HYDROGEN PEROXIDE. 103- Hydrogen Peroxide.
The oxidation of water to hydrogen peroxide takes place with absorption of
heat:
2 H 2 O (liquid) + O 2 + 44,320 cal. = 2 H 2 O 2 (liquid).
Hydrogen peroxide is thus endothermic as regards its formation from water
and oxygen. In accordance with the relation which exists between the equi-
librium constants of the mass-action law and the temperature (cf. p. 47), it
follows that the quantity of hydrogen peroxide in an equilibrium mixture must
increase with rise of temperature, and that therefore in order to obtain hydro-
gen peroxide synthetically by the above reaction it is necessary to work at as
high a temperature as possible. Even at 2000° there is but little peroxide pres-
ent in the equilibrium mixture.^1 At low temperatures hydrogen peroxide is
not stable at any appreciable concentration; in fact a liquid containing a high
percentage of it is explosive. The velocity at which hydrogen peroxide decom-
poses, when in the region of instability as regards temperature or concentration,
increases (as does the velocity of all reactions) with rise of tempera-
ture.^2 If then it is desired to obtain the hydrogen peroxide produced by a
reaction at a high temperature, it is necessary to cool the reaction products
very rapidly to a point where the decomposition velocity is inappreciable.
When thus chilled the hydrogen peroxide continues to exist, as it were, in a
supercooled condition. This end is accomplished, for example, when an oxy-
hydrogen flame comes in direct contact with a piece of ice and thereby an
extremely sudden drop in temperature is brought about. Cf. Cyanogen,
No. 59, and Acetylene, No. 64.
The principle of preserving the equilibrium concentrations, as they exist at
high temperatures, by means of sudden cooling was first introduced by Deville
in 1863 and applied in the construction of his " hot and cold tubes." A nar-
row silver tube cooled by running water was placed in the center of a white-hot
porcelain tube at the walls of which the gas-reaction to be measured reached its
high-temperature equilibrium; the reaction-products on coming in contact
with the inner tube were chilled, and thus prevented, partially at least, from
undergoing the reverse reaction. In recent years many equilibria, in which
concentrations of measurable magnitude are reached only at high tempera-
tures, have been studied in this manner. The mixtures are cooled, without
suffering change in concentration, to temperatures at which analytical meas-
urements are possible. The quantitative preservation of the concentration
fails, however, in the case of hydrogen peroxide on account of its great decom-
position-velocity.
The velocity of many reactions can be increased by means of catalyzers, as
well as by rise of temperature. The decomposition of hydrogen peroxide,
since it takes place with conveniently measurable rapidity at ordinary tem-
(^1) At 2000° according to Nernst (1903) less than 1% H
3 O 2 exists in a mixture
of water vapor and oxygen, each at 0.1 atmosphere pressure.
! In order to keep the temperature low during the distillation of hydrogen
peroxide, it is customary to work in a vacuum.