1. Electrochemistry involves studying chemical reactions that produce electricity or using electricity to cause non-spontaneous reactions. 2. Key concepts include galvanic cells which generate electricity from spontaneous redox reactions, electrolytic cells which use electricity to drive non-spontaneous reactions, and standard reduction potentials which quantify reaction tendencies. 3. Standard cell potentials allow calculation of cell voltage from half-reaction potentials under standard conditions.

1. Electrochemistry
“ Electrochemical methods”
References:
1- Analytical Chemistry, Christian, Gary D
2- Fundamentals of Analytical Chemistry, Skoog
Email: rajkailany@gmail.com

2. Introduction
Some qualitative and quantitative information can be obtained
by measuring one or more electrical properties.
When a solution is a part of an electrochemical cell, the cell
voltage, the current passing through it and its resistance
depends on the chemical structure of the solution.
As the name suggests, electrochemistry is the study of
changes that cause electrons to move. This movement of
electrons is called electricity.

3. Electroanalytical methods include two parts:
*Direct analysis methods where the electrical
properties of the sample solution as well as a standard
solution are measured.
*Indirect methods by following the electrical property during a
titration to find the end point
In electrochemistry, electricity can be generated by
movements of electrons from one element to another in a
reaction known as a redox reaction or oxidation-
reduction reaction
In Redox reactions, the oxidation number “The charge that an
atom (or group of atoms) have” changes.

4. Oxidation number
The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred.
1. Free elements (uncombined state) have an oxidation
number of zero.
Na, Be, K, Pb, H2, O2, P4, S8 = 0
2. In monatomic ions, the oxidation number is equal to
the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
3. The oxidation number of oxygen is usually –2. In H2O2
and O2
2- it is –1.

5. 4. The oxidation number of hydrogen is +1 except when
it is bonded to metals in binary compounds. In these
cases, its oxidation number is –1.
6. The sum of the oxidation numbers of all the atoms in a
molecule or ion is equal to the charge on the
molecule or ion.
5. Group IA metals are +1, IIA metals are +2 and fluorine is
always –1.
HCO3
-
O = -2 H = +1
3x(-2) + 1 + ? = -1
C = +4
Oxidation numbers of all
the atoms in HCO3
- ?

6. Redox Reactions
 A substance is reduced when it gains electrons from
another substance. GER
- gain of e- net decrease in charge of species
- Oxidizing agent (oxidant)
1- Redox Reaction involves transfer of electrons from one
species to another.
 A substance is oxidized when it loses electrons to
another substance. LEO
- loss of e- net increase in charge of species
- Reducing agent (reductant)
Fe3+ + Cu+ → Cu2 + + Fe2+

7. (Reduction)
(Oxidation)
Oxidizing
Agent
Reducing
Agent
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:

8. Balancing Redox Equations
1. Write the unbalanced equation for the reaction ion ionic form.
Fe2+ + Cr2O7
2- Fe3+ + Cr3+ ( acidic)
2. Separate the equation into two half-reactions.
3. Balance the atoms other than O and H in each half-reaction.
MnO4
- + I- MnO2 + IO3
- (basic)

9. Balancing Redox Equations
4. For reactions in acid, add H2O to balance O atoms and H+ to
balance H atoms.
5. Add electrons to one side of each half-reaction to balance the
charges on the half-reaction.
6. If necessary, equalize the number of electrons in the two half-
reactions by multiplying the half-reactions by appropriate
coefficients.
19.1

10. Balancing Redox Equations
7. Add the two half-reactions together and balance the final
equation by inspection. The number of electrons on both
sides must cancel.
8. Verify that the number of atoms and the charges are balanced.
19.1
9. For reactions in basic solutions, add OH- to both sides of the
equation for every H+ that appears in the final equation.

11. Introduction to Electrochemistry
The relation between chemical reactions and electricity.
“ Electrochemical cell”
Certain chemical reactions can create electricity.
“ Galvanic or Voltaic cell”
“ Spontaneous”
Electricity can make certain chemical reaction happen that
would not happen otherwise.
“ Electrolytic cell”
“Not Spontaneous”

12. Introduction to Electrochemistry
Units:
1.) Electric Charge (q)
 Measured in coulombs (C)
 Faraday constant (F) is the charge of a mole of electrons
2.) Electric current ( I )
 Quantity of charge flowing each second
through a circuit.
Unit of current is Ampere: 1A = (C/sec)
F
n
q 

Relation between
charge and moles:
Coulombs moles

e
mol
Coulombs

13. Introduction to Electrochemistry
3.) 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
- 1 V = J / C
Higher potential difference requires
more work to lift water (electrons) to
higher trough

14. Introduction to Electrochemistry
3.) Electric Potential (E)
 Combining definition of electrical charge and potential
q
E
w ork
G 




 F
n
q 

n F E
G 


Relation between free energy
difference and electric potential
difference:
Describes the voltage that can be generated by a chemical reaction

15. Introduction to Electrochemistry
4.) Ohm’s Law
Current (I ) is directly proportional to the potential difference (voltage)
across a circuit and inversely proportional to the resistance (R)
- I = E / R
5.) Resistance ( R) Ohm (W) R = E / I 1 ohm = 1V/1A
6.) Conductance (1/ R) mho
7.) Power (P) Watt ( W)
 It is the work done per unit time
- Units: joules per second W = J/S
We can measure the change in one (or more) of these
electrical properties in an electrochemical cell

16. Electrochemical cell
Components of electrochemical cell:
1- Two half cells
2- Two metal electrodes
3- One voltmeter
4- One salt bridge
5- Two aqueous solutions for each half cell

17. Cell Potential:
When a metal is immersed in a solution containing its
ions, a difference in potential between the metal and the
solution arises due to the tendency of the atoms of the
metal to move to the solution as ions by losing electrons.
M0 → M+ + e-
The former system is called a half-cell. The ion's tendency
to lose or gain electrons is known as electrode potential
In the electrochemical cell, the electrode potential is not
measured directly, but the potential difference between the
two electrodes, which is called cell potential, is measured.
Ecell = Ecathode − Eanode

19. Galvanic Cells
spontaneous
redox reaction
anode
oxidation
cathode
reduction

20. Galvanic Cells
The difference in electrical potential
between the anode and cathode is
called:
• Electromotive force emf
• Cell Voltage
• Cell potential
Cell Notation
Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)
[Cu2+] = 1 M & [Zn2+] = 1 M
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
Anode │ solution │ Cathode

21. Cell Notation
same example: Zn+2 & Cu2+ at 1.0 M
Zn + Cu2+ -----) Zn2+ + Cu
Zn | Zn2+ (1.0 M) || Cu2+(1.0 M) | Cu
Anode (-) (oxidation) Zn ---) Zn2+ + 2e-
Cathode (+) (reduction) Cu2+ + 2e- ---) Cu
Know the following Conventions:
Anode written on left
Cathode written on right
| = phase boundary
|| = salt bridge
Zn = anode electrode
Cu = cathode electrode
, = separates different ions in the same half cell
M = Molar concentration of ions
This battery develops 1.1 volts at 25 oC
Anode │ solution │ Cathode

22. Standard Reduction Potentials
Standard Reduction Potentials (E0): is the voltage
associated with a reduction reaction at an electrode when
all solutes are 1 M and all gases are at 1 atm.
E0 = 0 V
Standard hydrogen electrode (SHE)
2e- + 2H+ (1 M) H2 (1 atm)
Reduction Reaction

23. Standard Reduction Potentials
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
2e- + 2H+ (1 M) H2 (1 atm)
Zn (s) Zn2+ (1 M) + 2e-
Anode (oxidation):
Cathode (reduction):
Zn (s) + 2H+ (1 M) Zn2+ + H2 (1 atm)

24. E0 = 0.76 V
cell
Standard emf (E0 )
cell
0.76 V = 0 - EZn /Zn
0 2+
EZn /Zn = -0.76 V
0 2+
Zn2+ (1 M) + 2e- Zn E0 = -0.76 V
E0 = EH /H - EZn /Zn
cell
0 0
+ 2+
2
Standard Reduction Potentials
E0 = Ecathode - Eanode
cell
0 0
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)

25. Standard Reduction Potentials
Pt (s) | H2 (1 atm) | H+ (1 M) || Cu2+ (1 M) | Cu (s)
2e- + Cu2+ (1 M) Cu (s)
H2 (1 atm) 2H+ (1 M) + 2e-
Anode (oxidation):
Cathode (reduction):
H2 (1 atm) + Cu2+ (1 M) Cu (s) + 2H+ (1 M)
E0 = Ecathode - Eanode
cell
0 0
E0 = 0.34 V
cell
Ecell = ECu /Cu – EH /H
2+ +
2
0 0 0
0.34 = ECu /Cu - 0
0 2+
ECu /Cu = 0.34 V
2+
0

26. • E0 is for the reaction as
written
• The more positive E0 the
greater the tendency for the
substance to be reduced
• The half-cell reactions are
reversible
• The sign of E0 changes
when the reaction is
reversed
• Changing the stoichiometric
coefficients of a half-cell
reaction does not change
the value of E0

27. What is the standard emf of an electrochemical cell made
of a Cd electrode in a 1.0 M Cd(NO3)2 solution and a Cr
electrode in a 1.0 M Cr(NO3)3 solution?
Cd2+ (aq) + 2e- Cd (s) E0 = -0.40 V
Cr3+ (aq) + 3e- Cr (s) E0 = -0.74 V
Cd is the stronger oxidizer
Cd will oxidize Cr
2e- + Cd2+ (1 M) Cd (s)
Cr (s) Cr3+ (1 M) + 3e-
Anode (oxidation):
Cathode (reduction):
2Cr (s) + 3Cd2+ (1 M) 3Cd (s) + 2Cr3+ (1 M)
x 2
x 3
E0 = Ecathode - Eanode
cell
0 0
E0 = -0.40 – (-0.74)
cell
E0 = _______ V
cell

28. Galvanic vs Electrolytic

29. Galvanic vs Electrolytic
Galvanic
1. In galvanic cell, electrical energy
is produced.
2. In galvanic cell, reaction taking place
is spontaneous.
3. The two half cells are set up in different
containers and are connected through
salt bridge or porous partition.
4. In galvanic cell, anode is negative
and cathode is positive.
5. The electrons move from anode to
cathode in external circuit
Electrolytic
1. In electrolytic cell, electrical energy
is consumed.
2. In electrolytic cell, reaction taking place
is nonspontaneous.
3. Both the electrodes are placed in the solution
or molten electrolyte in the same container.
4. In electrolytic cell, the anode is positive
and cathode is negative.
5. The electrons are supplied by the external
source. They enter through cathode and
come out through anode

30. Electrolysis of Water
Electrolysis is the process in which electrical energy is used
to cause a nonspontaneous chemical reaction to occur.

31. Spontaneity of Redox Reactions
G = -nFEcell
G0 = -nFEcell
0
n = number of moles of electrons in reaction
F = 96,500
J
V • mol
= 96,500 C/mol
G0 = -RT ln K = -nFEcell
0
Ecell
0 =
RT
nF
ln K
(8.314 J/K•mol)(298 K)
n (96,500 J/V•mol)
ln K
=
=
0.0257 V
n
ln K
Ecell
0
=
0.0592 V
n
log K
Ecell
0
ln K = 2.303 log K

32. Spontaneity of Redox Reactions
G0 = -RT ln K = -nFEcell
0

33. 2e- + Fe2+ Fe
2Ag 2Ag+ + 2e-
Oxidation:
Reduction:
What is the equilibrium constant for the following reaction
at 250C? Fe2+ (aq) + 2Ag (s) Fe (s) + 2Ag+ (aq)
=
0.0257 V
n
ln K
Ecell
0
E0 = -0.44 – (0.80)
E0 = __-1.24____ V
0.0257 V
x n
E0
cell
exp
K =
n = 2
0.0257 V
x 2
-1.24 V
= exp
K = __1.23 x 10 -42___
E0 = EFe /Fe – EAg /Ag
0 0
2+ +

34. The Effect of Concentration on Cell Emf
G = G0 + RT ln Q G = -nFE G0 = -nFE 0
-nFE = -nFE0 + RT ln Q
E = E0 - ln Q
RT
nF
Nernst equation
At 298
-
0.0257 V
n
ln Q
E0
E =
-
0.0592 V
n
log Q
E0
E =
-
0.0592
n
log Q
E0
E =

36. Will the following reaction occur spontaneously at 250C if
[Fe2+] = 0.60 M and [Cd2+] = 0.010 M?
Fe2+ (aq) + Cd (s) Fe (s) + Cd2+ (aq)
2e- + Fe2+ Fe
Cd Cd2+ + 2e-
Oxidation:
Reduction:
n = 2
E0 = -0.44 – (-0.40)
E0 = __-0.04__ V
E0 = EFe /Fe – ECd /Cd
0 0
2+ 2+
-
0.0257
n
ln Q
E0
E =
-
0.0257
2
ln
-0.04 V
E =
0.010
0.60
E = __0.0126__ V
E > 0 _________________

37. ELECTROANALYTICAL METHODS
Based on the electrical properties of metals and their ionic solutions,
several analytical techniques have been developed. For example:
Potentiometry, Coductometry, Polarography,…..
Potentiometry
Potentiometry is based on the
measurement of the potential
of an electrode system (e.g.
electrochemical cell).
Potentiometric measurement
system consists of two
electrodes called reference
and indicator electrode,
potentiometer and a solution
of analyte.

38. Potentiometric Analysis
• Based on potential measurement of electrochemical
cells without any appreciable current
• The use of electrodes to measure voltages from
chemical reactions
Components of a Potentiometric Cell
1. Reference electrode
2. Salt bridge
3. Analyte
4. Indicator electrode
RE SB A IE
Eref Ej Eind

39. Reference electrode
• Half-cell with known potential (Eref)
• Left hand electrode (by convention)
• Easily assembled
• Insensitive to analyte concentration
▫ Reversible and obeys Nernst equation
▫ Constant potential

40. Indicator electrode
• Generates a potential (Eind) that
depends on analyte concentration
• Selective
• Rapid and reproducible response
Salt bridge
• Prevents mixing up of analyte components
• Generates potential (Ej) = negligible

42. A- Reference electrode
Reference electrode must:
1. Have a constant potential
2. Its potential must be definite
To express any electrode we have to mention:
1. Redox reaction at the electrode surface.
2. Half cell and Nernst equation.
3. Sketch of its design.
4. Any necessary conditions for its preparation.
5. Any necessary precautions for its use.

43. Reference Electrodes
1. Standard Hydrogen Electrode
2. Calomel Reference Electrode
3. Silver/Silver Chloride Reference
Electrode

44. It’s a primary reference electrode.
Its potential is considered to
be zero.
Electrode reaction:
Eo = zero
Limitation:
The SHE is rarely used because it
is difficult to prepare and
inconvenient to use.
Standard Hydrogen Electrode

45. Calomel is the common name for the
compound Hg2Cl2.
electrode reaction in calomel half-cell
Eo = + 0.268V
E = Eo – (0.05916/2) log[Cl–]2
Temperature dependent
At 25 C , E = 0.244 V
Saturated calomel electrode (S.C.E.)

46. Ag(s) | AgCl (sat’d), KCl (xM) | |
AgCl(s) + e = Ag(s) + Cl–
Eo = +0.244V
E = Eo – (0.05916/1) log [Cl–]
E (saturated KCl) = + 0.197 V (25oC)
Silver-silver chloride electrode

47. Disadvantage of silver-silver chloride
electrode
• It is more difficult to prepare than SCE.
• AgCI in the electrode has large solubility in
saturated KCl
Advantage of Ag/AgCI electrodes over SCE.
• It has better thermal stability.
• Less toxicity and environmental problems with
consequent cleanup and disposal difficulties.

48. Ecell=Eindicator-Ereference

50. Liquid Junction Potential
• Liquid junction - interface between two solutions
containing different electrolytes or different
concentrations of the same electrolyte
• A junction potential occurs at every liquid junction.
Caused by unequal mobilities of the + and -
ions.
Saturated KCl is normally used.

51. B- Indicator electrode
Its potential is sensitive to the concentration of analyte
Ecell=Eindicator-Ereference
It must be:
(a) give a rapid response and
(b) its response must be reproducible.
1- Metallic electrodes: where the redox reaction takes
place at the electrode surface.
2- Membrane (specific or ion selective) electrodes:
where charge exchange takes place at a specific
surfaces and as a result a potential is developed.

52. a- Electrodes of the First Kind
• Pure metal electrode in direct equilibrium with its
cation
• Metal is in contact with a solution containing its
cation.
M+n(aq) + ne-  M(s)
Metallic electrodes:

53. a- Electrodes of the First Kind:
• The electrode system can be represented by M/Mn+, in which
the line represents an electrode – solution interface . For silver
electrode , we have
• Ag/Ag+
• the half – reaction is
• Ag+ + e = Ag(s) Eo = + 0.800V
E = 0.800 – (0.05916/1) log {1/[Ag+]}
Metallic electrodes

54. Disadvantages of First Kind Electrodes
• Not very selective
– Ag+ interferes with Cu+2
• May be pH dependent
– Zn and Cd dissolve in acidic solutions
• Easily oxidized (deaeration required)
• Non-reproducible response

55. b) Electrodes of the Second Kind
• The general form of this type electrode is M/MX/Xn- , where
MX is a slightly soluble salt. An example is the silver – silver
chloride electrode ;
• Ag/AgCl(s)/Cl-
It responds to the change in negative ion concentration
The Nernst expression for this process at25°C is

56. c- Inert Metallic (Redox) Electrodes
• Inert conductors that respond to redox systems
• An inert metal in contact with a solution
containing the soluble oxidized and reduced
forms of the redox half-reaction.
• May not be reversible
• Examples:
– Pt, Au, Pd, C

57. Example:
Ag(s) | AgCl[sat’d], KCl[xM] | | Fe2+,Fe3+) | Pt
Fe3++e = Fe2+ Eo = +0.770V
Ecell = Eindicator – Ereference
= {0.770 – (0.05916/1) log [Fe2+]/[Fe3+]}
– {0.222 – (0.05916/1) log [Cl–]}
C. Electrodes of the third Kind Inert electrodes (Indicators electrodes for
redox reaction)
The inert metal used is usually platinum
The inert metal is contact with a solution containing the soluble
oxidized and reduced form of the redox half – reaction. “(Fe2+,Fe3+) “
And Ag/AgCl is used as ref. electrode:

58. 1-They are sensitive electrodes to certain ions depending on the
type of membrane material.
2- This sensitivity is produced by a difference in voltage through a
membrane separating two solutions with different concentrations of
the same ion, the first solution with a constant concentration (the
reference solution) and the second solution is the sample solution.
3- The potential developed at this type of electrode results from an
unequal charge buildup at opposing surface of a special
membrane. The charge at each surface is governed by the position
of an equilibrium involving analyte ions, which, in turn, depends on
the concentration of those ions in the solution.
2) Membrane indicator electrodes

59. 4- There are many types of membranes, each of which is sensitive
to a specific type of ions made from different materials depending
on the type of ions to be estimated.
5- Consist of a thin membrane separating two solutions of different
ion concentrations. Most common: pH Glass electrode
6- The concentration of ions using these electrodesis expressed in
the form of the p function (eg. pH, pCa, pLi, ..)
2) Membrane indicator electrodes

60. Glass Membrane Electrode ( pH electrode):
• It is a hydrogen ion sensitive electrode (used to measure pH).
• When a thin glass membrane is placed between two acid solutions of
different conc., a potential difference develops across the glass
membrane and a hydrogen half cell is created.
• The cell contains a glass membrane electrode (indicator
electrode) and the saturated calomel electrode (reference
electrode).
a- separeated:
b- combined:

61. A complete cell can be represented by

62. Glass Membrane Electrode
E = K + 0.059 (pH1 - pH2)
K= constant known by the asymmetry
potential.
PH1 = pH of the internal solution 1.
PH2 = pH of the external solution 2.
The final equation is:
E = L - 0.059 pH
Ex:????
•Asymmetry potential
•E of the 2 reference electrodes
•pH of the internal solution
•Liquid junction potential

63. Properties of Glass pH electrode
• Potential not affected by the presence of oxidizing or
reducing agents
• Operates over a wide pH range
• Fast response
• Functions well in physiological systems
• Very selective
• Long lifespan

64. • Due to the electrode ability to respond to other ions such
as Na +, K +, at lower concentration of the hydrogen.
• Such error is due to the capability of the membrane for
responding to other cations such as Na and K besides the
hydrogen ion. As the hydrogen ion activity becomes very
small, these other ions can compete successfully in the
potential-determining mechanism.
• The error is negligible at pH less than about 9; but at pH
values above this, the glass electrode senses other
cations besides H+. this leads to negative errors (pH is
less than the real value)
Alkaline error:

65. • At very low pH values (pH < 1), the gel layer of the pH-
sensitive glass membrane absorbs acid molecules. This
absorption decreases the activity of hydrogen ions and
results in a lower potential at the outer membrane phase
boundary.
• The pH measurement therefore shows a higher pH value
than the actual pH value of the sample solution.
Acid error:

66. Application of potentiometry
1. Direct potentiometric measurements:
E = L - 0.059 pH
2. potentiometric titration:
• Potentiometric titration is a volumetric method in which
the potential between two electrodes (referent and
indicator electrode) is measured as a function of the
added reagent volume.
• Involves measurement of the potential of a suitable
indicator electrode as a function of titrant volume
• The part of the curve that has the maximum change
marks the equivalence point of the titration.

67. Potentiometric titration
It is used for all types of volumetric analysis: acid
base, precipitimetry, complexometry and redox
It is used when it is not easy or impossible to detect
the end point by ordinary visual methods i.e:
1. For highly coloured or turbid solutions.
2. For very dilute solutions 10-3, 10-6 M.
3. When there is no available indicator .

68. Potentiometric titration

69. Titration Curves

70. a -Titration curve. b- 1st derivative.
c- 2 nd derivative. d- Gran plot.
End Point Determination