THERMAL IGNITION OF LITHIUM NITRATE-MAGNESIUM
July, 1956
867
the oxygen-phosphorus bonds, rendering them ‘more TABLE I11 THERMAL DECOMPOSITION OF DIHYDROGEN ORTHOPHOS- liable to thermal rupture and also that the lighter PHATES
Reactant
Reaction rate appreciable Cation at temp., radius,o A. OC.
Principal initial product b
atoms have a greater amplitude of thermal vibration leading to decomposition a t lower temperatures.
TABLE IV LiHzPO4 0.68 189 Li~HzPz07 THERMAL DECOMPOSITION OF MONOHYDROGEN ORTHOPHOSNaHZP04 0.97 200 Na2H~Pa07 PHATES KH~POI 1.33 208 [KzHzPz071 Reaction rate CsHzPOi 1.67 233 [CSZHZPZO~] Cation appreciable a t Reactant radius, A. temp., ‘C. Product 208 [PbHzPzOrI Pb O L J “ d 2 1.20 NazHP04 0.97 240 Na4P~07 1.34 243 [BaHzP2071 Ba(HzP04)z KzHPOa 1.33 282 K4Pz07 a From L. H. Ahrens, Geochzm. Cosmochim. Acta, 2, 155 CszHP04 1.67 339 CS4p207 (1952). Brackets indicate the product undergoes further decomposition a t the temperature listed.
contribution to the lattice structure such that quite different arrangements of atoms can have much the same stability and lattice free energy.”l’ A trend is apparent toward lower reaction temperatures with the smaller cations. This is the usual trend for the thermal decomposition of a series of salts of an oxy-acid. It is likely here that polarization of the oxygen atoms by the smaller cations weakens the oxygen-hydrogen bonds and (17) A. R. Ubbelohde and I. Woodward, Pvoe. R o g . Soe. (London), 8179, 399 (1942).
The degree of hydrogen bonding in the monohydrogen orthophosphates, from purely stoichiometric considerations, must be less than in the dihydrogen orthophosphates. I n the monohydrogen orthophosphates hydrogen bonding can then be expected to make a smaller contribution to the crystal properties. For a smaller range of cations, a greater range of decomposition temperatures is noted (Table IV). The trend in reaction temperature with cation size is the same as for the dihydrogen orthophosphates and for salts of oxyacids in general.
THE APPLICATION OF THE ABSOLUTE RATE THEORY TO THE lGNITION OF PROPAGATIVELY REACTING SYSTEMS. THE THERMAL IGNITION O F THE SYSTEMS LITHIUM XITRATE-MAGNESIUM, SODIUM NITRATEMAGNESIUM1 BY ELIS. FREEMAN AND SAULGORDON Pyrotechnics Chemical Research Laboratory, Picatinny Arsenal, Dover, N . J . Received September $1, 1966
Upon heating intimate mixtures of finely divided alkali nitrates with powdered magnesium a t temperatures greater than 480°, ignition ensues, accompanied by a highly exothermal self-propagating reaction. I n this investigation, the thermal ignition of the binary systems, lithium nitrate-magnesium and sodium nitrate-magnesium, was studied by measuring the time to ignition as a function of temperature. The experimental data indicate that, to a first approximation, the duration of the induction period prior to ignition, depends upon the accumulation of a specific minimum amount of heat. By the application of the technique of differential thermal analysis, it was shown that the increase in temperature during the induction or pre-ignition stage of the reaction is relatively small. The data appear to support a proposed theory for self-propagating exothermal reactions, expressed in terms of the absolute rate theory. The derived expressions relate the time to ignition, to the specific rate free energy, entropy and energy of activation of the pre-ignition reaction. The calculated and observed ignition times are in relatively good agreement.
Introduction The concept of pre-ignition reactions was formulated during the latter part of the 19th century.2 Semenov3 expressed the relationship between the time for ignition and the temperature of the system as t =
AfeEIRT
(definition of symbols in Appendix). The form of this equation is followed by a large number of (1) This paper was presented in part, before the Division of Physical and Inorganic Chemistry at the National Meeting of the American Chemical Society in New York, Sept., 1954. (2) Van’t Hob, “fitudes de Dynamique Chimique,” 1884. (3) N. Semenov, “Chemical Kinetics and Chain Reactions,” Oxford, Ch. 18, 1935.
propagatively reacting systems, such as explosive^,^ propellants and pyrotechnic compositions. The validity of applying this empirical relationship appears to be independently corroborated by the derivation presented in this paper, which is based upon the application of modern rate theory. The derivation also gives a physical significance to the coefficient, A , which is not apparent in the Semenov relationship. Eyring and Zwolinski5 used the absolute rate theory to predict the minimum ignition temperature of magnesium in oxygen, but did not consider the time to ignition. (4) H.Henkin a n d R. McGill, Ind. Eng. Chem., 44, 1391 (1962). (5) H. Eyring and B. J. Zwolinski, Ree. Chem. Pmg., 8,87 (1947).
ELIS. FREEMAN AND SAULGORDON
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Due to the importance of this factor, expressions Experimental have been derived relating the time to ignition to The atomized (spherical) magnesium powder, 60/80 the specific rate, the energy of activation, the free U. S. Standard Sieve fraction, specific surface 176 cm.2/g., energy of activation and the entropy of activation of was obtained from the Golwynne Chemical Corporation. the preignition reaction. The systems lithium The specific surface was measured by the air permeability method using the Picatinny Arsenal Particle Size Ap aratus.6 nitrate-magnesium and sodium nitrate-magnesium The lithium and sodium nitrate were purchased g o m the were studied by measuring the time to ignition as a Fisher Chemical Company and are of C.P. grade. Mixtures function of temperature. The theoretical and ex- of 200 mg. of ground sodium nitrate and 100 mg. of magnesium and 250 mg. of lithium nitrate and 25 mg. of magneperimental times t o ignition are presented.
80 '
6
70
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sium were weighed, placed in small copper blasting caps (Special, Type 2), No. 33 B & S gauge, 5 mm. in diameter and 3 cm. long, and kept in a desiccator prior to use. The blasting caps containing the composition were plunged into a molten lead bath maintained at the desired temperature to f 2 O . 7 The moment the sample entered the bath, a circuit was closed by means of a microswitch, automatically starting a timer. At the first appearance of light resulting from the reaction, the clock was stopped manually. The over-all timing error for these operations is estimated to be 3 ~ 0 . 2second, and is negligible with respect to the measured times. By determining the time re ulred to melt barium nitrate crystals, it was found that wxen the bath temperature is 600", it takes 0.2 second for the crystals a t the inner surface of the blasting cap to attain a temperature of 590". The equipment used for the differential thermal analyses (DTA) of the compositions, has been previously described.* The differential temperatures of the samples were recorded with respect to an inert substance, aluminum oxide, as the temperature of the furnace was increased a t a rate of ap roximately 15" per minute. The compositions studied by 6 T A were, 5 g. of lithium nitrate-1 g. of magnesium and 5 g. of sodium nitrate-l g. of magnesium.
8
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Results and Discussion From graphs of the time to ignition vs. temperatures, Figs. 1 and 2, it is seen that the minimum ignition temperatures of the lithium nitrate-magneTemp., "C. Fig. 1.-Time to ignition of mixture of 250 mg. of lithium sium and sodium nitrate-magnesium systems are nitrate and 25 mg. of magnesium as a function of tem- 480 and 570°,respectively. The largest dispersion perature. of points occurs a t these values, where changes in temperature as small as 1 to 2' result in either relatively large changes in ignition time or ignition failure. This illustrates the sensitivity of the sys36 n tems to the heat balance boundary condition which, when exceeded, results in ignition. The condition 32 is that the rate of heat produced by the pre-ignition reaction equals the rate of heat dissipation. n A considerable amount of information concern28 ing the thermal behavior of the systems during the pre-ignition reaction may be obtained from the differential thermal analyses (see Figs. 3 and 4). 24 Figure 3, represents the reaction of sodium nitrate and magnesium. At approximately 280" an endoE 20 I is observed, which is attributed to the rotatherm B tion of the nitrate ion in the crystal followed by a 0 lattic expansion.9 The endotherm, at 315", results from the melting of the sodium nitrate crys8 tals, after which a new reference level is attained c; due t o the greater thermal conductivity of the 12 melt with respect to the solid.8 No other prominent thermal effects are noted until a temperature 8 of 580" is attained, at which point there is a sharp increase in temperature, due to a highly exothermal self-propagating reaction which is accompanied by 4 ignition. Figure 4 shows a similar curve for lith10
a
8
s
.e 42
4,.
590 600 610 620 630 Temp., "C. Fig. Z-Time to ignition of mixture of 200 mg. of sodium nitrate and 100 mg. of magnesium as a function of temperature.
570
580
(6) B. Dubrow and M. Nieradka, Anal. Chem., 21, 302 (1955). (7) S. Gordon and C. Campbell, "Fifth Symposium (International) on Combustion," Reinhold Publ. Corp., New York, N. Y., 1954, p. 277. (8) 6. Gordon and C. Campbell, A n d . Chen., 21, 1102 (1955). (9) S. Glasstone, "Textbook of Physical Chemistry." D. Van Nostrand Co., Inc., New York, N. y., 1946, p. 423.
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July, 195G
THERMAL IGR'ITION OF LITHIUMNITRATE-MAGNESIUM
ium nitrate-magnesium. An endotherm representing a fusion is observed a t approximately 268O, and a t 457" there is a rapid increase in temperature, accompanied by ignition. For both compositions there is only a small increase in temperature (about 3") prior t o the rapid exothermal reaction, a t which point, within several tenths of a second, ignition is observed. Considering the time to ignition as a function of temperature, where the induction period was found to vary from 4 t o 95 seconds, it appears that a relatively long, slow pre-ignition reaction occurs prior to ignition (see Figs. 1 and 2). The thermal energy accumulated by the system as the pre-ignition reaction proceeds results in a small increase in temperature. During this induction period, the rate of the pre-ignition reaction increases to a relatively small extent with the increase in temperature. As the reaction proceeds a critical region is reached in which further small changes in the temperature produce relatively large changes in the reaction rate, which is accompanied by a corresponding increase in the rate of hea.t evolution. At this point, the temperature and the rate of reaction accelerate rapidly and ignition ensues. This behavior may be expected since the rate of reaction is an exponential function of temperature. A possible explanation for the relatively small increase in temperature just prior to the rapid reaction leading to ignition may be that the heat produced by the pre-ignition reaction goes into raising the surface of the metal particles t o the activated state, and that the rate of formation of activated molecules is more rapid than the rate a t which the activated complexes decompose to the final products. On this basis the net amount of heat evolved may be expected to be relatively small. However, since the rate of decomposition is accelerated as the concentration of activated molecules increases, it is reasonable to assume that a time will be attained when the rate of decomposition of the activated molecules t o products will be comparable to that of the rate of formation of the activated state. This results in a relatively rapid evolution of heat, and an increase in surface temperature. In order to have a self-propagating reaction the reaction rate must be rapid enough so that the net amount of heat gained due to the decomposition of an activated complex and the thermal conductivity of the composition is greater than that required to activate an adjacent molecule. Therefore, the duration of the induction period depends upon the time required to attain this condition. Since there is only a relatively small increase in temperature during the induction period, it is assumed in the following derivation that the pre-ignition reaction occurs at a constant temperature. Consider the general pre-ignition reaction aA
+ bB = [ a A . . .bB]*+products
(1)
The rate of reaction is given by the expression
(see Appendix for definition of symbols). By the appropriate substitutions'o in equation 2,
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Fig. 3.-Differential thermal antlysis: marks on curve appear a t 1 min. intervals; temp. C. indicated a t timing marks. Li NO,+ Mg (ATOMIZED 1
0' 10 N
Fig. 4.-Differential thermal anaiysis: marks on curve appear a t 1 min. intervals; temp. C., indicated a t timing marks; SN slight NO2 fumes.
equation 3 is obtained Multiplication of the rate of reaction and the heat of reaction gives the amount of heat evolved per unit time
2
X AH, = dH
(4)
Substituting equation 3 in equation 4
The net rate of heat gain in the system dQH/dt is the difference between the rate of heat produced, as a result of the pre-ignition reaction, dH/ dt, and the rate of heat dissipation dq/dt. Separating the variables in equation 5 and integrating from time equals zero to the time for ignition, equation 7 results
(10) S. Glasstone, K. J. Laidler and R. Eyring, "The Theory of Rate Processes," McGraw-Hill Book Co., New York, N. Y.,1941, p. 187.
ELI S. FREEMAN AND SAUL GORDON
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where H is the total heat produced by the pre-ignition reaction in the time interval zero to t. It is assumed, for reasons previously mentioned, that ignition occurs, when
heat of activation may be calculated from the slope of the line. A plot of log t us. T-' shows this linear relationship for the sodium nitrate-magnesium system. The heat of activation, which is approximately equal to QH = H' (8) the activation energy, was calculated to be 32 kcal. From equation 6 it may be seen that per mole. A linear relationship is also observed for the lithium nitrate-magnesium system and the QH=H-q (9) heat of activation was calculated to be 34 kcal. per The condition for ignition, then, is mole. If H - q 2 H ' ; ignition will occur (10) Using equation 13, theoretical values for the If H - q < H'; ignition will not occur (11) time to ignition were calculated and compared with The total heat produced by the pre-ignition re- the observed results (see Table I). The values of q) are approximately equal t o the heats of action must be greater than H' by the amount of (H' activation, 32 kcal. per mole and 34 kcal. per mole heat dissipated for the sodium nitrate-magnesium and lithium H = H ' + q (12) nitrate-magnesium systems, respectively. The Transposing and solving for time in equation 7 transmission coefficient, K , was taken to be unity. Since the oxidant is in the molten state, the activity of the magnesium surface was taken as the number of moles of oxidant on the surface of a particle per mole of magnesium (bulk), and were calculated t o be 3.94 X and 5.91 X 10-7 for the sodium nitrate-magnesium and lithium nitrateSubstituting in equation 13 magnesium systems. The activity of the sodium nitrate, which is in excess, is assumed to be unity. The calculations of the heats of reaction, AH,, a t 2.5" were based upon the reaction between metal Let and oxidants to form magnesium oxide and sodium B = HI+ q or lithium nitrite. The values are 125 kcal. per A H , a: a; mole magnesium or oxidant and 117 kcal. per mole t = B/k. (17) magnesium or oxidant for the reactions between Substituting the expression for the equilibrium sodium nitrate and magnesium and lithium nitrate and magnesium, respectively. Since the difference constant between the activated and normal states of a molecule is the conversion of a vibrational degree of 'freedom to a translational degree of freedom, the between the activated and normal states in equation partition function term (&/&*) was calculated to be 13, equation 19 results approximately unity. The values for Planck's and Bolteman's constants are 6.62 X erg see. (H' q)h t = (19) and 1.38 X ergper OK. The calculated results KAH,~;~;~TK*,, are in good agreement with the observed data conK*., = e-AF*/RT (20) sidering the magnitude of the numbers and assumptions involved. substituting in equation 19
+
+
AF* = AH* - TAS*
(22)
substituting in equation 21
TABLE I COMPARISON BETWEEN CALCULATED AND OBSERVEDTIMES TO IGNITION Temp., System
OC.
NaNOa-Mg
570
Let LiNOgMg
Substituting in equation 24 t = AeAH*/RT
580 600 630 480 520
590 (25)
The coefficient A is a function of the activity of the reactants and the absolute temperature. If the activities of the reactants do not change significantly during the pre-ignition reaction, the logarithm of the time t o ignition should be a linear function of the reciprocal of the absolute temperature over relatively short temperature ranges. It is also apparent from the relationship that the
Time,aec. Obsd. Calcd.
8 6 4 2 282 83 13
(Range)
15-3G 11-14 7-9 4 45-95 32-41 3-5
The logarithms of the times t o ignition for mixtures varying in composition from 20 to 50% magnesium having specific surfaces from 176 to 650 cm.2/g. were also found to be linear functions of T-I. The ignition times were effected to a relatively small extent, considering the type of measurement, increasing by 10 to 30%, with increases in magnesiun concentration and specific surface. The slopes remained essentially unchanged.
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July, 1956
HYDROGEN OVERPOTENTIAL AT IRON CATHODES IN SODIUM HYDROXIDE SOLUTIOKS 871
Acknowledgment.-Acknowledgment is made to A. Tomichek for obtaining the ignition temperature data presented in this investigation. The authors wish to express their gratitude to D. Hart for reviewing this manuscript. Appendix A’ A
= constant = constant,
AE* AH* h H H‘
= energy of activation
a
=
activity
heat of activation Planck’s constant total heat produced by pre-ignition reaction = min. amount of heat required to raise the temp. of system to the point of ignition = = =
IC. k
= specific rate
Boltzmann’s constant transmission coefficient length of coordinate of decomposition mass of activated complex m* = partition function = gas constant AS* = entropy of activation T = absolute t>emp. t = time AHr = heat of reaction per mole magnesium dH/dt = rate of heat produced by pre-ignition reaction dQa/dt = net rate of heat gain in system dnldt = no. of moles of activated complex traveling over the potential energy barrier per unit length of the decomposition coordinate per unit time $q/dt = rate of heat dissipation = designates activated state
K 1
= = = =
i
THE EFFECT OF SOME CORROSION INHIBITORS AND ACTIVATORS ON THE HYDROGEN OVERPOTENTIAL AT Fe CATHODES I N NaOH SOLUTIONS BY I. A. AMMAR AND S. A. AWAD Chemistry Department, Faculty of Science, Cairo University, Cairo, Egypt Received October 10, 1966
Hydrogen overpotential, 9, a t Fe cathodes is measured in pure NaOH solutions at 25”. Measurements are also carried out in 0.2 N NaOH solutions to which the following organic substances have been added: ethylamine, n-butylamine, dimethylamine, tri-n-propylamine, benzylamine, pyridine, quinoline, picric acid, p-benzoquinone and m-dinitrobenzene. The numerical increase of 11upon addition of the organic substance is ex lained on the basis of a decrease in the available surface area for hydrogen discharge. On the other hand, the numerical &crease of 7 is robably caused by the depolarization of the hydrogen evolution reaction. An alternative explanation of the results is base: on the assumption that either an attraction or a repulsion, between the adsorbed atomic hydrogen and the organic molecules, take place.
Introduction The inhibition, by organic compounds, of the corrosion of iron in acid solutions was studied by Mann, Lauer and Hu1tin.l They observed that with mono-aliphatic amines, the efficiency of corrosion inhibition increased with the length of the aliphatic chain. Furthermore, the efficiency increased as the number of radicals attached to the nitrogen increased t o three. With aromatic amines the efficiency was dependent on the nature and the size of the inhibitor. Various theories were put forward to explain the action of organic inhibitors. Thus Chappell, Roetheli and McCarthy2 concluded that the inhibition of corrosion was caused by a resistive film on the surface of iron. Bockris and Conway3 measured the hydrogen overpotential, q, at an iron cathode, in the presence and absence of corrosion inhibitors and activators, in acid solutions. They observed a rise in q toward more negative values in the presence of inhibitors. In the presence of activators q was numerically decreased. Furthermore, they observed that the direct and the indirect methods of measuring overpotential gave the same results. From this they concluded that the inhibition and activation of corrosion were directly related t o the hydrogen activation overpotential (1) C. A. Mann, B. E. Lauer and C. T. Hultin, I n d . Eng. Chem., 28, 159 (1936).
(2) E. L. Chappell, B. E. Roetheli and B. Y. McCarthy, I n d . Erie. Chem., 20, 582 (1928). (3) J. O’M. Bockris and B. Conway, THIBJOURNAL, 63,527 (1949).
rather than to the formation of a surface film. Hackerman and Makrides4 explained the effect of inhibitors on the basis of general adsorption, ie., weak physical adsorption, electrostatic attraction between the inhibitor ion and the metal and chemisorption resulting in the formation of a dative bond between the inhibitor and the metal. The aim of the present investigation is to study the effect of corrosion inhibitors and activators on the hydrogen overpotential at iron cathodes in NaOH solutions. Experimental Hydrogen oirerpotential was measured in an electrolytic cell similar to that described by Bockris.6 The cell consisted of four compartments: the anode compartment, the cathode compartment, the hydrogen electrode compartment and the inhibitor (or activator) compartment. A sintered glass disc was inserted between the first two compartments to minimize the diffusion of gaseous anodic products toward the cathode. Diffusion of atmospheric oxygen into the cell was hindered by the use of water-sealed taps and ground glass joints for the construction of the cell. Electrical contact between the cathode and the reference hydrogen electrode was made through a Luggin capillary (internal diameter 1 mm.). Hydrogen, previously purified by passing it over hot copper (450’) and then over soda lime, was introduced into the cell and was divided between the four compartments. The iron electrode, in the form of a pure iron strips sealed to glass, was introduced into the cathode compart(4) N. Hackerman and A. C. Makrides, I n d . Eng. Chem., 46, 523 (1954). (5) J. O’M. Bockris, Chem. Revs., 43, 525 (1948). (6) A ”Hilger” pure iron strip, prepared by Hilger and Co., London,
England.
