Lecture_overheads_-_Ch13

Lecture_overheads_-_Ch13 - CHEM 321 Quantitative Analysis...

Info iconThis preview shows pages 1–6. Sign up to view the full content.

View Full Document Right Arrow Icon

Info iconThis preview has intentionally blurred sections. Sign up to view the full version.

View Full DocumentRight Arrow Icon

Info iconThis preview has intentionally blurred sections. Sign up to view the full version.

View Full DocumentRight Arrow Icon

Info iconThis preview has intentionally blurred sections. Sign up to view the full version.

View Full DocumentRight Arrow Icon
This is the end of the preview. Sign up to access the rest of the document.

Unformatted text preview: CHEM 321 Quantitative Analysis Ch. 13 - Fundamentals of Electrochemistry 13-1 Basic Concepts Oxidation #s, Balancing Redox Reactions q, F, Work, G, Ohms Law, Power Law 13-2 Galvanic Cells Conventions, Reactions, E, G 13-3 Standard Potentials Standard Hydrogen Electrode Standard Reduction Potentials Tables of Standard Reduction Potentials 13-4 Nernst Equation Calculating Cell Potentials Examples 13-5 E and K eq See Text Appendix D on oxidation #s and balancing redox reactions See Text Appendix H for standard reduction potentials 13-1 Basic Concepts Redox reaction transfer of electrons from one species to another (involves both oxidation and reduction) Oxidation loss of electrons Oxidant oxidizes something (oxidizing agent) and in the process is reduced Reduction gain of electrons Reductant reduces something (reducing agent) and in the process is oxidized Examples Fe 3+ + V 2+ Fe 2+ + V 3+ Cu (s) + 4H + + 2NO 3- 2NO 2(g) + Cu 2+ +2H 2 O Fe 3+ + e- Fe 2+ NO 3- + 2H + + e- NO 2 + H 2 O V 2+ V 3+ + e- Cu Cu 2+ +2e- Fe 3+ is reduced (oxidizing agent) NO 3- is reduced (oxidizing agent) V 2+ is oxidized (reducing agent) Cu (s) is oxidized (reducing agent) Review of formal oxidation #s A convention used to give number of e- associated with a particular element 1. Oxidation # of an element when bonded to itself or in its metallic form = Examples: Cu metal, Na, Cl 2 , H 2 , C (diamond, graphite) 2. Oxidation # of H = +1 Exception: Hydrides, e.g. NaBH 4 , NaH, H charge is-1 3. Oxidation # of O = -2 Exception: Peroxides, e.g., HOOH, O charge is-1 4. Oxidation # of alkali metals (Li, Na, K) = +1 and alkaline earth metals (Be, Mg, Ca) = +2 Examples: LiCl, K 2 SO 4 5. Oxidation # of halogens (F, Cl, Br, ...) = -1 Examples: LiCl, CH 3 Cl Review on how to balance redox reactions General guidelines for balancing redox reactions : (involving both reduction and oxidation): 1. Assign oxidation numbers to elements that are oxidized or reduced 2. Break the overall reaction into two half reactions (one involving oxidation and the other reduction) and now for each half reaction: a. Balance the number of atoms of the element that is oxidized or reduced b. Account for the change in oxidation number by adding electrons to one side of the reaction c. Balance the number of O atoms by adding H 2 O to one side of the reaction d. Balance the number of H atoms by adding H + to other side of the reaction 3. Multiply each half reaction by the lowest common denominator so that electrons will cancel out in the overall reaction. The number of atoms of each element and the net charge on both sides of the reaction must be the same. Example of balancing redox reaction Balance the following reaction : MnO 4- + NO 2- Mn 2+ + NO 3- MnO 4- is reduced to Mn 2+ : NO 2- is oxidized to NO 3- : Mn oxidation state from +7 to +2 N oxidation state from +3 to +5 must gain electrons to do this must lose electrons to do this Balance half reaction:...
View Full Document

Page1 / 26

Lecture_overheads_-_Ch13 - CHEM 321 Quantitative Analysis...

This preview shows document pages 1 - 6. Sign up to view the full document.

View Full Document Right Arrow Icon
Ask a homework question - tutors are online