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PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 18 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university CHAPTER 18 ELECTROCHEMISTRY ELECTROCHEMISTRY - Is the study of the relations between chemical reactions and electricity - Electrochemical processes involve the transfer of electrons from one substance to another OXIDATION-REDUCTION REACTIONS - Also called redox reactions - Involve transfer of electrons from one species to another Oxidation - loss of electrons Reduction - gain of electrons - Ionic solid sodium chloride (Na+ and Cl- ions) formed from solid sodium and chlorine gas 2Na(s) + Cl2(g) → 2NaCl(s) - The oxidation (rusting) of iron by reaction with moist air 4Fe(s) + 3O2(g) → 2Fe2O3(s) NONREDOX REACTIONS - There is no transfer of electrons from one reactant to another reactant Examples BaCO3(s) → BaO(s) + CO2(g) Double-replacement reactions OXIDATION NUMBER (STATE) The concept of oxidation number - provides a way to keep track of electrons in redox reactions - not necessarily ionic charges Conventionally - actual charges on ions are written as n+ or n- oxidation numbers are written as +n or -n Oxidation - increase in oxidation number (loss of electrons) Reduction - decrease in oxidation number (gain of electrons) OXIDATION NUMBERS 1. Oxidation number of uncombined elements = 0 Na(s), O2(g), H2(g), Hg(l) 2. Oxidation number of a monatomic ion = charge Na+ = +1, Cl- = -1, Ca2+ = +2, Al3+ = +3 3. Oxygen is usually assigned -2 H2O, CO2, SO2, SO3 Exceptions: H2O2 (oxygen = -1) and OF2 (oxygen = +2) OXIDATION NUMBERS 4. Hydrogen is usually assigned +1 Exceptions: -1 when bonded to metals (+1in HCl, NH3, H2O and -1in CaH2, NaH) 5. Halogens are usually assigned -1 (F, Cl, Br, I) Exceptions: when Cl, Br, and I are bonded to oxygen or a more electronegative halogen (Cl2O: O = -2 and Cl = +1) 6. The sum of oxidation numbers for - neutral compound = 0 - polyatomic ion = charge (H2O = 0, CO32- = -2, NH4+ = +1) OXIDATION NUMBERS CO2 The oxidation state of oxygen is -2 CO2 has no charge The sum of oxidation states of carbon and oxygen = 0 1 carbon atom and 2 oxygen atoms 1(x) + 2(-2) = 0 x = +4 CO2 x -2 for each oxygen OXIDATION NUMBERS CH4 x +1 1(x) + 4(+1) = 0 x = -4 OXIDATION NUMBERS NO3x -2 1(x) + 3(-2) = -1 x = +5 HALF-REACTIONS - Just the oxidation or the reduction is given - The transferred electrons are shown Oxidation Half-Reaction - Electrons are on the product side of the equation Reduction Half-Reaction - Electrons are on the reactant side of the equation HALF-REACTIONS For the redox reaction Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) Zn is oxidized (oxidation number changes from 0 to +2) Cu is reduced (oxidation number changes from +2 to 0) The oxidation half-reaction is: Zn(s) → Zn2+(aq) + 2eThe reduction half-reaction is: Cu2+(aq) + 2e- → Cu(s) OXIDIZING AND REDUCING AGENTS Oxidizing Agent - Is the reduced species (accepts electrons from another species) Reducing Agents - Is the oxidized species (donates electrons to another species) For the redox reaction Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) Cu is reduced so is the oxidizing agent Zn is oxidized so is the reducing agent BALANCING REDOX EQUATIONS Half-Reaction Method Acidic Solutions BALANCING REDOX EQUATIONS - Write separate oxidation and reduction half-reactions - Balance all the elements except hydrogen and oxygen in each - Balance oxygen using H2O(l), hydrogen using H+(aq), and charge using electrons (e-) - Multiply both half-reactions by suitable factors to equalize electron count - Combine the balanced half-reactions BALANCING REDOX EQUATIONS Balance the following redox reaction in acid medium MnO4-(aq) + Fe2+(aq) → Fe3+(aq) + Mn2+(aq) Answer MnO4-(aq) + 5Fe2+(aq) + 8H+(aq) → 5Fe3+(aq) + Mn2+(aq) + 4H2O(l) BALANCING REDOX EQUATIONS Half-Reaction Method Basic Solutions BALANCING REDOX EQUATIONS - Balance equation as if it were acidic - Note H+ ions and add same number of OH- ions to both sides - Cancel H+ and OH- (=H2O) with H2O on other side BALANCING REDOX EQUATIONS Balance the following redox reaction in basic medium MnO4-(aq) + C2O42-(aq) → MnO2(s) + CO32-(aq) Answer 2MnO4-(aq) + 3C2O42-(aq) + 4OH-(aq) → 2MnO2(s) + 6CO32-(aq) + 2H2O(l) 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 ELECTRODE Chemically Inert Electrodes - Do not participate in the reaction Examples Carbon, Gold, Platinum, ITO Reactive Electrodes - Participate in the reaction Examples Silver, Copper, Iron, Zinc 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) VOLTAIC (GALVANIC) CELL - Spontaneous reaction - Produces electrical energy from chemical 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 VOLTAIC (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) - Each is called a half-cell - 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 VOLTAIC (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 VOLTAIC (GALVANIC) CELL For the overall reaction Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) e- Voltmeter - Zn electrode e- + Cu electrode ClK+ Zn2+ Salt bridge (KCl) Anode Oxidation Zn(s) → Zn2+(aq) + 2e- Cu2+ Cathode Reduction Cu2+(aq) + 2e- → Cu(s) POTENTIALS VOLTAIC CELL Voltage or Potential Difference (E) - Referred to as the electromotive force (emf) - 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 (larger emf) POTENTIALS OF VOLTAIC CELL Voltage or 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) POTENTIALS OF VOLTAIC CELL Charge 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) = 9.6485 x 104 C/mol = Faraday constant (F) q=nxF (n = number of moles) POTENTIALS OF VOLTAIC CELL Current - The quantity of charge flowing past a point in an electric circuit per second Units Ampere (A) = coulomb per second (C/s) Charge (C) = current (A) x time (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 atm 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 atm) Conventionally, Eo = 0 for SHE STANDARD POTENTIALS The Eo for Ag+(aq) + e- ↔ Ag(s) is +0.799 V Implies - 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 STANDARD POTENTIALS Silver does not react spontaneously with hydrogen 2H+(aq) + 2e- → H2(g) Ag+(aq) + e- → Ag(s) Eo = 0.000 V Eo = +0.799 V Reverse the second equation (sign changes) Ag(s) → Ag+(aq) + eEo = -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) Zn2+(aq) + 2e- → Zn(s) Eo = +0.339 V Eo = -0.762 V Reverse the second equation (sign changes) Zn(s) → Zn2+(aq) + 2eEo = +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 - Refer to series and Eo values in textbook - For combining two half reactions, the one higher in the series proceeds spontaneously as reduction under standard conditions - The one lower in the series proceeds spontaneously as oxidation under standard conditions Formal Potential - The potential for a cell containing a specified concentration of reagent other than 1 M STANDARD POTENTIALS - When a half reaction is multiplied by a factor Eo remains the same - For a complete reaction Ecell = E+ - Eand Eo = E+o - E-o E+ = potential at positive terminal E- = potential at negative terminal STANDARD POTENTIALS 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 (overall) Reaction Cu2+(aq) + Fe(s) ↔ Cu(s) + Fe2+(aq) STANDARD POTENTIALS - 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 ∆G AND Keq - Recall that spontaneous reaction has a negative value of ∆G ∆G = -nFE n = number of moles of electrons transferred F = Faraday constant = 9.6485 x 104 C/mol E = cell potential (V or J/C) Under Standard Conditions ∆Go = -nFEo NERNST EQUATION For the half reaction aA + ne- ↔ bB The half-cell potential (at 25 oC), E, is given by RT B b EE ln a nF A O b 0.0591 B E EO log a n A NERNST EQUATION Eo = standard electrode potential R = gas constant = 8.314 J/K-mol T = absolute temperature F = Faraday’s constant = 9.6485 x 104 C/mol n = number of moles of electrons transferred NERNST EQUATION - The standard reduction potential (Eo) is when [A] = [B] = 1M - [B]b/[A]a = Q = reaction quotient - Concentration for gases are expressed as pressures in atm - Q = 1 for [ ] = 1 M and P = 1 atm logQ = 0 and E = Eo - Pure solids, liquids, and solvents do not appear in Q expression NERNST EQUATION At cell equilibrium at 25 oC E = 0 and Q = Keq (the equilibrium constant) 0.0591 E logK eq n o Or Keq 10 nEo /0.05916 Positive Eo implies Keq > 1 Negative Eo implies Keq < 1 REFERENCE ELECTRODES - Provide known and constant potential Examples Silver-silver chloride electrode (Ag/AgCl) Saturated Calomel electrode (SCE) INDICATOR ELECTRODES - Respond directly to the analyte - Two Classes of Indicator Electrodes Metal Electrodes - Surfaces on which redox reactions take place Examples Platinum Silver INDICATOR ELECTRODES - Respond directly to the analyte - Two Classes of Indicator Electrodes - Ion-Selective Electrodes - Selectively binds one ion (no redox chemistry) Examples pH (H+) electrode Calcium (Ca2+) electrode Chloride (Cl-) electrode 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- Electroplating ELECTROLYSIS CELL - Consists of two electrodes in an electrolyte solution - Nonspontaneous reaction - Requires electrical energy to occur CORROSION - Oxidation of a metal to produce compounds of the metal Examples - Rusting of iron (Fe forms Fe2O3·xH2O - Copper in bronze forming copper(II) compounds (green color) Prevention - Painting or coating - Plating of iron with chromium - Anodic protection (metal is oxidized under controlled conditions) - Cathodic protection (a more reactive metal is placed in contact)