Fundamentals of Electrochemistry

1
Fundamentals of Electrochemistry

• Introduction
• 1.) Electrical Measurements of Chemical Processes
• Redox Reaction involves transfer of electrons
  from one species to another.
• Chemicals are separated
• Can monitor redox reaction when electrons flow
  through an electric current
• Electric current is proportional to rate of
  reaction
• Cell voltage is proportional to free-energy
  change
• Batteries produce a direct current by converting
  chemical energy to electrical energy.
• Common applications run the gamut from cars to
  ipods to laptops

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Fundamentals of Electrochemistry

• Basic Concepts
• 1.) A Redox titration is an analytical technique
  based on the transfer of electrons between
  analyte and titrant
• Reduction-oxidation reaction
• A substance is reduced when it gains electrons
  from another substance
• gain of e- net decrease in charge of species
• Oxidizing agent (oxidant)
• A substance is oxidized when it loses electrons
  to another substance
• loss of e- net increase in charge of species
• Reducing agent (reductant)

(Reduction)
(Oxidation)
Oxidizing Agent
Reducing Agent
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Fundamentals of Electrochemistry

• Basic Concepts
• 2.) The first two reactions are known as 1/2
  cell reactions
• Include electrons in their equation
• 3.) The net reaction is known as the total cell
  reaction
• No free electrons in its equation
• 4.) In order for a redox reaction to occur, both
  reduction of one compound and oxidation of
  another must take place simultaneously
• Total number of electrons is constant

½ cell reactions
Net Reaction
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Fundamentals of Electrochemistry

• Basic Concepts
• 5.) Electric Charge (q)
• Measured in coulombs (C)
• Charge of a single electron is 1.602x10-19C
• Faraday constant (F) 9.649x104C is the charge
  of a mole of electrons
• 6.) Electric current
• Quantity of charge flowing each second
• through a circuit
• Ampere unit of current (C/sec)

Relation between charge and moles
Coulombs
moles
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Fundamentals of Electrochemistry

• Basic Concepts
• 7.) Electric Potential (E)
• Measured in volts (V)
• Work (energy) needed when moving an electric
  charge from one point to another
• Measure of force pushing on electrons

Higher potential difference
Relation between free energy, work and voltage
Volts
Coulombs
Joules
Higher potential difference requires more work to
lift water (electrons) to higher trough
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Fundamentals of Electrochemistry

• Basic Concepts
• 7.) Electric Potential (E)
• Combining definition of electrical charge and
  potential

Relation between free energy difference and
electric potential difference
Describes the voltage that can be generated by a
chemical reaction
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Fundamentals of Electrochemistry

• Basic Concepts
• 8.) Ohms Law
• Current (I) is directly proportional to the
  potential difference (voltage) across a circuit
  and inversely proportional to the resistance (R)
• Ohms (W) - units of resistance
• 9.) Power (P)
• Work done per unit time
• Units joules per second J/sec or watts (W)

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Fundamentals of Electrochemistry

• Galvanic Cells
• 1.) Galvanic or Voltaic cell
• Spontaneous chemical reaction to generate
  electricity
• One reagent oxidized the other reduced
• two reagents cannot be in contact
• Electrons flow from reducing agent to oxidizing
  agent
• Flow through external circuit to go from one
  reagent to the other

Reduction
Oxidation
Net Reaction
AgCl(s) is reduced to Ag(s) Ag deposited on
electrode and Cl- goes into solution
Cd(s) is oxidized to Cd2 Cd2 goes into solution
Electrons travel from Cd electrode to Ag
electrode
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Fundamentals of Electrochemistry

• Galvanic Cells
• 1.) Galvanic or Voltaic cell
• Example Calculate the voltage for the following
  chemical reaction

DG -150kJ/mol of Cd
n number of moles of electrons
Solution
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Fundamentals of Electrochemistry

• Galvanic Cells
• 2.) Cell Potentials vs. DG
• Reaction is spontaneous if it does not require
  external energy

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Fundamentals of Electrochemistry

• Galvanic Cells
• 2.) Cell Potentials vs. DG
• Reaction is spontaneous if it does not require
  external energy

Potential of overall cell measure of the
tendency of this reaction to proceed to
equilibrium Larger the potential, the further
the reaction is from equilibrium and the
greater the driving force that exists
Similar in concept to balls sitting at different
heights along a hill
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Fundamentals of Electrochemistry

• Galvanic Cells
• 3.) Electrodes

Cathode electrode where reduction takes place
Anode electrode where oxidation takes place
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Fundamentals of Electrochemistry

• Galvanic Cells
• 4.) Salt Bridge
• Connects separates two half-cell reactions
• Prevents charge build-up and allows counter-ion
  migration

Salt Bridge

• Contains electrolytes not
• involved in redox reaction.
• K (and Cd2) moves to cathode with
• e- through salt bridge (counter
• balances charge build-up
• NO3- moves to anode (counter
• balances charge build-up)
• Completes circuit

Two half-cell reactions
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Fundamentals of Electrochemistry

• Galvanic Cells
• 5.) Short-Hand Notation
• Representation of Cells by convention start with
  anode on left

ZnZnSO4(aZN2
0.0100)CuSO4(aCu2 0.0100)Cu
Phase boundary Electrode/solution interface
anode
cathode
2 liquid junctions due to salt bridge
Solution in contact with anode its
concentration
Solution in contact with cathode its
concentration
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Fundamentals of Electrochemistry

• Standard Potentials
• 1.) Predict voltage observed when two half-cells
  are connected
• Standard reduction potential (Eo) the measured
  potential of a half-cell reduction reaction
  relative to a standard oxidation reaction
• Potential arbitrary set to 0 for standard
  electrode
• Potential of cell Potential of ½ reaction
• Potentials measured at standard conditions
• All concentrations (or activities) 1M
• 25oC, 1 atm pressure

Ag e- Ag(s) Eo 0.799V
Standard Hydrogen Electrode (S.H.E)
Pt(s)H2(g)(aH2 1)H(aq)(aH 1)
Hydrogen gas is bubbled over a Pt electrode
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Fundamentals of Electrochemistry

• Standard Potentials
• 1.) Predict voltage observed when two half-cells
  are connected

As Eo increases, the more favorable the reaction
and the more easily the compound is reduced
(better oxidizing agent). Reactions always
written as reduction
Appendix H contains a more extensive list
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Fundamentals of Electrochemistry

• Standard Potentials
• 2.) When combining two ½ cell reaction together
  to get a complete net reaction, the total cell
  potential (Ecell) is given by

Where E the reduction potential for the ½
cell reaction at the positive electrode E
electrode where reduction occurs (cathode) E-
the reduction potential for the ½ cell reaction
at the negative electrode E- electrode where
oxidation occurs (anode)
Place values on number line to determine the
potential difference
Electrons always flow towards more positive
potential
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Fundamentals of Electrochemistry

• Standard Potentials
• 3.) Example Calculate Eo, and DGo for the
  following reaction

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Fundamentals of Electrochemistry

• Nernst Equation
• 1.) Reduction Potential under Non-standard
  Conditions
• E determined using Nernst Equation
• Concentrations not-equal to 1M

For the given reaction
aA ne- bB Eo
The ½ cell reduction potential is given by
Where E actual ½ cell reduction potential Eo
standard ½ cell reduction potential n number
of electrons in reaction T temperature (K) R
ideal gas law constant (8.314J/(K-mol) F
Faradays constant (9.649x104 C/mol) A
activity of A or B
at 25oC
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Fundamentals of Electrochemistry

• Nernst Equation
• 2.) Example
• Calculate the cell voltage if the concentration
  of NaF and KCl were each 0.10 M in the following
  cell

Pb(s) PbF2(s) F- (aq) Cl- (aq) AgCl(s)
Ag(s)
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Fundamentals of Electrochemistry

• Eo and the Equilibrium Constant
• 1.) A Galvanic Cell Produces Electricity because
  the Cell Reaction is NOT at Equilibrium
• Concentration in two cells change with current
• Concentration will continue to change until
  Equilibrium is reached
• E 0V at equilibrium
• Battery is dead

Consider the following ½ cell reactions
aA ne- cC Eo dD ne- bB E-o
Cell potential in terms of Nernst Equation is
Simplify
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Fundamentals of Electrochemistry

• Eo and the Equilibrium Constant
• 1.) A Galvanic Cell Produces Electricity because
  the Cell Reaction is NOT at Equilibrium

Since EoEo- E-o
At equilibrium Ecell 0
Definition of equilibrium constant
at 25oC
at 25oC
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Fundamentals of Electrochemistry

• Eo and the Equilibrium Constant
• 2.) Example
• Calculate the equilibrium constant (K) for the
  following reaction

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Fundamentals of Electrochemistry

• Cells as Chemical Probes
• 1.) Two Types of Equilibrium in Galvanic Cells
• Equilibrium between the two half-cells
• Equilibrium within each half-cell

If a Galvanic Cell has a nonzero voltage then
the net cell reaction is not at equilibrium
Conversely, a chemical reaction within a ½ cell
will reach and remain at equilibrium.
For a potential to exist, electrons must flow
from one cell to the other which requires the
reaction to proceed ? not at equilibrium.
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Fundamentals of Electrochemistry

• Cells as Chemical Probes
• 2.) Example
• If the voltage for the following cell is 0.512V,
  find Ksp for Cu(IO3)2

Ni(s)NiSO4(0.0025M)KIO3(0.10
M)Cu(IO3)2(s)Cu(s)
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Fundamentals of Electrochemistry

• Biochemists Use Eo
• 1.) Redox Potentials Containing Acids or Bases
  are pH Dependent
• Standard potential ? all concentrations 1 M
• pH0 for H 1M
• 2.) pH Inside of a Plant or Animal Cell is 7
• Standard potentials at pH 0 not appropriate for
  biological systems
• Reduction or oxidation strength may be reversed
  at pH 0 compared to pH 7

Metabolic Pathways
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Fundamentals of Electrochemistry

• Biochemists Use Eo
• 3.) Formal Potential
• Reduction potential that applies under a
  specified set of conditions
• Formal potential at pH 7 is Eo

Eo (V)
Need to express concentrations as function of Ka
and H.
Cannot use formal concentrations!