The document provides an introduction to key concepts in electrochemistry including oxidation/reduction reactions, oxidation numbers, and definitions of terms like oxidizing agent and reducing agent. It then discusses rules for assigning oxidation numbers, types of redox reactions like disproportionation, electrochemical cells, and how to determine the potential of a cell.

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Introduction
to
Electrochemistry

3. Oxidation/Reduction
• Oxidation
• Gain of oxygen atoms
• Loss of hydrogen atoms
• LOSS OF ELECTRONS!
• Reduction
• Loss of oxygen atoms
• Gain of hydrogen atoms
• GAIN OF ELECTRONS
• There must be both oxidation and reduction processes for a
reaction to occur
Electrons are
transferred, not
‘lost’ or ‘gained
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4. Definition of terms
• Oxidizing agent (oxidant)
• Causes another species to be oxidized
• it is reduced!
• Reducing agent (reductant)
• Causes another species to be reduced
• it is oxidized!
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5. Oxidation Number (Oxidation State)
• Is the number that is assigned to each kind
of atom in a compound or ion of an
element.
• This represents the number of electrons that
have been gained, lost or shared by the
species.
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6. Rules of Assigning Oxidation Numbers
1. Any uncombined element or compound of same element is
assigned Oxidation Number of zero e.g. O, K, H2.
2. For a compound, the sum of all oxidation number of all atoms is
zero
3. For polyatomic ions, the sum of all oxidation numbers of all atoms
is equal to the charge on the ion.
4. All monoatomic ions are assigned oxidation number equal to the
charge on their ions.
5. When oxygen is present in compound or ion, it usually has an
oxidation number of -2 (exception include peroxides in which
oxygen has oxidation number of -1)
6. Hydrogen usually has oxidation number of +1 except in metal
Hydrides where H is -1)
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7. Rules of Assigning Oxidation Numbers
• To determine the oxidation number of an element
in a compound, the following procedures can be
followed.
• Example: Determine the oxidation number of
chromium in potassium dichromate (K2Cr2O7)
Atom Oxidation
number
K +1
Cr Unknown
O -2
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8. Rules of Assigning Oxidation Numbers
• Multiply the oxidation number of each element by
appropriate subscript shown in the formula. Write
these total oxidation number below the
corresponding symbol in the formula.
Formula K2 Cr2 O7
Product of OS
and subscript
+1 × 2 𝑥 × 2 -2 × 7
Total oxidation
states
+2 2x -14
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9. Rules of Assigning Oxidation Numbers
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10. Oxidation-Reduction Reaction
• These are sometimes called redox reactions.
• Oxidation and reduction reactions go hand in hand.
E.g. Ag+ + Fe2+  Ag(s) + Fe3+
• The oxidation number of Ag+ changes from +1 to
0, thus Ag+ is reduced. The oxidation number of
Fe2+changes from +2 to +3 thus Fe2+is oxidized.
• When one substance is being oxidized, the other
must be reduced.
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11. Oxidation – Reduction half reactions
• Oxidation-reduction reaction can be split into two
half-reactions
• The half-reactions show which species gains
electrons and which loses them.
Zn(s) + Cu2+
(aq) → Zn2+
(aq) + Cu(s)
• The half reactions for the above redox reaction
are;
Oxidation: Zn(s) →Zn2+
(aq) + 2e-
• Reduction: Cu2+
(aq) + 2e- → Cu(s)
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12. Balancing Oxidation – Reduction
Equations
• The redox reactions must be balanced in terms of number
of atoms and charges.
• It is a challenge to balance them by inspection
• To simplify the balancing redox reaction is separated into
its reduction and oxidation half reactions which are
balanced separately and then added together to obtain the
balanced equation for the overall reaction.
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13. The general procedure for balancing the redox
reactions
Example - Balance the following redox reaction
SO3
2-
(aq) + MnO4
-
(aq) → SO4
2-
(aq) + Mn2+
Step 1:
 Identify the species being oxidized and the species being reduced from the changes
in their oxidation numbers. In this reaction, oxidation number of sulphur increase
from
+4 in SO3
2- to +6 in SO4
2- and that of Mn decreases from
+7 in MnO4
- to +2 in Mn2+.
Step 2:
 Write the two skeletal (unbalanced) equations for the oxidation and reduction half
reactions.
Oxidation: SO3
2-
(aq) → SO4
2-
(aq)
Reduction: MnO4
-
(aq) → Mn2+
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Step 3:
Balance each half equation atomically in this order:
Start with atoms other than H and O (for this equation the
other atoms are already balanced)
Balance O atoms by adding H2O with appropriate coefficients.
— Oxidation: SO3
2-
(aq) + H2O → SO4
2-
(aq)
— Reduction: MnO4
-
(aq) → Mn2+ + 4H2O
Balance hydrogen atoms by adding H+ with appropriate
coefficients.
— Oxidation: SO3
2-
(aq) + H2O → SO4
2-
(aq) + 2H+
— Reduction: MnO4
-
(aq) + 8H+→ Mn2+ + 4H2O

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Step 4:
• Balance each half equation “electrically”. Add the number
of necessary electrons to get the same electric charge on
both sides of each half equation.
Oxidation: SO3
2-
(aq) + H2O → SO4
2-
(aq) + 2H+ + 2e-
Reduction: MnO4
-
(aq) + 8H+ + 5e- → Mn2+ + 4H2O
Step 5:
• Multiply the half reactions by the simplest set of whole
numbers to balance the electrons.
Oxidation: 5SO3
2-
(aq) + 5H2O → 5SO4
2-
(aq) + 10H+ + 10e-
Reduction: 2MnO4
-
(aq) + 16H+ + 10e- → 2Mn2+ + 8H2O

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Step 6:
Cancel electrons and equal amounts of any substance
that appear on both side of the equation
Oxidation: 5SO3
2-
(aq) + 5H2O → 5SO4
2-
(aq) + 10H+ +
10e-
Reduction: 2MnO4
-
(aq) + 16H+ + 10e- → 2Mn2+ +
8H2O
Step 7:
Write the net equation by adding the two equations
5SO3
2-
(aq) + 5H2O + 2MnO4
-
(aq) + 16H+ → 5SO4
2-
(aq) + 10H+ + 2Mn2+ + 8H2O

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Step 8:
• Simplify the net equation so that the equation should not contain
the same species on both sides: Therefore, subtract 5H2O from
each side and 10H+ from each side.
5SO3
2-
(aq) + 2MnO4
-
(aq) + 6H+ → 5SO4
2-
(aq) + 2Mn2+ + 3H2O
Step 9: Verify if the equation is balanced.
Note:
• The reaction above is balanced in acidic solution.
If the reaction is carried out in basic solution,
OH- is added to both sides of the net equation in
step 8. Then the step 8 is repeated

18. Disproportionation Reaction
• In some redox reactions, the same substance is
both oxidized and reduced.
• These kinds of redox reactions are called
disproportionation reactions.
For example;
2H2O2(aq) → O2(g) + 2H2O
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19. Electrochemical Cells
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• Electrochemical cells are either galvanic or electrolytic.
 A galvanic cell or voltaic cell is an electrochemical cell
in which a spontaneous chemical reaction is used to
generate an electric current.
 Technically, a battery is a collection of galvanic cells
joined in series.
 An electrolytic cell is an electrochemical cell in which
electrical energy causes nonspontaneous redox reactions
to occur
 It is a result of incorporating an external power source,
such as a battery, in the circuit to drive the reaction to
nonspontaneous direction.

21. Standard Potentials
• Standard potentials are also called standard
electrode potentials.
• Since they are always written for the reduction
half reactions, they also are sometimes called
standard reduction potentials.
• Different half-reactions have different
tendencies to occur.
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• To compare their tendencies to occur, the following conventions
have been developed:
 Since the tendencies for half-reactions to proceed depend on the
temperature, the concentrations of the chemical species involved,
and, if gases are involved, the pressure in the half-cell, the
defined standard conditions are a temperature of 25 ˚C, a
concentration of exactly 1 M for all dissolved chemical species
involved, and a pressure of exactly 1 atm.
 Because every cell consists of two half-cells, it is not possible to
measure the potential directly.
 However, if the tendency of a certain half-reaction is assigned to
be zero, then the tendencies of all other half-reactions can be
determined relative to this reference half-reaction.
 For that reason the half-reaction 2H+ + 2e- → H2 is the reference
half-reaction with the standard reduction potential of 0.0000 V.

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Half-reactions 𝑬 𝒐
(Volts)
F2(g) + 2e-
→ 2F-
(aq) +2.87
Au+
(aq) + e-
→ Au(s) +1.69
Ce4+
(aq) + e-
→ Ce3+
(aq) +1.61
MnO4
-
(aq) + 8H+
+ 5e-
→ Mn2+
(aq) + 4H2O(l) +1.51
Cl2(g) + 2e-
→ 2Cl-
(aq) +1.36
Cr2O7
2-
(aq) + 14H+
(aq) + 6e-
→ 2Cr3+
(aq) + 7H2O(l) +1.33
Ag+
(aq) + e-
→ Ag(s) +0.80
Fe3+
(aq) + e-
→ Fe2+
(aq) +0.77
Cu2+
(aq) + 2e-
→ Cu(s) +0.34
2H+
(aq) + 2e-
→ H2(g) 0
Fe3+
(aq) + 3e-
→ Fe(s) -0.04
O2(g) + H2O(l) 2e-
→ HO2
-
(aq) + OH-
(aq) -0.08
Pb2+
(aq) + 2e-
→ Pb(s) -0.13
Sn2+
(aq) + 2e-
→ Sn(s) -0.14
Fe2+
(aq) + 2e-
→ Fe(s) -0.44
Zn2+
(aq) + 2e-
→ Zn(s) -0.76
Al3+
(aq) + 3e-
→ Al(s) -1.66
Mg2+
(aq) + 2e-
→ Mg(s) -2.36

24. Observations
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25. Determination of Cell Potential
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26. Determination of Cell Potential
• The general steps for determining this potential
are presented and illustrated for the following cell
equation.
Cu(s) + Ag+
(aq)→ Ag(s)+ Cu2+
(aq)
Step 1: Write the equations representing the half- reactions
as extracted from the overall reaction given and label as an
oxidation and a reduction.
Oxidation: Cu(s)→Cu2+
(aq)+ 2e-
Reduction: Ag+
(aq) + e-→ Ag(s)
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27. Determination of Cell Potential
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28. Determination of Cell Potential
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29. Determination of Cell Potential
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30. Any Questions or Additions

31. 2/26/2020 31
Electrochemistry
&
Redox

32. Electrochemistry and Redox
• Oxidation-reduction: “Redox”
 electron transfer processes
 loss of 1 or more e-
• Oxidation numbers: imaginary charges (Balancing
redox reactions)
• Electrochemistry: study of the interchange between
chemical change and electrical work
• Electrochemical cells: systems utilizing a redox
reaction to produce or use electrical energy

33. Oxidation Numbers (O.N.)
1.Pure element O.N. is zero
2.Monatomic ion O.N. is charge
3.Neutral compound: sum of O.N. is zero
Polyatomic ion: sum of O.N. is ion’s charge
*Negative O.N. generally assigned to more
electronegative element

34. Oxidation Numbers (O.N.)
4.Hydrogen
assigned +1
(metal hydrides, -1)
5.Oxygen
assigned -2
(peroxides, -1; OF2, +2)
6.Fluorine
always -1

35. Oxidation-reduction
Oxidation is loss of e-
O.N. increases (more positive)
Reduction is gain of e-
O.N. decreases (more negative)
Oxidation involves loss OIL
Reduction involves gain RIG

36. Redox
Oxidation is loss of e-
causes reduction
“reducing agent”
Reduction is gain of e-
causes oxidation
“oxidizing agent”

38. Balancing Redox Reactions
1. Write separate equations (half-reactions) for
oxidation and reduction
2. For each half-reaction
a. Balance elements involved in e- transfer
b. Balance number e- lost and gained
3.To balance e-
multiply each half-reaction by whole numbers

39. Balancing Redox Reactions: Acidic
4.Add half-reactions/cancel like terms (e-)
5. Acidic conditions:
Balance oxygen using H2O
Balance hydrogen using H+
Basic conditions:
Balance oxygen using OH-
Balance hydrogen using H2O
6.Check that all atoms and charges balance

40. Examples
Acidic conditions:
Basic conditions:

  3
(aq)
2
(aq)
2
(aq)
-
4(aq) FeMnFeMnO acid

  2(aq)2(g)(aq)(s) Ag(CN)OCNAg base

41. Types of cells
Voltaic (galvanic) cells:
a spontaneous reaction generates electrical energy
Electrolytic cells:
absorb free energy from an electrical source to
drive a nonspontaneous reaction

42. Common Components
Electrodes:
conduct electricity between cell and
surroundings
Electrolyte:
mixture of ions involved in reaction or
carrying charge
Salt bridge:
completes circuit (provides charge balance)

43. Electrodes
Anode:
Oxidation occurs at the anode
Cathode:
Reduction occurs at the cathode
Active electrodes: participate in redox
Inactive: sites of ox. and red.

44. Voltaic (Galvanic) Cells
A device in which chemical energy
is changed to electrical energy.
Uses a spontaneous reaction.

45. 17_360
Porous disk
Reducing
agent
Oxidizing
agent
e –
e –
e – e –
e –
e –
CathodeAnode (b)(a)
Oxidation Reduction

49. Zn2+
(aq) + Cu(s)  Cu2+
(aq) + Zn(s)
Zn gives up electrons to Cu
— “pushes harder” on e-
— greater potential energy
— greater “electrical potential”
Spontaneous reaction due to
— relative difference in metals’ abilities to give e-
— ability of e- to flow

50. Cell Potential
Cell Potential / Electromotive Force (EMF):
The “pull” or driving force on electrons
Measured voltage (potential difference)
V
C
J
movedchargeofunit
energypotentialelectricalorwork
Ecell 

51. 17_363
e–
e– e–
e–
Zn 2+
SO4
2–
Zn(s)
1.0 M Zn 2+
solution
Anode
1.0 M Cu 2+
solution
Cathode
Cu 2+
SO4
2–
Cu(s)
Ecell = +1.10 V

52. Cell Potential, E0
cell
E0
cell
cell potential under standard conditions
elements in standard states (298 K)
solutions: 1 M
gases: 1 atm

54. Standard Reduction Potentials
E0 values for reduction half-reactions with
solutes at 1M and gases at 1 atm
Cu2+ + 2e  Cu
E0 = 0.34 V vs. SHE
SO4
2 + 4H+ + 2e  H2SO3 + H2O
E0 = 0.20 V vs. SHE

57. Any Questions or Additions

58. Electrochemistry
&
Pharmaceutical Applications

59. Industrial Pharmacy & Health-related use
of Electrolytic Processes
• At all stages of the development of electrochemistry, an intimate connection
existed between the development of theoretical concepts and the discovery
of solutions for a practical application of electrochemical processes and
phenomena.
• Theoretical investigations have been stimulated by the practical use of
various electrochemical phenomena and processes, and the theoretical
concepts that were developed have in turn contributed significantly to the
development of applied electrochemistry.
• Today, applied electrochemistry is of great value for the economy especially
in industrial pharmaceutical fields.
• Electrochemical phenomena and processes are useful for the quantitative
and qualitative chemical analysis of various substances and media, including
liquids, gases, and solids.
• The high accuracy of the electrochemical methods of analysis derives from
the fact that they are based on highly exact laws (e.g., those of Faraday).

60. • The methods of electrochemical analysis are instrumental.
• It is very convenient that electrical signals are used for the
perturbation: current, potential, and so on, and that the result (the
response) again is obtained as an electrical signal.
• This is the basis for the high speed and accuracy of the readings,
for the extensive possibilities of automated recording of the results,
as well as for automation of the entire analysis.
• Electrochemical methods of analysis are distinguished by their
high sensitivity, selectivity (the possibility of analyzing certain
substances in the presence of others), speed of the measurements,
and other advantages.
• In many cases extremely small volumes, less than 1 mL, of the test
solution will suffice for electrochemical analysis.

61. The following are the major groups of
electrochemical methods for chemical analysis:
1. Conductometry, which measures the electrical conductivity of the electrolyte
solution being examined
 nonselective method of analysis; all types of mobile ion present in the
solution (or other medium being examined) contribute to conductivity
 primarily useful when determining the concentrations in binary electrolyte
solutions (e.g., for determining the solubilities of poorly soluble
compounds)
 used in particular for the titration of acids with base (and vice versa) in
colored and turbid solutions or solutions containing reducing and oxidizing
agents
 Conductometric analysis is performed in both concentrated and dilute
solutions. Whose accuracy depends solution system that can monolayer,
binary or multicomponent system with corresponding accuracy

62. 2. Coulometry, which measures the amount of charge Q
consumed for the complete conversion (oxidation or
reduction) of the substance being examined
 regarded as an analog of titration where the substance being
examined is quantitatively converted to a reaction product not by
the addition of titrant, but by a certain amount of electric charge Q
 Electrochemical coulometers are based on the laws of Faraday;
with them the volume of gas or mercury liberated, which is
proportional to charge, is measured.
 For coulometric analysis, the substance being examined must react
in 100% current yields [i.e., other (secondary) reactions must be
entirely absent].

63. 3. Voltammetry, which determines the steady-state or
transient polarization characteristics of electrodes in
reactions involving the substance being examined
 In the transient voltammetric methods, one measures the
characteristic parameters on transient polarization curves after
some potential or current perturbation has been applied to the
electrode.
 Many versions of transient methods of voltammetric analysis
using single or repetitive potential or current signals with different
shapes and amplitude have been described.

64. 4. Potentiometry, which measures the open-circuit equilibrium
potential of an indicator electrode, for which the substance
being examined is potential determining
 suitable for the analysis of substances for which
electrochemical equilibrium is established at a suitable
indicator electrode at zero current
 An important condition for potentiometry is high selectivity;
the electrode’s potential should respond only to the substance
being examined, not to other components in the solution.
 This condition greatly restricts the possibilities of the version
of potentiometry described here when metal electrodes are
used as the indicator electrodes. The solution should be free of
ions of more electropositive metals and of the components of
other redox systems (in particular, dissolved air).

65. 5. The other pharmaceutical related uses of
electrochemistry include the following:
 Purification procedures – though not significant but
takes place such as purification of water by
electroflotation displays highly valuable
pharmaceutical use
 Medical applications - electrochemistry is able to
model a number of processes occurring in living
organisms, which in turn has led to progress in
fundamental medical science leading to widely use in
diagnosing various diseases.

66. Any Questions or Additions

67. THANK YOU

68.  Define the following terms:
[Oxidation, reduction, oxidant, reductant, redox, electrochemistry,
monatomic, polyatomic, anode, cathode, spontaneity, Conductometry,
Coulometry, Voltammetry, potentiometry, etc]
Respond to the following questions:
Give a detailed account of redox process in electrochemistry
Explain in details the rules of assigning oxidation number with chemical examples
With examples, illustrate three different oxidation-reduction chemical reaction
processes
With examples, explain the nine (9) principle steps that are considered for balancing
the redox reactions
Give a detailed account of standard potential as expressed in electrochemical reactions
With examples, explain the four (4) principle steps that are considered for the
determination of cell potential
In general, give an illustrated account of practical applications of electrochemical
procedures .
With particular reference to pharmaceutical applications, state and explain some of the
uses of electrochemical processes

69. Group work discussional questions:
With examples, illustrate three different oxidation-reduction
chemical reaction processes that are considered explaining
principle steps that are considered for balancing the redox
reactions
Give a descriptive account of different forms of interface types,
tensions, forces that affects material substances interactions
Illustrate the full understanding standard potential as it is used
in explaining the principle steps that are considered for the
determination of cell potential
With particular reference to pharmaceutical applications, state
and explain some of the uses of electrochemical processes