Posts Tagged ‘Mariotte’s law’

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The Honorable Robert Boyle FRS (1627-1691)

The fourteenth child of the immensely wealthy Richard Boyle, 1st Earl of Cork, Robert Boyle inherited land and property in England and Ireland which yielded a substantial income. He never had to work for a living, and following three years of travel and study as a teenager in Europe, Boyle decided at the age of 17 to devote his life to scientific research and the cultivation of what was called the “new philosophy”.

In Britain, Boyle was the leading figure in a move away from the Aristotelian view that knowledge was best obtained by the use of reason and logic. Boyle rejected this argument, and insisted that the path to knowledge was through empiricism and experiment. He won over many to his view, notably Isaac Newton, and in 1660 Boyle was a founding member of a society which believed that knowledge should be based first on experiment; we know it today as the Royal Society.

Boyle carried out a wealth of experiments in many areas of physics and chemistry, yet he seems to have been content with obtaining experimental results and generally stopped short of formulating theories to explain them.

Leibniz expressed astonishment that Boyle “who has so many fine experiments, had not come to some theory of chemistry after meditating so long on them”.

But what about Boyle’s law? you ask.

Well, it may surprise you to know that Robert Boyle did not originate the pressure-volume law commonly called Boyle’s law. A description of the reciprocal relation between the volume of air and its pressure does first appear in a book written by Boyle, but he refers to it as “Mr Towneley’s hypothesis”, for reasons we shall see.

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Torricelli leads the way

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Torricelli using a lot more mercury than necessary to demonstrate the barometer.

It was Evangelista Torricelli (1608-1647) in Italy who started it all in the summer of 1644 with the invention of the mercury barometer. It was an impressive device, made all the more impressive by the insight that came with it. For one thing, Torricelli had no problem accepting the space above the mercury as a vacuum, in contrast to the vacuum-denying views of Aristotle and Descartes. For another, he was the first to appreciate the fact that air had weight, and to understand that the column of mercury was supported by the pressure of the atmosphere. As he put it: “We live submerged at the bottom of an ocean of air, which by unquestioned experiments is known to have weight.”

This led Torricelli to surmise that atmospheric pressure should be less in elevated places like mountains, an idea which was put to the test in France by Blaise Pascal, or more precisely by Pascal’s brother-in-law Florin Périer, who happened to live in Clermont-Ferrand, which has Puy de Dôme nearby.

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Puy de Dôme in south-central France, close to Clermont-Ferrand.

On Saturday, September 19, 1648, Florin Périer and some friends performed the Torricelli experiment on the top of Puy de Dôme in south-central France. The height of the mercury column was substantially less – 85 mm less – than the control instrument stationed at the base of the mountain 1,400 metres below.

The Puy de Dôme experiment provided Pascal with convincing evidence that it was the weight of air, and thus atmospheric pressure, that balanced the weight of the mercury column. Torricelli’s instrument provided a convenient means of measuring this pressure. It was a barometer. The news quickly spread to England.

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Reaching new heights

The Torricellian experiment was demonstrated and discussed at the scientific centres of learning in London, Oxford and Cambridge from 1648 onwards. At Cambridge, there was an enthusiastic experimental scientist by the name of Henry Power, who was studying for a degree in medicine at Christ’s College. Power’s home was at Halifax in Yorkshire, and this gave him the opportunity to verify the Puy de Dôme experiment because unlike London, Oxford and Cambridge, the land around Halifax rises to significant heights.

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The hills around Halifax in northern England

On Tuesday, May 6, 1653, Henry Power carried the Torricellian experiment to the summit of Halifax Hill, to the east of the town, where he was able to verify Pascal’s observation. In further experiments, he began to investigate the elasticity of air – i.e. its expansion and compression characteristics. And it was this change of focus that was to characterise England’s contribution to the scientific study of air.

The pioneering work in Italy and France had been concerned with the physical properties of the atmosphere. In England, attention was turning to the physical properties of air itself.

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The Spring of the Air

In 1654, the year after Henry Power’s excursion on Halifax Hill, Robert Boyle arrived in Oxford where he rented rooms to pursue his scientific studies. With the assistance of Robert Hooke, he constructed an air pump in 1659 and began a series of experiments on the properties of air. An account of this work was published in 1660 under the title “New Experiments Physico-Mechanicall, Touching The Spring of the Air, and its Effects”.

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Boyle’s book was a landmark work, in which were reported the first controlled experiments on the effects of reduced air pressure. The experiments are divided into seven groups, the first of which concern “the spring of the air” i.e. the pressure exerted by the air when its volume is changed. It is clear that Boyle had an interest not only in demonstrating the elastic nature of air, but also in finding a quantitative expression of its elasticity. Due to inefficiencies in the air pump and the inherent difficulties of the experiment, Boyle failed in his first attempt. But it was not for want of trying.

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Enter Mr Towneley

Richard Towneley, whose home was near Burnley in Lancashire, northern England, was curiously similar to Robert Boyle in that the Towneley family estate also generated an income which meant that Richard did not have to work for a living. And just like Boyle, Towneley devoted his time to scientific studies.

Richard Towneley’s home, Towneley Hall in northern England, painted by JMW Turner in 1799

Now it just so happened that the previously-mentioned Henry Power of Halifax was the Towneley family’s physician. This gave Richard and Henry regular opportunities to share their enthusiasm for scientific experiments and discuss the latest scientific news.

In 1660 they both read Robert Boyle’s New Experiments Physico-Mechanicall, and this prompted further interest in the study of air pressure which Power had conducted on Halifax Hill seven years earlier. Not far to the north of Towneley Hall is Pendle Hill, whose summit is 1,827 feet (557 meters) above sea level, and it was here that Power and Towneley conducted an experiment that made history.

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Pendle Hill, photographed by Lee Pilkington

On Wednesday 27th April 1661, they introduced a quantity of air above the mercury in a Torricellian tube. They measured its volume, and then using the tube as a barometer, they measured the air pressure. They then ascended Pendle Hill and at the summit repeated the measurements of volume and air pressure. As expected, there was an increase in volume and a decrease in pressure.

Although their measurements were only roughly accurate and doubtless affected by temperature differences between the base and summit of the hill, the numerical data was sufficient to give them the intuitive realization of a reciprocal relationship between the pressure and volume of air

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On a high hill in northern England, nature revealed one of its secrets to Henry Power and Richard Towneley. Now they needed to communicate their finding.

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Boyle reports the news

It was not until two years after the Pendle Hill experiment that Henry Power eventually got around to publishing the results in Experimental Philosophy (1663). This delay explains why history has never associated the names of Power and Towneley with the gas law they discovered, because the news from Pendle Hill was first given to, and first reported by, Robert Boyle.

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The second edition of 1662

In the appendix to the second edition of his New Experiments Physico-Mechanical published in 1662, Boyle reports Power and Towneley’s experimental activities, which in error he ascribes solely to Richard Towneley. It was on the basis of this second edition of Boyle’s book that the reciprocal relationship between the volume of air and its pressure became known as “Mr. Towneley’s hypothesis” by contemporary authors such as Robert Hooke and Isaac Newton.

A complication was added by the fact that Power and Towneley’s experiment led Boyle to the idea of compressing air in a J-shaped tube by pouring mercury into the long arm and measuring the volume and applied pressure.

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From this experiment, conducted in September 1661, Boyle discovered that the volume of air was halved when the pressure was doubled. Curiously, he translated this finding into the hypothesis that there was a direct proportionality between the density of air and the applied pressure.

It didn’t occur to Boyle that the direct relation between density and pressure was the same thing as the reciprocal relation between volume and pressure discovered earlier by Power and Towneley. Boyle seems to have thought that compression was somehow different to expansion, and that his experiment broke new ground.

For this reason Boyle believed he had discovered something new, and since his name was far better known in scientific circles than that of Henry Power and Richard Towneley, it was inevitable that Boyle’s name would be associated with the newly-published hypothesis, which in time became Boyle’s law.

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Postscript : La loi de Mariotte

The pressure-volume law discovered by Power and Towneley, and confirmed by Boyle, is known as Boyle’s law only in Britain, America, Australia, the West Indies and other parts of what was once called the British Empire. Elsewhere it is called Mariotte’s law, after the French physicist Edme Mariotte.

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Mariotte stated the pressure-volume relationship in his book De la nature de l’air, published in 1679, some 17 years after the second edition of Boyle’s New Experiments Physico-Mechanical. Mariotte made no claim of originality, nor did he make any reference to Boyle. But he was an effective publicist for the law, with the result that his name became widely associated with it.

Some might argue that the attribution is not quite deserved. However, Edme Mariotte stated something immensely important that Robert Boyle neglected to mention. He pointed out that temperature must be held constant for the pressure-volume relationship to hold:

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In the view of CarnotCycle, the pressure-volume law so stated can with justification be called Mariotte’s law.

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P Mander January 2016

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Historical background

If you received formal tuition in physical chemistry at school, then it’s likely that among the first things you learned were the 17th/18th century gas laws of Mariotte and Gay-Lussac (Boyle and Charles in the English-speaking world) and the equation that expresses them: PV = kT.

It may be that the historical aspects of what is now known as the ideal (perfect) gas equation were not covered as part of your science education, in which case you may be surprised to learn that it took 174 years to advance from the pressure-volume law PV = k to the combined gas law PV = kT.

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The lengthy timescale indicates that putting together closely associated observations wasn’t regarded as a must-do in this particular era of scientific enquiry. The French physicist and mining engineer Émile Clapeyron eventually created the combined gas equation, not for its own sake, but because he needed an analytical expression for the pressure-volume work done in the cycle of reversible heat engine operations we know today as the Carnot cycle.

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The first appearance in print of the combined gas law, in Mémoire sur la Puissance Motrice de la Chaleur (Memoir on the Motive Power of Heat, 1834) by Émile Clapeyron

Students sometimes get in a muddle about combining the gas laws, so for the sake of completeness I will set out the procedure. Beginning with a quantity of gas at an arbitrary initial pressure P1 and volume V1, we suppose the pressure is changed to P2 while the temperature is maintained at T1. Applying the Mariotte relation (PV)T = k, we write

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The pressure being kept constant at P2 we now suppose the temperature changed to T2; the volume will then change from Vx to the final volume V2. Applying the Gay-Lussac relation (V/T)P = k, we write

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Substituting Vx in the original equation:

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whence

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Differences of opinion

In the mid-19th century, the ideal gas equation – or rather the ideal gas itself – was the cause of no end of trouble among those involved in developing the new science of thermodynamics. The argument went along the lines that since no real gas was ever perfect, was it legitimate to base thermodynamic theory on the use of a perfect gas as the working substance in the Carnot cycle? Joule, Clausius, Rankine, Maxwell and van der Waals said yes it was, while Mach and Thomson said no it wasn’t.

With thermometry on his mind, Thomson actually got quite upset. Here’s a sample outpouring from the Encyclopaedia Britannica:

“… a mere quicksand has been given as a foundation of thermometry, by building from the beginning on an ideal substance called a perfect gas, with none of its properties realized rigorously by any real substance, and with some of them unknown, and utterly unassignable, even by guess.”

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Joule (inset) and Thomson may have had their differences, but it didn’t stop them from becoming the most productive partnership in the history of thermodynamics

It seems strange that the notion of an ideal gas, as a theoretical convenience at least, caused this violent division into believers and disbelievers, when everyone agreed that the behavior of all real gases approaches a limit as the pressure approaches zero. This is indeed how the universal gas constant R was computed – by extrapolation from pressure-volume measurements made on real gases. There is no discontinuity between the measured and limiting state, as the following diagram demonstrates:

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Experiments on real gases show that

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where v is the molar volume and i signifies ice-point. The universal gas constant is defined by the equation

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so for real gases

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The behavior of n moles of any gas as the pressure approaches zero may thus be represented by

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The notion of an ideal gas is founded on this limiting state, and is defined as a gas that obeys this equation at all pressures. The equation of state of an ideal gas is therefore

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William Thomson, later Lord Kelvin, in the 1850s

Testing Mayer’s assumption

The notion of an ideal gas was not the only thing troubling William Thomson at the start of the 1850s. He also had a problem with real gases. This was because he was simultaneously engaged in a quest for a scale of thermodynamic temperature that was independent of the properties of any particular substance.

What he needed was to find a property of a real gas that would enable him to
a) prove by thermodynamic argument that real gases do not obey the ideal gas law
b) calculate the absolute temperature from a temperature measured on a (real) gas scale

And he found such a property, or at least he thought he had found it, in the thermodynamic function (∂U/∂V)T.

In the final part of his landmark paper, On the Dynamical Theory of Heat, which was read before the Royal Society of Edinburgh on Monday 15 December 1851, Thomson presented an equation which served his purpose. In modern notation it reads:

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This is a powerful equation indeed, since it enables any equation of state of a PVT system to be tested by relating the mechanical properties of a gas to a thermodynamic function of state which can be experimentally determined.

If the equation of state is that of an ideal gas (PV = nRT), then

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This defining property of an ideal gas, that its internal energy is independent of volume in an isothermal process, was an assumption made in the early 1840s by Julius Robert Mayer of Heilbronn, Germany in developing what we now call Mayer’s relation (Cp – CV = PΔV). Thomson was keen to disprove this assumption, and with it the notion of the ideal gas, by demonstrating non-zero values for (∂U/∂V)T.

In 1845 James Joule had tried to verify Mayer’s assumption in the famous experiment involving the expansion of air into an evacuated cylinder, but the results Joule obtained – although appearing to support Mayer’s claim – were deemed unreliable due to experimental design weaknesses.

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The equipment with which Joule tried to verify Mayer’s assumption, (∂U/∂V)T = 0. The calorimeter at the rear looks like a solid plate construction but is in fact hollow. This can be ascertained by tapping it – which the author of this blogpost has had the rare opportunity to do.

Thomson had meanwhile been working on an alternative approach to testing Mayer’s assumption. By 1852 he had a design for an apparatus and had arranged with Joule to start work in Manchester in May of that year. This was to be the Joule-Thomson experiment, which for the first time demonstrated decisive differences from ideal behavior in the behavior of real gases.

Mayer’s assumption was eventually shown to be incorrect – to the extent of about 3 parts in a thousand. But this was an insignificant finding in the context of Joule and Thomson’s wider endeavors, which would propel experimental research into the modern era and herald the birth of big science.

Curiously, it was not the fact that (∂U/∂V)T = 0 for an ideal gas that enabled the differences in real gas behavior to be shown in the Joule-Thomson experiment. It was the other defining property of an ideal gas, that its enthalpy H is independent of pressure P in an isothermal process. By parallel reasoning

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If the equation of state is that of an ideal gas (PV = nRT), then

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Since the Joule-Thomson coefficient (μJT) is defined

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and the second term on the right is zero for an ideal gas, μJT must also be zero. Unlike a real gas therefore, an ideal gas cannot exhibit Joule-Thomson cooling or heating.

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Finding a way to define absolute temperature

But to return to Thomson and his quest for a scale of absolute temperature. The equation he arrived at in his 1851 paper,

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besides enabling any equation of state of a PVT system to be tested, also makes it possible to give an exact definition of absolute temperature independently of the behavior of any particular substance.

The argument runs as follows. Given the temperature readings, t, of any arbitrary thermometer (mercury thermometer, bolometer, whatever..) the task is to express the absolute temperature T as a function of t. By direct measurement, it may be found how the behavior of some appropriate substance, e.g. a gas, depends on t and either V or P. Introducing t and V as the independent variables in the above equation instead of T and V, we have

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where (∂U/∂V)t, (∂P/∂t)V and P represent functions of t and V, which can be experimentally determined. Separating the variables so that both terms in T are on the left, the equation can then be integrated:

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Integrating between the ice point and the steam point

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This completely determines T as a function of t.

But as we have already seen, there was a catch to this argumentation – namely that (∂U/∂V) could not be experimentally determined under isothermal conditions with sufficient accuracy.

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The Joule-Thomson coefficient provides the key

Thomson’s means of circumventing this problem was the steady state Joule-Thomson experiment, which measured upstream and downstream temperature and pressure, and enabled the Joule-Thomson coefficient, μJT = (∂T/∂P)H, to be computed.

It should be borne in mind however that when Joule and Thomson began their work in 1852, they were not aware that their cleverly-designed experiment was subject to isenthalpic conditions. It was the Scottish engineer and mathematician William Rankine who first proved in 1854 that the equation of the curve of free expansion in the Joule-Thomson experiment was d(U+PV) = 0.

William John Macquorn Rankine (1820-1872)

William John Macquorn Rankine (1820-1872)

As for the Joule-Thomson coefficient itself, it was the crowning achievement of a decade of collaboration, appearing in an appendix to Joule and Thomson’s final joint paper published in the Philosophical Transactions of the Royal Society in 1862. They wrote it in the form

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where the upper symbol in the derivative denotes “thermal effect”, and K denotes thermal capacity at constant pressure of a unit mass of fluid.

The equation is now usually written

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By the method applied previously, this equation can be expressed in terms of an empirical t-scale and the absolute T-scale:

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where C’P is the heat capacity of the gas as measured on the empirical t-scale, i.e. C’P = CP(dT/dt). Cancelling (dT/dt) and separating the variables so that both terms in T are on the left, the equation becomes:

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Integrating between the ice point and the steam point

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This completely determines T as a function of t, with all the terms under the integral capable of experimental determination to a sufficient level of accuracy.

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P Mander May 2014