3 8o 



NATURE 



{August 1 6, 1888 



Christison, who devised a simple method of demonstrating the 

 fact. 



So long as the evidence in favour of the existence of gases in 

 the blood was so uncertain, the combustion theory of respiration 

 held its own. At last, in 1836, appeared the researches of 

 Heinrich Gustav Magnus, latterly Professor of Physics and 

 Technology in the University of Berlin. He first attempted to 

 drive off carbonic acid from the blood by a stream of hydrogen, 

 and thus obtained as much as 34 cubic centimetres of carbonic 

 acid from 62*9 cubic centimetres of blood. He then devised a 

 mercurial air-pump, by which it was possible to exhaust a re- 

 ceiver to a much greater extent than could be done by the ordinary 

 air-pump. When blood was introduced into such a vacuum, 

 considerable quantities of carbonic acid, oxygen, and nitrogen 

 were obtained. This research marks an epoch in physiological 

 discovery, as it threw a new light on the function of respiration 

 by demonstrating the existence of gases in the blood. 



In order to appreciate the value of this evidence, and the 

 method employed, let me direct your attention to the laws re- 

 gulating the diffusion of gases. As a mass of gaseous matter 

 has no independent form, like that of a solid body, nor a fixed 

 volume like that of a liquid, but consists of an enormous number 

 of molecules which, in consequence of their mutual repulsions, 

 endeavour more and more to separate from each other, it is 

 easy to see that if two masses of gas are brought into contact, 

 they will mix — that is, their molecules will interpenetrate, until 

 a mixture is formed containing an equal number of the molecules 

 of each gas. The force by which the molecules repel each other, 

 and by which they exercise pressure in all directions, is known 

 as the pressure or tension of the gas. It is evident that the 

 greater the number of gas molecules in a given space, the greater 

 will be the tension of the gas, and from this it follows that the 

 tension of a gas is in the inverse proportion to its volume (this is 

 known as Boyle's law). Suppose now that two gases are 

 separated by a porous partition ; the two gases will mix, and 

 the rapidity of the diffusion will vary according to the specific 

 weight of the gases. Thus light gases, like hydrogen or coal- 

 gas, will diffuse more quickly than air, or chlorine, or carbonic 

 acid. 



It is important also to note the laws regulating the absorption 

 of gases by fluids. If we allow a little water to come into con- 

 tact with ammonia gas above mercury, the gas is rapidly 

 absorbed by the water (1 volume of water absorbs 730 volumes 

 NH 3 ) all the gas above disappears, and in consequence of this 

 the pressure of outer air drives up the mercury in the tube. The 

 higher the temperature of the fluid the less gas it absorbs. At 

 the boiling-point of the fluid its absorption is = o, because at 

 that temperature, the fluid itself changes into gas. The power of 

 absorption of different fluids for the same gas, and the absorptive 

 power of the same fluid for different gases fluctuates between 

 wide limits. Bunsen defined the coefficient of absorption of a 

 fluid for a gas as that number which represents the volume of 

 gas (reduced to o° and 760 mm. barometric pressure) which is 

 taken up by 1 volume of the fluid. Thus 1 volume of distilled 

 water takes up the following volumes : — 



Again, 1 volume of distilled water at 0° C. absorbs 0*00193 

 volumes of hydrogen, while it can take up no less than 1180 

 volumes of ammonia ; again, I volume of water at o° C. absorbs 

 only 0-2563 volumes of olefiant gas, but I volume of alcohol, at 

 the same temperature, will take up as much as 3-595 volumes. 

 The volume of gas absorbed is independent of the pressure, and 

 the same volume of gas is always absorbed whatever the pressure 

 may happen to be. But as according to Boyle's law the density 

 of a gas, or in other words the number of molecules in a given 

 space, is in proportion to the pressure, and as the weight is 

 equal to the product of the volume and the density, so while the 

 volume absorbed always remains the same,- the quantity or 

 weight of the absorbed gas rises and falls in proportion to the 

 pressure (this is the law of Dalton and Henry). It therefore 

 follows that a gas is to be considered as physically absorbed 

 by a fluid, if it separates from it not in volumes but in 

 quantities, the weights of which are in proportion to the fall 

 of pressure. 



When two or more gases form an atmosphere above a fluid, 



the absorption takes place in proportion to the pressure which 

 each of the constituents of the mixture would exercise if it were 

 alone in the space occupied by the mixture of gases, because, ac- 

 cording to Dalton's law, one gas does not exercise any pressure 

 on another gas intermingled with it, but a space filled with one 

 gas must be considered, so far as a second gas is concerned, as a 

 space containing no gas, or in other words a vacuum. This 

 pressure, which determines the absorption of the constituents of 

 a gaseous mixture, is termed, according to Bunsen, the partial 

 pressure of the gas. The partial pressure of each single gas in 

 a mixture of gases depends, then, on the volume of the gas 

 in question in the mixture. Suppose atmospheric air to be 

 under a pressure of 760 mm. of mercury, then, as the air 

 consists of 21 volumes per cent, of O and 79 volumes per 



76o X 21 



cent, of N, '- = 1596 mm. of mercury, will be the 



100 



partial pressure under which the oxygen gas is absorbed, 



while the absorption of nitrogen will take place under a pres- 



?6o X 70 



sure of '- '-? = 600 mm. of mercury. Suppose, again, that 



100 

 above the fluid containing a gas, say carbonic acid, which has 

 been absorbed, there is an atmosphere of another gas, say at- 

 mospheric air, then as carbonic acid exists in the air only in 

 traces, its tension is equal to zero, and carbonic acid will escape 

 from the fluid until the difference of tension between the carbonic 

 acid in the water and the carbonic acid in the air above it has 

 been balanced — that is, until the carbonic acid which has escaped 

 into the air has reached a tension equal to that of the gas still 

 absorbed by the fluid. By the phrase " tension of the gas in a 

 fluid" is understood the partial pressure in millimetres of mer- 

 cury which the gas in question has to exercise in the atmosphere, 

 when no diffusion between the gas in the fluid and the gas in 

 the atmosphere takes place. 



The method followed by Magnus will now be understood. By 

 allowing the blood to flow into an exhausted receiver surrounded 

 by hot water, gases were set free. These were found to be oxy- 

 gen, carbonic acid, and nitrogen. He further made the important 

 observation that both arterial and venous blood contained the 

 gases, the difference being that in arterial blood there was more 

 oxygen and less carbonic acid than in venous blood. Magnus 

 concluded that the gases were simply dissolved in the blood, and 

 that respiration was a simple process of diffusion, carbonic acid 

 passing out and oxygen passing in, according to the law of 

 pressures I have just explained. 



Let us apply the explanation of Magnus to what occurs in 

 pulmonary respiration. Venous blood, containing a certain 

 amount of carbonic acid at the temperature of the blood and 

 under a certain pressure, is brought to the capillaries, which are 

 distributed on the walls of the air- vesicles in the lungs. In these 

 air-vesicles, we have an atmosphere at a certain temperature and 

 subject to a certain pressure. Setting temperature aside, as it 

 may be assumed to be the same in the blood and in the air-cells, 

 let us consider the question of pressure. If the pressure of the 

 carbonic acid in the blood be greater than that of the carbonic 

 acid in the air-cells, carbonic acid will escape until an equi- 

 librium is established between the pressure of the gas in 

 the blood and the pressure of the gas in the air-cells. Again, 

 if the pressure or tension of the oxygen in the air-cells be 

 greater than that of the oxygen in the venous blood, oxygen 

 will be absorbed until the tensions become equal. This 

 theory has no doubt the merit of simplicity, but it will 

 be observed that it depends entirely on the assumption that 

 the gases are simply dissolved in the blood. It was pointed 

 out by Liebig that, according to the experiments of Regnault and 

 Reiset, animals used the same amount of oxygen when breathing 

 an atmosphere composed of that gas alone as when they breathed 

 ordinary air, and that the vital processes are not much affected 

 by breathing the atmosphere of high altitudes where the amount 

 of oxygen taken in is only about two-thirds of that existing at 

 the sea level. It was also shown at a much later date, by Ludwig 

 and W. Muller, that animals breathing in a confined space of air 

 will use up the whole of the oxygen in the space, and it is clew 

 that as the oxygen is used up the partial pressure of the oxygen 

 remaining must be steadily falling. Liebig urged the view that 

 the gases were not simply dissolved in the blood, but existed in a 

 state of loose chemical combination which could be dissolved by 

 the diminished pressure in the vacuum, or by the action of other 

 gase^. He also pointed out the necessity of accurately deter- 

 mining the coefficient of absorption of blood for the gases — that 

 is, the amount absorbed under a pressure 0^760 mm. of mercury 



