Electrochemistry Part IV: Spontaneity & Nernst Equation

Slides:



Advertisements
Similar presentations
Electricity from Chemical Reactions
Advertisements

Electrochemical Cells. Definitions Voltaic cell (battery): An electrochemical cell or group of cells in which a product-favored redox reaction is used.
Electrochemistry II. Electrochemistry Cell Potential: Output of a Voltaic Cell Free Energy and Electrical Work.
Slide 1 of E cell, ΔG, and K eq  Cells do electrical work.  Moving electric charge.  Faraday constant, F = 96,485 C mol -1  elec = -nFE ΔG.
Chapter 18 Electrochemistry
ELECTROCHEMISTRY CHARGE (Q) – A property of matter which causes it to experience the electromagnetic force COULOMB (C) – The quantity of charge equal to.
The End is in Site! Nernst and Electrolysis. Electrochemistry.
Electrochemistry Chapter 19.
Redox Reactions and Electrochemistry
Redox Reactions and Electrochemistry
Chapter 20 Electrochemistry
Calculation of the standard emf of an electrochemical cell The procedure is simple: 1.Arrange the two half reactions placing the one with.
Electrochemistry Chapter 19 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Chapter 21: Electrochemistry II
Electrochemistry Mechanical work  G = Electrical work w = n = w = w max # moles e - F =charge on 1 mol e -  = electrical potential - P ext VV - nF.
Electrochemistry The study of the interchange of chemical and electrical energy. Sample electrochemical processes: 1) Corrosion 4 Fe (s) + 3 O 2(g) ⇌
Electrochemistry: Oxidation-Reduction Reactions Zn(s) + Cu +2 (aq)  Zn 2+ (aq) + Cu(s) loss of 2e - gaining to 2e - Zinc is oxidized - it goes up in.
Electrochemistry Chapter 3. 2Mg (s) + O 2 (g) 2MgO (s) 2Mg 2Mg e - O 2 + 4e - 2O 2- Oxidation half-reaction (lose e - ) Reduction half-reaction.
ELECTROCHEMISTRY Chap 20. Electrochemistry Sample Exercise 20.6 Calculating E° cell from E° red Using the standard reduction potentials listed in Table.
Inorganic chemistry Assistance Lecturer Amjad Ahmed Jumaa  Predicting whether a (redox) reaction is spontaneous.  Calculating (ΔG°)
Chapt. 18 Electrochemistry Sec. 5  G° from E° © University of South Carolina Board of Trustees.
Redox Reactions and Electrochemistry Chapter 19. Cell Potentials E cell  = E red  (cathode) − E red  (anode) = V − (−0.76 V) = V.
ELECTROCHEMISTRY CHARGE (Q) – A property of matter which causes it to experience the electromagnetic force COULOMB (C) – The quantity of charge equal to.
Electrochemistry Part Four. CHEMICAL CHANGE  ELECTRIC CURRENT To obtain a useful current, we separate the oxidizing and reducing agents so that electron.
Chapter There is an important change in how students will get their AP scores. This July, AP scores will only be available online. They will.
© 2015 Pearson Education, Inc. Chapter 20 Electrochemistry James F. Kirby Quinnipiac University Hamden, CT Lecture Presentation.
1 Electrochemistry Chapter 18 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Chapter 20: Electrochemistry. © 2009, Prentice-Hall, Inc. Electrochemical Reactions In electrochemical reactions, electrons are transferred from one species.
A redox reaction is one in which the reactants’ oxidation numbers change. What are the oxidation numbers of the metals in the reaction below? The.
CHAPTER SIX(19) Electrochemistry. Chapter 6 / Electrochemistry Chapter Six Contains: 6.1 Redox Reactions 6.2 Galvanic Cells 6.3 Standard Reduction Potentials.
Free Energy ∆G & Nernst Equation [ ]. Cell Potentials (emf) Zn  Zn e volts Cu e-  Cu volts Cu +2 + Zn  Cu + Zn +2.
Electrochemistry.
Chapter 20 Electrochemistry
Chapter 20 Electrochemistry
The study of the relationships between electricity and chem rxns
Electrochemistry.
Electrochemistry Chapter 18.
Electrochemistry Chapter 19.
Description of a Galvanic Cell.
Chapter 20 - Electrochemistry
Dr. Aisha Moubaraki CHEM 202
Redox Reactions and Electrochemistry
Electrochemistry the study of the interchange of chemical and electrical energy.
Electrochemistry Chapter 19
Unit 12 Electrochemistry
Cell Potential and the Nernst Equation
Chapter Thermodynamics and Electrochemistry
Chapter 17: Electrochemistry
Electrochemistry.
Thermodynamics Part II: Gibbs Free Energy & Equilibrium
Electrochemistry.
Chapter 20 Electrochemistry
Electrochemistry Part IV: Spontaneity, Nernst Equation & Electrolysis
Electrochemistry The study of the interchange
Electrochemistry Chapter 19
Electrochemistry Chapter 19
Chapter 20 Electrochemistry

From Voltage Cells to Nernst Equation
THE VOLTAIC (GALVANIC) ELECTROCHEMICAL CELL. 4/26
Electrochemistry.
Electrochemistry Part III: Reduction Potentials
EQUILIBRIUM AND SPONTANEITY
Chapter 20 Electrochemistry
Electrochemistry Chapter 19
Electrochemistry Part III: Reduction Potentials
Electrochemistry Chapter 19
ELECTROCHEMISTRY CHARGE (Q) – A property of matter which causes it to experience the electromagnetic force COULOMB (C) – The quantity of charge equal.
Presentation transcript:

Electrochemistry Part IV: Spontaneity & Nernst Equation Jespersen Chap. 20 Sec 4 & 5 Skipping Sec 6 & 8 Dr. C. Yau Fall 2014 1 1

Spontaneity of Reaction We know to have a spontaneous rxn… E > 0 ΔG < 0 How are these two related? ΔG = - n FEcell where n = moles of e- F = Faraday's constant 9.65x104 C/mol e- (remember 1 V = 1J/C) ΔGo = - n FEocell (under standard conditions) 2

NiO2 (s) + 2Cl(aq) + 4H+ (aq)  Cl2 (g) + Ni2+ (aq) + 2H2O (l) Example 20.7 p. 937 Calculate ΔGo for the reaction, given that its standard cell potential is 0.320 V at 25oC. NiO2 (s) + 2Cl(aq) + 4H+ (aq)  Cl2 (g) + Ni2+ (aq) + 2H2O (l) ΔGo = - n FEocell F = 9.65x104 C/mol e- (1 V = 1J/C, so Eo = 0.320 J/C) How do we figure out what n is? Ans. -61.8 kJ Do Pract Exer 13 & 14 p. 938

Calculating K from Cell Potential We know ΔGo = - n FEocell We also know ΔGo = - RT ln K (Chap. 19) so - n FEo = - RT ln K Example 20.8 p. 789 Calculate K for the reaction in Example 20.8. NiO2 (s) + 2Cl(aq) + 4H+ (aq)  2Cl2 (g) + Ni2+ (aq) + 2H2O (l) Collect all the constants we need. Do Pract Exer 15 & 16 p. 939

Derivation of the Nernst Eqn What happens when it is not under standard conditions? Divide both sides of eqn by (-nF) we get... ΔGo = - n FEocell ΔG = - n FEcell Nernst Equation

Common simplified version of the Nernst Equation for 25.0 oC: This version of Nernst Eqn will be given also, but remember it’s only for 25.0oC

Nernst Equation Eo is the cell potential under standard conditions (for aqueous soln, 1M) What if it is not 1 M? Example 20.9 p. 940 Suppose a galvanic cell employs the following: Ni2+ + 2e-  Ni Eo = - 0.25 V Cr3+ + 3e-  Cr Eo = - 0.74 V Calculate the cell potential when [Ni2+] = 4.87x10-4M and [Cr3+] = 2.48x10-3 M This type of quest will be on your final exam. Ans. +0.44 V

The rxn of tin metal with acid can be written as Example 20.10 p. 941 The rxn of tin metal with acid can be written as Sn (s) + 2H+ (aq)  Sn2+ (aq) + H2 (g) Calculate the cell potential (a) when the system is at standard state. (b) when the pH = 2.00 (c) when the pH is 5.00. Assume that [Sn2+] = 1.00 M and the partial pressure of H2 is also 1.00 atm. Do Pract Exer 17, 18, 20 p. 942 Ans. +0.02 V Ans. -0.16V

What we are skipping in Chap. 20: Concentration from E Measurements Sec 20.6 Electricity Batteries: Lead Storage Batteries Zinc-Manganese Dioxide Cells (LeClanche cell) Nickel-Cadmium Rechargeable Batteries Nickel-Metal Hydride Batteries Lithium Batteries Lithium Ion Cells Fuel Cells Photovoltaic Cells