A FRESHMAN EXPERIMENT? THE REACTION OF MAGNESIUM WITH ACIDS* MARTINKILPATRICK, JR.,UNIVERSITY OF PENNSYLVANIA, PHILADELPHIA, PENNA. I n accordance with the okier definition of acids one would say that the reaction of magnesium with acids i s a reactin between the metal and the hydrogen ions. This p o b l m i s discussed from the viewpoint of the Bronsted definition of a n acid and i t i s s h m that while the reaction i n dilute aqueous solutions of stronc acids i s with the acid H30+,the predominant reaction i n aqueous solutions of weak acids i s not with the hydrogen ion &0+ but with the other acids present. Experimental d m c e i s submitted to show that magnesium reacts urilh molecular chloroacetic, acetic, formic, and glycollic acids and with the ammonium ion, which i s an acid. In solutions of low hydrop&on concentration the reaction with the hydrogen in i s negligible. I t is pointed out that these results are not confined to the metal magnesium and that this point of wiew may be a fundamental one i n the study of corrosion. On the grounds that the new ideas lead to simplicity in one's concept of acids and bases i n non-aqueous and in aqueous solution, and to Progress in the.field of homogeneous catalysis and i n the study of the rate of solution of metals in acids, a plea is made for the teaching of acids and bases from the new point of view.
. . . . . .
The present paper is an attempt to bring to the attention of the teachers of chemistry the modem conception of acids by showing how this new viewpoint modifies our ideas on the rate of solution of metals in acids. If we were asked what happens when magnesium is added to a dilute aqueous solution of hydrogen chloride, we would probably reply that the metal reacts with the hydrogen ion and hydrogen is given off. If the same question were asked concerning a weak acid we would say the same and add that, as the hydrogen ions are removed, more appear, due to further ionization of the weak acid. An acid would therefore be a substance which forms hydrogen ions in aqueous solution. Let us now consider the more recent definition of an acid (I). Bronsted defines an acid by the formal equation A e B + H t
where A is an acid, irrespective of the charge, B is a base, and H + is a proton. In accordance with this definition the hydrogen ion, H30+, is an acid in aqueous solution, but we have other acids in aqueous solution such as the uncharged formic acid molecule or even the ammonium ion.
= =
+
HIO+ H,O H+ HCOOH HCOO H+ H+ NH4+ 6 NH8 H+ H,O 5 OH-
+ +
+
* Delivered before the Division of Chemical Education at the 81st meeting of the A. C. S., Indianapolis, Indiana, March 30 to April 3, 1931. 1566
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Of course, the ammonium ion is quite a weak acid since it does not give up its proton very readily. Water is both an acid and a base as indicated above, and the hydroxyl ion is only one of many bases. This definition furnishes a general hasis for consideration of acids and bases in aqueous and non-aqueous solution and has been a very profitable viewpoint in the study of homogeneous catalysis (2). In aqueous solution there exists a double acid-base equilihrium, due to the fact that the solvent itself is both an acid and a base. The reaction accompanying the solution of hydrogen chloride in water may be written add HC1
base
+ HaO
acid HaO+
base
+ Cl-
Hydrogen chloride is such a strong acid that it readily gives up its proton to the base, water, the equilibrium heing displaced practically completely to the right. The acid present besides water in a hydrochloric acid solution is the ion HJO+ or, as we shall call it, the hydrogen ion. The concentration of the proton is so small as to he of no kinetic significance. The case above is a general illustration of the state of affairs existing when any strong acid is dissolved in water and explains why all strong acids appear to he of equal strength in dilute aqueous solution. Let us now consider a weak acid. When formic acid is dissolved in water we have acid HCOOH
acid base + base H 2 0 = HaOf + HCOO-
The equilihrium is not displaced far to the right as in the case of a strong acid, so we have, besides water, two acids in appreciable concentrationfirst, the acid H30+, and second, HCOOH. The problem is to determine whether the magnesium reacts with the acid H30+,with the acid HCOOH, or with both. Our first experiments were so striking that they are worth describing. The experiments were camed out by rotating a cylinder of magnesium in a known volume of the solution and measuring the rate of evolution of the hydrogen. Our first experiment was with 0.01 M hydrochloric acid solution, that is to say, with a solution 0.01 M in hydrogen ion. Hydrogen was evolved a t a rate of 0.5 cc. per minute. We then made up a buffer solution 1.7 M in formic acid and 0.1 M in sodium formate. The hydrogen ion concentration of such a solution is approximately 0.005 M; if the reaction were with the hydrogen ion alone a rate of about 0.25 cc. per minute would be expected. The experiment was camed out with the same piece of magnesium and the rate of evolution of gas was much too rapid for measurement, heing of the order of 100 cc. per minute. From this we concluded that the reaction might he partly with the molecular formic acid unless the formate ion or the change in medium had some remarkable effect.
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The systematic experiments, carried out with Mr. Rushton ( 3 ) ,are given below. All experiments were performed a t 25'C. Figure 1 shows the proportionality between the rate of reaction and the hydrogen ion for dilute solutions of strong acids. The results also indicate that the electrolyte (salt) effect is small in this case. Figure 2 summarizes the experiments with formate-formic acid buffers. The rate of reaction is proportional to the formic acid concentration. In each series the H30+concentratiou was kept practically constant by diluting the buffer with sodium chloride solution so that the equivalent salt concentration was kept the same. According to the law of mass action the
hydrogen-ion concentration of the buffer solution is given by the equation [HaO']
=
K.[HCOOH]/IHCOO-I
where the brackets represent concentration and K, is the classical dissociation constant. Now K, is not a true constant but depends upon the concentration of electrolyte and to some extent upon the individual electrolyte present. By keeping the electrolyte concentration fixed, K, is sufficiently constant, however, so that the hydrogen-ion concentration may, for our purposes, he considered constant. From the fact that the hydrogenion concentration in the second series is 75 per cent less than in the first, and from the fact that the best line through the points passes practically through the origin we may conclude that the reaction between the H80+ and the magnesium is negligible a t these concentrations of hydrogen ion.
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Figure 3 shows the results with acetate-acetic acid and with glycollateglycollic acid buffer solutions. In both cases the rate of reaction is proportional to the concentration of the molecular acid. The rates with glycollic acid are faster than with acetic as might be expected from the fact that glycollic acid is a stronger acid than acetic, i. e., it gives up its proton more readily. Figure 4 gives the results with chloroacetate-chloroacetic acid buffers. Again proportionality between the rate and the acid concentration is established. Magnesium also reacts with the ammonium ion but the results do not
show good proportionality when experiments in ammonia-ammonium chloride buffer solution are considered. The rate seems to decrease as one increases the concentration of the ammonia. This effecthas also been observed by Bekier and Zablocki (4). This specific effect may be due to the formation of a film of magnesium hydroxide. The reaction of the metal with water has not been investigated but i t is well known that magnesium does react with water a t high temperatures. The next step is t o see if there is any relationship between acid strength and the velocity constant. Velocity constants for the experiments described were calculated by means of the equation
JOURNAL OF CHEMICAL EDUCATION
hr Liter M. ntd Na Salt in e . r d
Males of CHJICOOH 0.1
Arid
-
Chlorerctafa
F I G ~4E
AUGUST. 1933
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A FRESHMAN EXPERIMENT?
where V is the volume of the solution in cc. ; S, the surface of the magnesium in ern.%; t, the time in minutes; and co and c, are the initial and current concentrations of acid, respectively. In the second column of Table 1are given the average values of the velocity constants; in the third column are a number of velocity constants obtained by Tarle (5). TABLE I The Rate of Solution of Maenesium in Acids Add
Hydrogen ion, H,Ot Formic acid Acetic acid Glycollic acid Chlaroacetic acid 1.2.5-Dioayhenzoic acid 8-Resorcylic acid Tricarballylic acid Ammonium ion
1.20 0.460 0.324 0.375 0.390
1.23 0.390
0.774 0.511 0.489 0.0984
56 2.1 x 1.8 x 1.5 X 1.6 x 1.1 X 5 2 X 2 2 x 3.2 X
IUP 10-5 lo-"
IOF 10F 10P
lo-'"
Although an acid is defined by the equation A=B+H'
the acid strength of the substance, A, cannot he expressed by the equilibrium constant of the process above, K.' = [BI[H+I/[AI
(1)
for the concentration of the proton is too small to be measurable. This constant can, however, he related to the equilihrium constant Ko which defmes the acid property of the hydrogen ion, KO= [H~OI[H+I/[H~O+I
(2)
Dividing (1) by (2) one obtains [BI[HJO+]/[AI[HBOI= K.'/Kn
the ratio being equal to the equilibrium constant for the double acid-base equilihrium. Considering the concentration of the water as constant, and transposing i t to the other side, one has [B][HaO+I/[A] = K.'[HzOI/Ko = K,
where K,' is the customary dissociation constant of the acid, A. The dissociation constant may therefore be taken as a measure of acid strength in any given solvent. In Figure 5 the logarithm of the velocity constant is plotted against the logarithm of the dissociation constant of the reacting acid, and the best
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straight line has been drawn through the points. The ammonium ion point is uncertain due to the effectof ammonia, but the plot seems to indicate the same general relationship between the velocity constant and acid strength as found in certain cases of acid and basic catalysis (6). These ideas do not apply to the metal magnesium only but are quite generally applicable. King (7) has obtained reproducible results with zinc by the addition of potassium nitrate or hydrogen peroxide to the acid solution; he has shown that zinc reacts with acids in the same way as does magnesium. In this laboratory we are investigating the reacting between amal-
gams and acids from the same point of view with promising results (8). The work has also been extended to non-aqueous solutions. An examination of the literature indicates that this conception of acids may be a fundamental and a profitable one in the study of corrosion, especially the corrosion of tin and lead. The results with metals dissolving in acids, as illustrated above, and the results in homogeneous catalysis indicate that the new conception of acids and bases has made for progress in these fields. One might say, off-hand, that it has made acids and bases more complicated and that the old ideas are simpler and more teachable. This is not the case. The new ideas lead to simplicity in non-aqueous solution. The student who has not
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learned the old ideas takes to the new ideas much more readily than to the old ideas, even in aqueous solution. I believe that teachers of chemistry should begin to teach their freshmen from this point of view. After all, chemistry today is not confined to aqueous solution.
Literature Cited (1) For references, see HALL, J. CHBM.EDUC.,7, 782 (1930). (2) For references, see LIVINGSTON, ibid., 7, 2887 (1930). (3) KILPATRXCK AND RUSHTON, 3. Phyr. Chem., 39, 2180 (1930). 4 BEKIER AND ZABMCKI, Roczniki C h . ,10, 314 (1930). (5) TARLB,Dissertation, Leipzig, 1912. (6) BRONSTEDAND PEDERSBN, Z. physik. C h m . , 108, 185 (1924); BKONSTHD AND D w s , Z. physik. Chem., 117, 299 (1925); BR~NSTED ANO GUGGBNHBIM, 1.Am. Chem. Soc., 49, 2554 (1927); BRONSTED AND WYNNB-JONES, Tmm. Faradoy Soc.,
25, 59 (1929). (7) KIND, unpublished results, New York University. AND DUNNING, unpublished results. (8) KILPATRICK
