T H E EFFECT OF T H E ADDITION OF SOME ALKALOIDS ON T H E RATE OF SOLUTIOTU' OF IRON IN DILUTE HYDROCHLORIC ACID PAR?' I M. B. R A N E A N D MATA PRASAD
The action of acids on metals has been studied by various workers. Divers and Shimidzul investigated the action of mixtures of sulphuric and nitric acids on zinc and found that the rate of solution of zinc depended on the surface of the metal exposed and the concentration of the acid, but was independent of the quantity of the acid used. Spring and Aube12 examined the action of various acids on zinc containing lead and found that the rate of evolution of hydrogen was slow a t first, but increased rapidly afterwards till it reached a maximum and then decreased in such a way that the reaction was proportional to the decreasing concentration of the acids. Veley3 found that the rate of reaction of the acids on metal was affected by the dimunition in the concentration of the acids in the immediate neighborhood of the metals and by the gas bubbles that adhere to the surface. Friend and Dunnett4 also observed that the rate of solution of iron in dilute sulphuric acid was greatly altered when the reactants were kept in motion. A systematic study of the rate of solution of iron in hydrochloric acid was made by Conroy5. He measured the rate of solution of iron at successive intervals of time, by noting the volume of hydrogen evolved in a given time, and found that the rate rose to a maximum after a certain time depending upon the temperature and the concentration of the acid. Incidentally he found that the rate of solution of iron was diminished by the presence of small quantities of arsenious oxide. Arsenious oxide is a well-known poison and the inhibiting effect observed by Conroy may be attributed to the poisoning action of arsenic. The authors were, therefore, led to examine if other poisonous substances also affect the rate of solution of iron in hydrochloric acid in a similar manner as arsenic. The inhibiting effect of arsenious oxide has not been satisfactorily explained. It is likely that arsenious acid forms a thin metallic film on iron and thus protects it from being affected by the acid. The effect of poisons on the corrosion of iron and other metals and on the solution of iron by acids are somewhat analogous in character to the phenomenon of possivity. In order to obviate the likely formation of a metallic film a s postulated in the theory generally put forward to explain the inhibiting action of such subJ. Chem. SOC.47, 597 (1885). ZAnn.Chim. Phys. (6) 11, 505. 3 J. Chem SOC. 5 5 , 361 (1889). J. Chem. SOC.121, 41 (1922). J. SOC. Chem. Ind 20, 316 (1901).
2 50
M. B. RANE AND MATA PRASAD
stances as arsenic oxide in the above phenomenon, it was considered desirable to study the behaviour of such poisons as the organic alkaloids under the same circumstances. Experimental In the present investigation the effects of various alkaloids on the rate of solution of iron in hydrochloric acid have been studied. Two sets of experiments were undertaken in the investigation. The first consisted of the measurement of the rate of solution of pure iron in hydrochlorjc acid of various concentrations. These experiments had to be performed as there were no reliable data on the subject, the only work of importance being that of Conroy who unfortunately used commercial iron.
FIG.I
The second set of experiments was performed after the addition of a given quantity of a known alkaloid and the changes in the rate of solution of iron in hydrochloric acid were noted by measuring the quantities of hydrogen evolved at various intervals. The apparatus used is a modification of Conroy’s apparatus and is shown in Fig. I. The rubber stopper A of the Woulff’sbottle has two holes. Through one passes a small burette for adding the alkaloid when needed and through the other passes a thermometer T for registering the inside temperature of the bottle and for carrying a small glass hook for suspending the coil of iron wire. The burette B in the other stopper serves the purpose of adding strong hydrochloric acid as will be shown later. The gas is led out through a delivery tube C and is collected in burettes D over water under reduced pressure. A system of two burettes is connected to the delivery tube by a two-way stop-cock S so that the gas evolved can be measured continuously in one or the other burette. The bottle is placed in a thermostat maintained at a constant temperature of 45OC and kept suspended in it from a support fixed to the side of the thermostat. As the rate of evolutions of hydrogen by the action of acids on metals
RATE OF SOLUTION O F IRON
251
is greatly influenced by the motion of reactants’, the bottle is given a regular and constant shaking by means of a mechanical attachment to the crankshaft of a hot-air engine. As pure electrolytic iron was not available, the best iron wire that could be had was employed. The specifications of the wire were: Wire No. 2 0 S.W.G. Diameter 0.0927 cms. The analysis of the wire gave the following result :Iron 99.619 per cent. Manganese 0.310 per cent. Carbon 0.071 per cent. 400 cc. of the acid were placed in the bottle which was suspended in the thermostat as described above. It was kept there for sufficient time till it acquired the temperature of the bath. The thermometer T indicated that the temperature of the acid inside the bottle was the same as that of the bath. 87 cms of iron wire were coiled and inserted in the bottle through the glass hook. The gas evolved mas measured after an interval of every 2 0 minutes. The temperature and pressure a t which the gas collected were also noted and the volumes were reduced to N.T.P. Friend and Dunnett have observed some anomalies in the curves obtained by plotting the rate of solution of iron in sulphuric acid against the concentration of the acid. The effect of concentration of the hydrochloric acid on the rate of solution of iron is investigated by the authors and the results obtained are shown in Table I.
TABLEI Amount of Hydrogen evolved a t N.T.P. per hour Time
st hour 2nd ” 3rd ,’ 4th ” 5th ” 6th ” I
Concentration of HCI. 2N 2.5N
3N
4 s
N
1.5s
cc.
cc.
cc.
cc.
cc.
cc.
117.4 171.3
137.6
211 2
275.1 430.3 406.1 114.4
141.6 221.3 337.6 500.8 352’7
144.1 234.3 352.7 549 3 337.0
194.0 247.4 384.0
91 ’
7
143.6 177.4 224 3 286.7 248.4
305.3 388.6 325.6
200. I
It will be seen from the above table that the rate of solution of iron increases as the concentration of the acid is increased and no anomaly is observed. Also in each case the rate first rises, goes up to a maximum, and then decreases with a great rapidity. These results are similar to those obtained by Conroy and exhibit no periodicity2. ‘Cf. Friend and Dunnett: J. Chem. SOC.121, 141 (1922). Cf. Ostwald: Z. phpsik. Chem. 35, 3 3 , 204 (1900); Hedges and Myers: J. Chem. SOC. 125, 608 (1924).
M. E. RAKE AKD MATA PRAGAD
252
There are two factors which are affecting the rate of solution of iron and they are:I. The diminution in the percentage of the free acid in the solution due to‘ the reaction. 2. The increase in the surface of the wire which when taken out of the acid and examined, is found to be more porous than before the experiment. To counteract the effect of the first ___________~ -----factor, a strong solution of hydrochloric acid was prepared and placed in the burette B. After every 2 0 minutes the amount of the acid consumed was calculated and was added to the bottle. There mas certainly a very slight increment in volume thereby (never exceeding 3 cc. in all) but it has been shown by Veley, Divers and Shimidzu, and others that the volume of the acid is ineffective in influencing the rate of reaction. In subsequent experiments 300 cc. FIG 2 of the normal hydrochloric acid and j o cms. of the wire were used. The results of the variation in the rate of evolution of hydrogen with time on keeping the concentration of the acid constant are shown in Table 11. ~
i
’
TABLE
Time in Minutes
Volume of the gas evolved per 20’ at N. T. P.
Time in Minutes
cc. 0’-
20’-
40’-
60’80’-
20’
40’ 60‘ 80’ 100’
100’-
I 20’
120’-
140’
18.68 23 ’ 63
11
Volume of the gas evolved per 20’ at X. T. P.
Time in Minutes
cc.
27.11
140’-160’ 160’- 180’ 180’-220’
30.09 31.37 35.76 37.54
220’-240’ 240’-260‘ 260’-280’
200’-220’
Volume of the gas evolved per 20’ at X. T. P. cc.
41.21 45.67 53. 5 2 58.22 66.23 73.63
280’ - 300’ 300’- 320’ 320’ - 340’ 3 40’ - 360’ 360’-380’
76.20
400’-420’
380’-400’
77 ’ 72 75.27 67.28 58.08 42 * 90 21.68 8.29
By plotting the volume of the gas evolved each 2 0 minutes against time, a curve is obtained as shown in Fig. 2 . This curve is almost similar to the curve obtained by plotting the values given in Table I.
Effect of Alkaloids The effect of the alkaloids available in the laboratory was examined. The alkaloids tried in the preliminary experiments were codeine, cinchonine, brucine, strychnine, nicotine, and cocaine. I n all the cases the inhibiting effect was observed.
RATE O F SOLVTIOPU’ O F IRON
253
A detailed study of only the first three alkaloids has been made in this part of the effect of the addition of the varying amounts of alkaloids on the rate of solution of iron has been examined. I n most of the experiments the alkaloids were added initially to the acid and in some cases they were added after the reaction had proceeded for sometime. The solutions of all alkaloids were prepared in normal hydrochloric acid and the total volume of the acid and the alkaloid solution was ,500 C.C.in all cases. The results of the investigation are indicated in Tables 111-V. TABLEI11 Effect of Codeine on the Rate of Solution of Iron Volume of the gas evolved on the addition of the following quantities initially Grams Grame; Grams Grams Grams Grams Grams 0,0008 0,0002 O.OIj0 O.OIOO 0.0050 0.0020 0.0200
Time
cc.
cc.
cc.
cc.
2.91 10.61 17.36 18.62 19.29 21.89 23.17 23.22 23,90 24.80 25.31 26.93
6.16 12.99 16.02 22.10 24.33 24.42 24.55 25.29
6.53 16.50 21.94 23.37 26.46 28.23 29.38 30.50 31.31 33.08 33’50 35.17
8.27 17.89
cc. 20’
0’-
40‘ 60’ 80’
20’-
40’60’80’100’-
100’ 120’
120’-140’ 140’-160’ 160’- 180’
1.76 5.80 11.50
16.20 18.07 19.89 20.30 20.50
200’-220’
21.66 21.91 21.09
220’-240’
22.75
180’-200’
25.52
25.99 28.79 30.95
cc.
22.02
24.40 25.39 30.89 31.74 33.58 35.02 38.73 41.35 43.80
9.55 20.52
23 ‘ 24 26,63 28.77 31.83 34.93 36.94 38.95 46.22 48.08
cc. 12.81 23.15 29.55 32.39 35.03 37.20 40.91 41.16 46.70 51.12 54.71
TABLEI V Effect of Cinchonine on the Rate of Solution of Iron Volume OF the gas evolved on the addition of the following quantities initially 0 0200 0 0020 0 0002 o 00015 0 0001 grams grams grams grams grams
Time
cc. .07
0’-
20’
I
20’-
40’ 60‘ 80’
.41 3.31 5.45 6.20
40‘60’80’-
100’
I IO’ - 1 2 0 ’
-140 140’- 160’ 160’-180’ I 80 ’- 2 0 0 ‘ 120
200’
-220’
220‘-
240’
I
7.12
9.63 9.44 10.33 11 .67 11.73 13.71
cc.
cc *
cc.
cc.
2.63 7.55
5.81 14. I O 24.97 31.44 34.10 37.88 36.83 41.97 42 ’ 23 46.75
5.77
6.84 19.86
12.27
18.76 21.73 24.33 26.93 28.24
29.80
50.0 2
45 ’ 73
16.54 28.65 31.09 33.12 35.29 35.70 37.05 40.15 43.86 48.43 53.09
29.42
34.04 35.83 36.59 37.10 39.05 40.73 43.27 47.40 52.73
M. B. RANE AND MATA PRASAD
254
TABLE T.1 Effect of Brucine on the Rate of Solution of Iron Volume of the gas evolved on the addition of the following quantities initially
Time
0’-
‘40‘-
60’80‘-
20’
40’ 60’ 80’ 100‘
100’-120‘
120‘-140’ 140’- 160‘ 160‘- 180’ 180’ - 2 0 0 ’ 200’ -220‘ 220’-
240‘
0.0020
0.0002
0.00015
0.0001
grams
grams
grams
grams
cc.
cc.
cc.
5.61 13.90 17.80
12.33 19.18 23.74 26.79 28.35 31.48 29.37 32 ’ 54 35.45 39.64 45.44 50.57
12.66 18.35 25.02 29.54 32.57 36.66 38.77 41.80 45.53 50.68 55.00 59,46
20.25
23.33 24.32 28.19 28.57 31.03 37.47 43.30 47.10
cc. 12.68 21.08 25.70 29.46 33.84 36.09 38.23 42.91 47.18 50.90 57.93 63.20
Discussion of Results It will be seen from the Tables I and I1 that the evolution of hydrogen is similar in both the cases. I n the latter case the concentration of hydrochloric acid is maintained constant while in the former no such allowance is made. This indicates that the rise and fall in the rate of solution of iron in hydrochloric acid is not much affected by the slight variations in the concentration of the acid due to neutralisation. The rise in the rate of reaction is due to the iron which becomes more and more porous as the reaction proceeds and thus offers a larger and larger surface for reaction. The rate of reaction increases up to a certain limit when the iron reaches its maximum limit of porosity and consequently the largest surface. The curve, shown in Fig. 2 , gives an idea of the manner in which the surface of iron changes. After the maximum surface is reached, the further action of the acid lessens the active mass of the iron and the rate decreases as would be required by the law of mass action. It is clear from the Tables 111, IT7 and V that the rate of evolution of hydrogen falls off by the addition of alkaloids even in such small quantities as 0.0001grams. The inhibiting effect is greater as the amount of the alkaloid used is larger. The inhibiting effect varies with the nature of the alkaloid. The following table gives the effect observed when 0 . 0 2 grams of various alkaloids were added to the bottle after the reaction has proceeded for two hours. In these experiments the length of the wire used is 87 em., and the volume of the normal hydrochloric acid is 400 C.C. The alkaloids can, therefore, be arranged in the following series according to their poisonizlg action upon iron: brucine > strychnine > cinchonine > nicotine > codeine > cocaine > conine.
255
RATE O F SOLUTION O F I R O N
Effect of
Time
Volume of the gas evolved in each hour on the addition of alkaloids Without the Brucine Strychnine Cinchonine Nicotine Codeine Cocaine Conine alkaloid
cc. Ist hour 2nd hour grdhour 4thhour sthhour 6thhour
0.02
TABLE VI grams of Alkaloids on the Rate of Solution of Iron
gI,7
cc. .~
143.6 -35.78 2 2 4 . 3 35.78 286.7 36.28 248.4 3 7 . 7 9 177.4
cc. ~
cc. ~
cc.
cc.
__
-
cc.
_______
__
39.81 43.33 46.86 49.89
cc. __-
62.48 9 1 . 2 0 9 1 . 7 0 1 2 0 . 4 0 160.30 68.53 109.80 - I 2 j . 0 0 2 0 j . 1 0 72.06 110.40 9 3 . 7 4 138.00 276.70 72.56 133.00 99.79 145.60 2 5 2 . 0 0
To see if the above phenomena is any way similar to passivity of iron, the following test1 indicative of passivity was made. Two similar coils of iron wire were taken and dipped into two beakers, one containing normal hydrochloric acid and the other normal hydrochloric acid and alkaloid. They were taken out after sometime and dipped into beakers containing o.5yo copper sulphate solution. The deposition of copper was almost the same on the two coils. Kext, two similar coils after introducing them in hydrochloric acid for some time were taken out and dipped into two beakers, one containing the copper sulphate solution and the other the copper sulphate solution and the alkaloid. It was observed that the copper deposited exceedingly slowly on the coil which was placed in the mixed solution of copper sulphate and the alkaloid while the deposition was very rapid on the other coil. It would appear from these experiments that the effect of the alkaloid in inhibiting the evolution of hydrogen by the action of hydrochloric acid is somewhat analogous to the phenomena of passivity. It is intended to carry on further work in the direction of making measurements of the electromotive force and other physical properties induced in the iron by the action of the alkaloids. By this means it is hoped that a rational interpertation of the effects observed would be forthcoming. The authors take this opportunity of expressing their thanks to Dr. S. S. Bhatnagar for taking interest in the work. Chemzcal Laboratorzes, Benares H i n d u University, Benares, India. Dunstan and Hill: J. Chem. SOC.99, 1835 (1911).
