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ELECTROCHEMISTRY Chapter 21 Section 21.1 Electrochemical Cells • Objectives – Interpret an activity series and identify the elements that are most easily oxidized and those that are least easily oxidized – Name the type of reactions involved in electrochemical processes – Describe how a voltaic cell produces electrical energy • 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 processes either release or absorb energy – Energy is sometimes 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 – The field of chemistry that deals with these two situations is called electrochemistry Spontaneous Redox Reactions How do we know this reaction between zinc and copper is spontaneous? – Table J – Activity Series • Zinc is higher on the list than copper –For any two metals on table J, the more active metal is more readily oxidized Which are spontaneous & if spontaneous then Identify Reaction Products ? • Li + AlCl3 • • • • • • Spontaneous (Li above Al) Cs + CuCl2 Spontaneous (Cs / Cu) I2 + NaCl Nonspontaneous (Cl / I) Spontaneous (Cl / Br) Cl2 + KBr Fe + CaBr2 Nonspontaneous (Ca / Fe) Nonspontaneous (Sr / Mg) Mg + Sr(NO3)2 F2 + MgCl2 Spontaneous (F / Cl) Electrochemistry • Many applications – Flashlights – Automobile batteries – Manufacture of sodium and aluminum metal – Silver plating of table ware and jewelry – Biological systems • Carrying impulses Electrochemical Cells • An electrochemical cell is any device that converts chemical energy into electrical energy or electrical energy into chemical energy – Redox reactions occur in all electrochemical cells • A battery functions by transferring electrons through an external wire from the reducing agent to the oxidizing agent. Overview of Electrochemistry Two kinds of electrochemical cells (kind of opposites): 1. Voltaic • Use a spontaneous reaction to produce a flow of electrons (electricity) exothermic. 2. Electrolytic • Use a flow of electrons (electricity) to force a nonspontaneous reaction to occur - endothermic. Batteries are voltaic cells Animation of Voltaic Cell http://www.chem.iastate.edu/group/Greenbo we/sections/projectfolder/flashfiles/electroCh em/volticCell.html Voltaic Cells • Spontaneous redox reactions • Converts chemical energy into electrical energy – Electrical energy is produced in a voltaic cell by spontaneous redox reactions within the cell • Parts of a Voltaic Cell – 2 half cells – 2 electrodes (an anode and a cathode) – Aqueous solutions – Wire – Salt bridge Voltaic Cell Anode Cathode Aqueous solution must contain ions of same metal as electrode: here ions = Zn2+ ions. Aqueous solution must contain ions of same metal as electrode: here ions = Cu2+ ions. Solution might be Zn(NO3)3(aq) or ZnSO4(aq) Solution might be Cu(NO3)3(aq) or CuSO4(aq) Voltaic Cell Wire: connects the electrodes, carries electrons (electric current) Anode Cathode Salt bridge: Allows ions to pass from one cell to another but prevents solutions from mixing completely Anode and Cathode • Electrodes (the metals) – Anode • Electrode at which the oxidation occurs – Electrons are produced at the anode and it is labeled the negative electrode – Cathode • Electrode at which reduction occurs – Electrons are consumed at the cathode as a result the cathode is labeled the positive electrode – Neither electrode is really charged • All parts remain neutral • Moving electrons and ions balance any charge that might build up What happens at the electrodes? Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s) Anode: Zinc metal slowly dissolves (looses mass) Oxidation: Zn0 Zn2+ + 2eCathode: Copper atoms are deposited as metallic copper on top of zinc (gaining mass) Reduction: Cu2+ +2e- Cu0 An Ox Ate a Red Cat • Anode – Oxidation – The anode = location for the oxidation half-reaction. • Reduction – Cathode – The cathode = location for the reduction half-reaction. How do we determine which electrode is the anode and which electrode is the cathode? • Remember… – Table J • The more active metal is oxidized and is therefore the anode • Anode = oxidation = electron donor – Higher metal on table J • Cathode = reduction = electron acceptor – Lower metal on table J Zn is above Cu on Table J Zn is the anode Cu is the cathode Zn2+ Electrons flow from zinc to copper (through wire) Positive ions (Zn2+)flow from zinc to copper (through salt bridge) Negative ions (SO42-)flow from copper to zinc (through salt bridge) e- e- e- e- Anode e- Cathode e- Complete Electrochemical Cell Zn --> Zn2+ + 2e- Cu2+ + 2e- --> Cu Oxidation Anode Negative AN OX Reduction Cathode Positive <-- Anions Cations --> RED CAT How a Voltaic Cell Works … • The electrochemical process that occurs in a Zn-Cu voltaic cell can best be described in steps, but the steps actually occur at the same time 1. Electrons are produced at zinc anode according the oxidation half reaction: Zn(s) Zn2+(aq) + 2e-. Zinc is the anode and the negative electrode because it is oxidized Electrons leave zinc anode and pass through the external circuit (wire) to the copper rod Electrons enter the copper rod and interact with copper ions in solution. The following reduction half reaction occurs Cu2+(aq) + 2e- Cu(s). Copper ions are reduced. The copper is the cathode and the positive electrode because it is reduced To complete the circuit, both the positive and negative ions move through the aqueous solution via the salt bridge 2. 3. 4. Construct a Voltaic Cell with Al & Pb 1. Use Table J to identify anode & cathode. 2. Draw Cell, put in electrodes & solutions using nitrate as the negative ion 3. Label: a. anode b. cathode c. direction of electron flow in wire d. direction of positive ion flow in salt bridge e. positive electrode f. negative electrode. 4. Write out half reactions 5. Write the overall reaction Electron flow wire Positive ion flow Al = anode Pb = cathode Salt bridge Al+3 & NO3 -1 Pb+2 & NO3-1 What are the half-reactions? Al Al+3 + 3e- Al metal is the electrode – it is dissolving. Al+3 ions go into the solution. (loosing mass) Pb+2 + 2e- Pb Pb+2 ions are in the solution. They pick up 2 electrons at the surface of the Pb electrode & plate out. (gaining mass) Al is higher on Table J = anode = oxidation Pb is lower on Table J = cathode = reduction Overall Reaction 2(Al Al+3 + 3e-) 3(Pb+2 + 2e- Pb) ____________________________ 2Al + 3Pb+2 2Al+3 + 3Pb 2Al + 3Pb+2 2Al+3 + 3Pb • Which electrode is losing mass? Al • Which electrode is gaining mass? Pb • What’s happening to the [Al+3] in solution? Increasing • What’s happening to the [Pb+2]in solution? Decreasing Animation of Voltaic Cell Let’s see the animation again: http://www.chem.iastate.edu/group/Greenbo we/sections/projectfolder/flashfiles/electroCh em/volticCell.html Videos • http://www.youtube.com/watch?v=kfgtU9D DvdY&feature=channel (3:29 min) 21.3 Section 21.3 – Electrolytic Cells • Objectives – Distinguish between electrolytic and voltaic cells – Describe the chemical changes that take place during electrolysis – Name three ways that electrolysis is used in metal processing Electrolytic vs. Voltaic Cells • Voltaic cells convert chemical energy to electrical energy during a spontaneous redox reaction • Electrolytic cells use an electric current to make a nonspontaneous redox reaction go forward – This process of using electrical energy to bring about a chemical change is called electrolysis • Applications of electrolysis include: – Silver plating dishes and utensils – Gold-plating jewelry – Chrome plating automobile parts Voltaic vs Electrolytic Cells – Student version Answer Key Voltaic & Electrolytic Cells What’s the same? • Anode is always the site of oxidation – AN OX • Cathode is always the site of reduction – RED CAT • Electrons always flow from anode to cathode Voltaic & Electrolytic Cells What’s different? • Presence of a battery – Voltaic – no battery – Electrolytic – has battery • Which metal is the anode & which is the cathode – Voltaic – anode is more active metal spontaneous – Electrolytic – anode is less active metal non-spontaneous • Charge on anode & charge on cathode – Voltaic – anode is + – Electrolytic – cathode is - Electroplating • 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 platting is the anode (positive) Electroplating e- Example – copper plating a key • The key is the cathode • Copper is the anode Cu Anode oxidation Cu dissolves Cu2+ ions in solution combine with excess eon key to form Cu coating Electroplating • Usually, the object to be electroplated, such as a spoon, is cast of an inexpensive metal. It is then coated with a thin layer of a more attractive, corrosion-resistant, and expensive metal, such as silver or gold. Oxidation… so is Ag the anode anode or cathode? Oxidation Look at the direction of e- flow, what does that tell us about where oxidation occurs? Ag0 Ag+ + 1eCathode Ag+ + e- Ag0 Silver coats spoon Electrolysis of Molten Sodium Chloride • Our next example is the electrolysis of molten sodium chloride – This is the only way to produce sodium metal • Why do you think we need to melt the sodium chloride first? – So the ions are mobile – when you hear molten NaCl – think of ions floating around in liquid Let’s examine the electrolytic cell for molten NaCl. Which way do the electrons flow through the wire? Electrons always flow from anode to cathode…which is the anode? What is the charge on the anode? What is the charge on the cathode? What electrode are the Cl- attracted to? The Na+? + - - + Inert electrode e- ClNa+ Anode Ox: 2Cl- Cl2 + 2e- Na+ Cl- Inert electrode e- Cathode Red: Na+ + 1e- Na Chlorine ions combine to form chlorine gas Remember this is molten NaCl, when all the Cl escapes as gas, we are left with Na e- + - e- Inert electrode ClCl2(g) eCl- Anode Ox: 2Cl- Cl2 + 2e- 0 Na Na+ Na0+ e- Inert electrode - + Cathode Red: Na+ + 1e- Na Electrolysis of Molten NaCl Observe the reactions at the electrodes - battery + Cl2 (g) escapes Na (l) Clelectrode half-cell (-) ClCl- Na+ + e- Na NaCl (l) Na+ Na+ (+) Na+ electrode half-cell 2Cl- Cl2 + 2e- Predicting Spontaneous Redox Reactions In the 2 beakers a strip of Cu was placed in a solution Zn(NO3)2 or AgNO3. Which beaker had the Zn(NO3)2 & which had AgNO3? Beaker (a) was the solution of AgNO3 and beaker (b) was the solution of Zn(NO3)2 Because there is no outside source of electricity the only reaction that can occur is a spontaneous one – Cu is above Ag therefore spontaneous, Cu is below Zn, therefore a nonspontaneous reaction and would require an outside source of energy to force the reaction to occur End of Ch 21 Corrosion Prevention Zinc is more easily oxidized than iron, therefore the zinc will be oxidized as opposed to the iron prevent corrosion of the iron metal Application: Corrosion Charging a Battery When you charge a battery, you are forcing the electrons backwards (from the + to the -). To do this, you will need a higher voltage backwards than forwards. This is why the ammeter in your car often goes slightly higher while your battery is charging, and then returns to normal. In your car, the battery charger is called an alternator. If you have a dead battery, it could be the battery needs to be replaced OR the alternator is not charging the battery properly. Predicting Redox Reactions A + BX B + AX A & B are metals. If metal A is above metal B in Table J, the reaction is spontaneous. X + AY Y + AX • X & Y are nonmetals. If nonmetal X is above nonmetal Y in Table J, the reaction is spontaneous. Electrolysis of Molten NaCl At the microscopic level - e- battery + NaCl (l) cations migrate toward (-) electrode Cl- (-) cathode Na+ + e- Na+ ClCl- Na Na+ e(+) Na+ anions migrate toward (+) electrode anode 2Cl- Cl2 + 2e-