SHOCK TO THE SYSTEM! ELECTROCHEMISTRY

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1 SHOCK TO THE SYSTEM! ELECTROCHEMISTRY REVIEW I. Re: Balancing Redox Reactions. A. Every redox reaction requires a substance to be oxidized (loses electrons). a.k.a. reducing agent 2. reduced (gains electrons). a.k.a. oxidizing agent. 3. LEORA goes GEROA!!! B. The part of the reaction that focuses solely on the reduction or the oxidation is known as the halfreaction. C. Balancing using the half-reaction method (Use if you are given ions) 1. Identify what is being reduced and oxidized. 2. Break the reaction down into two half reactions. a. If a species is not part of the reaction, leave it out 3. Balance all atoms. 4. Write in the number of electrons gained or lost as a product or reactant. 5. Multiply each half-reaction until the electrons are equal for both. 6. Recombine the half-reactions and cancel out any repeating species (there should be no electrons left). 7. Example: Fe 3+ (aq) + Cu(s) Cu 2+ (aq) + Fe(s) D. For acidic solutions: 1. Complete steps C Balance the Os with H 2 O 3. Balance the Hs with H + 4. Complete steps C Example: Balance the following in acidic medium: Al(s) + MnO 4- (aq) Al 3+ (aq) + Mn 2+ (aq) E. For basic solutions: 1. Complete steps D For each H +, add an OH - to both sides. 3. Combine H + and OH - to form H 2 O 4. Cancel out H 2 Os 5. Complete step D4 6. Ex: Balance the following in a basic medium: Mg(s) + OCl - (aq) Mg(OH) 2 (s) + Cl - (aq) GALVANIC CELLS (BATTERIES!) I. An apparatus in which chemical energy is changed to electrical energy (electricity). II. Components of a galvanic cell. A. Need a redox reaction to occur. B. Each half-reaction must take place in a separate container 1. A wire must be present for electrons to move through. 2. If reactants were placed in the same container, reaction would occur without the movement of electrons (no electricity). 3. Usually, each compartment has a solid electrode. If the half reaction does not have a solid involved, platinum is usually used.

4. Problem with the cell below: a. The electrons will build up on one side and the reaction will stop. b. We need a way to keep the reaction going (prevent electron buildup).

Will produce a current (flow of electrons) for a long time. D. Other definitions: 1. Oxidation occurs at the anode. Negative charge 2.

Draw a galvanic cell for the following reaction: Cu 2+ + Zn(s) Zn 2+ (aq) + Cu(s) Show the direction of electron and ion flow. 1.

2 4. Problem with the cell below: a. The electrons will build up on one side and the reaction will stop. b. We need a way to keep the reaction going (prevent electron buildup). C. The cell NEEDS a salt bridge. 1. Can be a U-tube or a porous disk. 2. Allows ions to move without mixing the solutions. 3. Keeps the overall charge zero. 4. Will produce a current (flow of electrons) for a long time. D. Other definitions: 1. Oxidation occurs at the anode. Negative charge 2. Reduction occurs at the cathode. Positive charge. E. Draw a galvanic cell for the following reaction: Cu 2+ + Zn(s) Zn 2+ (aq) + Cu(s) Show the direction of electron and ion flow. 1. Electrons flow from the anode to the cathode. F. What if one or both of the reactions do not involve metal electrodes? Sometimes a cell requires a gas like H 2. III. What happens to the electrodes over time? A. Cathode electrode will grow (metal deposited) over time. B. Anode electrode will shrink (metal ionized/dissolved) over time.

IV. Line notation: How to identify cells: A. A single vertical line indicates a change in state. B. A double vertical line indicates the separation of half-cells. C.

Example: STANDARD ELECTRON REDUCTION POTENTIALS I. The most important (SERP) table (p. 481) in this chapter!!! A. Hints: 1. All of the reactions listed are REDUCTION half rxns!!! 2.

3 IV. Line notation: How to identify cells: A. A single vertical line indicates a change in state. B. A double vertical line indicates the separation of half-cells. C. Within each half-cell the reactants are listed before the products. D. The line notation for the anode is written before the cathode. E. Concentrations are in parentheses. F. Example: STANDARD ELECTRON REDUCTION POTENTIALS I. The most important (SERP) table (p. 481) in this chapter!!! A. Hints: 1. All of the reactions listed are REDUCTION half rxns!!! 2. Electron cell potential (E ) = volts (V) (intensive property) 3. You can use this table to help you balance redox reactions! 4. All of these potentials are measured in standard state: 1.0M, 1 atm, 298K 5. Hydrogen is the standard (0.00V) 6. You can use it to predict reactivity as well (like an activity series).

4 II. Calculating potentials: A. Flip the sign of the cell potential of the oxidation half-rxn. B. Just add them. C. If you need to multiply a half reaction to balance the electrons, DO NOT multiply the potential (intensive property). D. If the voltage is negative, reaction WILL NOT happen (not thermodynamically favored). E. Examples: 1. PbO 2 + Na Pb 2+ + Na + 2. Fe 3+ + Mg Mg 2+ + Fe 2+ III. Reduction strength A. The stronger the reduction potential, the stronger the oxidizing agent. B. Examples 1. Is H 2 (g) capable of reducing Ag + (aq)? 2. Is H 2 (g) capable of reducing Ni 2+ (aq)? 3. Is Fe 2+ (aq) capable of reducing VO 2+ (aq)? 4. Is Fe 2+ (aq) capable of reducing Cr 3+ (aq)? 5. Rank the following from the strongest oxidizing agent to the weakest oxidizing agent: Ce 4+, Fe 2+, Fe 3+, Mg 2+, Mg, Ni 2+, Sn CELL POTENTIAL AND FREE ENERGY I. Definitions: A. Volt (V): joule/coulomb (J/C) B. Coulomb (C): charge of something (an e- in this case) C. Faraday (F): charge of a mole of electrons: 96,500C/mol II. Free energy ( G) is related to cell potential (E ) A. G=-nFE 1. n=moles of electrons involved in the reaction. 2. F=Faraday s constant B. In standard conditions, G =-nfe C. Ex: What is the free energy change involved in following reaction: Cu 2+ (aq) + Zn(s) Cu(s) + Zn 2+ (aq)? D. Ex: Calculate the free energy change for the following reaction: Cu 2+ (aq) + Fe(s) Cu(s) + Fe 2+ (aq) E. This is why a positive voltage yields a spontaneous reaction. CELL POTENTIAL AND CONCENTRATION I. Remember our assumption that everything was at 1.0M (Thus, K=1) A. Electron potential DOES depend on concentration. B. Remember Q and K 1. Q=reaction quotient 2. K=equilibrium constant C. Cells will run until they reach equilibrium ( G=0 and E=0). Thus, if 1. If Q is further from equilibrium, the magnitude of the cell potential will increase 2. If Q is closer to equilibrium, the magnitude of the cell potential will decrease

Concentration Cells: A. Can you have a current with the same electrodes and a diﬀerence in concentration? B. YES!!! Example: CELL POTENTIAL AND EQUILIBRIUM I.

5 D. All batteries will discharge until it reaches equilibrium. When a galvanic cell reaches equilibrium, the battery dies. II. Ex: What would have to be done to increase the voltage of the following reaction? Cu2+ + Zn Zn2+ + Cu A. To increase voltage, [products]<[reactants] B. To decrease voltage, [reactants]<[products] V. Concentration Cells: A. Can you have a current with the same electrodes and a diﬀerence in concentration? B. YES!!! Example: CELL POTENTIAL AND EQUILIBRIUM I. Remember from the last chapter that G = -RTlnK. II. Thus, you can ﬁnd K once you ﬁnd G. A. Given E, use G = -nfe to ﬁnd G. B. Then, you can use G to ﬁnd K through G = -RTlnK. III. All batteries will discharge until it reaches equilibrium. When the galvanic cell reaches equilibrium, the battery dies. ELECTROLYSIS I. Electrolysis: Adding electricity to a solution to break up a substance. II. Electrolytic cell: The opposite of a galvanic cell. A. This is created by adding electrons into a nonspontaneous cell. B. E for electrolytic cells are negative (nonspontaneous) C. The reverse reaction takes place and it is caused by the power source, which pumps enough electrons to cause an opposite ﬂow D. These reactions do not need a salt bridge. III. Electrolysis problems A. First some deﬁnitions: 1. Ampere (A): number of Coulomb per second (C/s) 2. Remember the Faraday: 96, 500C/mol of eb. These problems are like stoichiometry problems (start dimensional analysis with time or mass, whichever is given). C. Below: using electrolysis to split water (will not work without salt water). D. Ex.: Calculate the amount of time required to produce 1000 g of magnesium metal by electrolysis of molten MgCl2 using a current of 50A. E. Ex: A Cr3+ (aq) solution is electrolyzed, using a current of 7.60A. What mass of Cr (s) is plated out after 2.00 days? 1. Known as electroplating. 2. Can also be used to produce pure metals. F. What amperage is required to plate out mol Cr form a Cr3+ solution in a period of 8.00 hours?

IV. Determining electrolytic cells. A. Reactions can be electrolyzed to yield chemical energy. B.

Examples: What reaction will take place at the cathode and the anode when each of the following is electrolyzed? 1. 1.0M KF solution 2. 1.0M CuCl 2 solution 3. 1.0M H 2 O 2 solution containing 1.

You can assume that 1 Faraday of charge =1 mole of electrons picked up. C.

0 faradays passed through the following electrolytic cells. How many moles of metal were plated? 1. AgNO 3 2. FeSO 4 3. FeCl 3 APPLICATIONS I.

battery acid is sulfuric acid. 2. spongy for more surface area. C. Alkaline battery 1. Cathode is a graphite rod. 2. The paste is the salt bridge. 3.

6 IV. Determining electrolytic cells. A. Reactions can be electrolyzed to yield chemical energy. B. You can determine what reaction will occur by finding out which half-reaction will yield the more positive voltage. C. Examples: What reaction will take place at the cathode and the anode when each of the following is electrolyzed? M KF solution M CuCl 2 solution M H 2 O 2 solution containing 1.0M HCl THE FARADAY I. What we have learned is the constant (conversion unit). II. It can be a measured unit. A. A Faraday is the charge on one mole of electrons. B. You can assume that 1 Faraday of charge =1 mole of electrons picked up. C. Therefore, it takes two Faradays to form one mole of Ni from Ni 2+ because there are two electrons present. D. Example: 3.0 faradays passed through the following electrolytic cells. How many moles of metal were plated? 1. AgNO 3 2. FeSO 4 3. FeCl 3 APPLICATIONS I. A battery is a galvanic cell or a series of galvanic cells put together. A. All batteries need an anode and a cathode. B. Car battery: 1. battery acid is sulfuric acid. 2. spongy for more surface area. C. Alkaline battery 1. Cathode is a graphite rod. 2. The paste is the salt bridge. 3. Anode/Cathode are solids. II. Corrosion: When iron rusts... A. Rusting happens when a mini-galvanic cell is formed. B. The iron dissolves as it is oxidized. C. How do you stop this? Get rid of the salt bridge 1. Do not expose iron to water (because it forms a solution). 2. Paint it! Prevents a solution forming with iron. D. Another way to do this is to use a sacrificial anode 1. Use a more active metal so that it will corrode first.

Electrochemistry Pearson Education, Inc. Mr. Matthew Totaro Legacy High School AP Chemistry

Electrochemistry Pearson Education, Inc. Mr. Matthew Totaro Legacy High School AP Chemistry 2012 Pearson Education, Inc. Mr. Matthew Totaro Legacy High School AP Chemistry Electricity from Chemistry Many chemical reactions involve the transfer of electrons between atoms or ions electron transfer