Cell EMF Eocell = Eored(cathode) - Eored(anode)

Slides:

Advertisements

Similar presentations

Inorganic chemistry Assiastance Lecturer Amjad Ahmed Jumaa  Calculating the standard (emf) of an electrochemical cell.  Spontaneity.

Advertisements

Oxidation Reduction Regents Review.

Galvanic Cells What will happen if a piece of Zn metal is immersed in a CuSO 4 solution? A spontaneous redox reaction occurs: Zn (s) + Cu 2 + (aq) Zn 2.

Electrochemical Cells. Definitions Voltaic cell (battery): An electrochemical cell or group of cells in which a product-favored redox reaction is used.

Chapter 17 Electrochemistry

Electrochemistry II. Electrochemistry Cell Potential: Output of a Voltaic Cell Free Energy and Electrical Work.

Chapter 20 Electrochemistry.

Lecture 263/30/07. E° F 2 (g) + 2e - ↔ 2F Ag + + e - ↔ Ag (s)+0.80 Cu e - ↔ Cu (s)+0.34 Zn e - ↔ Zn (s)-0.76 Quiz 1. Consider these.

Lecture 254/4/05 Seminar today. Standard Reduction Potentials 1. Each half-reaction is written as a reduction 2. Each half-reaction can occur in either.

ELECTROCHEMISTRY Chap 20.

Electron-Transfer Reactions Cu 2+ (aq) + Zn(s)  Cu(s) + Zn 2+ (aq)

Prentice Hall © 2003Chapter 20 Zn added to HCl yields the spontaneous reaction Zn(s) + 2H + (aq)  Zn 2+ (aq) + H 2 (g). The oxidation number of Zn has.

Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Electrochemistry The study of the interchange of chemical and electrical energy.

Electron-Transfer Reactions Cu 2+ (aq) + Zn(s)  Cu(s) + Zn 2+ (aq)

Electrochemistry Part 1 Ch. 20 in Text (Omit Sections 20.7 and 20.8) redoxmusic.com.

Chapter 17 Electrochemistry 1. Voltaic Cells In spontaneous reduction-oxidation reactions, electrons are transferred and energy is released. The energy.

Electrochemical Reactions

Predicting Spontaneous Reactions

Electrochemistry AP Chapter 20. Electrochemistry Electrochemistry relates electricity and chemical reactions. It involves oxidation-reduction reactions.

Electrochemistry Chapter 19.

Chapter Oxidation states 20.2 half reactions.

Chapter 20 – Redox Reactions One of the earliest recognized chemical reactions were with oxygen. Some substances would combine with oxygen, and some would.

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. 2Mg (s) + O 2 (g) 2MgO (s) 2Mg 2Mg e - O 2 + 4e - 2O 2- Oxidation half-reaction (lose e - ) Reduction half-reaction.

A Summary of Redox Terminology

Section 10 Electrochemical Cells and Electrode Potentials.

Chapter 21: Electrochemistry II

Electrochemistry  G = Electrical work w = n = w max # moles e - F =charge on 1 mol e -  = electrical potential - nF  w max =  H - T  S.

Electrochemisty Electron Transfer Reaction Section 20.1.

Oxidation-Reduction Reactions Chapter 4 and 18. 2Mg (s) + O 2 (g) 2MgO (s) 2Mg 2Mg e - O 2 + 4e - 2O 2- _______ half-reaction (____ e - ) ______________________.

Redox Reactions & Electrochemistry Chapter 19 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

ELECTROCHEMISTRY Chap 20. Electrochemistry Sample Exercise 20.6 Calculating E° cell from E° red Using the standard reduction potentials listed in Table.

John E. McMurry Robert C. Fay C H E M I S T R Y Chapter 17 Electrochemistry.

CHM 112 Summer 2007 M. Prushan Chapter 18 Electrochemistry.

Redox Reactions and Electrochemistry Chapter 19. Voltaic Cells In spontaneous oxidation-reduction (redox) reactions, electrons are transferred and energy.

Electrochemistry Electrochemical Cells –Galvanic cells –Voltaic cells –Nernst Equation –Free Energy.

19.4 Spontaneity of Redox Reactions  G = -nFE cell  G 0 = -nFE cell 0 n = number of moles of electrons in reaction F = 96,500 J V mol = 96,500 C/mol.

Chapter 17 Electrochemistry

The Nernst Equation Standard potentials assume a concentration of 1 M. The Nernst equation allows us to calculate potential when the two cells are not.

1 CELL POTENTIAL, E Electrons are “driven” from anode to cathode by an electromotive force or emf.Electrons are “driven” from anode to cathode by an electromotive.

Nernst Equation Walther Nernst

Electrochemistry Combining the Half-Reactions 5 C 2 O 4 2−  10 CO e − 10 e − + 16 H MnO 4 −  2 Mn H 2 O When we add these together,

Chapter 20: Electrochemistry Chemistry 1062: Principles of Chemistry II Andy Aspaas, Instructor.

Oxidation & Reduction Electrochemistry BLB 10 th Chapters 4, 20.

Chapter 17 Electrochemistry

For a half-reaction, the more (+) the E o red value, the greater the tendency for that reaction to “go” in that direction (i.e., reduction). Strongest.

Inorganic chemistry Assistance Lecturer Amjad Ahmed Jumaa  Predicting whether a (redox) reaction is spontaneous.  Calculating (ΔG°)

CHM 103 Lesson Plan 6/1/2004. Electrochemistry Introduction Voltaic Cells Standard Voltage: E° Relations Between K, ΔG°, & E° Effect of Concentration.

REDOX AND ELECTROCHEMISTRY SUROVIEC SPRING 2015 Chapter 18.

Chapter 20 Electrochemistry. Oxidation States electron bookkeeping * NOT really the charge on the species but a way of describing chemical behavior. Oxidation:

Redox Reactions and Electrochemistry Chapter 19. Cell Potentials E cell  = E red  (cathode) − E red  (anode) = V − (−0.76 V) = V.

Electrochemistry Part Four. CHEMICAL CHANGE  ELECTRIC CURRENT To obtain a useful current, we separate the oxidizing and reducing agents so that electron.

ELECTROCHEMISTRY Electrochemistry relates electricity and chemical reactions. It involves oxidation-reduction reactions (aka – redox) They are identified.

Chapter 19: Electrochemistry: Voltaic Cells Generate Electricity which can do electrical work. Voltaic or galvanic cells are devices in which electron.

Electrochemistry. #13 Electrochemistry and the Nernst Equation Goals: To determine reduction potentials of metals To measure the effect of concentration.

ELECTROCHEMISTRY CHEM171 – Lecture Series Four : 2012/01  Redox reactions  Electrochemical cells  Cell potential  Nernst equation  Relationship between.

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 20 Problem Set: p , 11, 13, 17, 23, 27, 31, 35, 43, 45, 47, 53, 61, 63, 69, 72, 75, 79, 85, 87, 95.

Electrochemistry. To obtain a useful current, we separate the oxidation and reduction half-reactions so that electron transfer occurs thru an external.

John E. McMurry Robert C. Fay C H E M I S T R Y Sixth Edition Chapter 17 Electrochemistry © 2012 Pearson Education, Inc.

Free Energy ∆G & Nernst Equation [ ]. Cell Potentials (emf) Zn  Zn e volts Cu e-  Cu volts Cu +2 + Zn  Cu + Zn +2.

Free Energy and Redox Reactions

Free Energy and Redox Reactions

Free Energy and Redox Reactions

Electrochemical cells

Chapter 19 Electrochemistry Semester 1/2009 Ref: 19.2 Galvanic Cells

Catalyst Calculate the standard reduction potential for the reaction:

Electrochemistry.

Presentation transcript:

Cell EMF Eocell = Eored(cathode) - Eored(anode) Example: Zn + Cu+2  Zn+2 + Cu E0cell = 0.34 V - (-0.76 V) = + 1.10 V

A voltaic cell is based on two half-reactions: Cd+2/Cd Sn+2/Sn Which half-reaction takes place at the cathode? Which half-reaction takes place at the anode? What is the standard cell potential?

A voltaic cell is based on two half-reactions: Cd+2/Cd Sn+2/Sn Which half-reaction takes place at the cathode? Which half-reaction takes place at the anode? What is the standard cell potential? Sn+2 + 2e-  Sn Cd  Cd+2 + 2e- 0.267 V

Spontaneity of Redox Reactions Eo = Eored(reduction) - Eored(oxidation) E0 = (+) spontaneous E0 = (-) nonspontaneous

Calculate the value of E0. Cu + 2H+  Cu+2 + H2 Calculate the value of E0.

Calculate the value of E0. Cu + 2H+  Cu+2 + H2 Calculate the value of E0. E0 = -0.34 V NOT SPONTANEOUS

EMF and Free-Energy Change G = -nFE n = a positive # (the # of electrons transferred) F = Faraday’s constant. 1F = 96,500 J/V-mol E = EMF

Use the standard reduction potentials to calculate the standard free-energy change, Go, for the following reaction: 4Ag + O2 + 4H+  4Ag+ + 2H2O

Go = -(4)(96,500J/V-mol)(+0.43V) Use the standard reduction potentials to calculate the standard free-energy change, Go, for the following reaction: 4Ag + O2 + 4H+  4Ag+ + 2H2O Go = -(4)(96,500J/V-mol)(+0.43V) - 170 kJ/mol

The Nernst Equation (at 298 K)

Cr2O7-2(aq) + 14H+(aq) + 6I-(aq)  2Cr+3(aq) + 3I2(s) + 7H20(l) Calculate the emf at 298K generated by the following cell, when (Cr2O7-2) = 2.0 M, (H+) = 1.0 M, (I-) = 1.0 M, and (Cr+3) = 1.0 x 10-5 M: Cr2O7-2(aq) + 14H+(aq) + 6I-(aq)  2Cr+3(aq) + 3I2(s) + 7H20(l)

Calculate the emf at 298K generated by the following cell, when (Cr2O7-2) = 2.0 M, (H+) = 1.0 M, (I-) = 1.0 M, and (Cr+3) = 1.0 x 10-5 M: Cr2O7-2 + 14H+ + 6I-  2Cr+3 + 3I2 + 7H20 E = 0.79 V - 0.0592V/6 log(5.0 x 10-11) E = 0.89 V