Please Note: This sheet was written in 1994/95 and needs to be updated.
Cells are chemo-electric generators which means that they convert chemical energy to electrical energy. In general the interaction between different combinations of chemical substances results in the transfer of electrons from one terminal of the cell to the other. As described under 1.3.1: The Chemical Cell the chemical action continues until the charge accumulated at the terminals produces sufficient counter-force to prevent further transfer of electrons.
Cells can be classified in several ways but the first is into the two categories storage cells and fuel cells. A storage cell stores all the chemical constituents within the cell and, when those chemicals are exhausted, the cell ceases to function. The chemicals are not necessarily consumed purely in producing electricity; unwanted side reactions cause the cells to have a limited shelf-life and some of them suffer also a destructive corrosion when they are left for long periods. A fuel cell uses chemicals which are fed to it as required from external containers. There are of course advantages and disadvantages to each type:
(a) Storage cells can be made in very small sizes and are fundamental to portable equipment but their limited storage means that they have a limited useful life and limited output current. Storage cells can be made in large sizes also when they can offer either a large output current or long duration.
(b) Fuel cells, while themselves not necessarily large, can be accompanied by storage containers whose size is chosen according to need and which can be placed as convenient; the size of these containers determines the life of the cells and, together with the cell construction, the current-capacity. This type of generator is used in small space-craft in preference to nuclear reactors because, apart from cost, they do not pose a radiation hazard in the event of an accident.
(c) Different chemicals used in a fuel cell are housed separately and this prevents inter-action between them until they enter the cell; this prevents the chemical side-reactions that so often limit shelf-life of storage cells and cause their destruction.
(d) In general the chemicals used in fuel cells, and the products of their reaction, need to be volatile so that the chemicals can be fed to the cell under proper control and so that the products of the reaction can be removed. Often the usefulness of storage cells is limited more by the failure to remove by-products of chemical reaction than by exhaustion of the basic chemicals.
Storage cells are sub-divided into two main categories:
(i) Primary Cells
(ii) Secondary Cells
Primary cells are those whose useful life ends when they have consumed all the chemicals stored within themselves. In theory all chemical reactions can be operated in reverse if the correct conditions are created and none of the by-products of the reaction are lost; sometimes however the conditions required produce negative returns and the only practical action once a primary cell fails is to throw it away.
These cells utilise chemical reactions that can be reversed under reasonable conditions. Once the cell becomes exhausted an electric current is passed through it in the “wrong” direction to force the chemical action to reverse and so restore the chemica-mix to the original starting condition. This charge-discharge cycle can be repeated many times but ultimately the cell becomes useless either because of mechanical failure in its construction or because abuses of the cell cause a change in the chemical mix.
This sub-division is partly a misnomer in that few, if any, cells will function in the dry state; indeed car batteries are usually sold in the dry condition because it improves their storage life and because there is less hazard from the acid contents.
If atoms and/or their electrons are to move freely about a cell to accomplish the necessary transfer of electrons from the positive terminal to the negative terminal then the various chemicals must be in intimate contact in a medium that permits movement; in other words they must be in solution.
Most of those designated as wet cells usually use either hydrochloric or sulphuric acid but some use Potassium Hydroxide; all these are destructive and dangerous when spilled and most certainly could not be used in portable equipment such as a torch or a radio-receiver.
So-called dry-cell construction prevents spillage by use of a thickening or gelling agent; however, the decreased mobility of the chemical atoms which results limits the current-generating capacity of a cell and also brings into being a problem called polarisation.
A product of the chemical reactions that take place in many primary cells is a gas and the escape of this explains why the reaction cannot be put into reverse; it also explains why such cells cannot be hermetically sealed to prevent spillage.or seepage. With a liquid electrolyte this gas simply forms larger and larger bubbles until their buoyancy causes them to break away from the electrode where they had formed, so producing an insulating layer, and to rise to the surface.
When the electrolyte is either a paste or a gel however such escape is denied and the bubbles remain around the electrode eventually to form an insulating layer around it. This layer has the same effect as removing the electrode and so the cell ceases to function even though it has plenty of potential life remaining.
Polarisation is prevented by surrounding the electrode with a porous mass of material that absorbs the gas; the same technique is used in gas masks. Not surprisingly this material is referred to as the depolariser. The depolariser offers resistance to electron-transfer within the cell and so reduces the current- generating capacity. Furthermore depolarisers need time in which to accomplish their appointed task and so a cell, which fails in use, will often recover after a period of rest and this process is sometimes accelerated by gentle heating.
WARNING — many modern cells are hermetically-sealed and it can be dangerous to heat them.
The phenomenon of polarisation and the presence of a depolariser in dry cells leads to a fundamental difference in the manner of checking them. The high degree of electron-mobility in wet cells enables them to generate very large currents (hundreds of amps) and this poses an obvious hazard. In dry cells the depolariser and its relatively slow action means that only small currents can be produced; for example (when new) the largest cells can produce up to 7 amps while a torch battery may limit at around 1 to 3 amps.
With reference to the previous Section wet cells have a low internal-resistance, about 0.02 ohms or less, while dry cells have internal-resistances between 0.2 and 1 ohm; as the depolariser and the chemicals become exhausted so this internal resistance rises.
Many dry cells have a terminal voltage which decreases steadily throughout the cell’s life; they are said to have a sloping discharge-curve. By contrast most wet cells have a fairly constant terminal voltage until the point of exhaustion when they fail very suddenly; these are said to have a flat discharge curve
To check a dry cell by measuring only the terminal voltage tells little about the internal state of the chemical mix or the depolariser both of which are crucial to the cell’s performance. The only criterion is the ability of the cell to produce current and so it is necessary to measure the short-circuit current by connecting an ammeter directly across the terminals. WARNING This test must be conducted as a “flash-test”; if the meter is connected for longer than about a second the cell IS useless. Do not perform such a test unless it is necessary.
By contrast a wet cell MUST NOT be subjected to a short-circuit; a device called a drop-tester exists whose purpose is to test the terminal voltage under short-circuit conditions. (Given that the resistance of the “short-circuit” is known the short-circuit current can be calculated.) Such a test is useful only on a wet cell which is believed to be at the end of its life; in the hands of an inexperienced operator this can be a very dangerous pastime as abuse may cause the electrolyte to boil. The proper way to test wet cells is to understand their operation and check the terminal-voltage (both on and off charge) and the specific gravity of the electrolyte; see later.
Many modern cells do not give access to the electrolyte and so an intelligent reading of terminal voltage is the only criterion for judging their health.
There are various designs of cells which provide constant terminal voltages and they were used as laboratory standard-references. These cells are not intended to supply current and are used in circuit arrangements where their emf is balanced against a p.d. Today they are largely supplanted by semiconductor reference devices.
An example of such a cell is the Weston cell which delivers 1.0183 volts at 20°C for an almost indefinite period. It uses Cadmium and Mercury for its electrodes and Cadmium Sulphate solution as the electrolyte. It has a temperature coefficient of 40 parts per million per degree C (40 ppm/ °C) but modern semiconductor references can produce temperature stabilities around 1 ppm/ °C.
This cell is the basis of the original dry cell which is now being displaced by the Alkaline cell and various forms of sealed and re-chargeable cells. It consisted of a glass vessel which contained ammonium chloride solution into which was dipped a zinc electrode and a carbon rod; the zinc formed the negative terminal and the carbon rod the positive terminal. A depolariser was used which consisted of manganese dioxide contained around the carbon rod by a porous pot. The chemical reaction consumed the zinc.
To make this into a dry cell the zinc is formed as the container inside which is placed a carbon rod packed around with manganese dioxide; the electrolyte is mixed as a paste instead of being in solution and it is packed between the zinc container and the manganese dioxide.
The problem with this type of cell as it becomes exhausted, either by use or through over-long storage, is that the zinc case becomes perforated and the contents leak.
The off-load terminal voltage (the emf) of the LeClanché cell is about 1.4 volts. It is the cheapest of the available cells but also has the lowest energy-storage density, it does not produce high currents and its life is greatly diminished if discharged at a too high rate. The sloping discharge curve indicates that the internal resistance rises as the cell is discharged. It does not perform at low temperatures.
The "zinc-chloride" cell is an improved version of the zinc-carbon cell which has a slightly better energy-storage capacity and a better performance at high currents and low temperatures. It does not leak and its reduced likelihood of sparking during installation makes it suitable for use in some hazardous-area applications.
Alkaline-manganese cells (known generally as “alkaline cells”) are of sealed construction and any attempt to recharge them is likely to end in an explosion. It is unwise to tamper with these cells because they contain highly toxic substances.
They offer a much longer service life than Leclanché types particularly in equipments which draw large transient currents. However their useful life increases still further at low current-drains. They have a long shelf-life because of their low self-discharge current. Although they have a sloping discharge curve they offer a better performance than LeClanché cells both at high current and low temperature.
The on-load voltage of a fully-charged alkaline cell is 1.5 volts which drops to 0.9 volts when discharged.
The (wet) secondary Nife cell is little used today but saw much service around the time of WWII. Its electrodes are Nickel (Ni) and Iron (Fe) and the electrolyte is Potassium Hydroxide . It had advantages in that it was mechanically very robust, could produce a large current, could be completely discharged and could be subjected to a high charge rate all without damage and had negligible self-discharge; its disadvantages were a low e.m.f., higher internal resistance than the lead-acid cell, a larger size for the same overall voltage and generally greater cost.
These are sealed primary cells. The silver-oxide types provide an e.m.f. of 1.5 volts and the mercuric-oxide types an e.m.f. of 1.35 volts. Both offer 3 to 4 times the energy storage of a LeClanché cell, they have flat discharge curves and excellent shelf life. Both types are good at high temperatures and the silver cell operates down to -20 °C. They are both relatively expensive.
These cells should NOT be re-charged, tampered with or disposed of in a fire.
These are hermetically-sealed primary cells intended generally as a back-up source for low-current circuits and for devices such as clocks. Their shelf-life is typically around 10 to 20 years and they are suited to low-discharge rate applications where they give a level voltage characteristic. They may outlive the equipment into which they are incorporated.
There are several different constructions for Lithium cells, in terms of chemistry, and the information at present in my possession is a bit confused. On-load terminal voltages are between 3.3V and 3.7V when new but they appear to drop to around 2.0 volts at the end of the cell’s life. It is important not to mix them up with similarly-dimensioned cells of other types. Many have been designed for specific applications and it is recommended that they are used only for those purposes.
They have the highest energy storage per unit weight, a flat discharge curve and operate in general from -55°C to 120°C depending on type.
The cells must NOT be subjected to either charging or forced discharging and no more than two should be connected in series. They should NOT be disassembled.
These wet secondary cells are widely used as batteries mainly in road-transport applications but they find application too where large-capacity storage is required with a high-current capability such as the electric car or float.
Their bulk and weight are a great disadvantage for use in operating portable equipment. For such usage the hazard from their sulphuric-acid electrolyte can be mitigated by gelling the electrolyte (which raises the internal resistance) or by holding it in sheets of fibre-glass wadding interleaved with the plates. Valves are sealed into the battery-cover to allow venting of any gases which may be formed by over-charging.
The electrodes are formed as two different oxides of lead and these, in paste form, are packed into recesses formed in the surfaces of lead plates. Positive and negative plates are interleaved (with suitable spacers) so that both sides of the plates are utilised. Too-high charging/discharging rates can result in the formation of hot spots or pockets of gas which dislodge the oxide packings and so destroy the storage capacity of the plates. Additionally the dislodged oxides fall to the bottom of each cell where, as the debris builds up, it eventually short-circuits the cell.
When a lead-acid cell is fully discharged both sets of plates tend to a dirty-grey colour but, when fully charged, one set assumes a chocolate-brown colour and the other a sharp grey. This is a very useful guide to the state of a cell but, unfortunately, many manufacturers now hide them from view with a device meant to make it easier to top-up the electrolyte.
The electrolyte is diluted sulphuric acid which is a very nasty substance. If spilled it can in theory be neutralised with an alkaline substance such as washing soda but in practice it seems able to go on causing corrosion despite all efforts to clean-up. These cells/batteries should always be handled with care.
The degree of dilution is specified in terms of the specific gravity (S.G.) of the solution; in loose terms this compares the weight of a given quantity with the weight of an equal quantity of water. For a fully-charged lead-acid cell the S.G. (as measured with a Hydrometer) should be 1.280 but generally quoted as twelve-eighty. When the cell is discharged its S.G. is around 1.100. Thus the hydrometer provides a very good check on the condition of a cell.
The e.m.f. of a lead-acid cell is nominally 2-volts with an internal resistance that can be as low as 0.02 ohms. (Note that a short-circuit will evoke a current in excess of 100-amps — but that short-circuit is unlikely to last for long!) When fully-charged and still on charge the terminal voltage may rise above 2.4 volts per cell; when left for a while this will generally drop to around 2.2 volts. On load this rapidly falls to the nominal value of 2-volts which is maintained until the battery-charge is exhausted when it falls rapidly (in seconds) toward zero. The end of its useful life is when the terminal voltage falls to around 1.7 volts. However the e.m.f. depends on the specific-gravity of the electrolyte as shown in Table I below.
During the re-charging process the terminal voltage rises fairly quickly to just above 2.2 volts but, as the fully-charged state is reached, this quickly increases to the value given above. Once the oxide coatings of the plates have been restored to their original states any further passage of current is supported by electrolysing the water content of the electrolyte; i.e. the water is broken down into its constituent atoms namely Oxygen and Hydrogen which are lost to the environment.
The escaping bubbles of gas carry with them fine globules of acid, which accounts for the distinctive smell of a charging battery, but the acid causes damage to any surrounding materials. The two gases are present in the correct proportions for them to chemically recombine to form water for which process all they need is heat. The recombination process however releases a great deal of energy and it is not unknown for unventilated battery rooms to be destroyed in an explosion following either a carelessly drawn arc or the entrance of a smoker.
| (showing the Specific Gravity of the sulphuric-acid electrolyte at various temperatures) | |||
| Fully-charged | S.G. | Temp F | Temp C |
| 1268 | 100 | 38 | |
| 1272 | 90 | 32 | |
| 1276 | 80 | 27 | |
| 1280 | 70 | 21 | |
| 1284 | 60 | 16 | |
| 1288 | 50 | 10 | |
| 1292 | 40 | 4 | |
| 1296 | 30 | -1.5 | |
| Fully-discharged | S.G. | Temp F | Temp C |
| 1098 | 100 | 38 | |
| 1102 | 90 | 32 | |
| 1106 | 80 | 27 | |
| 1110 | 70 | 21 | |
| 1114 | 60 | 16 | |
| 1118 | 50 | 10 | |
| 1122 | 40 | 4 | |
| 1126 | 30 | -1.5 | |
Over-charging represents a loss of water from the electrolyte whose concentration is thus increased (its S.G. is increased) and this results in an increase in the cell’s e.m.f. It follows that the battery-charger becomes less successful in forcing current back through the cell and so, because it is unable to achieve a full charge, the cell may appear to have lost some of its storage capacity. The obvious cure is to replace the water (a process known as topping-up) but acid nay be lost also during gassing and so constant topping-up leads to a steady fall in the S.G.
Clearly overcharging is not desirable but inevitable and so occasionally a battery may need to be topped up with acid so as to restore both the S.G. and the terminal voltage to the correct value when fully-charged.
Lead-acid cells should be charged from a constant-voltage source so that, as the terminal voltage reaches its fully-charged value, so charging automatically ceases. However, where many cells are charged in series, constant-current charging may be effective in ensuring an even distribution of charge. A constant-current source however does not provide an automatic end to charging. (More information on this subject is available on request).
The life of a lead-acid cell is limited by the break-up of the plates which both reduces the available plate area (the storage capacity of the cell) and produces debris which may short-circuit the unit. This process is accelerated by high charging and discharging rates. The cells usefulness is also terminated by the formation of a layer of sulphur which insulates the plates from the electrolyte; this is most often caused by allowing the cell to stand for a lengthy period with the electrolyte in place when it will eventually reach a discharged state, For storage purposes lead/acid cells should be fully charged, emptied, given a distilled-water charge and then stored empty and dry.
Little can be done about a cell which has suffered break-up but it is sometimes possible to recover a large part of a cell’s capacity which has been lost through sulphation by use of a chemical process. Briefly this involves cleaning out the cell, washing and then charging it with distilled water and then charging it with a solution of Glaubers Salt. There is a problem however in that there is very little room inside a lead-acid cell and it may be difficult to remove the debris. (Further information on request.)
Lead-acid cells and batteries are available today which really are sealed and safe to use within equipment. It is claimed that they are good for 250-1,000 charge/discharge cycles but this does depend on the use/abuse to which they are subjected.
These secondary (re-chargeable) cells derive their name from the Nickel and Cadmium metals used in their construction. They are sealed and may be operated in any position but it is important not to exceed the charge-rates recommended by the manufacturer. They are fitted with a safety vent to release gases formed during overcharging.
Their nominal on-load voltage is 1.25 volts and they should be regarded as fully discharged when this voltage reaches 1.0 volt. They have internal resistances measured in milliohms and can produced discharge currents up to 60 amps depending on size. On charge they produce terminal voltages between 1.4 and 1.5 volts.
Their discharge time is generally less (about one-half) than that of comparable alkaline types but of course they offer a medium-term saving in operating costs.
There is a belief that these cells should not be recharged unless first fully discharged because they develop a memory such that their capacity falls to the low value of the partial discharge. I cannot find any evidence to corroborate this belief but, if the cells are abused (e.g. by overcharging), internal chemical changes raise the terminal voltage and so preclude proper charging. Many circuits have appeared for so-called cyclic chargers which, if a cell is left permanently attached, constantly discharge and then re-charge the cell thus supposedly keeping it in good condition.
It is worth noting that, unlike lead-acid cells, these cells can be stored for periods in excess of five years without the need for periodic re-charging. Furthermore manufacturers give the cell-life in terms of the number of charge/discharge cycles. It follows that a cyclic charger, far from preserving a nicad, may rapidly reduce a new and unused cell to useless junk. Note too that a cyclic charger may well deliver a completely-discharged cell if called upon at short notice.
Nicad cells should be charged from a constant-current source that, as a rule-of-thumb, produces a complete charge in about 16 hours. Charge-rates are usually expressed in terms of a quantity C which is the storage capacity of the cell expressed in ampere-hours. Thus a charging-rate of C/10 means that the battery would take 10 hours to reach a full charge (from the fully-discharged state). For a 1 A-hr cell this would mean charging at 100 mA; however, at this rate, a discharged cell should be charged for 14 to 16 hours.
Once a nicad cell has become fully-charged it is essential to reduce the charging current to a maximum value about one-tenth the ampere-hour rating (C/1O) else a cell will be permanently damaged. The fully-charged state is indicated by the terminal voltage reaching a maximum value but to accurately determine this point is a complex operation that involves monitoring the voltage level, the rate at which that voltage changes, cell temperature and/or charging time.
(Further information on request).
These are a recent development that offer the advantages over Nicad cells of 50% greater energy storage and that they do not use the toxic Cadmium metal. They offer too the disadvantages of higher cost, need strict control of the charging process to avoid explosion risks of over-charging and that their life is impaired by prolonged trickle-charging.
These lithium-based cells are not yet available at the time of writing but are expected to come into production later in 1995. The many advantages of the Lithium type are offset by the difficulties in constructing a re-chargeable version that does not risk explosion. It is claimed that an SSS cell offers around twice the storage capacity of the standard Nicad cell and nearly three-times the terminal voltage (around 3.5 volts).
The anode is formed of a carbonaceous material and the cathode of lithium-manganese-oxide. The conventional liquid electrolyte is replaced by a conductive plastic membrane which contains a lithium salt although this is not quite so conductive as would be a liquid electrolyte.
The cell may be charged either with a constant-current or a constant-voltage charger but a strict upper voltage-limit must be observed (IEE Review May 1995).
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