M. C. Ball
University of Technology Loughborough, Leicestershire England
Chemical Transport Reactions
In discussions of mass trans~ortin the gas phase, processes often considered are sublimation, distillation, and gaseous diffusion, usually in terms of the Clapeyron-Clausius equation, the Phase Rule, and the kinetic theory of gases, respectively. However, a large number of chemical reactions have been studied in which reversible heterogeneous reactions can lead to mass transfer through the gas phase. Schafer (1) has suggested the term "chemical transport" for these reactions in which a gas reacts with a solid to produce a gaseous reaction product which is itself thermally unstable. I n more general terms: where A has a negligible vapor pressure a t the reaction temperature. One of the best illustrations of this type of reaction has been used commercially for many years; this is the Mond-Langer process for the purification of nickel via the tetracarbonyl: Ni
+ 4C0 = Ni(CO)r
I n this process, the forward reaction is carried out a t 5OoC,and the reverse reaction a t 200°C (2). With such reversible reactions and with deliberate temperature gradients, it is possible to have the forward and backward reactions predominating in different parts of the system and hence the solid A will be transported. Under such conditions, the general reaction is composed of three processes, any or all of which may be rate dekrmining: the heterogeneous forward reaction A + B C at Tr the homogeneous transport from one temperature zone to the other (the gas motion) the heterogeneous reverse reaction C A B at Tn -t
-
+
The Gas Motion
At very low pressures the mean free path of any molecule becomes very large and molecule collisions will be infrequent. Under these conditions the intermediate species (C(g)) will reach the hot zone (TB) where a considerable proportion will be decomposed. The amount of A transported is therefore proportional to the surface area of the hot zone. As the pressure increases, the mean free path decreases and gaseous diffusion becomes rate determining; the amount of A transported is practically independent of the surface area of the hot zone. This diffusion control of the transport reaction has a pressure range from about to 3 atmospheres, i.e., over most experimental pressures, but these limits are also dependent to some extent on the design of the apparatus. At pressures greater than about 3 atmospheres, diffusion control of the reaction breaks down and thermal convection must also be taken into account.
The Reaction Rate
I n most reactions studied, Conditions are deliberately made such that the heterogeneous reactions proceed so rapidly that equilibrium is established at the solid material, and the transport observed agrees with that calculated (see next section), assuming that gaseous step, diffusionis the rate Only a few reactions have been found where the heterogeneous reaction is the rate determining step. One of the most important of these is the Boudouard reaction
in which the experimental transport is about 100 times smaller than that calculated for a given temperature gradient (I). The very high bond energy of carbon monoxide, 256 kcals, can account for its low rate of decomposition and hence the low transport, of carbon. The formation of any thermally stable compound either in the gas phase as above, or as a solid reaction product can slow down the rate of transport. I n the well-known Van Arkel-De Boer iodide method for the purification of zirconium, as the temperature of the crude metal rises, the formation of pure zirconium goes through a maximum and then a minimum (8). This effecthas not been fully explained, but it has been suggested that a t the Lower temperature, volatile ZrL is formed and that its concentration in the gas phase a t first increases with temperature giving rise to transport. As the temperature rises, the more stable lower iodides of Zr are formed. These solids cover the metal surface, a t the same time removing the transporting agent from the gas phase and reducing the amount of transport. When a temperature is reached a t which these lower iodides disproportionate, 2ZrI.
-
Zr
+ ZrL
the iodide content of the gas phase- increases with a consequent increase in transport. Thermodynamic ~onsiderkions
I n addition to the transport of gas molecules already mentioned, another prerequisite is that the partial pressure of the gascous reaction product a t the starting solid should differ from that a t the final solid phase. The number of moles of A transported, N A , is given by the following relationship (I,4) : i Dqt P , N*= 7 -
3
1RT
where i and j are the coeEcients of the reaction eqn. (I), PCis the difference in partial pressures, q(cm2)and l(cm) are the cross-section and length of the diffusion zone, t is the reaction time, and R and T are the gas ~ o l u m e45, Number 10, October 1968
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651
constant and absolute temperature. D is the diffusion C under coeficient in cm2/sec for the gas mixture B the experimental temperature and pressure conditions. Values of the diffusion coefficients for a large number of gas mixtures are known; a t 273°K and atmospheric pressure D has an average value of 0.1 cm2/sec in hydrogen-free systems. This value increases with increasing temperature and pressure (5). For maximum transport therefore, PC must be a maximum, and this will happen when the equilibrium constants for both the forward and backward reactions are large'.
+
Forward reaction A
+B = C
Backward reaction C
A
+B
KB
=
Ke =
[Cl at T F [A1 LBI [A1 P I at T B [CI
Because the forward and backward reactions are deliberately carried out a t different temperatures, K p will not equal K,. For transport to reach a maximum, then KF/KB 1, i.e., In Kp/KB must he positive. If Kp/KB < 1then more gaseous product is formed than is decomposed, and the transport of solid material is decreased. Now In Kp = (-AH/RTp) - (ASIR) and similarly for in Kg.
>
If AH is negative, -AH/R is positive, and for h K p / KB to he also positive then TF < TB. This means that transport occurs towards the higher temperature zone. When AH is positive, then -AH/R is negative, and Tp > TB to make the whole expression positive. Therefore transport takes place towards the lower temperature zone. This simple explanation of the effect of AH on transport direction ignores both the magnitude of AH and also the effects of entropy changes during the reaction. These effects can he summarized, however: For transport to occnr, AH must not he too large (-50 kcals4). 2. The sign of AH governs the transport direction. Exolhermic reactions transport to the bigher temperature aone, and endothermic reactions t,o the lower tempemtnre zone. When A H = 0, no transport can take place. 3. When AS = 0, there will be a temperat,ure difference het,ween bhe hot and cold zones which gives maximum transport. 4. When AS is large, then AS and AH should have the same sign. 1.
For a much fuller account of these effects in experimental systems, see the recent paper by Alcocli and Jcffes. (6) Experimental Mefhods
A4ost reactions are controlled by gaseous diffusion, and for these, closed reaction vessels must be used. Suitable materials are glass or quartz tubing, depending upon the final reaction temperat,ure; convenient dimensions for the reaction tuhe are 10 X 150-200 mms. For preparative reactions, the loading of the tuhe is very simple: the solid mat,erial t,o be transported is placcd in a tuhe closed a t one end and the transporting agent,, if solid or liquid, is added and a constriction made in the tuhe a t an appropriate point. The tube is then evacuated (mm Hg) cooling the contents if 652
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Journal o f Chemical Education
necessary, and then sealed off a t the constriction. The quantity of transporting agent added should generate a pressure between a few mm Hg and 2-3 atmospheres a t the reaction temperature. A common transporting agent, iodine, added in quantities of 0.5 -5.0 mg/cm3 of tube volume, will generatc pressures of 0.16-1.6 atmospheres a t 700°C. If the transporting agent is gaseous a t room temperature, the reaction tube is filled with gas a t pressures up to atmospheric before sealing off. If gas motion by convection can he added to diffusion during the reaction, transport can he greatly increased. This is helped by the use of wider reaction tubes (up to 30 mm) and tilting the reaction tube in the temperature gradient so that the bottom of the tuhe is a t the higher temperature. Gas motion can also he achieved by passing a current of reacting gas over the reactant solid; this is particularly useful where the gas is commercially available from cylinders, e.g., 02, CL, HCl, Hz. The solid can he loosely packed into the reaction tuhe over a distance of a few centimeters and heated, while another part of the tube is maintained a t a different temperature to reverse the reaction. The heating arrangements can be equally simple: ordinary tribe furnaces are most commonly used, advantage being taken of their natural temperature gradients. A silica or aluminia tube, 1-in. diameter, wound with two consecutive heating coils, each with its own voltage regulator, is very suitable for most purposes. Two very simple laboratory experiments can be used to demonstrate the phenomenon of transport (1): A nickel mirror can be produced from granular nickel by means of the following reactions
Copper foil exposed to a normal 1a.horatory atmosphere will always have, in addition to cuprous oxide, some halide on the surface. This can he demonstrated by heating in a flame, when the typical green color is obtained. If pieces of such foil are heated in evscuated tubes so t,hst the metal is at %bout900PC, then coprous oxide is transported according to the equation
The cuprous oxide can occasionally be deposited crystals large enoagh to be seen with bhe naked eye.
8.8
reddish
Applications
A very large number of transport reactions are known, leading to transport of elements, oxides, sulfides, selenides, tellurides, halides, oxyhalides, arsenides, phosphides, and mixed oxides. Transporting agents include oxygen, hydrogen, water, carbon monoxide, halogens, and many volatile halides. Although some of these reactions have been known for many years, it is only recently that the potential of this type of reaction has been appreciated. Important applications are: purification of materials; synthesis of compounds and enhancement of solid state reactions; crystal growt,h; and determination of thermodynamic properties. These applications will now he covered in . . more detail. Purification. Two methods for the purification of metals have already been mentioned, viz., the RIond~
~
Langer and the Van Arkel-De Boer processes. The latter method, using iodine as transporting agent, has been used to purify Ti, Zr, Hf, V, Nh, Ta, Ni, Cu, Fe, Cr, U, and Si. These reactions are invariably exothermic and therefore transport to the hot wire. Other halogens may be used but iodine is preferred because of the lower thermal stabilities of the iodides. The iodide method is extremely useful in that the interstitial impurities which are normally troublesome (carbon, nitrogen) are not transported. However metallic impurities might easily be transported along with the hulk metal. These transport reactions are the most closely studied and further reference should he made to the standard work (7). An interesting small scale application of the Van Arkel-De Boer method is in the so-called "quartziodine" lamp (8). Normal tungsten filament lamps are run a t temperatures of about 2000°C and if the temperature is raised in an attempt to increase the illumination, the filament breaks because tungsten evaporates onto the cooler parts of the bulb. If a small amount of iodine is int,roduced into the bulb, then this reacts with t.he evaporated tungsten, which is a t the lower temperature, with the formation of volatile iodides which transport to the hot filament and decompose. Thus t,he evaporation is compensated for by the transport reaction, and the filament can be run a t temperatures up to 2SO0°C. Under optimum conditions an increase of 60% in luminous efficiency can be achieved. Synthesis. I t is in the field of synthesis that most developments are taking place. No attempt has been made to cover all the examples; those given are of particular interest to the author. The following reactions proceed much faster when reducing agents (Hz,CO) are present:
pared similarly by altering the composition of the starting mixture since the oxyiodide NbOIs has a second mode of decomposition as follows: 2NbO18 e NbOa
+ NbIs + '/&
The net enthalpy change for the series of reactions involving Nh02 must he nearly zero, since hTb02does not transport in a temperature gradient. An increasingly important use of transport reactions is in the formation of mixed oxides, many of which are ferroelectric. These can he made from mixtures of the constituent oxides with hydrogen chloride as the transporting agent: MO
MCls
+ 2HC1 e MCll + H 2 0
+ 6HCl + ZFeCI, + 3H,O Fe%08 + 2FeCI, + 4H%0= MFelOl + RHCl
Manganese, cobalt, and nickel ferrites have been prepared in this way, as have several tungstates, yttrium iron garnet, and lanthanum niobate (LaNh04) (12). For a full account of the many other synthetic applications, see the main reference (1). Crystal Growth. Single crystals are normally grown from solution, from various melts, or by sublimation. Frequently, however, these methods cannot he used, either because the material is insoluble or unstable below its melting point. I n such cases transport reactions are useful, because lower temperatures can often be used, and better crystals result. Many very large crystals have been grown by transport methods, e.g., CdS single crystals up to 50 g in weight (IS). Transport has also proved very useful in the preparation of single crystals of many oxychlorides which are normally amorphous or microcrystalline. Occasionally application of a temperature gradient 900'C leads to separation of the constituents rather than 2Ca0 + SnOl Ca1SnO4 crystal growth; attempts to grow single crystals of soooc i\fgWOa lead to transport of WO1 preferentially, leaving SrO + SnO. SrSnOa MgO as residue (12). Even when the reactants are not in contact, 8110% Determination of Thermodynamic Properties. It was migrates to the other oxide (9). This is explained by shown earlier that the direction and extent of transport the formation of gaseous SnO and its regeneration accould be calculated from a knowledge of the thermocording to the following equations: dynamics of the reaction. Conversely, it is possible to obtain information on the thermodynamic properties from studies on chemical transport. If both exo- and endothermic equilibria are involved The usual method for the preparation of lower oxides in the transport of the same solid material, then in a of metals is to heat the metal plus a higher oxide, in the temperature gradient, transport may take place to a correct proportions, to a high temperature (lo),e.g., higher temperature, but with a change in conditions 17OO0C also to a low temperature. This can happen where NblOj + 3Nb -+ 5Nb0 more than one compound exists in the gas phase. This Most oxides, however, have appreciable vapor presinversion point in the transport direction (the critical sures a t high temperatures, leading to volatilization and decomposition point) has been used to determine alteration in the stoichiometry of the mixture. Other equilibrium constants, particularly where the hulk oxides are unstable, producing oxygen and a lower oxide, composition of the gas phase can be fixed. As an or the metal. example, the determination of the enthalpy of formation The same mixture, in the presence of iodine, will of carbon monosulfide will he discussed. produce NhO a t 900°C, as shown in the following If a carbon filament is heated a t 1500°C in an atmoequations: sphere of CS2 a t a measurable pressure, then the following reactions take place:
-
+ Sdg) CS2(g) C + C&(g) = 2CS(g) C
This synthesis is best carried out over a temperature gradient from 900 to llOO°C (11). NbOz can be pre-
exotherrnio endothermic
and the amount of carbon taken reversibly into the gas phase, P , = I/, P , - Ps,. Volume 45, Number 10, October 1968
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653
Now the total pressure and temperature are known, and thc equilibrium pressure of S2 in the exothermic reaction can be calculated from known thermodynamic data. Therefore Pcs is also known. Since all the other values are known, it is possible to calculate the enthalpy of formation of carbon monosulfide (14) : AH(CS(g)) = 58
* 5 keals
An independent value (1) derived from measurement of the equilibrium for the reaction: gave AH (CS(g)) = 55.0 * 0.5 l~cals. The transport mcthod does not give data of the highest accuracy, but the experiments are rapid and fairly simple to carry out. Literature Cited (1) SCHAFER,H., "Chemical Transport Reactions" Academic Press, Inc., New Ynrk, 1964. (2) DENNIS,W. H., "Aletallorgy of Non-ferrous Metals" (Snd Ed.) Pitman. Londm 1961.. . D. 379. (3) MILLER, G. L., i i Z i r ~ ~ n i ~ r(2nd n " Ed.) Bulterworths,
654
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Journal of Chemical Education
London, 1957, p. 79. (4) SCHAFER, H., J-~COII, H., A N D ETZEL,K., Z. amrg. Chem., 286, 27 and 42 (1956). 151 T.. "The Pruoerties of . . See REID. R.. AND SHERWOOD. Gases and ~iquids,"M c ~ r a w - ~ i l 1 ' ~ oCo., o k ~ e York, w 1958, Ch. 8; "Internationsl Critical Tables" Vol. 5, p. 62, McGrew-Hill Book Co., Now York, 1929. (6) ALCOCK, C. B., A N D JEFFES, J. H. E., Trans. Inst. M . M., 76, C246 (1967). (7) ROLSTEN,R. F., "Iodide Metals and lMetal Iodides" John Wiley & Sons, Inc., New York, 1961. (8) VAN TIYEN, J. W., Philips Technical Review, 23, 237 (1961/2). (9) SPANDAU, H., A N D KORLMEYER, E. J., Z. anorg. Chem., 254, fii .. 11947)~ \---.,(10) BRAUER,G., "Handbook of Prepmrttive Inorganic Chemistry", Vol. 2 (2nd Ed.) Academic Press, Inc., New York, 1965. H., A N D HUESKER, M., Z. anovg. Chem., 317, 321 ' (11) SCH~FER, (1962). (12) CURTIS, B. J., ANT) WILEINSON? J. A,, J. Amer. Cemm. Soc., 4 8 . 4 9 (1965). (13) REY&OL&, I).'c., .ANDGREENE,L. C.,J. Appl. Phyr., 29, 559 ...1195XI \-., .-,(14) SCHAFER,H., AND WIEDEMEIER, H., Z. dnorg. Chem., 296, 241 (1958).
