Electrochemistry

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Electrochemistry

• PHAR 526 Physical Chemistry
• R. Gary Hollenbeck

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Objectives

• Characterization of Oxidation-Reduction Reactions
• Terms and conventions
• Ion Specific Electrodes
• Thermodynamics of Oxidation-Reduction Reactions
• Effects of temperature and the activity of the
  species involved
• Oxidative Degradation of Pharmaceuticals
• Rational selection of antioxidants

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Electrochemical Cell
Each compartment is know as a half cell
Zinc Electrode
Copper Electrode
Zinc Sulfate Solution
Copper Sulfate Solution
Porous Diaphragm
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Electrochemical Cell
Galvanometer
e
Anode (-) (oxidation)
Cathode () (reduction)
e
Zinc has a greater tendency to ionize than Copper
e
Zn2
SO4-2
Zn2
Cu2
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Electrochemical Cell

• A spontaneous reaction occurs when the two
  half-cells are connected with a wire
• Atoms of the zinc electrode go into solution as
  Zn2 ions
• Electrons left behind negatively charged
  electrode anode
• Electrons flow to the copper electrode
• Cu2 ions take on the electrons and deposit
  copper metal at the electrode surface
• Electrons given up to the solution positively
  charged cathode

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Standard Half-cell Reduction Potentials at 25 oC

• By convention, any half-cell reaction is written
  as a reduction
• A single line represents the boundary between an
  electrode and its solution.
• A comma is used to separate two species that are
  present together in the same phase
• If the electrode is acting as an anode
  (oxidation), then the sign of Eo, as given by the
  convention, must be changed.

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Standard Half-cell Reduction Potentials at 25 oC
Increasing tendency for reduction

• Reduction Reaction Reduction
  Electrode Eo (volts)
• ½ Cl2 e- Cl- Cl-Cl2, Pt
  1.360
• ¼ O2 H e- ½ H2O HO2, Pt
  1.229
• Hg2 e- ½ Hg22 Hg2, Hg22 Pt
  0.907
• Ag e- Ag AgAg
  0.799
• ½ Hg2 e- Hg Hg22 Hg
  0.854
• Fe3 e- Fe2 Fe3,
  Fe2Pt 0.771
• ½ I2 e- I- I- I2
  0.536
• Fe(CN)63- e- Fe(CN)64- Fe(CN)63-,Fe(CN)64-
  Pt 0.356
• ½ Cu2 e- ½ Cu Cu2Cu
  0.337

Increasing tendency for oxidation
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Standard Half-cell Reduction Potentials at 25 oC
Increasing tendency for reduction

• Reduction Reaction Reduction
  Electrode Eo (volts)
• ½ Hg2Cl2 e- Hg Cl- Cl-Hg2Cl2, Hg
  0.268
• AgCl e- Ag Cl- Cl-AgCl,Ag
  0.223
• AgBr e- Ag Br- Br-AgBr,Ag
  0.071
• H e- ½ H2 HH2, Pt
  0.000
• ½ Pb2 e- ½ Pb Pb2 Pb
  -0.126
• AgI e- Ag I- I- AgI,Ag
  -0.156
• ½ Ni2 e- ½ Ni Ni2 Ni
  -0.250
• ½ Cd2 e- ½ Cd Cd2 Cd
  -0.403
• ½ Fe2 e- ½ Fe Fe2 Fe
  -0.440

Increasing tendency for oxidation
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Standard Half-cell Reduction Potentials at 25 oC
Increasing tendency for reduction

• Reduction Reaction Reduction
  Electrode Eo (volts)
• ½ Zn2 e- ½ Zn Zn2 Zn
  -0.763
• Na e- Na Na Na
  -2.714
• K e- K K K
  -2.925
• Li e- Li Li
  Li -3.045

Increasing tendency for oxidation
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Electrochemical Cell
Galvanometer
e
Anode (-) (oxidation)
Cathode () (reduction)
Zinc has a greater tendency to ionize than Copper
e
e
Zn2
SO4-2
Porous Diaphragm
Zn2
Cu2
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Electrochemical Cell

• A spontaneous reaction occurs when the two
  half-cells are connected with a wire
• The zinc at the anode is oxidized
• ½ Zn ½ Zn2 e- Eleft
• 0.763 volts
• The copper ions at the cathode are reduced
• ½ Cu2 e- ½ Cu Eright
• 0.337 volts

Eleft and Eright are referred to as electrode
potentials
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Electrochemical Cell

• The two half-reactions can be added together to
  represent the overall cell reaction
• ½ Zn ½ Cu2 ½ Zn2 ½ Cu
• Zn Cu2 Zn2 Cu
• Ecell Eleft Eright
• Ecell voltage or emf of the cell sum of the
  electrode potentials
• Ecell 1.10 volts (at standard state)

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Electrochemical Cell
1.10 volts
e
Anode (-) (oxidation)
Cathode () (reduction)
e
e
Zn2
SO4-2
Zn2 (1M)
Cu2 (1M)
Zn Zn2 (cZn2) Cu2 (cCu2) Cu
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Galvanic Cell

• An electrochemical cell in which a spontaneous
  reaction occurs at the electrode surfaces and
  that can be used to provide electric energy from
  the chemical reaction occurring within it, is
  know as a galvanic cell.

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Example of a glass reference electrode
Ag, AgCl Cl-
This electrode can detect potentials at the
glass/solution interface.
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Metal-Insoluble Salt Electrode

• Ag, AgCl Cl- (c, moles/liter)
• Electrode process
• Ag Ag e-
• Ag Cl- AgCl
• Ag Cl- AgCl e-

Used as a reference electrode since its potential
does not vary with changes in solution
concentration. Eo -0.223 volts
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Thermodynamics of Electrochemical Cells

• ?G -n F E
• ?G Free energy change or work done by the
  electrochemical cell operating reversibly
• n number of equivalents of ions reacting or the
  number of electrons transferred
• F Faraday approximately 96500 coulomb/Eq of
  ions
• E Electromotive Force

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Thermodynamics of Electrochemical Cells

• ?Go -n F Eo
• ?Go Free energy change under standard
  conditions (i.e., fixed temperature and pressure)
• n number of equivalents of ions reacting or the
  number of electrons transferred
• F Faraday approximately 96500 coulomb/Eq of
  ions 23000 cal volt-1 Eq-1
• Eo Electromotive Force under standard conditions

Activities of all reactants and products are
unity.
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What is the standard free energy change for the
following cell reaction in which Eo 1.100 volts?

• Zn Cu2 Zn2 Cu
• ?Go -n F Eo
• ?Go -2 Eq/mole x 23000 cal volt-1 Eq-1 x 1.100
  volts
• ?Go - 50,600 cal/mole

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Determination of Equilibrium Constants

• ?Go -n F Eo
• ?Go -RT ln K -2.303 RT log K
• log K (n F Eo)/(2.303 R T)

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Determination of the Equilibrium Constant

• Zn Cu2 Zn2 Cu
• aZn2 aCu aZn2
• K
• aZn aCu2 aCu2
• Activities of the solid phases Zn and Cu are
  taken as unity.

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Determination of the Equilibrium Constant

• n F Eo
• log K
• 2.303 RT
• log K 50600 cal/mole
• 2.303 (1.987 cal mole-1 deg-1) 298
  deg
• log K 37.1

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Determination of the Equilibrium Constant

• Zn Cu2 Zn2 Cu
• aZn2
• aCu2
• This reaction essentially goes to completion

K 1.26 x 1037
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The Nernst Equation

• By convention, any half cell reaction is written
  as a reduction
• ?(Ox) n e- ?(Rd)
• (reactants) (products)
• ? moles of the oxidized species (Ox) is reduced
  by a reaction involving n electrons to ? moles of
  the reduced species (Rd)

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The Nernst Equation
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The Nernst Equation

• At 25 oC

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The Nernst Equation

• At 25 oC, when the activities of all reactants
  and products are unity.
• Thus, the Nernst Equation determines the
  electromotive force when the activities are not
  unity.

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Determination of an electrode potential
(cathode)

• What is the reduction potential at 25 oC of a
  platinum wire electrode immersed in an acidic
  solution of ferrous ions, at a concentration of
  0.50 m, and of ferric ions, at a concentration of
  0.25 m?
• The activity coefficient for the ferrous ion is
  0.435 and that for the ferric ion is 0.390 under
  these conditions.

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Determination of an electrode potential
(cathode)

• Pt Fe2 (c1, moles/L), Fe3 (c2, moles/L)
• Fe2 Fe3 e- written as an oxidation
• Fe3 e- Fe2 written as a reduction
• activity activity coefficient x concentration
• aFe2 0.435 (0.5) 0.218
• aFe3 0.390 (0.25) 0.0975

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Determination of an electrode potential
(cathode)
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Determination of an electrode potential (anode)

• What is the oxidation potential at 25 oC of a
  platinum wire electrode immersed in an acidic
  solution of ferrous ions, at a concentration of
  0.50 m, and of ferric ions, at a concentration of
  0.25 m?
• The activity coefficient for the ferrous ion is
  0.435 and that for the ferric ion is 0.390 under
  these conditions.

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Determination of an electrode potential (anode)

• Pt Fe2 (c1, moles/L), Fe3 (c2, moles/L)
• Fe2 Fe3 e- written as an oxidation
• Fe3 e- Fe2 written as a reduction
• activity activity coefficient x concentration
• aFe2 0.435 (0.5) 0.218
• aFe3 0.390 (0.25) 0.0975

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Determination of an electrode potential (anode)
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emf Calculation

• Consider a cell consisting of the following
  half-cells
• CuCu2 (a0.1)
• ZnZn2 (a0.2)
• What will be oxidized?
• What will be reduced?
• Determine the emf of the cell.

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CuCu2 (a0.1)
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CuCu2 (a0.1)
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ZnZn2 (a0.2)
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emf calculation

• Consider a cell consisting of the following
  half-cells
• CuCu2 (a0.1) E 0.3074 volts
• ZnZn2 (a0.2) E -0.7837 volts
• The copper electrode has the higher reduction
  potential and will be reduced.
• The zinc will be oxidized.

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emf Calculation

• Convention is to write the oxidation (anode) on
  the left
• ZnZn2 (a0.2) Cu2 (a0.1)Cu
• Eleft (oxidation) 0.7837 volts
• Eright (reduction) 0.3074 volts
• Ecell 1.091 volts

Note If the cell is configured incorrectly, E
will be negative, corresponding to a positive ?G
and an indication that electrons will not flow
spontaneously in the direction predicted.
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Pharmaceuticals subject to oxidation

• Steroids
• Vitamins
• Antibiotics
• Epinephrine

These reactions are mediated either by free
radicals or by molecular oxygen.
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Preventing the Oxidation of Pharmaceuticals

• Effect of pH
• Antioxidants
• Free Radical Scavengers
• Chelating agents

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Oxidation

• Oxidation is loss of electrons
• An oxidizing agent must be able to take on
  electrons
• In organic chemistry, oxidation is synonymous
  with dehydrogenation (loss of hydrogen)
• The release of H means that the oxidation
  potential is sensitive to pH
• Oxidation often involves the addition of oxygen.
• Heavy metals catalyze oxidative deterioration

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Oxidation of Hydroquinone
OH
O
2 H 2e-
OH
O
Hydroquinone
Benzoquinone
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Increasing the activity of H decreases Oxidation
Potential
Lowering pH is an effective way to reduce
oxidation
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Antioxidants

• Most antioxidants function by providing
• electrons (reducing agents)
• labile H which will be accepted by any free
  radical to terminate the chain reaction
• Preferential antioxidants
• The antioxidant must be oxidized more easily than
  the drug
• The antioxidant acts as a reducing agent to
  consume the oxygen that is present

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Examples of Water Soluble Antioxidants

• Sodium bisulfite
• Ascorbic acid
• Sodium sulfite
• Sodium metabisulfite
• Cysteine hydrochloride
• Thioglycolic acid
• Sulfur dioxide

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Examples of Oil Soluble Antioxidants

• Ascorbyl palmitate
• Butylated hydroxyanisole (BHA)
• Butylated hydroxytoluene (BHT)
• Lecithin
• Propyl gallate
• ?-tocopherol

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Sulfite Ion as an Antioxidant

• SO42- 2 e- H2O SO32- 2OH-
• Reduction Potential Eo - 0.93 volt
• SO32- 2OH- SO42- 2 e- H2O
• Oxidation Potential Eo 0.93 volt
• Sulfite is useful as an antioxidant for drugs
  undergoing redox reactions with smaller positive
  oxidation potentials.
• If the drug has a higher oxidation potential than
  sulfite, sulfite will not work as an antioxidant.

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Oxidation of Ascorbic Acid
OH
HO
O
HO
H e-
HO
HO
HC
HC
O
O
O
O
CH2OH
CH2OH
Eo (oxidation potential) 0.383 volt at 25 oC
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Example 1

• Will Sulfite ion prevent the oxidation of
  Epinephrine?
• Based on Standard Oxidation Potentials
• Eo (Epinephrine) - 0.808 volt
• Eo (Sulfite ion) 0.93 volt
• Yes the sulfite ion is more easily ionized than
  epinephrine.

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Example 2

• Will Sulfite ion prevent the oxidation of
  Ascorbic Acid?
• Based on Standard Oxidation Potentials
• Eo (Ascorbic Acid) 0.383
• Eo (Sulfite ion) 0.93 volt
• Yes the sulfite ion is oxidized easier than the
  ascorbic acid.

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Qualification

• Note that the actual oxidation potentials are
  significantly influenced by temperature and the
  activities of the various species involved such
  that a conclusion based on the standard state
  oxidation potentials may not be valid.

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Example 2 (See Example 9-13 in Martins Physical
Pharmacy

• Will a solution containing 10-2 M sulfite ion
  prevent the oxidation of 10-3 M ascorbic acid at
  pH 7?
• Based on calculated Oxidation Potentials
• ED (Ascorbic Acid) 0.915
• ED (Sulfite ion) 0.67 volt
• No the ascorbic acid is oxidized easier than
  the sulfite ion.

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Oxidative decomposition

• Autoxidation
• Reaction of any material with molecular oxygen
• Free radical chain process
• Free radicals are produced by reactions involving
  homolytic bond fission of a covalent bond so that
  each atom or group of atoms involved retains one
  of the electrons of the original covalent bond
• A B A. B.
• CH3 CH3 2CH3.

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Free Radical Chain Reaction

• Initiation
• RH R. (H)
• Propagation
• R. O2 RO2.
• RO2. RH ROOH R.
• Hydroperoxide Decomposition
• ROOH RO. .OH
• Termination
• RO2. X inactive products
• RO2. RO2. inactive products

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Controlling Autoxidation

• Free radicals can generally be terminated by a
  free radical inhibitor
• Sodium metabisulfite
• Thiourea
• Cysteine hydrochloride

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Chelating agents

• Form complexes with trace amounts of heavy metal
  ions, inactivating the catalytic effect
• EDTA (ethylenediamine tetraacetic acid)
• Dihydroxyethyl glycine
• Citric acid
• Tartaric acid

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Stabilizing Pharmaceuticals subject to Oxidation

• Purge the system of oxygen / flush with nitrogen
• Buffer to lower pH
• Use preferential antioxidants
• Use free radical scavengers
• Use chelating agents