The reaction of hydrogen chloride with liquid sodium and a sodium-barium alloy

J. inorg, nucl. Chem., 1970, Vol. 32, pp. 1443 to 1454.

THE

Pergamon Press.

Printed in Great Britain

REACTION OF HYDROGEN CHLORIDE WITH LIQUID SODIUM AND A SODIUM-BARIUM ALLOY* M. R. H O B D E L L and L. N E W M A N Brookhaven National Laboratory, Upton. N.Y. 11973 (Received 18 September 1969)

A b s t r a c t - T h e reaction of hydrogen chloride with ligand sodium and sodium containing 4 At% dissolved barium was studied in the temperature range 100-250°C. The reaction took place in two distinct steps: (1) HCI---> ½H2+ CI- and (2)½H2--> H-, with only the first step involving hydrogen chloride. The effect of barium was to decrease the rate of the hydrogen chloride reaction, step (1), but to increase that of hydrogen, step (2). However, in both cases molecular hydrogen was produced quantitatively in the gas phase, indicating the absence of net hydrogen uptake in the presence of unreacted hydrogen chloride. The initial reaction of hydrogen chloride with pure sodium exhibited autocatalytic behavior. A reaction scheme was suggested in which the rate controlling process would be the growth of crystalline sodium chloride on the metal surface. In the presence of barium, however, barium chloride was the predominant product, and the hydrogen chloride reaction rate was found to decrease with time. The rate controlling process in this case was considered to involve diffusion of hydrogen chloride molecules through pores created in an otherwise protective barium chloride layer. Thus in each case, the overall reaction of hydrogen chloride was controlled by processes involving the solid reaction products. In contrast, under the conditions employed, the kinetics of the hydrogen reaction reflects the nature of the reactants. INTRODUCTION

THE USE of liquid sodium as a coolant in fast breeder reactors has created an interest in sodium chemistry. One important class of reactions exhibited by sodium is the displacement of hydrogen from compounds in which the hydrogen atom 8+ is bonded to a more electronegative atom or group: N a + H-X--> [H] + NaX. Hydrogen chloride is such a compound and this paper describes its reaction with excess liquid sodium in the temperature range 100-200°C. Under these conditions thermal dissociation of sodium hydride can be neglected[l] and the overall reaction will be: 2 N a + HC1 --> N a i l + NaCI. No previous kinetic study of this reaction has been reported, but in other recently studied reactions of the same type, gaseous hydrogen was formed as an intermediate[2, 3]. It is not immediately obvious why hydrogen should be produced in the presence of excess sodium and under conditions where sodium hydride is stable. One reason for reacting hydrogen chloride with liquid sodium, therefore, was that it might provide a convenient means of studying hydrogen displacement and molecular hydrogen formation in the presence of excess metal. *This work was performed under the auspices of the United States Atomic Energy Commission. 1. C. C. Addison, R. J. Puiham and R. J. Roy,J. chem. Soc. 4895 (1964). 2. C. C. Addison, M. R. Hobdell and R. J. Pulham, The Alkali Metals, Spec. Pubis chem. Soc. No. 22, p. 270 (1967). 3. C. C. Addison andJ. A. Manning,J. chem. Soc. 4887 (1964). 1443

1444

M . R . H O B D E L L and L. N E W M A N

This reaction has also been considered from the standpoint of liquid metal reactivity. It has recently been shown, for example, that dissolved barium stimulates sodium reactivity towards such compounds as nitrogen[4], hydrogen[5], and ethylene [6]. In contrast, other work [6] would suggest that reactivity towards acetylene is diminished. The latter reaction is complex, but the first step involves 8+ dissociation of the polar ~ C - - H bond, which is formally analogous to the C1-s+ H bond. It was therefore of interest to consider the effect of barium in the case of hydrogen chloride. The reaction of a gas with a liquid metal is a heterogeneous process, and the initial step will involve adsorption of gas molecules at the metal surface. As the reaction proceeds its overall rate may feasibly depend on (1) the chemical nature of the reactants, or (2) the rate at which reactive species are transported through a product surface layer. Should the physical process (2) prevail, the nature of the reaction will depend largely on the properties of the reaction product. A study of the kinetic aspects of liquid metal chemistry requires a situation where this is not the case. In the experiments described, a fixed volume of hydrogen chloride was reacted with a large excess of vigorously stirred liquid sodium, or sodium-barium alloy. The course of the reaction was followed by measuring pressure change as a function of time. EXPERIMENTAL Apparatus The apparatus used was similar to that previously described[2, 6], and is shown in Fig. 1. The reaction vessel, constructed from pyrex glass, consisted of two concentric cylindrical vessels, A and B, which could be independently isolated. The former was an inner reservoir for HCI (Matheson Co., GAS 4

~ T O

SUPPLY

IANOMETER

I

Fig. 1. Apparatus for reaction of gases with excess liquid sodium. 4. C. C. Addison, B. M. Davies, R. J. Pulham and D. P. Wallace, The Alkali Metals, Spec. Pubis chem. Soc. No. 22, p. 290 (1967). 5. M. R. Hobdell and L. Newman,J. inorg, nucl. Chem. 31, 1843 (1969). 6. M. R. Hobdeli, Unpublished work, Ph.D. Thesis, Nottinghman, England (1967).

Liquid Na- and Na/Ba-HCI reaction

1445

Inc. electronic grade purity 99-99 per cent) while the latter contained the sodium. The entire unit could be connected to a simple mercury manometer (of negligible volume compared to B) and a vacuum line. The vessel was heated by immersion in a silicone oil bath, whose temperature was carefully controlled (_ I°C). When molten, the sodium could be vigorously stirred over an internal baffle by means of the external rotating magnet C.

Sodium filling The sodium used was of high purity, containing < 10 ppm oxygen, and was stored in a steel pot (600 ml) prior to use. When sodium was required, this was connected to the reaction vessel, a vacuum line, and an argon supply. A thin steel delivery tube from the storage pot passed through the stopcock barrel of B, which together with the reaction vessel and sodium pot was heated to ~ 120°C by means of electrical heating tapes. After evacuation, a slight argon pressure was applied to the bulk sodium forcing it through the delivery tube into B. When sufficient sodium had been added (estimated by visual observation), the argon pressure was removed, and the sodium supply isolated. The reaction vessel was then "back-filled" with pure argon before disconnecting from the filling unit.

Procedure The important features of the technique employed were as follows: ( 1) The sodium surface was observed throughout the reaction. (2) Pure HCI from A was allowed into B (after previous evacuation) then A quickly isolated. The reactants were at the same temperature when brought into contact, and pressure changes due to thermal expansion eliminated. The use of capillary tubing outside B reduced the external gas which was at room temperature, to a negligible volume. (3) All HCI in B was available for reaction and the hydrogen produced was readily accessible to the sodium. (4) The induced "eddy" currents set up in the liquid metal by the rotating magnet caused vigorous stirring over the internal baffle; agitation of the metal surface was expected to prevent the formation of protective product layers and to continuously expose fresh metal. (5) In experiments involving barium, the metal was filed clean in an argon-filled dry box and added to the reaction vessel prior to sodium filling. in a typical experiment, the volume B (Fig. 1) was - 140 ml and when - 40 g of sodium were used there remained a fixed volume of reacting gas of ~ 100 ml. Pressure values were measured to the nearest 0.5 mm Hg. RESULTS

Pure sodium

The variation of pressure, P, with time for successive additions of HC1 to the same sodium sample at 160°C is shown in Fig. 2(a). Each pressure-time curve exhibited a sharp discontinuity at a pressure P', after which the pressure fell slowly to zero (not shown). A golden film, darkening with time, was formed on the metal surface on contact with HCI. Despite vigorous stirring, this film remained continuous in each case as the pressure fell from its initial value po to P'. At the stage corresponding to P' the surface product layer became mobile and would break up exposing clear metal. Continued stirring would cause the surface products to coagulate behind the baffle, thus exposing a similar metal surface area prior to each HCI reaction. The reaction of the first addition of HCI (Curve 1, Fig. 2(a)) exhibited an initial slow stage after which the reaction rate steadily increased (to P' ); this behavior was found to a slight extent with the second addition also (Curve 2, Fig. 2(a)). However, the third addition at 160°C (Curve 3, Fig. 2(a)) gave rise to an initially rapid but constant reaction rate. Thereafter, successive HC1 reactions

1446

M.R.

PI

I

I

I

H O B D E L L and L. N E W M A N

I

I

I

160

80

p 120 E E

A 60

A

"I-

I

I

I

I

2

-

I,LI (r

(/1 hi a.

_

P'

-

p ~ : ~

4C

2O

I 5

I I I I 0 15 20 TIME(rains)

I 25

0

0

5

I0 15 20 TIME (rains)

25

30

Fig. 2. Hydrogen chloride reaction with liquid sodium at 160"C.

produced identical linear P-time curves. [In Fig. 2(a) only one (Curve 3) is shown.] Linear P-time curves were also obtained when the same sodium was further reacted at a series of temperatures in the range 106-160°C. In these cases, however, the reaction was arrested at the stage P'. The variation of pressure change Ap, (i.e. W-P), with time for the curves of Fig. 2(a) is shown in Fig. 2(b). Here the first and subsequent HCI reactions are more conveniently compared. The slope of a AP vs. time curve is a measure of the overall reaction rate for the po ~ p, stage; in Fig. 3 the variation of log (dAP/dt) with 1/T(°K) for the linear "reproducible" reaction stage is shown. An apparent activation energy of 12.7 kcal/mole was obtained from the slope of this line. In Table 1 values of initial pressures po and "break" pressures P' are listed with values for the ratio po/p,; in every case (po/p,) _ 2. Hence the sudden decrease in reaction rate took place when the pressure was reduced to one-half its initial Value. Other experiments were conducted where the gas phase was sampled after arresting the reaction at P'. Analysis by mass spectroscopy failed to detect any HCI; the gas phase consisted entirely of hydrogen.

Sodium-barium alloy Typical pressure-time curves are shown in Fig. 4 for the reaction HCI with sodium containing - 4 At% barium. In all cases a break in the curve (at P') again occurred when the pressure had fallen to one-half of its initial value. However, the first addition of HC1 did not show the initially increasing reaction rate with time characteristic of pure sodium. In separate experiments, the gas phase was sampled at the stage corresponding

Liquid Na- and Na/Ba-HCI reaction

40

I

I

I

1

1447

I

30

20

B

160*C

e

0.

8

Io

7 6

~o

5 4

3-

2-

~

i 2'2

I 2.3

I 2.4

I 2"5

f 2.6

6 °¢

--

2.7

2-8

103/T=K

Fig. 3. Variation of the sodium-hydrogen chloride reaction rate with temperature.

I

I

I

~0

80{ p O p"

"r

60

-

-

80 -r

tN

40

-

20

--

-

E E --

I,g nr aO O3 hi n.O-

40 -

180°C

"~%, • P' p,"~ ~. ~240°C

~

"',.

o

"X ~

I

I

I

I

0.5

I

1.5

2

~ ~-J, 2-5 .3

20

0.5

1.0 TIME

(hrs)

Fig. 4. Hydrogen chloride reaction with sodium-barium alloy.

1.5

1448

M . R . HOBDELL and L. NEWMAN Table 1 Temp. P° P' (°C) (mmHg)(mmHg) 106 117 128 136 148 156 160 160 160

126 135 130 118 134 86 110 72 171

63 70 67 60 69 43 57 36 89

P°/P' 2"0 1"92 1"94 1"96 1"94 2"0 1"92 2"0 1"9

to P' (Fig. 4) and only hydrogen was detected (mass spectroscopy). After evacuation and cooling to room temperature, the solid product was carefully removed from the metal surface, and remaining traces of free sodium and barium removed by repeated mercury washing. The residue was analyzed (by atomic absorption) for barium and sodium; an atom ratio barium: sodium of 5 : 1 was found, i.e. the product appeared to contain some 83 per cent BaCI2. The amount of barium consumed during the experiments was negligible compared to the total amount available so that the alloy composition remained effectively constant throughout. DISCUSSION

Pure sodium Hydrogen production. The results showed that each HCI reaction took place in two separate steps: N a + HCI ~ ½H2 + NaCI

(1)

Na+½H2 ~ Nail.

(2)

In the first step where po ~ p,, HCI reacted with sodium to produce gaseous hydrogen. The second step, where po __. O, was due to the reaction of hydrogen to produce sodium hydride, which had a negligible dissociation pressure at the temperatures involved[l]. The fact that po/p, was always close to two (Table 1) indicated that the first reaction (1) proceeded effectively to completion before step (2) took place, i.e. all hydrogen appeared in the gas phase as H2 prior to sodium hydride formation, and only the first stage involved HC1. This observation was of interest since: (1) In the process: H - C I -~ ½H-H + CI-, "reactive" hydrogen intermediates (e.g. H .) would be formed in contact with, or in the vicinity of the sodium surface. (2) The stable form of hydrogen was the hydride ion, H - (as sodium hydride) in the presence of excess sodium at the temperatures concerned. In view of its slow reaction rate, negligible uptake of molecular hydrogen was to be expected during the time taken for complete HCI reaction(step 1). However,

Liquid Na- and Na/Ba-HCI reaction

1449

it was surprising to observe that there was also no net uptake of the atomic hydrogen presumably produced during the reaction. There are two feasible reaction mechanisms which would lead to a quantitative production of molecular hydrogen in this way: (1) If it is assumed that hydrogen atoms are transiently adsorbed on the sodium surface, a rapid secondary reaction: H-(ad) + H ~ I ---> H2 + Cl-(ad) could occur; thus there would be no net uptake of hydrogen until all HCI had reacted. Alternatively, (2) because of their high electron affinity, chlorine atoms might be expected to readily accept electrons from the metal surface to form an adsorbed chloride layer: HCI + e --~ Cl-(ad) + H. The metal surface would be rendered inactive to released hydrogen atoms (until chlorine was desorbed as NaC1), i.e. the process H. + e ---> H - would be inhibited and hydrogen atoms would combine to form hydrogen molecules: 2H. ---> Hz. If the initial adsorption of HC1 (via CI) is assumed to always be favored over that of hydrogen, then again there could be no net hydrogen uptake until all HCI had been consumed. The initial acceleration stage. The shape (" concave upwards") of the APtime curve for the first HCI addition (Curve 1, Fig. 2b) strongly suggested that the reaction was autocatalytic. An initial accelerating stage is characteristic of a heterogeneous reaction in which the rate controlling step involves the formation and growth of product nuclei at a phase boundary [7]. After the first HCI reaction the sodium surface always contained some reaction products, so the absence of the accelerating stage in subsequent reactions could have feasibly been attributed to the presence of N a i l or NaCI. Alternatively the initial reaction of HCI could have been inhibited by trace impurities (e.g. Na~O) on the metal surface. (Gas phase impurities were not considered, since if they played a significant role their effect would have been seen in each reaction.) In separate experiments, therefore, HCI was reacted with stirred sodium whose surface contained visible layers of (i) Nail, (ii) Na20, (iii) NaCI. The following observations were made: (i) Nail had no measurable effect, indicating that HCI still reacted preferentially with the sodium. The initial accelerating stage was observed and also the usual (po/p,) __ 2. Sufficient N a i l was present for the complete reaction, N a i l + HCI ~ H2 + NaCI, which would have produced no overall pressure change. (ii) The presence of Na~O gave rise to an extremely rapid initial reaction (a transient red glow appeared on the surface) and no accelerating stage; in addition, (po/p,) > 2. This suggested that the reaction Na20+HCI---> N a O H + N a C I had initially predominated, eliminating the initial slow stage, and preventing some of the hydrogen from appearing. 7. W. E. Garner, Chemistry of the Solid State. Butterworth, London (1955).

1450

M.R.

H O B D E L L and L. N E W M A N

(iii) The presence of NaCI eliminated the initial accelerating stage, and gave rise to curves (at 160°C) identical to Curve 3 in Fig. 2. These simple experiments suggested that the difference between the first and subsequent HCI reactions was associated with the presence of NaCl rather than N a i l or absence of Na20. The absence of further information obviously prevents any firm conclusions in the present case. The implication is, however, that the initial reaction rate of HCI, or rate of hydrogen production, is controlled by the formation and growth of NaCl crystals on the metal surface. In Fig. 5, the data of Fig. 2, Curve 1 is plotted as log AP vs. time. The plot is linear over the initial region, indicating an exponential increase in reaction rate with time. Possible interpretation of reaction curves. Instantly upon contact of HCI with pure sodium there is produced a visible surface film which is assumed to be protective and composed of amorphous sodium chloride. No pressure change is detected at this stage. (The smallest pressure change measurable was 0.5 mm Hg which would correspond to a uniform layer of sodium chloride of - 103 ,~ thick). Visual observation did indicate that the initial film was cohesive and continuous in spite of the stirring. In the absence of other change, continued reaction will only be possible via some diffusion process e.g., Na + migration from the metal to the NaCI-HCI interface. However, it is proposed that nuclei of crystalline sodium chloride begin to form and then grow at an increasing rate as the interfacial area of the amorphous and crystalline phases increases. If the crystalline sodium chloride so forming is considered to be nonprotective (e.g., is continuously broken up by surface agitation), then reactive metal surface will be made available to the hydrogen chloride at a steadily increasing rate. More amorphous NaC1 will form and the above cycle will be continuously repeated. Thus a rate of pressure drop increasing with time is observed. Eventually, for a fixed metal surface area, the interfacial area of amorphous and crystalline sodium chloride will become constant, resulting in a constant reaction rate. I00

I

I

i

I

I

]

I

I

50

D.. o

I

4

8 12 16 TIME (rains)

]

20

Fig. 5. Initial reaction of hydrogen chloride with sodium at 160°C.

Liquid Na- and Na/Ba-HCI

reaction

1451

The absence of accelerating behavior with subsequent additions of hydrogen chloride is then expected since there would already be present sufficient crystalline sodium chloride to produce an initially rapid reaction rate. The linear curves obtained are a result of the protective amorphous chloride being consumed by the crystallization process at a constant rate, by the time pressure changes are detected. In the above interpretation the reaction of hydrogen chloride with sodium and the subsequent formation of a protective sodium chloride layer is considered to be rapid at all times. The rate controlling step for pressure decrease is assumed to be the conversion of this protective layer to a nonprotective one. If as suggested above the protective and nonprotective layers are composed of amorphous and crystalline sodium chloride respectively, then the activation energy obtained from Fig. 3 will be that for the process: NaCl (amorphous) --> NaCl (crystalline).

Sodium-barium alloy Comparison with pure sodium. The reaction again took place in two quite distinct steps: HCi ~ CI-+½Hz

(1)

½H2 ~ H-.

(2)

Again all hydrogen appeared in the gas phase prior to subsequent reaction. In this case, however, the rate of the second stage (P' ~ O) at 200-240°C was comparable to the rate of the first (po ~ p,) and at 160-180°C was even more rapid than the first stage (Fig. 4). Comparison of typical reaction curves at 160°C produced by pure sodium and sodium containing 4 At% barium is made in Fig. 6. The effect of barium is

I

~o

I

I

I

I

I

160 pO 120 _- ~ ~ HCL'*I/2 . . ~ Hz÷ Ct"

E

m ~ w

CURVE l: PURE No 2: No/Bo

HCL-~I/2 H2+Cl-

80

r~ a_

-

_

~

o

\1/2 H2--~ H"

.

I

2

3 4 5 6 TIME (hrs) Fig. 6. Hydrogen chloride reaction with sodium and sodium-barium alloy at 160~C.

1452

M . R . H O B D E L L and L. N E W M A N

clearly illustrated. The uptake of HCI was markedly retarded, but the reaction with hydrogen was considerably enhanced. In other words, the effect of barium was to decrease the rate of hydrogen production, while increasing that of hydrogen uptake. Hydrogen production. The experiments with barium, made it readily apparent that the reaction of molecular hydrogen with sodium (as well as atomic) was inhibited in some way in the presence of unreacted HCI. (i.e. prior to P'). As the first stage of the reaction (po .._} p,) neared completion the partial pressure of hydrogen would greatly exceed that of HCI, but in spite of its potentially rapid reaction rate none reacted. This implied that hydrogen was not produced by the reaction: H - ( a d ) + HCI ---> H2+ CI-. It was also unlikely that solid reaction products would physically restrict access of hydrogen, since they were still present as the reaction proceeded beyond P'. The implications were (l) the initial adsorption of HC1 was far more rapid than that of H. or Hv (2) Only initially formed chloride prevented electron transfer to hydrogen (H. or H2) to form hydride (H-). The presence of unreacted HC1 thus effectively prevented the reaction of any form of hydrogen. Analysis of the reaction curves. The shape of the pressure-time curve from po to P' at 160°C in Fig. 4 indicated an initially rapid pressure drop which suddenly decreased. This suggested that the fall in pressure varied logarithmically with time, i.e. AP = klog(a+bt). Using the data of Fig. 4, AP values were calculated and plotted against log(1 + t) in Fig. 7. Linear plots were obtained. (The discontinuities correspond to P'). Other relationships were tested: neither the power law (log AP vs. log t), exponential law (log AP vs. t) or inverse log law (I/AP vs. log t) would give linear relationships over an extended range of values. Thus, the law, log AP = k( 1 + t) was considered to uniquely express the curve of the overall reaction occurring between po and P'. (The data at 200°C was not in harmony with that at other temperatures but still produced a linear curve). Laws of this type are met in kinetic studies of the oxidation of metals [8-10], but most models proposed to explain them involve transport of reactive species across a very thin oxide film (< 102 ~). It is highly unlikely that such models would provide an explanation in the present case which involves chloride layers approximating to 104,~ thick. However, Evans[8] and Harrison[l 1] have discussed oxidation mechanisms which will allow for a logarithmic oxidation law where thick films (> 10a A) are involved. The present results will be considered in this light. Possible interpretation. As indicated earlier, the main reaction is: B a + 2HCI "-~ BaCI2 + H2. 8. U. R. Evans, The Corrosion and Oxidation of Metals. Arnold, London ( 1961 ). 9. O. Kubaschewski and B. E. Hopkins, Oxidation of Metals and Alloys. Butterworth, London (1962). 10. P. Kofstad, High Temperature Oxidation of Metals. Wiley, New York (1966). 11. P. L. Harrison,J. electrochem. Soc. 112,235 0965).

Liquid N a- and Na/Ba-HCI reaction 500

I

I

I

I

1453

I

I I00 160" 50 +

z~

A" 200*

I0

i1

o

I

I

20

I

I

40

60

AP (mm Hg)

Fig. 7. Variation of AP with Iog(l + t ) where t is time in minutes for the reaction of hydrogen chloride with sodium-barium alloy.

A possible reaction scheme might involve the same basic steps as considered in the pure sodium case: (1) Adsorption of HCI on the metal surface with the formation of chloride ions and liberation of hydrogen, (2)formation of amorphous BaCI~, and (3) the growth of a layer of crystalline BaCl2. However, barium chloride differs from sodium chloride in that the volume ratio (molecular volume of BaCIJatomic volume of barium) for BaCl2 on barium or sodium-barium alloy is greater than unity, and ionic conduction in BaCl~ takes place via Cl- migration [7]. Consequently, a growth of a continuous BaCl2 layer might be expected to occur by diffusion of Cl- ions inward to the metal surface. The Pilling and Bedworth rule would therefore predict highly protective layers of BaCI,. In practice, when the metal was unstirred, no pressure change was detected, although the metal surface did acquire a visible film. None of the above reaction steps is therefore considered to be rate controlling, and no autocatalytic behavior is observed. However, because of the stirring, external work was being done on the system. It is proposed that this causes the continuous formation of "leakage paths" [8, 11] (e.g. cracks and pores) in the surface layer~ The reaction of HCI is considered to take place only at leakage points, and as discussed above does so at the expense of H. or H2. Following the treatment by Evans [8], if it is assumed that

1454

M.R. HOBDELL and L. NEWMAN

the increase in volume associated with the formation of BaCI~ at each leakage point can result in the blocking of others nearby, then one can arrive at a logarithmic growth law for the BaCI2 layer. Since the quantity of BaCI2 produced was proportional to AP, a possible interpretation of the form of the reaction curves is obtained. When all HC1 had been consumed (at P') no more BaCI2 could be produced, and the production of "leakage paths" would proceed unchecked. Continued stirring would then cause a complete break-up or coagulation of the chloride (visual observation) and the reaction of H~ would take place.

The hydrogen reaction The second stage of the reaction curves from P -- P' to P = 0 has been shown to be due to the uptake of hydrogen alone. The reaction of hydrogen with liquid sodium and the effect of dissolved barium has been the subject of a separate investigation by the present authors [5]. Under the conditions described above for hydrogen chloride, the reaction rate was first order with respect to hydrogen pressure in both cases. The presence of barium enhanced the reaction rate by two orders of magnitude at 200°C. In contrast to hydrogen chloride, the rate controlling step for hydrogen uptake was considered to involve only the reactants and not the reaction product.