HCM INTERNATIONAL UNIVERSITY
SCHOOL OF BIOTECHNOLOGY
Course: ANALYTICAL CHEMISTRY
Lecturer: Dr. NGUYEN TUAN KHOI
MEMBERS OF GROUP
T ôn Th ị H ồng Th ảo
Ph ạm Nguyễn Huệ Nh ân
Nguy ễn Vi ệt Th ư
Thái Văn Ch í
Nguy ễn Ng ọc Y ên Nhi
Nguyễn Th ị Phương Thùy
Đ oàn T ây Nguy ên
Nguy ễn Duy Trung
Nguy ễn Đ ức Thanh Long
Trần Đỗ Ngọc Oanh
V õ Hoàng Lâm
Nguyễn Th ị Thu Cúc
Trương Thị Ngọc Nhi
V ũ Ng ọc C ư ơng
Lê Trần Khánh Trang
Nguyễn Vũ Nh ất Th ịnh
Đỗ Vân Khanh
Outline
I. Introduction
II. Electrochemical cells
a.
b.
Galvanic cells
Electrolytic cells
III. Electrochemical cell applications
a.
b.
c.
Battery
Corrosion
Electrolysis
IV. Electrochemical methods
a.
Nernst equation
b.
Potentiometry
c.
Coulometry
d.
Voltammetry
I. INTRODUCTION
• Electrochemistry is the study of reactions in which
charged particles (ions or electrons) cross the interface
between two phases of matter, typically a metallic
phase (the electrode) and a conductive solution, or
electrolyte. This reaction is simple oxidation-reduction
process.
I. INTRODUCTION
Redox reaction
(reduction-oxidation reactions)
• Are reactions that mention to the transfer of
electrons between species.
• Describe all chemical reactions in which atoms
have oxidation number change.
Ox1 + red2  red1 + ox2
II.
Electrochemical
cells
a.
Galvanic
cells
b.
Electrolytic
cells
II. Electrochemical cell
• Transform energy from chemical reaction to
electrical energy or vice versa.
• An electrochemical cell includes:
– Two electrodes: half redox reactions occur
o Anode: oxidation reaction occur
o Cathode: reduction reaction occur
– Electrolyte solution(s)
II. Electrochemical cell
• Conditions for generating electricity flow:
̶ The electrodes must be externally connected by a
metal wire to permit electron flow.
̶ The electrolyte solutions are in contact to allow
movement of ions.
II. Electrochemical cell
• There are two types of electrochemical cells:
– Galvanic cells (or Voltaic cells): spontaneous
reactions occur.
– Electrolytic cells: nonspontaneous reactions occur
(electrical energy supply).
a. GALVANIC CELL
What about the sign of
the electrodes?
What about half-cell
reactions?
-
+
cathode half-cell
Cu+2 + 2e-  Cu
Cu
plates out
or deposits
on
electrode
Why?
anode half-cell
Zn  Zn+2 + 2e-
Cu
1.0 M CuSO4
What
happened at
each
electrode?
Zn electrode
erodes
or dissolves
Zn
1.0 M ZnSO4
a. Galvanic cells
HALF REACTION
Anode (Ox) :
Zn(s) = Zn2+ + 2e
Cathode (Red) : Cu2+ + 2e = Cu (s)
Net reaction : Zn (s) + Cu2+ = Zn2+ +Cu (s)
Zn (s) + Cu2+ (aq)
Cu (s) + Zn2+ (aq)
Salt bridge
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
cathode
anode
b. ELECTROLYTIC CELL
CELL POTENTIAL
• cell potential: Electrons flow from one
electrode to the other in one direction, there is
a potential difference between the electrodes.
• Cell potential is calculated in voltage (V) by the
formula:
Cell potential (E cell) = E cathode – E anode
E cathode: reduction potential (V)
 E anode: oxidation potential (V)
Baterry
Electrochemical
cell applications
Corrosion
Electrolysis
Energy
storage
Fuel
cells
Applications
of Galvanic
and
electrolytic
cells
PRIMARY BATTERY
Primary battery has long been known as dry-cell. It
cannot be recharged. It’s widely used to power
flashlight and some other similar devices.
The first practical battery consisted of a stack of small
electrical cell, each consisting of a silver plate and a
zinc plate separated by a sheet of cardboard which
had been soaked in salt water
A TYPICAL STRUCTURE OF
A PRIMARY BATTERY
The electrode reactions are
Zn → Zn2+ + 2e–
2 MnO2 + 2H+ + 2e– → Mn2O3 +
H2O
SECONDARY BATTERIES
A secondary or storage battery is capable of being
recharged. Its electrode reactions can proceed in
either direction.
During charging, electrical work is done on the cell to
provide the free energy needed to force the
reaction in the non-spontaneous direction.
Fuel cell
Conventional batteries supply electrical energy from the
chemical reactants stored within them. When these
reactants are consumed, the battery is "dead". An
alternative approach would be to feed the reactants
into the cell as they are required, so as to permit the
cell to operate continuously. In this case the reactants
can be thought of as "fuel" to drive the cell, hence the
term fuel cell.
MODERN HYDROGEN-OXYGEN
FUEL CELL
anode: H2(g) → 2 H+ + 2e–
cathode: ½ O2 + 2 H+ + 2e– → H2O(l)
net: H2(g) + ½ O2(g) → H2O(l)
E° = 0 v
E° = +1.23 v
E° = +1.23 v
ELECTROCHEMICAL
CORROSION
Corrosion is the deterioration
of materials by chemical
processes. Of these, the most
important
by
far
is
electrochemical corrosion of
metals,
in
which
the
oxidation
process
M → M+ + e– is facilitated by
the presence of a suitable
electron acceptor
Sacrificial
coating
Anticorrosion
Cathodic
protection
Sacrificial coating
One way of supplying this negative charge is to apply a
coating of a more active metal
a very common way of protecting steel from corrosion is
to coat it with a thin layer of zinc
Cathodic protection
A more sophisticated strategy is to
maintain a continual negative electrical
charge on a metal, so that its dissolution
as positive ions is inhibited. The entire
surface is forced into the cathodic
condition.
ELECTROLYSIS
Electrolysis refers to the decomposition of a
substance by an electric current
ELECTROLYSIS OF WATER
• Water is capable of undergoing both oxidation and
reduction
• Pure water is an insulator and cannot undergo signifigant
electrolysis without adding an electrolyte.
• Electrolysis of a solution of sulfuric acid or of a salt such as
NaNO3 results in the decomposition of water at both
electrodes:
• cathode:
([OH–] =
M)
• anode:
• net:
H2O + 2 e– → H2(g) + 2 OH–
E =+0.41 v
10-7
2 H2O → O2(g) + 4 H+ + 2 e–
E° = -0.82 v
2 H2O(l) → 2 H2(g) + O2(g) E = -1.23 v
ELECTROLYSIS
Chloralkali
industry
Refining of
Aluminum
THE CHLORALKALI
INDUSTRY
• The electrolysis of brine is carried out on a huge scale for
the industrial production of chlorine and caustic soda
(sodium hydroxide). Because the reduction potential of
Na+ is much higher than that of water, the latter substance
undergoes decomposition at the cathode, yielding
hydrogen gas and OH–.
• 2 NaCl + 2 H2O → 2 NaOH + Cl2(g) + H2(g)
A modern industrial chloralkali plant
Schematic diagram of a cell for
the production of chlorine
ELECTROLYTIC REFINING OF
ALUMINUM
• The Hall-Hérault process
takes advantage of the
principle that the melting
point of a substance is
reduced by admixture with
another substance with which
it forms a homogeneous
phase.
• The net reaction is
2 Al2O3 + 3 C → 4 Al + 3 CO2
NERNST
EQUATION
Nernst equation allows one unknown concentration to be
determined from a measurement of the cell voltage.
aOx + ne- ↔ bRed
E = E0 – (2.3026RT/nF)log ([Red]b/[Ox]a)
E: the reduction potential at the specific concentration
n: the number of electrons
R: the gas constant (8.3143 V coul deg-1 mol -1)
T: the absolute temperature
F: the Faraday constant (96487 colul eq-1)
At 25oC, the value of 2.3026RT/F is 0.05916
ELECTROCHEMICAL ANALYSIS
a. Potentiometry
b. Coulometry
c. Voltammetry
a. POTENTIOMETRY
• Potentiometry passively measures the potential of a solution between
two electrodes, affecting the solution very little in the process. The
potential is then related to the concentration of one or more analytes.
• In potentiometry, there are no current, or only negligible current flows,
so the compound in the solution remain unchanged. It is used for
measure the cell potential and for determine the analytical quantity of
interest. Potentiometry is a useful quantitative method.
Potentiometric titration
Mechanism
Ecell= Eind - Eref
Difference in potential
Reference electrode
Indicator electrode
(Constant potential)
(Change in potential)
Mobilities of ions
Solution
Electrode
Metal
Indicator
Electrode
Membrane
Potentiometry
Electrode
Reference
Electrode
Reference
electrodes
Calomel Reference
Electrodes
Silver/ Silver
Chloride Reference
Electrodes
Indicator electrodes
Metallic
Membrane
Application
Used in pH meter, by using
glass electrode
In environment, used to analyse
ion -CN-, F-, NH3, and NO3in water and in wastewater.
b. COULOMETRY
• Coulometry: electrochemical method based on
the quantitative oxidation or reduction of analyte
- Measure amount of analyte by measuring amount
of current and time required to complete
reaction
• charge = current (i) x time in coulombs
Q = ite
APPLICATION OF
COULOMETRY
COULOMETER:
• A coulometer is a
device used for
measuring the
quantity of electricity
required to bring
about a chemical
change of the
analyte.
• It is usual practice in
coulometry to
substitute the
ammeter
c. Voltammetry
• Measures current as a
function of applied
potential under
conditions that keep a
working electrode
polarized
c. Voltammetry
• Include 3 electrodes
1. Working electrode: which
the analyte is oxidizes or
reduce
2. Counter electrode: which
is often a coil of platinum
wire or a pool of
mercury.
3. Reference electrode:
potential remains
constant (Ag/AgCl
electrode or calomel)
Application of voltammetry in
Diabetes diagnostic
• Diabetes is a serious disease and is the
fourth leading cause of death by disease in
US. Its causes are unknown, and there is
no cure.
Testing blood: A Crucial Tool