Transcript Bendroji chemija

BASICS OF
ELECTROCHEMISTRY
J.Daniell and M.Faraday – English
scientists who gave basics for
electrochemistry (18-19 century)
Electrochemistry is the study of interchange of chemical
and electrical energy
Electrochemical processes – redox processes
1. GALVANIC processes
2. ELECTROLYTIC processes
COROSSIVE processes
1. GALVANIC processes
Electrochemical cell – derives electrical energy from
spontaneous redox reactions taking place within the cell
In an electrochemical cell – an electric potential is created
between two dissimilar metals
It is also called a Galvanic cell or a Voltaic cell, named after Luigi Galvani,
or Alessandro Volta respectively
in Galvanic cell:
Chemical energy is converted to Electrical energy
Chemical reaction is spontaneous and produces electricity
METAL ELECTRODE
Rod of metal immersed in a salt solution is called an electrode
Metal electrode potential will be either positive or negative
HALF-CELL
Metal electrode potential is
measured in comparison with
Hydrogen electrode potential
HYDROGEN ELECTRODE POTENTIAL
Standard hydrogen electrode
allows H2 gas molecules to interact
directly with H+ dissolved in water and
with the electrons from the external
circuit simultaneously
Pt, H2 /H+ - such special electrode
is assigned a potential of ZERO volts
METAL ELECTRODE POTENTIAL
Standard metal electrode potential is determined at
solute concentrations of 1 Molar, gas pressure of 1 atm, and a
standard temperature which is usually 25°C
Standard half-cell potential is denoted by a degree sign as
a superscript

o
Me
1.Measured against – standard hydrogen electrode
2.Concentration – 1 Molar
3.Pressure – 1 atmosphere
4.Temperature – 25°C
METAL ACTIVITY ROW
Li-K-Ca-Mg-Al-Zn-Cr-Fe-Cd-Pb-H-Cu-Hg-Ag-Pt-Au
-3.0
-2.4 -0.76
-0.4
0.0V +0.8
+1.7
GALVANIC CELLS
Galvanic cell – two half-cells of
different metals connected with a
conductor
Such galvanic cell can be
formulated:
(-)Zn│ZnSO4 ││ CuSO4│Cu(+)
Anode (negative electrode, active metal) is written on the left side, and
Cathode (positive electrode, less active metal) – on the right side
GALVANIC PROCESSES
It is customary to visualize the cell reaction in terms of two halfreactions:
an oxidation half-reaction and
a reduction half-reaction
A (-)
C (+)
– ne → Men+ (oxidation)
Less active Men+ + ne → Me0 (reduction)
More active Me0
By definition:
• Anode – the electrode where oxidation (loss of electrons) takes place;
the anode attracts anions (Zinc electrode is the anode)
• Cathode – the electrode where reduction (gain of electrons) takes place;
the cathode attracts cations (Copper electrode is the cathode)
GALVANIC PROCESSES
Anode (-) Zn0 – 2e → Zn2+ (oxidation)
Cathode (+) Cu2+ + 2e → Cu0 (reduction)
Total reaction: Zn0 + CuSO4 → Cu0 + ZnSO4 (redox)
ELECTROMOTIVE FORCE
Cell potential – Ecell
(also called electromotive force or EMF) has a contribution:
• from the anode which is a measure of its ability to lose
electrons (called as oxidation potential), and
• from the cathode which has a contribution based on its
ability to gain electrons (called as reduction potential)
Cell potential can then expressed:
Ecell = reduction potential – oxidation potential
ELECTROMOTIVE FORCE
Ecell = oreduction – ooxidation
0reduction – cathode standard potential, V
0oxidation – anode standard potential, V
Standard potential for the cell is equal:
to more positive
In our case:
ovalue
minus more negative
ovalue
Ecell = o (Cu2+/Cu) – o (Zn2+/Zn)
Ecell = 0.34 – (- 0.76) ═ 1.1 V
TYPES of GALVANIC CELLS
There are different number of
battery cell types which include:
•
•
•
•
•
•
wet cell battery
dry cell battery
molten-salt cell battery
reserve cell battery
concentration cell battery
fuel cell battery
Batteries can also be made up
for different types of materials :
•
•
•
•
•
zinc-carbon battery
alkaline battery
nickel-zinc battery
lithium-ion baterry
lead-acid battery
Wet cell battery has a liquid electrolyte, while dry cell has an electrolyte that is
immobilized as a paste, which is only liquid enough to allow the flow of electrons
Molten salt batteries use molten salt as an electrolyte – used in electric vehicles
Reserve battery – a battery that is often stored in an unassembled form and only
activates when the inner parts are assembled. Can be stored for long periods of
time and are often a part of emergency kits. Provide power for a few minutes
CONCENTRATION CELL
Is a kind of a wet galvanic cell that has two equivalent half-cells of the
same material differing only in concentrations
Rods of Cu immersed in salt solution (CuSO4) of different concentration
(A) Cu | CuSO4 (0.05 M) || CuSO4 (2.0 M) | Cu (C)
One can calculate the potential developed
by such a cell using the Nernst Equation
Potential of half-cell:
(Me) = o(Me) – (0.0592/n)·log10[Men+]
(A) Cu0 – 2e → Cu 2+ (0.05 M)
(C) Cu 2+ (2.0 M) + 2e → Cu 0
DRY HERMETIC CELL
Zinc-carbon cell battery
Electrolyte – immobilized as a paste (mixture of NH4Cl and starch or C powder)
(–) Anode – zinc box (cell is encased in Zinc casing)
(+) Cathode – carbon (graphite) rod, surrounded by a layer of MnO2
(-A) Zn  NH4Cl  MnO2  C (+C)
Cells are often encased in a plastic or
metal casing, with two points of a
negative (-) and a positive (+) sides
ALKALINE CELL
REDOX REACTIONS:
(-A) Zn + 2OH− - 2e → ZnO + H2O (oxidation)
(+C) 2MnO2 + H2O + 2e− →Mn2O3 + 2OH− (reduction)
Anode – zinc powder
Cathode – MnO2
Electrolyte – KOH
The usual voltage for an
alkaline battery is 1.5 V
FUEL CELL
Fuel cell is an electrochemical energy conversion device – most fuel cells
in use today use hydrogen and oxygen as the chemicals
Hydrogen fuel cell converts the chemicals hydrogen and oxygen into
water, and in the process it produces electricity
Battery has all of its chemicals stored inside, and converts those chemicals
into electricity. This means that a battery eventually “goes dead”
With a fuel cell, chemicals constantly flow into the cell so it never “goes dead”
As long as there is a flow of chemicals into the cell, the electricity flows out…
FUEL CELL
Hydrogen fuel cell operates similar to a battery – it has two electrodes, an
anode and a cathode, separated by a membrane. Oxygen passes over one
electrode and hydrogen over the other
Hydrogen reacts to a catalyst on the anode that converts the hydrogen gas
into negatively charged electrons (e-) and positively charged ions (H+).
Electrons flow out of the cell to be used as electrical energy
Hydrogen ions move through the electrolyte membrane to the cathode
where they combine with oxygen and the electrons to produce water
Reactions involved in a fuel cell are as follows:
2H2 - 4e- => 4H+
(+)Cathode side: O2 + 4H+ + 4e- => 2H2O
Net "redox" reaction: 2H2 + O2 => 2H2O
(-) Anode side:
Unlike batteries, fuel cells never run out
ALKALINE FUEL CELL
Alkaline fuel cells (AFCs) were one of the first
fuel cell technologies developed, and they
were the first type widely used in the U.S.
space program to produce electrical energy
and water on-board spacecrafts
These fuel cells use a solution of potassium
hydroxide in water as the electrolyte and can
use a variety of metals as a catalyst at the
anode and cathode
The reactions involved in a fuel cell are as follows:
Anode Reaction (Oxidation):
2H2 + 2O2− → 2H2O + 4e−
Cathode Reaction (Reduction): O2 + 4e– → 2O2−
Overall Cell Reaction (Redox):
2H2 + O2 → 2H2O
FUEL CELLS – PROBLEMS TO BE SOLVED
PRODUCTION: Hydrogen can be produced using diverse, domestic resources
including fossil fuels, such as natural gas and coal (with carbon
sequestration); nuclear; biomass; and other renewable energy technologies,
such as wind, solar, geothermal, and hydro-electric power.
The overall challenge to hydrogen production is cost reduction
STORAGE: On-board hydrogen storage for transportation applications
continues to be one of the most technically challenging barriers to the
widespread commercialization of hydrogen-fueled vehicles. Specific
hydrogen storage material classes: on-board reversible metal hydrides,
hydrogen adsorbents, and chemical hydrogen storage materials
Regenerative Fuel Cells: produce electricity from H2 and O2 and generate
heat and water as byproducts, just like other fuel cells. However,
regenerative fuel cell systems can also use electricity from solar power or
some other source to divide the excess water into oxygen and hydrogen
fuel by "electrolysis"
Difference between Cell and Battery
A battery is basically nothing but a
stack or pile of galvanic cells – battery
consists of multiple cells
Volta was the inventor of the voltaic
pile, the first electrical battery
Usual voltage for a single alkaline
cell is 1.5 V and the voltage can be
increased by adding on more cells
Depending on application of the
battery, the cells are combined to
provide a higher voltage, for
example a 9-Volt battery would have
6 alkaline cells with a 1.5 V charge
2. ELECTROLYTIC processes
Electrolytic cell is an electrochemical cell that undergoes a redox
reaction when electrical energy is applied
It is most often used to decompose chemical compounds, in a process
called electrolysis (the Greek word lysis means to break up)
Important examples of electrolysis are the decomposition of water into
hydrogen and oxygen, and bauxite into aluminium and other chemicals
in Electrolytic cell:
Electrical energy is converted to Chemical energy
ELECTROLYSIS
Electrolytic cell has three component parts: an electrolyte and two
electrodes (a cathode and an anode)
The electrolyte is usually a solution of water in which ions are dissolved
Molten salts such as sodium chloride are also used
as electrolytes
Anode is positive electrode +(A)
Cathode – negative electrode - (C)
Molten salt:
NaCl → Na+ + ClElectrolyte provides ions
that flow to the electrodes,
where charge-transferring
reactions take place
(-C) Na+ + e → Nao
(reduction)
(+A) Cl- - e → Clo
(oxidation)
(Cl + Cl → Cl2)
ELECTROLYSIS – application
1. Rechargeable battery – cell battery, which acts as a galvanic cell when
discharging (converting chemical energy to electrical energy), and an
electrolytic cell when being charged (converting electrical energy to
chemical energy
2. Electrorefining – process of electrolytic refining of metals is used to
extract impurities from crude metals. Here in this process a block of crude
metal is used as anode, a diluted salt of that metal is used as electrolyte
and plates of that pure metal is used as cathode
3. Electrolysis – used to produce acids, alkalis, non-metals (H2, Cl2, O2, As)
ELECTROLYSIS – application
4. Electroforming – reproduction (making copies) of objects by electrodeposition. The surface of the wax mould which bears exact impression of
the object, is coated with graphite powder in order to make it conducting.
Then the mould is dipped into the electrolyte solution as cathode and
metal will be deposited on the graphite coated surface of the mould
5. Electroplating – cathode is an object on which the electroplating to be
done (e.g. iron being copper-platted by using copper electroplating)
RECHARGEABLE BATTERY
A rechargeable battery (e.g. lead-acid battery) acts as:
• Galvanic cell – when discharging (converting chemical
energy to electrical energy), and
• Electrolytic cell – when being charged (converting
electrical energy to chemical energy)
Discharging
In the discharged state both the positive and negative plates become lead(II)
sulfate (PbSO4) and the electrolyte loses much of its dissolved sulfuric acid and
becomes primarily water. The discharge process is driven by the conduction of
electrons from the negative plate back into the cell at the positive plate in the
external circuit
Negative plate reaction (Anode Reaction):
Pb(s) + HSO4− (aq) → PbSO4(s) + H+(aq) + 2ePositive plate reaction (Cathode Reaction):
PbO2(s) + HSO4− (aq) + 3H+(aq) + 2e- → PbSO4(s) + 2H2O(l)
Total reaction can be written:
Pb(s) + PbO2(s) + 2H2SO4(aq) → 2PbSO4(s) + 2H2O(l)
Charging
In the charged state, each cell contains negative plates of elemental lead (Pb)
and positive plates of lead(IV) oxide (PbO2) in an electrolyte of approximately
33.5% v/v sulfuric acid (H2SO4). The charging process is driven by the forcible
removal of electrons from the positive plate and the forcible introduction of them
to the negative plate by the charging source
Negative plate reaction:
PbSO4(s) + H+ (aq) + 2e- → Pb(s) + HSO4− (aq)
Positive plate reaction:
PbSO4(s) + 2H2O(l) → PbO2(s) + HSO4− (aq) + 3H+(aq) + 2eOvercharging with high charging voltages generates oxygen and hydrogen gas
by electrolysis of water, which is lost to the cell. Periodic maintenance of lead
acid batteries requires inspection of the electrolyte level and replacement of any
water that has been lost
Ion motion
During discharge, H+ produced at the negative plates and from the electrolyte
solution moves to the positive plates where it is consumed, while HSO4− is
consumed at both plates. The reverse occurs during charge
Due to the freezing-point depression of the electrolyte, as the battery
discharges and the concentration of sulfuric acid decreases, the electrolyte is
more likely to freeze during winter weather when discharged