Hydrogen Isotope with Mass 2

G. S. Landsberg, Moscow

The January issue of Physical Review contains a brief communication, a letter to the editor, signed by Urey, Brickwedde, and Murphy,* describing experiments which apparently establish, rather reliably, the existence of an isotope of hydrogen with a mass close to 2.

The question of the existence of such an isotope is not new. As early as 1920 Rutherford, from observations of the artificial disintegration of nitrogen and oxygen nuclei, considered it possible to conclude that there existed a formation consisting of two protons and one electron, united into a compact whole having the dimensions of atomic nuclei. Such a formation could be the nucleus of an isotope of hydrogen. However, the supposition that such an isotope existed rested on observations that were thought to prove the presence of an atom with mass 3 and two outer electrons, i.e., an isotope of helium. Rutherford regarded the nucleus of this atom as containing the nucleus of an isotope of hydrogen. Thus the existence of the latter was proved in a very indirect way. Moreover, subsequent observations cast doubt on the existence of the helium isotope as well, and therefore the problem of hydrogen isotopy that interests us could in no way be considered solved. Later investigations by Aston and others made it possible to refine the measurements of the masses of individual atoms to such an extent that determinations of atom-

* H. C. Urey, F. G. Brickwedde, G. M. Murphy, Phys. Rev. 39, 164, 1932.

of weight by the mass-spectrograph method proved capable, in accuracy, of competing with the methods of the most precise chemical determinations. At the same time Aston’s method gives, as is known, the exact value of the atomic weight of individual isotopes, whereas chemical determinations refer to average values characterizing those mixtures of isotopes which constitute “chemically pure” substances. Therefore, when comparing the results of the one method with those of the other, we encounter a characteristic difficulty.

Determinations of atomic weights, both by the chemical method and by Aston’s method, are relative determinations, in which the mass of the atom under study is compared with the mass of the oxygen atom, conventionally taken to be equal to an integer, namely 16. If oxygen had no isotopes, as had been assumed on the basis of Aston’s measurements, then the conventional unit of both systems would be the same and comparison of the results obtained by the two methods could be made by a simple juxtaposition of the numbers. However, in recent times* studies of the band spectrum of oxygen have led to the undoubted conclusion that oxygen, besides \(O^{16}\), exists also in the form of two isotopes \(O^{18}\) and \(O^{17}\), the latter being present in exceedingly small quantities. The method of studying band spectra proved not only more sensitive than Aston’s method in the question of detecting isotopes, but it even made it possible to make reliable determinations of the relative weights of the oxygen isotopes.

Taking the weight of the principal isotope as 16.000, Mecke and Wurm determined the weight of the next isotope in magnitude to be \(17.981 \pm 0.010\).** Thus, proceeding from Aston’s unit of atomic weights (\(O^{16} = 16.000\)), we must estimate the weight of chemical oxygen, i.e. of the mixture in which, along with \(O^{16}\), \(O^{18}\) and \(O^{17}\) occur, by a somewhat larger number, which can be determined by knowing the percentage content of all the isotopes in oxygen. Comparison of the two methods of determination

* W. F. Giauquè and H. W. Johnston, Nature 123, 318, 1929.
** R. Mecke u. K. Wurm, Z. Physik 61, 37, 1930.

of atomic weights—chemical and physical—becomes, consequently, possible only in the case where the ratio of the isotopes in the mixture has been precisely established.

A good deal of attention has been devoted to this important task in recent years. Soon after the discovery of the isotopy of oxygen, Babcock* attempted to determine this ratio by comparing the intensity of the lines of two oxygen bands \(A\) and \(A'\), belonging to the molecules \(O^{16}\cdot O^{18}\) and \(O^{16}\cdot O^{16}\). His estimates gave \([O^{16}]:[O^{18}]=1250:1\). Later, Naudé** repeated the same attempt, studying the intensity of the lines of the band spectrum of NO, whose complex structure is due to the presence of three isotopes of oxygen and two isotopes of nitrogen \(N^{14}\) and \(N^{15}\). Naudé’s measurements led him to the figures: \([O^{16}]:[O^{18}]:[O^{17}]=1075:1:0.12\). Finally, Mecke and Childs*** again undertook a careful quantitative study of the intensity of the lines of the band spectrum of oxygen and arrived at somewhat different and, apparently, more reliable results. According to their measurements
\[ [O^{16}]:[O^{18}]:[O^{17}]=(630\pm20):1:0.2. \]

If, proceeding from the last ratio, one calculates the atomic weight of chemical oxygen in Aston units, one obtains the value 16.0035. Thus, between the chemical unit of atomic weights and the physical unit the relation is established
\[ A_{\mathrm{phys.}}=A_{\mathrm{chem.}}\cdot1.00022. \tag{a} \]

The indicated relation apparently agrees quite well with a whole series of measurements that make it possible to compare the chemical and physical methods of comparing atomic weights for a number of elements.

Thus, for carbon we have \(C_{\mathrm{chem.}}=12.0025\). According to Aston’s determinations, \(C=12.0036\) (if \(O^{16}=16.0000\) is taken as the unit). Recalculating by means of relation (a), we find \(C=12.0010\). In order to reconcile this value with the chemical determination, it is necessary to admit the existence of the carbon isotope \(C^{13}\), at a quantit—

* H. D. Babcock, Proc. Nat. Acad. U. S. A. 15, 471, 1929.
** S. M. Naudé, Phys. Rev. 36, 333, 1930.
*** R. Mecke u. W. H. J. Childs. Z. Physik 68, 362, 1931.

in the ratio \([C^{12}] : [C^{13}] = 650 : 1\). Spectroscopic data do indeed indicate the existence of two such isotopes, with a rough estimate of the concentration \(500 : 1\).*

For nitrogen we have: \(N_{\text{chem.}} = 14.008\); according to Aston \(N = 14.008\). Recalculation by relation \((a)\) gives: \(N = 14.0042\), and agreement with the chemical determination presupposes the presence of the isotope \(N^{15}\), with the ratio \([N^{14}] : [N^{15}] = 320 : 1\). The already-mentioned investigations by Noda, with the Mecke and Childs correction, give, from spectroscopic data, the presence of nitrogen isotopes \([N^{14}] : [N^{15}] = 400 : 1\). For He we have, according to chemical determinations, \(4.0018\), according to Aston: \(4.00216\), with recalculation by \((a)\) \(4.0013\)—a value within the limits of observational error indistinguishable from the chemical one.

Analogous results are obtained for F: chemical \(18.995\), Aston—\(19.000\), with recalculation \(18.996\)—there is no isotope. The same for I, etc.

Of particular interest is such a comparison for hydrogen.

According to chemical-analysis data \(H = 1.00777 \pm 0.00002\). According to Aston’s measurements: \(H = 1.00778 \pm 0.00015\). Taking relation \((a)\) into account, \(H = 1.00756\) is found, a value different from that measured chemically. The unusual accuracy of both the chemical and the physical determinations of the weight of the hydrogen atom makes it possible to regard this discrepancy as real. For agreement of the results one would have to suppose the presence of a hydrogen isotope with mass close to 2, admixed to the principal isotope in negligible quantities, namely:
\([H^{1}] : [H^{2}] == 4500 : 1\).

Precisely such a hypothesis was put forward by Birge and Menzel,** who suggested that the establishment of the indicated isotope could also be obtained by the method of band spectra. Taking into account the simple character of the hydrogen spectrum and the exceedingly large difference in the nuclear masses of the two presumed isotopes (dif-

* A. S. King and R. T. Birge, Astroph. J. 72, 19, 1930.
* R. T. Birge and D. H. Menzel, Phys. Rev. 37*, 1671, 1931.

difference by a factor of two) one might have hoped to detect the isotope of interest to us by direct observation of the Balmer spectral series. Indeed, the phenomenon of nuclear motion makes it possible, by Bohr’s well-known method, to calculate that the difference in the wavelengths of the lines of the Balmer series of the two isotopes should amount to from 1 to 2 Å (see the table given below). This quantity can easily be detected with the aid of a diffraction grating, for example. However, photographing the hydrogen spectrum with such a grating (a concave diffraction grating) of radius 21 feet, with a dispersion of 1.31 Å per mm, carried out by Urey and his collaborators, showed that, although traces of lines do appear in the expected positions, the ratio of intensities of the principal and secondary lines is so unfavorable that it is difficult to be certain of the reality of the observed lines. This ratio of intensities corresponds, incidentally, to the assumed concentration ratio. Thus there arises an urgent need to enrich hydrogen with the heavier isotope in order to make the spectral conclusions more certain. Urey, Brickwedde, and Murphy chose, for such enrichment, the method of prolonged evaporation of liquid hydrogen. As calculations based on quite probable assumptions show, the vapor pressure of solid unevaporated hydrogen will depend on the molecular weight of the hydrogen molecules as follows: \(p_{11}:p_{12}=1:0.37\), where \(p_{11}\) is the pressure above solid hydrogen composed of molecules \(\mathrm{H}'\mathrm{H}'\), and \(p_{12}\) is that above hydrogen consisting of molecules \(\mathrm{H}'\mathrm{H}^2\) (it goes without saying that the probability of molecules of the type \(\mathrm{H}^2\mathrm{H}^2\) may be regarded as practically negligible). Thus, as a result of prolonged evaporation, one may expect to obtain the final portion richer in the less mobile component. The experiment was carried out with two samples. One contained 6 l of liquid hydrogen evaporating at normal pressure. The second consisted of 4 l of liquid hydrogen, boiling under a pressure of several mm, i.e. at a temperature close to the solidification temperature of hydro-

of hydrogen. The last portion of the last cubic centimeter is introduced into the discharge tube, and the spectrum of the gas was studied with the aid of the aforementioned diffraction grating (with a dispersion of \(1.31\ \text{Å}/\text{mm}\)).

The study of the first sample gave no improvements in comparison with ordinary hydrogen; traces of the isotope lines could be obtained only by increasing the exposures, relative to normal, approximately 4500-fold, i.e., near the strongly overexposed lines of the principal isotope and, consequently, rather uncertainly. The study of the second sample, however, gave a quite noticeable (about fivefold) increase in the intensity of the lines of the presumed isotope. This increase in intensity not only facilitates the ascertainment and measurement of the sought lines, but is in itself an important argument in favor of an isotopic interpretation of these lines. For the intensity of accidental lines of the diffraction grating (“ghosts”) could not depend on the history of the hydrogen introduced into the tube. Likewise, lines of the molecular spectrum, whose intensity changes when the conditions of excitation of the tube are changed, can thereby be distinguished from the lines under study. Thus, in the experiment indicated, we have rather reliable proof of the isotopic origin of the observed lines. Measurements of their wavelengths reveal good agreement with theoretical calculations, as the following table shows.

In outward appearance the lines of \(\mathrm{H}^2\) seem somewhat less broadened than the lines of \(\mathrm{H}^1\), which could be explained by a smaller Doppler effect, since the difference in the masses of the ato-

of atoms is sufficiently large. The principal part of the broadening arises, however, from the fact that the hydrogen lines belong, as is known, to the type of doublet lines (as a consequence of the electron spin). These doublets, generally speaking, were not resolved by the grating used, with the exception of one line, \(H^{2} a\), which was resolved into a doublet with a component separation of \(0.16\) Å, in agreement with the same value found for \(H^{1}\).

Despite the considerable difficulty of the measurements described and the certain uncertainty in the interpretation of the results connected with it, nevertheless, on the basis of the totality of the data, the existence of the isotope \(H^{2}\) in quantities of about \(1/4000\) of the isotope \(H^{1}\) may apparently be considered proven, in accordance with the assumptions of Birge and Menzel.

The measurements described do not make it possible to determine the atomic weight of the new isotope with the same accuracy as was done for the principal one. But, of course, its small concentration shows that all conclusions concerning the mass defect in the packing of protons into more complex nuclei remain unchanged.

Direct confirmations of the discovery of the isotope \(H^{2}\) were not slow to follow. Walker Bleakney* made a successful attempt to analyze hydrogen by means of a mass spectrograph. The difficulty of the investigation is increased by the circumstance that an ion of mass 3 may have a dual origin: \((H^{1}.H^{2})+\) and \((H^{1}.H^{1}.H^{1})+\). There is, however, a possibility of recognizing these two ions by observing the dependence of their concentrations on pressure. The ion \((H^{1}.H^{2})+\) is formed from molecules \((H^{1}.H^{2})\), and its concentration is proportional to the concentration of these molecules, i.e. to the pressure of hydrogen. The ion \((H^{1}.H^{1}.H^{1})+\), on the other hand, is formed as a result of the collision of two molecules \((H^{1}.H^{1})\), i.e. its concentration is proportional to the square of the pressure of hydrogen and must decrease rapidly as the pressure is lowered. By studying the dependence of the concentration of the ion of mass 3 on the pressure of hydrogen in the instrument, one may hope to distinguish with which

* Walker Bleakney, Phys. Rev. 39, 535, 1932.

ion we are dealing with. Bleakney studied in his apparatus (in which reliable measurements could be made when the pressure was varied from \(10^{-5}\) to \(10^{-6}\) mm), on the one hand, ordinary hydrogen, and on the other, hydrogen enriched with the isotope by means of evaporation. A portion of such hydrogen was prepared for these experiments by Brickwedde, who also took part in Urey’s experiments. The investigation fully confirmed that the hydrogen obtained as a result of evaporation is enriched with the isotope \((\mathrm{H}^{2})\), and Bleakney estimates the ratio of the concentrations of the two isotopes at 1100 (with a possible error of 10%), i.e., it agrees with the estimate made by the spectroscopic method \(\left(\dfrac{1}{5}\cdot 4500 = 900\right)\).

In conclusion it should be noted that the investigation mentioned, like a number of others referred to in the present note, has revealed the power of the spectroscopic method in this most difficult problem of studying isotopes. In sensitivity, accuracy, and the possibility of quantitatively estimating the percentage ratio of isotopes, the spectroscopic method not only is not inferior to, but apparently may even surpass, the famous Thomson–Aston method*.

* Note added in proof. While this note was in press, a detailed communication by Urey, Brickwedde and Murphy appeared in Phys. Rev.