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PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 18

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

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PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 18

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  1. PRINCIPLES OF CHEMISTRY II CHEM 1212CHAPTER 18 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university

  2. CHAPTER 18 ELECTROCHEMISTRY

  3. ELECTROCHEMISTRY - Is the study of the relations between chemical reactions and electricity - Electrochemical processes involve the transfer of electrons from one substance to another

  4. 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)

  5. NONREDOX REACTIONS - There is no transfer of electrons from one reactant to another reactant Examples BaCO3(s) → BaO(s) + CO2(g) Double-replacement reactions

  6. 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)

  7. 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)

  8. 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)

  9. 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

  10. OXIDATION NUMBERS CH4 x +1 1(x) + 4(+1) = 0 x = -4

  11. OXIDATION NUMBERS NO3- x -2 1(x) + 3(-2) = -1 x = +5

  12. 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

  13. 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) + 2e- The reduction half-reaction is: Cu2+(aq) + 2e- → Cu(s)

  14. 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

  15. BALANCING REDOX EQUATIONS Half-Reaction Method Acidic Solutions

  16. 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

  17. 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)

  18. BALANCING REDOX EQUATIONS Half-Reaction Method Basic Solutions

  19. 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

  20. 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)

  21. 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

  22. 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

  23. 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)

  24. 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

  25. 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)

  26. 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

  27. 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

  28. 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)

  29. 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)

  30. 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)

  31. 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 = n x F (n = number of moles)

  32. 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)

  33. 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

  34. 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

  35. 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

  36. 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

  37. 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

  38. 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

  39. 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

  40. STANDARD POTENTIALS - 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

  41. 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)

  42. 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

  43. ∆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

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

  45. 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

  46. 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

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

  48. REFERENCE ELECTRODES - Provide known and constant potential Examples Silver-silver chloride electrode (Ag/AgCl) Saturated Calomel electrode (SCE)

  49. INDICATOR ELECTRODES - Respond directly to the analyte - Two Classes of Indicator Electrodes Metal Electrodes - Surfaces on which redox reactions take place Examples Platinum Silver

  50. 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

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