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Electrochemical Cells
Contained systems in which oxidation-reduction reactions occur
Three fundamental cell types
Electrolytic: nonspontaneous reactions
Galvanic : spontaneous reactions
Concentration : spontaneous reactions
Electrochemical Cells: Ion flow and electrodes
Anodes: where oxidation takes place
Cathode: where reduction takes place
Current (I) runs from cathode to anode
Electrons move from anode to cathode
Electromotove Force (emf)
corresponds to the voltage or electrical potential difference of the cell
If positive, cell is able to release energy and is spontaneous
If negative, cell must absorb energy (nonspontaneous)
Galvanic (volcaic) Cells
Charge flows as a result of oxidation-reduction reaction between the two half-cells half cells
connecting the two half cell solutions is a salt bridge
As rxn proceeds to equilibrium, movement of electrons converter electrical energy to potential energy; this energy can be used to do work
are spontaneous
Half Cells
two distinct electrodes are placed in separate compartments and connected via conductive material
Each electrode is surrounded by aqueous electrolyte solution (of cations/anions)
Distribution of Charge (galvanic cells)
Wire itself would result in buildup of charge on anode; salt bridge allows exchange of ions to balance out buildup of charge
Salt bridge contains inert electrolyte that will not react with electrodes (anions from salt bridge diffuse towards anode; cations towards cathode)
Cell Diagram
a shorthand rotation representing the reactions in an electrochemical cell
Reactants and products are always listed from left to right
Single vertical line indicates phase boundary
Double vertical line indicates presence of salt bridge or some other barrier
Electrolytic Cells
House nonspontaneous reactions that require input of energy in an electrolysis reaction (ΔG>0) ; chemical compounds are decomposed
Faraday theorized the amount of chemical change induced in an electrolytic cell is directly proportional to the number of moles of electrons during oxidation reduction reaction
Can be determined from the balanced half-reaction
Chemical Change of Electrolytic cell
Equation 12.1 :
Mn+ + n e- → M (s)
n= moles of electrons
M = mole of metal
Moles of Element Deposited on a Plate
Equation 12.2: Number of moles of Element Deposited on a Plate
Mol M = It/nF
I = Current
M = mole of metal iob being deposited at a specific electrode
t = time
n = number of electrons equivalents for a specific metal ion
F = Faraday constant
F is used to compare the number of moles of electrons to the measurable electrical property of charge (C/mol e-)
Equation can be used to determine amount of gas liberated during electrolysis
Concentration Cells
A special type of galvanic cell in which the electrodes are chemically identical
Current is generated as a function of a concentration gradient established between the two solutions surrounding the electrodes
Results in potential difference between two electrodes and drives electrons in direction that establishes equilibrium
Current stops moving when concentrations of ionic species in half-cells are equal, implies that the voltage or emf is 0 when concentrations are equal
Nernst Equation allows calculation of voltage as function of concentration
Ex: neuron cell membrane
Rechargeable Cells
one that functions as both a galvanic and electrolytic cell
Lead-acid Batteries
When fully charged, as a voltaic cell, consists of two half cells (Pb anode and porous PbO2 cathode)
When fully discharged, consists of two PbSO4 electroplated lead electrodes w/ dilute concentration of H2SO4
The cell is part of an electrolytic circuit when charging
External source reverses electroplating process and concentrations acid solution
Discharging
describes how both half reactions cause electrodes to plate with lead sulfate and dilute the acid electrolyte
ex.: lead acid batteries
Energy Density
a measure of a battery’s ability to produce power as a function of its weight
Lead-acid batteries require a heavier amount of battery material to produce a certain output as compared to other batteries
Nickel -Cadmium Batteries
Rechargeable batteries that consist of two half-cells made of solid cadmium (anode) and nickel (III) oxide-hydroxide (cathode) connected by a conductive material
Like lead battery, charging reverses electrolytic cell potentials
Have higher energy density than lead-acid batteries
Provide higher Surge currents: periods of large current (amperage) early in the discharge cycle
Electrode Charge Designations
In galvanic cells
anode is negatively charged (source of electrons) and Cathode is positively charged
Electrons move from negative (low electric potential) to positive (high electric potential)
Current (flow of positive charge) is from high electric potential to low potential
In electrolytic cells
Anode is positively charged and attracts anions from solution while cathode is negatively charged and attracts cations from solution
Tenets of Oxidation-reduction Rxns in Electrochemical Cells
In spite of charge designation, oxidation always takes place at anode and reduction at cathode no matter the cell type
Regardless of charge designation, cathode always attracts cations and anode always attracts anions
Electrons always flow from anode to cathode and current from cathode to anod
Reduction Potentials
Measured in volts and defined relative to the standard hydrogen electrode (has a potential of 0V by conventio)
Can determine which species in a reaction will be oxidized/reduced
The more positive the potential, the greater the tendency to be reduced
Standard Reduction Potential
measured under standard conditions; relative reactivities of different half-cells can be compared to predict the direction of electron flow
Less positive Ered° means greater relative tendency for oxidation to occur
More positive Ered° means greater relative tendency for reduction to occur
Reduction Potentials of Galvanic and Electrolyic Cells
Galvanic cells
Cathode has greater Ered°
Anode has lesser Ered°
Electrolytic cells
Anode has greater Ered°
Cathode has lesser Ered
Reduction and Oxidation Potentials
Reduction and oxidation are opposite processes
To obtain the oxidation potential of a given half reaction, both the reduction half reaction and the sign of the reduction potential are reversed
Standard Electromotive Force
Standard reduction potentials are used to calculate the standard electromotive force
Difference in potential (voltage) between two half-cells under standard conditions
Equation 12.3: Standard Electromotive Force
E°cell = E° red, cathode - E°red, anode
When subtracting, do no multiply potentials by number of moles oxidized or reduced
Electromotive Force and Thermodynamics
In an electrochemical cell, the work done is dependent on the number of Coulombs of charge transferred and the energy available
Equation 12.4: Standard Free Energy of Electrochemical Cell
ΔG° = -nFE°cell
ΔG° = standard change in free energy
n= number of moles of electrons exchanged
F = Faraday constant
E°cell = standard emf of the cell
Reaction Quotient of Electrochemical Cell
Ecell = E°cell - RT/nF (lnQ)
Ecell = emf of cell under nonstandard conditions
E°cell = emf of cell under standard conditions
R = ideal gas constant
n= number of moles of electrons
F = Faraday constant
Q = reaction quotient for the reaction at a given point in time
Assuming T = 298 K, reaction can be simplified to
Equation 12.6
Ecell = E°cell - 0.0592/n (logQ)
An even more simplified version converts natural logarithm to base logarithm
Reaction Quotient of Electrochemical Cells (using products and reactants)
Equation 12.7
Q = [C]c [D]d / [A]a [B]b
Number of terms (C,D, A, B) is dependent on number of reactants and products
Equilibria and Thermodynamics of Electrolytic Cells
Recall equation:
ΔG° = -RTlnKeq
for redox reactions, with equilibrium constants less than 1, the E°cell will be negative because the natural logarithm of any number between 0 and 1 is negative
Property is characteristic of electrolytic cells
Equilibria and Thermodynamics of Galvanic Cells
If equilibrium constant for the reaction is greater than 1, E°cell will be positive because natural logarithm of any number greater than 1 is positive
Property is characteristic of galvanic cells
Thermodynamics of Electrochemical Cells at Equilibrium
Recall equation:
ΔG° = -RTlnKeq
If equilibrium constant is equal to 1, E°cell will be zero
Change in Free Energy w/ Varying Concentration of Electrolytic Cells
Equation 12.9: Change in Free Energy of Cell w/ Varying Concentrations
ΔG = ΔG° + RTlnQ
ΔG = free energy under nonstandard conditions
ΔG° = free energy on standard conditions
You know the rest of the variables
