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Walther Nernst and the Last Law William H. Cropper St. Lawrence University, Canton, NY 13617 According to a story current in Berlin in the early 1900's, God decided one day to create a superman. He worked first and subtle mind". on the brain, fashioning a "most But he had other business, and the job had to he put aside. One of God's assistants. the Archaneel Gabriel. saw this marvelous brain and could not resist tLe temptation to try to create the complete man. He overestimated his abilities, however, and succeeded only in creating "a rather unimpressive lookina little man". Discouraged by his failure, he left his creation inanimate. T h e ~ e v i i c a m ealong, looked with satisfaction upon this unique h u t lifeless being, a n d breathed life into it. "That was Walther Nernst."
The Authenllc Nernst This story is told in a fine biography of Nernst by Kurt Mendelssohn (I),who supplies us with a further (possibly more authoritative) picture of Nernst: "There is no record of hereditary genius [in Nernst's family] or even of outstanding enterprise. I t seemed that Walther owed his brilliance to a lucky throw of the genetic dice" ( I ) . At one time Nernst considered becoming an actor and he realized this ambition to some extent by wearing throughout his life the mask of a trusting and credulous little man. His favoriteexpression of innocent astonishment could he underlined by a twitch of the nose which removed [his]pince-nez. There was alsousually anote ofastonishmentin hisvoiceand theoutrageous and sarcastic comment of which he was the master was never accompanied by a change in his voice or asmile. He just remained genuinely serious and mildly surprised ( I ) . As a student, Nernst traveled, according to the custom, amone the universities where the great men of science lived and tRughr. His peripatetic educ&ion took him to Zurich, Iltrlin ,where Hermann t l ~ h h o l t zlectured un thrrmodynamirsi. hack t t ~Zurich, then to Grnz (to study under l a d wi6 Uoltzn~ann,,and finally t o M'uwburg (where work u,ith Frledrich Kohlra~~sch inspired a lifelong intrrest in electroch~mistry,.1Ir paused Ion:: enough in Grtu to write a docror.a1 thriii imd learn lessons in "irritation . ohssics" from Alhert . von Ettinghausen, a former student of Boltzmann's and Nernst's collaborator in his thesis research. Nernst, who could never conquer his impatience, had endless admiration for Ettinehausen's easv acceptance of experimental frustrationz. Alter a dismnl inilurt. of an experiment, Ettinghausen ~nizhrsns calml\.. "Well, the experiment was not successful. a t Least not entGely." Lelpzlg, GMtIngen, and Berlin Nernst's professional career was a story of almost unmitigated success. In the late 1880's, while a t Wiirzhurg, Nernst met Wilhelm Ostwald and became his assistant when Ostwnld wepl pled a prnfessorship at Leipzi~.Nernsr lost no time in Findine ocrupntion for his talents in Ostwald's endeavors, all concerned with building the foundations for the new subject of physical chemistry. Nernst's first publication from Leipzig became one of the classics of the literature of electrochemistry; it presented to the world anequation which came to be known to generations of physical chemistry students as the "Nernst equation" (2). In 1891, Nernst was appointed Assistant Professor of Physics a t the University of Gottingen. Two years later he published one of the first physical chemistry texts, in fact,
the second text in the field after Ostwald's Lehrbuch der allgemeinen Chemie. Nernst's text had the title Theoretische Chemie, and i t was dedicated to Ettinghausen. Nernst built his view of physical chemistry from thermodynamic foundations laid by Helmholtz, and the molecular hypothesis ("Avogradro's hypothesis") advocated by Boltzmann (and strenuously opposed a t the time by Ostwald). Still in use 30 vears later (in its 15th edition), the Nernst text was the most influential in the field. ~ m o its n successors ~ were texts hv Eucken and Eaeert, -- . both students of Nernst's. In three more years, Nernst had so impressed the Ministries of Education, not only in Prussia (where Gottingen was located), hut also in Bavaria, he was offered the Professorship of Theoretical Physics a t the University of Munich as Boltzmann's successor. The Prussian Minister, Friedrich Althoff, was not to he outdone, however. Mendelssohn tells of the further bureaucratic bargaining, masterfully manipulated by Nernst: If Althoff wanted to keep Nernst in Prussia, he now had to make aneffortthat would goa bit beyond his owndepartmentalresponsibility. Nernst's price was the creation of a new chair of physical chemistry st Gbttingen, and to go with it an electrochemical laboratory. Althoff could produce the new chair from the funds at his disposal, but for the laboratory he had to get the money from the Ministry of Finance-and that would take time. Nemst, who was certain that he held the whip hand and who always knew how to drive a hard bargain, farced Althoff into an unheard of act. It was the promise, to be given in writing, that should the laboratory in GBttingen not materialize, Nernst would get a chair of physics st Berlin. Althoff yielded, possibly because he had every reason to believe the Minister of Finance would play, asindeed he did. That was in 1894 and Berlin would have to wait another eleven years (3).
Nernst's scientific talent extended to applied problems, especially those that had economic possibilities. While a t Gottingen he invented an electric lamp that he hoped would compete with the Edison lamp, then not fully developed. Nernst's design was an application of his studies of ionic conduction. He first tried to persuade Siemens, an estahlished German electric firm, to buy the patent on the invention. Siemens was not interested either in the technical possibilities of the lamp ur in Nernst's Financial demands. The m t w t was next oiiert>dto Allaemeine Elektrilitsts Gesellschaft (A.E.G.), a newer and more adventurous company. After extended haraainina in which Nernst demanded a lump sum and refused royalties, he got what he wanted, a million marks, enough to make him a wealthy man. Although it was ingeniously developed by A.E.G., with much of the technical work done by two of Nernst's students, the Nernst lamp finally lost out in the competition with the Edison lamp. This financial disaster for A.E.G. seems not to have discouraged their confidence in Nernst's technical abilities. Emil Rathenau, the A.E.G. chairman, remained friendly with Nernst for the rest of his life. Although he had acquired wealth, become influential with those highly placed in the political and business worlds, and reached a position of eminence second only to Ostwald in the new science of physical chemistry, Nernst had not quite reached the pinnacle of success. There was one more academic world to conquer, and his scientific efforts, although extensive and impressive, did not yet rank with the work of his most distinguished predecessors. His next move took Volume 64 Number 1 January 1987
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him, in Mendelssohn's words, from his Gottingen "place in the sun", to an academic and scientific "summit" a t the University of Berlin. In 1905, strongly supported by Max Planck, Nernst was appointed as Professor of Physical Chemistrv a t Berlin. In the snrine . of 1905, Nernst drove his family frdm Gdttingen to Berlin in an openmotorcar, accompanied by his favorite mechanic in case of breakdowns. That same year Nernst found the clue he needed to formulate his statement of the Third Law of Thermodynamics. Chemlcal Thermodynamics Accordlng to Nernst
We might pause here, with Nernst about to make his great discovery, and look briefly a t the principles of thermodynamics as he viewed them. Nernst was not convinced that thermodynamics needed the entropy concept. Anything he wanted to do. he said. could be done iust as well, or better, with entropy'replaced by the free-energy-temperature derivative -(aAlaT)~.When the free energy definition A = U - T S was rewritten with this derivat&e substituted, an equation without entropy factors was derived,
Equations of this form have come to he known as "GibbsHelmholtz" equations, although no such equation was mentioned by Gibbs (he preferred to work with the entropy concept); Helmholtz derived an equation like eq 1 (see ref. 4), hut his derivation was not the first. William Thomson, in one of his brilliant insights not exploited beyond a special application, arrived a t an equation of the "Gibbs-Helmholtz" kind in 1855,27 years before the free energy concept was discussed and named by Helmholtz (5).In the interests of historical accuracy we shall call equations of the form of eq 1"Thomson-Helmholtz" equations. We shall not use Thomson's derivation because it concerns thermoelasticitv. a tonic not relevant to Nernst's story. A broader equation,-derived by Nernst in his pioneering physical chemistry text (61, is worth mentioning, however, as a further lesson in how thermodynamics can be done sans entropy. One form of the equation given by Nernst was AA, = AU,
+ T (d(d"TA' ))v
with AU, and AA, representing internal energy and free energy changes for any chemical reaction. This was Nernst's version of the Thomson-Helmholtz equation. It was valid for reacting systems and had the same form as eq 1. An analogous equation for AG,, the change in Gihhs free energy for a reaction, could also be derived,
The question of the nature of the forces which come into play in the chemical union or decomposition of substances was discussed long before scientific chemistry existed. The Greek philosophers e the "love and hate" of atoms of matter. . .We themselves s ~ o k of retain anthr~pomorphieviews like the ancients, changing the names only when we seek the cause of chemical changes in the changing affinity of the atoms. To be sure, attempts to form more definite ideas have never been wanting. All gradationsof opinion are found, from the crude notions of Borelli and Lemery, who regarded the tendency of the atoms to unite firmly with each other as being due to their hookshaped structure.. .to the well-conceived ideas of Newton, Bergman and Berthollet, who saw in the chemical process phenomena of attraction comparable with the fall of a stone to earth.. . It is not too much to sav that there is no discoverv of anv
ingenuity displayed (7). Two of Nernst's more recent oredecessors in the study of chemical affinity were the pioneering thermochemists Julius Thomsen and Marcellin Berthelot. who both believed that chemical affinities were measured by the heats, that is, A?f,'s, of chemical reactions. The affinity principle asserted by Thomsen in 1853, for example, was that "every simple or complex action of a purely chemical nature is accompanied by an evolution of heat" (8).In other words, all spontaneous chemical reactions had to be of the exothermic kind, for which AH, < 0. The Thomsen-Berthelot principle was criticized by thermodynamicists, including Gibhs, Helmholtz, and Boltzmann, who cited instances of endothermic spontaneous reactions. Their analysis, based on the First and Second Laws of Thermodynamics, showed that only the free energy change AG, or AA, could indicate the spontaneous direction of a reaction and serve as a reliable measure of chemical affinity. The Thomsen-Berthelot principle was not entirely worthless, however. It did, in fact, agree with experimental ohservations in a large number of cases. Nernst appreciated these successes and thought they might be as important as the failures: It would he as absurd to give it [the Thomsen-Berthelot principle] complete neglect, as to give it absolute recognition.. . It is never to be doubted in the investiaation of nature, that a rule which holds good in many cases, b i t which fails in a few eases, contains a genuine kernel of truth-a kernel which has not yet been 'shelled' from its enclosing hull (9). Nernst was oartieularlv cognizant of the fact that the Thomsen-~ertheiot principle was more likely to he successful when it was applied to reactions involving solid-phase components. Equilibrium Constants
Nernst's fundamental aim was to develop a useful theory that exolained, and showed how to measure, the "affinity" chemical components have for each other when they react. This was an ancient problem, as Nernst noted in his Theoretical Chemistry:
In addition to the general problem of chemical affinity, Nernst had a strong interest in the practical problem of calculating equilibrium constants. His aim, like that of several others a t the time (oarticularlvFritz Haber), was to find a method for c a l c u l a t i ~equilibri;m ~ constants of gas-phase reactions. beginning with calorimetric data. Nernst's 1906 paper in which he Formulated his "heat theorem", later to become the Third Law, had the title, "On the Calculation of Chemical Equilibrium from Thermal Measurements" (10). Franz Simon, a colleague of Nernst's during the 1920's, tells of the prevailing concern with gas-phase reactions that inspired Nernst's work. (Simon's remarks were written in 1956, on the occasion of the 50th anniversary of the puhlication of the heat-theorem paper.)
I Nernst, and others at the time, were not always careful in distinguishing between AA, and AG,. Note the criticism by Lewis and Randall in Ref 20. p 158 (footnote).
Fifty years ago there was an intense interest in chemical gas reactions, partly because of the relative simplicityof the problem involved which seemed to lend itself to treatment by physical methods, and partly because of the economic possibilities. Gas reactions had already played an important role in the growth of
in which AHH,is the enthalpy change for the reaction. I n the further discussion we shall use this equation in preference to eq 2 because i t is more in keeping with modern usage and also because two equations introduced later (eqs 7 and 8 below), which are necessarily written in terms of AG, and AH,, have to he combined with eq 3.' The Conceol of Chemlcal Afflnltv
4
Journal of Chemical Education
chemical heavy industry, and it was realized that the ammonia synthesis in particular had became very important indeed for the German economy, both in peace and war.. . (11). Integration Problems . One of Nernst's goals was to use the Thomson-Helmholtz equation to calculate affinities for chemical reactions as measured by free energy changes. But the equation had to be integrated to derive a usable free energy equation, and there was a long-standing problem with the required integration constant. The integrated equation, expressed in terms of the Gihbs free energy change AG,, was (12)
in which AH, was the reaction's enthalpy change and TOwas the lower temperature limit chosen for the integration. An integration by parts reduced this equation to
with ACp, the heat capacity change for the reaction and J(T0) the integration constant, dependent as indicated on TO.The quantities AHr and ACp, could both be evaluated with calorimetric data, as Nernst required, hut the integration constant J(TO)could not he approached calorimetrically. It could be calculated in terms of the temperature derivative of the free energy change,
but no calorimetric method had been found to calculate the derivative. Nernst also hoped to integrate the van't Hoff equation,
a
was the standard enthalpy change for the in which reaction, to make the all-important equilibrium constant calculation for gaseous reactions. This equation, too, had an integration problem, in fact essentially the same integration problem as that of the Thomson-Helmholtz equation. The van't Hoff equation integrated between the temperature limits T and To could he expressed
As Nernst saw it, the theory ofaffinitv calculations hinged and J ' ( T d . If they on the t a u integr~ttrmconstants /(To) could be evaluated. methurls for determinine affinities and equilibrium constants from calorimetric data could he develooed. The inteeration constants could he determined if. in t i r n , the derivative (a(AG,)laT)p,~,could he calculated. The prohlem was solved if the derivative could he evaluated a t any temperature, but it was clear to Nernst, and to others a t the time, that the calculation could be done simply and directly only a t a certain temperature. According to Simon, "Nernst had a hunch that [nature] could reveal her intentions only a t absolute zero, the one point of special significance in the whole range of temperature" (13), so Nernst proceeded with his analysis assuming that TO= 0 K. But this choice led to another stubborn prohlem. Equations 5 and 9 could apparently not be applied to gas-phase reactions because for them there was no reason to expect zero as the temoerature avthat ACD. . . or A& . , aonroached .. proarhrd zero. That mrnnt that the integrands in eqsSand 9 became infinite at the, h e r temverature limit if T I = U K was assumed. Simon tells us that' "this had really @en the root of the urohlem in the vrevious a t t e m ~ tto s calculate the integration'constant" (13f "Now comes Nernst's decisive step." Simon writes. "Leaving aside the gas reactions for the moment he switched to the condensed phases, which so far had not aroused the attention of the chemists" (13). According to the classical theory available to Nernst, i t was expected that ACp, for solidphase reactions, unlike gas-phase reactions, should approach zero a t low temperatures. For reactions of that kind eqs 5 and 9 could apparently he applied. Continuing with Simon's account:
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Now we are prepared for the final step.. . According to Berthelot's principle spontaneous reactions develop heat, that is, they take place when the total energy decreases [AHr < 01. Of course, we know from the [Thomson-Helmholtz]relation that this cannot be strictlv true: it is the free enerev ... which determines the direction of the process S\'PVPTL~&CF rhtre is n oreat amoonr of w drnce sh.rxing that it is often a very good nppraxmation, pnrticulady at not too high temperatures. Therefore Nernst argued that AG, and AHrcannot be very different from each other, at least at roam temperatures and often at higher temperatures, and that the AG, curve [vs. TI is not likely to start at a steep angle at absolutezero. The simplest assumption was that the AG, and AH, curvesshould eventually [at T = 0 K] run together, andasa(AGJ1 d T can be finite only if B(AHJ/dT = ACp, becomes zero, Nernst arrived at his famous formulation, at first restricted to condensed states: ~
~
~
~
, The equilibrium constant K ( T ) was related to A G ~ ( T )the standard free energy change for the reaction, according to It was as simple as that ( 1 4 ) , To summarize, Nernst's assumption was simply that for solid-phase reactions the derivative (a(AG,)/aT)p had to he finite or zero a t 0 K. This fundamental assumption led to the several conclusions mentioned by Simon. According to eq 5, Nernst's assumption was not permitted a t all unle& Integration of this by parts reduced it to
with $(To) another integration constant whose value was determined by the temperature TO chosen for the lower integration limit. Like J(To), this constant was dependent on the temperature derivative of a free energy change, in this case the standard free energy change A G ~ :
(solid-phase reactions; T = 0 K). Then if (a(AG,)Ic3T)p was finite or zero a t 0 K, as assumed, eq 3 guaranteed that AG, = AH.
(12)
(solid-phase reactions; T = 0 K), and therefore that
(solid-phase reactions; T = 0 K). Combining eqs 11 and 13 Volume 64
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Nernst could finally conclude that
(solid-phase reactions; T = 0 K). The combined effect of eqs 11,12, and 14 was to predict that AH, and AG, data should approach each other asymptotically as T approached 0 K. Nernst and others found impressive evidence for this kind of behavior in a variety of solid-phase reactions (15). Thus Nernst concluded that, providing they were applied to reactions whose components condensed into solid phases a t some low temperature (so eq 14 could he applied), eqs 5 and 9 could he written with the integration constants given values of zero, and with the lower integration limit To = 0 K,
chose to set the conditions so the reaction of interest took placeat a temperature low enough to permit the components to he in equilihrium with their (separate) solid phases. For a simple reaction involving the reactant A and the product R, Nernst's two-phase reaction system was
The two free energy changes AG, and AG; for the gasphase and solid-phase reactions were the same in this system,
and
AG, = AG:
(reaction components condensible into solid phases). Equation 14 was the embodiment of Nernst's principle, which he at first called the "new heat t h e ~ r e m " .For ~ the enlightenment of readers familiar with modern Third Law literature, Nernst's statement can he revised so it concerns the entropy change AS, for the reaction a t 0 K. The entropyfree-energy connection,
is introduced, and Nernst's principle becomes (solid-~hasereactions: T = 0 K). Most modern thermodvnamic; texts call this Nernst's heat theorem. Hut that practire would certninlv nut have pleased thr inventor ofthe heat theorem. Nernst never made use of the entropy limiting condition, except as an afterthought. The free energy concept was always more tangible to him and therefore preferable to the entropy concept. Chemical Conslants (True)
Nernst had no doubt that he had discovered a fundamental principle, but his theorem had little to say about the original problem, calculation of equilihrium constants for gas-phase reactions. He now had to find his way hack to the gaseous reactions. As before, he began his equilihrium constant calculation by integrating the van't Hoff equation, eq 7, (see ref. 16). But gas-phase reactions needed a different procedure; the integration constant could not be expected to vanish for the gaseous reactions and alower integration limit of zero was not allowed. Here the integrated equation could he written
with the integral put in indefinite form and I representing the integration constant that Nernst was primarily interested in calculating. The integration constant I could not he given a zero value, hut it was independent of temperature, so it could he evaluated once and for all, under convenient conditions. Nernst
so AG;, which Nernst could calculate with eq 15, derived from his heat theorem, could then he used to determine the free energy AG, for the gaseous reaction of interest. According to eq 15,
The reaction isotherm could he used to determine AG,
with the standard free energy change AG: related to the desired equilihrium constant K according to eq 8. The quantity lnK was evaluated with eq 18, and the two pressuresp~a n d p ~since , they were determined by suhlimation equilibria, could also he calculated with eq 18. The sublimation equilihrium for A was A(s) = A(d for which K = PA, and eq 18 could be written
where @A was the standard enthalpy change for the suhlimation of A and iAwas another integration constant. Similarly,
Combining eqs 18-23, Nernst could arrive at
The enthalpy terms on opposite sides of this equation very nearly cancelled each other, since Hrrepresented the direct @represented solid-phase reaction while the same process taking place indirectly in the three steps
a,+
A(s)
Nernst's original statement was made in terms of the temperature derivative of AA ,rather than AG, (and Nernst used Me symbol A for AA ,). Here, as elsewhere in this paper, AA ,is replaced by AG,, and AU, by AH,. See remarks following eq 3. 6
Journal of Chemical Education
(19)
- - A(g)
R(g)
wR
R(s)
The remaining equation held identically for all values of T , and therefore coefficients of the same powers of T on the two sides of the equation could he equated. Since none of the integrals involved constant terms (all such terms having
been removed to the integration constants), the only linear terms in the equation were those with -IR, ~ R Rand , - ~ A Ras coefficients. These coefficients had to total zero,
so the desired integration constant I was simply the difference between the integration constants belonging to the product R and the reactant A,
Similar formulas could he derived for more complicated gas-phase reactions. The desired equilibrium constant could he calculated using eq 18,and the general recipe for calculating the integration constant I was
Nernst called the i integration constants "true chemical constants". They had to be calculated with equations of the eq 22 form, to which vapor-pressure data obtained in the study of sublimation equilihria could he fit. Chemlcal Constants (Conventional) Nernst's study of gas-phase equilihria demonstrates, if any proof is needed, that the paths of theoretical research can he devious. As we have seen, he had started out by attempting to integrate the Thnmsnn-Helmholtz and van't Hoff equations hut a t first conld find no way to do that directly for gaseous reactions. He had next turned to the little-knownthermodynamics of solid-phase reactions and uncovered the experimental basis for his heat theorem, with imnlications reachine far hevond the realm of eas eauilihria. ~ e t u r n i n gto the gaseous "reactions and fo;mula'ting his problem as always in terms of temperature-independent integration constants, Nernst had then managed to express the inteeration constants he wanted as summations of seoa" rate constants, one for each reaction component. In arriving at this result. Iternst had done what chemical thermodvnamicists have always done best: he had dissected a compl~cated calculation, possibly involving mans chemical components, into R component-by-component analysis. For each component achernical constant could he calculated and rahulacd "once for all .. . most directly from vapor-pressure curves of the substances concerned in t h e . . .solid state. . ." 117). But Nernst's task was still not complete. The vapor-pressure and heat capacity data required by eq 22 for accurate calculation of chemical constants were not available when Nernst first formulated his theory in 1906. He was aware of these limitations and soon embarked on one of the first programs of experimental work aimed a t obtaining the necessary heat capacity data. But as an interim measure he developed an empirical vapor pressure formula for suhlimation equilihria which had the form @
a
b
lnp = - - + - l n ~ - - ~ + i ' RT R R
(25)
(see ref. IS), in which a and b were constants, and i' was an approximation to the i term in ea 22, which Nernst called a "c&ventional chemical constani". erns st found he could assign the value 712 to the constant a while b and i'had to he calculated for each gaseous component of interest. On about the same level of approximation Nernst also derived an empirical equation for calculating equilibrium constants,
in which AH, was an average enthalpy change for the reac-
tion at ordinary temperatures, An, was the change in stoichiometric coefficients of gaseous components for the reaction, and A(ni'), was the change in the quantity nP, also for the reaction's gaseous components. Nernst calculated a tahle of values for his conventional chemical constants i' and used his approximate eq 26 to determine equilibrium constants. "Surveying the whole materialavailable at that time,"Simon tells us, "he showed that results of his calculations agreed with the experimental facts within a rather generous limit of error" (19). In a typical application, eq 26 and the tahle of conventional chemical constants calculated K = 1.82 for the water gas reaction,
at 1073 K. The observed value for the equilibrium constant at this temperature was K = 0.93. T o some of Nernst's colleagues this kind of agreement was not verv imnressive. G. N. Lewis. a former student of Nernst'; a n d one of his most impokant successors in the development of the methods of chemical thermodynamics, credited Nernst's efforts to obtain the data needed for accurate calculations with eqs 18 and 22 but deplored "the rapidly growing use of [conventional] chemical constants". Lewis found dismavina "the various efforts which have been made to square the cafculations based on these constants with the results of measurements.. . [They] constitute a regrettable episode in the history of chemistry" (20). Lewis would note with approval that the theory and practice of chemicalconstant calculations are nowhere to he found in the modern literature of chemical thermodynamics. But Simon reminds us that the "eenerous limit of error" with which Nernst measured his success "was infinitely preferable to the comnlete ienorance which existed before.. . INernst'sl anproximations were very useful to chemical industry, making it possible to get very quickly a rough idea which reactions were thermodynamically feasible in complex reaction patterns" (21). The Theorem Is a Law Nernst made it clear that his heat theorem was fundarnentally a "lau," and not just another formula or mathematical rccipr. Ilr insisted not only that his theorem belonged with the two established laws of thermodynamics but also that there could newr br anuther thermodsnarnic law. This conclusion followed from an extrapolation of the facts that the discoverers were three (Mayer, Joule, and Helmholtz) for the First Law, two (Carnot and Clausius) for the Second Law, and iust one (W. Nernst) for the Third Law. With no one to discover it, a Fourth Law of Thermodynamics could not exist. The Third Law was the Last Law. With all his immodesty, self-glorification, and sarcastic wit, Nernst continued to expand his influence not only among the high and mighty hut also within the intimate circle of graduate students. James Partington, an Englishman who worked in Nernst's Berlin laboratory, writes of Nernst's kind attention to a "very young man, with little ex~erience". Unlike manv scientific potentates then (and now) Nernst did not ignore the daily iabors of his research students. leaving them to sink or swim. Partindon found his researchdifficuk, hut Nernst's presence was-an incentive: ". . . one felt that he conld do the work easily himself, and that perseverance would remove lark of skill; a fault which l ~ application.. y . His true kindness is somecould i ~ cured e thing I remember with gratitude. . ." (22). At least one visitor to Berlin in the 1930's had initial reservations about Nernst and his unusual manner. Hendrik Casimir gives his impressions of Nernst and his research colloquium: K u r t hlendclswhn hasdescribed t h i i i n i t i t u t i o n Ithecolloquium] i n mthusiarti, terms a\ the place w l ~ r ethe most prominent pltysici,ts uf t l w dn, prono~nccdjudrrncnt on the most recent
Volume 64 Number 1 January 1987
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developments. It did not strike me that way at all. . . discussions were both formal and perfunctory. But it was an experience to listen to Walther Nernst. In a fairly soft, yet penetrating, rather high-pitched voice he could proclaim that he had already said some of the things presented at the colloquium in his book, and complain that people apparentlydid not recognize that as a publication. He struck me at the time as a ridiculous figure.. .Later, I realized that same of his remarks had contained a rather subtle point. In 1964, the centenary of his birth was celebrated at Gettineen and I was invited to eive the main talk. On that occasion. I studird hi; published work mureowr;tnmuyhout n remnrkaoly rlror and often prophetic vision. And so I had an opportunity to atone in public for an error of judgment I had never voiced (23). Joy and Sorrow Nernst's nrivate life was almost as extraordinarilv fortunate as his professional career. His wife Emma was, among of domestic efficiencv and manv other thinas. - . a naraaon . hard work. She customariiy rose a t 6 a.m. and kept the Nernst households, which were never simple or quiet, in perfect order. Not long after their arrival in Berlin, the Nernst's were known as the most hospitableacademic family in Berlin. There were five children, three daughters and two sons, and family life was an important part of the Nernsts' existence. But no one, not even Nernst, could escape the tragedies of two world wars. In the first war both Nernst sons were killed. Long before the armistice, Nernst could see that Germany was beaten and nearlv ruined. In vain. he tried to use his , others) to prevent furconnections (with t h e ~ a i s e ramong ther devastation. In 1917, he found escape in the peaceful realm of science by gathering in amonograph his work on the Third Law. The opening sentences of this book tell of the solace he found in Bcienci:
In times of trouble and distress, many of the old Greeks and Romans sought consolation in philosophy, and found it. Today we may well say that there is hardly any science so well adapted as theoretical physics to divert the mind from the mournful present
. . . (24).
Peace finally returned and miraculously Germany recovered. For a time there was political and economic chaos, hut German science emerged as strong and active as ever; Nernst and his Berlin collea~uescould reconstruct scientifically. Now there were conceptual revolutions to he fought; both the quantum theory and relativity had come over the scientific horizon. Nernst did not contribute extensively to these endeavors, although he understood and appreciated what was happening. More recognition came his way; he was offered (but refused) an ambassadorship to the United States,
he was elected Rector of Berlin University, became a Helmholtz successor as Professor of Physics, and won the Nobel Prize in Chemistry in 1920. Dammerung
But Germanv had not comnletelv recovered from the nolitical ruin broight on by the k t war. In the 1930's the Nazi influence beaan to crow. and then Hitler was in Dower. With incredible ignoranEe and insane anti- emiti ism, the Nazis utterly destroyed German science. Only a handful of the dozens of scientific leaders populating German universities remained. A list of the scientists who fled Germany and the Nazi influence reads like a scientific Who's Who of the 1930's and 1940's. Nernst was strenuously opposed to the Nazi policies, but lacked the energy and influence to act. Mercifully, he retired to his country home and found a measure of peace in the last seven years of his life. During his final days, Mendelssohn tells us, Emma sat with him, and recorded his last words. "True to his whole character," Mendelssohn writes, Nernst told Emma just before he died, "I have already been to Heaven. It was quite nice there, hut I told them they could have it even better" (25). Lnerature Cited 1.
Mendel~sohn.K. The Worldol Wolther Nernat: Uniu.Pittabugh: Pittnburgh,
3. Ref l . p p 4 1 4 2 .
He1mholh.H. "On theTherrnodynnmier of Chemical Prorasres".Engliahfranrlstion in Phyr. Mem.. Phys.Soc. Lmdon 1888.1.53. 5. Thornson, W. Mothamoticol and Physic01 Poprrs: Cambridge Univ.: Cambridge. 19n;Val. 1, p 297. Also see Partington. J. R. An Advonred Tr~oiisaon Physical Chrmirliy: Longmans: London. 1949; Vol 1.p 182. 6. Nernst. W. Thromicai Chemistry; 7th ed., Enelish translation by Tizard. H.T.; Marmillan: London, 1916; pp 22-24. 7. Ref6, p 473. 8. Quoted in Partington, J. R. J . Chem. Soc. 1953.2860. 9. R d 6 . p 744. 10. Nernst. W. Nochrichten uon der Koniglichrn Gesebchn/t der Wimenscholten i u 4.
Cnllinnrn ~.. .".. 1 9 0 6 1 ~~
11. Simon F. YearbookPhys. Soc. 1956.2. 12. sse c l a r ~ m n eS. , ~ e r t b o o ko i ~ h y s i r n~hamiatry: l 2nd ed.; van ~ o e t r a n dN: ~ WYWL. 1946; pp 859.864. ior a good coverage of derivations of equations quoted in this section and later. 1%. Ref 1 , . p 4 . 14. Ref 11, pp 4-5. Free enorgyand internal enorgy changes AA,and AU. (Simon uses the qvmbols Aand U for these qusntitiesl have heen changed to the Gihhlfree e n e x y change A0, and enthalpy change AH,throughout this passage. 15. See Nernst, W. TheNeu~HeaITheorem:Dover: New York,1969: Chapter IX. 16. For Normstis moreabbreviated treatment of thissubjectsee Ref15, pp 125-126. 17. Ref 15.p 126. 18. See Ref 12,p 863. for sgaod summaryof Nernrt'sspprorimatemethoda foresleulsiing and using his chemical mnstants.
Professional Issues in Chemical Education The American Institute of Chemists will present an invitational symposium on "Professional Issues in Chemical Education" during the 1987 AIC national meeting, Sunday, April 26, 1987, at the Four Seasons Hotel in Philadelphia. Features speakers include Bassam Shakhashiri (National Science Foundation),Moses Passer (American Chemical Soeiety), James N. BeMiller (Purdue University), Nina Roseher (American University), K. R. Fountain (Northeast Missouri State University), Arthur Breyer (Beaver College), David Katz (Corninunity College of Philadelphia), and William R. Blean, Jr. (Radnor High School). The full-day,two-session program addresses professional challenges facingchemists teaching at universities,colleges, community colleges, and high schools. Topics include teacher training, professional development,undergraduate research, chemical education for the handicavved. and contributions to the elementarv school science curriculum.
i 8
Journal of Chemical Education
1973:p
2. Nernd, W. 2.phys.Chem. 1883.4.129.
