One-Dimensional Organic Conductors TTF-TCNQ and other Organic Semi-metals," Energy and Charge Transfer in Organic Semiconductors, (Eds. K. Masuda and M. Silver), Plenum Press (1974). 39-f. O. Poehler, A. N. Bloch, T. F. Corruthers, and D. O. Cowan, "The Organic Metallic State: Some Physical Aspects and Chemical

Robert W. Hart (left) and Samuel N. Foner

SAMUEL N. FONER is Vice-Chairman of the Milton S. Eisenhower Research Center and Supervisor of its Electronic Physics Group. Born in New York City (1920), he studied physics and mathematics at what is now the Carnegie-Mellon University, where he received his D.Sc. degree in physics in 1945. He was employed as an instructor in the Physics Department and later as a research associate of the Manhattan Project, working in the laboratory of the Nobelist Otto Stern who instilled in him the use of conceptually simple experiments to answer complex questions.

Dr. Foner joined APL in 1945 and became associated with the Research Center as Supervisor of the Mass Spectrometry Group (1947-52) and the Electronic Physics Group (1953-present). He has made many noteworthy contributions to the mass spectrometry of free radicals

RESEARCH RETROSPECTIVES

THE STRUCTURE OF FLAMES

Flames have been the most important source of heat, light, and power since the earliest days of civilization. At present, the combustion of fuels is, by far, the largest chemical operation under human control. Yet, until quite recently, detailed knowl­edge of what goes on within a flame did not exist. Although the complexity of combustion is not en­tirely understood, even today, what was virtually terra incognita has been opened up during the past 25 years by the classic studies at APL by Robert M. Fristrom, Arthur A. Westenberg, and their col­leagues.

What does one need to know about a flame? Chemists want a detailed accounting of the steps by which fuels (such as oil, natural gas, or coal) and oxidizers (such as air) are converted into products of combustion (water, oxides of carbon, soots, and

January- March 1980

Trends," Proc. NATO Conference on Chemistry and Physics of One­Dimensional Metals (Ed. H. J. Keller), Plenum Press (1977). 4OR. S. Potember, T. O. Poehler, and D. O. Cowan,"Electrical Switch­

-ing and Memory Phenomena in Cu-TCNQ Thin Films," Appl. Phys. Lett. 34, p. 405 (1979).

and reaction intermediates, to the detection of free radicals stabilized at very low temperatures, and to the ionization of substances by electron impact.

He has served as a member of the NAS/NRC Adviso­ry Committee for the Army Research Office and an Ad­visor to NATO's Scientific Affairs Division. In 1954, Dr. Foner received the Physical Sciences Award of the Washington Academy of Sciences for work in free radical chemistry and physics. He is a member of the Combustion Institute, the Philosophical Society of Washington, and a Fellow of the AAAS, the American Physical Society, and the Washington Academy of Sciences.

ROBERT W. HART is Chairman of the Milton S. Eisenhower Research Center and Assistant Director of APL for Exploratory Development. Born in Yankton, SD, in 1922, he studied at the University of Iowa and received his Ph.D. degree in physics from the University of Pittsburgh in 1949. After a year of teaching at the Catholic University in Washington, he joined APL in 1950. He has been a member of the Research Center ever since.

During the 1960s, Dr. Hart developed a detailed theory of the complex combustion behavior of solid pro­pellants in rockets in collaboration with the late Frank T. McClure and as member of the Joint Armed Services Committee on Combustion Instability. As Supervisor of the Special Problems Research Group (1954-1975), his interests covered theoretical aspects of wave scattering, the structure of the eye, and other phYSical and biophysical topics.

Dr. Hart is a member of the American Physical Socie­ty and of the Combustion Institute. Outside of profes­sional activities, he is interested in the origin and evolu­tion of civilization and science.

ash) as well as the intermediate reaction paths that are involved in this transformation. They want to understand why and how inhibitors can extinguish flames or prevent engine "knock" and know how rapidly these transformations can take place. Physicists, on the other hand, are interested in temperature effects, radiation, the flow fields set up by the gases moving into and out of the flames, and countless other physical properties.

What sets flames apart from more conventional chemical transformations is that one is dealing with a very intricate situation in which chemical reac­tions are closely coupled with the physical flow of substances into and out of a reaction zone, accom­panied by a steep rise in temperature, abrupt changes in composition, and numerous optical and electrical phenomena that may be important under specific circumstances. In a distance of less than 1 mm, temperatures can change by thousands of

33

degrees and gas concentrations can rise or fall abruptly. The thinness of this transformation zone held back experimenters in exploring the detailed structure of flames.

A fundamental advance in understanding the simplest combustion case (where gaseous reactants are mixed prior to combustion and no solid reac­tion products are formed), came about in the early 1950's. First, a theoretical analysis of flames was made l whereby simplified models of combustion processes could be analyzed in detail. This led to suggestions of how one could, in principle, separate the chemical transformations from the simulta­neous physical processes (mainly diffusional). Sec­ond, it was found that most flames, when stabi­lized at pressures well below atmospheric, would widen in thickness without altering the sequence of the chemical transformations. This made it possible to introduce sampling probes and thermocouples into the reaction zone and obtain point-by-point samples for subsequent analysis of the concentra­tion of reactants, intermediates, and products as well as temperature profiles.

APL's work pioneered in the development of these experimental tools and their application to the analysis of simple flame systems. Mass spec­trometers were found to be useful in identifying the chemical species that survived the sampling process. For highly reactive intermediates such as radicals and atoms, stablizing techniques were developed to preserve them for later analysis. Tiny microprobes and thermocouples were built to obtain spatial resolution in flames whose thickness now extended over several millimeters. 2

Once the value and workability of these tech­niques were recognized, they were quickly adopted by others. A vigorous worldwide exploration of many flame systems was started. Much was learned about the intricate interplay by which stable reac­tants transform into stable products by way of in­termediate steps and, in particular, about the cru­cial role played by free radicals in mediating the rapid transformations.

The picture that has emerged3 of the structure of, for example, a hydrogen-oxygen flame with ex­cess hydrogen is that initially, as the reactant gases enter the combustion region, they are heated some­what by conduction as heat flows toward the in­coming gases from the hot combustion products. A little later, heat-producing chemical reactions com­mence. They are initiated by reactive free radicals that diffuse from the hot side of the flame, where they are generated, toward the incoming gases and attack hydrogen molecules, leading to the forma­tion (as well as the disappearance) of the hydro­peroxo radical (H02 ) and water. At a still later (and hotter) stage, the relatively slow reaction 0 + H2 ~ OH + H becomes important. Because more free radicals are generated than are consumed, this reaction stage is the source of the reactive free radicals that are active earlier in the flame. Finally,

34

after the main rapid reactions have taken place, a sorting-out zone follows where the free radicals that were formed in excess amounts in the previous stage recombine relatively slowly until the system settles down to an equilibrium state in which reac­tions respond only to the slow temperature changes brought about by heat losses.

For the "simple" hydrogen-oxygen flame, where most of the possible intermediate steps have now been identified and their reaction rates measured, the overall behavior of the flame can be described by a quite limited number of reaction steps. 4 For hydrocarbon flames, many more chemical interac­tions are possible because of the presence of carbon in the fuel molecule. An overall description of hy­drocarbon flame behavior based on individual steps is now nearly in hand. Here, too, the role of free radicals as crucial mediators (hydroxyl radicals if the flames are deficient in fuel for complete com­bustion, or hydrogen atoms if there is an excess of fuel) is beyond doubt.

During the combustion of hydrocarbons with ex­cess air, the oxidation of hydrogen and carbon monoxide (which appear as intermediates in the flame) is of particular importance. The hydrogen oxidation furnishes the oxygen and hydrogen atoms and hydroxyl radicals that attack the hydrocarbon molecule to form methyl (CH3 ) radicals. The latter subsequently interact with oxygen to form formal­dehyde, which reacts in further steps to produce carbon monoxide and, ultimately, carbon dioxide. This sequential formation and disappearance of in­termediates (carbon monoxide, hydrogen, formal­dehyde, and atoms and radicals) was clearly shown in the experiments of Fristrom and WestenbergS

and is fully supported by the detailed reaction scheme proposed for hydrocarbon oxidation in flames 6 (Fig. 1).

In fuel-rich flames, many more intermediate steps are possible. Polymerization reactions that lead to the temporary appearance of hydrocarbons that are of higher molecular weight than the initial fuel become important. They give rise, subsequent­ly, to aldehydes and unsaturated hydrocarbons. A complete description of all the individual reaction steps is not yet possible because of the lack of reac­tion rate data for all of the numerous elementary reactions involved.

One goal of the flame structure work was to ob­tain quantitative information about individual reac­tion steps. This turned out to be difficult in prac­tice because so many reactions are proceeding si­multaneously even in the simplest flame that it is nearly impossible to single out anyone for detailed analysis-with one important exception. In the hot stream of combustion products beyond the active reaction zone, free radicals are still present in measurable amounts. This region can be used as a "hot bath" where the reactions of compounds with free radicals can be studied. It became evident that it was possible to inject traces of well-known flame

Johns HopkinsAPL Technical Digest

c: o 0.15

.~ c: <I> (,,)

c: o (,,)

v 0.10 I u '0

:ii 6 IN 0.05

6 u

0' u

N I '0

~ 0.004

1jN 0.0021:1....--"---:

0 H2O

cO2

°2

CO CH4

H2 CH20

2000 rr---- r-------,---- --r-----,------,

Q -; 1500

~ ~ 1000 ~

~ 500 u.

Distance along f lame (mm)

0.90

*-'0

0.85 ..s c: 0

.~ c: <I>

0.80 g 8 0'

0.75

Fig. I-In 1960, a low pressure fuel-lean flame of methane and air was analyzed by Fristrom et als for the appearance and disappearance of chemical species and for local temperature. The reaction zone was approximately 4 mm thick. Experimental results are shown as individual data points. Recently, J . War­natz6 was able to calculate composition and temperature pro­files of the same flame from known reaction rates of approx­imately 80 of the important individual steps that are postulated to occur during the oxidation of methane (solid lines) . The agreement between prediction and experiment is excellent.

inhibitors (methyl halides) into this bath, determine their rates of reaction with hydrogen atoms at the high temperature, and compare the results with ex­trapolations obtained with entirely different tech-

THE SEARCH FOR H02

In more than two hundred years, chemists have isolated and identified about 100 chemical elements and millions of compounds into which these elementary building blocks can be combined. This continuing and unending quest (the number of potential combinations of elements into compounds is virtually limitless) has been accompanied by an intensive effort to learn more about the bonds that hold these building blocks together in recognizable structures and shapes and to discover the rules that determine the rates and the pathways by which one chemical structure changes into another.

Until about 50 years ago, the study of chemistry was based on a belief in stability. To be sure, many

January-March 1980

niques at lower temperature. 7,8 For the first time it

was demonstrated that these measurements give concordant results.

Many intriguing and important problems remain, especially in applied areas where fuel! oxidizer mix­ing limitations, catalytic surface effects, soot for­mation, and many other subtle interactions may lead to undesirable end effects. However, for the central problem of gas phase combustion, 200 years of flame research have, at long last, brought about a remarkable confluence of theory and experiment. Taking into account the dominant physical process­es of diffusion and heat conduction and the numerous interacting chemical reaction steps, the structure of flames can now be viewed in its full in­tricacy. Complex flame systems can be constructed out of the many individual reactions that proceed within a flame and the overall behavior of such flames can, in principle, be predicted.

WALTERG.BERL

REFERENCES

IJ. O. Hirschfelder, C. F. Curtiss, and D. E. Campbell, "The Theory of Flame Propagation," J . Phys. Chern. 57, pp. 403-4 14 (1953) .

2R. M. Fris trom and A. A. Westenberg, Flame Structure, McGraw- H ill (1965) .

3G. Dixon Lewis, "Kinetic Mechanisms, Structure, and Properties of Premixed Flames in Hydrogen-Oxygen-Nit rogen Mixtures," Phil. Trans. Roy. Soc. 292, pp. 45-99 (1979).

4 J . Warnatz, "Calculation of the Structure of Laminar Flat Flames II: Flame Velocity and Structure of Freely P ropagating Hydrogen-Oxygen and Hydrogen-Air Flames," Ber. Bunsenges. Phys. Chern. 82, pp. 643-649 (1978) .

SR. M . Fristrom, C. Grunfelder, and S. Favin, "Methane-Oxygen Flame Structure I: Characteristic P rofiles in the Low-Pressure, Laminar, Lean, Premixed Methane-Oxygen Flame," 1. Phys. Chern. 64 pp.1386-1392 (1960)

6J . Warnatz, "Flame Velocity and Structure of Laminar Hydrocarbon Flames," Proc., Seventh International Col/oq. on Gas Dynamics of Ex­plosions and Reactive Systems (to be published) .

7L. W. Hart , C. G run fe lder , and R. M. Frislrom, "The ' Point Source' Technique Using Upstream Sampling fo r Rate Constant Determi nat ions in Flame Gases," Combust. Flame 23, pp. 109-119 (1974)

8A. A. Westenberg and N. de Haas, " Rates of H + C H3X Reactions," J. Chem. Phys. 62, pp.3321-3325 (1975).

levels of stability were identified and techniques were developed to move from one level to another. Some substances were so labile that they would barely survive at room temperature. At high temperatures most compounds would change into a relatively small number of stable products. A few elements like radium and polonium showed signs of instability. But once substances such as hydrogen (H2) and oxygen (02) molecules reacted with each other, they were expected to form only water or, on occasion, hydrogen peroxide (H20 2). The details of how such transformations take place were but dimly perceived. It was generally assumed that a direct reaction (commonly written as 2H2 + O2 -. 2H20) occurs that involves no other chemical species.

35

A change in this simplistic view came about in the 1920's. Researchers working in the area be­tween physics and chemistry became familiar with the details of explosions and flames where chemical reactions proceed at speeds well beyond the leisure­ly pace of the conventional chemical reaction. Puzzling observations were made. For example, if a container is filled with hydrogen and oxygen at a high enough temperature, its contents would always explode during mixing. But by lowering the temper­ature below a critical level, the reaction between the two gases settles down to a slow rate of water formation that might take hours to go to comple­tion. However, when one starts to withdraw such a slowly reacting mixture from the container by means of a vacuum pump, the remaining mixture of hydrogen, oxygen, and water vapor suddenly (and always at a precisely reproducible pressure designated the "explosion limit") explodes with a bright flash and the reaction is completed instantly.

This extraordinary behavior of a sudden transi­tion from a slow reaction to a very fast one by a mere change in pressure became a turning point in the interpretation of chemical reactions. A general explanation was provided by the Russian chemist, N. N. Semenov, 1 who proposed that highly reactive atoms and so-called "free" radicals mediate the transformation among the reactants. These in­termediaries, whose concentration depended on the particular conditions of the experiment, were, in fact, crucial participants in the reaction. Their very reactivity made it impossible (at the then-existing state of the art) to detect or to isolate them by con­ventional chemical means. Because of their inac­cessibility, they had not been thought of by chem­ists as being important in chemical reactions even though physicists had been aware of their existence in many experiments in which gases were exposed to either high temperatures or electric discharges.

For the hydrogen-oxygen system, not many choices of intermediaries are possible. Hydrogen atoms (H), oxygen atoms (0), and hydroxyl radicals (OH) were already known in physics. One could postulate that water is formed by the elemen­tary reaction

OH + H2 -+ H20 + H, (1) in which one reactive substance (OH) disappears but another one (H) appears. This atom, in turn, would react with oxygen

H + O2 ~ OH + O. (2) The OH formed in (2) would react as in (1) with more hydrogen to form another molecule of water. The oxygen atom, on the other hand, can react with a hydrogen molecule in the following way:

o + H2 ~ OH + H. (3) According to this scheme, once a reactive

substance is provided to initiate the process, water can be formed in a simple series of chain reactions in which the key radicals are regenerated endlessly. In fact, in reactions (2) and (3) the number of "chain carriers" increases in each reaction step.

36

However, these three reactions are inadequate by themselves to explain the observed events. Since more chain carriers are formed than are used up, the formation of water would be expected to become progressively faster with time and always end up in an explosion. A way had to be proposed to remove the chain carriers as quickly as they are formed in the region where the reaction proceeds slowly and to let them build up at a rapidly in­creasing rate only at the explosion limit.

A careful analysis of the experimental observa­tions on how the explosion limit is affected by gas temperature, container size, pressure, mixture ratio, and the presence of inert diluents led to a reaction scheme that forced the postulation of one addi­tional intermediary, the hydroperoxo radical (H02), which has the extraordinary properties of being a stable enough free radical to compete with the chain reaction involving hydrogen atoms (reac­tion (2» but also unstable enough to be destroyed once it reaches the wall of the reaction vessel. Only by postulating its formation by the reaction

H+02 + M-+ H02 + M (4) (where M can be any of the stable participants in the reaction) was it possible to account for the observed system behavior in detail. One of the pro­ponents of H02 said about its existence "These conclusions are inescapable.,, 2

When this "inescapable" interpretation was pro­posed in the 1930's, there was no experimental evidence from any source for the physical existence of H02 nor would there be for 20 years. Yet after studying reactions involving the oxidation of hy­drogen-containing materials that included the entire family of hydrocarbon fuels and of myriad other organic compounds, physical chemists had few doubts that the proposed free radical must exist. Without H02, the observed system behavior simply could not be interpreted; with it, the observations became understandable.

In the middle 1940's there was a brief report from the Shell Research Laboratories that when a hydrogen/ oxygen flame is placed in front of the sampling port of a mass spectrometer, it would produce a molecular fragment of the correct mass (33) attributable to H02. But the evidence was flimsy. The system was too complex and too dif­ficult to analyze. Convincing experimental evidence for H02 was still lacking.

It was another 10 years before the definitive ex­periments that would identify H02 were made by Samuel N. Foner and Richard L. Hudson at APL in 1953.3 This was done by combining a sensitive mass spectrometer with a reliable "molecular beam" inlet and a reacting system that could leave little doubt about the events that were occurring. The reaction scheme was to generate hydrogen atoms in a separate electric discharge, mix them rapidly with oxygen molecules before all the hydrogen atoms had time to recombine, and analyze the reaction products for H02. In this in-

Johns Hopkins A PL Technical Digest

itial effort, the mass spectrometer, despite its sen­sitivity, was hard pressed to show a peak at mass 33 caused by H02 • The peak was there, never­theless, disappearing when the hydrogen atoms were shut off and reappearing when they were turned on again (Fig. 1).

Once such a trailbreaking discovery is made, a veritable flood of additional information common­ly comes to light quickly. 4 Other reactions were found (mainly involving hydrogen peroxide) that proved to be more convenient sources of H02 in much higher concentration. Many rates of H02

reactions with itself, with other simple molecules, and with hydrocarbons have now been measured. 5

The spectra of H02 were obtained, from which many of its structural properties have been estimated. 6 Its terrestrial existence is beyond ques­tion. Only its presence in interstellar gas clouds, where so many other free radicals have recently been discovered, has yet to be established. 7

Nearly 20 years later, H02 made one more ap­pearance in the APL research effort. The reaction of H02 with carbon monoxide was proposed at one time as an important link in atmospheric chemistry research. If the reaction proceeded rapid­ly, it would offer an attractive pathway for eliminating the carbon monoxide that is generated when fossil fuels are burned and for whose scavenging from the atmosphere no really satisfac-

Discharge on

T ime (one minute intervals)

Fig. I-The mass spectrometric ion intensities for (a) mass number 32 (due to 0 160 16), (b) mass number 33 (due to 0 160 17

and HO I6016) , (c) mass number 34 (due to 0 160 18 and HH0160 16) show that the pronounced plateaus for mass 33 must have been due to a molecular fragment (H02) that is pro­duced only during the period when hydrogen atoms generated in an electric discharge subsequently react with oxygen. Much of the intensity background at mass 33 and 34 comes from the ever-present oxygen isotopes 0 16 0 17 and 0 160 18. Since no notable changes at mass 34 are observed with the discharge on, hydrogen peroxide was not formed in appreciable amounts . Consequently, its fragments could not have been responsible for any of the peaks at mass 33 .

January-March 1980

tory mechanism had been known. In a series of elegant experiments, 8 A. Westen berg and Newman de Haas measured free radical reactions at low pressure in simple systems in which the rates of ap­pearance or disappearance of these radicals could be determined with great precision by using elec­tron spin resonance techniques. The hydrogen atom/ oxygen molecule reaction was a seemingly straightforward pathway that, in the presence of carbon monoxide, should permit rate measurements of the latter with H02 • The results indicated a very fast reaction, quite in contrast to earlier, more in­direct results deduced from more conventional ex­plosion experiments.

In order to resolve this conflict, people elsewhere (particularly at the University of Maryland) carried out independent measurements or the H02 -CO reaction using isotope tracer techniques and higher pressures. 9 There is now a consensus that H02 is not as effective a scavenger of carbon monoxide as was believed by Westenberg, even though its role in reacting quickly with other atmospheric con­taminants such as nitric oxide and sulfur dioxide is important. H02 may have formed in Westenberg's experiment so as to produce a nonequilibrated species with unusually high energy content and great reactivity. If this is so, a new chapter of reac­tion research is opening where reactants are not in equilibrium with their surroundings. Such non­equilibrium conditions may prevail at high altitudes (low pressure) or in combustion situations where reactions occur at very high speed.

The detection of H02 validated a theoretical prediction of great subtlety. For two decades ex­perimenters were challenged to devise experiments that would provide convincing evidence. Foner and Hudson met the challenge.

WALTERG. BERL

REFERENCES

IN. N. Semenov, Chemical Kinetics and Chain Reactions, Oxford (1935).

28. Lewis and G. Von Elbe, Combustion, Flames, and Explosions of Gases, Cambridge (1938).

3S. N. Foner and R. L. Hudson, "Detection of the H02 Radical by Mass Spectrometry," J. Chern. Phys. 21, pp. 1608-1609 (1953).

4S. N. Foner and R. L. Hudson, "Mass Spectrometry of the H02 Free Radical," J. Chern. Phys. 36, pp. 2681-2688 (1962).

5A. C. Lloyd, "Evaluation and Estimated Kinetic Data for Phase Reac­tions of the Hydroperoxyl Radical," Int. J. Chern. Kin. 6, pp. 169-228 (1974).

6T. T. Pankert and H. J. Johnson, " Spectra and Kinetics of Hydroperoxy1 Free Radicals in the Gas Phase," J. Chern . Phys. 56, pp. 2824-2838 (1972).

7S. Saito, "Microwave Spectroscopy of Short-Lived Molecules," Proc. XXVIIth International Congress of Pure and Applied Chemistry: Vol. 2, Physical Chemistry, pp. 1239-1250 (1977).

8A. A. Westenberg and N. de Haas, "Steady State Intermediate Concen­tration and Rate Constants, Some H02 Results," J . Phys. Chern. 76, pp. 1586-1593 (1972) .

9D. D. Davis, W. A. Payne, and L. J . Stief, "The Hydroperoxyl Radical in Atmospheric Chemical Dynamics: Reaction with Carbon Monoxide," Science 179, pp. 280-282 (1973).

37