1 / 22

Redox Reactions

Redox Reactions. Electrochemical Cells: The Voltaic Cell. Mr. Shields Regents Chemistry U14 L03. Half cell Nomenclature. We’re now going to discuss some practical applications of Redox reactions - Recall that Redox reactions involve electron transfers

wynona
Download Presentation

Redox Reactions

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. Redox Reactions Electrochemical Cells: The Voltaic Cell Mr. Shields Regents Chemistry U14 L03

  2. Half cell Nomenclature We’re now going to discuss some practical applications of Redox reactions - Recall that Redox reactions involve electron transfers - Participating atoms either provide or accept these e- Example: Zn0 Zn+2 + 2e- Oxidation half-cell Cu+2 + 2e-  Cu0 Reduction half-cell These reactions can also be written in “Shorthand” nomenclature As Follows: Zn0 |Zn+2 Cu+2 |Cu0 oxidation Reduction

  3. Electrochemical cells The electrical nature of a Redox reaction allows one the ability to construct several types of electrical cells An apparatus that uses Redox reactions to produce electrical Energy OR uses electrical energy to cause a chemical reaction is Known as an ELECTROCHEMICAL CELL

  4. Electrochemical cells There are two types of electrochemical cells: - VOLTAIC (also known as GALVANIC) CELLS - ELECTROLYTIC CELLS Devices that convert CHEMICAL ENERGY into ELECTRICAL ENERGY through a Spontaneous REDOX reaction are called VOLTAIC CELLS They represent EXOTHERMIC reactions. Note Well: NYS Regents Refers to only the VOLTAIC CELLS as an ELECTROCHEMICAL cell (even though this is not quite accurate, remember it!)

  5. Electrochemical cells In an ELECTROLYTIC CELL electrical energy is provided to Force a NON-SPONTANEOUS Redox reaction to happen. These Cells are ENDOTHERMIC reactions. Electrolytic cells are typically used to - plate metals on other metals - obtain pure metal from it’s compounds - to recharge batteries We’ll talk about GALVANIC CELLS first and later we’ll discuss ELECTROLYTIC CELLS

  6. Remember … We used Table J to determine if a Rxn will occur spontaneously. A Rxn will be spontaneous if the substance to be oxidized is above the substance to be reduced. Example 1: Will the reaction Cu+2 + K  Cu + K+ be spontaneous? Cu+2  Cu (reduction) and K  K+ (oxidation) Since K is oxidized & above Cu on the table the Rxn is spontaneous.

  7. Example 2: Will the reaction Li+ + Al  Li + Al+3 be spontaneous? Li+ Li (reduction) Al  Al+3 (oxidation) Since Al is oxidized but below Li on the table, the Rxn is non-spontaneous.

  8. Electrochemical Cells Most Voltaic cells share certain common features. They all have: - Two physically separated half cells - A Cathode (+) in one cell (Reduction occurs here) - An Anode (-) in the other cell (Ox. Occurs here) - Two different metal electrodes - A solution in each cell containing a dissolved salt made from the metal in the electrode - An electrical connection between Anode and Cathode - A physical connection between two separated cells containing ions that can move freely into each cell

  9. Memory Jogger for Galvanic cells For Redox reactions you’ve remembered the phrase “OIL RIG” For Voltaic cells there is another phrase to help you remember how Voltaic cells are constructed. “AN Ox ate a Red Pussy Cat” AnodeCathode OxidationReduction NegativePositive A- oxidation C+ Reduction Fe  Fe+2 + 2e- Cu+2 +2e-  Cu

  10. The basics of a Voltaic electrochemical cell Anode Cathode

  11. Electrochemical Cells Voltaic (galvanic) Cells generate usable electricity Let’s see how this happens. Consider the following reaction: Fe + CuCl2 FeCl2 +Cu This is a spontaneous Redox reaction. Fe is being oxidized & since it’s higher in Table J than Cu, Cu will be reduced. What are the 2 half cell reactions? Fe  Fe+2 +2e-oxidation half cell Cu+2 +2e-  Cu0 reduction half cell

  12. Electrochemical Cells In this redox Reaction the Transfer of e- Between Fe And Cu is Direct. In Voltaic cells electron Transfer is indirect. One way to do this is to separate the Fe and Cu half cells and connect the metals by a wire.

  13. Two separated Half cells + - Electrochemical cells If we do this, The reaction quickly Comes to a Screeching halt. Why? The build up of charge In both cells stops the Reaction from Continuing. ButWhy does Charge build up? Fe+2 goes into sol’n as Fe is oxidized releasing 2e- & leaving behind pos. ions The Cu+2 consumes the 2 electrons leaving behind an excess of Cl-. Rxn can no longer proceed.

  14. Salt Bridge • We need a way to stop the build up of charge in each half cell. • This is done using what is called aSALT BRIDGE • - A salt bridge allows ions to flow from cell to cell but • still maintains a separation • between them • It contains a salt in • solution. The Cation (+) is • be chosen such that it • won’t react with the metal • Electrodes (Na salts are • typical)

  15. Salt Bridge Ion diffusion Pos. Ion Build Up Neg. Ion Build Up A membraneon each side of the salt bridge keeps the NaNO3 Inside but allows ions to diffuse into the half cells to keep them Electrically neutral (i.e no build up of ions). In this case, for each Cu reduced 2 Na+ ions must migrate into The reduction half cell and at the same time2NO3- ions Must migrate into the Oxidation cell (why?)

  16. Changing Concentration of Half cells Notice that as the Cu+2 is reduced the concentration of Copper salt solution in that half cell decreases. On the other hand, as Fe is oxidized the concentration of Iron salt solution in that half cell increases. WHY?

  17. Direction of electron Flow In this voltaic cell e- transfer is from Fe to Cu. Therefore the Fe Electrode is negative (excess of e-) and Cu is + (needs electrons). Therefore electrons flow through the wire from Fe to Cu. But how do we know which metal is the One to loose e- ? Table J provides the Information needed. Metals higher on the List are more easily Oxidized. - +

  18. Changing Electrode Mass The electrode at which oxidation occurs loses mass Fe  Fe+2 +2e- The electrode at which reduction occurs gains mass Cu+2 +2e- Cu Since iron electrode is negative it is called the And since Cu elec. is positive it is the called the ANODE (-) - + CATHODE (+) Gains Mass Looses Mass

  19. Voltage depends Upon the metal Electrodes used. + - A standard Zn/Cu Electrochemical Cell produces about 1.1V A standard Li/Ag Voltaic Cell produces about 3.8V

  20. Equilibrium As an electrochemical cell is used the voltage continuously Decreases until the cell reaches equilibrium. When the oxidized electrode is completely used up the cell voltage is zero(0) and the cell has reached equilibrium since the concentrations are no longer changing. The cell is now said to be “dead”.

  21. PROBLEM: The following overall reaction occurs in a galvanic cell. 4AgNO3(aq) + Sn(s)  Sn(NO3)4(aq) + 4Ag(s) What’s oxidized? What’s Reduced? What are the balanced REDOX half cell reactions? 4Ag+ +4e-  4Ag Reduction Half cell Sn  Sn+4 + 4e- Oxidation Half Cell What metal electrode gains mass? What metal electrode loses mass? What electrode is negative, which is positive? Which metal is the cathode, Which is the anode? Which Sol’n Conc. Increases? Sn Ag+ Ag Sn Sn, Ag Ag, Sn Sn+4

  22. Problem: Draw and Fully label an Al & Zn Electrochemical cell. (Make sure you Indicate the direction of e- and ion flow; use nitrate solutions ). Lastly, write a balanced chemical equation that describes this cell NaNO3 2Al + 3Zn(NO3)2 2Al(NO3)3 + 3Zn Salt Bridge

More Related