Joseph William Mellor.

A comprehensive treatise on inorganic and theoretical chemistry (Volume 2) online

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number of reacting or active molecules is small. The
increase in the number of active molecules with temp,
increases with the sp. ht. C v and the increase in the
speed of the reaction is greater the greater the value
of C v .

The proportion of hydrogen iodide dissociated
decreases with rise of temp, so long as the temp, does
not exceed about 320 ; above that critical temp., the higher the temp, the
greater the amount of hydrogen iodide dissociated. This is illustrated by the
graph, Fig. 4. The thermal value of the reaction changes sign at about the
same critical temp. ; for instance, at 18, the union of hydrogen and iodine is an
endothermal reaction : H 2 +I 2 =2HI 6'1 Cals. ; and at 520, exothermal : H 2 +I 2
=2HI +4*4 Cals. Experience shows that a rise of temp, always favours endothermal
reactions. When a system is in physical or chemical equilibrium a rise of temperature
promotes the formation of those products which are formed with an absorption of heat ;
a rise of temperature resists the formation of those products formed with an evolution
of heat ; and a change of temperature has no effect on the equilibrium thermally neutral
J. H. van't Hoff's equilibrium law (1884). 9 The law is a special case of the
great principle of reversibility. If an exothermal reaction becomes endothermal
at a high temp., there is a curious paradox: A compound may be stable at
temperatures exceeding that at which it dissociates. The case of hydrogen
iodide is particularly instructive. The change in the thermal value of the reaction
corresponds with a change in the effect of a rise of temp, on the equilibrium,
quantitative side of this rule has already been discussed, and W. Nernst 1(
shown for the reaction H 2 +I 2 =2HI, the equation : log A'=540'4T~
+2'35, represents the relation between the equilibrium and the absolute temp. 1 .
in agreement with the observations of G. Lemoine (1877), P. Stegmiiller (1907),
M. Bodenstein (1894), and K. V. von Falckenstein (1910). For the equilibrium
H 2 +Br 2 ^2HBr, the equation log ff=5223T- 1 -0'533 log T-2'72 is similarly

VOL. II.



250' 300"

Temperature

FIG. 4. Effect of Heat on
the Dissociation of Hydro-
gen Iodine.



146



INORGANIC AND THEORETICAL CHEMISTRY



employed for the observations of K. V. von Falckenstein (1910), and M. Bodenstein
and A. Geiger (1904) ; and for the equilibrium H 2 +C1 2 ; F^2HC1, the equation
log A'=9554T- 1 0'553 log T+2'42 is used for the observations of M. Bodenstein
and A. Geiger (1904). The comparison of the percentage degree of dissociation of
the hydrogen compounds of the three halogens in Table III, calculated by

TABLE III.- PERCENTAGE DISSOCIATION OF THE HYDROGEN HALIDES AT DIFFERENT

TEMPERATURES.



Percentage dissociation.



Temperature



TK.


HC1


HBr


HI


290


2-51 xlO~ 15


4-14 x 10- 8


6-2


500


l-92xlO- 8


2-91x10-*


15-5


700


1-12 xlO~ 5


9-93 x 10~ 3


22-2


900


3-98 xlO~ 4


7-18 x 10~ 2


27-0


1000


1 -34x10-3 0-144


29-0


1500


6-10 x!0~ 2


1-19


.


2000


0-41


3-40


.


2500


1-30









F. Pollitzer, emphasizes how the tendency of the hydrogen halides to dissociation
increases as the at. wt. of the halogen increases. At the higher temp., the results
are complicated by the dissociation of the two-atom molecules of the halogen into
one-atom molecules. The free energy of the reaction JH 2 +Jl2 gas r= HI J a t T K.,
according to F. Haber, is -89'575 1'575T log T +0' 00549 T*+RT log (P^Pif^
+2*67T. The equation is not consistent with results obtained in other ways
possibly, as F. Haber suggested, owing to dissociation. G. N. Lewis and M. Randall
give 1340+0-000725T 2 2-48T; and if solid iodine be used, 7110+3'35T
log T 0-00275T 2 41;845T.

The principle applies to physical equilibria. When anhydrous sodium sulphate
is dissolved in water, heat is evolved, and its solubility decreases with a rise of temp. ;
hydrated sodium sulphate dissolves in water with an absorption of heat, and its
solubility increases with rise of temp. The vaporization of water is an endothermal
reaction, and hence a rise of temp, favours vaporization, for it increases the con-
centration of the vapour phase.

The effect o! pressure on equilibria. The principle is also applicable to other
forms of energy. Increasing the press, of a dissociating compound decreases
the amount of dissociation, and this presumably relieves the strain set up by the
increased press. Thus : when a system is in a state of physical or chemical equilibrium,
an increase of pressure favours the system formed with a decrease in volume ; a reduction
of pressure favours the system formed with an increase in volume ; and a change of
pressure has no effect on a system formed without a change in volume G. Robin's
law (1879). n Thus, hydrogen iodide is formed from hydrogen and iodine without
a change in volume, and the state of equilibrium is not affected by variations of
press. When ice melts, the liquid occupies a smaller volume than an eq. amount of
ice ; and experiment shows that the m.p. of ice is lowered by press, in agreement
with the law. With sulphur, the converse is true. The m.p. of sulphur is raised
by press., but the liquid phase has a greater sp. vol. than the solid phase.

The principle of least effort. The principle of least action underlies all these
rules, and it is of great service, and of wide application. P. L. M. Maupertius fore-
shadowed the idea in 1747 : All natural changes take place in such a way that the
existing state of things suffers the least possible change ; or, as W. D. Bancroft
(1911) expressed it : A system tends to change so as to minimize the effects
of an external disturbing force. This has been called the principle of the



THE COMPOUNDS OF THE HALOGENS WITH HYDROGEN 147

opposition of reactions to further change, and it was stated in general terms by
H. le Chatelier : // a system in physical or chemical equilibrium be subjected to a stress
involving a change of temperature, pressure, concentration, etc., the state of the system
will automatically tend to alter so as to undo the effect <\f the stress H. le Chatelier's
law (1888). 12 For instance, if the temp, of a system in equilibrium be raised a few
degrees, the state of the system will change so as to induce the formation of that
component or phase which absorbs most heat, and accordingly tend to lower the
temp. If the -> reaction be exothermal the change will proceed in the reverse
direction ; and if the > reaction be endothermal, the system will change in the
same direction. Again, if the press, of the dissociating iodine I 2 (one vol.) =21 (2 vols.)
be increased, the state of the system will change so that the volume is diminished ;
and conversely, if the press, be reduced, the state of the system will change so that
the volume is increased, that is, the less the press, the greater the amount of iodine
dissociated. In the case of soln., an increase of concentration will induce the
formation of that component or phase which will lower the concentration of the
solute added ; and an increase of vap. press, will lead to the formation of that
component or phase which will reduce the vap. press., etc.

Again, in virtue of H. le Chatelier's rule that all compounds tend to change in
such a way as to relieve the disturbing effects of the strain, W. D. Bancroft showed
that light will tend to destroy all substances which absorb it e.g. chlorine in a
mixture of hydrogen and chlorine and just as all compounds absorb light rays of a
certain wave-length, so must all compounds be sensitive to such rays. The rule,
however, gives no inkling whether the given stimulus will actually produce a change.
Thus, no appreciable change takes place with copper sulphate ; nor with chromium
sulphate unless a catalytic agent be present ; there is a visible decomposition with
silver salts ; and in some cases fluorescent or phosphorescent phenomena occur.



REFERENCES.

1 S. Cooke, Chem. News, 58. 105, 1888; J. F. L. Meslens, Compt. Rend., 83. 145, 1876;
M. Berthelot and A. Guntz, ib., 99. 7, 1884.

2 M. Bodenstein, Zeit. phys. Chem., 29. 295, 1899.

3 J. J. Thomson, Phil Mag., (5), 18. 233, 1884 ; J. W. Mellor, Higher Mathematics, London,
509. 1920.

4 A. Kekule, Liebig's Ann., 106. 129, 1858 ; A. Michael, Journ. Amer. Chem. Soc., 32. 990, 1910.

5 E. Mitscherlich, Pogg. Ann., 59. 94, 1843; M. Faraday, Phil. Trans., 114. 55, 1834;
J. W. Mellor, Chemical Statics and Dynamics, London, 1904.

6 N. Menschutkin, Zeit. phys. Chem., 6. 41, 1890 ; H. von Halban and A. Kirsch, ib., 82. 325,
1913 ; J. T. Cundall, Journ. Chem. Soc., 59. 1076, 1891 ; 67. 794, 1895 ; W. Ostwald, Lehrbuch
der allgemeinen Chemie, Leipzig, 2. ii, 602, 1892 ; W. D. Bancroft, Journ. Phys. Chem., 21. 573,
644, 734, 1917 ; 22. 22, 1918 ; J. H. van't Hoff, Vorlesungen uber theoretische und physikalische
Chemie, Braunschweig, 1. 210, 1898 ; London, 1. 215, 1898 ; W. Nernst, Theoretical Chemistry,
London, 617, 1916 ; 0. Dimroth, Liebig's Ann., 399. 91, 1913.

7 S. V. Sjanoschentzky, Journ. Russian Phys. Chem. Soc., 40. 1676, 1908.

8 D. Amato, Gazz. Chim. Hal., 14. 57, 1884 ; J. H. Kastle and W. A. Beatty, Amer. Chem.
Journ., 20. 159, 1898.

9 J. H. van't Hoff, fitudes de dynamique Chimique, Paris, 126, 1884.

10 M. Bodenstein and A. Geiger, Zeit. phys. Chem., 49. 72, 1904 ; M. Bodenstein, ib., 13. 56,
1894 ; 22. 1, 1897 ; K. V. von Falckenstein, ib., 68. 270, 1910 ; 72. 113, 1910 ; G. Lemoine, Ann.
Chim. Phys., (5), 12. 145, 1877 ; W. Nernst, Zeit. Elektrochem., 15. 689, 1909 ; G. N. Lewis and
M. Randall, Journ. Amer. Chem. Soc., 36. 2259, 1914 ; F. Haber, Thermodynamik technischer
Gasreaktionen, Miinchen, 1905 ; F. Pollitzer, Die Berechnung chemischer Affinitdten nach dem
Nernstschen Wdrmetheorem, Stuttgart, 95. 1912 ; P. Stegmuller, Beitrag zur Kenntnis der Bildung*-
wdrme von Jodwasserstoff aus den Elementen, Karlsruhe, 1907 ; H. Gottlob, Beitrag zur Kenntnis
der Reaktionsenergie bei der Vereinigung von Jod und Wasserstoff, Karlsruhe, 1906.

11 G. Robin, Bull. Soc. Philomath., (3), 4. 24, 1879 ; J. Moutier, ib., (3), 1. 96, 1877 ; H. C. Sorby,
Proc. Roy. Soc., 12. 538, 1863 ; F. Braun, Wied. Ann., 30. 250, 1887 ; 33. 337, 1888 ; Zeit. phys.
Chem., 1. 259, 1887.

12 H. le Chatelier, Recherches experimental et theoriques sur les equilibres chimiques, Paris,
48, 1888 ; Compt. Rend., 99. 786, 1884 ; W. D. Bancroft, Journ. Phys. Chem., 12. 209, 1908.



148



TNOEGANIC AND THEORETICAL CHEMISTRY



5. The Union of Hydrogen and Chlorine in Light

On August 10th, 1801, W. Cruickshank 1 noticed the gradual combination of
oxygenated muriatic acid (i.e. chlorine) with hydrogen, hydrocarbons, and carbon
monoxide. He said :

If the pure oxigenated muriatic acid, in the form of a gas, be mixed in certain proportions
with any of these inflammable gases and introduced into a bottle filled with and inverted
over water, though no immediate action may be at first perceptible, yet, in twenty-four hours
a complete decomposition will be found to have taken place, the products varying according
to the nature of the gases employed. ... I introduced into a phial with a glass stopper,
filled with and inverted over water, one measure of pure hydrogen and afterwards two
measures of very pure oxigenated muriatic gas, this nearly filled the bottle ; the stopper,
was then introduced very tight under water, and before the stopper was introduced, a
whitish cloud appeared in the mixture yet very little or no diminution could be observed . . . ;
at the end of twenty -four hours when the stopper was withdrawn the whole of the gas
instantly disappeared except about one-tenth of a measure, which was found to be azote,
and must have been originally contained in the two measures of oxigenated muriatic acid
and water, for the water in the phial contained common muriatic acid, but did not in the
least smell of the oxigenated acid.

On February 27th, 1809, J. L. Gay Lussac and L. J. Thenard confined a mixture
containing equal volumes of hydrogen and chlorine in darkness and another mixture
in light for several days. The characteristic colour of the chlorine disappeared in




FIG. 5. Hydrogen-Chlorine Actinometer.

less than a quarter of an hour in the vessel exposed to light, while the green colour
of the chlorine in the other appeared to have suffered no change.

Being now no longer able, after these experiments, to doubt the influence of light in
the combination of these two gases, and judging from the rapidity with which it has
operated, that if the light had been much more vivid it would have operated much more
quickly, we made new mixtures and . . . exposed them to the sun. Scarcely had they been
exposed when they suddenly inflamed with a loud detonation, and the jars were reduced
to splinters, and projected to a great distance. Fortunately, we had provided against such
occurrences, and had taken precautions to secure ourselves against accident.

John Dalton also noted in June of the same year that on repeating W. Cruickshank's
experiment, " the gases after being put together over water seemed to have no
effect for one or two minutes, when suddenly the mixture began to diminish with
rapidity," the hydrogen chloride of course was absorbed by the liquid in contact
with the gas as fast as it was formed. J. W. Draper constructed an instrument
which he first called the tithonometer and afterwards changed the name to the chlor-
hydrogen photometer intended to measure the rate of combination of the two gases
under the influence of light from the contraction which occurred as the hydrogen
chloride which is formed is absorbed by the confining liquid. The contraction was
indicated by the movement of the liquid along an index.

The tithonometer or actinometer of J. W. Draper was considerably improved by
R. Bunsen and H. E. Koscoe. Fig. 5 illustrates the principles of the instrument devised by



THE COMPOUNDS OF THE HALOGENS WITH HYDROGEN 149



the latter ; but modified in some details. The mixture of hydrogen and chlorine from the
electrolytic cell enters the tube A, and bubbles through the water in the flattened bulb
of thin glass, F ; the gases pass along the tube E, and after bubbling through the water in D,
escape into the fume-chamber. The tube B with the 3-way cock G is convenient for intro-
ducing other gases into the system. The funnel-tube C and the cock H are convenient for
introducing liquids into F. There is a ground joint at E useful for disconnecting the index
tube EID ; and the cock 7 is useful for arresting the liquid in D. With the cocks H closed, G
and / open, the gases from the- electrolytic cell are passed through the apparatus long enough
to expel all the air, and saturate the liquids in the system. When all is in equilibrium, and
a mixture of hydrogen and oxygen in equal volumes is slowly passing through the system,
the cock G is closed. The position of the liquid in the index tube is noted when movement
has ceased. If the bulb F be illuminated by the light from an oil-lamp or gas-flame, as
hydrogen chloride is formed, it is absorbed by the liquid in the lower part of F, and the
resulting contraction is measured by the left-to-right motion of the liquid in the index
tube IE.

Starting with a mixture of hydrogen and chlorine prepared by the electrolysis
of hydrochloric acid in darkness no movement was observed for 600 seconds after
the commencement of exposure, and after that the time occupied by the liquid
in moving over the 1st, 2nd, 3rd, 4th, and 5th divisions of the scale was respectively
480, 165, 130, 95, and 93 seconds, and thereafter it moved regularly at the same rate.
These results are graphed in Fig. 6, as average velocities per second. Starting from
the moment the mixture is illumi-
nated by a steady source of light,
there is (i) what V. H. Yeley 2
afterwards called a period of inert-
ness, during which there is no
visible sign of chemical action ;
(ii) a period of acceleration, during
which the rate of combination
gradually increases to a maxi-
mum ; and (iii) a steady state
where the rate of the reaction is
uniform and regular. If hydro-
gen and chlorine be exposed to a
bright flash of light there is a



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Period of induction

FIG. 6. The Rate of Combination of Hydrogen and
Chlorine during the Insolation of the M ixture when
the product of the Reaction is removed from the
system as fast as it is formed,
momentary expansion called the

Draper effect, because it was first noted by J. W. Draper. This phenomenon is
different from the Budde effect, for it is a secondary result of the heat liberated
during the reaction ; after a large number of Draper effects the amount of chemical
change is measurable. It is therefore probable that there is no real period of
inertness, but the amount of chemical action during the earlier stages of the reaction
is too small to be detected. The initial period occupied by the reaction in
assuming the steady state was called by R. Bunsen and H. E. Roscoe 3 the period
of photochemical induction :

At one time bodies are enabled to follow the s attraction of their chemical forces, whilst
at another time they are prevented from doing so by forces acting in an opposite direction.
These opposite attractions which must be overcome in order that the chemical combination
should take place, may be presented to the mind under the image of a resistance similar
to that occurring in friction, in the passage of electricity through conductors, in the dis-
tribution of magnetism in steel, or in the conduction of heat. We overcome this resistance
when we quicken the formation of a precipitate by agitation, or when, by increase of temp.,
by catalytic action, or by insolation, we cause a chemical action to occur. . . . The act by
which the resistance to combination is diminished, and the combining power thus brought
into greater activity, we called chemical induction ; and we specify this as photo-chemical,
thermo-chemical, electro-chemical, and ideo-chemical, according as light, heat, electricity,
or pure chemical action is the active agent concerned in overcoming the chemical resistance.

It was recognized that the initial period is characteristic of many other reactions
not dependent on light, and it was called the period of induction. Such a period
has been observed in the action of acids on zinc ; of nitric acid on copper ; of



150 INOKGANIC AND THEORETICAL CHEMISTRY

sulphurous on iodic acid ; of bromine on organic acids ; and M. Berthelot and
L. P. de St. Gilles also noticed what they called une acceleration initiate in the action
of acids on alcohol. J. W. Draper believed that the first action of light was to
induce a more active allotropic modification of the chlorine ; E. Pringsheim,
J. W. Mellor, P. V. Bevan believed that some intermediate compound is formed in
the initial stages of the reaction ; but C. H. Burgess and D. L. Chapman showed
that the system probably contains some impurity which must be destroyed by the
light before the chlorine and hydrogen can react.

There is no general cause for the period of induction applicable to all reactions.
Excluding disturbances arising from the imperfect mixing of the reacting substances,
from the heat of the reaction, and from the preliminary saturation of a liquid by
a gas before the speed of a reaction can be measured by the rate of evolution of a

tas, etc., three causes of a period of induction have been recognized. (1) The
nal products of the reaction are produced after certain intermediate products have
been formed. (2) The main reaction may be accelerated catalytically by the
products of the reaction e.g. A. V. Harcourt and W. Esson showed that the
manganese sulphate produced during the reaction between potassium permanganate
and oxalic acid in the presence of sulphuric acid accelerates the reaction. (3) The
overcoming of a passive resistance of some kind, or the destruction of a negative
or inhibitory catalyst as is supposed by C. H. Burgess and D. L. Chapman to occur
when light acts on a mixture of hydrogen and chlorine containing a trace of nitrogen
chloride.

A period of induction is characteristic of chemical reactions which take place
in a series of intermediate stages. This is a necessary consequence of the law of
mass action. The duration of the period depends on the relative magnitude of the
velocity constants of the intermediate reactions. For example, with the reaction
A=M=B, the rate of formation of the intermediate compound, M in the A M
reaction, will be quickest at the start, and the rate of formation of B by the destruction
of M in the M B reaction will be slowest at first, and increase with time as the
amount of M accumulates in the system. At first, during the period of acceleration,
the speed of the A=M reaction exerts a preponderating influence and M accumulates
in the system ; but the increasing speed of the M=B reaction gradually neutralizes
the effect of the first reaction, until the rate of formation of M by the A=M reaction is
equal to the rate of its destruction by the M=B reaction, and finally, the M is con-
sumed faster than it is formed. The amount of M in the system at any moment
thus determines the rate of formation of B, so that the curves showing (i) the rate
of formation of B, and (ii) the amount of M in the system at different moments, are
similar in shape. This is illustrated by the dotted line in Fig. 6. The duration of
the period of induction naturally depends upon the relative speeds of the two
reactions. If the rate of formation of the intermediate compound is immeasurably
fast, there will be no appreciable period of induction.

In the case of the hydrogen- chlorine reaction, the intermediate compound M,
by J. W. Draper's hypothesis, is allotropic or active chlorine ; by E. Pringsheim's
hypothesis, 4 chlorine monoxide or some analogous compound. Having shown that
the additive chlorine monoxide does not accelerate the reaction or abbreviate the
period of induction, an imaginary intermediate compound was postulated of a more
indefinite and vague form : ieCl 2 .2/H 2 .zH 2 . One naturally shirks vague hypotheses
of this type, but we are always confronted with the fact that the presence of the
third component, water, seems necessary for this reaction. There are also many
other reactions in which it seems necessary to assume either the formation of a
complex intermediate compound of this type, or else a sequence of consecutive
chemical reactions in which water plays an essential part. The reason for including
water in the formula of the imaginary intermediate compound is to indicate that the
formation of hydrogen chloride from the mixed gases, in light, seems to be dependent
on its presence. In the absence of water, the dry insolated gases unite with great
difficulty, if at all.



THE COMPOUNDS OF THE HALOGENS WITH HYDROGEN 151

If the formation of an intermediate compound be the source of the activity of
the chlorohydrogen mixture, this compound must be moderately stable in the
presence of chlorine and hydrogen gases because J. W. Draper found evidence that
it was not decomposed immediately the illuminated mixture of hydrogen and
chlorine is darkened. He says that he kept a pre-illuminated mixture of hydrogen
and chlorine in darkness for ten hours, and on re-exposure to light, the movement of
the liquid in the index tube commenced in a few seconds, whereas in the non-
illuminated mixture 600 seconds elapsed before any movement was visible. This
showed that the change which occurs in the mixture is not transient, but can persist
for some hours. R. Bunsen and H. E. Roscoe denied that previously insolated
chlorine instantly gives rise to hydrogen chloride on exposure to light ; P. V. Bevan
showed that J. W. Draper's observation is accurate, but the increased power so
acquired by chlorine is lost when the gas is passed through water. Whatever is
formed during the action in light is therefore either decomposed or washed out of
the gas when bubbled through water. R. Bunsen and H. E. Roscoe's failure to
verify J. W. Draper's observation was due to their having bubbled the gas through
water before it was tested. If X Q be the activity of a mixture of hydrogen and



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