Ch 17: Instrumental Methods in Electrochemistry Principle parts of a personal glucose monitor (covered in the section on Amperometry) The following chapter.

Slides:

Advertisements

Similar presentations

Electrochemical Energy Systems (Fuel Cells and Batteries)

Advertisements

Electrolysis & Understanding Electrolytic Cells : When a non-spontaneous redox reaction is made to occur by putting electrical energy into the system.

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

19.2 Galvanic Cells 19.3 Standard Reduction Potentials 19.4 Spontaneity of Redox Reactions 19.5 The Effect of Concentration on Emf 19.8 Electrolysis Chapter.

Chemistry 1011 Slot 51 Chemistry 1011 TOPIC Electrochemistry TEXT REFERENCE Masterton and Hurley Chapter 18.

Dry Cell Battery Anode (-) Zn ---> Zn e- Cathode (+) 2 NH e- ---> 2 NH 3 + H 2 Common dry cell Copyright © 1999 by Harcourt Brace & Company.

Electrochemistry Batteries. Batteries Lead-Acid Battery A 12 V car battery consists of 6 cathode/anode pairs each producing 2 V. Cathode: PbO 2 on a metal.

Oxidation and reduction reactions occur in many chemical systems. Examples include the rusting of iron, the action of bleach on stains, and the reactions.

Fundamentals of Electrochemistry

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.

Prentice Hall © 2003Chapter 20 For the SHE, we assign 2H + (aq, 1M) + 2e -  H 2 (g, 1 atm) E  red = 0.

Fundamentals of Electrochemistry Introduction 1.)Electrical Measurements of Chemical Processes  Redox Reaction involves transfer of electrons from one.

Electrochemistry Chapter 11 Web-site:

Representing electrochemical cells The electrochemical cell established by the following half cells: Zn(s) --> Zn 2+ (aq) + 2 e - Cu 2+ (aq) + 2 e - -->

Electrochemistry The first of the BIG FOUR. Introduction of Terms  Electrochemistry- using chemical changes to produce an electric current or using electric.

Unit 2 A Coulometry and Electrogravimetry

Chemistry. Session Electrochemistry - 2 Session Objectives Electrolysis Faradays Laws of electrolysis Electrode Potential Electromotive force Electrochemical.

Electrochemistry Chapter 19.

Chapter 21: Electrochemistry III Chemical Change and Electrical Work 21.6 Corrosion: A Case of Environmental Electrochemistry 21.7 Electrolytic Cells:

The End is in Site! Nernst and Electrolysis. Electrochemistry.

Instrumental Analysis Fundamentals of Electrochemistry

Other Electrochemical Methods

Coulometric Methods A.) Introduction:

Notes on Electrolytic Cells An electrolytic cell is a system of two inert (nonreactive) electrodes (C or Pt) and an electrolyte connected to a power supply.

Activity Series lithiumpotassiummagnesiumaluminumzincironnickelleadHYDROGENcoppersilverplatinumgold Oxidizes easily Reduces easily Less active More active.

Electrical and Chemical Energy Interconversion

Electrochemisty Electron Transfer Reaction Section 20.1.

Electrochemistry - The relationship between chemical processes and electricity oxidation – something loses electrons reduction – something gains electrons.

Chapter 8 Bulk Electrolysis: Electrogravimetry and Coulometry.

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

Unit 5: Everything You Wanted to Know About Electrochemical Cells, But Were Too Afraid to Ask By : Michael “Chuy el Chulo” Bilow And “H”Elliot Pinkus.

Electrochemistry Electrolysis Electrolytic Cells An electrolytic cell is an electrochemical cell that undergoes a redox reaction when electrical energy.

Electrochemistry - Section 1 Voltaic Cells

Electrolysis & Applications Since chemical oxidation-reduction involves the transfer of electrons from one substance to another, it should be.

Electrochemistry Mr. Weldon. 1. Definition: Field that deals with chemical changes caused by electric current and the production of electricity by chemical.

Electrolysis Chapter 17 Section 7 Electrochemistry e-

Electrochemistry Electrolysis.

REDOX Part 2 - Electrochemistry Text Ch. 9 and 10.

18.8 Electrolysis 2 Types of electrochemistry 1.Battery or Voltaic Cell – Purpose? 2.Electrolysis - forces a current through a cell to produce a chemical.

A.) Introduction : 1.) Coulometry: electrochemical method based on the quantitative oxidation or reduction of analyte - measure amount of analyte by measuring.

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

1 © 2006 Brooks/Cole - Thomson Electrolysis Using electrical energy to produce chemical change. Sn 2+ (aq) + 2 Cl - (aq) ---> Sn(s) + Cl 2 (g)

Chapter 19 Last Unit Electrochemistry: Voltaic Cells and Reduction Potentials.

Electrolytic Cells Section 9.2. Vocabulary Electrolysis: electrical energy used to bring about a non-spontaneous redox reaction Electrolyte: any substance.

CHE1102, Chapter 19 Learn, 1 Chapter 19 Electrochemistry Lecture Presentation.

Corrosion of Iron Since E  red (Fe 2+ ) < E  red (O 2 ) iron can be oxidized by oxygen. Cathode: O 2 (g) + 4H + (aq) + 4e -  2H 2 O(l). Anode: Fe(s)

2 Types of Electrochemical Cells 1)Voltaic Cells  Spontaneous reaction  Reaction itself creates electric current  Main concept for batteries 2)Electrolytic.

 Anything that uses batteries: › Cell phones › Game boys › Flash lights › Cars  Jewelry—electroplating.

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

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

Electrochemistry Chapter 20. oxidation: lose e- -increase oxidation number reduction: gain e- -reduces oxidation number LEO goes GER Oxidation-Reduction.

18.8 Electrolysis: Driving Non-Spontaneous Chemical Reactions with Electricity.

Stoichiometry of Cells Faraday’s Law. The mass deposited or eroded from an electrode depends on the quantity of electricity. Quantity of electricity –

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

Electrochemistry - The relationship between chemical processes and electricity oxidation – something loses electrons reduction – something gains electrons.

Bulk Electrolysis: Electrogravimetry and Coulometry

Electrolysis 3.7 Electrolysis…. Electrolysis Use of electrical energy to produce chemical change...forcing a current through a cell to produce a chemical.

Instrumental Analysis Electrogravimetry , Coulometry

Electroanalytical – Coulometry and Conductivity Ch 24, 7th e, WMDS)

Ch. 20: Electrochemistry Lecture 4: Electrolytic Cells & Faraday’s Law.

H.W. # 24 Study pp (sec ) Ans. ques. p. 883 # 93a,95c,97,104,105,

Notes on Electrolytic Cells

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

Chapter 15 Oxidation and Reduction

Chapter 20 Electrochemistry

An electrolytic cell uses electricity to do a chemical reaction.

Chapter 20 Electrochemistry

Electrogravimetry in electrogravimetry the product is deposited quantitatively on an electrode by an electrolytic reaction and the amount of the product.

from a battery or other external energy source

Presentation transcript:

Ch 17: Instrumental Methods in Electrochemistry Principle parts of a personal glucose monitor (covered in the section on Amperometry) The following chapter material is taken from your text and another textbook by Harris: Daniel C. Harris, "Quantitative Analysis, 7 th ed", Chapter 17, p You have been given a copy of the reading from this additional reference.

Electrolysis is the process in which electrical energy is used to cause a nonspontaneous chemical reaction to occur by "pumping" electrons into the cathode. - + Fundamentals of Electrolysis 2 Na Cl - → 2 Na(l) + Cl 2 (g) E o = is the minimum voltage that must be applied to the cathode to reduce Na +

Cathode: Cu e - = Cu(s) Anode: H 2 O = ½ O 2 (g) + 2 H + + 2e - Net: H 2 O + Cu 2+ = Cu(s) + ½ O 2 (g) + 2 H + Harris, 7 th ed q = I x t Conversion between moles e- and Coulombs is Faraday's constant: 1 mole e - 96,500 C Current (I) is measured in Coulombs second Coulombs = Coulombs x seconds

Electrolysis and Mass Changes Example (p. 350, Harris 7 th ed): If a current of 0.17 A flows for 16 min through the cell in Figure 17-1, how many grams of Cu(s) will be deposited? Harris, 7 th ed

Electrogravimetric Analysis Practical application of electrolysis: reduce a metal at the cathode, and weigh the mass of the cathode before and after the reaction. Cu 2+ (aq) + 2e - → Cu(s) deposited on cathode (Pt gauze) Harris, 4 th ed

Coulometry or Coulometric Titration analyte titrant generated at the anode from Br - Like electrolysis since both "pump" electrons into the cathode to cause reduction there. The solution initially contains Br -. At the anode (oxidation) the Br 2 is generated by oxidation of Br - 2 Br - → Br 2 + 2e - As soon as Br 2 is produced, it reacts with the C 6 H 10 When all the C 6 H 10 is consumed, the excess Br 2 generated at the anode signals the end of the titration (solution turns orange-brown and the detector electrodes voltage reaches a maximum) The moles of C 6 H 10 are calculated from the current and the time to the endpoint. Harris, 7 th ed

Br 2 generated here by oxidation of Br - C 6 H 10 reduced here to trans-1,2- dibromocyclohexane Harris, 7 th ed

Example (p. 356, Harris 7 th ed): A mL volume containing mg of cyclohexene/mL is to be titrated in Figure With a current of mA, how much time is required for complete titration?