Each cell delivers about 2 V, a battery consisting of six cells in series gives about 12 V 37
Commercial Voltaic Cells Nickel-cadmium cell (nicad cell) Cd(s) + 2OH (aq) Cd(OH) 2 (s) + 2e (anode) NiOOH(s) + H 2 O(l) + 2e Ni(OH) 2 (s) + OH (aq) (cathode) These half-reactions are reversed when the cell is recharged. Nicad batteries can be recharged and discharged many times. 38
Commercial Voltaic Cells: Hydrogen oxygen fuel cell Battery with a continuous supply of energetic reactants (fuel). 2H 2 (g) + O 2 (g) 2H 2 O(l) H 2 (g) 2H + (aq) + 2e (anode) O 2 (g) + 4H + (aq) + 4e 2H 2 O(l) (cathode) Potential = 0.7 V 39
Rusting is also an Electrochemical Process O 2 (g) + 2H 2 O(l) + 4e 4OH (aq) Fe(s) Fe 2+ (aq) + 2e Fe 2+ (aq) + 2OH (aq) Fe(OH) 2 (s) 4 Fe(OH) 2 (s) + O 2 (g) 2Fe 2 O 3 H 2 O(s) + 2H 2 O(l) 40
Cathodic Protection against Corrosion f 41
Electrolysis of Molten Salts: NaCl In electrolytic cell electric current drives an otherwise nonspontaneous reaction. The process of producing a chemical change in an electrolytic cell is electrolysis. Many important substances, including aluminum and chlorine, are produced commercially by electrolysis. Na + (l) + Cl (l) Na(l) + ½Cl 2 (g) Downs cell Li, Mg, and Ca metals are obtained by the electrolysis of their chlorides. 42
Aqueous Electrolysis: NaCl The electrolysis of aqueous NaCl is the basis of the chlor-alkali industry the major commercial source of Cl 2 and NaOH. Chlor-alkali membrane cell reduction: 2H 2 O(l) + 2e H 2 (g) + 2OH (aq) oxidation: 2H 2 O(l) O 2 (g) + 4H + (aq) + 4e possible cathode half-reactions: Na + (aq) + e Na(s); E o = 2.71 V 2H 2 O(l) + 2e H 2 (g) + 2OH (aq); E o = 0.83 V 2H 2 O(l) + 2Cl (aq) H 2 (g) + Cl 2 (g) + 2OH (aq) 43
Electroplating of Metals Purification of copper by electrolysis 44
Stoichiometry of Electrolysis Michael Faraday (1831): the amounts of substances released at the electrodes during electrolysis are related to the total charge that has flowed in the electric circuit. 1 faraday (9.6485 10 4 C) = the charge on one mole of electrons Electric charge = electric current time lapse Ampere (A) is the SI unit of current. Coulomb (C) is the SI unit of electric charge, equivalent to an ampere-second. A current of 0.5 amperes flowing for 84 seconds gives a charge of 0.5 A 84 s = 42 A s, or 42 C. From the amount of substance produced at an electrode and the time of electrolysis, we can determine the current. From the current and the time of electrolysis, we can calculate the amount of substance produced at an electrode. 45
ww mmmmmm = GG oo oo = nnnnee cccccccc oo EE cccccccc oo = EE ccaaaaaaaaaaa oo EE aaaaaaaaaa Summary oo EE cccccccc = EE cccccccc 0.0592 nn lloooooo (iiii vvvvvvvvvv aaaa 25 oo CC) Oxidation-reduction reactions involve a transfer of electrons from one species to another. Electrochemical cells are of 2 types: voltaic and electrolytic. Voltaic cells use a spontaneous chemical reaction to generate an electric current. Electrolytic cells use an external voltage source to push a reaction in a nonspontaneous direction. Electrons flow in the external circuit from the anode to the cathode. Cell potential is the maximum voltage of a voltaic cell. The standard free-energy change, standard cell potential, and equilibrium constant are all related. Electrochemical measurements can provide equilibrium or thermodynamic information. Electrode potential depends on concentrations of the electrode substances, according to the Nernst equation. Voltaic cells are used commercially as batteries. The basic principle of the voltaic cell is employed in the cathodic protection of buried pipelines and tanks. The electrolysis of an aqueous solution often involves the oxidation or reduction of water. 46