CHAPTER 2 PRIMARY AND SECONDARY CELLS

2.1       ORIGINAL VOLTAIC CELL

About the year 1799 Alessandro Volta, an Italian scientist (after whom the unit ‘volt’ is 

named), discovered that, if two dissimilar metals were separated by a certain liquid which was in contact with both, an ‘electromotive force’ (emf), or ‘voltage’, appeared across the metals and, when they were connected together, an electric current flowed.  Volta used zinc and copper discs separated by a piece of cloth moistened with a solution of common salt.   Later the voltage was found to be 1.0V no matter what the size of the metal discs.  Of course at that time no such unit as the ‘volt’ was known.  This was the original ‘Voltaic Cell’ and was used, with modifications, for 150 years as almost the sole source of electric currents for the experimenters of the day. 

Figure 2.1 PRIMARY CELLS

If several such cells are arranged in a pile as in the lower part of Figure 2.1 (as Volta did), each cell is in fact in series with its neighbour, and the voltages of each individual cell add together to give a much higher voltage than could be achieved with only one.  Such a series collection is referred to as a ‘battery of cells’ (by analogy with a battery of guns) - or nowadays simply as a ‘battery’.

2.2       PRIMARY CELLS

Although Volta used zinc and copper as his main elements, almost any two dissimilar 

metals will produce a like result.  Only the voltages would be different, depending on the metals actually used.

Such voltaic cells, whatever the materials used, are termed ‘primary cells.’  They do not have to be charged, as in a modern battery, but derive their electrical energy from chemical action between the metals and the surrounding conducting liquid, called the ‘electrolyte’.  As power is extracted, one of the two metals gradually erodes away as the chemical action proceeds, and the process is not reversible - that is to say, the cell cannot be recharged.  It will continue to give power until the erosion is complete.

An example of the application of primary cells was in the domestic bell systems of some old houses, which used a battery of ‘Leclanche’ cells.  Each consisted of a glass jar filled with sal ammoniac and containing a zinc rod and a carbon plate inside a porous pot.  After prolonged but intermittent use the zinc rod eventually eroded away and had to be replaced.

Certain types of primary cell are still used in laboratories as a voltage reference; they are called ‘standard cells’, and their voltage is very accurate and constant.  One such is the ‘Weston Standard Cell’.  Its electrodes are mercury (positive) and mercury-cadmium amalgam (negative) with an electrolyte of a solution of cadmium sulphate.  The emf of this cell is 1.0183V at 20^oC, varying by only one part in 25 000 per ^oC change.  Such a cell is never used as a source of current, it is purely a very accurate voltage reference.

A common use for the primary cell today is the so-called ‘dry battery’.  This uses zinc and carbon for its two elements, the zinc forming the outer case and the carbon taking the form of a rod down the centre.  The electrolyte is a stiff moist paste packed into the space between the zinc case and carbon rod.  It is not strictly dry, but it cannot spill and can be sealed in the case.

Except in such special applications primary cells are little used in power plants today, being replaced by the ‘secondary cell’, which is rechargeable.

2.3       SECONDARY CELLS

Like the primary cell, the secondary cell, also sometimes called an ‘accumulator’ or ‘storage battery’, consists of two metal plates immersed in a conducting liquid electrolyte.  They fall into two types: ‘lead-acid’ and ‘alkaline’.

Lead-acid Cells

In the original lead-acid cell both plates were of pure lead, immersed in dilute sulphuric acid.  In this state the metals are not dissimilar, and no emf (i.e. voltage) is developed between them.  But after an external ‘charging’ current has been passed through the cell from one plate to the other, a chemical action takes place; one plate becomes covered with spongy lead and the other with lead peroxide.  In this state the cell behaves like a primary one, the spongy lead plate being the negative and the lead peroxide the positive.  When these plates are connected together externally, a current will flow, driven by the cell’s emf or internal voltage.

This construction was soon found to be rather unsatisfactory, as the plates’ coatings rapidly fell off.  Now each plate consists of a grid of lead-antimony alloy, into which the active materials are inserted in the form of a hard paste.

As shown in Figure 2.2, positive and negative groups of plates are interleaved.  The capacity of the cell is determined by the number and total surface area of the plates, which are kept from touching each other by ‘separators’.  These were originally of wood, but nowadays porous plastic is often used.  If, due to too heavy a rate of discharge, some of the active material becomes dislodged from its grid, it will gradually accumulate in the bottom of the cell until its level rises to touch the bottoms of the plates.  This will short-circuit them and rapidly discharge the battery.  If the cell is sealed, its life is finished, but if the plates can be removed, the cells can be flushed out, refilled and used again.

  Figure 2.2 TYPICAL LEAD ACID CELL

The sulphuric acid electrolyte is diluted to a relative density between 1.200 and 1.300, depending on the make and duty of the battery.  When fully charged, an open-circuit voltage of about 2.1V appears across the cell’s terminals.

As power is taken out, derived from the chemical action, the sulphuric acid begins to form lead sulphate at both plates and in doing so loses density.  The cell voltage also falls steadily.  Once the acid density has fallen to 1.150 (the figure differs somewhat with different makes of cell), the cell is regarded as ‘discharged’, and further extraction of power could damage it.

It must then be recharged (not possible with primary cells).  The current is reversed by applying an external voltage to the cell higher than the cell’s own voltage.  The chemical action is thereby reversed, the lead sulphate returning to the electrolyte and raising its density again to its original level, at the same time restoring the cell voltage.

If a cell is left in a discharged state for long periods, the lead sulphate on both plates will harden; the cell is then said to be ‘sulphated’.  In this condition attempts to recharge it may not remove the sulphate coatings, in which case the cell has become unfit for further use.  

Lead-acid cells must never be left for long periods in a discharged state.

The ability to discharge and recharge is a property only of the secondary cell.  Its efficiency is such that it takes about 1.4 times the power to recharge it as can be obtained out of it on discharge.  The lost power appears as heat.

Another property of the secondary cell is that, if discharged too quickly, not only is active material liable to be dislodged from the plates but also the normal chemical processes are upset and limit the output obtainable.  For this reason the rated discharge current is always specified as a certain maximum rate - for example 50A at a 5-hour rate, or 70A at a 10-hour rate.  If these rates are exceeded, the cells will not give their full output for the stated times.

The capacity of a cell is expressed in ‘ampere-hours’ (Ah).  For example, if a battery is rated 300Ah, it should sustain a discharge rate of 1A for 300 hours, or 5A for 60 hours, or 10A for 30 hours, so long as the discharge rate does not exceed the stated rating.  If it does, the rated ampere-hour capacity will not be achieved.

The easiest way to test a lead-acid cell’s state of charge is to measure the electrolyte density with a hydrometer.  The full-charge density varies somewhat between makes of battery, but it is always stated by the maker.  A typical figure is 1.250.

Since a cell can only lose liquid by evaporation or by ‘gassing’, it is only water that is lost, not the acid.  Therefore liquid should only be replaced by topping up with pure (distilled) water, never by acid (unless the acid has actually been spilt).

When a secondary cell is recharged, the power going into it is at first absorbed by causing chemical changes.  But once these changes are complete the power, if continued, starts to electrolyse the water, forming hydrogen and oxygen at the two plates; the cell is said to be ‘gassing’.  Gas bubbling is a sign of a completed charge, but if allowed to continue it will release a considerable quantity of hydrogen and oxygen mixed in its most explosive proportions.

Ventilation of battery rooms is therefore very important, and charging should on no account continue, at least at full rate, unless ventilation is available.  Most platforms now have ventilation monitors which stop, or at least reduce, charging on failure of air flow.

Alkaline Cells

A different class of secondary cell is known as ‘alkaline’.  It was invented by Edison and originally consisted of plates of nickel hydrate and iron oxide immersed in an electrolyte of a solution of potassium hydroxide (with some lithium hydroxide added) in distilled water.  These were the original nickel-iron (or ‘NiFe’) cells; they had a cell voltage of about 1.2V.

The chief advantages of the alkaline call over the lead-acid are:

·         longer life

·         greater reliability

·         less maintenance

·         lighter weight per Ah

·         greater robustness against vibration and shock

·         good high-rate discharge performance

·         ability to accept high rates of charge

·         better charge retention during long periods of rest

·         rapid voltage recovery after heavy discharge

·         immunity from harm if over-discharged.

  

Alkaline cells differ from lead-acid not only in their voltage but also because their electrolyte does not lose density as the cell discharges.  The electrolyte is necessary to sustain the chemical actions but is not itself affected - it is in fact a ‘catalyst’.  The only way to know the state of charge is to measure the cell voltage.  When it has fallen to about 0.8V, the cell is considered to be discharged.

 Figure 2.3 TYPICAL ALKALINE (NICKEL-IRON) CELL

A typical alkaline cell is shown in Figure 2.3.  During charging, the positive nickel hydrate plates become heavily oxidised, whereas the negative iron oxide plates are reduced to pure iron, and the cell voltage rises to about 1.2V.  This amounts simply to the transfer of oxygen from one plate (the positive) to the other and does not call for any chemical changes in the electrolyte.  It is for this reason that in an alkaline cell the electrolyte density does not change with the state of charge.  On discharge the chemical process is reversed, oxygen returning to the iron to re-form it into iron oxide, and in so doing the cell voltage is reduced.

Like the lead-acid cells, alkaline cells are rated in ampere-hours at a specified maximum discharge rate.  However, they are very robust and can stand heavy discharges without damage, though in that case they will not give their full rated ampere-hour output.

More recently the iron in alkaline cells was replaced by cadmium to give the nickel-cadmium (‘Nicad’) cell.  Most platforms and shore installations use this type exclusively rather than the lead-acid or nickel-iron.  It has about 20% higher voltage per cell.

The foregoing description of the nickel-iron cell applies in large part also to the nickel-cadmium, except that the cell’s open-circuit voltage lies between 1.4 and 1.5V, and at 1.1V the cell is regarded as discharged.

The nickel-cadmium cells are manufactured in either plastic or steel containers.  The latter have greater strength against severe vibration and shock and also have advantages when operating in extreme climates.  Plastic containers on the other hand are completely free from corrosion, especially in salt-laden atmospheres, and, being translucent, allow the electrolyte level to be checked at a glance and the electrolyte topped up if necessary.

2.4       CHARGING OF SECONDARY CELLS

 

The apparatus for charging secondary cells and the voltage variations during charge are described in the manual ‘Electrical Distribution Equipment’.