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Reduction- Oxidation Reactions

Reduction- Oxidation Reactions. 2 nd lecture. Learning Objectives What are some of the key things we learned from this lecture?. Types of electrochemical cells. √ How to measure Standard electrode potential. √ reference electrode. √

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Reduction- Oxidation Reactions

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  1. Reduction- Oxidation Reactions 2nd lecture

  2. Learning ObjectivesWhat are some of the key things we learned from this lecture? • Types of electrochemical cells. √ • How to measure Standard electrode potential. √ • reference electrode. √ • How to measure the electrode potential from the Nernst equation. √ • Factors affecting the oxidation potential. √

  3. Electrochemical cell An electrochemical cell is a device used for generating an electromotive force (voltage) and current from chemical reactions, (galvanic, voltaic cell). or the reverse, inducing a chemical reaction by a flow of current, (electrolytic). The current is caused by the reactions releasing and accepting electrons at the different ends of a conductor.

  4. There are two types of electrochemical cells: • Electrolytic cell, (electrical Energy chem. Energy) • Galvanic cell (voltaic), (chemical energy) electrical energy)

  5. Electrolytic cell An electrolytic cell decomposes chemical compounds by means of electrical energy, in a process called electrolysis. The result is that the chemical energy is increased. An electrolytic cell has three component parts: an electrolyte and two electrodes (a cathode and an anode). The electrolyte is usually a solution of water or other solvents in which ions are dissolved. Molten salts such as sodium chloride are also electrolytes.

  6. When driven by an external voltage applied to the electrodes, the electrolyte provides ions that flow to and from the electrodes, where charge-transferring, or redox, reactions can take place. Only for an external electrical potential (i.e. voltage) of the correct polarity and large enough magnitude can an electrolytic cell decompose a normally stable, or inert chemical compound in the solution.

  7. At anode: 2Cl- Cl2 +2e-(oxidation) At cathode: 2Na+ + 2e- 2Na (reduction) Cell reaction: 2Na+ + 2Cl- 2Na + Cl2(redox reaction)

  8. 2- Galvanic Cells • In redox reaction, electrons are transferred from the oxidized species to the reduced species. • Imagine separating the two 1/2 cells physically, then providing a conduit through which the electrons travel from one cell to the other. 8H+ + MnO4- + 5e Mn2+ + 4H2O Fe2+ Fe3+ + e- x 5

  9. We need to “complete the circuit” by allowing positive ions to flow as well. We do this using a “salt bridge” which will allow charge neutrality in each cell to be maintained. Salt bridge/porous disk: allows for ion migration such that the solutions will remain neutral.

  10. Galvanic Cell Electrochemical cell in which chemical reactions are used to create spontaneous current (electron) flow

  11. Standard electrode potential (E0 ) • If the concentration of the components of the cell are unity (1M), measured at 250c • The potential is called standard potential. • It is not possible to measure the absolute potential of an electrode, but it could be measured relative to a standard reference electrode.

  12. How to measure E0 standard electrode potential? • Suppose we have Cu rod immersed in its ions (half cells). • In order to measure potential of an electrode, it must be connected to another electrode (half cell) having known constant potential called reference electrode and we use “Normal (standard) Hydrogen Electrode“. • So an electric current will flow from that having a higher potential to the other one, this current is called e.m.f. Which is the algebric difference of the electrode potential of the two half cells. • If the reference electrode has a potential equal to zero, therefore, so the e.m.f. measured will be equal to E0

  13. Reference Electrodes

  14. What is Normal Hydrogen Electrode (NHE)? • It consists of a piece of platinium foil coated electrolytically with platinum black and immersed in a solution of HCl (normal solu). • Hydrogen gas at pressure of 1 atmosphere is passed. • Platinum black layer absorbs a large amount of H2 and can be considered as a bar of H2. • Under fixed conditions: 1 atm pressure Normal solu of HCl • The potential is assumed to be zero. • By connecting the electrode of unknown potential to “NHE“, the E0 can be measured by a potentiometer. Pt wire H+ + e- ½ H2 Glass tube to contain H2 (g) (1 atm) (1 N) Bubbles of H2 (g)

  15. Standard electrode potential E0: Is the e.m.f. produced when a half cell consisiting of an element immersed in one molar solution of its ions is coupled with NHE. By measuring E0 for all elements and arrange into a series known as “Electrochemical Series “:Starting from E0 = -ve value, to hydrogen where E0 = zero, then to E0 = +ve value From this series we can conclude that: 1- The sign of E0 is similar to the charge on the metal electrode. 2- The greater the –ve value of E0, the greater the tendency of the metal to pass into ionic state (oxidized) i.e. E0 is a quantitative measure of the ability of the element to lose electrons giving its ions. 3- Metal with more negative potential will displace any other metal below it in the series from its solution. System E0 Zn/Zn2+ -0.76 Fe/Fe2+ -0.44 Sn/Sn2+ -0.13 H2(pt)/H+ 0 Cu/Cu2+ +0.34

  16. Reduction Potentials Electrode Couple "E0, V" Na+ + e- --> Na -2.7144 Mg2+ + 2e- --> Mg -2.3568 Al3+ + 3e- --> Al -1.676 Zn2+ + 2e- --> Zn -0.7621 Fe2+ + 2e- --> Fe -0.4089 Cd2+ + 2e- --> Cd -0.4022 Tl+ + e- --> Tl -0.3358 Sn2+ + 2e- --> Sn -0.141 Pb2+ + 2e- --> Pb -0.1266 2H+ + 2e- --> H2(SHE) 0 S4O62- + 2e- --> 2S2O32- 0.0238 Sn4+ + 2e- --> Sn2+ 0.1539 SO42- + 4H+ + 2e- --> H2O + H2SO3(aq) 0.1576 Cu2+ + e- --> Cu+ 0.1607 S + 2H+ + 2e- --> H2S 0.1739 AgCl + e- --> Ag + Cl- 0.2221 Saturated Calomel (SCE) 0.2412 UO22+ + 4H+ + 2e- --> U4+ + 4H2O 0.2682

  17. Reduction Potentials Hg2Cl2 + 2e- --> 2Cl- + 2Hg 0.268 Bi3+ + 3e- --> Bi 0.286 Cu2+ + 2e- --> Cu 0.3394 Fe(CN)63- + e- --> Fe(CN)64- 0.3557 Cu+ + e- --> Cu 0.518 I2 + 2e- --> 2I- 0.5345 I3- + 2e- --> 3I- 0.5354 H3AsO4(aq) + 2H+ + 2e- -->H3AsO3(aq) + H2O 0.5748 2HgCl2 + 4H+ + 2e- -->Hg2Cl2 + 2Cl- 0.6011 Hg2SO4 + 2e- --> 2Hg + SO42- 0.6152 I2(aq) + 2e- --> 2I- 0.6195 O2 + 2H+ + 2e- --> H2O2(l) 0.6237 O2 + 2H+ + 2e- --> H2O2(aq) 0.6945 Fe3+ + e- --> Fe2+ 0.769 Hg22+ + 2e- --> Hg 0.7955 Ag+ + e- --> Ag 0.7991 Hg2+ + 2e- --> Hg 0.8519 2Hg2+ + 2e- --> Hg22+ 0.9083 NO3- + 3H+ + 2e- -->HNO2(aq) + H2O 0.9275

  18. If an inert electrode such as platinum is placed in a redox system e.g. Fe3+/Fe2+ it will assume a definite potential (oxidation potential) indicative of the position of equilibrium. • Oxidation potential is the pot. difference arround an inert electrode immersed in solu. of oxidized and reduced form. If the system has reducing properties, electrons will be given up to platinum and it will acquire a negative charge. If the oxidizing tendency predominate, the system will take electrons from platinum leaving it positively charged The magnitude of the potential will be a measure of oxidizing or reducing character of the system. If +ve oxidizing and if –ve reducing.

  19. To measure the oxidation potential on pt wire, the system should be connected to NHE. Standard Oxidation potential (SOP): Is the e.m.f. produced when a half cell consisting of an inert electrode (pt) dipped in equal concentration of both oxidized and reduced form of a system is connected with NHE. Depending on the SOP the systems were arranged in a series from this series we can conclude that: 1- The higher the SOP of the system, the stronger the oxidizing power of its oxidized form and the weaker the reducing power of its reduced form. 2- the most powerful oxidizing agents are those at the top of the series with higher +ve potential, and the most powerful reducing agents occupy the bottom with higher –ve potential. There should be adifference in oxidation potential between the two systems of about 0.4 v for oxidation to occur.

  20. Nernst Equation FOR OXIDATION POTENTIAL • Compensated for non unit activity (not 1 M) • Relationship between cell potential and activities • aA + bB +ne- <--> cC + dD • At 298K,2.3RT/F = 0.0592 • What is potential of an electrode of Zn(s) and 0.01 M Zn2+ • Zn2+ +2e- <--> Zn °= -0.763 V • activity of metal is 1

  21. Eind.electrode =E0 - 0.059 log [products] n [reactants] • Eind.electrode =E0M/M+ - 0.059 log 1 n [M+]

  22. Factors affecting the oxidation potential: 1- Common ion effect: If we determine a sample of Fe2+ with KMnO4 in presence of Cl- ions we have 3 systems: MnO4-/Mn2+ (1.52) Cl2/Cl- (1.36) Fe3+/Fe2+ (0.77) KMnO4 will oxidize both Cl- and Fe2+ so more KMnO4 will be consumed. To overcome this problem the oxidation potential of MnO4-/Mn2+ should be reduced below Cl2/Cl-HOW? By adding Zimmermann´s reagent It contains [MnSO4, H3PO4, H2SO4] E25 = E0 - 0.059/n log [Mn2+] / [MnO4-] • Presence of MnSO4 increases the conc of Mn2+ ion. • So the oxidation potential will decrease below that Cl2/Cl-

  23. 2-Effect of H+ ion concentration: [H+] has a deciding effect on the oxidation potential of oxidizing agents containing oxygen and so [H+] should be included in Nernest equation. E increases with increasing [H+] i.e.increasing acidity. MnO4- + 8H+ + 5e ↔ Mn2++ 4H2O EMnO4-/Mn2+ = E0 - 0.059/5 log [Mn2+] / [MnO4-] [H+]8 At pH 5: KMnO4 oxidizes I- but not Br- and Cl- At pH 3: KMnO4 oxidizes I- and Br- but not Cl- At lower pH : KMnO4 oxidizes I- , Br- and Cl-

  24. AsO43- + 2H+ + 2e  AsO33- + H2O EAsO43-/ AsO33-= E0 - 0.059/2 log [ASO33-] / [AsO43-] [H+]2 If [H+] increases: The oxidation potential of AsO43- increases. If [H+] decreases: The oxidation potential of AsO43-decreases and the reducing power of AsO33- increases. Therefore: AsO43- oxidizes I- to I2 in acid medium. While: I2 oxidizes AsO3-3 in alkaline medium with NaHCO3 AsO4-3 + 2 I- + 2H+ AsO3-3 + I2 + H2O acid alkaline

  25. 3-Effect of complexing agents: E.g.1- Oxidation potential of I2/2I- system increases in presence of HgCl2 WHY? Because HgCl2 forms a complex with the I- ions [HgI4]2- I2 + 2 e  2I- EI2 /2I-= E0- • removing the I- ions from the share of the reaction. • minimizing its concentration. • decreasing the ratio of [I-]2 / I2 . • increasing the oxidation potential of I2 /2I- system .

  26. E.g.2- Oxidation potential of Fe3+/ Fe2+ system decreases on addition of F- or PO43-WHY? Because F- and PO43-form stable complexs with Fe3+ which are [FeF6]3- or [Fe (PO4)2]-3 Fe3++ e  Fe2+ EFe3+/Fe2+ = E0 - • removing the Fe3+ ions from the share of the reaction. • minimizing its concentration. • Increasing the ratio of Fe2+/ Fe3+ . • lowering the oxidation potential of Fe3+/ Fe2+ system .

  27. 4-Effect of precipitating agents: Oxidation potential of Ferricyanide/ Ferrocyanide system increases on addition of Zn2+ WHY? Because Zn2+ will precipitate Zn2[Fe(CN)6] Fe (CN)63- + e  Fe(CN)64- E = Eo - • removing the Ferrocyanideions from the share of the reaction. • minimizing conc of ferrocyanide. • decreasing the ratio of [Fe(CN)6]4-/ [Fe (CN)6]3- . • increasing the oxidation potential of [Fe(CN)6]3-/ [Fe (CN)6]4-

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