100 words about diffusion
Answers
diffusion, process resulting from random motion of molecules by which there is a net flow of matter from a region of high concentration to a region of low concentration. A familiar example is the perfume of a flower that quickly permeates the still air of a room.
Heat conduction in fluids involves thermal energy transported, or diffused, from higher to lower temperature. Operation of a nuclear reactor involves the diffusion of neutrons through a medium that causes frequent scattering but only rare absorption of neutrons.
Light molecules have higher average speeds than do heavy molecules at the same temperature. This result follows from kinetic theory, as explained below, but it can also be seen by noting that the speed of sound is greater in a light gas than in a heavy gas. This is the basis of the well-known demonstration that breathing helium causes one to speak with a high-pitched voice. If a light and a heavy gas are interdiffusing, the light molecules should move into the heavy-gas region faster than the heavy molecules move into the light-gas region, thereby causing the pressure to rise in the heavy-gas region. If the diffusion takes place in a closed vessel, the pressure difference drives the heavy gas into the light-gas region at a faster rate than it would otherwise diffuse, and a steady state is quickly reached in which the number of heavy molecules traveling in one direction equals, on the average, the number of light molecules traveling in the opposite direction. This method, called equimolar countercurrent diffusion, is the usual manner in which gaseous diffusion measurements are now carried out.
The steady-state pressure difference that develops is almost unmeasurably small unless the diffusion occurs through a fine capillary or a fine-grained porous material. Nevertheless, experimenters have been able to devise clever schemes either to measure it or to prevent its development. The first to do the latter was Graham in 1831; he kept the pressure uniform by allowing the gas mixture to flow. The results of this work now appear in elementary textbooks as Graham’s law of diffusion. Most of these accounts are incorrect or incomplete or both, owing to the fact that the writers confuse the uniform-pressure experiment either with the equal countercurrent experiment or with the phenomenon of effusion (described below in the section Kinetic theory of gases). Graham also performed equal countercurrent experiments in 1863, using a long closed-tube apparatus he devised. This sort of apparatus is now usually called a Loschmidt diffusion tube after Loschmidt, who used a modified version of the tube in 1870 to make a series of accurate diffusion measurements on a number of gas pairs.
A quantitative description of diffusion follows. A composition difference in a two-component gas mixture causes a relative flow of the components that tends to make the composition uniform. The flow of one component is proportional to its concentration difference, and in an equal countercurrent experiment this is balanced by an equal and opposite flow of the other component. The constant of proportionality is the same for both components and is called the diffusion coefficient, D12, for that gas pair. This relationship between the flow rate and the concentration difference is called Fick’s law of diffusion. The SI units for the diffusion coefficient are square metres per second (m2/s). Diffusion, even in gases, is an extremely slow process, as was pointed out above in estimating molecular sizes and collision rates. Gaseous diffusion coefficients at one atmosphere pressure and ordinary temperatures lie largely in the range of 10-5 to 10-4 m2/s, but diffusion coefficients for liquids and solutions lie in the range of only 10-10 to 10-9 m2/s. To a rough approximation, gases diffuse about 100,000 times faster than do liquids.
Diffusion coefficients are inversely proportional to total pressure or total molar density and are therefore reported by convention at a standard pressure of one atmosphere. Doubling the pressure of a diffusing mixture halves the diffusion coefficient, but the actual rate of diffusion remains unchanged. This seemingly paradoxical result occurs because doubling the pressure also doubles the concentration, according to the ideal gas equation of state, and hence doubles the concentration difference, which is the driving force for diffusion. The two effects exactly compensate.