basic of cells and batteries

1. Module-6 Basic concepts of cells and batteries-nominal
voltage, operating voltage, capacity, self
discharge, depth of discharge, energy
density, service life, shelf life.
Electrochemistry of Primary cells –
Comparative features and applications of
Lechlanche, alkaline and Li-primary cells.

2. Electrochemical energy systems
Unit -V
Batteries

4. Electrochemical cell
• An electrochemical cell is a device in which a
redox reaction (chemical reaction) is utilized to
get electrical energy
• Commonly referred to as voltaic or galvanic cell
•
• The electrode where oxidation occurs is called
anode while the electrode where reduction occurs
is called cathode

5. An example for Electrochemical cell or
galvanic cell - Daniel Cell
Zinc Electrode dipped in ZnSO4 solution
Oxidation
Copper Electrode dipped in CuSO4 solution
Reduction
Each electrode is referred to as half cell
which are connected through a salt bridge
Electrode reactions:
At anode
At cathode
Zn Zn2+ + 2e-
Cu2+ + 2e- Cu

6. The Electrochemical Cell
salt bridge
4CuSO4ZnSO
reduction
at copper
cathode
oxidation
at zinc
anode
consumer

e
Half Cell I Half Cell II

8.  A battery is a package that consist of one or more
galvanic cells used for the production and storage of
electric energy.
 The power is generated because of the reaction that
happens between the anode and the cathode in the
galvanic cells.
 Each half cell contains an electrode and an electrolyte
solution. The electrolyte solution usually has ions of the
electrode in them.
 A Galvanic Cell is also known as voltaic cell which
produces electrical energy spontaneously
Battery

9. Basic Concepts of Cells and Batteries
 When two dissimilar metals or metallic compounds are immersed in an
electrolyte, there will be a potential difference produced between these
metals or metallic compounds.
 If two different kinds of metals or metallic compounds are immersed in the
same electrolyte solution, one of them will gain electrons and the other will
release electrons.
 These electrons come out into the electrolyte solution and are added to the
positive ions of the solution.
 In this way, one of these metals or compounds gains electrons and another
one loses electrons.
 As a result, there will be a difference in electron concentration between these
two metals.
 This difference of electron concentration causes an electrical potential
difference to develop between the metals or compounds.
 This electrical potential difference or emf can be utilized as a source of
voltage in any electronics or electrical circuit.
 All battery cells are based only on this basic principle.

10. Representation of galvanic cell
1. Anode is written on the left-hand side: Cathode is written on the
right hand side
2. Electrode on the left
Metal (or solid phase) Electrolyte (whole formula or ion)
ZnZnSO4 (1M)
3. Electrode on the right
Electrolyte Metal Example - CuSO4 (1M)Metal
4. A salt bridge is indicated by two vertical lines, separating the two
half cells
ZnZn2+ (1M) ║ Cu2+ (1M)Cu

11. Nernst equation
EMF of an electrochemical cell
ZnZnSO4 (aq.) ║CuSO4 (aq.)Cu
e. m. f. of an electrochemical cell
Algebraic sum of single electrode potentials (including
sign)
 Ecell = Eright-Eleft

12. Electrolytic Cell
• A device in which the electrical energy is
converted to chemical energy and resulting in a
chemical reaction
An Electrolytic cell is one kind
of battery that requires an
outside electrical source to
drive the non-spontaneous
redox reaction. Rechargeable
batteries act as Electrolytic
cells when they are being
recharged

16. Differences between a Galvanic cell and an Electrolytic cell
Electrochemical cell (Galvanic Cell) Electrolytic cell
A Galvanic cell converts chemical
energy into electrical energy.
An electrolytic cell converts electrical
energy into chemical energy.
Here, the redox reaction is spontaneous
and is responsible for the production of
electrical energy.
The redox reaction is not spontaneous
and electrical energy has to be
supplied to initiate the reaction.
The two half-cells are set up in different
containers, being connected through
the salt bridge or porous partition.
Both the electrodes are placed in a
same container in the solution of
molten electrolyte.
Here the anode is negative and cathode
is the positive electrode. The reaction
at the anode is oxidation and that at
the cathode is reduction.
Here, the anode is positive and
cathode is the negative electrode. The
reaction at the anode is oxidation and
that at the cathode is reduction.
The electrons are supplied by the
species getting oxidized. They move
from anode to the cathode in the
external circuit.
The external battery supplies the
electrons.

17. Battery Basics
 The anode or negative electrode
The reducing electrode— which gives up electrons to the
external circuit and is oxidized during the
electrochemical reaction.
 The cathode or positive electrode
— which accepts electrons from the external circuit and is
reduced during the electrochemical reaction.
 The electrolyte
The ionic conductor—which provides the medium for transfer of
charge, as ions, inside the cell between the anode and cathode.
 The electrolyte is typically a liquid, such as water or other solvents,
with dissolved salts, acids, or alkalis to impart ionic conductivity. Some
batteries use solid electrolytes, which are ionic conductors at the
operating temperature of the cell.

18.  Cell vs. Battery:
 A cell is the basic electrochemical unit providing a source of
electrical energy by direct conversion of chemical energy.
 The cell consists of an assembly of electrodes, separators,
electrolyte, container and terminals.
 A battery consists of one or more electrochemical cells,
electrically connected in an appropriate series/parallel
arrangement to provide the required operating voltage and
current levels
Battery Basics

19. Nominal voltage :
 The nominal voltage of a cell is the potential difference between the positive
and negative electrodes in fully charged condition (secondary) or freshly
prepared condition (primary).
Eg. Nominal voltage of dry cell : 1.5 V
Nominal voltage of lead‐acid cell : 2.0 V
Operating voltage :
 The sustained voltage that the cell can maintain during discharge at a
particular current is the operating voltage before reaching the cut‐off voltage
.
Eg. While discharging, the operating voltage of a dry cell for a typical discharge
duration lasting 20 hours will be of the order of 1.2 V
o When the voltage falls below this, the battery is assumed to be dead (it can no
longer be capable of taking any load)
The end voltage, or cutoff voltage (COV), is defined as a point along the
discharge curve below which no usable energy can be drawn for the specified
application.
Typically 0.9 V has been found to be the COV for a 1.5‐V cell when used in a
flashlight.
Battery Basics

20. Capacity :
 A battery's capacity is the amount of electric charge it can deliver at the rated
voltage.
 The more electrode material contained in the cell, the greater its capacity.
 The capacity of the cell = discharge current x discharge duration in hours
during which the cell will maintain voltage above a specified terminal voltage
(above the specified cut off voltage)
 Capacity is measured in units such as ampere hours (Ah).
 Eg. If a lead‐acid cell is discharged at a current of 5 amperes and it lasts for 10
hours before reaching the cut‐off voltage, then the capacity of the cell is
5 amp. x 10 hours = 50 Ah
 A small cell has less capacity than a larger cell with the same chemistry,
although they develop the same open‐circuit voltage.
Battery Basics

21. Self discharge:
 The loss in capacity of a cell when stored at specified
temperature conditions without actually discharging, is called
self discharge.
This is estimated by storing the cell at the specified temp. (eg. 25oC
for 1 year) and then discharged to check the residual capacity.
For an ideal cell (battery), the self discharge should below low.
capacity after storage
capacity before storage
x 100Percentage of self discharge =
Battery Basics

22. Depth of discharge (DOD):
DOD, is used to describe how deeply the battery is discharged.
If we say a battery is 100% fully charged, it means the DOD of this
battery is 0%,
If we say the battery have delivered 30% of its energy, here are 70%
energy reserved, we say the DOD of this battery is 30%.
And if a battery is 100% empty (discharge), the DOD of this battery
is 100%.
DOD always can be treated as how much energy that the battery
has delivered.
Battery Basics

23. Energy density:
The power that a cell can deliver at different discharge currents is
expressed as energy density of the cell
The discharge voltage considered for the calculation is the midpoint
voltage during discharge.
Battery Basics

24. Service life or Cycle life:
 Cycle Life is defined as the number of complete charge ‐ discharge
cycles a cell can perform before its nominal capacity falls below 80%
of its initial rated capacity.
Key factors affecting cycle life are:
o Storage temperature
o Depth of discharge
o Charge voltage and current
o Number of discharge cycles
Shelf life: (Calendar life)
 It is the elapsed time before a battery becomes unusable whether it
is in active use or inactive.
There are two key factors influencing shelf life are :
o temperature
o time
Battery basics

25.  Batteries use a chemical reaction to do work on
charge and produce a voltage between their output
terminals.
 The basic element in a battery is called an
electrochemical cell and makes use of an
oxidation/reduction reaction.
 An electrochemical cell which produces an external
current is called a voltaic cell. Voltages generated
by such cells have historically been referred to as
e.m.f (electromotive force).
Batteries

26. EMF is the maximum potential difference
between two electrodes of a galvanic or voltaic cell
Batteries are devices where several
electrochemical systems are connected together in
series.
Can store chemical energy for later release as
electricity
It is a source of direct electric current at a
constant voltage

27. Types of batteries
Primary battery (Primary cells)
The cell reaction is not reversible. When all the
reactants have been converted to product, no more
electricity is produced and the battery is dead.
Secondary battery (secondary cells)
The cell reactions can be reversed by passing
electric current in the opposite direction. Thus it
can be used for a large number of cycles.
The materials (reactants, products, electrolytes)
pass through the battery, which is simply an
electrochemical cell that converts chemical to
electrical energy.
Flow battery and fuel cell

29. Primary batteries
Dry or Leclanche cell
Alkaline battery
Lithium batteries

30. DRY(or LECLANCHE) CELL
•The venerable
carbonzinc cell or
Lechlanche' cell was
invented in 1866 by
Georges Lechlanche and
was the most common
small battery throughout
most of the 20th century

32. Dry cell contains
Zn, NH4Cl, ZnCl2 and MnO2
Anodic reaction
Zn(s) Zn2+
(aq) + 2e-
Cathodic reaction
2NH4
+
(aq) + 2MnO2(s) + 2e-
Mn2O3(s) + H2O(l) + 2NH3(aq)
Some of the complexity of this reaction comes from the fact
that the reduction of the ammonium ion produces two
gaseous products
Which must be absorbed to prevent the buildup of gas pressure
ZnCl2 (aq) + 2NH3 (g) Zn(NH3)2Cl2 (s)
2MnO2 (s) + H2(g) Mn2O3(s) + H2O(l)
2NH4
+
(aq) + 2e-
2NH3(aq) + H2 (g)
Zn (s) + 2 MnO2 (s) + 2 NH4Cl (aq) → ZnCl2 + Mn2O3 (s) + 2NH3+ H2O
xx

33. Applications
Disadvantages of dry cell
 The voltage of this cell is initially about 1.5 volts, but decreases as energy is taken
from the cell due to the accumulation of the products on electrodes.
 It also has a short shelf life (because of its acidic medium) and deteriorates
rapidly in cold weather.
 Oxidation of the zinc wall eventually causes the contents to leak out, so such
batteries should not be left in electric equipment for long periods.
 While these batteries have a long history of usefulness, they are declining in
application since some of their problems are overcome in ALKALINE
BATTERIES.
Flash lights, transistor radios, calculators etc

34. ALKALINE DRY CELLS
• Alkaline cells overcome some of the problems
with carbon-zinc batteries by using potassium
hydroxide (KOH) in place of ammonium chloride
as the electrolyte.
• Potassium hydroxide is a base or alkaline material,
hence "alkaline" batteries. The active materials
used are the same as in the Leclanché cell – zinc
and manganese dioxide.

35. Chemistry
The zinc anode is in the form of a
powder instead of metal , giving a large
surface area. The following half-cell
reactions take place inside the cell:
At the anode
Zn (s) + 2OH-
(aq) Zn(OH)2(s) + 2e-
2MnO2 (s) + H2O(l) + 2e- Mn2O3(s) + 2OH-(l)
At the cathode
Overall
Zn + 2MnO2 (s) + H2O(l) Mn2O3(s) + Zn(OH)2 (s)

36. Construction
This cell is “inside out” compared to the Leclanché cell

37. Advantages and Uses:
• Zinc does not dissolve as readily in
alkaline medium
• Long life
• Used in calculators and watches

38. Lithium batteries
Li cannot be used with the traditional aqueous electrolytes due to the
very vigorous corrosive reaction between Li and water, which will
results in flammable hydrogen as the product.
In the 1980s progress was made in the use of Li as an anode material
with MnO2, liquid SO2 or thionyl chlorides as the cathode, and
hexaflurophosphate dissolved in propylene carbonate as a typical organic
electrolyte.
Li cells are generally properly sealed against contact with air and
moisture
Main attractions of lithium as an anode material is its
position as the most electronegative metal in the
electrochemical series combined with its low density, thus
offering the largest amount of electrical energy per unit
weight among all solid elements.

39. The Electrochemical Series
Most wants to reduce
(gain electrons)
• Gold
• Mercury
• Silver
• Copper
• Lead
• Nickel
• Cadmium
• Iron
• Zinc
• Aluminum
• Magnesium
• Sodium
• Potassium
• Lithium
Most wants to oxidize
(lose electrons)

40. System/ Nominal Cell Voltage (V)/ Advantages/
Disadvantages Applications
• Li/SOCl2 3.60 V High Energy density; long shelf life. Only low to
moderate rate applications. Memory devices; standby electrical power
devices
• Li/SO2 3.00 V High energy density; best low-temperature
performance; long shelf life. High-cost pressurized system, Military and
special industrial needs
• Li/MnO2 3.00 V High energy density; good low-temperature
performance; cost effective. Small in size, only low-drain applications,
Electrical medical devices; memory circuits
• Li/I2 2.80 V Highly stable and mainly used in pace makers
2Li + 2SO2 Li2S2O4
4Li + 2SOCl2 4LiCl + S + SO2

41. Chemistry
The cell is represented as
Li/Li+(nonaqueous)/KOH(paste)/MnO2,Mn(OH)2,C.
The anode is lithium. The cathode is carbon in contact
with manganese (III), Manganese(IV) electrode. The
electrolyte is a paste of aqueous KOH
At anode
At cathode
MnO2 + 2H2O + 2e-
Mn(OH)2 + 2OH-
Li + MnO2 + 2H2O Li+
+ Mn(OH)3 + OH-
The overall reaction is
Li  Li+ + e-
3

42. Advantages and uses
High electron density
Long shelf life
Low self discharge
Need less maintenance
Can provide very high current
Used in auto focus cameras, mobiles

43. How batteries work
Conduction mechanisms
Development of voltage at plates
Charging, discharging, and state of charge
Key equations and models
The Nernst equation: voltage vs. ion concentration
Battery model
Battery capacity
Energy efficiency, battery life, and charge
profiles
Battery life vs. depth of discharge
Charging strategies and battery charge controllers