1 / 29

Understanding Redox Reactions: Key Concepts and Examples

Learn the fundamentals of oxidation and reduction, including oxidation numbers, identifying agents in equations, electrochemical cells, and more. Explore how to balance equations for redox reactions with special techniques.

jwaldrop
Download Presentation

Understanding Redox Reactions: Key Concepts and Examples

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Ch. 19Oxidation and Reduction REDOX

  2. 3 ways of looking at oxidation and reduction • 1. oxidation is a gain of oxygen atoms, reduction is a loss of oxygen atoms • 2. oxidation is a loss of hydrogen atoms, reduction is a gain of hydrogen atoms • 3. oxidation is a loss of electrons, reduction is a gain of electrons • Most fundamental explanation, what we will be dealing with the most

  3. Oxygen Hydrogen e-

  4. Leo the Lion! • LEO the lion says GER • Loss of Electrons is Oxidation • Gain of Electrons is Reduction • OIL RIG • Oxidation Is the Lost of electrons • Reduction Is the Gain of electrons

  5. Examples • Is the reactant oxidized or reduced? • Pb  PbO3 • SnO2  SnO • KClO3 KCl • C2H6O  C2H4O • C2H2  C2H6

  6. Pertaining to LEO… • Mg + S  MgS • Mg + S  Mg2+ + S2- • Magnesium is oxidized • Said to be the reducing agent • Substance in the reaction that loses electrons • Sulfide sulfur atom is reduced • Said to be the oxidizing agent • Substance in the reaction that gains electrons

  7. Oxidation Numbers • A count of the electrons transferred or shared in the formation or breaking of chemical bonds • You must assign each element in the reaction an oxidation number • Follow a set of rules…

  8. Oxidation Number Rules 1. The total of the oxidation numbers of all the atoms in a neutral molecule, an isolated atom, or a formula unit is 0 2. In their compounds, the Group 1A metals all have an oxidation number of +1, and the Group 2A metals have an oxidation number of 2+

  9. Rules Con’t • In its compounds, hydrogen has an oxidation number of +1 (except in metal hydrides such as NaH, where it is -1) 4. In its compounds, oxygen has an oxidation number of -2 (except in peroxides such as H2O2, where it is -1) 5. In their binary compounds with metals, Group7A elements have an oxidation number of -1. Group 6A elements have an oxidation number of -2, and Groups 5A elements have an oxidation number of -3.

  10. Problems • What is the oxidation number of each element? • I2 • Cr2O3 • AlCl3 • Na2SO4 • CaH2

  11. Identifying Redox Reactions 0 +3 -2 0 +3 -2 • 2 Al + Fe2O3 2 Fe + Al2O3 • Al increases from 0 to +3, it is ______ • Oxidized! • Fe decreases from +3 to 0, it is _______ • Reduced!

  12. Oxidizing and Reducing Agents • Now the confusing part… • CuO + H2 Cu + H2O • Cu goes from +2 to 0 • Cu is reduced, therefore it is called an oxidizing agent because it causes some other substance to be oxidized • H goes from 0 to +1 • H is oxidized, therefore it is called a reducing agent because it causes some other substance to be reduced.

  13. Identifying Agents in an Equation • CuO + H2 Cu + H2O Reduction: CuO is the oxidizing agent Oxidation: H2 is the reducing agent

  14. Electrochemical Cells • An apparatus that allows a redox reaction to occur by transferring electrons through an external connector. • Product favored reaction ---> voltaic or galvanic cell ----> electric current • Reactant favored reaction ---> electrolytic cell ---> electric current used to cause chemical change. Batteries are voltaic cells

  15. Basic Concepts of Electrochemical Cells Anode Cathode

  16. CHEMICAL CHANGE --->ELECTRIC CURRENT • Zn is oxidized and is the reducing agent Zn(s) ---> Zn2+(aq) + 2e- • Cu2+ is reduced and is the oxidizing agentCu2+(aq) + 2e- ---> Cu(s) With time, Cu plates out onto Zn metal strip, and Zn strip “disappears.”

  17. GOOD AFTERNOON!  PRE-AP: Please turn in your redox problems to the bin. Grab the sheets from the corner. Homework DUE: FRIDAY REGULAR Please grab a worksheet from the corner Take out the homework from last night – on desk REMINDERS: Project Description and Outline DUE MAY 2nd. Project DUE May 23rd (or by May 9th for +5 pts)

  18. Zn --> Zn2+ + 2e- Cu2+ + 2e- --> Cu Oxidation Anode Negative Reduction Cathode Positive •Electrons travel thru external wire. • Salt bridge allows anions and cations to move between electrode compartments. <--Anions Cations--> RED CAT

  19. Standard reduction Potential • E° = standard potential of the cell. • Represents the voltage of the cell when the electrons create a current by passing through the wire. • E°= E (reduction) + E (oxidation)

  20. Electrolysis of Water • Electrolysis of water is the breaking apart of water from H20 into its ions by running an electrical charge through it. • Example of the first use and type of a electrochemical cell

  21. CHEMICAL CHANGE --->ELECTRIC CURRENT • To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire. This is accomplished in a GALVANIC or VOLTAIC cell. A group of such cells is called a battery. http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf

  22. Terms Used for Voltaic Cells

  23. Balancing Equations for Redox Reactions Some redox reactions have equations that must be balanced by special techniques. MnO4- + 5 Fe2+ + 8 H+ ---> Mn2+ + 5 Fe3+ + 4 H2O Mn = +7 Fe = +2 Mn = +2 Fe = +3

  24. Balancing Equations Consider the reduction of Ag+ ions with copper metal. Cu + Ag+ --give--> Cu2+ + Ag

  25. Step 1: Divide the reaction into half-reactions, one for oxidation and the other for reduction. Ox Cu ---> Cu2+ Red Ag+ ---> Ag Step 2: Balance each element for mass. Already done in this case. Step 3: Balance each half-reaction for charge by adding electrons. Ox Cu ---> Cu2+ + 2e- Red Ag+ + e- ---> Ag

  26. Step 4: Multiply each half-reaction by a factor so that the reducing agent supplies as many electrons as the oxidizing agent requires. Reducing agent Cu ---> Cu2+ + 2e- Oxidizing agent 2 Ag+ + 2 e- ---> 2 Ag Step 5: Add half-reactions to give the overall equation. Cu + 2 Ag+ ---> Cu2+ + 2Ag The equation is now balanced for both charge and mass.

  27. Balancing Equations Balance the following in acid solution— VO2+ + Zn ---> VO2+ + Zn2+ Step 1: Write the half-reactions Ox Zn ---> Zn2+ Red VO2+ ---> VO2+ Step 2: Balance each half-reaction for mass. Ox Zn ---> Zn2+ Red 2 H++ VO2+ ---> VO2+ + H2O Add H2O on O-deficient side and add H+ on other side for H-balance.

  28. Balancing Equations Step 3: Balance half-reactions for charge. Ox Zn ---> Zn2+ + 2e- Red e- + 2 H+ + VO2+ ---> VO2+ + H2O Step 4: Multiply by an appropriate factor. Ox Zn ---> Zn2+ +2e- Red 2e-+ 4 H+ + 2 VO2+ ---> 2 VO2+ + 2 H2O Step 5: Add balanced half-reactions Zn + 4 H+ + 2 VO2+ ---> Zn2+ + 2 VO2+ + 2 H2O

  29. Tips on Balancing Equations • Never add O2, O atoms, or O2- to balance oxygen. • Never add H2 or H atoms to balance hydrogen. • Be sure to write the correct charges on all the ions. • Check your work at the end to make sure mass and charge are balanced. • PRACTICE!

More Related