ELECTROCHEMISTRY Oxidation-reduction or REDOX reactions result in the To play the movies and simulations included, view the presentation in Slide Show Mode. Oxidation-reduction or REDOX reactions result in the generation of an electric current (electricity) or recharging of a battery.

Fireflies Glow to attract mates Anglerfish Emit light to attract prey Squid, jellyfish, bacteria, and shrimp Luminous How? Redox reactions Transfer of electrons in a redox reaction provides energy

Electrochemical Processes Chemical reactions either release or absorb energy Energy very often in the form of heat, light and sound but sometimes it is in the form of electricity Electron transfer reactions or redox reactions result in the generation of an electric current spontaneously or are caused by imposing an electric current nonspontaneously

Spontaneous Redox Reactions How do we know if a reaction is spontaneous? Use the Activity series chart (Reference TABLE J) Any metal or nonmetal in the series will displace a metal BELOW it from a compound or solution of that compound. Consequently, Metals higher are easier to react and therefore LOSE their electrons (thus OXIDIZED) and the metal that is displaced from compound is REDUCED Displacement of a metal or nonmetal by a metal/ nonmetal higher in the series corresponds to a SPONTANEOUS reaction For example: Would zinc spontaneously displace copper from a copper II sulfate (CuSO4) solution? Yes, Zinc is higher on the list than copper and would SPONTANEOUSLY displace copper Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s) Zinc is oxidized while the copper is reduced (a redox reaction)

Which are spontaneous & if spontaneous then Identify Reaction Products ? Yes Li + AlCl3  Cs + CuCl2  I2 + NaCl  Cl2 + KBr  Fe + CaBr2  Mg + Sr(NO3)2  F2 + MgCl2  Al + LiCl Yes Cu + CsCl NonSpontaneous Yes Br + KCl NonSpontaneous NonSpontaneous Yes Cl2 + MgF2

Predicting Spontaneous Redox Reactions In the 2 beakers a strip of Cu was placed in a solution AgNO3. Cu + AgNO3  Is the reaction Spontaneous? Yes, Ag is below Cu, therefore a Spontaneous reaction would occur Cu + AgNO3 Cu(NO3)2 + Ag (Note: the fuzzy material is the silver) So Who cares!!!!!!! Why is this useful? Well, this reaction is a redox reaction 0 +1 +2 0 Cu + AgNO3 Cu(NO3)2 + Ag Since copper loses electrons to the silver, if we separate the copper and silver from each other by a wire, we can create BATTERIES Oxidation Cu ---> Cu2+ + 2e- Reduction Ag+ + e- ---> Ag

Electrochemical Cell: Batteries A battery is a voltaic cell where a chemical or redox reaction allows flow of electrons (electricity) SPONTANEOUSLY through an external wire Voltaic (AKA Electrochemical) Chemical energy changed into electrical energy SPONTANEOUS – happens naturally or on its own A chemical reaction PRODUCES an electric current

Oxidation Cu ---> Cu2+ + 2e- Reduction Ag+ + e- ---> Ag --- Ag A strip of copper is in a copper II nitrate solution on one side and silver strip is in a silver nitrate solution on the other side with the silver and copper connected by an external wire. According to our equation before, Cu(s) + AgNO3(aq)  Cu(NO3)2(aq) + Ag(s) Remember Copper is oxidized (loses e-) while the silver is reduced (gains e-) Cu + Ag+ ---> Cu2+ + Ag Ag Ag+ NO3- --- Oxidation Cu ---> Cu2+ + 2e- Reduction Ag+ + e- ---> Ag

Two Parts to a Voltaic or Battery Cell Oxidation on one side and Reduction at the other side - Consists of two containers (each half cells) – each side with a piece of metal (AKA an electrode) that is placed in aqueous solution of a salt of that metal. Wire or External Conductor – connects electrodes. Salt bridge (or U-tube) that connects the aqueous solutions. Ag Ag+ NO3- --- Wire: connects the electrodes, carries electrons (electric current) Salt bridge: Allows positive and negative ions to pass from one cell to another but prevents solutions from mixing completely (maintain neutrality)

An Ox Ate a Red Cat Anode – Oxidation The anode = location for the oxidation (loss of electrons) half-reaction. Cu ---> Cu2+ + 2e- Reduction – Cathode The cathode = location for the reduction (gain of electrons) half-reaction. Ag+ + e- ---> Ag Surface at which oxidation or reduction half-reaction occurs. Anode loses electrons (losses mass or gets smaller) Oxidation: Cu0  Cu2+ + 2e- while Cathode or place where e- gained (gains mass). Reduction: Ag+1 +1e-  Ag0 Ultimately a battery dies when anode is completely gone

Salt Bridge: Pathway for flow of positive & negative IONS between half-cells. Necessary to maintain electrical neutrality. Reaction will not proceed without salt bridge. (Positive ions toward electrode where are being gained and negative ions vice versa) The metals (Ag & Cu in this case) or ELECTRODES are called: CATHODE (+)(reduction) & ANODE (-) (oxidation) : Electrodes may be connected by a wire and symbol shown as

Electron Flow thru wire from Anode to Cathode Anode to Cathode http://www.youtube.com/user/MarkRosengarten#p/c/65159266CFC74682/26/-bxJXt_69yM (3:29 min) http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/electroChem/volticCell.html e- e- e- Anode Cathode

Electrolytic cells (rechargeable) or Electrolysis Use electrical energy to force a chemical reaction to occur (Nonspontaneous). Electrolytic cells can be identified by the a battery or a power supply. Apply an external voltage (using a battery) and Reverse the direction of the electrons from a Redox rxn. Recogonize by seeing that this usually occurs in one container without a salt bridge

Electrolysis - Reversing a Voltaic Cell Voltmeter is here, but a power source (battery) must be here to recharge (reverse) the reaction Here’s a voltaic cell that must be reversed (recharged) when it no longer works. This time: Ag is oxidized & Cu is reduced. Ag AgNO3(aq)

Electrolytic cells - Electroplating How do we plate silver onto a copper spoon if copper is more likely to oxidize than silver? Electroplating (47 seconds): http://www.youtube.com/watch?v=D99qgd8XlV0 Mark Rosengarten Electrolysis (2:56min) http://www.youtube.com/user/MarkRosengarten#p/c/65159266CFC74682/27/-OVwt6Bx0eg Copper(Cu) Use a battery or power supply to force electrons to move in the reverse or nonspontaneous direction. Ag + AgNO3  Nonspontaneous Website showing Electrolysis: http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/electroChem/electrolysis10.html

Electroplating Example of Electrolysis Ag Example of Electrolysis One metal is plated onto another Used to protect the surface of the base metal from corrosion Or to make objects, such as tableware and jewelry more attractive An electric current is used to produce a chemical reaction The object to be plated is the cathode (negative) The metal used for plating is the anode (positive) Solution contains ions of element to be plated. AgNO3, or AgClO3, or AgCl etc. Oxidation: Ag ---> Ag+ + e- Silver ions (Ag+) move toward negative electrode (copper key) & effectively stick or plate to the outside of the key.

Electroplating: Ag onto Fe Fe is above Ag in Table J. This reaction will not happen spontaneously like in a volatic cell Thus, use an external energy source – a power supply or a battery – to force reaction to occur. Symbols to look for:

An Ox ate a Red Cat Anode is still electrode at which oxidation occurs. Cathode is still electrode at which reduction occurs. Polarity of electrodes is DIFFERENT between Voltaic Cells & Electrolytic Cells! Anode is positive. Cathode is negative. Look at drawings in Regents questions.

A POX on Electrolytic Cells Anode – Positive – Oxidation In an electrolytic cell, the polarity is determined by the outside power supply. The anode is hooked up to the positive terminal and the cathode is hooked up to the negative terminal. Look at drawings in Regents questions. ANODE CATHODE

Na+ Cl- Electrolysis of Molten Sodium Chloride Electrolysis of sodium chloride is the only process for the production of sodium metal What chemical species would be present in a vessel of molten sodium chloride, NaCl (l)? Na+ Cl- Let’s examine the electrolytic cell for molten NaCl.

At the microscopic level Molten NaCl At the microscopic level - + battery e- NaCl (l) cations migrate toward (-) electrode anions migrate toward (+) electrode Na+ Cl- Na+ e- Cl- (-) (+) cathode anode Cl- Na+ 2Cl-  Cl2 + 2e- Na+ + e-  Na

Electrolysis of Molten NaCl Why does the NaCl have to be molten? Which electrode will the Na+ ions be attracted to? Which electrode will the Cl- ions be attracted to? Ions are mobile Negative electrode, in an electrolytic cell the cathode Positive electrode, in an electrolytic cell the anode

Electrolysis of Molten NaCl Cathode half-cell (-) REDUCTION Na+ + e-  Na Anode half-cell (+) OXIDATION 2Cl-  Cl2 + 2e- Overall reaction 2Na+ + 2Cl-  2Na + Cl2 X 2 Non-spontaneous reaction!

Electrochemistry: http://www. youtube. com/watch Electrochemistry: http://www.youtube.com/watch?v=aPnx1GHskAY (4:29 min) Cleaning sterling silver jewelery: http://www.youtube.com/watch?v=KVWClb3pQk0&feature=related (5:48=7 min)

Construct a Voltaic Cell with Al & Pb Use Table J to identify anode & cathode. Draw Cell, put in electrodes & solutions Use Al(NO3)3 and Pb(NO3)2 for aqueous solutions Label anode, cathode, direction of electron flow in wire, direction of negative ion flow in salt bridge, positive electrode, negative electrode. Negative electrode is where electrons originate. Positive electrode attracts electrons. Write half reactions for the Al electrode and Pb electrode

 - Electron flow  wire Pb = Positive ion flow  Al = cathode anode Salt bridge  - Pb+2 & NO3-1 Al+3 & NO3-1

What are the half-reactions? Al  Al+3 + 3e- Pb+2 + 2e-  Pb Al metal is the electrode – it is dissolving. Al+3 ions go into the solution. Pb+2 ions are in the solution. They pick up 2 electrons at the surface of the Pb electrode & plate out.

Overall Reaction 2(Al  Al+3 + 3e-) 3(Pb+2 + 2e-  Pb) OXIDATION: Half Reaction REDUCTION: Half Reaction _____________________________ 2Al + 3Pb+2  2Al+3 + 3Pb

2Al + 3Pb+2  2Al+3 + 3Pb Al Pb Which electrode is losing mass? Which electrode is gaining mass? What’s happening to the [Al+3]? What’s happening to the [Pb+2]? Al Pb Increasing Decreasing