Oxidation-Reduction Chemistry

Slides:



Advertisements
Similar presentations
1 Electrochemistry Chapter 18, Electrochemical processes are oxidation-reduction reactions in which: the energy released by a spontaneous reaction.
Advertisements

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 /28/07.
Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Electrochemistry The study of the interchange of chemical and electrical energy.
Predicting Spontaneous Reactions
ELECTROCHEMISTRY REDOX REVISITED! 24-Nov-97Electrochemistry (Ch. 21) & Phosphorus and Sulfur (ch 22)1.
Section 18.1 Electron Transfer Reactions 1.To learn about metal-nonmetal oxidation–reduction reactions 2.To learn to assign oxidation states Objectives.
Oxidation-Reduction Chemistry Redox. Definitions Oxidation: Reduction: Oxidizing Agent: Reducing Agent:
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.
Electrochemistry Chapter 19 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Redox Reactions and Electrochemistry
Redox Reactions and Electrochemistry
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.
Electrochemistry Chapter 19 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
CHEM 163 Chapter 21 Spring minute review What is a redox reaction? 2.
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 - ) ______________________.
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.
Redox Reactions & Electrochemistry Chapter 19 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
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.
Oxidation-Reduction Chemistry Redox. Definitions Oxidation: Reduction: Oxidizing Agent: Reducing Agent:
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.
1 Electrochemistry Chapter 18 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Electrochemistry Terminology  Oxidation  Oxidation – A process in which an element attains a more positive oxidation state Na(s)  Na + + e -  Reduction.
Electrolysis 3.7 Electrolysis…. Electrolysis Use of electrical energy to produce chemical change...forcing a current through a cell to produce a chemical.
NCEA Chemistry 3.7 Redox AS
Chapter 17: Electrochemistry
Electrochemistry.
Third Edition Lecture Power Points
Oxidation-Reduction Reactions
Engineering Chemistry CHM 406
17.1 Galvanic Cells (Batteries)
Electrochemistry Chapter 19.
Chemsheets AS006 (Electron arrangement)
Galvanic Cells Electrochemistry: The area of chemistry concerned with the interconversion of chemical and electrical energy Galvanic (Voltaic) Cell: A.
Electrochemistry Chapter 20.
Chapter 20 - Electrochemistry
Chapter 21: Electrochemistry
Electrochemistry Applications of Redox.
Redox Reactions and Electrochemistry
Batteries and Galvanic Cells
Electrochemical cells
Electrochemistry The batteries we use to power portable computers and other electronic devices all rely on redox reactions to generate an electric current.
Electrochemistry the study of the interchange of chemical and electrical energy.
Chemsheets AS006 (Electron arrangement)
Electrochemistry Chapter 19
Cell Potential and the Nernst Equation
Chapter 15 Oxidation and Reduction
Electrochemistry Applications of Redox.
Batteries and Galvanic Cells
Electrochemistry Chapter 19
Electrochemistry Chapter 18.
ELECTROCHEMISTRY Chapter 18
Electrochemistry Chapter 19

Chapter 20: Electrochemistry
Chapter 18 Electrochemistry Lesson 2
Galvanic Cell or Voltaic
Electrochemistry.
Electrochemistry Chapter 19
Chapter 21 Thanks to D Scoggin Cabrillo College
A. Oxidation-Reduction Reactions
Electrochemistry Chapter 19
Galvanic Cell or Voltaic
Presentation transcript:

Oxidation-Reduction Chemistry Redox

Definitions Oxidation: Reduction: Oxidizing Agent: Reducing Agent:

Examples: 2 Ag+(aq) + Cu(s)  2 Ag(s) + Cu2+(aq) MnO4-(aq) + 5 Fe2+(aq) + 8 H+(aq)  Mn2+(aq) + 5 Fe3+(aq) + 4 H2O(l)

Electrochemical Cells: Batteries Anode: Cathode:

Reduction potential: measure of how much the reactant wants to gain electrons to form product. Rank oxidizing and reducing agents.

Reduction Potentials and Cell Potentials Eocell = Eoanode – Eocathode more positive = more favored What is the standard cell potential for a cell using: Al  Al3+ and Hg2+  Hg

When is a cell reaction favored? Go = -nFEo F = 96485 J/mol When is a cell reaction favored? What is Go for a cell using Cu2+  Cu and Zn  Zn2+

Balancing in Basic Solution SiO32- + I-  Si + I2

Balancing in Basic Solution SiO32- + I-  Si + I2

Predicting favored reactions from the table Eo = Eo cathode - Eo anode Eo = Eo forward half reaction - Eo reversed half reaction

If Eo relates to Go, it must relate to K. Go = -RT lnK Go = -nFEo At 298 K, RT/F = 0.0257 V, so:

Cell potentials under nonstandard conditions Nernst Equation: When concentrations differ Recall Q Predict E vs. [products] or [reactants]

Concentration Cells Cu2+ + Cu  Cu + Cu2+ What is the potential of a cell that contains a Cu wire and [Cu2+] = 1.5 M in the cathode compartment and a Cu wire and [Cu2+] = 0.00058 M in the anode compartment? Cu2+ + Cu  Cu + Cu2+

Concentration Cells: pH meter Real use of concentration cells: measure E to determine concentration. 𝐸 cell = 𝐸 cell o − 0.059V 𝑛 log𝑄 If the pH meter electrode has a measured voltage of 1.43 V, what is the pH? H2(g) + 2 H+(aq, cathode)  2 H+(aq, anode) + H2(g)

Electrolysis: Coulometry–Counting electrons Use electrical energy to effect chemical change: Use an applied voltage to make a cell run “backwards.” One of only three methods to get products from thermodynamically unfavorable reactions.

Electrolysis: Coulometry– Counting electrons 1 C = 1 Amp x 1 sec 96485 C = 1 mol e- How many grams of Cu can be deposited from a Cu2+ solution by applying a current of 1.68 A for 22 minutes?

Electrolysis of Molten Salts NaCl melts at about 801 oC. Use a mixture of CaCl2 and NaCl is used. Melts at 580 oC.

Electrolysis of Molten Salts A current of 7.06 A is passed through an electrolysis cell containing molten CaCl2 for 12.9 minutes. a. Predict the products of the electrolysis and the reactions occurring at the cathode and anode.

Electrolysis of Molten Salts A current of 7.06 A is passed through an electrolysis cell containing molten CaCl2 for 12.9 minutes. b. Calculate the quantity or volume of products collected at the anode and cathode (assume gases are collected at 298 K and 1.00 atm).

Batteries

Alkaline Battery Cathode: 2 MnO2(s) + H2O(ℓ) + 2 e−  Mn2O3(s) + 2 OH−(aq) Anode: Zn(s) + 2 OH−(aq)  ZnO(s) + H2O(ℓ) + 2 e−   Net reaction: 2 MnO2(s) + Zn(s)  Mn2O3(s) + ZnO(s) Ecell = 1.5 V

Watch Battery Cathode: HgO(s) + H2O(ℓ) + 2e−  Hg(ℓ) + 2 OH−(aq) Anode: Zn(s) + 2 OH−(aq)  ZnO(s) + H2O(ℓ) + 2 e−   Net reaction: HgO(s) + Zn(s)  ZnO(s) + Hg(ℓ) Ecell = 1.3 V

Lead Acid Battery Cathode: PbO2(s) + 2 H2SO4(aq) + 2 e−  PbSO4(s) + 2 H2O(ℓ) Anode: Pb(s) + SO42-(aq)  PbSO4(s) + 2 e−   Net reaction: PbO2(s) + Pb(s) + H2SO4(aq)  2 PbSO4(s) + 2 H2O(ℓ) Recharging: 2 PbSO4(s) + 2 H2O(ℓ)  PbO2(s) + Pb(s) + H2SO4(aq)

Rechargeable Batteries Nickel Metal Hydride Batteries   Cathode: NiO(OH)(s) + H2O(ℓ) + e−  2 Ni(OH)2(s) + OH−(aq) Anode: MH(s) + OH−(aq)  M(s) + H2O(ℓ) + e− Net reaction: MH(s) + NiO(OH)(s)  M(s) + Ni(OH)2(s) Lithium Ion Batteries   Cathode: CoO2(s) + Li+ + 1 e−  LiCoO2(s) Anode: Li-graphite  C(graphite) + Li+ + e− Net reaction: Li(s) + CoO2(s)  LiCoO2(s)

Rechargeable Batteries What is required for a battery to be rechargeable?