Posts Tagged ‘Nobel Prize’


Taking a break from studying thermodynamics: Mme. Curie measuring picoamp leakage currents

Having read biographies of the Curies, Marie Curie’s doctoral thesis, and a number of scholarly articles about the radium phenomenon, I have come to the conclusion that the Marie Curie legend in popular culture tends to sideline her scientific achievements, focusing more on her imagined saintliness and perceived role as bringer of medical marvels than on her pioneering work as a physical chemist, in which her husband Pierre (above left) played an important facilitating role.

To my mind, the popular press of her day was largely responsible for the misconstruction of the Marie Curie legend, filling the public’s mind with the discovery of a miracle cure for cancer brought about by what it portrayed as an angelic young foreign-born mother slaving away in a dark shed in Paris for no wages.

The media frenzy around radium had consequences. Once radium production was established on a commercial scale, ignorant and unscrupulous marketers quickly morphed the Curies’ discovery of an element that glowed in the dark into revitalizing radium baths, radium drinking water, radium chocolate, radium toothpaste, radium cigarettes and even radium suppositories for restoring male potency while eradicating hemorrhoids:

Also splendid for piles and rectal sores. Try them and see what good results you get!

The dreadful damage these products must have caused doesn’t bear thinking about. Sadly, the Curies themselves seemed carried along similarly radium-dazzled tracks. They failed to connect Pierre’s rapidly deteriorating health with exposure to radioactive emissions, while stoically accepting the painful damage to Marie’s hands as a price worth paying for the greater good they somehow imagined radium to represent.

Irène Curie points to her mother’s radiation-damaged hands

The sensationalist aspects of the Curie legend, while an education in themselves, are however not the subject of this post. Physics and chemistry are the subjects here. When you look at Marie Curie as a physical chemist, and examine her contributions to the science of natural radioactivity, it is clear how crucial a role was played by the miracle machine designed and developed by Pierre Curie.

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The Curie quartz electrometer apparatus

The above diagram is taken from Méthodes de Mesure employées en Radioactivité published in Paris in 1911 by Albert Laborde, a graduate engineer who became Pierre Curie’s assistant in 1902. It shows the quartz electrometer apparatus developed by Pierre Curie and his brother Jacques for the precise measurement of very weak currents (of the order of tenths of picoamperes) following their discovery of piezoelectricity in 1880. It was this discovery that prompted the brothers Curie to build a calibrated electrostatic charge generator using a thin quartz lamella (center) to compensate and thus measure the leakage current from a charged capacitor (left) using a quadrant electrometer (right).

This apparatus was later adapted by Pierre Curie to allow accurate quantification of the tiny leakage currents produced in an ionization chamber by samples of radioactive material.

This is the experimental set-up that Marie Curie can be seen using in the header photograph, which dates from 1898.

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Amazing coincidences

When you consider the train of coincidences that led to Marie Curie’s choice of subject for her doctorate (Recherches sur les substances radioactives) it is nothing short of amazing. At the time she was looking around for suitable topic, five years after having journeyed from Poland to Paris to enroll as a Sorbonne student, Henri Becquerel had just accidentally discovered mysterious rays emanating from uranium which had the property of weakly ionizing air. This was in 1896. Just a year previously, Marie had married Pierre Curie who happened to possess the one instrument capable of accurately measuring small ionization currents, following his discovery of piezoelectricity sixteen years earlier.

Because uranic rays were a new phenomenon, Marie was saved the task of first researching the topic which otherwise would have entailed reading a lot of academic papers in unfamiliar French. This saving of time and effort attracted her to choose to study Becquerel’s uranic rays, something she admitted in later life. Furthermore she had no competition since Becquerel had shown little interest in pursuing his original finding – the big news at the time was X rays, discovered by Wilhelm Röntgen in 1895. No fewer than 1044 papers on X rays were published in 1896, when Becquerel first announced his discovery. Not surprisingly, nobody took any notice. Marie Curie had the field to herself.

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The weight, the watch and the light spot

This magnified portion of the header image shows Marie Curie as she sits at the quartz electrometer apparatus. Her right hand can be seen holding an analytical balance weight in a controlled manner while in her left hand the edge of a stopwatch can be seen. Her eyes are looking fixedly at a horizontal measuring scale above a light source (square hole) mounted on a wooden pedestal.

The light source is shining a beam onto the mirror of a quadrant electrometer out of view to the left. The light spot is reflected onto the horizontal scale (cf. diagram above) and Marie is endeavoring to keep the light spot stationary. She does this by gradually releasing the weight which is attached to the quartz lamella, thereby generating charge to compensate the ionization current produced by the radioactive sample in the ionization chamber also out of view to the left. The entire process of weight release is timed by a stopwatch. Once the weight is fully released the watch is stopped. The weight generates a specific amount of charge Q* on the quartz lamella during the measured time T. Hence Q/T is equal to the ionization current, which is directly proportional to the intensity of the ionizing radiation emitted by the sample, or to use the term Marie Curie coined, its radioactivity.

*The amount of charge Q is calculated from Q = W × K × L/B where W is the applied weight, K is the quartz specific constant, L is the lamella length and B is the lamella thickness.

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On Thursday 25th June 1903, at La Faculté des Sciences de Paris, Marie Curie presented her doctoral thesis to the examination committee, two of whose members were later to become Nobel laureates. The committee was impressed; in fact it expressed the view that her findings represented the greatest scientific contribution ever made in a doctoral thesis.

At the outset, Curie coined a new term – radioactivity – to describe the ionizing radiation emitted by the uranium compounds studied by Henri Becquerel. She announced her discovery that the element thorium also displays radioactivity. And she presented a method, using the quartz electrometer apparatus developed by Jacques and Pierre Curie, by which the intensity of radioactive emissions could be precisely quantified and expressed as ionization currents. This was a game-changing advance on the essentially qualitative methods that had been used hitherto e.g. electroscopes and photographic plates.

As one would expect, Curie began her experimental work with a systematic study of uranium and its compounds, measuring and tabulating their ionization currents. There was a considerable range from the largest to the smallest currents, and within the limits of experimental error it was evident that the ionization currents were proportional to the amounts of uranium present in the sample. The same was true for thorium.

From the chemist’s perspective this was a puzzling result. The properties of chemical compounds of the same element generally depend on what it is compounded with and the arrangement of atoms in the molecule. Yet here was a very different finding – the radioactivity Curie measured was independent of compounding or molecular structure.

Curie drew the conclusion that radioactivity was a property of the atom – une propriété atomique she called it. She wasn’t referring to the uranium atom or the thorium atom, but to the atom as a generalized material unit with an implied interior from which radioactive emissions issued. That is a profound conception, with which Marie Curie made a significant contribution to the advancement of physics.

And at this point in her thesis she hadn’t even mentioned radium.

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New elements

For the next part of her thesis, Marie Curie turned her attention to the study of uranium-containing minerals, one of which was the mineral pictured above. Today we call it uraninite but in Curie’s day it was called pitchblende. The sample she obtained was from a uranium mine near the town of Joachimsthal in Austria, now Jáchymov in the Czech Republic. She measured its ionization current and found it to be considerably higher than its uranium content warranted. If her radioactivity hypothesis was correct, there was only one explanation: pitchblende contained atoms that emitted much more intense radiation than uranium atoms, which meant that another radioactive substance must be present in the ore. Curie now had the task of finding it, and was joined in this quest by her husband Pierre and the chemist Gustave Bémont.

The quartz electrometer demonstrated its value yet again, since the various fractions derived from the pitchblende sample during chemical analysis could be tested for radioactivity. In this way, the radioactivity was followed to two fractions: one containing the post-transition metal bismuth* and another containing the alkaline earth metal element barium. The Curies announced their findings in July 1898, stating their belief that these fractions contained two previously unknown metal elements, and suggesting that if the existence of these metals were confirmed, the bismuth-like element should be called polonium and the barium-like element radium.

*unknown to the Curies, the uranium decay series actually produces two radioisotopes of bismuth along with the isotopes of polonium, so the presence of radioactivity in this fraction did not solely indicate the presence of a new element.

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The shed

The fabled shed where Marie Curie labored for four years, chemically processing tons of pitchblende to produce a tiny spoonful of radium.

The heroic work which made Marie Curie a legend took place in a shed at the back of a Grande École* in Rue Vauquelin. In order to produce sufficient quantities to isolate the new elements and determine their atomic weights, tons of pitchblende were needed. This was because the maximum amounts of radium and polonium that can coexist in secular equilibrium with uranium are in the ratios of their respective half lives. This fact and the limited human resources available rendered any attempt to isolate polonium impossible, and the situation with radium was not much better. At best, a quantity of uranium ore containing 3 metric tons of elemental uranium is needed to extract 1 gram of radium at a yield of 100%. In the primitive conditions of the shed, obtaining a gram of radium meant processing 8 or 9 tons of uranium ore. One can only wonder at how Marie Curie found the physical and mental strength for such an arduous task.

*At the time, it was called École supérieure de physique et de chimie industrielles de la ville de Paris. Today it is called ESCPI Paris.

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Did she deserve two Nobel Prizes?

Marie Curie was awarded a quarter of the Nobel Prize for Physics in 1903 for her work on radioactivity. In 1911 she was the sole recipient of the Nobel Prize for Chemistry, awarded for the discovery of radium and polonium.

There can be no doubt about her credentials for the 1903 award, but some biographers have questioned whether the 1911 Prize was deserved, claiming that the discoveries of radium and polonium were part of the reason for the first prize.

As described in this post, the experimental evidence which Marie Curie set forth to reason that radioactivity is an atomic property was based solely on her experiments with uranium and thorium. Neither radium nor polonium had anything to do with it. On these grounds the claims of those biographers can be rejected.

Which leaves the question of under what circumstances the discovery of a new element qualifies for a Nobel Prize in chemistry. Clearly the discovery of a naturally radioactive element is not sufficient, otherwise Marguerite Perey – who worked as Marie Curie’s lab assistant and discovered francium in 1939 – would have qualified. Other aspects of the discovery need to be taken into account, and in 1911 there were many such aspects to Marie Curie’s discovery of radium and polonium, and the isolation of radium.

Reading the award citation, what comes across to me – albeit between the lines – is a recognition of the monumental personal effort and dedication involved in the discovery and characterization of these remarkable elements that led to the modern science of nuclear physics.

The miracle machine

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P Mander June 2017


JH van ‘t Hoff’s laboratory in Amsterdam

The 1880s were important years for the developing discipline of physical chemistry. The gas laws of Mariotte and Gay-Lussac (Boyle and Charles in the English-speaking world) had reached a high point of refinement in Europe following the work of Thomas Andrews and James Thomson in Belfast, and Johannes van der Waals in Leiden. The neophyte science was now poised to discover the laws of solutions.

The need for this advance was clear. As future Nobel Prize winner Wilhelm Ostwald put it in his Lehrbuch der allgemeinen Chemie (1891), “A knowledge of the laws of solutions is important because almost all the chemical processes which occur in nature, whether in animal or vegetable organisms, or in the nonliving surface of the earth, and also those which are carried out in the laboratory, take place between substances in solution. . . . . Solutions are more important than gases, for the latter seldom react together at ordinary temperatures, whereas solutions present the best conditions for the occurrence of all chemical processes.”

In France, important discoveries concerning the vapor pressures exerted by solutions were already being made by François-Marie Raoult. In Germany, the botanist Wilhelm Pfeffer had developed a rigid semipermeable membrane to study the effect of temperature and concentration on the osmotic pressures of solutions. And in the Netherlands, a talented theoretician by the name of Jacobus Henricus van ‘t Hoff (note the space before the apostrophe) was busy writing up his research on chemical kinetics in a work entitled “Studies in Chemical Dynamics”, which contained all that was previously known as well as a great deal that was entirely new.

Then one day in 1883, while van ‘t Hoff was writing the last chapter of the Studies on the subject of chemical affinity, in which he demonstrates an exact relation between osmotic pressure and the vapor pressures of pure solvent and solvent in solution, a chance encounter with a colleague in an Amsterdam street misdirected his thinking and diverted him onto the wrong conceptual road.


On page 233 of the Studies in Chemical Dynamics, van ‘t Hoff showed that osmotic pressure (D) has a thermodynamic explanation in the difference of vapor pressures of pure solvent and solvent in solution. Yet having discovered this truth, he promptly abandoned it in favor of an idea which seemed to possess greater aesthetic appeal. It was one of those wrong turns we all take in life, but in van ‘t Hoff’s case it seems particularly wayward.

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Jumping to conclusions

Writing in the Journal of Chemical Education (August 1986), the American Nobel Prize winner George Wald relates how van ‘t Hoff had just left his laboratory when he encountered his fellow professor the Dutch botanist Hugo de Vries, who told him about Wilhelm Pfeffer’s experiments with a semipermeable membrane, and Pfeffer’s discovery that for each degree rise in temperature, the osmotic pressure of a dilute solution goes up by about 1/270.


Hugo de Vries (1848-1935) and Wilhelm Pfeffer (1845-1920)

In an instant, van ‘t Hoff recognized this to be an approximation of the reciprocal of the absolute temperature at 0°C. As he himself put it:

“That was a ray of light, and led at once to the inescapable conclusion that the osmotic pressure of dilute solutions must vary with temperature entirely as does gas pressure, that is, in accord with Gay-Lussac’s Law [pressure directly proportional to temperature]. There followed at once however a second relationship, which Pfeffer had already drawn close to: the osmotic pressure of dilute solutions is proportional also to concentration, i.e., alongside Gay-Lussac’s Law, that of Boyle applies. Without doubt the famous mathematical expression pv = RT holds for both.”

And thus was born, in a moment of flawed inspiration on an Amsterdam street, the Gaseous Theory of Solutions. It even had a mechanism. Osmotic pressure, according to van ‘t Hoff, was caused by one-sided bombardment of a membrane by molecules of solute and was equal to the pressure that would be exerted if the solute occupied the space by itself in the form of an ideal gas. For van ‘t Hoff, this provided the answer to the age-old mystery of why sugar dissolves in water. The answer was simple – it turns into a gas.

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Compounding the error

The law of osmotic pressure, and the gaseous theory that lay behind it, was published by van ‘t Hoff in 1886. Right from the start it was viewed with skepticism in several quarters, and it is not hard to figure out why. As the above quotation shows, van ‘t Hoff had convinced himself in advance that the law of dilute solutions was formally identical with the ideal gas law, and the theoretical support he supplies in his paper seems predicated to a preordained conclusion and shows little regard for stringency.

In particular, the deduction of the proportionality between osmotic pressure and concentration is analogy rather than proof, since it makes use of hypothetical considerations as to the cause of osmotic pressure. Moreover, mechanism is advocated – an anathema to the model-free spirit of classical thermodynamics.

Before long, van ‘t Hoff would distance himself from claims of solute molecules mimicking ideal gases, thanks to a brilliant piece of reasoning from Wilhelm Ostwald – to which I shall return. But van ‘t Hoff’s equation for the osmotic pressure of dilute solutions


where Π is the osmotic pressure, kept the association with the ideal gas equation firmly in place. And it was this formal identity that led those influenced by van ‘t Hoff along the wrong track for several years.

One such was the wealthy British aristocrat Lord Berkeley, who developed a passion for experimental science at about this time, and furnished a notable example of how one conceptual error can lead to another.

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Misguided research


Lord Berkeley (1865-1942)

It was known from existing data that the more concentrated the solution, the more the osmotic pressure deviated from the value calculated with van ‘t Hoff’s equation. The idea circulating at the time was that the refinements of the ideal gas law that had been shown to apply to real gases, could equally well be applied to more concentrated solutions. As Lord Berkeley put it in the introduction to a paper, On some Physical Constants of Saturated Solutions, communicated to the Royal Society in London in May 1904:

“The following work was undertaken with a view to obtaining data for the tentative application of van der Waals’ equation to concentrated solutions. It is evidently probable that if the ordinary gas equation be applicable to dilute solutions, then that of van der Waals, or one of analogous form, should apply to concentrated solutions – that is, to solutions having large osmotic pressures.”

And so it was that Lord Berkeley embarked upon a program of research which lasted for more than two decades and failed to deliver any meaningful results because his work was founded on false premises. It is in the highest measure ironic that van ‘t Hoff, just before he was sidetracked, had found his way to the truth in the Studies, in an equation which rendered in modern notation reads


where Π is the osmotic pressure and V1 is the partial molal volume of the solvent in the solution. This thermodynamic relationship between osmotic pressure and vapor pressure is independent of any theory or mechanism of osmotic pressure. It is also exact, provided that the vapor exhibits ideal gas behavior and that the solution is incompressible.

If van ‘t Hoff had realized this, Lord Berkeley’s research could have taken another, more fruitful path. But history dictated otherwise, and it would have to wait until the publication in 1933 of Edward Guggenheim’s Modern Thermodynamics by the methods of Willard Gibbs before physical chemists in Europe would gain a broader theoretical understanding of colligative properties – of which the osmotic phenomenon is one.

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Brilliant reasoning


Wilhelm Ostwald (1853-1932)

But to return to van ‘t Hoff’s change of stance regarding mechanism in osmosis. By 1892 he was no longer advocating his membrane bombardment idea, and in stark contrast was voicing the opinion that the actual mechanism of osmotic pressure was not important. It is likely that his change of mind was brought about by a brilliant piece of thinking by his close colleague Wilhelm Ostwald, published in 1891 in the latter’s Lehrbuch der allgemeinen Chemie. Using a thought experiment worthy of Sadi Carnot, Ostwald shows that osmotic pressure must be independent of the nature of the membrane, thereby rendering mechanism unimportant.

Ostwald’s reasoning is so lucid and compelling that one wonders why it didn’t put an end to speculation on osmotic mechanisms. Here is how Ostwald presented his argument:


“… it may be stated with certainty that the amount of pressure is independent of the nature of the membrane, provided that the membrane is not permeable by the dissolved substance. To understand this, let it be supposed that two separating partitions, A and B, formed of different membranes, are placed in a cylinder (fig. 17). Let the space between the membranes contain a solution and let there be pure water in the space at the ends of the cylinder. Let the membrane A show a higher pressure, P, and the membrane B show a smaller pressure, p. At the outset, water will pass through both membranes into the inner space until the pressure p is attained, when the passage of water through B will cease, but the passage through  A will continue. As soon as the pressure in the inner space has been thus increased above p, water will be pressed out through B. The pressure can never reach the value P; water must enter continuously through A, while a finite difference of pressures is maintained. If this were realized we should have a machine capable of performing infinite work, which is impossible. A similar demonstration holds good if p>P ; it is, therefore, necessary that P=p; in other words, it follows necessarily that osmotic pressure is independent of the nature of the membrane.”

(English translation by Matthew Pattison Muir)

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For van ‘t Hoff, his work on osmosis culminated in triumph. He was awarded the very first Nobel Prize in Chemistry in 1901 for which the citation reads:

“in recognition of the extraordinary services he has rendered by the discovery of the laws of chemical dynamics and osmotic pressure in solutions”.

But van ‘t Hoff did not have long to enjoy the accolade. “Something seems to have altered my constitution,” he wrote on August 1, 1906, and on March 1, 1911, he died of tuberculosis aged 58.


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Mouse-over links to works referred to in this post

Jacobus Henricus van ‘t Hoff Studies in Chemical Dynamics

Wilhelm Ostwald Lehrbuch der allgemeinen Chemie (1891) [English Version – see page 103]

Lord Berkeley On some Physical Constants of Saturated Solutions

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P Mander June 2015

New kids on the science block: Wilhelm Ostwald and Svante Arrhenius in the late 1800s. Photo credit

New kids on the science block: Wilhelm Ostwald and Svante Arrhenius in the 1800s. Photo credit

For a binary solution in equilibrium with its own vapor, the Phase Rule tells us that the system possesses two degrees  of freedom. So if temperature is held constant, there will be a relation at equilibrium between pressure and composition, corresponding to the observed reduction of vapor pressure by solutes. And if the solution is in equilibrium with solid solvent, there will be a relation at constant pressure between temperature and composition, corresponding to the observed depression of freezing point by solutes.

Back in the 1870s, before the thermodynamics of colligative properties had been placed on a theoretical footing, these relations had been discovered in Grenoble, France, by physicist François-Marie Raoult in connexion with his work on solutions, which occupied the last two decades of his life.

François-Marie Raoult (1830-1901), whose work on the freezing-point depression of solutes had an  unexpected influence on the history of physical chemistry. Raoult's Law is named after him. Photo credit

François-Marie Raoult (1830-1901), whose work on the freezing-point depression of solutes had an unexpected influence on the history of physical chemistry. Raoult’s Law is named after him. Photo credit

Raoult’s first paper describing solute-mediated depression of freezing point was published in 1878. It was a fertile area of study which Raoult appears to have had to himself, and not surprisingly he was the first to discover empirical relationships between quantities. One such relation he found was between the depression of freezing point and the molality of the solute:

Freezing-point depression = ΔTfp = Kfp x msolute

where Kfp is the cryoscopic constant for the solvent in degrees per molal. Raoult conducted many measurements of solutes in various solvents, and over time built up a body of data in support of the above relation. But in 1884, Raoult discovered a curious exception to the rule when he used sodium chloride as a solute – the effect on the freezing point of water was nearly twice as large as it should be. There was something peculiar about the behavior of common salt in solution.

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Svante Arrhenius in 1884, the year he submitted his doctoral dissertation on electrolytic conductivity that would ultimately lead to a Nobel Prize. The adjudicating committee at Uppsala University gave him the lowest possible grade for his work. Photo credit

Svante Arrhenius in 1884, the year he submitted his doctoral dissertation on electrolytic conductivity that would ultimately lead to a Nobel Prize. The adjudicating committee at Uppsala University gave him the lowest possible grade for his work. Photo credit

In the same year that Raoult in France discovered the anomalous colligative actions of sodium chloride as a solute, a 24-year-old student in Sweden named Svante Arrhenius submitted to Uppsala University a 150-page doctoral dissertation on electrolytic conductivity in aqueous solution. In it, he advanced the thesis that the conductivity of solutions of salts in water was due to the dissociation of the salt into oppositely charged particles, to which Faraday had given the name “ions” fifty years earlier. But while Faraday believed that ions were only produced during electrolysis, Arrhenius asserted that even in the absence of an electric current, solutions of salts contained ions. And he went as far as proposing that chemical reactions in solution were reactions between ions.

The professors at Uppsala were incredulous, and duly gave Arrhenius and his far-fetched ideas the minimum mark. But they underestimated the self-belief and resolve of their young student. Arrhenius followed up on their rebuttal by sending his dissertation to cutting-edge figures in Europe such as Wilhelm Ostwald and Jacobus Henricus van ‘t Hoff, who were actively developing the new science of physical chemistry. They were far more impressed, and following the award of a travel grant from the Swedish Academy of Sciences, Arrhenius was able to study with Ostwald in Riga, Kohlrausch in Germany, Boltzmann in Austria, and van ‘t Hoff in Amsterdam.

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J H "Haircut" van't Hoff (1852-1911), a founding figure in physical chemistry, a pioneer in chemical thermodynamics, and the first winner of the Nobel Prize in Chemistry (1901). Photo credit

J H “Haircut” van ‘t Hoff (1852-1911), a founding figure in physical chemistry, a pioneer in chemical thermodynamics, and the first winner of the Nobel Prize in Chemistry (1901). Photo credit

Raoult’s discovery of the anomalous effect of common salt on the freezing point of water attracted the interest of JH van ‘t Hoff, who in 1887 subjected this peculiar result to detailed study. Investigating a number of ‘misbehaving’ salts, van ‘t Hoff found in each instance that the ratio of the measured freezing-point depression to the expected value approached a whole number as the solutions became increasingly dilute.

In the case of common salt, sodium chloride, the limiting ratio was 2. For sodium sulfate on the other hand, the ratio was 3, and for aluminum sulfate it was 5.

At the time when van ‘t Hoff found these whole number relations, Svante Arrhenius just happened to be visiting Amsterdam on his study tour. Arrhenius saw in these results the affirmation of his doctoral thesis, and could immediately supply the explanation. In aqueous solution, sodium chloride dissociates into sodium and chloride ions, so there are really two sets of solutes. Thus the total molality of the fully dissociated (ionised) solute will be double its undissociated value, and the freezing-point depression will be twice the expected amount.

By parallel reasoning, sodium sulfate dissociates into 3 aqueous ions (2 sodium ions and one sulfate ion) and in the case of aluminum sulfate there are 5 ions (2 aluminum ions and 3 sulfate ions).

It was compelling logic; the truth of Arrhenius’ thesis of ions in solution, and its implications for the understanding of chemical reactions and bonding, could no longer be denied.

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François-Marie Raoult’s work on solutions, and his discovery of the uncommon effect of common salt on the depression of freezing point, marked the start of a chain of circumstances that directly contributed to the founding of physical chemistry as a modern science. Not only did it provide affirmation of electrolytic dissociation and the existence of ions in solution, it also brought together the bright minds of Svante Arrhenius and Jacobus Henricus van ‘t Hoff, who with Wilhelm Ostwald were to propel physical chemistry into the modern age.

nobel vh

The Nobel Prize in Chemistry 1901 was awarded to JH van ‘t Hoff “in recognition of the extraordinary services he has rendered by the discovery of the laws of chemical dynamics and osmotic pressure in solutions”.

nobel sa

The Nobel Prize in Chemistry 1903 was awarded to Svante Arrhenius “in recognition of the extraordinary services he has rendered to the advancement of chemistry by his electrolytic theory of dissociation”.

nobel wo

The Nobel Prize in Chemistry 1909 was awarded to Wilhelm Ostwald “in recognition of his work on catalysis and for his investigations into the fundamental principles governing chemical equilibria and rates of reaction”.

photo credits:

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P Mander November 2013