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Oxidation/Reduction Reactions In Cells

Electrochemistry ( 2 lectures) Mr. Zaheer E. Clarke 1:00 p.m. / 2:15 p.m. ½ full question on C10K Paper 1. Oxidation/Reduction Reactions In Cells. Most chemical reactions involve the transfer of electrons between atoms & molecules? Not always clearly seen! Eg.

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Oxidation/Reduction Reactions In Cells

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  1. Electrochemistry(2 lectures)Mr. Zaheer E. Clarke1:00 p.m. / 2:15 p.m.½full question on C10K Paper 1

  2. Oxidation/Reduction Reactions In Cells • Most chemical reactions involve the transfer of electrons between atoms & molecules? • Not always clearly seen! • Eg. • 2Mg (s) + O2(g) 2MgO(s) • May be written: • 2Mg 2Mg2+ + 4e- and • 4e- + O2 2O2-

  3. 2Mg 2Mg2+ + 4e- Mg has been oxidized (lost electrons) OIL (Oxidation Is Lost) Mg is the reducing agent • 4e- + O2 2O2- O2 has been reduced (gained electrons) RIG (Reduction Is Gain) O2 is the oxidizing agent

  4. The steps involved in the electron transfer when metallic elements (CONDUCTING ELECTRODES) & solutions of their salts (CONDUCTING SOLUTIONS) are combined can be isolated & observed clearly • oxidation/reduction • When the reactions occur spontaneously, the separation of the oxidation/reduction sites gives rise to potential difference which can drive e- through external resistive circuit • Eg. Zn-Cu couple

  5. What happens when a Zn metal strip is inserted in a CuSO4 solution? • Solution will lose its blue colour – Cu metal is being deposited on the Zn strip • Zn is going into solution as Zn2+ ions while Cu2+ is coming out of solution as metallic Cu • How can this be written in in terms of chemical equations? • Overall Reaction Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

  6. Reaction proceeds spontaneously (Gθ is –ve)& will continue until it reaches equilibrium • Half Equations Zn(s) Zn2+(aq) + 2e- Cu2+(aq) + 2e-Cu(s) Oxidation Reduction

  7. Galvanic Cell • Presently both oxidation & reduction occurs at the same site! – Zn metal strip • If we SEPARATE these sites where oxidation & reduction occurs we would have what is called a Galvanic Cell • What is a Galvanic Cell? • A Galvanic or Voltaic Cell is one in which a spontaneous chemical reaction drives electrons from Anode to Cathode in an external circuit

  8. Galvanic Cell • Example of a Galvanic Cell is a Daniell Cell (Early)

  9. Galvanic Cell Vs Electrolytic Cell • In electrolytic cells an external source of electricity is used to drive a chemical reaction e.g. electrolysis of a salt solution • Electrolysis – Anode is +ve (external electricity), cathode is –ve • Galvanic Cell – Anode is -ve, cathode is +ve • Always holds true - Anode=Oxidation Cathode= Reduction

  10. Cell Potential • EMF of a Daniell cell is 1.10V but this is not seen I practice • Factors that affect the measured potential in a cell (e.g. Daneill Cell)? • Thickness & porosity of the porous pot • Cleanness of the electrodes • Electrical Resistance of the Measuring Device • Internal Resistance • Combat • Porous pot must be as thin & porous as possible • Electrodes clean • High Resistance Voltmeter

  11. Liquid Junction Potential • The liquid junction, porous pot, is also a source of “lost” potential • Why is this? • Build-up of charge results from the different mobilities of the ions as they move across the wall of the porous pot to neutralize the charge • Potential difference exists between the inner & outer surfaces of the wall of the porous pot • This potential that results is called a LIQUID JUNCTION POTENTIAL

  12. Liquid Junction Potential How do we overcome the LIQUID JUNCTION POTENTIAL? • Use a salt bridge to reduce the effect of liquid junction potentials (i.e. potentials which arise because of the difference in mobilities of the ions) • Once these precautions are taken the EMF of the cell depends solely on the concentrations of the solutions & the metals used as the electrodes

  13. Daniell Cell with Salt Bridge instead of Porous Pot • Salt bridge consists of 5% agar jelly mixed with a saturated solution of KCl or KNO3 (K+, Cl- and NO3- have similar mobilities) • The salt bridge reduces the LJP because of the large difference in concn. of the ions in the bridge compared to in the electrolyte solutions • Effects due to mobility & availability of ions at the interface between the bridge & the solutions becomes negligible • EMF of a Daniell Cell is 1.10V when the solutions are 1M

  14. Daniell Cell • The spontaneous reaction that drives this cell is: Cu2+(aq) + Zn(s) Cu(s) + Zn2+(aq) • Cu2+ has a greater tendency to pull electrons than Zn2+ and that difference in electron pulling potential is what appears as a difference in electrical potential • Cu2+ is a better oxidizing agent than Zn2+

  15. Electrode Potential & Half Cells • If a copper/silver or a zinc/silver cell was constructed a different potential would be observed for each cell (i.e. not 1.10 V) • In a copper/silver cell, the silver is the +ve electrode and the copper is the –ve electrode • Ag+ has a greater tendency to pull electrons than Cu2+ • Ag+ is a better oxidizing agent than Cu2+ • The spontaneous reaction Cu + 2Ag+ Cu2+ + Ag EMF = 0.46V

  16. Electrode Potential & Half Cells • Each electrode, i.e. the ion & its neutral atom, [Ag+/Ag] • Contributes a characteristic potential to the overall cell potential • Independent of the other electrode in the pair • Cu | Cu2+ half cell has a characteristic potential Zn | Zn2+ half cell has a characteristic potential Ag | Ag+ half cell has a characteristic potential • To assign a potential to each half cell one must assign an electrode as a “standard electrode” & measure each electrode relative to this standard electrode

  17. Standard Hydrogen Electrode (SHE) • The standard to which all electrodes are compared is the Standard Hydrogen Electrode • Its characteristic potential is ZERO at ALL temperatures • Potentials measured against the SHE are called Reduction Potentials and are represented by Eθin Volts • The SHE is represented as: Pt(s) | H2(g) | H+(aq)

  18. Standard Electrode Potentials • Standard potentials are measured with the test electrode on the right hand side • The measured potential is +ve if the electrode has a greater tendency to pull electrons than the H2 electrode (SHE) and –ve if it has a lower tendency • Reduction Potentials • Cu2+ + 2e- Cu Eθ = + 0.34V • Zn2+ + 2e- Zn Eθ = - 0.76V • Ag+ + e- Ag Eθ = + 0.80V • Pb2+ + 2e-Pb Eθ = - 0.13V • Pb4+ + 2e-Pb2+Eθ = + 1.67V

  19. Standard Cell Notation • A vertical line represents a phase boundary while double vertical lines represent the salt bridge (no liquid junction potential) • Standard Notation for Cells is based on this assumption: • Right hand electrode is the cathode (where reduction occurs) • Daniell Cell can be written as Zn(s) | ZnSO4(aq) || CuSO4(aq) | Cu(s) or Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

  20. Standard Cell Notation & Eθ • (R) Cu2+ + 2e- Cu Eθ = + 0.34V • (L)Zn2+ + 2e- Zn Eθ = - 0.76V • Overall (R) - (L) Cu2+(aq) + Zn(s) Cu(s) + Zn2+(aq) • Overall Eθ= EθR – EθL = 0.34 – (-0.76) = 1.10V • When Eθ is +ve the reaction is spontaneous in the direction written • If the Zn electrode was written as the cathode, the Eθ would be –ve & the reaction would be spontaneous in the opposite direction

  21. Eθ– Indicator of Spontaneity • Gθ is the maximum non-expansion (useful) work available from the reaction • Gθ can be equated to the electrical work done (assuming constant pressure & temp.) as the cell runs down & reaches equilibrium • Gθ = - (electrical work that can be done by the system) = - (charge transferred) x (potential against which the charge is transferred)

  22. Eθ– Indicator of Spontaneity • Work done = -ν e- NAEθ ν– number of electrons transferred for each single oxid./red. e- – charge on each electron NA – is the single reactions per mole of reaction/Avogadro’s constant e- NA = Faraday constant = 9.6485 x 104 C mol-1 • Gθ = -νF Eθ • work is done reversibly • constant pressure & temperature

  23. Gθ & Eθ • Gθ = -νF Eθ • When Eθ is +ve, Gθ is -ve = reaction is spontaneous • When Eθ is -ve, Gθ is +ve = reaction is not spontaneous • Gθ for the Daniell Cell Gθ = -(2)(96485)(1.10) = 212267 J mol-1 = 212.3 kJ mol-1

  24. Nernst Equation • Recall Gat any stage of rxn = G + RT ln Q • -νF E = -νF Eθ + RT ln Q • E = Eθ – (RT/ νF) ln Q Nernst Equation • At unit activity of the components (a = 1), ln Q = 0 & E = Eθ • At equilibriumln Q = ln K & E = 0 (G = 0)

  25. Nernst Equation • If we have a Cu2+/Cu electrode in one half & the SHE in the other Pt(s) | H2(g) | H+(aq) || Cu2+(aq)| Cu(s)Eθ = + 0.34V (R) Cu2+ + 2e- CuEθ = + 0.34 V (L)2H+ + 2e- H2Eθ = 0 V (R) - (L) Cu2+(aq) + H2(g) Cu(s) + 2H+(aq) Eθ = + 0.34 V • Q = [aH+]2 [aCu]/[aH2][aCu2+]

  26. Nernst Equation • Q = [aH+]2 [aCu]/[aH2][aCu2+] • Perfect gas: a = p / p • Pure liquids and solids , a= 1 • For solutions at low concentration: a= [conc.]/ [1 mol dm-3] • Q = [1.00/1.00]2 [1]/[1.0/1.0][1.00/1.00] = 1 E = Eθ – (RT/ νF) ln Q E = 0.763 – 0 = 0.763 V

  27. Nernst Equation - pH • When H+ concentration is NOT 1.00 M but everything remains the same Q = [aH+]2 E = Eθ – (RT/ νF) ln{[aH+]2} E = 0.34 – (RT/ νF) ln{[aH+]2} • The measured potential is related to the activity/ concentration of H+ and Eθ of the cell • pH 4.00, E = 0.577V or pH 7.00, E = 0.754V • pH can be measured electrically • E.g. pH meter

  28. Applications • Galvanic cells are used in flashlights, clocks, watches, remote controllers as Dry Cells • Rechargeable batteries are used in cars to start engines, cell phones, video cameras, computers • Fuel Cells

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