Archive for the ‘thermodynamics’ Category

William Nicholson and Anthony Carlisle

May 1800: Carlisle (left) and Nicholson discover electrolysis

The two previous posts on this blog concerning the leaking of details about the newly-invented Voltaic pile to Anthony Carlisle and William Nicholson, and their subsequent discovery of electrolysis, are more about the path of temptation and birth of electrochemistry than about classical thermodynamics. In fact there was no thermodynamic content at all.

So by way of steering this set of posts back on track, I thought I would apply contemporary thermodynamic knowledge to Carlisle and Nicholson’s 18th century activities, in order to give another perspective to their famous experiments.

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The Voltaic pile


Z = zinc, A = silver

In thermodynamic terms, Alessandro Volta’s fabulous invention – an early form of battery – is a system capable of performing additional work other than pressure-volume work. The extra capability can be incorporated into the fundamental equation of thermodynamics by adding a further generalised force-displacement term: the intensive variable is the electrical potential E, whose conjugate extensive variable is the charge Q moved across that potential




At constant temperature and pressure, the left hand side identifies with dG. For an appreciable difference therefore


where E is the electromotive force of the cell, Q is the charge moved across the potential, and ΔGrxn is the free energy change of the reaction taking place in the battery.

For one mole of reaction, Q = nF where n is the number of moles of electrons transferred per mole of reaction, and F is the total charge on a mole of electrons, otherwise known as the Faraday. For a reaction to occur spontaneously at constant temperature and pressure, ΔGrxn must be negative and so the EMF must be positive. Under standard conditions therefore


The redox reaction which took place in the Voltaic pile constructed by Carlisle and Nicholson was


ΔG0rxn for this reaction is –146.7 kJ/mole, and n=2, giving an EMF of 0.762 volts.

We know from Nicholson’s published paper that their first Voltaic pile consisted of “17 half crowns, with a like number of pieces of zinc”. We also know that Volta’s method of constructing the pile – which Carlisle and Nicholson followed – resulted in the uppermost and lowest discs acting merely as conductors for the adjoining discs. Thus there were not 17, but 16 cells in Carlisle and Nicholson’s first Voltaic pile, giving a total EMF of 12.192 volts.

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External work

On May 1st, 1800, Carlisle and Nicholson set up their Voltaic pile, gave themselves an obligatory electric shock, and then began experiments with an electrometer which showed “that the action of the instrument was freely transmitted through the usual conductors of electricity, but stopped by glass and other non-conductors.”

Electrical contact with the pile was assisted by placing a drop of water on the uppermost disc, and it was this action which opened the path to discovery. Nicholson records in his paper that at an early stage in these experiments, “Mr. Carlisle observed a disengagement of gas round the touching wire. This gas, though very minute in quantity, evidently seemed to me to have the smell afforded by hydrogen”.

The fact that gas was formed “round the touching wire” indicates that the contact was intermittent: when the wire was in contact with the water drop but not the zinc disc, a miniature electrolytic cell was formed and hydrogen gas was evolved at the wire cathode, while at the anode the zinc conductor was immediately oxidised as soon as the oxygen gas was formed.

In thermodynamic terms, the electrochemical cells in the pile were being used to do external work on the electrolytic cell in which the decomposition of water took place


ΔG0rxn for this reaction is +237.2 kJ/mole. So it can be seen that the external work done by the pile consists of driving what is in effect the combustion of hydrogen in a backwards direction to recover the reactants.

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Carlisle and Nicholson were intuitive physical chemists. They knew that water was composed of two gases, hydrogen and oxygen, so when bubbles which smelled of hydrogen were observed in their first experiment, it immediately set them thinking. Nicholson wrote of being “led by our reasoning on the first appearance of hydrogen to expect a decomposition of water.”


William Nicholson (1753-1815)

Nicholson used the term decomposition, so it seems safe to assume they formed the notion that just as water is composed from its constituent gases, it can be decomposed to recover them. That is a powerful conception, the idea that the combustion of hydrogen is a reversible process.

Whether Carlisle and Nicholson extended this thought to other chemical reactions, or even to chemical reactions in general, we do not know. But their demonstration of reversibility, beneath which the principle of chemical equilibrium lies, was an achievement of perhaps even greater moment than the discovery of electrolysis by which they achieved it.


Anthony Carlisle (1768-1840)

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Redox reactions

Carlisle and Nicholson’s discovery of electrolysis was made possible by the fact that the decomposition of water into hydrogen and oxygen is a redox reaction. In fact every reaction that takes place in an electrolytic cell is a redox reaction, with oxidation taking place at the anode and reduction taking place at the cathode. The overall electrolytic reaction is thus divided into two half-reactions. In the case of the electrolysis of water, we have


These combined half-reactions are not spontaneous. To facilitate this redox process requires EQ work, which in Carlisle and Nicholson’s case was supplied by the Voltaic pile.

Redox reactions also take place in every voltaic cell, with oxidation at the anode and reduction at the cathode. The difference is that the combined half-reactions are spontaneous, thereby making the cell capable of performing EQ work.

The spontaneous redox reactions in voltaic cells, and the non-spontaneous redox reactions in electrolytic cells, can best be understood by looking at a table of standard oxidation potentials arranged in descending order, such as the one shown below. Using such a list, the EMF of the cell is calculated by subtracting the cathode potential from the anode potential.

[Note that if you use a table of standard reduction potentials, the signs are reversed and the EMF of the cell is calculated by subtracting the anode potential from the cathode potential.]

For voltaic cells, the half-reaction taking place left-to-right at the anode (oxidation) appears higher in the list than the half-reaction taking place right-to-left at the cathode (reduction). The EMF of the cell is positive, and so ΔG will be negative, meaning that the cell reaction is spontaneous and thus capable of performing EQ work.

The situation is reversed for electrolytic cells. The half-reaction taking place left-to-right at the anode (oxidation) appears lower in the list than the half-reaction taking place right-to-left at the cathode (reduction). The EMF of the cell is negative, and so ΔG will be positive, meaning that the cell reaction is non-spontaneous and that EQ work must be performed on the cell to facilitate electrolysis.

The half-reactions of Carlisle and Nicholson’s Voltaic pile, and their platinum-electrode electrolytic cell, are indicated in the table below.


Table of standard oxidation potentials

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The advent of the fuel cell

Anthony Carlisle and William Nicholson

If Carlisle and Nicholson had disconnected their platinum-wire electrolytic cell after bubbles of hydrogen and oxygen had formed on the respective electrodes, and then connected an electrometer across the wires, they would have added yet another momentous discovery to that of electrolysis. They would have discovered the fuel cell.

From a thermodynamic perspective, it is a fairly straightforward matter to comprehend. Under ordinary temperature and pressure conditions, the decomposition of water is a non-spontaneous process; work is required to drive the reaction shown below in the non-spontaneous direction. This work was provided by the Voltaic pile, the effect of which was to increase the Gibbs free energy of the reaction system.


Upon disconnection of the Voltaic pile, and the substitution of a circuit wire, the reaction would spontaneously proceed in the reverse direction, decreasing the Gibbs free energy of the reaction system. This system would be capable of performing EQ work.

The reversal of reaction direction transforms the electrolytic cell into a voltaic cell, whose arrangement can be written

H2(g)/Pt | electrolyte | Pt/O2(g)

As can be seen from the above table, the EMF of this voltaic cell is 1.229 volts. We know it today as the hydrogen fuel cell.

Carlisle and Nicholson most surely created the first fuel cell in May 1800. They just didn’t apprehend it, nor did they operate it as a voltaic cell – at least we have no record that they did. So we must classify Carlisle and Nicholson’s fuel cell as an overlooked actuality; an unnoticed birth.

It would take another 42 years before a barrister from the city of Swansea in Wales, William Robert Grove QC, developed the first operational fuel cell, whose essential design features can clearly be traced back to Carlisle and Nicholson’s original.

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Mouse-over link to the original papers mentioned in this post

Nicholson’s paper (begins on page 179)

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CarnotCycle would like to say thank you to everyone who has visited this blog since its inception in August 2012.

Thermodynamics may be a niche topic on WordPress, but it’s a powerful subject with global appeal. CarnotCycle’s country statistics show that thermodynamics interests many, many people. They come to this blog from all over the world, and they keep coming.

It’s wonderful to see all this activity, but perhaps not so surprising. After all, thermodynamics has played – and continues to play – a major role in shaping our world. It can be a difficult subject, but time spent learning about thermodynamics is never wasted. It enriches knowledge and empowers the mind.

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У меня имеется цифровая метеостанция с беспроводным датчиком, расположенным вне помещения. На фотографии: в верхнем правом квадранте отображается температура и относительная влажность вне помещения (6,2°C/94%) и в помещении (21,6°C/55%).

Я считаю, что эта разница (в помещении и вне) очень важна для определенных целей. Давайте посмотрим на цифры. Когда я смотрю на показания, то всегда задаюсь вопросом о том, различается ли количество водяного пара в воздухе внутри и вне помещения? Простой вопрос, а ответ потребует некоторых рассуждений. За основание возьмем уравнение идеального газа; для вычисления абсолютной влажности по температуре и относительной влажности необходим еще специальный алгоритм расчета давления насыщенного пара как функции от температуры. А это не очень простая вещь.

Формула для вычисления абсолютной влажности

В формуле ниже, температура (Т) измерена в градусах Цельсия, относительная влажность (rh) — в %, а е — это основание натурального логарифма 2,71828 [возведенное в степень, указанную в скобках]:

Абсолютная влажность (г/м3) =
6,112 x e^[(17,67 x T)/(T+243,5)] x rh x 18,02
(273,15+T) x 100 x 0,08314

что упрощается до

Абсолютная влажность (г/м3) =
6,112 x e^[(17,67 x T)/(T+243,5)] x rh x 2,1674

Точность этой формулы в пределах 0,1% на диапазоне температур от –30°C до +35°C

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формат gif


формат jpg


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Дополнительные примечания для студентов

Стратегия вычисления абсолютной влажности, определяемой как плотность водяного пара (г/м3) по температуре (Т) и относительной влажности (rh):

1. Водяной пар — это газ, поведение которого при обычной температуре атмосферы приближено к поведению идеального газа.

2. Применимо уравнение идеального газа PV = nRT. Газовая постоянная R и переменные T и V в этом случае известны (Т измерена, V = 1 m3). Для вычисления n необходимо рассчитать Р.

3. Чтобы получить значение Р можно применить следующий вариант формулы [см. eq.10] Магнуса-Тетенса, которая дает давление насыщенного пара Psat (гектопаскали) как функцию от температуры Т (в градусах Цельсия):

Psat = 6,112 x e^[(17,67 x T)/(T+243,5)]

4. Psat — это давление при относительной влажности 100%. Для вычисления давления P при любом значении относительной влажности, выраженном в %, мы умножаем выражение для Psat на коэффициент (rh/100):

P = 6,112 x e^[(17,67 x T)/(T+243,5)] x (rh/100)

5. Теперь мы знаем P, V, R, T и можем вычислить n, а это и есть количество водяного пара в молях. Значение затем умножается на 18,02 — это молекулярный вес воды. Ответ получается в граммах.

6. Обобщение:
Формула абсолютной влажности получена из уравнения идеального газа. Она выражает n всего через две переменные: температуру (Т) и относительную влажность (rh). Давление вычисляется как функция от обеих этих переменных; объем указан (1 m3), а газовая постоянная R известна.

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YouTube: Умный гараж, часть 3, Управление вытяжным вентилятором в подвале

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Игорь пользуется моей формулой, чтобы поддерживать ячейку погреба сухой.


Октябрь 2016: Я впечатлился системой управления влажностью основания здания, разработанной Игорем, и даже опубликовал отчет на форуме

Внутри короткой трубки установлен вентилятор с круговым уплотнением, распечатанным на 3D-принтере. Вентилятор замещает воздух, находящийся в основании, на воздух снаружи. Он включается, если абсолютная влажность в ячейке выше, чем на улице на 0,5 г/м3. Предполагается, что температура снаружи ниже. Это как раз и гарантирует, что вода в ячейке превращается в пар и вытягивается, а обратный процесс не может произойти.


Полное описание с набором отличных фотографий → здесь

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Формула позволяет измерять AH по данным от высокоточного датчика RH и T


Датчик SHT75 RH и T от SENSIRION

Апрель 2016: Проф. Антониетта Франи (Prof. Antonietta Frani) на основе моей формулы создала миниатюрный прибор для измерения абсолютной влажности. Миниатюрный микроконтроллер Arduino Uno оборудован датчиком SHT75 RH и T и подключается к компьютеру по кабелю USB. Системный интегратор Роберто Валголио (Roberto Valgolio) разработал интерфейс для передачи данных в листы Excel и отображения графиков.

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Формула позволила создать калькулятор RH←→AH


Март 2016: Немецкий веб-сайт использует мою формулу для своего онлайн-калькулятора RH/AH

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Формулу процитировали в академической исследовательской публикации


Январь 2016: Исследовательская публикация в Landscape Ecology (октябрь 2015) посвящена микроклиматическим образцам в городской среде США. Там для вычисления абсолютной влажности по температуре и относительной влажности использована моя формула.

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Формула нашла применение и в блоках управления влажностью

Август 2015: ПО с открытым исходным кодом (проект Arduino) также использует в микроконтроллере управления влажностью основания здания мою формулу для расчета абсолютной влажности:


«Вся идея состоит в том чтобы измерить температуру и относительную влажность в подвале и на улице, на основании температуры и относительной влажности рассчитать абсолютную влажность и принять решение о включении вытяжного вентилятора в подвале. Теория для расчета изложена здесь –»

Дополнительные фотографии по ссылке

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Процедура вычисления AH применена в калибровке спутника погоды Американского Национального Космического Агентства (NASA)

Июнь 2015: Моя общая процедура расчета AH по RH и T применена для абсолютной калибровки Глобальной Спутниковой Системы Навигации Циклона (CYGNSS), причем именно в отношении данных RH, предоставленных Системой Непрерывного Анализа и Прогноза Климата (CFSR). Единственное изменение в моей формуле Psat состоит в том, что используется выражение Августа-Роше-Мангуса, а не Болтона.

Система CYGNSS имеет сеть из восьми спутников. Она предназначена для улучшенного прогноза силы ураганов. Запущена 15 декабря 2016 г.

Ссылка (Технический отчет “An Antenna Temperature Model for CYGNSS”, июнь 2015)

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Sir Joseph Banks, President of the Royal Society in London, sat in the splendour of his office in Somerset House. It was an April morning in the year 1800. The clatter of horse-drawn carriages in the Strand rose to his window, but he did not notice; his attention was elsewhere. Staring into space, he clutched a letter that had been delivered to him that very morning. Dated March 20th and written in French, it had been sent from Como in Lombardy by an Italian professor of experimental physics named Alessandro Volta.

Professor Volta’s letter was clearly attended by some haste, since he had dispatched the first four pages in advance of the remainder, which was to follow. The subject matter was experiments on electricity, and the first pages of Volta’s letter to Banks described the invention of an apparatus “which will no doubt astonish you”.


Alessandro Volta (1745-1827) and his amazing invention

On that April morning in London, Banks read the letter and was duly astonished. Volta’s apparatus, consisting of a series of discs of two different metals in contact separated by brine-soaked pasteboard, was capable of generating a continuous current of electricity. This was a world apart from the static electricity of the celebrated Leyden jar and indeed a most astonishing discovery; no wonder Volta was so anxious to communicate it without delay to Banks and thereby to the Royal Society – of which Volta was also a fellow.

Still clutching the letter, Joseph Banks regained his composure and collected his thoughts. He must of course arrange for the letter to be read to the Society, after which it would duly appear in print in the Society’s Philosophical Transactions.

In the meantime, Banks was naturally obliged to keep Volta’s discovery confidential. But then again, with a such an astonishing discovery as this, it was sorely tempting to show Signor Volta’s letter – in the strictest confidence of course – to certain individuals in his large circle of scientific acquaintances, who would surely be fascinated by its contents. What could be the harm in that?

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Yielding to temptation

The state of mind into which Sir Joseph Banks was propelled by Professor Volta’s letter was pertinently observed (albeit almost a century later) by Oscar Wilde in his play Lady Windermere’s Fan, in which Lord Darlington famously quips “I can resist everything but temptation.”


32 Soho Square (right), the London home of Sir Joseph Banks

The London home of Joseph Banks, in Soho Square, was the centre of bustling scientific activity and attracted all the leading members of the scientific establishment. Within a month of receiving Volta’s letter, Banks had yielded to temptation and shown it to a number of acquaintances.

Among them was Anthony Carlisle, a fashionable London surgeon who was shortly to display remarkable abilities in the realm of physical chemistry. Having perused the letter, Carlisle immediately arranged for his friend the chemist William Nicholson to look over the pages with him, after which Carlisle set about constructing the apparatus according to Volta’s instructions – the fabled instrument we now call the Voltaic Pile.


Sir Anthony Carlisle (1768-1840), painted by Henry Bone in 1827

So within a month of Volta’s hastened communication to Banks, the details of the construction of the Voltaic Pile had been leaked to, among others, Carlisle and Nicholson, enabling the latter to begin experiments with Volta’s apparatus that would lead to their privileged discovery of electrolysis, before Volta’s letter had even been read to – let alone published by – the Royal Society.

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The chronology of the case


March 20th
Volta sends a letter (in French) from Como, Lombardy, to Sir Joseph Banks at the Royal Society in London, announcing his invention of the Voltaic pile.

Banks leaks the contents of Volta’s letter to several acquaintances, including Anthony Carlisle, who arranges for William Nicholson to view the letter.

Carlisle and Nicholson construct a Voltaic Pile according to Volta’s instructions. With this apparatus they discover the electrolysis of water into hydrogen and oxygen.

June 26th
Volta’s letter is read to the Royal Society.

William Nicholson publishes a paper in The Journal of Natural Philosophy, Chemistry & the Arts, announcing the discovery of electrolysis by Anthony Carlisle and himself, using the Voltaic Pile.

Volta’s letter announcing his invention of the Voltaic pile is published in French in the Philosophical Transactions of the Royal Society, and in English in The Philosophical Magazine.

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Mouse-over links to original papers mentioned in this post

Volta’s letter to Banks (begins on page 289)

Nicholson’s paper (begins on page 179)

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8 octobre 1850 – 17 septembre 1936


Le Châtelier’s principle is unusual in that it was conceived as a generalization of a principle first stated by someone else.

In 1884, the Dutch theoretician JH van ‘t Hoff published a work entitled Etudes de Dynamique Chimique [Studies in Chemical Dynamics]. In it, he stated a principle drawn from observations of different forms of equilibrium:

“Lowering the temperature displaces the equilibrium between two different conditions of matter (systems) towards the system whose formation produces heat.”

The converse statement was also implied, leading van ‘t Hoff to the realization that application of the principle made it possible “to predict the direction in which any given chemical equilibrium will be displaced at higher or lower temperatures.”

A few months after the publication of the Etudes, the following note appeared on page 786 of volume 99 of Comptes-rendus de l’Academie des Sciences:


The note covers two pages, but the crucial paragraph is the one shown immediately above, in which Le Châtelier extends van ‘t Hoff’s recently published principle to include pressure and (in modern terms) chemical potential. Rendered in English, the paragraph reads

“Any system in stable chemical equilibrium, subjected to the influence of an external cause which tends to change either its temperature or its condensation (pressure, concentration, number of molecules in unit volume), either as a whole or in some of its parts, can only undergo such internal modifications as would, if produced alone, bring about a change of temperature or of condensation of opposite sign to that resulting from the external cause.”

Just as van ‘t Hoff used inductive reasoning to relate temperature change to displacement of equilibrium, so Le Châtelier adopts the same technique to extend the principle to changes of pressure and potential.

Having arrived at a generalized principle – that systems in stable equilibrium tend to counteract changes imposed on them – Le Châtelier then sought to deduce this result mathematically from equations describing systems in equilibrium. During this quest, he discovered that the American physicist Josiah Willard Gibbs had done a good part of the groundwork in his milestone monograph On The Equilibrium of Heterogeneous Substances (1876-1878). In 1899, Le Châtelier translated this hugely difficult treatise into French, thereby helping many scientists in France and beyond to access Gibbs’ powerful ideas.

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Early misunderstandings

Le Châtelier’s principle, first stated in 1884 and extended as the Le Châtelier-Braun principle in 1887, has stood the test of time. Today we view it as a very useful law, but that was not how it was viewed by some of the academic establishment in the early 20th century. Critics including the illustrious Paul Ehrenfest and Lord Rayleigh regarded the principle as vaguely worded and impossible to apply without ambiguity. As late as 1937, Paul Epstein in his Textbook of Thermodynamics wrote that this criticism “has been generally accepted since”.

This was news to me; when I was taught Le Châtelier’s principle at school, the wording was the same as in Epstein’s day but we had no issues with vagueness or ambiguity. I wondered what this criticism was all about, so I delved into the online archive of ancient journals. And came up with this:


From J Chem Soc, 1917; vol 111. CarnotCycle hopes that the misspelling of Braun in the title was a genuine typo, and not the deliberate use of irony to mock the authors of the principle.

It is clear from the first paragraph that the charge of ambiguity by Ehrenfest and Rayleigh arose from a failure to distinguish between cause and effect. Perturbations of systems in stable chemical equilibrium are caused by changes in generalized forces which, as Le Châtelier documents, are intensive variables. The ‘response of the system’, or generalized displacements, are the extensive conjugates. This answers Rayleigh’s question as to why we are to choose the one (pressure) rather than the other (volume) as the independent variable.

What surprised me was that this misunderstanding persisted for three decades. It just goes to show that in thermodynamics, even the most perspicacious individuals can have enduring blind spots.

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The Principle behind the Principle


In the Etudes of 1884, van ‘t Hoff stated his principle on the basis of different observations of equilibrium displacement with temperature. But while reaching his conclusion inductively, he still managed to give a precise mathematical expression of the principle. In modern notation it reads:


This famous equation, sometimes called the van ‘t Hoff isochore, was stated without proof in the 1884 edition, but in the second edition of 1896 a proof was provided which is based – as with many proofs of that era – on a reversible cycle of operations involving heat and work.

Although thermodynamically exact, the equation provides little insight into why a system in stable equilibrium tends to resist actions which alter that state. Not that this would have bothered van ‘t Hoff, who was much more interested in practicality than philosophical pondering.

But in the early 1900s, physical chemists began to look for an explanation. In A Textbook of Thermodynamics with special reference to Chemistry (1913), J.R. Partington remarked that Le Châtelier’s principle is an expression of “a very general theorem … called the Principle of Least Action. We can state it in the form that, if the system is in stable equilibrium, and if anything is done so as to alter this state, then something occurs in the system itself which tends to resist the change, by partially annulling the action imposed on the system.”

Partington was hinting at a more general notion underlying Le Châtelier’s original description. That notion was more concisely expressed in another volume entitled A Textbook of Thermodynamics, written by Frank Ernest Hoare in 1931, in which he stated “every system in equilibrium is conservative”.

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Interlude : Mapping chemical reactions


It is one the conditions of stable equilibrium in thermodynamic systems that for a given temperature and pressure, the Gibbs free energy is a minimum. In the context of a chemical reaction, it means that the Gibbs free energy of the reaction mixture will decrease in the manner shown above, where the difference between P (pure products) and R (pure reactants) is the standard free energy of reaction and E is the equilibrium point at the minimum point of the curve.

If the reactants are initially present in stoichiometric proportions, the x-axis represents the mole fraction of products in the reaction mixture. In 1920, a Belgian mathematician and physicist called Théophile de Donder proposed another name for this dimensionless extensive variable. He called it “the degree of advancement of a chemical reaction”, and represented it by the Greek letter ξ (xi).

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Defining conservative behavior

In 1937, Professor Mark Zemansky – at the time an associate professor of physics at what was then called the College of the City of New York – published a textbook entitled Heat and Thermodynamics.

In the last section of the last chapter of the book, Zemansky turns his attention to Le Châtelier’s principle. He considers a heterogeneous chemical reaction which is in phase equilibrium but not chemical equilibrium; under these circumstances the Gibbs free energy G is a function of temperature T, pressure P and degree of advancement ξ.


When the chemical reaction reaches stable equilibrium at temperature T and pressure P, it follows that ∂G/∂ξ = 0. Zemansky then considers a neighboring equilibrium state at temperature T+dT and pressure P+dP. The new degree of reaction will be ξ+dξ, but the change in the slope of the curve during this process is zero. Therefore


Zemansky thus arrives at a mathematical definition of conservative behavior for a thermodynamic system consisting of a reaction mixture in stable equilibrium with respect to the reaction to which ξ refers.

The next task is to use the operations of calculus to find expressions for the derivatives ∂ξ/∂T and ∂ξ/∂P in terms of ΔS (=ΔH/T) and ΔV respectively. The first step is to write out fully the condition on dT, dP and dξ required to maintain conservative behavior:


Zemansky then employs a neat device to introduce S and V into the calculation. The order of differentiation of a state function is immaterial, so he reverses the order of differentiation in the first two terms


Since (∂G/∂T)P,ξ = –S and (∂G/∂P)T,ξ = V,


For the sake of brevity, I will introduce at this point a shortcut that Zemansky did not use, but which does not in any way alter the results of his reasoning.

For any extensive property X which varies according to the degree of advancement of a chemical reaction ξ at constant temperature and pressure, the slope of the curve has the following property


Applying this fact to the above equation, we find that in order to maintain the equilibrium condition ∂G/∂ξ=0, dT, dP and dξ must be such that


Setting dP=0 yields the result


When ΔG=0, the denominator is positive. At equilibrium therefore, (∂ξ/∂T)P and ΔH have the same sign. So for an endothermic reaction (positive ΔH) the degree of reaction advancement at equilibrium increases as the temperature increases. This accords with Le Châtelier’s principle.

Setting dT=0 yields the result


When ΔG=0, the denominator is positive. At equilibrium therefore, (∂ξ/∂P)T and ΔV have opposite signs. For a reaction resulting in a reduction of volume, the degree of reaction advancement at equilibrium increases as the pressure increases. This accords with Le Châtelier’s principle.

Zemansky thus demonstrates that deductions from a mathematical definition of conservative behavior for a thermodynamic system consisting of a reaction mixture in stable equilibrium result in equations which “express in a rigorous form that part of Le Châtelier’s principle which concerns chemical reaction in heterogeneous systems”.

Le Châtelier never got to see this deduction of his principle. He died in 1936, just a year before Zemansky’s book was published.

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