Sec 5.8 Electrochemical Cells

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1 Sec 5.8 Electrochemical Cells Demonstration (Cu/Zn Cell) Cu Definitions 1 M Zn(NO 3 ) 2 1 M Cu(NO 3 ) 2 Electrochemical cell A device which converts chemical energy into electrical energy Electrode A conductor (usually a metal) at which a half-cell reaction (oxidation or reduction) occurs. Anode - The electrode at which oxidation occurs. (A & O are both vowels) LEOA Cathode The electrode at which reduction occurs. (R & C are both consonants) GERC Half-cell reactions Anode Oxidation half-rxn Zn(s) à Zn 2+ (aq) + 2e - Metal atoms are changed to + ions. Metal dissolves and anode loses mass as the cell operates. Cathode Reduction half-rxn Cu e - à Cu (s) + ions are changed to metal atoms. New metal is formed so the cathode gains mass. Flow of Electrons -Since the anode loses e - s, (LEOA) and the cathode gains e - s (GERC) Electrons flow from the anode toward the cathode in the wire (conducting solid) V e - e - Anode Zn à Zn e - Zn Cu Cathode Cu2+ + 2e - à Cu e - s flow from A à C in the wire 1

2 Flow of Ions in the salt bridge -Salt bridge contains any electrolyte (conducting solution) K + NO 3 - Porous barrier which ions can pass through Most common U-tube containing a salt such as KNO 3 or NaNO 3 Common things which contain electrolytes, e.g.) potatoes, apples, oranges, lemons, frogs, people Anode Cathode Zn à Zn e - If there was no SB the + ions (cations), Zn +2, would build up at the anode 1 M Zn 2+ Cations (+) flow (migrate) toward the cathode and Anions ( ) flow toward the anode 1 M Cu 2+ Cu e - à Cu the + ions are used up, so if there was no SB, there would be an excess of ions (anions), NO 3 -, at the cathode In the salt bridge Show with arrows the direction of flow of all ions (Zn 2+, NO3 -, K +, Cu 2+ ) Also show the direction of flow of electrons in the wire Cu Zn 2+ Cu 2+ 2

3 Identifying the anode and cathode - Look at the reduction table - All half-rxn s are reversible (can go forward or backward) - All are written as reductions (GERC) - Their reverse would be oxidations (LEOA) - The half rxn which is lower on the table will be a stronger reducing agent therefore it will be oxidized = anode - The half-rxn which is higher on the table will be a weaker reducing agent, therefore it will be reduced = cathode So the higher half-rxn is the cathode (HIC) So the lower half-rxn is the anode (Notice Cu e - à Cu is higher than Zn e - à Zn so Cu gets to be the cathode) Also notice that the Anode reaction Is Reversed (AIR) (Anode rxn: Zn à Zn e - ) Question. Fill in the following table. Use your reduction table: Metal/ion Metal/ion Cathode (HIC) Cathode Halfrxn Anode Anode (AIR) Half-rxn Ag/Ag + Fe/Fe 2+ Ag (higher) Ag + + e - à Ag Fe (lower) Fe à Fe e - Zn/Zn 2+ Pb/Pb 2+ Ni/Ni 2+ Al/Al 3+ Au/Au 3+ Ag/Ag + Mg/Mg 2+ H 2 /H + Co/Co 2+ Sn/Sn 2+ 3

4 Summary of Electrochemical Cells (ECC s) so far 1) Electrochemical cells convert _CHEMICAL energy into _ELECTRICAL_ energy. 2) ANODE is the electrode where oxidation occurs. 3) Electrons are _LOST at the anode. 4) _CATHODE_ is the electrode where reduction occurs. 5) In the half-rxn at the cathode, e - s are on the _LEFT side of the equation. 6) Electrons flow from the _ANODE toward the _CATHODE in the COPPER WIRE. 7) The higher half-rxn on the table is the one for the _CATHODE and is not reversed. 8) The lower half-rxn on the table is the one for the _ANODE and is reversed. 9) Electrons do not travel through the SALT BRIDGE, only through the COPPER WIRE. 10) Ions (cations & anions) do not travel through the wire but only through the SALT BRIDGE. 11) The salt bridge can contain any _ELECTROLYTES. 12) The anode will LOSE (gains/loses) mass as it is _OXIDIZED (oxidized/reduced). 13) The cathode will _GAIN mass as it is REDUCED (oxidized/reduced). Read p Ex 34 a-e & 35 a-e p

5 Sec 5.8 Electrochemical Cells Demonstration (Cu/Zn Cell) Cu Definitions 1 M Zn(NO 3 ) 2 1 M Cu(NO 3 ) 2 Electrochemical cell Electrode Anode - Cathode Half-cell reactions Anode Oxidation half-rxn Metal atoms are changed to + ions. Cathode Reduction half-rxn + ions are changed to metal atoms. Flow of electrons -Since the anode loses e - s, (LEOA) and the cathode gains e - s (GERC) V e - e - Anode Zn à Zn e - Zn Cu Cathode Cu2+ + 2e - à Cu e - s flow from A à C in the wire 5

6 Flow of Ions in the salt bridge -Salt bridge K + NO 3 - Porous barrier which ions can pass through Most common U-tube containing a salt such as KNO 3 or NaNO 3 Common things which contain electrolytes, e.g.) potatoes, apples, oranges, lemons, frogs, people Anode Cathode Zn à Zn e - Cu e - à Cu 1 M Zn 2+ Cations (+) flow (migrate) toward the cathode and Anions ( ) flow toward the anode 1 M Cu 2+ In the salt bridge Show with arrows the direction of flow of all ions (Zn 2+, NO3 -, K +, Cu 2+ ) Also show the direction of flow of electrons in the wire Cu Zn 2+ Cu 2+ 6

7 Identifying the anode and cathode - Look at the reduction table - All half-rxn s are reversible (can go forward or backward) - All are written as (GERC) - Their reverse would be (LEOA) - The half rxn which is on the table will be a therefore it will be = - The half-rxn which is on the table will be a, therefore it will be = So the half-rxn is the (HIC) So the half-rxn is the (Notice Cu e - à Cu is higher than Zn e - à Zn so Cu gets to be the cathode) Also notice that the Anode reaction Is Reversed (AIR) (Anode rxn: Zn à Zn e - ) Question. Fill in the following table. Use your reduction table: Metal/ion Metal/ion Cathode (HIC) Cathode Halfrxn Anode Anode (AIR) Half-rxn Ag/Ag + Fe/Fe 2+ Ag (higher) Ag + + e - à Ag Fe (lower) Fe à Fe e - Zn/Zn 2+ Pb/Pb 2+ Ni/Ni 2+ Al/Al 3+ Au/Au 3+ Ag/Ag + Mg/Mg 2+ H 2 /H + Co/Co 2+ Sn/Sn 2+ 7

8 Summary of Electrochemical Cells (ECC s) so far 1) Electrochemical cells convert energy into energy. 2) is the electrode where oxidation occurs. 3) Electrons are at the anode. 4) is the electrode where reduction occurs. 5) In the half-rx at the cathode, e - s are on the side of the equation. 6) Electrons flow from the toward the in the. 7) The higher half-rx on the table is the one for the and is not reversed. 8) The lower half-rx on the table is the one for the and is reversed. 9) Electrons do not travel through the, only through the. 10) Ions (cations & anions) do not travel through the wire but only through the. 11) The salt bridge can contain any. 12) The anode will (gains/loses) mass as it is (oxidized/reduced). 13) The cathode will mass as it is (oxidized/reduced). Read p Ex 34 a-e & 35 a-e p

9 Sec 5.9 Standard reduction potentials and voltages Voltage The tendency for e - s to flow in an electrochemical cell. (Note: a cell may have a high voltage even if no e - s are flowing. It is the tendency (or potential) for e - s to flow. -Can also be defined as the potential energy per coulomb. (Where 1C = the charge carried by 6.25x10 18 e - ) 1 Volt = 1 Joule/Coulomb Reduction potential of half-cells -The tendency of a half-cell to be reduced (take e - s) Voltage only depends on the difference in potentials not the absolute potentials. The voltage of a cell depends only on the difference in reduction potentials of the two half-cells. -Relative potentials of half-cells can only be determined by connecting with other half-cells and reading the voltage. Bubbles of H 2 - A standard half-cell was arbitrarily chosen to compare other half-cells with. - It was assigned a reduction potential of V - It is: 2H + (aq) + 2e - H 2(g) E o = V Connecting wire H 2 (g) (at standard pressure kpa) 1.0 M H + (HCl) solution at 25 o C - Since potential depends on gas pressure, temp. and [H + ] this half cell is assigned a value of V at standard state which is 25 o C, kpa and [H + ] =1.0 M Means standard state Platinum electrode Potential E o = 0.000V Surface for metal, gas & solution to contact -The standard half-cell acts as an anode (LEOA) or cathode (GERC) depending on what it is connected to: 9

10 Using the reduction table to find the initial voltage of ECC s at standard state 1) Find the two metals on the reduction table. Higher one is the cathode. (HIC) 2) Write the cathode half-rxn as is on the table (cathode - reduction) Include the reduction potential (E o ) beside it. 3) Reverse the anode reaction (AIR) (anode à oxidation) Reverse the sign on the E o (If (+) à (-) or If (-) à (+)) 4) Multiply half-rxn s by factors that will make e - s cancel. DON T multiply the E o s by these factors. 5) Add up half-rxn s to get overall redox reaction. 6) Add up E o s (as you have written them) to get the initial voltage of the cell. Example: A cell is constructed using Nickel metal and 1M nickel (II) nitrate along with Fe metal and 1M Iron (II) nitrate. a) half-rxn at the cathode (with the E o ) - b) half-rxn at the anode (with the E o ) - c) Write the balanced equation for the overall reaction (with the E o ) d) What is the initial cell voltage? V Another example: A cell is constructed using aluminum metal, 1M Al(NO 3 ) 3 and lead metal with 1M Pb(NO 3 ) 2. Use the method in the last example to write the overall redox reaction and find the initial cell voltage. Overall redox reaction Initial cell voltage: volts. 10

11 Example; A student has 3 metals: Ag, Zn and Cu; three solutions: AgNO 3, Zn(NO 3 ) 2, and Cu(NO 3 ) 2, all 1M. She also has a salt bridge containing KNO 3 (aq) wires and a voltmeter. a) Which combination of 2 metals and 2 solutions should she choose to get the highest possible voltage? Metal: Solution: Metal: Solution: b) Draw a diagram of her cell labeling metals, solutions, salt bridge, wires, and voltmeter. c) Write an equation for the half-rx at the cathode. (with E o ) d) Write an equation for the half-rx at the anode (with E o ) e) Write a balanced equation for the overall redox reaction in the cell (with E o ) f) The initial voltage of this cell is volts. g) In this cell, e - s are flowing toward which metal? In the h) Positive ions are moving toward the solution in the. i) Nitrate ions migrate toward the solution in the. j) metal is gaining mass metal is losing mass As the cell operates. 11

12 The student now wants to find the combination of metals and solutions that will give the lowest voltage. k) Which metals and solution should she use? Metal Solution Metal Solution l) Find the overall redox equation for this cell. m) Find the initial cell voltage of this cell volts. 12

13 Voltages at non-standard conditions Note: When cells are first constructed, they are not at equilibrium. All the voltages calculated by the reduction table are initial voltages. -As the cells operate, the concentrations of the ions change: eg) For the cell: Cu(NO 3 ) 2 /Cu and Zn/Zn(NO 3 ) 2 the cathode ½ reaction is: Cu e - à Cu the anode ½ reaction is: Zn à Zn e - the overall reaction is: Cu 2+ + Zn à Cu + Zn 2+ E o = v E o = v E o = v All electrochemical cells are exothermic (they give off energy) Initially: Cu 2+ + Zn(s) Cu(s) + Zn 2+ + energy Voltage = 1.10 V - As the cell operates [Cu 2+ ] decreases (reactants used up) & [Zn 2+ ] increases (products formed). Both these changes tend to push the reaction to the left (Le Chateliers Principle) Cu 2+ + Zn(s) Cu(s) + Zn 2+ + energy Voltage < 1.10 V Eventually, these tendencies will be equal. At this point, the cell has reached equilibrium. At equilibrium the cell voltage becomes 0.00 V. Predicting spontaneity from E o of a redox reaction If E o for any redox (overall) reaction is positive the reaction is Spontaneous. If E o is negative the reaction is Non-spontaneous Example: a) Find the standard potential (E o ) for the following reaction: 2MnO H 2 O + 3Sn 2+ à 2MnO 2 + 8OH - + 3Sn 4+ b) Is this reaction as written (forward rx) spontaneous? c) Is the reverse reaction spontaneous? E o = 13

14 Solution: a) Find the two half-rx s which add up to give this reaction. The half-rx for MnO 4 - à MnO 2 in basic soln. is at To keep MnO 4 - on the left, this ½rx is written as it is on the table. MnO H 2 O + 3e - à MnO 2 + 4OH - E o = The rest of the overall rx involves Sn 2+ changing to Sn 4+. The ½ reaction for that must be reversed as well as its E o. Since Sn 2+ must stay on the left side, the half-rx on the table must be reversed as well as its E o. Sn 2+ à Sn e - E o = V (MnO H 2 O + 3e - à MnO 2 + 4OH - ) 2 (Sn 2+ à Sn 4+ 2e - ) 3 2MnO H 2 O + 3Sn 2+ à 2MnO 2 + 8OH - + 3Sn 4+ E o = V E o = V E o = V So E o for the overall redox reaction = v b) Since E o is positive, this reaction is spontaneous as written. c) The E o for the reverse reaction would be 0.45 v so the reverse reaction is non-spontaneous. Question a) Calculate E o for the reaction: 3N 2 O 4 + 2Cr H 2 O à 6NO Cr + 12H + b) Is the forward rxn spontaneous? The reverse rxn? Read SW. p Do Ex. 36 a-d & on p

15 Sec 5.11 Practical Applications of Electrochemical Cells -See textbook. p The Lead-Acid Storage Battery (Automobile battery) Anode (Pb Plate) Cathode (PbO 2 plate) Pb PbO 2 Electrolyte H 2 SO 4 (aq) -This cell is rechargeable (Reactions can be reversed) Anode ½ reaction (Discharging or operating) Pb (s) + HSO 4 - (aq) à PbSO 4 (s) + H + (aq) + 2e - Ox # = 0 Ox # = +2 You can tell this is oxidation because ox # of Pb creases and electrons are. (LEOA) Cathode ½ reaction (Discharging or operating) PbO 2 (s) + HSO 4 -(aq) + 3H + (aq) + 2e - à PbSO 4 (s) + 2H 2 O (l) Ox # of Pb = Ox # of Pb = GERC Write the balanced equation for the overall redox reaction (discharging) The overall redox reaction: (discharging or operating) Pb (s) + PbO 2(s) + 2H + (aq) + 2HSO 4 - (aq) à 2PbSO 4 (s) + 2H 2 O (l) + electrical energy Anode Cathode White solid forms on plates as battery discharges Notes: As the cell discharges, the anode (Pb) and cathode (PbO 2 ) disintegrate and the white solid (PbSO 4 ) forms on both plates. Originally, [H + ] & [HSO 4 - ] is high. i.e.) H 2 SO 4 is denser than H 2 O therefore the density (specific gravity) of the electrolyte is high to start with. As the cell discharges, H 2 SO 4 (H + & HSO 4 - ) is used up and H 2 O is formed. Therefore, electrolyte gets less dense as the battery discharges. 15

16 The condition of the battery can be determined using a hydrometer or battery tester. (Higher the float, the denser the electrolyte) Adding electrical energy to this reaction will reverse it (recharging the battery) Charging Reaction: Electrical energy + 2PbSO 4(s) + H 2 O (l) à Pb (s) + PbO 2(s) +2H + (aq) + 2HSO 4 - (aq) Supplied by alternator or generator The E o for the discharging rxn, is volts. A typical car battery has of these in (Series/parallel) to give a total voltage = V. The Zinc-Carbon battery (LeClanche-Cell, Common Dry cell, or regular carbon battery. Often called Heavy Duty ) Carbon Cathode MnO 2 (s) Steel Case Carbon Rod Zinc Anode NH 4 Cl & ZnCl 2 paste electrolyte Paste is acidic (Causes rust when they leak) Cathode ½ reaction: + (GERC) 2MnO 2 (s) + 2NH 4 (aq) +2e - à 2MnO(OH) (s) + 2NH 3 (aq) Or simplified: Mn 4+ + e - à Mn 3+ Anode ½ reaction: (LEOA) Zn (s) + 4NH 3 (aq) à Zn(NH 3 ) e - Or simplified: Zn (s) à Zn e - -Not rechargeable -Doesn t last too long especially with large currents -Fairly cheap 16

17 The alkaline dry cell -Operates under basic conditions - Delivers much greater current - Voltage remains constant - More expensive - Lasts longer MnO 2 (cathode) KOH (Alkaline electrolyte) + Powdered Zn anode Cathode ½ reaction (GERC) 2MnO 2 + H 2 O +2e - à Mn 2 O 3 + 2OH - Anode ½ reaction (LEOA) Zn (s) +2OH - à ZnO + H 2 O +H 2 O +2e - Ox # = Ox # = ***What is the effect of the oxidation potential of the cell if the [OH - ] is decreased around the anode? Fuel cells -Fuel cells are continuously fed fuel and they convert the chemical energy in fuel to electrical energy -more efficient (70-80%) than burning gas or diesel (30-40%) to run generators -no pollution only produces water - can use H 2 and O 2 or hydrogen rich fuels (e.g. methane CH 4 ) and O 2. -used in space capsules H 2 & O 2 in tanks H 2 O produced used for drinking + - H 2 out H 2 in O2 out O 2 in Anode Cathode KOH (aq) electrolyte Carbon electrode impregnated with catalysts Anode ½ reaction: 2H 2(g) + 4OH - (aq) à 4H 2 O (l) + 4e - Cathode ½ reaction: O 2(g) + 2H 2 O (l) + 4e - à 4OH - (aq) Overall reaction: 2H 2(g) + O 2(g) à 2H 2 O (l) 17

Applied electrochemistry The Breathalyzer Test After drinking, breath contains ethanol C 2 H 5 OH. Acidified dichromate (at E o = 1.23V on the reduction table) will oxidize alcohol.

18 Applied electrochemistry The Breathalyzer Test After drinking, breath contains ethanol C 2 H 5 OH. Acidified dichromate (at E o = 1.23V on the reduction table) will oxidize alcohol. The unbalanced formula equation is: C 2 H 5 OH + K 2 Cr 2 O 7 + H 2 SO 4 à CH 3 COOH + Cr 2 (SO 4 ) 3 + K 2 SO 4 + H 2 O Ox # of C= Ox # of Cr= Ox # of C= Ox # of Cr= K 2Cr 2O 7 is yellow (orange at higher concentrations) Cr 2(SO 4) 3 is green - Exhaled air is mixed with standardized acidified dichromate - Put in a spectrophotometer set at the wavelength of green light - More alcohol produces more green Cr 2 (SO 4 ) 3 (green) - Machine is calibrated with known concentration samples of alcohol to ensure accuracy Question: In the reaction above name a) The oxidizing agent b) The reducing agent c) The product of oxidation d) The product of reduction 18

CHAPTER 5 REVIEW. C. CO 2 D. Fe 2 O 3. A. Fe B. CO

CHAPTER 5 REVIEW. C. CO 2 D. Fe 2 O 3. A. Fe B. CO CHAPTER 5 REVIEW 1. The following represents the process used to produce iron from iron III oxide: Fe 2 O 3 + 3CO 2Fe + 3CO 2 What is the reducing agent in this process? A. Fe B. CO C. CO 2 D. Fe 2 O 3