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Transcript of Electrochemistry Guaranteed to give you a jolt. Electrochemical cells A chemical system in which...
![Page 1: Electrochemistry Guaranteed to give you a jolt. Electrochemical cells A chemical system in which oxidation and reduction can occur – often a single.](https://reader035.fdocuments.us/reader035/viewer/2022062720/56649efb5503460f94c0d6b1/html5/thumbnails/1.jpg)
Electrochemistry
Guaranteed to give you a jolt
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Electrochemical cells
A chemical system in which oxidation and reduction can occur – often a single displacement reaction
Zn + Cu+2 Zn+2 + Cu Oxidation reaction and reduction
reaction are physically separated so that useable work can be obtained from the reaction
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Electrochemical cells Voltaic cells (galvanic cells) – redox
occurs spontaneously Anode – where oxidation takes place
(solid metal becomes aqueous positive ions)
Cathode – where reduction takes place (metal ions deposit as solid metal)
AN OX RED CAT
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Electrochemical cells
anode cathode
Salt bridge
Electrons flow from anode to cathode through wire
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Electrochemical cells
Salt bridge – usually KNO3 or a semipermeable membrane– completes circuit by providing mobile ions
Conditions for spontaneity – one metal must lose electrons more easily than another (metals can be identical if one is warmer) p. 288 (activity series)
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Electric current
Rate of flow – amperes1 amp = 1 coulomb/sec
1 coulomb = 1/96490 of a mole of electrons (6.24x1018e-)
Faraday’s number = 96490coul/mole Potential – volts (joules/coulomb)
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Cell potential
Half cell potential (E) – the voltage contribution of a half reaction to the cell
Standard half cell potential (Eº) - voltage when solutions are 1M, gases are 1atm and temp. is 25ºC.
Half cell potentials are given as reductions relative to the reduction of H+ to H2, which is assigned 0 volts.
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Cell potential
Half cell potential is a measure of the tendency of a particle to gain electrons.
The cell potential is the sum of the two half cell potentials.
For oxidation, the sign of the reduction potential is reversed.
A positive cell potential means a spontaneous reaction, i.e. a galvanic cell.
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Cell potential
Example. Find the cell potential for a cell with a mercury/mercury (II) cathode and a lead/lead (II) anode.
Standard reduction potentials:Hg+2 +2e- Hg Eº = 0.851V
Pb+2 + 2e- Pb Eº = -0.1262V Lead is the anode, so it is oxidized
(reduction equation is reversed)
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Cell potential
Hg+2 +2e- Hg Eº = 0.851VPb Pb+2 + 2e- Eº = 0.1262V
Cell potential is the sum of the half cell potentials
Hg+2 + Pb Pb+2 + Hg Eº = 0.851V + 0.1262V = 0.977V
Potential is intensive, so it’s not affected by coefficients
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Cell potential
A positive cell potential means a spontaneous reaction, i.e. a galvanic cell.
Galvanic cell calculations Cell notation
Zn|Zn+2║Cu+2|Cu anode salt bridge cathode
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Galvanic cell calculations
Vertical lines represent the barrier between two different states of matter.
Two different materials in the same part of the cell are separated by commas.
H2,Pt|H+║Ag+|Ag+
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Galvanic cell calculations
Calculate the voltage of this cell:Mg|Mg+2║Au+3|AuMg Mg+2 + 2e-
E º = +2.37Vcathode: Au+3 + 3e- Au
E º = +1.50V Total cell voltage = 3.87V
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Current calculations
Extent of oxidation/reduction and amount of material oxidized or reduced is directly related to the number of electrons transferred
Moles e- = current x time / Faraday’s # = It/F F =
96485coul/mole
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Current calculations A zinc anode with mass 2.30g is used
in a copper/zinc cell. The cell has a current of 0.00140 amps. How long will the electrode last?
Solution: 2.30 g Zn is 0.0352 mole zinc.
Each mole zinc requires 2e-, so 0.0704 mole e- are required to completely oxidize the anode.
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Current calculations
There are 96,485 coulombs/mole, so 6790 coulombs are needed.
0.0704mol e- x 96485 coul/mol = 6790c At 0.00140 coulombs/sec, the
electrode will last 4.85x106 sec, or 1350 hours.
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Types of cells Concentration
cell – uses identical electrodes – potential difference due to concentration differences in the cell
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Nonstandard cells
Ecell = Eºcell – (RT/nF )lnQ Q = reaction quotient For the reaction aA + bB cC + dD,
Q = [C]c[D]d/[A]a[B]b
Find the voltage for a zinc/copper cell at 55ºC where the concentrations of zinc and copper are 0.10 and 0.75M respectively.
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Types of cells
Leclanché cell (dry cell) MnO2, water, NH4Cl in a paste around a graphite cathode, with a zinc anode
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Types of cells Overall Leclanché cell equation:
Zn+MnO2+NH4ClZnCl2+Mn2O3+NH3+H2O
Balance it!Zn Zn+2 + 2e-
2H+ + 2MnO2 + 2e- Mn2O3 + H2O
2H++2MnO2+ZnZn+2+Mn2O3+H2O
Add spectators (NH3 and Cl-):
2NH4Cl+2MnO2+ZnZnCl2+Mn2O3+H2O+2NH3
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Types of cells Alkaline
manganese cell – used in alkaline batteries – uses KOH as electrolyte
Zn+MnO2+H2OZn(OH)2+
Mn2O3
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Types of cells
Electrolytic cells – nonspontaneous redox reaction is forced to proceed by application of an electric current
Electrolysis of water – Hoffman apparatus
2H2O 2H2+ O2
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Electrolysis of aluminum oxide
Alumina (Al2O3) from bauxite is dissolved in molten cryolite (Na3AlF6).
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Electrolysis of aluminum oxide The steel container is coated with
carbon (graphite) and serves as the negative electrode (anode).
Electrolysis of the alumina/cryolite solution gives aluminum at the cathode and oxygen at the anode.
Aluminum is more dense than the alumina/cryolite solution, and so falls to the bottom of the cell.
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Electrolysis of aluminum oxide Aluminum can be tapped off the bottom as
pure liquid metal. The overall reaction is:
2Al2O3(l) 4Al(l) + 3O2(g)
Oxygen is discharged at the positive carbon (graphite) anode.
Oxygen reacts with the carbon anode to form carbon dioxide gas. The carbon anode slowly disappears as carbon dioxide and needs to be replaced regularly.
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Anodic protection