Avogadro: voice in the wilderness


Avogadro's hypothesis was key to solving many problems facing the chemical sciences in the 1800s. But his idea was initially rejected. Colin Russell reports

In Short
  • Amedeo Avogadro's hypothesis stated that equal volumes of all gases, under the same temperature/pressure conditions, contain equal numbers of molecules 
  • Avogadro was born in Turin, Italy, in 1776 
  • He began his career as a lawyer and later became a professor of mathematical physics  

The demise of a little-known Italian lawyer 150 years ago hardly seems a matter of concern to 21st century chemists. Yet this man, for all his obscurity, held the key to one of the most central chemical problems of his day and ours. Without the solution that he suggested, chemistry would be unthinkably different today. His name was Amedeo Avogadro.   

Avogadro

Amedeo Avogadro was enthralled by science
 

The beginning of the 19th century was an exciting time for chemistry. In the previous two or three decades the age-old myth of phlogiston (a weightless entity given off during combustion) was exploded by Antoine Lavoisier's recognition of oxygen, and something like a modern list of elements became possible. Moreover, in Britain the ancient belief in particulate matter (atoms) had been transformed by John Dalton who, around 1808, proposed that the atoms of each element were all the same as each other but different from those of all other elements. In particular, each element had its own distinctive atomic weight. It was even possible to guess at what these atomic weights might be.   

Weighing had always been important for chemistry, never more so than when applied to the calcination of metals. By admittedly crude gravimetric techniques it became possible to determine the relative masses of elements in simple binary compounds like water, hydrogen sulphide and ammonia. For example, the relative weights of oxygen, sulphur and nitrogen, respectively, combining with one unit weight of hydrogen were shown to be roughly 8, 16 and 5. On the assumption that only one atom of each element was involved, Dalton concluded that these numbers were the actual relative masses of the atoms, their 'atomic weights'.   

Giant leap forward?   

Looking back we might well suppose that chemistry was now poised for a giant leap forward. For the first time in history it could know atomic weights. At this very time a new science of organic chemistry was beginning to emerge, and shortly afterwards quite sophisticated methods would be developed by Justus von Liebig and others for accurate gravimetric analysis of organic compounds. Given the atomic weights of carbon, nitrogen, oxygen and a few other elements it would be possible to calculate the molecular formula of each new compound, and chemists everywhere would be agreed.   

Alas, this happy prospect is an illusion, and chemistry was to become stuck in the mud for over half a century. People who write glibly about the universal progress of science would do well to consider this particular piece of chemical history. The precise problem was this.   

Although Dalton was widely venerated in his lifetime, his proposals for atomic weights were received with scepticism in many quarters. There were all kinds of reasons, not least that it was highly arbitrary to make the assumption that all binary compounds contained only two atoms. Dalton believed this on the grounds that the simplest solution was usually the right one, but that was more intuitive than anything else and was widely challenged. In its simplest form the problem could be illustrated by water: did it have one or two hydrogen atoms attached to one oxygen atom? In the former case it would be HO and oxygen would have an atomic weight of 8, and in the latter it would be (in our formulae) H2O and the value would be 16. But how could one know?    

It was clearly impossible to know, given the current state of chemical science, and within a few years confusion reigned supreme. Early in the 19th century there appeared to be no way out of this intractable problem. No matter how many gravimetric analyses were done it was hard to see any effective way forward. They could tell you how many atoms of each kind might be present if you assumed (say) that   O = 8, but not whether that assumption was right in the first place. It was like arguing in a circle the whole time. Yet, as the century wore on, the solution was staring chemists in the face but they simply could not see it. It lay in a paper by an almost unknown figure, and in a periodical that did not enjoy the widest circulation. This was an article by Amedeo Avogadro published in French in the Journal de Physique for 1811. 

Volumetric analyses   

What Avogadro really did was to switch attention from gravimetric to volumetric analyses. Two years earlier Joseph Gay-Lussac had published his law of combining volumes, which posited a simple relationship between the volumes of combining gases and also of the product if gaseous. Thus two volumes of hydrogen combined with one volume of oxygen to yield two volumes of (gaseous) water. Totally unremarkable today because of its sheer familiarity, this discovery held the clue that Avogadro needed.    

He suggested that the explanation for Gay-Lussac's law lay in the possibility that equal volumes of all gases, under the same temperature/pressure conditions, contain equal numbers of molecules (using the term in its modern sense). This is the famous Avogadro hypothesis. Thus, two molecules of water are synthesised from two of hydrogen and one of oxygen. If so, one molecule of water would be formed from half a molecule of oxygen. This cannot be the same as an atom for, as Dalton himself used to say, you cannot split an atom. The molecule must therefore be at least diatomic. Similar conclusions can be reached for chlorine and hydrogen when they have been allowed to react. So the simplest representations would be:   

2 H2+ O2= 2 H2O and H2+ Cl2= 2HCl   

If that is correct, there must be a distinction between atoms and molecules, elementary molecules could be polyatomic and the atomic weight of oxygen must be 16, not 8 as Dalton had it. Not only did this offer another self-consistent method for determining atomic weights, it also yielded results for metals that were in conformity with the later laws of Pierre Dulong and Alexis Petit (1819) and Eilhardt Mitscherlich (1822). We shall see something of its fate in a moment. Meanwhile, who was this legal interloper into the chemical scene, Amedeo Avogadro?   

Distinguished lawyer   

Avogadro was born on 9 August 1776 in Turin, being generously named Lorenzo Romano Carlo Avogadro di Quaregua di Cerreto. He came from a noble family, his father being Count Filippo Avogadro. Like many of his family his father was a distinguished lawyer, even his surname being perhaps a corruption of the Italian word for advocate.    

Early 19th century equipment to study gases

Avogadro experimented with gases

© Istituto Museo di Storia della Scienza, Firenze, Italy / Mondadoripress, Milano
 

Turin was the largest town in the Piedmont region, in the far northwest corner of modern Italy. Here young Amedeo grew up in troubled times. His homeland was part of the Kingdom of Sardinia but (when he was 22) this became occupied by the French revolutionary army. Incursions from Austria added to the general instability but the Piedmontese leant more towards France than the rest of Italy, and the region only merged into a new united Italy in 1861. 

The young Avogadro went to school in Turin and, true to family tradition, undertook legal training. In 1796 he started to practise law, later becoming a member of the local civil service.    

Like many young men of his day, however, he became enthralled by science and embarked on his own course of experiments. In a country flooded by French literature he could freely learn of the achievements of French chemistry, with the work of Gay-Lussac and Claude Berthollet especially prominent. Avogadro began experimenting with electricity, for which there was a strong tradition at the University of Turin. It is not likely to be a coincidence that this was at just the time that European science was being rocked by the discoveries of Alessandro Volta, with the first electric battery and a new focus on the relations between chemistry and electricity. Volta also lived in northern Italy.   

The outcome of this new interest in electricity was a couple of essays that Avogadro submitted to the Academy of Sciences of Turin, and, possibly as a result, he was made a corresponding member. In 1806-07 he submitted some theoretical ideas on dielectrics to the Journal de Physique. By this time he had abandoned his career in law and taken up teaching, first at a small college in Turin, and then at another in Vercelli. From here he launched his famous hypothesis of molecular volumes. Further work followed on the specific heats of gases and, in 1820, he returned to Turin as professor of mathematical physics. Two years later his post was closed for political reasons. In 1832 it was reopened, occupied briefly by the mathematician Augustin Cauchy and then, after his resignation, restored to Avogadro. He remained here for the rest of his working life, producing further papers on electricity and a four-volume textbook on theoretical physics. At the age of 74 he retired, and he died four years later, in 1856.   

We know little of Avogadro's private life. He married in 1818 and was father to seven children. His family brought him great pleasure, and they spent summer holidays together at his country house in Quaregna. He was said to have been religious without being bigoted, and seems not to have been keen on social functions. Maybe he was naturally shy, perhaps he felt he had to keep his head down in the world of constant political intrigue, and certainly he endured a good deal of scientific isolation. He did not travel Europe like Humphry Davy and was even more withdrawn than Dalton.   

Minimal impact   

This may well have been one explanation for the minimal impact of his great idea. There are others, of course. He was not a fine experimenter and the lack of hard data in his 1811 paper cannot have commended it to working chemists in other European countries. Also, the great Jöns Berzelius had proposed an electrochemical theory in which even the simplest molecules were polar, with one part positive and the other negative. That effectively excluded diatomic molecules of the same atoms, like hydrogen, chlorine or oxygen.   

Cannizzaro

Avogadro's hypothesis was the basis for Cannizzaro's work

© Royal Society of Chemistry / Library and Information Centre
 

The rejection of Avogadro's hypothesis is reminiscent of the response to the revolutionary ideas of Nicolaus Copernicus in the 16th century. For the next 50 years the number of convinced Copernicans in Europe was a few dozen only. It may be said of both men that their ideas were simply too far ahead for their time.    

This view is supported by the case of the much better known Andrè Ampère. When he supported Avogadro's hypothesis his words, too, fell largely on deaf ears. The bottom line was probably that, though coming from a lawyer, it was not a law but a hypothesis, not a statement of empirical generalisation but a hunch. It could no more be proved to be true at the time than (for over three centuries) could the sun-centred system of Copernicus. 

In the years that followed the atomic theory was adopted by some, like Berzelius and Thomas Graham, but its value was strictly limited since atomic weights were merely a matter of opinion.    

The organic chemists were particularly affected, and brave efforts to come near to a theory of structure were invariably frustrated by simple ignorance of how many atoms there might be in a new compound; everything depended on correct atomic weights. Thus, a textbook of organic chemistry by Friedrich Kekulé could offer a choice of no less than 19 different formulae for acetic acid available in the late 1850s. Each of them was wrong, because each was predicated on the assumption C = 6 and   O = 8, and no principles for structure determination then existed.    

Meanwhile the very existence of atoms was being questioned during impassioned debates at the Chemical Society in London. Clearly something had to be done. An international conference convened at Karlsruhe, Germany, in 1860 got nowhere, but as delegates were leaving they were presented with a pamphlet from another scientist from northern Italy, Stanislao Cannizzaro of Genoa. He had spoken at the conference and this was an outline for his course on chemical philosophy, where Avogadro's hypothesis was used to give a self-consistent system of atomic weights.    

Reforming chemistry   

Lothar Meyer read it on the train going home, and confessed that as he did so 'the scales fell from my eyes, doubts disappeared and a feeling of calm certainty took their place'. After such a Damascus road experience he took a lead in reforming chemistry and uniting it around agreed atomic weights. Such innovations as a clear understanding of valency, a theory of structure and a periodic table all followed from this reform. So did recognition of what came to be called Avogadro's number (NA). As Cannizzaro was quick to point out, the basis of his work was Avogadro's hypothesis.   

Not long ago this hypothesis was a familiar part of the liturgy of chemical instruction: 'Under the same conditions of temperature and pressure, equal volumes.'. Perhaps some can still recall that incantation from their days at school. Modern chemists who ignore history may be blissfully unaware of their debt to the lawyer from Piedmont.    

Colin Russell is emeritus professor, department of history of science, technology and medicine, Open University, and affiliated research scholar, department of history and philosophy of science, University of Cambridge, UK.

Further Reading

  • M Morselli, Amedeo Avogadro, 1984, Kluwer Academic 
  • M P Crosland, Dict. Sci. Biography, 1970, 1, 343 
  • J Bradley, Before and after Cannizzaro, 1992, Bradley 
  • J N Murrell, Helv. Chim. Acta., 2001, 84, 1314 
  • J H Brooke, Hist. Sci., 1981, 19, 235 
  • N Fisher, Hist. Sci. 1982, 20, 77 and 212 
  • R M Hawthorne, J. Chem. Educ., 1970, 47, 751