Chapter 5. Oxidation and reduction us1 Chapter 5. Oxidation and reduction Chemistry 481(01) Spring 2009 Chemistry 481(01) Spring 2009 Chemistry 481(01) Spring 2009 Chemistry 481(01) Spring 2009 Instructor: Instructor: Dr. Upa Dr. Upali i Sir Siriwardane wardane Red Reduct Red Reduct ctio ctio ion p ion p n potent n potent tentia tentia ials ials ls ls e-mail: upali@chem.latech.edu e-mail: upali@c .latech.edu 5.1 Redox 5.1 Redox half half- -reactions reactions Office: CTH 31 Office: CTH 311 Phone 257-4941 1 Phone 257-4941 5.2 Standard potentials 5.2 Standard potentials Office Hours: Office Hours: 5.3 Trends in standard potentials 5.3 Trends in standard potentials M,W 8:00-9:00 & 11:00-12:00 am; 5.4 The electrochemical series 5.4 The electrochemical series Tu,Th,F 10:00 - 12:00 a.m. 5.5 The Nernst equation 5.5 The Nernst equation 5.6 Reactions with water 5.6 Reactions with water March 26 , 2009(Test 1): Chapters 1,2, 3 5.7 Oxidation by atmospheric oxygen 5.7 Oxidation by atmospheric oxygen April 28, 2009 (Test 2): Chapters 5,7 & 8. 5.8 Disproportionation 5.8 Disproportionation and and comproportionation comproportionation May 19, 2008 (Test 3) Chapters 18 & 19 5.9 The influence of complexation 5.9 The influence of complexation May 21, Make Up: Comprehensive covering all Chapters 1 5 - 1 5 - 2 Oxidation /Reductions Conventions Oxidation /Reductions Conventions Electrochemistry Review Electrochemistry Review Electrochemical Cells Electrochemical Cells Voltaic Cells Voltaic Cells Standard Cell Potentials Standard Cell Potentials Effect of Concentration on Cell Potentials Effect of Concentration on Cell Potentials Free Energy and Cell Potential Free Energy and Cell Potential Batteries Batteries Corrosion Corrosion Electrolytic Cells Electrolytic Cells Stoichiometry of Electrochemical Reactions of Electrochemical Reactions Stoichiometry Practical Application: pH Electrode Practical Application: pH Electrode 5-3 5 - 3 5 - 4 1
Slide 2 us1 Upali Siriwardane, 3/26/2008
Rules for assigning oxidation numbers Rules for assigning oxidation numbers Assigning Oxidation Numbers Assigning Oxidation Numbers The oxidation number of a free element = 0. The oxidation number of a monatomic ion = charge on the ion. K 2 CO 3 The oxidation number of hydrogen = + 1 and rarely - The sum of all the oxidation numbers in this 1. formula equal 0. Multiply the subscript by the The oxidation number of oxygen = - 2 and in oxidation number for each element. peroxides - 1. To calculate O.N. of C The sum of the oxidation numbers in a polyatomic K = (2) ( + 1 ) = + 2 ion = charge on the ion. O = (3) ( - 2 ) = - 6 Elements in group 1, 2, and aluminum are always as therefore, C = (1) ( + 4 ) = + 4 indicated on the periodic table. 5 - 5 5 - 6 Reducing Agents and Oxidizing Agents Reducing Agents and Oxidizing Agents Balancing Redox Balancing Redox Equations by the Equations by the Half- -reaction Method reaction Method Half Reduc Reducing agent ng agent - the reactant that gives up Reducing agent Reduc ng agent electrons. Decide what is reduced (oxidizing agent) and what is oxidized (reducing agent). The reducing agent contains the element that is oxidized (looses electrons). Oxidizing agent - Write the reduction half-reaction. the reactant that gains electrons. Write the oxidation half-reaction. Oxid idiz izing ing ag agent ent - contains the element that is Oxid idiz izing ing ag agent ent The number of electrons gained must equal the reduced (gains electrons). number of electrons lost. If a substance gains electrons easily, it is said to Add the two half-reactions. be a strong oxidizing agent.If a substance Simplify the equation. gives up electrons easily, it is said to be a Check to see that electrons, elements, and total strong reducing agent charge are balanced. 5 - 7 5 - 8 2
Types of electrochemical cells Types of electrochemical cells Types of electrochemical cells Types of electrochemical cells Not a Not all reactions are l reactions are revers reversible. Non e. Non Galvanic or Voltaic Galvanic or Voltaic rechargeab rechargeable le The ‘spontaneous’ reaction. ∆ G=- nFE Examp amples of es of non non- -re re revers versible reactions ble reactions Examp amples of es of non non revers versible reactions ble reactions Produces electrical energy. If a gas If a gas is is produced produced which escape which escapes. Electrolytic 2H + + 2 Electrolytic 2H + 2 e e - H 2 (g (g) Non-spontaneous reaction. If one or more of If one or more of the e sp spec ecies ies Requires electrical energy to occur. decomposes. decomp oses. Some me rever reversible sible and rechargeable and rechargeable For reversible cells, the galvanic reaction Examp Examples es of of rever reversib sible reactions e reactions can occur spontaneously and then be Examp Examples es of of rever reversib sible reactions e reactions 2+ + N Cd Cd 2+ + Ni (s) Cd Cd (s) (s) + Ni + Ni 2+ 2+ reversed electrolytically - rechargeable (s) batteries. Pb (s) + PbO 2 (s) + 2H + (aq) + 2HSO 4 - (aq) 2PbSO 4 (s) + 2H 2 O 5-9 5-10 Zinc- -carbon dry cell carbon dry cell Zinc- -carbon dry cell carbon dry cell Zinc Zinc The e The electrolyte, aque ectrolyte, aqueous ous NH NH 4 Cl Cl is made int made into a a paste paste by adding a by adding an inert inert filler. filler. Seal Electr ectrochemica chemical reaction reaction Electr ectrochemica chemical reaction reaction Carbon rod Zn (s) Zn (s) + 2 + 2MnO 2 (s) (s) + 2 NH + 2 NH 4 - (aq) (aq) Paste Zinc Zn Zn 2+ 2+ (aq) (aq) + Mn + Mn 2 O 3 (s) (s) + 2 + 2NH 3 (aq) (aq) + H + H 2 O (l) (l) This cell This cell has a pote has a potential of 1.5 V when ntial of 1.5 V when new. new. 5-11 5-12 3
Lead storage battery Lead storage battery Lead storage battery Lead storage battery These are used when a large capacity and Electrochemical reaction. Electrochemical reaction. moderately high current is need. 2PbSO 4 (s) + 2H 2 O (l) It has a potential of 2 V. Unlike the zinc-carbon dry cell, it can be Pb (s) + PbO 2 (s) + 2H + (aq) + 2HSO 4 - (aq) recharged by applying a voltage. Car battery. Car battery. Note. Note. Lead changes from a +2 to 0 and +4 This is the most common application. oxidation state when a lead storage battery Most cars are designed to use a 12 V battery. is discharged. As a result, six cells connected in a series are needed. Lead also remains in a solid form. 5-13 5-14 Lead storage battery Voltaic cells Lead storage battery Voltaic cells Electrochemical cells Electrochemical cells Each half reaction is put in a separate ‘half cell.’ They can then be connected A series of A series of electrically. 6 cells in 6 cells in series are series are used to used to produce the produce the 12 volts that 12 volts that most cars most cars require. require. This permits better control over the system. 5-16 5-15 4
Voltaic cells Voltaic cells Voltaic cells Voltaic cells Cu 2+ + Zn (s) Cu (s) + Zn 2+ e - e - e - e - Electrons are Electrons are To complete the To complete the transferred from Cu transferred from Zn circuit, a salt circuit, a salt one half-cell to one half-cell to bridge is used. bridge is used. the other using the other using an external metal an external metal conductor. conductor. Zn 2+ Cu 2+ salt bridge 5-17 5-18 Voltaic cells Voltaic cells Cell diagrams Cell diagrams For our other half cell, we have copper metal being produced. Rather than drawing an entire cell, a type of shorthand can be used. Reduction at Cu electrode The electrode is the cathode is For our copper - zinc cell, it would be: negative (-) Zn | Zn 2+ (1M) || Cu 2+ (1M) | Cu Oxidation at Zn electrode The electrode is the anode is ∆ G= - nFE The anode is always on the left. positive (+) Cu 2+ + 2e - Cu; E o = +0.337V | = boundaries between phases || = salt bridge Zn 2+ + 2e - Zn; E o = -0.763V Other conditions like concentration are E o cell = E o half-cell of reduction - E o half-cell of oxidation listed just after each species. E o cell = 0.34 - ( -0.763V) = 1.03 V 5-19 5-20 5
Cell diagrams Cell diagrams Electrode potentials Electrode potentials Other examples A measure of how willing a species is to gain or Other examples lose electrons. + (1M) || Pt, H 2 (1atm) | H + (1M) || Pt, H 2 (1atm) | H This is the SHE. Pt is used to maintain Standard potentials Standard potentials electrical contact so is listed. The pressure of H 2 is given in atmospheres. Potential of a cell acting as a cathode compared to a standard hydrogen electrode. Pt, H 2 (1atm) | HCl HCl (0.01M) || Ag+ (sat) | Ag (0.01M) || Ag+ (sat) | Ag Pt, H 2 (1atm) | A saturated silver solution (1.8 x 10 -8 M) Values also require other standard based on the K SP AgCl and [Cl - ] conditions. 5-21 5-22 Standard hydrogen electrode Electrode potentials Standard hydrogen electrode Electrode potentials Standard potentials are defined using Hydrogen electrode (SHE) Hydrogen electrode (SHE) H 2 specific concentrations. The ultimate reference electrode. All soluble species are at 1 M H 2 is constantly bubbled into a 1 M HCl solution Slightly soluble species must be at saturation. Pt | H 2 (1atm), 1M H + || Pt black Any gas is constantly introduced at 1 atm plate E o = 0.000 000 V Any metal must be in electrical contact Other solids must also be present and in All other standard potentials 1 M HCl contact. are then reported relative to SHE. 5-23 5-24 6
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