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PROJECT REPORT SHEET
PROJECT DESCRIPTION: BATTERY
NAME: HYEONGRAE KIM
STUDENT ID: C05723
1. General Features of Batteries
A battery is a combination of cells. The chemical battery has always been important as a
dc voltage source for the operation of radios and other electronic equipment.
The reason is that a transistor amplifier needs dc operating voltages in order to conduct
current.
However, now batteries are used more than ever for all types of portable electronic
equipment.
From the old days of radio, dry batteries are still called A, B, and C batteries, according
to their original functions in vacuum-tube operation. The A emission of electrons from a
heated cathode. A typical rating is 4.5V or 6V with a load current of 150mA or more.
The C battery was used for a small negative dc bias voltage at the control grid, typically
1.5V, with practically no current drain.
a. Advantages. The use of a single storage battery gives the common-battery system
several important advantages over the local-battery system.
(1) The storage battery used in the common-battery system is much more economical to
maintain than the dry cells of the local-battery system. Dry cells deteriorate and must be
replaced periodically, whereas storage batteries can be recharged when necessary.
(2) The storage battery of the common-battery system gives a voice signal more uniform
in amplitude, because it maintains a fairly constant voltage--more constant than the dry
cells of the local-battery system.
(3) In the common-battery system, signaling is performed automatically when the
receiver is lifted, eliminating the need for a hand generator and manual cranking, and
making the equipment of the telephone set much simpler.
(4) In the common-battery system, the operator is signaled automatically upon
completion of calls. This reduces the amount of supervision required, and allows a single
operator to handle many more lines than is possible in a local-battery system.
(5) Finally, because the single storage battery is located at the telephone central office,
inspectors are not required to make periodic visits to the associated telephone stations, as
they must do in a local-battery system, to test dry cells for deterioration.
b. Limitations. In spite of its many advantages, the common-battery system has
limitations and disadvantages that must be considered before giving it preference over the
local-battery systems in certain applications.
(1) The common-battery system requires line construction of much higher quality than
that required for the local-battery system, because current for the operation of the
transmitter at the telephone set and supervisory relays at the central office must be
supplied over the line.
(2) The lines of the common-battery system must be well balanced electrically, since
unbalance in the wires of the outside plant impairs the quality of transmission and the
distance over which transmission can be effected.
2. Primary Cells
This type cannot be recharged. After it has delivered its rated capacity, the primary cell
must be discarded. The reason is that the internal chemical reaction cannot be restored.
A primary cell is any kind of battery in which the electrochemical reaction is not
reversible. A common example of a primary cell is the disposable battery. Unlike a
secondary cell, the reaction cannot be reversed by running a current into the cell; the
chemical reactants cannot be restored to their initial position and capacity. Primary
batteries use up the materials in one or both of their electrodes.
Anode and cathode
The plate which carries the positive terminal (usually carbon) is termed the cathode and
the plate which carries the negative terminal (usually zinc) is termed the anode. This is
the reverse of the terminology used in an electrolytic cell. The reason is that the terms are
related to the passage of electric current through the electrolyte, not the external circuit
Inside the cell the anode is the electrode where chemical oxidation occurs, as it accepts
electrons from the electrolyte. The cathode is defined as the electrode where chemical
reduction occurs, as it donates electrons to the electrolyte.
Outside the cell, different terminology is used. Since the anode accepts electrons from the
electrolyte, it becomes negatively charged and is therefore connected to the terminal
marked "−" on the outside of the cell. The cathode, meanwhile, donates electrons to the
electrolyte, so it becomes positively charged and is therefore connected to the terminal
marked "+" on the outside of the cell.[1]
Old textbooks sometimes contain different terminology that can cause confusion to
modern readers. For example, a 1911 textbook by Ayrton and Mather[2] describes the
electrodes as the "positive plate" and "negative plate" in a way that contradicts modern
usage.
Polarization
A primary cell becomes polarized when in use. This means that hydrogen accumulates at
the cathode and reduces the effectiveness of the cell. To remove the hydrogen, a
depolarizer is used and this may be mechanical, chemical or electrochemical.
Attempts have been made to make simple cells self-depolarizing by roughening the
surface of the copper plate to facilitate the detachment of hydrogen bubbles. These
attempts have had little success. Chemical depolarization utilizes an oxidizing agent, such
as manganese dioxide (e.g. Leclanché cell and Zinc-carbon cell) or nitric acid (e.g.
Bunsen cell and Grove cell), to oxidize the hydrogen to water. Electrochemical
depolarization exchanges the hydrogen for a metal, such as copper (e.g. Daniell cell), or
silver (e.g. Silver-oxide cell).
3. The Voltaic cell
A voltaic cell is created whenever dissimilar metals, connected in some way, are
immersed in a conductive fluid.The tendency for the electrons to flow from one chemical
to another, such as from the zinc metal to the copper ion as shown here, is something that
can be channeled and controlled. Channeling the flow of electrons is what we will take up
here. The tendency for the electrons to flow from one chemical to another, such as from
the zinc metal to the copper ion as shown here, is something that can be channeled and
controlled. Channeling the flow of electrons is what we will take up here.
Zn Zn2+ + 2e- 0.76v
Cu2+ + 2e- Cu +0.34v
____________________________
Zn + Cu2+ Zn2+ + Cu 1.10v
Copper-Zinc Voltaic Cell
The apparatus shown here is used to channel electron flow in this reaction and is called a
voltaic cell. The zinc metal is on the right side and it is providing electrons to the blue
copper ions in the beaker to the left. However, to get there the electrons have to travel
through a wire. We can see that they are doing so because the voltmeter that they must
pass through shows a reading.
This voltaic cell has copper metal and copper sulfate solution in the left hand beaker; zinc
metal and zinc sulfate in the right hand beaker; probably potassium sulfate in the salt
bridge. With the wires connected to the voltmeter, you can see that there is a voltage
reading; electricity is flowing through this cell. It's a voltaic cell because the chemical
reaction is causing the flow of electric current. Here you can see the importance of the
salt bridge. In the top picture, the salt bridge has been removed and and the voltage
reading is zero showing the there is no current flowing. In the bottom picture, the salt
bridge is replaced and the voltmeter again shows a reading. This same voltaic cell is
diagrammed for you here (and in example 20 in your workbook). In the diagram, copper
is on the right and zinc is on the left.
Zinc metal is reacting with the copper ion to form zinc ion and copper metal. The zinc
metal is on one side and the copper ion in solution is on the other side in the other beaker.
Consequently, the two chemicals are not in direct contact with one another.
Even so, the tendency for the zinc metal to lose (or transfer) electrons to the copper ion
still exists. That transfer is made possible by connecting a wire between the zinc metal
and the copper metal. The electrons go from the zinc over to the copper metal where they
can react with the copper ions in solution.
Over a period of time, the zinc electrode will dissolve and increase the concentration of
the zinc ion solution. The copper ion will plate onto the copper electrode and thus the
concentration of the copper ion in solution will decrease. As this happens, the reaction
slows down and the voltage decreases.
Zinc is the anode because that is where oxidation is occurring. The oxidation halfreaction is Zn Zn2+ + 2e-. Those electrons go over to the copper side where they react
with copper ion and change it into copper metal (Cu2+ + 2e- Cu), which is the
reduction half-reaction that takes place at the cathode.
The voltage for such a cell can be calculated
using the standard oxidation potential list.
Working that through, as we have done
before, the voltage is 1.10 volts.
Zn Zn2+ + 2e- 0.76v
Cu2+ + 2e- Cu +0.34v
____________________________
Zn + Cu2+ Zn2+ + Cu 1.10v
However, if you measure the voltage of a cell like this using a voltmeter, you will likely
not get that particular voltage. The reasons are several.
One is that quite often solutions are not kept at 25oC, particularly in this lab.
Another reason is that the concentrations of all the solutions are probably not one mole
per liter.
Also, the voltage measured by a voltmeter depends on the electrical resistance of the cell.
So the electrical resistance in the connections (especially bad ones), the wire, the
voltmeter, and the solutions will tend to cause the voltage to be less than what you
calculate it to be under standard or ideal conditions.
We have a cell like this set up in the lab. You should look at it, experiment with it,
remove and replace the salt bridge, check the voltage, and so forth, when you are in the
lab.
When two different conducting materials are immersed in an electrolyte, as illustrated
in Fig. 12-3a, the chemical action of forming a new solution results in the separation of
charges. This method for converting chemical energy into electric energy is a voltaic
cell. It is also called a galvanic cell, named after Luigi Galvani.
Referring to Fig. 12-3a, the charged conductors in the electrolyte are the electrodes or
plates of the cell. They serve as the terminals for connecting the voltage output to an
external circuit, as shown in Fig. 12-3b. Then the potential difference resulting from the
separated charges enables the cell to function as a source of applied voltage. The voltage
across the cell’s terminals forces current to flow in the circuit to light the bulb.
4. Electromotive Series
The fact that the voltage output of a cell depends on its elements can be seen from
Table 12-2. This list, called the electrochemical series or electromotive series, gives the
relative activity in forming ion charges for some of the chemical elements.
The potential for each element is the voltage with respect to hydrogen as a zero
reference. The difference between the potentials for two different elements indicates the
voltage of an ideal cell using these electrodes. It should be noted, though, that other
factors, such as the electrolyte, cost stability, and long life, are important for the
construction of commercial batteries.
5. Carbon-zinc Dry Cell
A zinc-carbon dry cell or battery is packaged in a zinc can that serves as both a container
and negative terminal. It was developed from the wet Leclanché cell (pronounced
/lɛklɑːnˈʃeɪ/). The positive terminal is usually a carbon rod or graphite rod surrounded
by a mixture of manganese dioxide and carbon powder . The electrolyte used is a paste of
zinc chloride and ammonium chloride dissolved in water. Zinc chloride cells are an
improved version from the original ammonium chloride variety. Zinc-carbon batteries are
the least expensive primary batteries and thus a popular choice by manufacturers when
devices are sold with batteries included. They are commonly labeled as "General
Purpose" batteries. They can be used in remote controls, flashlights, clocks, or transistor
radios, since the power drain is not too heavy.
History
By 1876, the wet Leclanché cell was made with a compressed block of manganese
dioxide. In 1888 Dr. Carl Gassner built the first "dry" version by using a zinc cup as the
anode and making the electrolyte with a paste of plaster of Paris (and later, wheat flour)
to gel and immobilize the electrolyte. In 1900 Gassner demostrated dry cells for portable
lighting at the World's Fair in Paris. Continual improvements were made to the stability
and capacity of zinc-carbon cells throughout the 20th Century; by the end of the century
the capacity of a zinc-carbon cell had increased four-fold over the 1910 equivalent.
A primary cell
A zinc-carbon dry cell is described as a primary cell because as the cell is discharged, it is
not intended to be recharged and must be discarded. "Battery Rejuvenators" were once
marketed to restore partially discharged zinc-carbon cells by applying a reverse current to
them. However the effects of such devices were only temporary and prone to cause the
cell to leak or burst. [1] Zinc-carbon cells are more likely to leak as the anode is the
container.
Chemical Reactions
The container of the zinc-carbon dry cell is a zinc can. This contains a layer of NH4Cl
with ZnCl2 aqueous paste separated by a paper layer from a mixture of powdered carbon
& manganese (IV) oxide (MnO2) which is packed around a carbon rod.
In a dry cell, the outer zinc container is the negative terminal. The zinc is oxidised
according to the following half-equation.
Zn(s) → Zn2+(aq) + 2 eA graphite rod surrounded by a powder containing manganese(IV) oxide is the positive
terminal. The manganese dioxide is mixed with carbon powder to increase the electrical
conductivity. The reaction is as follows:
2MnO2(s) + H2(g)→ Mn2O3(s) + H2O(l)
The H2 comes from the NH4+(aq):
2NH4+(aq) + 2 e- → H2(g) + 2NH3(aq)
and the NH3 combines with the Zn2+.
In this half-reaction, the manganese is reduced from an oxidation state of (+4) to (+3).
There are other possible side-reactions, but the overall reaction in a zinc-carbon cell can
be represented as:
Zn(s) + 2MnO2(s) + 2NH4+(aq) → Mn2O3(s) + Zn(NH3)22+(aq) + H2O(l)
The battery has an e.m.f. of about 1.5 V. The approximate nature of the e.m.f is related to
the complexity of the cathode reaction. The anode (zinc) reaction is comparatively simple
with a known potential. Side reactions and depletion of the active chemicals increases the
internal resistance of the battery, and this causes the e.m.f. to drop.
The carbon-zinc dry cell is a very common type because of its low cost. It is also called
the Leclanche cell, named after the inventor. Examples are shown in Fig. 12-1, and Fig.
12-4 illustrates internal construction for th D-size round cell.
Voltage output for the carbon-zinc cell is 1.4 to 1.6V, with a nominal value of 1.5V.
Suggested current range is up to 150mA for th D size, which has a height of 2 and 1/4
in, and volume of 3.18 in. The C, A, AA, and AAA sizes are smaller, with lower current
ratings. The larger No. 6 cell has a height of 6 in.. a diameter of 2 and 1/2 in. And a
current range of up to 1500 mA.
6. Alkaline Cell
Alkaline batteries and alkaline cells (a battery being a collection of multiple cells) are a
type of disposable battery or rechargeable battery dependent upon the reaction between
zinc and manganese dioxide (Zn/MnO2).
Compared with zinc-carbon batteries of the Leclanché or zinc chloride types, while all
produce approximately 1.5 volts per cell, alkaline batteries have a higher energy density
and longer shelf-life. Compared with silver-oxide batteries, which alkalines commonly
compete against in button cells, they have lower energy density and shorter lifetimes but
lower cost.
The alkaline battery gets its name because it has an alkaline electrolyte of potassium
hydroxide, instead of the acidic ammonium chloride or zinc chloride electrolyte of the
zinc-carbon batteries which are offered in the same nominal voltages and physical size.
Other battery systems also use alkaline electrolytes, but they use different active
materials for the electrodes.
Chemistry
In an alkaline battery, the anode (negative terminal) is made of zinc powder (which
allows more surface area for increased rate of reaction therefore increased electron flow)
and the cathode (positive terminal) is composed of manganese dioxide. Alkaline batteries
are comparable to zinc-carbon batteries, but the difference is that alkaline batteries use
potassium hydroxide (KOH) as an electrolyte rather than ammonium chloride or zinc
chloride.
The half-reactions are:
Zn (s) + 2OH− (aq) → ZnO (s) + H2O (l) + 2e−
2MnO2 (s) + H2O (l) + 2e− →Mn2O3 (s) + 2OH− (aq)
Capacity
Capacity of an alkaline battery is larger than an equal size Leclanché or zinc-chloride cell
because the manganese dioxide anode material is purer and denser, and space taken up by
internal components such as current collectors is less. An alkaline cell can provide
between three and five times as much operating time.[2]
The capacity of an alkaline battery is strongly dependent on the load. An AA-sized
alkaline battery might have an effective capacity of 3000 mAh at low power, but at a load
of 1000 mA, which is common for digital cameras, the capacity could be as little as 700
mAh.[3] The voltage of the battery declines steadily during use, so the total usable
capacity depends on the cut-off voltage of the application. Unlike Leclanche cells the
alkaline cell delivers about as much capacity on intermittent or continuous light loads. On
a heavy load, capacity is reduced on continuous discharge compared with intermittent
discharge, but the reduction is less than for Leclanche cells.
Voltage
The nominal voltage of a fresh alkaline cell is 1.5 V. Multiple voltages may be achieved
with series of cells. The effective zero-load voltage of a non discharged alkaline battery
varies from 1.50 to 1.65 V, depending on the chosen manganese dioxide and the contents
of zinc oxide in the electrolyte. The average voltage under load depends on discharge and
varies from 1.1 to 1.3 V. The fully discharged cell has a remaining voltage in the range of
0.8 to 1.0 V.
Construction
Alkaline batteries are manufactured in standardized cylindrical forms interchangeable
with zinc-carbon batteries, and in button forms. Several individual cells may be
interconnected to form a true "battery", such as those sold for use with flashlights and the
9 volt transistor-radio battery.[4]
A cylindrical cell is contained in a drawn steel can, which is the cathode current collector.
The cathode mixture is a compressed paste of manganese dioxide with carbon powder
added for increased conductivity. The paste may be pressed into the can or deposited as
pre-molded rings. The hollow center of the cathode is lined with a separator, which
prevents mixing of the anode and cathode materials and short-circuiting of the cell. The
separator is made of a non-woven layer of cellulose or a synthetic polymer. The separator
must conduct ions and remain stable in the highly alkaline electrolyte solution.
The anode is composed of a dispersion of zinc powder in a gel containing the potassium
hydroxide electrolyte. To prevent gassing of the cell at the end of its life, more
manganese dioxide is used than required to react with all the zinc.
When describing standard AAA,AA, C, sub-C and D size cells, the anode is connected to
the flat end while the cathode is connected to the end with the raised button.
Another popular type is the manganese-zinc cell shown in Fig. 12-5, which has an
alkaline electrolyte. It is available as either a primary or a secondary cell but the
primary type is more common. Output is the same 1.5V as a carbon-zinc cell but the
alkaline cell lasts much longer. The electrochemical system consists of a powdered
zinc anode and a manganese dioxide cathode in an alkaline electrolyte.
The electrolyte is potassium hydroxide, which is the main difference between the
alkaline and Leelanche cells. Hydroxide compounds are alkaline with negative
hydroxyl(OH) ions, whereas an acid electrolyte has positive hydrogen(H) ions.
Voltage output from the alkaline cell is 1.5V.
The alkaline cell has man applications because of its ability to work with high
efficiency with continuous high discharge rates. Depending on the application,
an alkaline cell can provide up to seven times the service of a Leclanche cell.
As examples, in a transistor radio an alkaline cell will normally have twice the
service life of a general purpose carbon-zinc cell; in toys the alkaline cell typically
provides seven time more service.
7. Aditional Types of Primary Cells
The miniature button construction shown in Fig. 12-6 is often used for the mercury cell
and the silver-oxide cell. Diameter of the cell is 3/8 to 1 in.
8. Mercury Cell
Zinc-mercuric oxide or .mercury. cells take advantage of the high electrode potential of
mercury to offer a very high energy density combined with a very flat discharge curve.
Mercuric oxide forms the positive electrode, sometimes mixed with manganese dioxide.
The negative electrode is metallic zinc powder, and the electrolyte is usually potassium
hydroxide, absorbed in a multi-layer separator.
The nominal terminal voltage of a mercury cell is 1.35V, and this remains almost
constant over the life of the cell. They have a low internal resistance, which is again fairly
constant.
Although made only in small .button. sizes, mercury cells are capable of reasonably high
pulsed discharge current levels and are thus suitable for applications such as quartz
analog watches and hearing aids as well as voltage references in instruments, etc. Silver
oxide cells
The electrochemical system consists of a zimc anode, a mercury compound for the
cathode, and an electrolyte of potassium or sodium hydroxide. Mercury cells are
available in flat, round cylinder, and miniature but-ton shapes. It should be noted,
though, that some round mercury cells have the top button as the negative terminal and
the bottom terminal positive. The open circuit voltage is 1.35V when the cathode is
mercuric oxide (HgO) and 1.4V or more with mercuric oxide/manganese dioxide.
9. Silver Oxide Cell
The zinc-silver oxide cell takes advantage of the high electrode potential of silver to
again provide a high energy density combined with a very flat discharge curve. The silver
oxide forms the positive electrode, again sometimes mixed with a small amount of
manganese dioxide. The negative electrode is powdered metallic zinc, mixed into a gel
with the electrolyte . which is usually either potassium hydroxide or sodium hydroxide. A
separator membrane stops the negative electrode gel from mixing with the positive
electrode. The nominal terminal voltage of a silver oxide cell is slightly over 1.5V, and
remains almost flat over the life of the cell . which is not as long as that of a mercury cell
of the same size and weight. The internal resistance is also low and relatively constant.
Low temperature performance is quite good. The silver oxide cell is again made only in
small .button. sizes of modest capacity, but has a relatively good pulsed discharge
capability. It.s again mainly used in watches, hearing aids, pagers and test instruments.
The electrochemical system consists of a zinc anode, a cathode of silver oxide
(AgO2) with small amounts of manganese dioxide, and an electrolyte of potassium or
sodium hydroxide. It is commonly available in the miniature button shape shown in
Fig. 12-6. The open circuit voltage is 1.6V, but the nominal output with a load is
considered to be 1.5V. Typical applications include hearing aids, cameras, and
electronic watches, which use very little current.
10. Dry Cells
The most common type of battery used today is the "dry cell" battery. There are many
different types of batteries ranging from the relatively large "flashlight" batteries to the
minaturized versions used for wristwatches or calculators. Although they vary widely in
composition and form, they all work on the sample principle. A "dry-cell" battery is
essentially comprised of a metal electrode or graphite rod (elemental carbon) surrounded
by a moist electrolyte paste enclosed in a metal cylinder as shown below.
In the most common type of dry cell battery, the cathode is composed of a form of
elemental carbon called graphite, which serves as a solid support for the reduction halfreaction. In an acidic dry cell, the reduction reaction occurs within the moist paste
comprised of ammonium chloride (NH4Cl) and manganese dioxide (MnO2):The types of
dry cells include carbon-zinc, zinc chloride (heavy-duty),and manganese-zinc (alkaline).
Actually, the alkaline cell is better for heavy-duty usethan the zinc-chloride type. They
are commonly used in the round, cylinder typeslisted in Table 12-3, for the D, C, AAA
sizes. The small button cells generally useeither mercury or silver oxide. All these dry
cells are the primary type that cannot berecharged. Each has output of 1.5V except for
the 1.35V mercury cell.
11. Lithium Cell
Lithium batteries are di sposable (primary) batteries that have lithium metal or lithium
compounds as an anode. Depending on the design and chemical compounds used, lithium
cells can produce voltages from 1.5 V to about 3.7 V, over twice the voltage of an
ordinary zinc-carbon battery or alkaline cell battery.[1] Lithium batteries are widely used
in products such as portable consumer electronic devices.
Lithium batteries find application in many long-life, critical devices, such as artificial
pacemakers and other implantable electronic medical devices. These devices use
specialized lithium-iodide batteries designed to last 15 or more years. But for other, less
critical, applications such as in toys, the lithium battery may actually outlast the toy. In
such cases, an expensive lithium battery is not cost-efficient.
Lithium batteries can be used in place of ordinary alkaline cells in many devices, such as
clocks and cameras. Although they are more costly, lithium cells will provide much
longer life, thereby minimizing battery replacement. However, attention must be given to
the higher voltage developed by the lithium cells before using them as a drop-in
replacement in devices that normally use ordinary cells.
Small lithium batteries are very commonly used in small, portable electronic devices,
such as PDAs, watches, thermometers, and calculators, as backup batteries in computers
and communication equipment, and in remote car locks. They are available in many
shapes and sizes, with a common variety being the 3 volt "coin" type manganese variety,
typically 20 mm in diameter and 1.6–4 mm thick. The heavy electrical demands of many
of these devices make lithium batteries a particularly attractive option. In particular,
lithium batteries can easily support the brief, heavy current demands of devices such as
digital cameras, and they maintain a higher voltage for a longer period than alkaline cells.
Some other lithium batteries use a platinum-iridium alloy instead of more usual
compounds. These batteries are generally not preferred, as their cost is high and they tend
to be fragile.
The lithium cell is a relatively new primary cell. However, its high output voltage, long
shelf life, low weight, and small volume make the lithium cell an excellent choice for
special applications. The open circuit output voltage is either 2.9V or 3.7V, depending
on the electrolyte. A lithium cell can provide at least 10 times more energy than the
equivalent carbon-zinc cell However, lithium is a very active chemical element. Many of
the problems in construction have been solved, though, especially for small cells
delivering low current. One interesting application is a lithium cell as the dc power
source for a cardiac pacemaker. The long service life is important for this use. Two
forms of lithium cells have obtained widespread use. These are the lithium-sulfur
dioxide (LiSO2) type and the lithium-thionyle chloride type. Output is approximately
3V.
12. Lead-Acid Wet Cell
Where high values of load current are necessary, the lead-acid cell is the type most
common only used. The electrolyte is a dilute solution of sulfuric acid (H2SO4).
In the application of battery power to start the engine in an auto-mobile, for example,
the load current to the starter motor is typically 200 to 400A. One cell has a nominal
output of 2.1V, but lead-acid cells are often used in a series combination of three for a 6V battery and six for a 12-V battery. The lead-acid type is a secondary cell or storage
cell, which can be recharged. The charge and discharge cycle can be repeated many
times to restore the output voltage, as long as the cell is in good physical condition.
Lead-acid batteries, invented in 1859 by French physicist Gaston Planté, are the oldest
type of rechargeable battery. Despite having a very low energy-to-weight ratio and a low
energy-to-volume ratio, their ability to supply high surge currents means that the cells
maintain a relatively large power-to-weight ratio. These features, along with their low
cost, make them attractive for use in motor vehicles to provide the high current required
by automobile starter motors.
Electrochemistry
In the charged state, each cell contains electrodes of elemental lead (Pb) and Lead(IV)
Oxide (PbO2) in an electrolyte of approximately 33.5% v/v (4.2 Molar) sulfuric acid
(H2SO4).
In the discharged state both electrodes turn into lead(II) sulfate (PbSO4) and the
electrolyte loses its dissolved sulfuric acid and becomes primarily water. Due to the
freezing-point depression of water, as the battery discharges and the concentration of
sulfuric acid decreases, the electrolyte is more likely to freeze during winter weather.
The chemical reactions are (discharged to charged):
Anode (oxidation):
Cathode (reduction):
Because of the open cells with liquid electrolyte in most lead-acid batteries, overcharging
with high charging voltages generates oxygen and hydrogen gas by electrolysis of water,
forming an explosive mix. The acid electrolyte is also corrosive.
Practical cells are usually not made with pure lead but have small amounts of antimony,
tin, calcium or selenium alloyed in the plate material to strengthen it and make simplify
manufacture.
Voltages for common usages
These are general voltage ranges for six-cell lead-acid batteries:




Open-circuit (quiescent) at full charge: 12.6 V to 12.8 V (2.10-2.13V per cell)
Open-circuit at full discharge: 11.8 V to 12.0 V
Loaded at full discharge: 10.5 V.
Continuous-preservation (float) charging: 13.4 V for gelled electrolyte; 13.5 V for
AGM (absorbed glass mat) and 13.8 V for flooded cells
1. All voltages are at 20 °C (68 °F), and must be adjusted -0.022V/°C for
temperature changes.
2. Float voltage recommendations vary, according to the manufacturer's
recommendation.
3. Precise (±0.05 V) float voltage is critical to longevity; insufficient (sulfation) is
almost as bad as excessive (corrosion and electrolyte loss)




Typical (daily) charging: 14.2 V to 14.5 V (depending on manufacturer's
recommendation)
Equalization charging (for flooded lead acids): 15 V for no more than 2 hours.
Battery temperature must be monitored.
Gassing threshold: 14.4 V
After full charge, terminal voltage drops quickly to 13.2 V and then slowly to
12.6 V.
Portable batteries, such as for miners' cap lamps (headlamps) typically have two cells,
and use one third of these voltages.
Applications
Wet cell stand-by (stationary) batteries designed for deep discharge are commonly used
in large backup power supplies for telephone and computer centers, grid energy storage,
and off-grid household electric power systems [4]. Lead-acid batteries are used in
emergency lighting in case of power failure.
Traction (propulsion) batteries are used for in golf carts and other battery electric vehicles.
Large lead-acid batteries are also used to power the electric motors in diesel-electric
(conventional) submarines and are used on nuclear submarines as well. Motor vehicle
starting, lighting and ignition (SLI) batteries (car batteries) provides current for starting
internal combustion engines.
Valve-regulated lead acid batteries cannot spill their electrolyte. They are used in back-up
power supplies for alarm and smaller computer systems (particularly in uninterruptible
power supplies) and for electric scooters, electrified bicycles, marine applications, battery
electric vehicles or micro hybrid vehicles, and motorcycles.
Lead-acid batteries were used to supply the filament (heater) voltage (usually between 2
and 12 volts with 2 V being most common) in early vacuum tube (valve) radio receivers.
13. Charging the Lead-Acid Battery
The requirements are illustrated in Fig. 12-10. An external dc voltage source is
necessary to produce current in one direction. Also, the charging voltage must be
more than the battery emf. Approximately 2.5V per cell is enough to overcome the
cell emf so that the charging voltage can produce current opposite to the direction
of discharge current.
14. Additional Types of Secondary Cells
A secondary cell is a storage cell that can be recharged by reversing the internal
chemical reaction. A primary cell must be discarded after it has been completely
discharged. The lead-acid cell is the most common type of storage cell.
However, other types of secondary cells are available.
The secondary cell wall is a structure found in many plant cells, located between the
primary cell wall and the plasma membrane. The cell starts producing the secondary cell
wall after the primary cell wall is complete and the cell has stopped expanding.
The secondary cell wall consists mainly of cellulose, but also other polysaccharides,
lignin, and glycoproteins. It sometimes consists of three distinct layers - S1, S2 and S3 where the direction of the Cellulose microfibrils differs between the layers.] Apparently
there are no Structural proteins or enzymes in the secondary wall.
The secondary cell wall have different ratios of wall constituents compared to the
primary wall. An example of this is that wood secondary walls contain xylans, whereas
the primary wall contain xyloglucans and the cellulose fraction is higher in the secondary
wall. Pectins may also be absent from the secondary wall and apparently it contain no
Structural proteins or enzymes.
The Cellulose microfibrils give tensile strength, whereas lignification in addition to
making the secondary wall impermeable to water also give a "brittle" texture.
Conceptually this give lignified secondary wall properties resembling armored concrete,
where the cellulose microfibrils act as the armoring and the lignin as concrete.
Lignification of the secondary wall confer resistance to pathogens by two mechanisms.
As ligning repel water, hydrolytic enzymes are less likely to attack and successfully
penetrate the wall and it lowers the nutritional value of the wall, providing less energy to
pathogens.
The secondary wall usually is absent under the regions of the primary wall, which contain
pit pairs, giving rise to a pit cavity. (this is somewhat simplified and someone who know
more about it than the author of this sentence should expand the section, e.g. describing
simple and bordered pit cavities).
Wood consists mostly of secondary cell wall, and holds the plant up against gravity.
Some secondary cell walls store nutrients, such as those in the cotyledons and the
endosperm. These contain little cellulose, and mostly other polysaccharides.
15. Series and Parallel Cells
An applied voltage higher than the emf of one cell can be obtained by connecting cells in
series. The botal voltage available across the battery of cells is equal to theum of the
individual values for each cell. Parallel cells have the same voltage as one cell but have
more current capacity. The combination of cells is a battery.
Single cell applications
Single cell batteries are used in watches, memory back up and cell phones. The nickelbased cell provides a nominal cell voltage of 1.2V; alkaline is 1.5V; silver-oxide 1.6V,
lead-acid 2V; primary lithium 3V and lithium-ion 3.6V. Spinel, lithium-ion polymer and
other lithium-based systems sometimes use 3.7V as the designated cell voltage. This
explains the unfamiliar voltages such as 11.1V if three cells are connected in series.
Modern microelectronics makes it possible to operate cell phones and other low power
portable communications devices from a single 3.6V lithium-ion cell. Mercury, a popular
cell for light meters in the 1960s has been discontinued because of environmental
concerns.
Nickel-based cells are either marked 1.2V or 1.25V. There is no difference in the cells
but only preference in marking. Most commercial batteries are identified with 1.2V/cell;
industrial, aviation and military batteries are still marked with 1.25V/cell.
Serial connection
Portable equipment with high-energy needs is powered with battery packs in which two
or more cell are connected in series. Figure 1 shows a battery pack with four 1.2-volt
cells in series. The nominal voltage of the battery string is 4.8V.
Serial connection
Portable equipment with high-energy needs is powered with battery packs in which two
or more cell are connected in series. Figure 1 shows a battery pack with four 1.2-volt
cells in series. The nominal voltage of the battery string is 4.8V.
Figure 1: Serial connection
of four cells.
Adding cells in a string increases the
voltage but the current remains the
same.
High voltage batteries have the advantage of keeping the conductor and switch sizes
small. Medium-priced industrial power tools run on 12V to 19.2V batteries; high-end
power tools go to 24V and 36V to get more power. The car industry will eventually
increase the starter-light-ignition (SLI) battery from 12V (14V) to 36V, better known as
42V. These batteries have 18 lead-acid cells in series. The early hybrid cars are running
on 148V batteries. Newer models feature batteries with 450-500V; mostly on nickelbased chemistry. A 480-volt nickel-metal-hydride battery has 400 cells in series. Some
hybrid cars are also experimenting with lead acid.
42V car batteries are expensive and produce more arcing on the switches than the 12V.
Another problem with higher voltage batteries is the possibility of one cell failing.
Similar to a chain, the more links that are connected in series, the greater the odds of one
failing. A faulty cell would produce a low voltage. In an extreme case, an open cell could
break the current flow. Replacement of a faulty cell is difficult because of matching. The
new cell will typically have a higher capacity than the aged cells.
Figure 2 illustrates a battery pack in which cell 3 produces only 0.6V instead of the full
1.2V. With the depressed operating voltage, the end-of-discharge point will be reached
sooner than with a normal pack and the runtime is severely shortened. Once the
equipment cuts off due to low voltage, the remaining three cells are unable to deliver the
stored energy. Cell 3 could also exhibit a high internal resistance, causing the string to
collapse under load. A weak cell in a battery string is like a blockage in a garden hose
that restricts water flow. Cell 3 could also be shorted, which would lower the terminal
voltage to 3.6V, or be open and cut off the current. A battery is only as good as the
weakest cell in the pack.
Figure 2: Serial connection
with one faulty cell.
Faulty cell 3 lowers the overall
voltage to 4.2V, causing the
equipment to cut off prematurely.
Parallel connection
To obtain higher ampere-hour (Ah) ratings, two or more cells are connected in parallel.
The alternative to parallel connection is using a larger cell. This option is not always
available because of limited cell selection. In addition, bulky cell sizes do not lend
themselves to build specialty battery shapes. Most chemistries allows parallel connection
and lithium-ion is one of the best suited. Figure 3 illustrates four cells connected in
parallel. The voltage of the pack remains at 1.2V but the current handling and runtime are
increased four fold
Figure 3: Parallel connection
of four cells.
With parallel cells, the voltage
stays the same but the current
handling and runtime increases.
A high resistance or open cell is less critical in a parallel circuit than the serial
configuration but the parallel pack will have reduced load capability and a shorter
runtime. It's like an engine running only on three cylinders. An electrical short would be
more devastating because the faulty cell would drain the energy from the other cells,
causing a fire hazard. Figure 4 illustrates a parallel configuration with one faulty cell.
Figure 4: Parallel connection
with one faulty cell.
A weak cell will not affect the
voltage but provide a low runtime.
A shorted cell could cause
excessive heat and create a fire
hazard.
Serial/parallel connection
Figure 5 illustrates a parallel/serial connection. This allows good design flexibility and
attains the wanted voltage and current ratings by using a standard cell size. It should be
noted that the total power does not change with different configurations. The power is the
product of voltage times current.
Figure 5: Serial/ parallel
connection of four cells.
The configurations will not affect
the overall power but provide the
most suitable voltage and current
source for the application.
Serial/parallel connections are common with lithium-ion. One of the most popular cells is
the 18650 (18mm diameter; 650mm long). Because of the protection circuit, which must
monitor each cell connected in series, the maximum practical voltage is 14.4V. The
protection must also monitor strings placed in parallel.
16. Current Drain Depends on Load Resistance
It is important to note that the current rating of batteries, or any voltage source, is only
a guide to typical values permissible for normal service life.
The actual amount of current produced when the battery is connected to a load
resistance is equal to I = V / R, by Ohm’s law.
17. Internal Resistance of a Generator
Internal resistance is a concept that helps model the electrical consequences of the
complex chemical reactions inside a battery. When a current is flowing through a cell, the
measured e.m.f. (voltage output) is lower than when there is no current delivered by the
cell.
Any source that produces voltage output continuously is a generator. It may be a cell
separating charges by chemical action or a rotary generator converting motion and
magnetism into voltage output.
The internal resistance of a battery can not be measured using the "resistance" or "ohms"
setting on a conventional multimeter, since it requires a current to be observed. However,
it can be calculated from current and voltage data measured from a test circuit containing
the battery and a load resistor RL. Since both the internal resistance and load resistor are
in series with the ideal voltage source, Kirchhoff's Laws and Ohm's law give
. This equation can be solved for internal resistance:
where




RB is the internal resistance of the battery
VB is the battery voltage without a load L
IL is the current supplied by the battery with this load L
RL is the resistance of this load L.
Internal resistance increases with the age of a battery, but for most commercial batteries
the internal resistance is on the order of 1 ohm.
It should be noted that the above only applies to ideal batteries under ideal load
conditions and does not directly relate to real world internal resistance of batteries due to
the chemical nature of the cells.
18. Why the Terminal Voltage Drops with More Load Current
Figure 12-19 illustrates how the output of a 100-V source can drop to 90V because of
the internal 10-V drop across ri. In Fig. 12-19a, the voltage across the output terminals
is equal to the 100V of VG because there is no load current on an open circuit. With no
current, the voltage drop across ri is zero. Then the full generated emf, open-circuit
voltage, or no-load voltage. We cannot connect the test leads inside the source to
measure VG. However, measuring this no-load voltage without any load current
provides a method of determining the internally generated emf. We can assume the
voltmeter draws practically no current because of its very high resistance.
19. Constant- Voltage and Constant- Current Sources
A generator with very low internal resistance is considered to be a constant-voltage
source.
Then the output voltage remains essentially the same when the load current changes.
This idea is illustrated in Fig. 12-21a, for a 6-V lead-acid battery with an ri of
0.005Ω is less than 0.5V. A current source is an electrical or electronic device that
delivers or absorbs electric current. A current source is the dual of a voltage source. The
term constant-current sink is sometimes used for sources fed from a negative voltage
supply. Figure 1 shows a schematic for an ideal current source driving a resistor load.
Most sources of electrical energy (mains electricity, a battery, ...) are best modeled as
voltage sources. Such sources provide constant voltage, which means that as long as the
amount of current drawn from the source is within the source's capabilities, its output
voltage stays constant. An ideal voltage source provides no energy when it is loaded by
an open circuit (i.e. an infinite impedance), but approaches infinite power and current
when the load resistance approaches zero (a short circuit). Such a theoretical device
would have a zero ohm output impedance in series with the source. A real-world voltage
source has a very low, but non-zero output impedance: often much less than 1 ohm.
Conversely, a current source provides a constant current, as long as the load connected to
the source terminals has sufficiently low impedance. An ideal current source would
provide no energy to a short circuit and approach infinite energy and voltage as the load
resistance approaches infinity (an open circuit). An ideal current source has an infinite
output impedance in parallel with the source. A real-world current source has a very high,
but finite output impedance. In the case of transistor current sources, impedances of a few
megohms (at DC) are typical.
An ideal current source cannot be connected to an ideal open circuit because this would
create the paradox of running a constant, non-zero current (from the current source)
through an element with a defined zero current (the open circuit). Nor can an ideal
voltage source be connected to an ideal short circuit (R=0), since this would result a
similar paradox of finite non zero voltage across an element with defined zero voltage
(the short circuit).
Because no ideal sources of either variety exist (all real-world examples have finite and
non-zero source impedance), any current source can be considered as a voltage source
with the same source impedance and vice versa. These concepts are dealt with by
Norton's and Thévenin's theorems. Temperature changes will change the output current
delivered by the circuit of Figure 3 because VBE is sensitive to temperature. Temperature
dependence can be compensated using the circuit of Figure 4 that includes a standard
diode D (of the same semiconductor material as the transistor) in series with the Zener
diode as shown in the image on the left. The diode drop (VD) tracks the VBE changes due
to temperature and thus significantly counteracts temperature dependence of the CCS.
Resistance R2 is now calculated as
Since VD = VBE = 0.65
Therefore,
(In practice VD is never exactly equal to VBE and hence it only suppresses the change in
VBE rather than nulling it out.)
and R1 is calculated as
(the compensating diode's forward voltage drop VD appears in
the equation and is typically 0.65 V for silicon devices.[3])
This method is most effective for Zener diodes rated at 5.6 V or more. For breakdown
diodes of less than 5.6 V, the compensating diode is usually not required because the
breakdown mechanism is not as temperature dependent as it is in breakdown diodes
above this voltage.