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Early Voltaic Batteries: an Evaluation in Modern Unitsand Application to the Work of Davy and FaradayAllan A. Mills aa Department of Physics and Astronomy , University of Leicester , Leicester LE1 7RH, UKPublished online: 04 Jun 2010.

To cite this article: Allan A. Mills (2003) Early Voltaic Batteries: an Evaluation in Modern Units and Application to the Workof Davy and Faraday, Annals of Science, 60:4, 373-398, DOI: 10.1080/00033790110117566

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A S, 60 (October, 2003), 373–398

Early Voltaic Batteries: an Evaluation in Modern Units and Applicationto the Work of Davy and Faraday

A A.M

Department of Physics and Astronomy, University of Leicester,Leicester LE1 7RH, UK

Received 19 June 2001. Revised paper accepted 15 November 2001

SummaryClassic voltaic batteries of the silver/zinc and copper/zinc types are the ancestorsof today’s primary cells, and facilitated the development of many aspects ofelectrical technology. Nevertheless, they appear never to have been studied andevaluated in a quantitative manner, with results recorded in terms of volts, amps,ohms, and watts. Research of this nature is reported here, and has been conductedfor the most part with copper/zinc cells. Log–log graphs of voltage versus loadand current, and power versus load, are presented for many electrolyte systems.It has been shown that, although the textbook electrolyte of dilute sulphuric aciddoes work, it is an order of magnitude inferior to a solution containing someadditional nitric acid. The latter diminishes the current-limiting phenomenon ofpolarization, and was in fact used by Davy, Faraday, and other early investigators.A quantitative consideration of Nicholson and Carlisle’s discovery of the electro-lysis of water with a silver/zinc voltaic pile is followed by examination of theelectrolysis of pure water, trough batteries, and Davy’s isolation of potassiumand sodium. Every battery gives maximum power when its resistance is adjusted(by appropriate series/parallel connections) to match the resistance of the load:the maximum output of the ‘Great Battery’ of the Royal Institution is assessedat no more than 3 kW. The paper concludes with a note on the recognized hazardof long-term exposure to mercury vapour (produced by amalgamation of zincelectrodes in batteries) and its possible relevance to the health of Michael Faraday.

Contents1. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3742. Some early problems . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 374

2.1. Mode of operation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3742.2. The ‘Baghdad Battery’ . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3752.3. Polarization . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 376

3. Quantification . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3763.1. Volts, amps, resistance, and power . . . . . . . . . . . . . . . . . . . . . . . . . . 3763.2. Electrodes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 377

4. Apparatus . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3774.1. Practical electrodes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3774.2. Instruments . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3774.3. Cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 378

5. Preliminary trials. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3786. Cu/Zn open-circuit potentials with various electrolytes . . . . . . . . . . . . . 3797. Cells in series: the influence of various electrolytes . . . . . . . . . . . . . . . . 379

7.1. Voltage versus load . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3817.2. Voltage versus current . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3827.3. Power . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 382

8. The maximum power theorem . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 382

Annals of Science ISSN 0003-3790 print/ISSN 1464-505X online © 2003 Taylor & Francis Ltdhttp://www.tandf.co.uk/journals

DOI: 10.1080/00033790110117566

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9. Cells in parallel: influence of electrode area . . . . . . . . . . . . . . . . . . . . . . . 3839.1. Sulphuric acid . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3859.2. Nitro-sulphuric acid . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3869.3. Modern ‘dry’ cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 387

10. Some well-known early batteries and applications . . . . . . . . . . . . . . . . . 38810.1. Nicholson and Carlisle’s voltaic pile . . . . . . . . . . . . . . . . . . . . . . . . 38810.2. Humphry Davy: use of nitric acid in voltaic cells . . . . . . . . . . . . 38910.3. Humphry Davy: the composition of water . . . . . . . . . . . . . . . . . . 39010.4. Trough batteries . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 39010.5. Davy’s isolation of potassium and sodium . . . . . . . . . . . . . . . . . . 39110.6. The Royal Institution battery . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 39510.7. Nitric acid and load-matching . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 396

11. Faraday and mercury poisoning . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 39612. Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 397

1. IntroductionThe announcement by Volta1 in 1800 that he had succeeded in obtaining electri-

city from a pile of pairs of silver coins and zinc discs separated by cloth moistenedwith salt water generated immense world-wide interest. His ‘voltaic pile’ was widelyreplicated, and research aimed at confirming the identity of ‘ordinary’ (i.e. frictional )electricity and this new form was soon successfully under way. However, even beforethe formal publication of Volta’s discovery, a lucky accident whereby the conductorsof a pile were bridged by a drop of water led to the discovery of electrolysis byNicholson and Carlisle.2

A desire to increase the amount of electricity generated was soon felt, and thebest pathway to achieve this appeared to be development of Volta’s second invention,the ‘crown of cups’. The basic unit (cell ) of this arrangement was a pair of dissimilarmetals, usually copper/zinc rather than the expensive silver/zinc, immersed in brinewithin a small glass vessel. A number of individual cells were then connectedsequentially in what we now term ‘series’ to form a ‘battery’. It was soon foundthat dilute inorganic acids performed better than brine, dilute sulphuric acid beingmost often mentioned by early reviewers such as Roget.3 This has been copied fromtextbook to textbook right up to the present day, so the ‘elementary cell’ is commonlyconsidered to consist of ‘electrodes’ of copper and zinc immersed in dilute sulphuricacid. It is sometimes stated to give an electromotive force (e.m.f.) of about 1 V.

2. Some early problems2.1. Mode of operation

In the early years there was, of course, no accepted theory behind the observa-tions. It was known that ordinary zinc dissolved in dilute sulphuric acid, the reactionbeing symbolized as

Zn+H2SO4�ZnSO4+H2It was suspected that the reaction probably occurred in steps, and that the chemical

1 Alessandro Volta ‘On the Electricity Excited by the Mere Contact of Conducting Substances ofDifferent Kinds’, Philosophical Transactions, 90 (1800), 403–31 (in French; English trans. in PhilosophicalMagazine, 7 (1800), 289–311).2William Nicholson, ‘Account of the New Electrical or Galvanic Apparatus of Sig. Alex. Volta, and

Experiments Performed with the Same’, Nicholson’s Journal, 4 (1800), 179–87. W. Nicholson, A. Carlisleand W. Cruickshank, ‘Experiments in Galvanic Electricity’, Philosophical Magazine, 7 (1800), 337–47.3 P. M. Roget, Electricity, Galvanism and Electro-Magnetism (London, 1832).

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Early Voltaic Batteries: an Evaluation in Modern Units 375

equation merely represented their summation. This view was given greater credencewhen it was found that pure zinc was not attacked by the acid unless another plateof a more noble—and therefore unattacked—metal (such as copper, silver, orplatinum) was also placed in the solution. When connected by a wire to the zinc itallowed a ‘charge’ to flow and the reaction could proceed, with dissolution of thezinc and evolution of hydrogen at the noble metal electrode.4 The matter wascarefully studied, quantified and systematized by Faraday.5 For modern ideas,reference must be made to appropriate textbooks.6

The fact that an electrode of commercial zinc in a voltaic cell slowly dissolved,even when the cell was not in use, proved a considerable nuisance. It was ascribedto impurities such as copper and iron setting up a multitude of tiny cells on thesurface of the metal, and was termed ‘local action’ (see note 5). This wastefulreaction could be prevented by lifting out the electrodes,7 but then of course the cellcould not be used in an ‘on-demand’ situation. However, it was discovered thatamalgamation with mercury8 allowed only zinc to migrate through and to reach theinterface with the ‘electrolyte’. Amalgamation did not affect the potentials andcurrents generated: only the life of the cell when not being called on to generateelectricity.

2.2. The ‘Baghdad Battery’Consideration9 of an ancient artefact that, in my view, has been mistakenly

identified as an electrochemical cell led to an attempt to operate a nominal 2.5 Vtorch lamp with a five-cell Cu/Fe battery containing dilute acetic acid. The freshlyprepared series arrangement registered an open-circuit potential of 2.6 V on anelectrometer (a voltmeter of near-infinite internal resistance) but this fell to nearzero as soon as the lamp was connected across the terminals. It proved impossibleto obtain even a transient glow of the filament. Much the same result was obtainedwith Cu |Zn |sulphuric acid cells. So, the textbook battery registers a satisfactory‘standing potential’ on open circuit, but appears hardly capable of supplying suffi-cient power to do significant work. How then could Davy10 have electrolysed fused

4William H. Wollaston, ‘Experiments on the Chemical Production and Agency of Electricity’,Philosophical Transactions, 91 (1801), 427–34.5Michael Faraday, Experimental Researches in Electricity, 2 vols (London: Quaritch, 1839).6 Peter Atkins, The Elements of Physical Chemistry, 2nd edn (Oxford: Oxford University Press, 2001),

Chapter 9.7 C. H. Wilkinson, ‘Description of an Improved Galvanic Trough’, Philosophical Magazine, 29

(1807), 243–44.8 Humphry Davy, ‘On the Relations of Electrical and Chemical Changes’, Philosophical Transactions,

116 (1826), 383–422; K. T. Kemp, ‘Description of a New Kind of Galvanic Pile, and also of AnotherGalvanic Apparatus in the Form of a Trough’, Edinburgh New Philosophical Journal, 6 (1828), 70–77;Auguste de la Rive, ‘Note relative a l’action qu’exerce sur le zinc l’acide sulfurique etendu d’eau’,Bibliotheque Universelle, Sciences et Arts, 43 (1830), 391–411; William Sturgeon, ‘On Electromagnets’,Philosophical Magazine, 11 (1832), 194–205; K. T. Kemp, ‘Voltaic Batteries with Amalgamated Zinc’,Sturgeon’s Annals of Electricity, 1 (1837), 81–88.9 Allan A. Mills, ‘The Baghdad Battery’, Bulletin of the Scientific Instrument Society (68) (2001) 35–37.10 Humphry Davy, ‘On Some New Phenomena of Chemical Changes Produced by Electricity,

Particularly the Decomposition of the Fixed Alkalies, and the Exhibition of the New Substances WhichConstitute their Bases; and on the General Nature of Alkaline Bodies’, Philosophical Transactions, 98(1808) 1–44; idem, ‘Electrochemical Researches in the Decomposition of the Earths; with Observationson the Metals obtained from the Alkaline Earths, and on the Amalgam Procured from Ammonia’,Philosophical Transactions, 98 (1808), 333–70; idem, ‘On Some New Electrochemical Researches, onVarious Objects, Particularly the Metallic Bodies, from the Alkalies and Earths, and on Some Combinationof Hydrogene’, Philosophical Transactions, 100 (1810), 16–74.

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sodium and potassium hydroxides to release the alkali metals, Faraday11 run thefirst electric motor, or Sturgeon12 operated a powerful electromagnet?

2.3. PolarizationThe problem of the ‘pressure’ of electricity almost disappearing as soon as a

load was connected to a battery was encountered very early in the history of electro-chemistry. It was ascribed to the film of hydrogen which was observed to form onthe silver or copper electrode of a working cell, obstructing access of reactants andincreasing cell resistance. It was also found that a failing cell which had just beendisconnected from a load could generate a potential in the reverse direction for abrief period. The phenomenon was therefore termed ‘polarization’, and attempts tocounter it were made by adding oxidizing agents to the electrolyte, using sine andtangent galvanometers as monitoring instruments.13

Nowadays, polarization is mostly avoided by arranging that the ion depositedis not hydrogen, but at high discharge rates some loss of e.m.f. still occurs as aresult of the time taken for ions to diffuse through the electrolyte. It remains animportant factor in the technical design of successful dry cells.14

3. QuantificationThe history of electricity is considered in most works on the history of science,

and early cells and batteries are specifically discussed in a number of reviews.15 Aweakness is that the subject appears to be treated in a qualitative descriptive mannerbased on literature and museum artefacts. Many unrecognized anomalies thereforeexist, such as those quoted above. A quantitative assessment, in modern units,appeared to be long overdue. This necessitated the construction of working voltaiccells and replica batteries, for obviously the few existing historical examples couldnot be filled with acids.

3.1. Volts, amps, resistance, and powerThe earliest experimenters16 recognized that the intensity or pressure of a flow

of electric charge must be distinguished from its overall quantity, just like the flow

11Michael Faraday, Faraday’s Diary, ed. by Thomas Martin (London: Bell & Sons, 1932), , 50 ff(4 September 1821 onwards); idem, Quarterly Journal of Science, 12 (1821), 186–283; idem, ExperimentalResearches in Electricity, 3 vols (London: Quaritch, 1839–55), , 147–51.12William Sturgeon (note 8).13 J. B. Rogers and J. Green, ‘Experiments with the Elementary Voltaic Battery’, Silliman’s Journal,

28 (1835), 33–42; C. M. S. Pouillet, ‘Memoire sur la pile de Volta et sur la loi generale de l’intensite queprennent les courants’, Comptes rendus de l’Academie des Sciences, Paris, 4 (1837), 267–79; E. Branly,‘Mesure de la polarisation dans l’element voltaique’, Comptes rendus de l’Academie des Sciences, Paris,74 (1872), 528–31.14 C. A. Vincent and B. Scrosati, Modern Batteries: An Introduction to Electrochemical Power Sources

(London: Arnold, 1997).15 John Bostock, An Account of the History and Present State of Galvanism (London, 1818); P. M.

Roget (note 3); Henry Noad, Lectures on Electricity (London, 1844); F. C. Bakewell, A Manual ofElectricity, Practical and Theoretical (London: Griffin, 1859); J. J. Fahie, ‘Magnetism, Electricity andElectromagnetism up to the Time of the Crowning Work of Michael Faraday in 1831’, Journal of theInstitute of Electrical Engineers, 69 (1931), 1331–57; J. King, ‘The Development of Electrical Technologyin the 19th Century. I: The Electrochemical Cell and the Electromagnet’ (Bulletin 228, US NationalMuseum, Washington, DC, 1962).16W. H. Wollaston for example. See note 4.

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of water in a stream. Davy17 and Children18 suspected that the former varied withthe number of cells in a series-connected battery, while the latter was related to thecumulative area of the electrodes.

Today we measure intensity (potential difference, p.d.) in ‘volts’, V, as a tributeto Volta. The amount (or current) is measured in ‘amperes’, A, to commemorateAmpere. It took years, however, before Ohm was able to convince the scientificcommunity that the ratio volts:amps was constant for a simple circuit. The constantof proportionality is now called the ‘ohm’, V, in his honour, and is a measure ofthe resistance of a circuit. Power in d.c. circuits is given by the product ofvolts×amps, termed ‘watts’, W.

It is not intended to develop the history of how these units were originallyidentified and measured: reference must be made to standard textbooks. None ofthese units had, of course, been defined when the earliest batteries were evolving.

3.2. ElectrodesIt was recognized very early on19 that an electrochemically active couple consisted

of a combination of a comparatively noble with a relatively reactive metal. Copper,silver, and platinum are representative of noble metals (to which electrical groupmust be added graphitic carbon) while zinc was the most reactive of the metalscommonly available in the early nineteenth century. Tables of ‘standard electrodepotentials’ were later established, but are no more than a guide to possible combina-tions. This is partly because the situation in a practical working cell is far fromstandard or reversible, but also because the purity, surface conditions, and pasthistory of the electrodes exert a very considerable influence.

4. ApparatusThe only way to begin to understand the performance and evolution of early

batteries appeared to be to carry out experimental measurements, using modernunits, on Cu/Zn and other couples. These must be immersed in the various electro-lytes that had been proposed. Not only should these measurements include ‘open-circuit’ potentials, but also the voltages, currents, and work done when connectedto loads of known resistance.

4.1. Practical electrodes$ Commercial zinc and copper sheet, 0.03 inch thick, as used for roofing.$ Sterling silver nineteenth-century coins.$ Graphitized carbon rod extracted from modern dry cell batteries.

4.2. Instruments$ Digital multimeter (Black Star 3225) reading from milliamps to 10 A, and

millivolts to 10 V.

17 Humphry Davy, ‘On some Chemical Agencies of Electricity’, Philosophical Transactions, 97(1807), 1–56.18 J. G. Children, ‘An Account of Some Experiments, Performed with a View to Ascertain the Most

Advantageous Method of Constructing a Voltaic Apparatus, for the Purpose of Chemical Research’,Philosophical Transactions, 99 (1809), 32–38.19W. H. Wollaston (note 4); Humphry Davy, ‘An Account of Some Galvanic Combinations, Formed

by the Arrangement of Singly Metallic Plates and Fluids, Analogous to the New Galvanic Apparatus ofMr Volta’, Philosophical Transactions, 91 (1801), 397–402.

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$ Electrometer (Levell TM9B) reading from microvolts to 10 V.$ A specially constructed resistance box containing selectable wire resistors in

the logarithmic series 1, 3, 10, 30, 100, 300 and 1000 V. These were capableof dissipating 20 W of power.

4.3. CellsA number of identical cells were constructed to the pattern diagrammed in

Figure 1, the polyvinyl chloride (PVC) spacers being attached with a spot of epoxyadhesive. Imperial units were used, to be concordant with historical cells. A colour-coded insulated flexible wire was soldered to each electrode. For use, each pair wasplaced in an individual Pyrex beaker, and a chosen electrolyte added to a depth of2 inches.

5. Preliminary trialsOne published recipe for ‘battery acid’ stipulates concentrated sulphuric acid

diluted with 12 times its weight of water, giving a concentration of 4.6% by volume.This solution was made by slowly stirring 230 ml of ‘Analar’ acid into 5 litres ofcold tap water, and then allowing the mixture to cool overnight to room temperaturebefore storing in a polythene carboy. This electrolyte was used in some prelimin-ary trials:

$ Evolution of hydrogen. When one of the above commercial zinc electrodes wasplaced in the dilute sulphuric acid, hydrogen was evolved from its surface,while a copper electrode appeared inert. However, when the pair were joinedby a wire, or touched together beneath the surface of the electrolyte, hydrogenevolution transferred to the copper. The zinc became coated with a finelydivided black coating of copper, and a much lower amount of hydrogen wasevolved from its surface. Present-day safety regulations concerning mercurymade trials with amalgamated zinc electrodes impracticable. (See below.)

$ Variation of p.d. with time. A freshly abraded virgin Cu/Zn pair made up asabove was connected to the digital multimeter, the latter being switched to

Figure 1. Construction and dimensions of copper/zinc electrode pairs.

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Early Voltaic Batteries: an Evaluation in Modern Units 379

indicate millivolts. The electrodes were then plunged into the standard dilutesulphuric acid, and the p.d. from time zero was noted. An initial voltage of1.05 V fell rapidly along a descending exponential curve, appearing to bottomout at 0.88 V between the fifth and tenth minutes. However, this voltage wasnot stable: agitation dislodged bubbles and caused it to rise temporarily. After24 h without disturbance, when a uniform black film coated the zinc, the p.d.appeared fairly stable at 0.96 V.

$ Series connection. Cu/Zn pairs were connected one by one in series, and theopen-circuit potential was measured in each case with the electrometer. Alwaysthe overall p.d. was additive, and corresponded to 0.96 V per cell. This wasindependent of electrode area.

$ Shared electrolyte. Placing a number of series-connected cells into a commonbath of electrolyte caused the potential to fall to close to that of a single cell.The interior electrodes are shorted out, and only the extremities remain effect-ive. It is therefore essential to isolate each metallic couple within its ownwatertight container of electrolyte: all historical batteries did this.

$ Other electrode pairs. Silver |zinc |H2SO4 gave an open circuit p.d. of 1.01 V,while carbon |zinc |H2SO4 climbed to 1.34 V.

6. Cu/Zn open-circuit potentials with various electrolytesA number of electrolytes were made up in tap water (Table 1). Faraday (note 5)

initially used 4.5% v/v sulphuric acid, later preferring 2.5% v/v plus 2% nitric acid.Exact quantities of other solutes are rarely given in the older literature, so concentra-tions of 4.6% by volume for liquids, and 4.6% by weight for solids, were employed.

The open-circuit potentials given in the final column were derived from measure-ments on five Cu/Zn cells in series, using the circuit given below. This was judgedto give a more accurate figure, averaging out any cell-to-cell variation. It was alwayshigher than when a load was connected across the terminals of the battery. Itappeared to be comparatively stable and reproducible for a given assembly, butsometimes apparently identical single cells gave slightly different numerical values.

7. Cells in series: the influence of various electrolytesFive of the Cu/Zn couples shown in Figure 1 were secured at intervals along a

Perspex bar and connected in series. The assembly was then wired into the circuitshown in Figure 2.

A chosen electrolyte was then added to each cell with the switch open, so thatthe electrometer indicated the open-circuit p.d. of the battery. One-fifth of thisvoltage is shown in Table 1. The resistance box was then set to 1000 V, the switchclosed, and both meters were read. Alternative resistors of 300, 100, 30, 10, 3, and1 V were then inserted in this order, and the corresponding voltage and currentsnoted. Periods of at least 2 min were allowed for the initially falling values tostabilize, but reproducibility was hard to achieve. A second run (i.e. after a cell orbattery had been connected to the lowest resistance) always gave lower—but morereproducible—results. These values were adopted for the work recorded here, thenumerical data being concisely presented as log–log plots in Figures 3–5.

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Table 1. Electrolytes used.

References P.d. per Cu/ZnElectrolyte Composition (Note numbers) couple (V )

1 Sulphuric acid 3–5,15,20,21 0.962 Hydrochloric acid 22 0.823 Nitric acid 21–23 0.724 (1)+sodium dichromate 24–26 1.025 (1)+potassium permanganate 21 1.006 (1)+nitric acid 5,12,18,19,27 0.827 (2)+nitric acid 22 0.648 (1)+alum 17 0.949 (3)+alum 28,29 0.62

10 Alum 17 1.0211 Potassium hydroxide 5,21 1.1212 Citric acid — 0.9213 Acetic acid — 1.0014 Sodium chloride 1,3,17,21 0.7815 Ammonium chloride 30 0.76

Figure 2. Circuit for measuring the characteristics of various batteries.

20 Robert Hare, ‘Calorimotor’, Silliman’s Journal, 1 (1818), 413–23.21 James P. Joule, ‘On the Intermittent Character of the Voltaic Current in Certain Cases of

Electrolysis; and on the Intensities of Various Voltaic Arrangements’, Philosophical Magazine, 24(1844), 106–15.22 J. B. Rogers and J. Green (note 13).23William Cruickshank, ‘Description of Mr Pepys’ Large Galvanic Apparatus’, Philosophical

Magazine, 15 (1803), 94–96.24 Robert Bunsen, ‘Ueber die Anwendung der Kohle zur Volta’schen Batterie’, Annalen der Physik,

54 (1841), 417–20.25 R. Warrington, ‘On the Employment of Chromic Acid as an Agent in Voltaic Arrangements’,

Philosophical Magazine, 20 (1842), 393–95.26 J. C. Poggendorff, ‘Uber die mit Chromsaure konstruirten galvanischen Ketten’, Annalen der Physik,

57 (1842), 101–11.27 J. G. Children, ‘An Account of Some Experiments with a Large Voltaic Battery’, Philosophical

Transactions, 105 (1815), 363–74.28 H. Davy (note 10, 1808).29 H. Davy (note 10, 1810).30W. Cruickshank, ‘Additional Remark on Galvanic Electricity’, Nicholson’s Journal, 4 (1801),

254–64.

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Early Voltaic Batteries: an Evaluation in Modern Units 381

7.1. Voltage versus loadThe variation of voltage of the five-cell Cu/Zn batteries with external load is

shown in Figure 3. The following can be observed.

(a) Potentials always dropped as soon as a load was connected, the degree ofdiminution increasing as the load resistance decreased.

(b) Three families of curves may be distinguished. In terms of decreasing e.m.f.we have the following: electrolytes 3–7, strong mineral acids, either them-selves oxidizing or with oxidizing additives; electrolytes 1 and 2, dilutesulphuric and hydrochloric acids; electrolytes 8–15, organic acids, salts, bases.

(c) The oxidizing electrolytes 3–7 are able to remove nascent hydrogen byoxidizing it to water (see note 5) so can act as depolarizers to remove ordiminish the film on the copper electrode. These electrolytes therefore exhibitcomparatively good voltage versus load characteristics at all loads. A mixtureof sulphuric and nitric acids discovered empirically by Davy (see note 19; hereferred to the second component as ‘nitrous’ acid) was especially favouredby early electricians.

(d) Dilute sulphuric (1) and hydrochloric acid (2) are comparable electrolytes.Their curves tend to be an order of magnitude below those of the oxidizing

Figure 3. Voltage versus load behaviour of a five-cell Cu/Zn battery filled with variouselectrolytes identified in Table 1.

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mixtures, presumably being dependent on diffusion and aerial oxidation fordepolarization.

(e) Organic acids and bases may show a reasonable open-circuit potential, butvoltage falls rapidly even when resistances as high as 1000 V are connectedacross the electrodes. The curves tend to be two orders of magnitude inferiorto oxidizing strong acids.

(f ) Alum (10) shows anomalous behaviour, holding up rather better than mostsalts. Perhaps this is why it was sometimes used by Davy (see notes 17, 28,and 29) but it was found here that adding it to sulphuric or nitric acidsalways reduced their performance. (Curves omitted for clarity.)

(g) The salts sodium chloride (14) and ammonium chloride (15) are extremelypoor electrolytes with the Cu/Zn couple. Nevertheless, their potential doesnot fall completely to zero, so they could be (and were) used with high-resistance loads and/or sensitive detectors. (A frog’s leg is a remarkablysensitive detector.) Corrosion currents can be important in the degradationof structures built of more than one metal, especially in the marine environ-ment. Thus, Davy’s idea of fixing zinc plates in contact with the coppersheathing of wooden ships certainly protected the copper against dissolutionas cupric ions—but had the unfortunate result of then allowing barnaclelarvae to settle and flourish.

7.2. Voltage versus currentFigure 4 summarizes the voltage versus current behaviour of the five-cell Cu/Zn

battery with various electrolytes. It will be seen that the disturbing effect of varyingpolarization makes these parameters far from proportional, one reason why Ohmhad so much difficulty getting his relationship accepted. Three families of curves areagain apparent, with oxidizing strong acids much the best.

7.3. PowerPower supplied is often the most important criterion, and milliwatts versus load

are plotted on log–log axes in Figure 5. As would be expected, the oxidizing strongacids are much the best, the five-cell sulphuric acid/sodium dichromate batterypeaking with some 2 W into a 3–10 V external load. Nitric acid alone, or its mixturewith sulphuric acid (‘nitro-sulphuric acid’), is, however, satisfactory, and much morepleasant to use. Pure dilute sulphuric and hydrochloric acids work, but tend to bean order of magnitude less efficient. Other electrolytes are four orders of magnitudeinferior, and are typified by sodium chloride in Figure 6.

8. The maximum power theoremIt was soon observed that a particular arrangement of batteries best suited a

given load. Thus, in 1831, Henry31 reported that the pull of an electromagnet couldbe increased by winding more turns of wire upon it—but only up to a certain limit.After that point, either the additional turns had to be connected in parallel withthose already present, or a battery with more pairs of plates had to be utilized.

According to King (note 15) the existence of an optimum match was first clearly

31 Joseph Henry, ‘On the Application of the Principle of the Galvanic Multiplier to ElectromagneticApparatus, and also to the Development of Great Magnetic Power in Soft Iron, with a Small GalvanicElement’, American Journal of Science, 19 (1831), 400–08.

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Early Voltaic Batteries: an Evaluation in Modern Units 383

Figure 4. Voltage versus current characteristics of five-cell Cu/Zn battery with electrolytes1–15, omitting 8 and 9.

recognized by Jacobi in 1844. Now known as the maximum power theorem,32 itstates that the power dissipated in an external load is greatest when its resistance isequal to that of the cell, battery, or other source of e.m.f.

Figure 5 indicates that the effective electrolytes 1–5, when used in this particularsize of Cu/Zn battery, give maximum power with loads of a few ohms. Nitro-sulphuric acid (6) was exceptional in peaking off the graph at about 0.5 V. Thesevalues for internal cell resistance are in general agreement with those deduced byapplication of Kirchhoff ’s laws.33

9. Cells in parallel: influence of electrode areaThe simplest voltaic cells contain rectangular copper and zinc electrodes of

identical size, disposed parallel to one another and separated by a narrow gap0.25–0.5 inches wide.34 There is, however, no a priori reason why this should be theoptimum arrangement: early researchers simply copied Volta, who in turn was(unconsciously?) influenced by his pile of discs. In the course of time dissimilar

32 G. R. Noakes, A Textbook of Electricity and Magnetism (London: Macmillan, 1948), pp. 147–48.33 Ibid., pp. 141–46.34 F. C. Bakewell (note 15).

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Figure 5. Power versus external load characteristics for a five-cell Cu/Zn battery with electro-lytes 1–7.

dimensions and other dispositions were considered,35,36 but the classic system ofparallel electrodes of identical size stood the test of practicality and economy.

The wide-ranging curiosity of William Wollaston did, however, cause him toquestion the contribution of the side of the zinc facing away from the copper. Hesuggested placing a second copper anode, electrically connected to the first, on thefar side of the zinc.37 The latter would therefore be suspended between two copperanodes, care being taken that an electrolyte-filled gap was maintained around allthree electrodes. Subsequent workers tried this ‘double-copper’ arrangement, andreported apparent improvements in performance (notes 5 and 37).

Quantitative investigations of these area-related parameters were made with newsupplies of the 0.03-inch zinc and copper sheet used in the ‘series’ investigationsabove. The metals were abraded with medium-grade alumina paper, and thenguillotined into strips 14 inches long by 1, 2, 4, and 6 inches wide. Flexible insulatedwires were soldered to a narrow edge of each. The electrodes were then assembledinto the following configurations using 0.25-inch PVC spacers and nylon screws:

35 J. B. Rogers and J. Green (note 13).36 C. Binks, ‘On Some of the Phaenomena and Laws of Action of Voltaic Electricity, and on the

Construction of Voltaic Batteries’, Philosophical Magazine, 11 (1837), 68–89.37 J. G. Children (note 27).

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Early Voltaic Batteries: an Evaluation in Modern Units 385

Figure 6. Power versus external load for a five-cell Cu/Zn battery with a sodium chlorideelectrolyte (14).

(a) simple Cu–Zn pairs;(b) ‘double-copper’ assemblies, Cu–Zn–Cu;(c) Cu–Zn pairs, but with the outside face of the zinc covered with waterproof

self-adhesive vinyl sheet.

Each assembly was clamped vertically in an individual glass measuring cylinder, andan electrolyte added to give an immersed depth of 12 inches. Two electrolytes werechosen from Table 1: dilute sulphuric acid (1) and nitro-sulphuric acid (6). Thecircuit diagrammed in Figure 2 was then used to measure e.m.f. and current atvarious loads.

9.1. Sulphuric acidThe numerical results are again too numerous to quote, but are summarized in

Figure 7. The susceptibility of this electrolyte to current-dependent polarization gavea rather complex picture, but the following were found.

(a) The greater the area (as would also be obtained by connecting identicalsingle cells in parallel ), the better the e.m.f. (and so current) holds up as theload resistance decreases.

(b) A ‘double-coppered’ zinc strip is nearly equivalent in performance to a‘single-copper’ cell of twice its width (e.g. a 2-inch ‘double copper’ cellapproximates a 4-inch ‘single copper’ cell ).

(c) Covering the far face of the zinc electrode in a ‘single-copper’ cell with aninsulating film produces only a slightly deleterious effect.

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Figure 7. Voltage versus load for Cu |Zn |H2SO4 cells of constant length (12 inches) butvarious widths, showing the influence of surface area: curve a, 6-inch double copper;curve b, 6-inch single and 4-inch double copper; curve c, 4-inch single copper; curved, 2-inch double copper; curve e, 1-inch single copper.

It does then appear that Wollaston’s speculation was correct: the far face of the zinccontributes only a little (estimated at 10%) to the current-holding capacity of thesimple cell. ‘Double coppers’ greatly improve performance. Nevertheless, the factthat a black film appears on both faces of the zinc in the elementary cell shows thatcopper ions do reach the distal face.

9.2. Nitro-sulphuric acidThis oxidizing electrolyte gives a simpler picture (Figure 8). Elimination of

polarization at the lower current levels results in a horizontal line of constant voltage

Figure 8. Voltage versus load for Cu/Zn electrodes of constant length (12 inches) but varyingwidth immersed in nitro-sulphuric acid (6): curve a, 4-inch single and 2-inch doublecopper; curve b, 1-inch double copper; curve c, 1-inch single copper. Also shows thebehaviour of modern dry cells.

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Early Voltaic Batteries: an Evaluation in Modern Units 387

(820 mV ) for both ‘single’ and ‘double’ electrodes of all widths. However, when theexternal load is less than 30 V and the current taken exceeds some 50 mA, the e.m.f.delivered by the 1 inch × 12 inch single electrodes begins to fall. The wider anddoubled arrangements maintain the potential for the greater currents passed by 10 Vloads, up to about 150 mA with 4-inch ‘double coppers’. Figures 9 and 10 showhow power continues to be delivered into still lower external resistance loads(10–1 V), but the e.m.f. can drop to as low as 600 mV.The advantages of increased area and ‘double coppers’ in maintaining the voltageare again illustrated. Of course, larger (and thicker) zincs also last longer when thecell or battery is delivering a given power.

9.3. Modern ‘dry’ cellsThese consist of a graphitic carbon rod surrounded by a concentric cylinder of

zinc foil. The annular gap is filled with an immobilized moist electrolyte containingcomplex zinc–ammonium ions. For comparison with the early cells, ‘Duracell’ brandcells of various sizes were tested. (It is common to refer to single cells as ‘batteries’nowadays.)

Results for a new D-size cell (the largest commonly available) and the tiny AAAcell (which had been in storage for many months) have been added to Figure 8. Itwill be seen that, as before, the larger the physical size the better the e.m.f. standsup under conditions of increasing current drain. Storage resulted in a generaldepression of the e.m.f. available.

Figure 9. Current versus voltage for Cu/Zn electrodes of various widths immersed 12 inchesin nitro-sulphuric acid (6): curve a, 2-inch double copper; curve b, 2-inch singlecopper; curve c, 1-inch double copper; curve d, 1-inch single copper.

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Figure 10. Power versus load for Cu/Zn electrodes of various widths immersed 12 inches innitro-sulphuric acid (6): curve a, 2-inch double copper; curve b, 2-inch single copper;curve c, 1-inch double copper; curve d, 1-inch single copper.

10. Some well-known early batteries and applicationsConventional historians (historians of events) cannot travel back in time to

investigate their subjects firsthand, or alter their environment, so are necessarilygreatly dependent on documentary evidence. Historians of science are potentiallymore favourably placed: scientific experiments should (by definition) be repeatableby anyone at any time, with identical results. Regrettably, this valuable additionalmode of investigation is rarely practised. Some classic experiments in the history ofelectrochemistry will now be examined, applying the data gained in the aboveexperimental sections to supplement the information given in well-known biographiesand reviews.

10.1. Nicholson and Carlisle’s voltaic pileVery soon after learning of Volta’s work, the surgeon Anthony Carlisle observed

the evolution of minute bubbles of gas around an iron wire placed on one plate ofa voltaic pile when a drop of water was used to improve the connection. This pileconsisted of 36 sterling silver half-crowns separated by corresponding 1.25-inchdiameter discs of zinc and pasteboard, the latter moistened with saline. WilliamNicholson suspected the gas to be hydrogen (from its smell ).

Together, Nicholson and Carlisle38 pursued this chance observation withimproved apparatus. This incorporated brass wires inserted through corks closingthe ends of a 0.5-inch glass tube filled with water. The tips of the wires werepositioned about 1.75 inches apart. A fine stream of bubbles was observed to risefrom the wire connected to the silver termination of the pile, while the other brass

38 See note 2.

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Early Voltaic Batteries: an Evaluation in Modern Units 389

wire became tarnished and oxidized. The evolved gas was confirmed to be hydrogenby its inflammability. Repetition with platinum wires gave gaseous oxygen in additionto the hydrogen previously observed.

A quantitative experiment was then conducted with a freshly assembled pilecontaining 68 Ag/Zn pairs connected to platinum electrodes in water. This time,however, arrangements were made to collect both the evolved gases. The authorsstate that the hydrogen released over a period of 13 h displaced 142 grains of water,equivalent to a volume of 9.2 ml at room temperature.

These observations prove that the currently common idea that the voltaic pile isincapable of work is not entirely true. I found the open-circuit potential between aVictorian (1894) half-crown and a freshly abraded zinc plate, when separated bycloth moistened with 5% saline, to be initially 1.00 V. I also found that a minimumpotential of 3.0 V (from a modern power pack) was required to produce visibleevolution of gases from platinum wires positioned about 1 mm apart in a brightlyilluminated drop of tap water. (This figure is higher than might be expected becauseof the phenomenon of overvoltage when a gaseous product is involved.) So, althoughthe 36 or 68 V open-circuit potentials expected from the Nicholson and Carlisle pilesmust have dropped when the electrolysis apparatus was connected across theirterminal plates (terminals), it must still have exceeded 3 V.

It is possible to estimate the average current flowing from the data given by theexperimenters. First, from the gas laws, 9.2 ml of hydrogen at room temperature (say17 °C) correspond to 8.7 ml at 0 °C. Now, hydrogen has a density of 0.0899 g/litreat standard temperature and pressure (STP), so 8.7 ml must weigh 0.78 mg. Faraday’slaw states that 1.00 g of hydrogen is released by the passage of 96500 C (ampereseconds), so 0.78 mg of the gas is liberated by 75.3 C. Thirteen hours contain 46800 s,so the average current flowing was 1.6 mA. The total resistance of the circuit (pileplus water drop) would have to have exceeded 1875 V. The measurements reportedabove confirm that these figures are feasible for a freshly assembled pile of comparat-ively large silver and zinc discs connected to a high-resistance load. Because ofcorrosion and drying-up, such a pile is not expected to remain active for more thana few days. The indefatigable Wollaston reported in a letter to a friend39 that anextensive pile of silver and zinc discs was not necessary to accomplish the electrolysisof water. He claimed that three (or even two) shillings were sufficient. This would notbe expected from my measurements above, and I have been unable to confirm hisextreme claim. A five-coin pile gave an open-circuit potential of 4.8 V on the electro-meter, but when connected to two 0.3 mm platinum wires dipping into a drop ofwater the e.m.f. fell to 3.7 V. A current of 0.3 mA flowed, and after a minute or twoit was possible to see tiny bubbles adhering to the platinum wires. There can be nodoubt that the prompt discovery of the electrolysis of water was facilitated by thefact that its gaseous decomposition products are readily visible at the microgram level.

10.2. Humphry Davy: use of nitric acid in voltaic cellsThe invention of the voltaic pile led Humphry Davy to begin his brilliant

researches in electrochemistry.40 In particular, his initial studies led him to emphasize

39William H. Wollaston. Letter to Henry Hasted, 6 June 1800. Quoted by M. C. Usselman,‘Wollaston’s Microtechniques for the Electrolysis of Water and Electrochemical Incandescence’, inElectrochemistry, Past and Present, ACS Symposium 390, ed. by J. T. Stock and M. V. Orna (AmericanChemical Society, 1988), Chapter 2, pp. 20–31.40 Harold Hartley, Humphry Davy (London: Nelson, 1966); Science and the Sons of Genius: Studies

on Humphry Davy, ed. by Sophie Forgan (London: Science Reviews Ltd, 1980); David Knight, HumphryDavy: Science and Power (Cambridge: Cambridge University Press, 1992).

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the importance of oxidation of the zinc (a pile would not work if isolated from theair) and to discover the improved performance consequent on moistening the porousbarriers with dilute nitric acid.41 Although he was later to modify his theoreticalviews,42 this is doubtless why he incorporated this oxidizing reagent in electrolytesfor batteries, with the order-of-magnitude improvement quantified above.

10.3. Humphry Davy: the composition of waterDavy was puzzled by the traces of alkali left on evaporation of water that had

been extensively electrolysed to yield hydrogen and oxygen. He believed that purewater consisted solely of these gases in combination, so hypothesized that the solidsmust have come from the glass still used to purify the water, or the vessels employedto contain it. He therefore constructed a silver still and (realizing the possibility ofcontamination by spray generated by bursting bubbles) operated it in film-boilingmode. When the resulting very pure water was subjected to electrolysis it gavehydrogen and oxygen in the expected ratios, and complete evaporation left nosolid residue.43 Knight (see note 40) has pointed out the anomaly that really purewater would have such a low conductivity that the experiment should not haveworked. That it did was almost certainly due to the dissolution of atmosphericcarbon dioxide giving carbonic acid in solution. (It is ‘equilibrium water’, not‘conductivity water’.)

10.4. Trough batteriesAn early form of battery was described in 1801 by Davy and Cruickshank (see

notes 19 and 30). It consisted of a wooden trough with grooved sides, into whichwere slid 10–40 plates of copper (or silver) and zinc soldered back to back andconnected in series. The assembly was made watertight and acid resistant with pitch,and electrolyte added to each cell.

In 1807 Wilkinson (note 7) proposed the improvement of using wooden partitionsto create individual compartments, so that the appropriately supported electrodescould be raised when out of use. The dimensions of these plates varied from 1D inches× 1D inches to 12 inches × 12 inches (notes 40 and 41) or even 48 inches ×24 inches in Children’s huge battery referenced in note 27. Various electrolytes wereused in the early years (see Table 1) but an aqueous mixture of sulphuric and nitricacids had proved itself generally superior by 1809. A simplified diagram of such atrough has become the standard symbol for a battery, and is used as such in Figure 2.

When moulded gutta-percha became available, ten-cell troughs containing4-inch-square electrodes became popular on account of their durability and reducedweight. Such batteries were used, for example, by Faraday (Figure 11).

A working replica of a ten-cell trough battery was set up by making a woodentrough to hold ten rectangular polythene vessels. Each contained a pair of 4-inch-square copper and zinc electrodes (Figure 12). It was first filled with nitro-sulphuricacid (electrolyte 6) and tested against external loads diminishing logarithmicallyfrom 1000 to 0.5 V in the procedure detailed above. The electrolyte became veryhot. As usual, the second series of results, obtained after the battery had oncedelivered its maximum power, was preferred. The procedure was then repeated with

41 See the reference to Davy in note 19.42 Colin A. Russell, ‘The Electrochemical Theory of Sir Humphry Davy’, Annals of Science, 15

(1959), 1–25.43 See the reference to Davy in note 17.

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Early Voltaic Batteries: an Evaluation in Modern Units 391

Figure 11. ‘Faraday mystifies Daniell with a slick demonstration of the 3-cup trick.’ Note theplunge battery in the background. From an untitled and undated calotype positive in thepossession of the Royal Institution. It must date from before 1845, when Daniell died.Source: TRI99023 Portrait photograph of Michael Faraday (1791–1867) and JohnFrederic Daniell (1790–1845) (photo). The Royal Institution, London, UK/Bridgeman Art Library.

dilute sulphuric acid (electrolyte 1) in the cells. With this, they became merely warm.In both cases, an unpleasant fine mist was generated by the bursting bubbles ofhydrogen. Results are displayed in Figures 13–15. They confirm that nitro-sulphuricacid is much superior to sulphuric acid alone, the former enabling the battery tomaintain 8.0 V out to 50 mA, and 7.8 V with a current of 100 mA. The maximumpower delivered was 15 W into a 0.5 V load, when the battery delivered 8.3 A at1.8 V. A similar amount of energy was dissipated in the electrolyte as heat. Thesulphuric acid electrolyte gave a maximum power of only 1.6 W into a 1 V load,represented by a current of 1.2 A at 1.3 V.

10.5. Davy’s isolation of potassium and sodiumPerhaps the best known of Davy’s chemical accomplishments is his isolation of

metallic potassium and sodium by electrolysis of their fused hydroxides. In hissecond Bakerian Lecture,44 published in 1808, he explains that in one set-up a smallpiece of KOH (which had been exposed to the atmosphere to make the surfaceconductive) was placed on an insulated disc of platinum connected to the negative

44 See note 10.

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Figure 12. Reconstruction of a ten-cell Cu/Zn battery, with 4-inch-square plates.

Figure 13. Voltage versus load for a ten-cell, 4-inch-square, Cu/Zn trough battery withsulphuric (1) and nitro-sulphuric (6) electrolytes.

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Early Voltaic Batteries: an Evaluation in Modern Units 393

Figure 14. Voltage versus current for the ten-cell trough battery. Sulphuric (1) and nitro-sulphuric (6) electrolytes.

side of a battery of 100 (or more) plates in a nitric acid and alum electrolyte. Aplatinum wire connected to the positive pole was then brought into contact with theupper surface of the alkali. All was in the open atmosphere. Under these circum-stances a vivid action was soon observed. The KOH began to fuse at its points ofcontact with the platinum electrodes owing to resistive heating, and there was violenteffervescence at the upper surface (anode). At the same time small globules with ahigh metallic lustre appeared at the lower surface in contact with the cathode. Someburnt immediately with explosion and a bright flame, but others persisted down toroom temperature as tarnished liquid metal globules soon covered by a white film.Davy states that much the same result was obtained if the potash were fused withan external flame. By subsequent experiments he proved to his own satisfaction thatthese globules consisted of the desired potassium, but Knight (see note 40) haspointed out that their liquid nature at ambient temperatures points to an alloy ofpotassium with sodium derived from impure KOH. The fragments of this compoundelectrolysed by Davy were stated to be about 4–70 grains (2.5–4.5 g) in weight, andB inch thick. His battery consisted of 100 Cu/Zn plates, each 6 inches square, in anitric acid plus alum electrolyte.

Davy had more trouble with NaOH, being obliged to use smaller pieces only15–20 grains (1.0–1.3 g) in weight, with the electrodes no more than A inch apart.However, globules of sodium that survived to room temperature were solid, aswould now be expected for the pure element. Electrolysis of fused NaOH was thebasis of the Castner process for the industrial production of sodium metal,45 which

45 C. H. Lemke and V. H. Markant, ‘Sodium and Sodium Alloys’, in Kirk-Othmer Encyclopaedia ofChemical Technology, 4th edn, ed. by J. I. Kroschwitz and M. Howe-Grant (New York: Wiley, 1997),, pp. 327–54.

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Figure 15. Power versus load for the ten-cell trough battery, with sulphuric (1) and nitro-sulphuric (6) electrolytes.

was used into the 1930s. There is therefore more information concerning mechanism,with the overall reaction being represented as:

2NaOH�2Na+H2+O2agbgccathode anode

Individual stages are complex, with the yield of sodium never exceeding 50% oftheoretical on the basis of current consumed.

No information is, of course, available on the electrical parameters of Davy’sexperiments. Table 1 gives 0.62 V for a Cu/Zn couple immersed in nitric acid/alumelectrolyte (9), so his 100-cell battery should have given some 60 V standing potential.This would have been important in initiating a sufficient current flow to begin fusion,but would then have fallen drastically depending on the final resistance of the moltenalkali. To give some further semi-quantitative estimates, reconstructions of Davy’sexperiments were attempted. A battered platinum–gold crucible was available asscrap, and its rim provided a length of wire of the same alloy. The latter was nippedin the end of a short length of narrow copper tubing by hammering, and theprotruding wire coiled into a flat spiral. The crucible was supported by a silicatriangle, and the wire electrode clamped centrally within it with the horizontal spirala short distance above the bottom. Sufficient pellets of pure NaOH were then addedthat, when melted with a small bunsen flame, the spiral electrode was covered. An

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Early Voltaic Batteries: an Evaluation in Modern Units 395

ohmmeter indicated a resistance of 0.3–0.5 V, depending on the temperature of themolten alkali. This is so low that, once it was operating, the voltage across theplatinum electrodes in Davy’s apparatus is thought (from Section 7) unlikely tohave exceeded 10 V, with perhaps 80 W as the maximum power dissipated.

This voltage and power are readily available from a modern 12 V car battery,so to give further data the crucible assembly above was connected in circuit withsuch a source, a 20 A meter, a large 0.4 V variable resistance capable of handlingthe currents involved, and a voltmeter between the central positive electrode andthe walls of the crucible. No current flowed until the alkali was remelted with asmall flame, but then electrolysis commenced. Adjustment of the rheostat to give acurrent of 15 A, with 4 V across the crucible assembly (60 W ), maintained the alkalimolten in the absence of external heating, and produced vigorous evolution of gasesaccompanied by yellow flashes and cracks. This appears to match Davy’s description,although regrettably I was unable to find any globules of metallic sodium in theblackened electrolyte when it had cooled. A similar result was obtained with potas-sium hydroxide. The platinum–gold crucible was blackened and corroded, and theexperiments had to be terminated when perforation of the base allowed moltenalkali to leak out.

10.6. The Royal Institution batteryThe ‘Great Battery’ of the Royal Institution46 was funded by public subscription

for Humphry Davy. It was made up of 200 compartmented ceramic troughs, eachconstituting a plunge battery of ten pairs of 4-inch-square copper and zinc electrodes(i.e. 2000 elementary cells in total ). It was set up in a cellar of the premises inAlbemarle Street (Figure 16), and could be connected in various series/parallelcombinations.

The measurements made above on a single ten-cell battery of comparable con-struction, and filled with nitro-sulphuric acid, indicate that an all-series connectionshould achieve a standing potential of 8 V×200=1600 V. However, once connectedto a load, it could only deliver a current below 50 mA (corresponding to 80 W into32 kV) before the voltage would start to fall. Nevertheless, the all-series connectionenabled experiments on the electrical discharge in vacuum to be initiated andmaintained.

Higher power outputs could be achieved with lower resistance loads. The ten-cellbattery tested above gave a maximum power of about 15 W delivered into a 0.5 Vload. So, 200 batteries connected in an appropriate series/parallel combinationshould put up to 3 kW into a matching load, at the same time dissipating an equalamount of energy as heat within the total electrolyte. Some combinations achievingthis might be those which resulted in the following:

Volts Amps Load (V)

100 30 3.350 60 0.817 176 0.112 250 0.05

46 See H. Davy (note 10, 1810).

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Allan A. Mills396

Figure 16. The 2000-cell ‘Great Battery’ built by Humphry Davy in a cellar of the RoyalInstitution, London.

The last line of figures above shows that the maximum power output of the hugeRoyal Institution battery was comparable with that put out by a modern lead–acidsecondary battery when starting a car on a cold morning.

10.7. Nitric acid and load-matchingDavy, Faraday, and Sturgeon achieved their epoch-making results by rarely

using the textbook electrolyte of dilute sulphuric acid alone, but commonly addeda vital proportion of nitric acid. This acted as an oxidant to reduce polarization,and promoted an order of magnitude greater output of electrical power.

A second factor in their success was an unrealized application of the maximumpower theorem. Trough batteries have a low internal resistance (a few hundredthsof an ohm per cell ) and therefore give maximum output into low-resistance loads.Most of the early electrical experiments embodied these high-current, low-voltagecharacteristics, and were connected up with thick copper wire. If the experiment hada higher inherent resistance, then groups of trough batteries were empirically con-nected in mixed series/parallel arrangements to approach an optimal match.

11. Faraday and mercury poisoningIt has been explained in Section 2 that amalgamation of zinc electrodes with

metallic mercury was employed to extend the life of Cu/Zn batteries by preventingthe ‘local action’ experienced with metal of commercial quality. Davy examined the

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Early Voltaic Batteries: an Evaluation in Modern Units 397

electrochemical properties of zinc amalgam in 1826, but its general application tobatteries was first promoted by Kemp in 1837 (see note 8). The procedure involvedno more than pouring a few drops of mercury upon the faces of an abraded zincelectrode, and rubbing with a cloth moistened with dilute sulphuric acid. Setting upjust one ten-cell battery with 6-inch plates would generate 5 square feet of glisteningmercury-rich surfaces, their cleanliness facilitating evaporation of the element intothe atmosphere. Evolution of mercury vapour would presumably be minimized whenthe battery had been filled with electrolyte but, ironically, would re-commence assoon as it was emptied and placed on a shelf for storage. The poorly ventilatedbasement laboratory of the Royal Institution must have been heavily loaded withmercury vapour whenever (after 1837) batteries were being constructed, ‘re-charged’,or stored, and the contamination would have been exacerbated if (as so easilyhappens) globules of mercury ran into inaccessible cracks, particularly near thefurnace.

It was not then appreciated that mercury vapour is toxic because of absorptionvia the lungs and concentration in the brain. The fact that the liquid metal wasadministered as a recognized medical treatment was probably taken to indicaterelative harmlessness, but is misleading since little absorption occurs in the gut.Concern about mercury in the environment has led to much more critical investi-gation in recent decades, and it is now recognized that a variety of subjectivecomplaints such as headache, confusion, depression, feelings of weakness, andimpaired memory may be associated with chronic exposure to sub-acute levels ofmercury vapour.47 Once removed from the source of the latter, the body can slowlyeliminate combined mercury and substantial recovery follows.

It is well known that Michael Faraday experienced the above symptoms, particu-larly from 1839 to 1844, and that a long holiday away from his beloved laboratoryworked wonders on all but his loss of memory.48 One might well speculate that anaturally poor memory was exacerbated by long-term exposure to mercury vapour.49Hare50 disagreed with this hypothesis, saying that it should have been manifestduring Faraday’s early years at the Royal Institution, but appears unaware thatlarge-scale amalgamation of zinc in multi-cell batteries only became general from1838. He also confused the toxicology of mercury ingested as its compounds withthat of the element inhaled as vapour.

12. ConclusionsQuantitative studies of the performance of voltaic cells (mainly Cu/Zn) have

confirmed the superiority of aqueous solutions of strong inorganic acids as activatingelectrolytes. The addition of oxidizing agents (e.g. nitric acid, dichromate ions) muchreduces polarization, and enables significant power to be delivered: Davy and

47W. V. Farrar and A. R. Williams, ‘Mercury as a Poison’, and K. H. Falchuk, L. J. Goldwater andB. L. Vallee, ‘The Biochemistry and Toxicology of Mercury’, both in The Chemistry of Mercury, ed. byC. A. McAuliffe (New York: Macmillan, 1977), pp. 36–45 and p. 272 respectively.48 L. Pearce Williams, Michael Faraday (London: Chapman & Hall, 1965); E. Hare, ‘Michael

Faraday’s Loss of Memory’, Proceedings of the Royal Institution, 49 (1976), 33–52; Geoffrey Cantor,Michael Faraday, Sandemanian and Scientist (London: Macmillan, 1991), Chapter 10.49 Alfred Stock, ‘Die Gefahrlichkeit des Quecksilbers’, Zeitschrift fur angewandte Chemie, 15 (1926)

461–66.50 See Hare (note 48).

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Early Voltaic Batteries: an Evaluation in Modern Units398

Faraday were therefore able to make their pioneering discoveries in electrochemistryand electromagnetism.

Series connection of single cells results in cumulative addition of their individualpotentials to the resulting battery. Parallel connections may enable higher currentsto be drawn, as does Wollaston’s ‘double-copper’ arrangement. The criterion formaximum power is that the overall internal resistance of the resulting assemblyshould equal the resistance of the load. Many of the classic electrical experimentsand devices were high-current/low-voltage applications. Amalgamation of the zincelectrodes with mercury improved the life of batteries, and became general afterKemp’s publication in 1837, but may well have been deleterious to the health oflong-term experimenters such as Michael Faraday.

The early discovery of the electrolysis of water by Nicholson and Carlisle, madewith a classic silver/zinc voltaic pile, was aided by the ready visibility of microgramquantities of gaseous hydrogen and oxygen. Davy’s proof that very pure water wasmade up from these elements alone was made possible by solution of atmosphericcarbon dioxide raising its conductivity sufficiently to permit electrolysis.

A ten-cell Cu/Zn battery with 4-inch-square plates immersed in the much-usednitro-sulphuric acid electrolyte was found by experiment to deliver a maximumpower of 8.3 A at 1.8 V (15 W ). The 100-cell 6-inch-square Cu/Zn battery used byDavy for the isolation of sodium and potassium is estimated to have delivered some60 W (15 A at 4 V ) under his experimental conditions. The giant 2000-cell batteryof the Royal Institution probably gave no more than 3 kW into a matching load,with another 3 kW being dissipated as heat within the battery itself.

It is concluded that quantitative basic measurements combined with carefulreconstruction of significant experiments can usefully complement documentarystudies in the history of science.

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