ANALYTICAL CHEMISTRY CHEM 3811 CHAPTER 14 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university

CHAPTER 14 ELECTRODE POTENTIALS

REDOX CHEMISTRY - Electron transfer occurs in redox reactions Oxidation - Loss of electrons Reduction - Gain of electrons Oxidizing agent (oxidant) is the species reduced Reducing agent (reductant) is the species oxidized

REDOX CHEMISTRY Oxidizing Agent - The species that gains electrons - The species that is reduced - Causes oxidation Aox + ne- ↔ Ared Cu2+(aq) + 2e- ↔ Cu(s)

REDOX CHEMISTRY Reducing Agent - The species that loses electrons - The species that is oxidized - Causes reduction Bred ↔ Box + ne- Fe(s) ↔ Fe2+(aq) + 2e-

REDOX CHEMISTRY The Overall Reaction - Both an oxidation and a reduction must occur in a redox reaction - The oxidizing agent accepts electrons from the reducing agent Aox + Bred ↔ Ared + Box Cu2+(aq) + Fe(s) ↔ Cu(s) + Fe2+(aq) - Reducing agent - Oxidized species - Electron loss - Oxidizing agent - Reduced species - Electron gain

REDOX CHEMISTRY Half Reactions - Oxidation half reaction Bred ↔ Box + ne- Fe(s) ↔ Fe2+(aq) + 2e- - Reduction half reaction Aox + ne- ↔ Ared Cu2+(aq) + 2e- ↔ Cu(s)

ELECTRODE - Conducts electrons into or out of a redox reaction system Examples platinum wire carbon (glassy or graphite) indium tin oxide (ITO) Electroactive Species - Donate or accept electrons at an electrode

REDOX CHEMISTRY Charge (q) of an electron = - 1.602 x 10-19 C Charge (q) of a proton = + 1.602 x 10-19 C C = coulombs Charge of one mole of electrons = (1.602 x 10-19 C)(6.022 x 1023/mol) = 9.649 x 104 C/mol = Faraday constant (F) q = n x F

CURRENT - The quantity of charge flowing past a point in an electric circuit per second Units Ampere (A) = coulomb per second (C/s)

VOLTAGE Potential Difference (E) - Work done by or on electrons when they move from one point to another Units: volts (V or J/C) Work (J) = E (V) x q (C)

CHEMICAL CHANGE Spontaneous Process - Takes place with no apparent cause Nonspontaneous Process - Requires something to be applied in order for it to occur (usually in the form of energy)

ELECTROLYSIS - Voltage is applied to drive a redox reaction that would not otherwise occur Examples - Production of aluminum metal from Al3+ - Production of Cl2 from Cl-

ELECTROLYSIS CELL - Nonspontaneous reaction - Requires electrical energy to occur

GALVANIC CELL - Spontaneous reaction - Produces electrical energy - Can be reversed electrolytically for reversible cells Example Rechargeable batteries Conditions for Non-reversibility - If one or more of the species decomposes - If a gas is produced and escapes

GALVANIC CELL - A spontaneous redox reaction generates electricity - One reagent is oxidized and the other is reduced - The two reagents must be separated (cannot be in contact) - Electrons flow through a wire (external circuit)

VOLTAIC (GALVANIC) CELL Oxidation Half reaction - Loss of electrons - Occurs at anode (negative electrode) - The left half-cell by convention Reduction Half Reaction - Gain of electrons - Occurs at cathode (positive electrode) - The right half-cell by convention

GALVANIC CELL Salt Bridge - Connects the two half-cells (anode and cathode) - Filled with gel containing saturated aqueous salt solution (KCl) - Ions migrate through to maintain electroneutrality - Prevents charge buildup that may cease the reaction process Preparation - Heat 3 g of agar and 30 g of KCl in 100 mL H2O - Heat until a clear solution is obtained - Pour into a U-tube and allow to gel - Store in a saturated aqueous KCl

VOLTAIC (GALVANIC) CELL For the overall reaction Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) e- Voltmeter e- - + Cu electrode Zn electrode Cl- K+ Zn2+ Salt bridge (KCl) Cu2+ Anode Oxidation Zn(s) → Zn2+(aq) + 2e- Cathode Reduction Cu2+(aq) + 2e- → Cu(s)

GALVANIC CELL Voltage or Potential Difference (E) - Is the voltage measured - Measured by a voltmeter (potentiometer) connected to electrodes Greater Voltage - More favorable net cell reaction - More work done by flowing electrons

GALVANIC CELL Line Notation Phase boundary: represented by one vertical line Salt bridge: represented by two vertical lines Fe(s) FeCl2(aq) CuSO4(aq) Cu(s)

STANDARD POTENTIALS Electrode Potentials - A measure of how willing a species is to gain or lose electrons Positive Voltage (spontaneous process) - Electrons flow into the negative terminal of voltmeter (flow from negative electrode to positive electrode) Negative Voltage (nonspontaneous process) - Electrons flow into the positive terminal of voltmeter (flow from positive electrode to negative electrode) Conventionally - Negative terminal is on the left of galvanic cells

STANDARD POTENTIALS Standard Reduction Potential (Eo) - Used to predict the voltage when different cells are connected - Potential of a cell as cathode compared to standard hydrogen electrode - Species are solids or liquids - Activities = 1 - We will use concentrations for simplicity Concentrations = 1 M Pressures = 1 bar

STANDARD POTENTIALS Standard Hydrogen Electrode (SHE) - Used to measure Eo for half-reactions - Connected to negative terminal - Pt electrode - Acidic solution in which [H+] = 1 M - H2 gas (1 bar) is bubbled past the electrode H+(aq, 1 M) + e- ↔ 1/2H2 (g, 1 bar) Conventionally, Eo = 0 for SHE

STANDARD POTENTIALS The Eo for Ag+ + e- ↔ Ag(s) is 0.799 V Implies that if a sample of silver metal is placed in a 1 M Ag+ solution, a value of 0.799 V will be measured with S. H. E. as reference Pt(s) H2(g, 1 bar) H+(aq, 1 M ) Ag+ (aq, 1 M) Ag(s) SHE Ag+ (aq, 1 M) Ag(s)

STANDARD POTENTIALS Silver does not react spontaneously with hydrogen 2H+(aq) + 2e- → H2(g) Eo = 0.000 V Ag+(aq) + e- → Ag(s) Eo = +0.799 V Reverse the second equation (sign changes) Ag(s) → Ag+(aq) + e- Eo = -0.799 V Multiply the second equation by 2 (Eo is intensive so remains) 2Ag(s) → 2Ag+(aq) + 2e- Eo = -0.799 V Combine (electrons cancel) 2Ag(s) + 2H+(aq) → 2Ag+(aq) + H2(g) Eo = -0.799 V

STANDARD POTENTIALS Consider Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) Cu2+(aq) + 2e- → Cu(s) Eo = +0.339 V Zn2+(aq) + 2e- → Zn(s) Eo = -0.762 V Reverse the second equation (sign changes) Zn(s) → Zn2+(aq) + 2e- Eo = +0.762 V Combine (electrons cancel) Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) Eo = +1.101 V Eo is positive so reaction is spontaneous Reverse reaction is nonspontaneous

STANDARD POTENTIALS - Half-reaction is more favorable for more positive Eo Formal Potential - The potential for a cell containing a specified concentration of reagent other than 1 M

STANDARD POTENTIALS Eo (V) Half Reaction 2.890 F2 + 2e- ↔ 2F- 1.507 1.280 1.229 0.799 0.339 0.000 -0.402 -0.440 -0.763 -1.659 -2.936 -3.040 Half Reaction F2 + 2e- ↔ 2F- MnO4- + 5e- ↔ Mn2+ Ce4+ + e- ↔ Ce3+ (in HCl) O2 + 4H+ + 4e- ↔ 2H2O Ag+ + e- ↔ Ag(s) Cu2+ + 2e- ↔ Cu(s) 2H+ + 2e- ↔ H2(g) Cd2+ + 2e- ↔ Cd(s) Fe2+ + 2e- ↔ Fe(s) Zn2+ + 2e- ↔ Zn(s) Al3+ + 3e- ↔ Al(s) K+ + e- ↔ K(s) Li+ + e- ↔ Li(s) Oxidizing agents Reducing agents Increasing oxidizing power Increasing reducing power

The half-cell potential (at 25 oC), E, is given by NERNST EQUATION For the half reaction aA + ne- ↔ bB The half-cell potential (at 25 oC), E, is given by

NERNST EQUATION Eo = standard electrode potential R = gas constant = 8.314 J/K-mol T = absolute temperature F = Faraday’s constant = 9.649 x 104 C/mol n = number of electrons

NERNST EQUATION - The standard reduction potential (Eo) when [A] = [B] = 1M - [B]b/[A]a = Q = reaction quotient - Concentration for gases are expressed as pressures in bars - Q = 1 for [ ] = 1 M and P = 1 bar logQ = 0 and E = Eo - Q is omitted for pure solids, liquids, and solvents

NERNST EQUATION - When a half reaction is multiplied by a factor Eo remains the same - For a complete reaction Ecell = E+ - E- and Eo = E+o - E-o E+ = potential at positive terminal E- = potential at negative terminal

NERNST EQUATION For the Cu – Fe cell at standard conditions Cu2+ + 2e- ↔ Cu(s) 0.339 V Fe2+ + 2e- ↔ Fe(s) -0.440 V Ecell = 0.779 V Galvanic Reaction Cu2+(aq) + Fe(s) ↔ Cu(s) + Fe2+(aq) Fe Fe2+ (1M) Cu2+ (1 M) Cu

NERNST EQUATION - Positive E implies spontaneous forward cell reaction - Negative E implies spontaneous reverse cell reaction If cell runs for long - Reactants are consumed - Products are formed - Equilibrium is reached - E becomes 0 - Reason why batteries run down

NERNST EQUATION At cell equilibrium at 25 oC E = 0 and Q = K (the equilibrium constant) Or Positive Eo implies K > 1 Negative Eo implies K < 1

REFERENCE ELECTRODES Indicator Electrode - Responds directly to the analyte Reference Electrode - Provides known and constant potential Examples Silver-silver chloride electrode (Ag/AgCl) Saturated Calomel electrode (SCE)

REFERENCE ELECTRODES Saturated Calomel electrode (SCE) - Saturated with KCl 1/2Hg2Cl2(s) + e- ↔ Hg(l) + Cl- E = + 0.241 V In this case, the reference is not 0.000 V (SHE) but 0.241 V (SCE)

REFERENCE ELECTRODES Saturated Calomel electrode (SCE) - Different KCl concentrations can be used - 0.1 M KCl is least temperature sensitive - Saturated KCl solution is easier to make and maintain

REFERENCE ELECTRODES Silver-Silver Chloride Electrode (Ag/AgCl) - Saturated with KCl AgCl(s) + e- ↔ Ag(s) + Cl- E = + 0.197 V

REFERENCE ELECTRODES Emeasured = Eo - 0.241 (SCE) Emeasured = Eo - 0.197 (Ag/AgCl) Eo(SHE) E(SCE) E(Ag/AgCl) Cu2+ + 2e- ↔ Cu(s) 0.339 V 0.098 V 0.142 V Fe2+ + 2e- ↔ Fe(s) -0.440 V -0.681 V -0.637 V