Powerpoint in electrochemistry
description
Transcript of Powerpoint in electrochemistry
1
Dr. Paul CharlesworthMichigan Technological UniversityDr. Paul Charlesworth
Michigan Technological University
C h a pte rC h a pte r 1818ElectrochemistryElectrochemistry
Chemistry 4th EditionMcMurry/Fay
Chemistry 4th EditionMcMurry/Fay
Chapter 18 Slide 2Prentice Hall ©2004
Kinds of electrochemical cellsKinds of electrochemical cells
• Galvanic (voltaic)Spontaneous chemical reaction generates electricity
with which to do worke.g. batteries
• ElectrolyticElectricity used to drive nonspontaneous chemical
reaction
Chapter 18 Slide 3Prentice Hall ©2004
Redox Reactions 01Redox Reactions 01
2
Chapter 18 Slide 4Prentice Hall ©2004
Redox reaction are those involving the oxidation and reduction of species.
OIL – Oxidation Is Loss of electrons.
RIG – Reduction Is Gain of electrons.
Oxidation and reduction must occur together.
They cannot exist alone.
Redox Reactions 01Redox Reactions 01
Chapter 18 Slide 5Prentice Hall ©2004
Redox Reactions 02Redox Reactions 02
Chapter 18 Slide 6Prentice Hall ©2004
• Oxidation Half-Reaction: Zn(s) → Zn2+(aq) + 2 e–.• The Zn loses two electrons to form Zn2+.
Redox Reactions 02Redox Reactions 02
3
Chapter 18 Slide 7Prentice Hall ©2004
Redox Reactions 03Redox Reactions 03
• Reduction Half-Reaction: Cu2+(aq) + 2 e– → Cu(s)• The Cu2+ gains two electrons to form copper.
Chapter 18 Slide 8Prentice Hall ©2004
• Overall: Zn(s) + Cu2+(aq)→ Zn2+(aq) + Cu(s)
Galvanic cell, example 1Galvanic cell, example 1
Chapter 18 Slide 9Prentice Hall ©2004
Electrochemical Cells 01Electrochemical Cells 01
• Electrodes: are usually metal strips/wires connected by an electrically conducting wire.
• Salt Bridge: is a U-shaped tube that contains a gel permeated with a solution of an inert electrolyte.
• Anode: is the electrode where oxidation takes place.
• Cathode: is the electrode where reduction takes place.
4
Chapter 18 Slide 10Prentice Hall ©2004
Galvanic cell, example 2Galvanic cell, example 2
Al(s) + Pb2+(aq) Al3+(aq) + Pb(s)
Chapter 18 Slide 11Prentice Hall ©2004
Galvanic cell, example 3Galvanic cell, example 3
Fe(s) + 2Fe3+(aq) 3Fe2+(aq)
Chapter 18 Slide 12Prentice Hall ©2004
Fe(s) + Fe3+(aq) 2Fe2+(aq)
5
Chapter 18 Slide 13Prentice Hall ©2004
Electrochemical Cells 02Electrochemical Cells 02
• Convention for expressing the cell:
Anode Half-Cell || Cathode Half-Cell
Electrode | Anode Soln || Cathode Soln | Electrode
Zn(s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu(s)
Fe(s) | Fe2+(aq) || Fe3+(aq), Fe2+(aq) | Pt(s)
• Electrons flow from anode to cathode. Anode is placed on left by convention.
Chapter 18 Slide 14Prentice Hall ©2004
Cell potential, Free-energy change
Chapter 18 Slide 15Prentice Hall ©2004
Driving force in electrochemical cells
Name: electromotive force (emf)cell potentialcell voltage
Symbol: E
SI units: Volt (V)
(1 J = 1 C x 1 V)
6
Chapter 18 Slide 16Prentice Hall ©2004
Relationship between ∆G and ERelationship between ∆G and E
∆G = -nFE
F = 96,500 C/mol e-
faraday (Faraday contant)
Chapter 18 Slide 17Prentice Hall ©2004
Standard cell potential, EoStandard cell potential, Eo
Standard conditionsaqueous solutes: 1 Mgases: 1 atmsolids and liquids: pure
∆Go = -nFEo
Chapter 18 Slide 18Prentice Hall ©2004
Example Problem
The standard cell potential (Eo) for following reaction is 1.10 V at 2.5oC. What is ∆Go?
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
7
Chapter 18 Slide 19Prentice Hall ©2004
Standard oxidation and reduction potentials
Eo = Eoox + Eo
red
Zn(s) Zn2+(aq) + 2e– Eoox = 0.76 V
2H+(aq) + 2e– H2(g) Eored = 0 V
Zn(s) + 2H+(aq) Zn2+(aq) + H2(g) Eo = 0.76 V
Example:
Chapter 18 Slide 20Prentice Hall ©2004
Electrochemical Cells 05Electrochemical Cells 05
Chapter 18 Slide 21Prentice Hall ©2004
Electrochemical Cells 04Electrochemical Cells 04
• The standard half-cell potentials are determined from the difference between two electrodes.
• The reference point is called the standard hydrogen electrode (S.H.E.) and consists of a platinum electrode in contact with H2 gas (1 atm) and aqueous H+ ions (1 M).
• The standard hydrogen electrode is assigned an arbitrary value of exactly 0.00 V.
8
Chapter 18 Slide 22Prentice Hall ©2004
H2(g) 2H+(aq) + 2e– Eoox =
Cu2+(aq) + 2e– Cu(s) Eored =
H2(g) + Cu2+( aq) 2H+(aq) + Cu(s) Eo = 0.34 V
Another example
0.34 V
0 V
Chapter 18 Slide 23Prentice Hall ©2004
Standard reduction potentialsStandard reduction potentials
Zn(s) Zn2+(aq) + 2e- Eoox = 0.76 V
Zn2+(aq) + 2e- Zn(s) Eored = -0.76 V
Eoox = -Eo
red
Chapter 18 Slide 24Prentice Hall ©2004
Electrochemical Cells 06Electrochemical Cells 06
9
Chapter 18 Slide 25Prentice Hall ©2004
Learning Goal
Determine Eo for a cell from a table of standard reduction potentials.
Chapter 18 Slide 26Prentice Hall ©2004
Electrochemical Cells 07Electrochemical Cells 07
• When selecting two half-cell reactions the more negative value will form the oxidation half-cell.
• Consider the reaction between zinc and silver:Ag+(aq) + e– → Ag(s) E° = 0.80 VZn2+(aq) + 2 e– → Zn(s) E° = – 0.76 V
• Therefore, zinc forms the oxidation half-cell:Zn(s) → Zn2+(aq) + 2 e– E° = – (–0.76 V)
Chapter 18 Slide 27Prentice Hall ©2004
Electrochemical Cells 08Electrochemical Cells 08
• What is the standard emf of an electrochemical cell made of a Cd electrode in a 1.0 M Cd(NO3)2solution and a Cr electrode in a 1.0 M Cr(NO3)2solution?
10
Chapter 18 Slide 28Prentice Hall ©2004
Electrochemical Cells 08Electrochemical Cells 08
• What is the standard emf of an electrochemical cell made of a Mg electrode in a 1.0 M Mg(NO3)2solution and a Ag electrode in a 1.0 M AgNO3solution?
Chapter 18 Slide 29Prentice Hall ©2004
What is Eo for the following cell?
Fe(s) | Fe2+(aq) || Pb2+(aq) | Pb(s)
Chapter 18 Slide 30Prentice Hall ©2004
Spontaneity of a Reaction 01Spontaneity of a Reaction 01
• The value of E˚ is related to the thermodynamic quantities of ∆G˚ and K.
∆G˚ = –nFE˚ and ∆G˚ = –RT ln K
combine to give:
E˚ = –(RT/nF) ln Kn = moles of electrons
RT/F = 0.0257 V (at 25oC)
11
Chapter 18 Slide 31Prentice Hall ©2004
Spontaneity of a Reaction 02Spontaneity of a Reaction 02
Equivalent forms of equation:
E˚cell = (0.0257/n) ln K
E˚cell = (0.0592/n) log K
Chapter 18 Slide 32Prentice Hall ©2004
Learning Goal
Determine equilibrium constant for a redox reaction from Eo.
Chapter 18 Slide 33Prentice Hall ©2004
Spontaneity of a Reaction 03Spontaneity of a Reaction 03
• Calculate the standard free energy change (∆G˚) and the equilibrium constant (K) for the following reactions at 25°C:
1. Sn(s) + 2 Cu2+(aq) Sn2+(aq) + 2 Cu+(aq)
2. Fe2+(aq) + 2 Ag(s) Fe(s) + 2 Ag+(aq)
3. 4 Fe2+(aq) + O2(g) + 4 H+(aq) 4 Fe3+(aq) + 2 H2O(l)
12
Chapter 18 Slide 34Prentice Hall ©2004
• Cell potentials can be modified by temperature and composition changes according to the equation:
∆G = ∆G° + RT lnQ
∴ –nFE = –nFE° + RT lnQ
• Giving:
The Nernst Equation 01The Nernst Equation 01
QnFRTEE o ln−=
Chapter 18 Slide 35Prentice Hall ©2004
Learning Goal
Determine a cell potential under non-standard-state conditions using Nernstequation (given: table of standard reduction potentials).
Chapter 18 Slide 36Prentice Hall ©2004
• Consider the reaction of metallic zinc with hydrochloric acid. Calculate the cell potential at 25°C when [H+] = 1.0 M, [Zn2+] = 0.0010 M, and PH2
= 0.10 atm.
The Nernst Equation 03The Nernst Equation 03
13
Chapter 18 Slide 37Prentice Hall ©2004
What is E for the following cell?
Zn(s) | Zn2+(2.0 M) || Cu2+(0.010 M) | Cu(s)
Chapter 18 Slide 38Prentice Hall ©2004
Practical Use: pH Meter
What is pH of the solution in the anode half-cell below?Given: E is measured to be 0.78 V, and
Eo[Cu+(aq) Cu(s)] is 0.52 V.
Pt(s) | H2(g,1 atm) | H+(aq,? M) || Cu+(aq,1 M) | Cu(s)
Reference electrodeSolution with unknown pH
Chapter 18 Slide 39Prentice Hall ©2004
Actual pH meter
Ag(s) | AgCl(s) | Cl–(aq)
One electrode (“glass electrode”):
Reference electrode (“calomel electrode”):
Hg2Cl2(s) | Hg(l), Cl–(aq)
Ag(s) + Cl–(aq) AgCl(s) + e– Eoox = -0.22 V
Hg2Cl2(s) + 2e– 2Hg(s) + 2Cl–(aq) Eored = 0.28 V
14
Chapter 18 Slide 40Prentice Hall ©2004
E depends on pH via potential across glass membrane, as well as on Eo
ox and Eored.