The Fontana History of Chemistry

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From Humboldt and Gay-Lussac’s analyses of water, it was known that 87.4 parts by weight of oxygen combined with 12.6 parts of hydrogen to form water. This ratio, H : O :: 12.6 : 87.4, must also be the ratio of the individual weights of hydrogen and oxygen atoms that make up the binary atom of water. Since hydrogen is the lightest substance known, it made sense to’ adopt it as a standard and to compare all heavier chemical objects with it. If the hydrogen atom is defined as having a weight of 1, the relative atomic weight of an atom of oxygen will be roughly 7. (Dalton always rounded calculations up or down to the nearest whole number.) Similarly, Dalton assumed ammonia to be a binary compound of azote (nitrogen) and hydrogen. From Berthollet’s analysis he calculated the relative atomic weight of nitrogen to be 5 or, after further experiments in 1810, 6.

TABLE 4.1 Some of Dalton’s relative weights.

Dalton was well aware of the arbitrary nature of his rules of simplicity. In the second part of the New System in 1810 he allowed the possibility that water could be a ternary compound, in which case oxygen would be 14 times heavier than hydrogen; or, if two atoms of oxygen were combined with one of hydrogen, oxygen’s atomic weight would be 3.5. This uncertainty was to plague chemists for another fifty years.

From the beginning, Dalton symbolized his atoms

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… by a small circle, with some distinctive mark; and the combinations consist in the juxta-position of two or more of these.

The synthesis of water and ammonia were represented as:

Such symbols referred to the atom and were therefore conceptually very different from alchemical symbols or those of Hassenfratz and Adet, which only had a hazy or qualitative meaning. Earlier symbols had been a shorthand; Dalton’s circles conveyed a theoretical meaning as well as being a convenient abbreviation.

Dalton was never to become reconciled to the symbols introduced by Berzelius, even though he himself used alphabetical abbreviations within circles for elements such as iron, sulphur, copper and lead. In 1837, soon after the British Association for the Advancement of Science had persuaded British chemists to adopt Berzelius’ symbols, Dalton wrote a testimonial for Thomas Graham’s application for the Chair of Chemistry at University College London.

Berzelius’s symbols are horrifying: a young student in chemistry might as soon learn Hebrew as make himself acquainted with them. They appear like a chaos of atoms. Why not put them together in some sort of order?…[They] equally perplex the adepts of science, discourage the learner, as well as to cloud the beauty and simplicity of the Atomic Theory.

Clearly Dalton felt strongly about his innovation and was prepared to criticize a professorial candidate with one hand while supporting him with another. Indeed, Dalton suffered the first of his two strokes in April 1837 after angrily discussing symbols with a visitor.

Dalton’s symbols did not survive, mainly one suspects because they were an additional printing expense, but both they, as well as Berzelius’ simplification, encouraged people to acquire a faith in the reality of chemical atoms and enabled chemists to visualize relatively complex chemical reactions. As in mathematics, chemistry could advance only to a certain degree without an adequate symbolism for its deeper study. Between them, Lavoisier and Dalton completed a revolution in the language of chemistry.

Dalton’s hieroglyphs also reveal that he had a three-dimensional geometrical model of combination in mind. When three or more particles combined, he conceived that like particles stationed themselves as far apart as possible. This conception offers not only an important clue concerning the origins of Dalton’s atomic theory, but an explanation of his opposition to the notion derived from the volumetric combination of gases that equal volumes contained equal numbers of particles.

THE ORIGINS OF DALTON’S THEORY (#ulink_27da9a08-7a0c-57b4-b8ab-a7f06a48fd80)

How did Dalton come to think of weighing atoms? There have been many different attempts by chemists and historians to explain this. Dalton supplied three, mutually inconsistent, accounts of his voyage of discovery. Reconstruction has been made difficult by the fact that most of Dalton’s surviving papers were destroyed during the Second World War, and, but for the fact that Henry Roscoe and Arthur Harden quoted from them in a historical study published in 1896, historians would have been hard-pressed for evidence. Although the debate over influences remains unresolved, all historians agree that Dalton must have come to his ‘new views’ through the study of the physical properties of gases, which in turn depended upon his youthful interest in meteorology. For, once air had been shown to be heterogeneous, and not a homogeneous element, the question arose whether oxygen, nitrogen, carbon dioxide and water vapour were chemically combined in air (perhaps a compound actually dissolved in the water vapour?) or merely mixed together. The fact that atmospheric air appeared to be homogeneous and that its gaseous components were not stratified according to their specific gravities (itself an indication that chemists like Priestley were prepared to think in terms of the specific weights of gas particles long before Dalton) made most late-eighteenth-century chemists believe that atmospheric gases were chemically combined.

Dalton thought differently. His long study of Newton’s Principia had made him familiar with Newton’s demonstration that Boyle’s law relating pressure and volume could be derived from a model in which homogeneous air particles were self-repulsive with a power inversely proportional to the distance. As a result of his meteorological studies, Dalton had become convinced by 1793 that water vapour could not possibly be chemically combined in air; instead, it was diffused among the other aerial particles and so freely available for precipitation or condensation as rain or dew. But if water was not chemically combined, why should the other constituents of air be?

If Newton’s model of the self-repulsion of air particles was translated into a model of self-repulsive constituents of air, what would be the consequences? Provided particles of one kind did not repel particles of a different gas, as Dalton showed, each gas or vapour would behave as if in a vacuum. The net effect would be a homogeneous mixture as the different gas particles repelled their own kinds. As to the cause of their self-repulsion, Lavoisier’s model of a gas supplied a satisfactory candidate: caloric. By imagining that the particles of oxygen, nitrogen, carbon dioxide and water vapour were surrounded by atmospheres of heat, Dalton arrived at a theory of mixed gases and, incidentally, a law of partial pressures that proved essential in quantitative work in gas analysis and barometry.

Dalton’s ‘New theory of the constitution of mixed aeriform fluids and particularly of the atmosphere’ was published in Nicholson’s Journal in 1801. It proved controversial, but this was no bad thing for Dalton’s British and European reputation. Most chemists who believed in the chemical theory of air wondered how it was that caloric atmospheres in different particles did not repel one another. Why suppose that there were ‘as many distinct kinds of repulsive powers, as of gases … and that heat was not the repulsive power in any one case’?

Dalton’s ingenious reply to this difficulty was published in full in the second part of the New System and was premised on differences of size of gaseous particles, the size being a function of both the atom’s volume and the radius of its atmosphere of heat. Using diagrams that look like the later magnetic force diagrams popularized by Faraday, Dalton showed visually that ‘no equilibrium can be established by particles of different sizes pressing against each other’. It followed that different particles would ‘ignore’ one another even when surrounded by the repulsive imponderable of heat. Such a static model remained the only satisfactory explanation of gaseous diffusion, partial pressures and atmospheric homogeneity until it was replaced in the 1850s by the kinetic theory of gases.

As historians of chemistry have shown, this second model of mixed gases, which was dependent on the sizes of atoms, was first developed by Dalton in September 1804, a full year after he had developed the first list of particle weights. The question of size offers a clue to his thinking during the previous year.

One of Dalton’s few supporters for the first theory of mixed gases was his Mancunian friend, William Henry (1774–1836), the owner of a chemical works for the manufacture of the pharmaceutical, milk of magnesia, used in the treatment of digestive complaints. Henry had at first opposed Dalton, only to be converted when he found that ‘water takes up the same volume of condensed gas as of a gas under ordinary pressure’. Henry’s law that the solubility of a gas at a given temperature depended upon pressure, which he discovered in 1803, was powerful evidence that solution was a purely mechanical effect. If chemical affinity was not involved, it seemed equally unlikely to be involved in the atmosphere. Moreover, as Henry found, a mixture of gases dissolved in water was ‘retained in its place by an atmosphere of no other gas but its own kind’.

Henry’s experiments were intriguing. Why, Dalton wondered, did different gases have different solubilities in water? Why were light and elementary gases such as hydrogen and oxygen least soluble, whereas heavier compound gases such as carbon dioxide were very soluble? If his first theory of mixed gases was correct, why should gases have different solubilities? Was solubility proportional to density and complexity? At this stage Dalton clearly thought solubilities were a function of the sizes of particles

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I am nearly persuaded that the circumstance depends upon the weight and number of the ultimate particles of the several gases: those whose particles are lightest and single being least absorbable, and others more according as they increase in weight and complexity.

One can see how this line of reasoning would lead automatically to ‘an inquiry into the relative weights of the ultimate particles of bodies … a subject as far as I know, entirely new’. It is important to realize, however, that Dalton really needed to know the weights of particles only because he wanted an estimate of their sizes from the simple relationship, density = weight/volume.

As we have seen, at the end of 1803 Dalton estimated the weights of gas atoms using known chemical analyses and the rule of simplicity. From these he derived a number of atomic volumes and radii, but was unable to find any simple or regular correlation with solubility. Even so, as late as 1810 in the New System he continued to record atomic sizes alongside atomic weights.

It was evidently not until 1804 that Dalton realized that relative atomic weights were a useful explanation of the law of constant composition and that the simple rules of chemical synthesis from which he had derived them explained and predicted that, when elements combined to form more than one compound, the weights of one element that combined with a fixed weight of the other were bound to be small whole numbers. For example if

then the weights of A combined with weight B are in the simple ratio of 1:2. Dalton drew attention to the fact that this was the ratio of hydrogen to carbon in methane (CH

) and ethane (C

H

), and that in the difficult and complicated cases of the oxides of nitrogen the ratio of oxygen to nitrogen was 1 : 2 and 1:3.

There were a large number of cases of this ‘law of multiple proportions’ that had been reported in the literature as a result of the dispute between Proust and Berthollet. When Berthollet accompanied Napoleon’s expedition to Egypt in 1798, he was surprised to find huge deposits of soda by the shores of salt lakes. Mineralogical analysis showed that the soda arose from a reaction between salt and limestone in the lake bottom, in complete contradiction to the usual laboratory reaction in which soda [sodium carbonate] and calcium chloride reacted to form salt and limestone [calcium carbonate]. He concluded that the enormous concentration of salt in the lakes had forced the reversal of the usual reaction. In other words, the action of mass (concentration) could overcome the usual play of elective affinities between substances. In modern terms, Berthollet had stumbled upon an equilibrium reaction:

CaCl

+ Na

CO

CaCO

+ 2NaCl

It was this awareness of the role of mass in reactions that caused Berthollet during the next few years to challenge the usual implicit presumption of chemists that substances combined together in fixed proportions, or that constant saturation proportions always characterized chemical union. Instead, Berthollet proposed that compounds combined together in variable and indefinite proportions, and he pointed to solutions and alloys, and what would today be defined as mixtures, as empirical evidence for his claim.

This radical reconceptualization of composition was immediately challenged by a fellow Frenchman, Joseph-Louis Proust (1754–1826), who worked as an analyst in Spain. In a long series of meticulous analyses and friendly challenges to Berthollet, Proust argued that there was overwhelming evidence that regular compounds were formed from their constituents in fixed and definite proportions. There might well be more than one compound of the same two substances, but their proportions were regular.

Although neither contender gained a definite victory (for Berthollet was perfectly correct in his position over some of the difficult substances like glasses that he examined), by 1810, and in the light of Dalton’s theory, it was seen that the laws of definite and multiple proportions offered a securer foundation for quantitative chemistry. For the time being Berthollet’s views, which were eventually to illuminate physical chemistry and the theory of semiconductors, served only to confuse and handicap the development of analytical chemistry, which was so beguilingly explained by Dalton’s theory.

In 1808 William Hyde Wollaston and Thomas Thomson provided further convincing experimental examples of multiple proportions when they showed that there was a 2 : 1 ratio of CO

in the bicarbonate and carbonate of potassium (viz. KHCO

or K

O.2CO

.H

O and K

CO

or K

O.CO

), and, for the same amount of acid in the normal and acid oxalates of potash and strontia, there was double the amount of base in the acid oxalates.

It was in the context of this experimental work with oxalates that Thomson recommended Dalton’s chemical atomism, having already briefly referred to it in the third edition of his important textbook, System of Chemistry, in 1807. Thomson’s initial account was based directly on conversations Thomson had had with Dalton in Manchester in 1804. Soon afterwards Dalton had read an account of his first list of atomic weights in a paper read to the Manchester Literary and Philosophical Society in October 1803 (published with differences in 1805). Thomson was also directly responsible for inviting Dalton to Scotland in 1807 to lecture on his views on air, heat and chemical synthesis to audiences at the Universities of Edinburgh and Glasgow. It was these lectures that led to the New System.

Although Dalton’s brilliant insight was developed by others, it is worth emphasizing that he retained other imaginative insights that remained undeveloped for decades. In particular, it is clear from surviving remnants that Dalton built models of atoms and compounds in order to illustrate his theory. This model-building followed directly from his first thoughts on mixed gases in 1803

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When an element A has an affinity for another B, I see no mechanical reason why it should not take as many atoms of B as are presented to it, and can possibly come into contact with it (which may probably be 12 in general) except so far as the repulsion of the atoms of B among themselves are more than a match for the attraction of an atom of A. Now this repulsion begins with 2 atoms of B to 1 of A, in which case the 2 atoms of B are diametrically opposed; it increases with 3 atoms of B to 1 of A, in which case the atoms of B are only 120° asunder; with 4 atoms of B it is still greater as the distance is then only 90°; and so on in proportion to the number of atoms. It is evident then from these positions, that, as far as powers of attraction and repulsion are concerned (and we know of no other in chemistry) binary compounds must first be formed in the ordinary course of things, then ternary, and so on, till the repulsion of the atoms of B (or A, whichever happens to be on the surface of the other), refuse to admit any more.

This statement shows that Dalton’s apparently intuitive appeal to a principle of simplicity in chemical synthesis was backed up by a geometrical force model – a model that, in a radically different setting, was to be used by ligand-field theorists a century and a half later. But it was entirely speculative, and, although it gave ‘order’ to Dalton’s symbols, it was not a path that the empirically minded Berzelius was to follow in his own symbolic language.

ELECTRIFYING DALTON’S THEORY (#ulink_6b54e1da-1169-5e2b-98d9-a66163eaa952)

Dalton presented his theory within the context of ideas concerning heat at a time when the chemical world had become excited by the news of galvanic or current electricity. In 1800 the Italian physicist, Alessandro Volta (1774–1827), described his ‘voltaic pile’ or battery in a paper published by the Royal Society. This simple machine made from a ‘pile’ or ‘battery’ of alternating zinc and silver discs gave chemists a powerful new analytical tool. As Davy said later, its use caused great excitement and it acted as ‘an alarm-bell to experimenters in every part of Europe’. Almost immediately it was found that the battery would decompose water into its elements. While there was nothing extraordinary about this further confirmation of Lavoisier’s chemistry, the puzzling fact was that hydrogen and oxygen were ejected from the water at different poles – the hydrogen at what Volta designated as the negative pole, and the oxygen at the positive pole. Two chemists who particularly concerned themselves with this galvanic phenomenon (the term ‘electrolysis’ was not coined by Faraday until 1832) were Davy and Berzelius.

Humphry Davy (1778–1829) was born at Penzance in Cornwall and educated locally. Intending to qualify as a doctor, he was apprenticed to a surgeon in 1795 and began to read Lavoisier’s Elements of Chemistry in French in his spare time. Though ignorant and completely self-taught, like Priestley before him, Davy began to repeat, correct and devise new experiments. Apart from this growing interest in chemistry, he wrote poetry (for this was the era of Romanticism when young men poured forth their individual feelings in verse), he admired the rich Cornish scenery and he fished. Through a friendship with Gregory Watt, the tubercular son of James Watt, Davy came to the attention of Thomas Beddoes, a pupil of Joseph Black and a former lecturer in chemistry at the University of Oxford, who had resigned from ‘that place’ because of his support for the French Revolution and his suspiciously radical politics. In 1798 Beddoes, convinced that the many gases that Priestley had discovered might prove beneficial in the treatment of tuberculosis (TB) and other urban diseases, founded a subscription-based Pneumatic Institute in Bristol. He persuaded Davy, whom he recognized as a man of talent, to join him as a research assistant. Davy probably still expected to qualify as a doctor, perhaps by saving sufficient money to enter Edinburgh University as a result of this experience. In the event, he became a chemist.

Davy’s risky and foolhardy experiments at Bristol, in which he narrowly escaped suffocation on several occasions, brought him fame and notoriety in 1800 when he published his results in Researches, Chemical and Philosophical; Chiefly Concerning Nitrous Oxide … and its Respiration. None of his inhalations demonstrated chemotherapeutic benefits – though his results with nitrous oxide (laughing gas) were to be the cause of regular student ‘saturnalia’ in chemical laboratories throughout the nineteenth century. Not until 1846 was the gas used as an anaesthetic. This inhalation research, and some further essays published in 1799, which included an attack on Lavoisier’s notion of caloric and the substitution of light for caloric in gaseous oxygen (phosoxygen), brought Davy’s name to the attention of another patron, Benjamin Thompson, who had also denied that heat was an imponderable fluid.

Count Rumford, as he is better known, had founded the Royal Institution in London in 1799 as a venue for publicizing ways in which science could help to improve the quality of life of the deserving poor and for the rising middle classes. By 1801 Rumford needed a new Professor of Chemistry. Davy’s appointment coincided with the wave of contemporary interest in electrolytic phenomena and, although he lectured, dazzlingly, on many other subjects at the Royal Institution, it was his research on electrochemistry that captured the public’s imagination and ensured the middle-class success of the Institution.

By building bigger and more powerful batteries, and by using fused electrolytes rather than electrolytes in solution, Davy confirmed Lavoisier’s hunch that soda and potash were not elementary by isolating sodium and potassium in 1807. In the next few years he demonstrated that Lavoisier’s alkaline earths were also compounds and prepared calcium, strontium and barium electrolytically. Later still, Davy argued convincingly against the view that muriatic acid contained oxygen, and for the opinion that oxymuriatic acid, which he renamed chlorine, was an undecompounded elementary body – a point supported by his isolation of its conjoiner, iodine, in 1813.

This succession of corrections to Lavoisier’s chemistry has led some historians to feel that Davy set out systematically to destroy French chemistry. Indeed, by 1815 he had critically and effectively questioned most of the assumptions of the antiphlogistic chemistry – that acidity was due to oxygen, that properties were due to ‘principles’ rather than arrangement, that heat was an imponderable fluid rather than a motion of particles, and that Lavoisier’s elements were truly elementary. Although Davy was often bold in his speculations and use of analogical reasoning, in stripping Lavoisier’s system to its empirical essentials he did not replace it with any grand system of his own, except to suggest that chemical affinity was, in the final analysis, an electrical phenomenon.

In the early 1800s there were two different opinions on the cause of electrolysis. According to the ‘contact theory’ advocated by Volta, electricity arose from the mere contact of different metals; an imposed liquid merely acted as a conductor. Since this theory did not easily account for the fact that the conducting liquid was always decomposed, the alternative ‘chemical theory’ argued that it was the chemical decomposition that produced the electric current. Davy found fault with both theories and as so often in the history of science, he drew a compromise: the contact theory explained the ‘power of action’ of, say, zinc becoming positively charged when placed in contact with copper; this power then disturbed the chemical equilibrium of substances dissolved in water, leading to a ‘permanent action’ of the voltaic pile. As to the cause of the initial ‘power of action’, Davy was in no doubt that it was chemical affinity

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Is not what has been called chemical affinity merely the union or coalescence of particles in naturally opposite states. And are not chemical attractions of particles and electrical attractions of masses owing to one property and governed by one simple law?

If Davy was the first chemist to link chemical reactivity with electrolytic phenomena, it was the Swede, Berzelius, who created an electrical theory of chemistry. Davy had concluded from his long and accurate work on electrolysis that, in general, combustible bodies and bases tended to be released at the negative pole, while oxygen and oxidized bodies were evolved at the positive pole

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It will be a general expression of the facts in common philosophical language, to say, that hydrogen, the alkaline substances, the metals, and certain metallic oxides, are attracted by negatively metallic surfaces [i.e. electrodes]; and repelled by positively electrified metallic surfaces; and contrariwise, that oxygen and acid substances are attracted by positively electrified metallic surfaces, and repelled by negatively electrified metallic surfaces; and these attractive and repulsive forces are sufficiently energetic to destroy or suspend the usual operation of elective affinity.

Berzelius, with his patron-collaborator, William Hisinger, had reached the same conclusion independently in 1804, but only developed the important and influential electrochemical theory, which was to leave a permanent mark on chemistry, in 1810 after he had learned of Dalton’s atomic theory. Jöns Jacob Berzelius (1779–1848), after being brought up by his stepfather, studied medicine at the University of Uppsala. Here he read Fourcroy’s Philosophie chimique (1792) and became convinced of Lavoisier’s new system. A competent reader and writer of English, French and German, and alert to the latest developments outside Sweden, his graduation thesis in 1802 was on the medical applications of galvanism. This brought him to the attention of Hisinger, a wealthy mine owner, who invited Berzelius to use the facilities of his home laboratory in Stockholm. Together they not only drew important conclusions about electrolysis, but discovered a new element, ‘ceria’, in 1803, which later turned out to be the parent of several ‘rare-earth’ elements (see chapter 9).

By 1807 Berzelius had become independent of Hisinger’s patronage when he was elected to a Chair of Chemistry and Pharmacy at the Carolian Medico-Chirurgical Institute in Stockholm. His light lecturing duties allowed him plenty of time to research in the Institute’s modest laboratory. Elected a member of the Swedish Academy of Sciences in 1808, in 1818 he became one of its joint secretaries. The appointment included a grace-and-favour house in which he built a simple laboratory adjacent to the kitchen. Here he took occasional pupils, such as Mitscherlich and Wöhler.

Berzelius first learned of Dalton when planning his own influential textbook, Larbok i kemien, the first volume of which was published in 1808. Somehow Berzelius had acquired a copy of Richter’s writings on stoichiometry (he remarked on how unusual this was) and so learned of the law of reciprocal proportions and of the idea of equivalents. He saw immediately how useful these generalizations were for analytical chemistry. An avid follower of British chemical investigations, Berzelius learned of Dalton’s theory when he read a reference to it in Wollaston’s report on multiple proportions in Nicholson’s Journal. Because of the European wars, which made scientific communication difficult, he was unable to obtain a copy of Dalton’s New System (from Dalton himself) until 1812. Nevertheless, just from Wollaston’s brief account he saw immediately that a corpuscular interpretation of these analytical regularities was ‘the greatest step which chemistry had made towards its completion as a science’.

His own analytical results more than confirmed that, whenever substances combined together in different proportions, they were always, as Dalton had already concluded, in the proportions A + B, A + 2B, 2A + 3B, A + 4B, etc. Berzelius reconciled this regularity with Berthollet’s views on the influence of mass in chemical reactions. He agreed that Berthollet was right in supposing that substances could combine together in varying proportions; but these proportions were never continuously variable, as Berthollet had argued against Proust, but fixed according to Dalton’s corpuscular ratios.

Berzelius’ teaching duties included the training of pharmacists. He was, therefore, conscious of the fact that the Swedish Pharmacopoeia had not been revised since the days of phlogiston chemistry and that by 1810 its language had become embarrassingly out of date. In 1811, in an attempt to persuade the government to make a sensible decision on its pharmaceutical nomenclature, Berzelius devised a new Latin classification of substances, which exploited the electrochemical phenomena that he and Davy had studied, and firmly founded the organization of ponderable matter on the dualistic system that lay at the basis of Lavoisier’s antiphlogistic nomenclature.

Ponderable bodies were divided into two classes, ‘electropositive’ and ‘electronegative’ according to whether during electrolysis they were deposited or evolved around the positive or negative pole. Since these definitions reversed

FIGURE 4.1 Berzelius’ classification of substances. (Based on C. A. Russell, Annals of Science, 19 (1963): 124.)

the convention that Davy had already introduced, Berzelius was soon obliged to conform to the definition that electropositive substances were attracted to the negative pole. It was because of the theoretical implications of galvanic language that Faraday, in 1832, introduced the valueneutral nomenclature of electrodes, cathodes, anodes and so on. Berzelius’ electropositive and electronegative substances then became anions and cations respectively.

Oxygen, according to Berzelius, was unique in its extreme electronegativity. Other, less electronegative substances, like sulphur, could be positive towards oxygen and negative towards metals. On combination, a small residual contact charge was left, which allowed further combination to occur to form salts and complex salts. Thus, electropositive metals might form electropositive (basic) oxides (as electrolysis demonstrated), which would combine with electronegative acidic oxides to form neutral salts. The latter, however, might still have a residual charge that allowed them to hydrate and to form complex salts:

The scheme allowed the elements to be arranged in an electrochemical series from oxygen to potassium, based upon the electrolytic behaviour of elements and their oxides. Because salts were defined as combinations of oxides, Berzelius had to insist for a long time that chlorine and iodine were oxides of unknown elements, and that ammonia was similarly an oxide of ‘ammonia’. It was not until the 1820s that Berzelius finally capitulated and agreed that chlorine, iodine and bromine (which he placed in the special category of forming electronegative ‘haloid’ salts) were elements and that ammonia was a compound of nitrogen and hydrogen only.

It was this electrochemical system which was to have far-reaching analogical implications for the classification of organic substances. It also allowed Berzelius in 1813 to introduce a rational symbolism based upon the Latin names of the elements. Compounds were denoted by a plus sign between the constituents, as in copper oxide, Cu + O, the electropositive element being written first. Later, Berzelius dispensed with the plus sign and set the two elements side by side as in algebra. Different numbers of elements were then indicated by superscripts, e.g. S

O

a molecule of ‘hyposulphuric acid’. These joined symbols, which were criticized initially for being potentially confusing with algebraic symbolism, only began to be used in the 1830s. It was Liebig who, in 1834, introduced the subscript convention we still use today, though French chemists went on using superscripts well into the twentieth century. Because of the importance of oxygen in Berzelius’ system, he abbreviated it to a dot over its electropositive congener, i.e. Cu = Cu + O. In 1827 he extended this to sulphur, which was indicated by a comma, i.e. copper sulphide, Cú.

In a further ‘simplification’, which in practice wrought havoc in the classification of organic compounds and in communication between chemists, Berzelius in 1827 introduced ‘barred’ or underlined symbols to indicate two atoms of an element. (Since the bar was one-third up the stem of the symbol it involved printers making a special type, thereby losing one advantage over Dalton’s symbols; hence the use of underlined symbols in some texts.) The symbols for water and potash alum thus became, respectively:

Although Berzelius introduced symbols as a memory aid to chemical proportions, they were initially adopted by few chemists. Berzelius himself virtually ignored his own suggestions until 1827, when he published the organic chemistry section of his textbook, which appeared in German and French translations soon afterwards. Indeed, the development of organic chemistry was undoubtedly the key factor into pushing chemists into symbolic representations. Following the determination of a group of younger British chemists to introduce Continental organic research into Britain, Edward Turner employed Berzelius’ symbols in the fourth edition of his Elements of Chemistry in 1834. From then on, together with chemical equations, whose use in Britain was pioneered by Thomas Graham, symbols became an indispensable part of chemical communication.

TABLE 4.2 The development of the chemical equation.

As we have seen, Dalton angrily rejected Berzelius’ symbols mainly on the grounds that they did not indicate structure but were merely synoptic. Nor was he at all pleased with the way Berzelius had taken over his creation and transformed it electrochemically. On his part, Berzelius, after struggling for years to obtain a copy of Dalton’s New System, expressed deep disappointment with the book when he eventually read it in 1812

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I have been able to skim through the book in haste, but I will not conceal that I was surprised to see how the author has disappointed my hopes. Incorrect even in the mathematical part (e.g. in determining the maximum density of water), in the chemical part he allows himself lapses from the truth at which we have the right to be astonished.

Berzelius’ extensive account of his interpretation of Dalton’s theory was published in English in Thomas Thomson’s monthly Annals of Philosophy in 1813. These articles were criticized by Dalton on at least five grounds. Whereas Dalton could see no good reason geometrically why atoms had to be spherical or all the same size, these were cardinal assumptions of Berzelius, who put them to good use in 1819 when he explained the isomorphism of crystals that Mitscherlich had discovered when studying with him in Stockholm. (Isomorphism refers to the fact that a family of salts containing different metals tend to have similar or identical crystal shapes.) Again, unlike Dalton, Berzelius refused to allow combinations of the type 2A + 2B or 2A + 3B on the grounds that, logically, nothing would prevent such ‘atoms’ from being divided. Dalton disagreed, since self-repulsions could be appealed to. Only after a lifetime’s analysis, in 1831, did Berzelius accept that occasionally two atoms of an element could combine with two or more other atoms. Before then this had led Berzelius to assume that all metallic oxides had the form MO. In the cases of the alkali metals and of silver, which are actually M