Physiology > Gases > Intro
About this section: Gas exchange physiology will be covered in detail in lecture. There is also some good information in Campbell on this topic. However, I don't feel that Campbell provides enough background on the basic physics of gas exchange processes, so I made these pages on gas exchange physiology. We'll do a lab (more like a problem set) on gas exchange, and these pages should help you complete that. The information here should also help you on the midterm.
We couldn't live without gases. All eukaryotes (including plants!) must continually take in oxygen and eliminate carbon dioxide in order to live. In order to understand how gases influence biological processes (gas physiology), it's helpful to first review some basic gas physics.
Gas molecules in a closed container (animated gif). Image from Wikipedia article on the Kinetic Theory of Gases.
Gases consist of molecules that are widely spread out-- in other words, the distance between gas molecules is large compared to the size of the molecules themselves. The molecules fly around, occasionally colliding with each other or with solid objects (such as the walls of the container that holds the gas). The kinetic energy of the moving molecules is also called heat energy. The higher the temperature, the faster the molecules move.
According to the kinetic theory of gases, the collisions are elastic -- the molecules just collide and bounce off. Collisions with the walls of the container produce the pressure of the gas. Gas pressure is simply a measure of how often and how hard gas molecules collide with solid objects.
The chemical properties of the gas molecules aren't very important in determining the physical behavior of gases. Water vapor (H2O in gas form) behaves more or less like oxygen gas (O2), even though they are chemically very different.
Units of pressure
Pressure is defined as force/area. For a gas, the pressure comes from all those gas molecules colliding with a solid object. Pressure can be expressed in various units:
- Pascal (Pa): The SI unit of pressure is the Pascal (Pa; 1 Pa=N/m2), but Pascals are rarely used in physiology.
- Atmosphere (atm): a unit of pressure intended to represent the atmospheric air pressure on a typical day at sea level. Atmospheres are not widely used scientifically, but they're still conceptually useful for Bio 6A.
- Torr, or mm Hg: Pressure has traditionally been measured by the ability of a gas to push a column of mercury (Hg) upward into a tube; the greater the pressure, the higher the mercury is pushed upward. Blood pressure is typically measured in mm Hg, as are other physiologically relevant aspects of gas pressure.
- Conversion factors: 1 atmosphere = 760 torr = 760 mm Hg = 101,325 Pa
There are also other units of pressure; for an exhaustive list, see the Wikipedia article on Pressure.
The total pressure of air in a particular container (or in the atmosphere) is the sum of the pressures of all the individual gases that make up the air. The pressure of each gas is called its partial pressure. Since N2 makes up about 78% of the molecules in air, it accounts for 78% of the pressure. If the total pressure of atmospheric air is 1 atm, the partial pressure of N2 is 0.78 atm. This principle is called Dalton's law.
Gases diffuse down their concentration gradients
Gas molecules move in random directions, depending on what they crash into. Any individual molecule of gas has an equal probability of moving in any direction. The net effect of the random movements of gas molecules is that the molecules end up being randomly dispersed throughout their container. This spreading out is called diffusion.
Suppose you have a closed container, and you carefully place some gas molecules (represented by purple dots) at one end, as shown here.
At first, there is a strong concentration gradient. Concentration gradient is defined as the difference in concentration between two locations, divided by the distance between the locations. A large difference in concentration over a short distance is a strong concentration gradient.
The gas molecules begin to move randomly due to heat energy. Each individual molecule is equally likely to move in any direction. Statistically, however, there is a very high probability of gas molecules moving away from an area of high concentration (the left side of the container) and toward an area of lower concentration (the right side). In other words, the gas molecules will tend to move down their concentration gradient.
After diffusion proceeds for a while, the molecules are randomly distributed; there is no more concentration gradient. The individual molecules still move just as fast, but there is no net movement in any direction.
In general, we can assume that biologically important gases behave like ideal gases -- in other words, the individual molecules of the gas are spread out and don't interact with each other. (This isn't true for some non-biological situations, such as gases at very high pressures or with very large gas molecules, such as refrigerants).
Ideal gases obey the ideal gas law:
- P = pressure of the gas
- V = volume
- n = the number of gas molecules, usually given in moles
- R = the gas constant
- T = temperature
This is one of the best laws in physics -- simple and powerful. For the purposes of Bio 6A, you won't need to do any calculations with this, and the ideal gas law can be simplified in a couple of useful ways. For example, suppose the number of gas molecules (n) and the temperature (T) remain constant. Of course, the gas constant (R) also remains constant, so therefore,
PV = constant
Or, if you increase the pressure, you decrease the volume:
For a closed container of gas, if the pressure is doubled, the volume must be cut in half. This part of the ideal gas law is called Boyle's Law.
Similarly, if P, n, and R remain constant, there is an inverse relationship between temperature and volume:
For a closed container of gas, if the absolute temperature (in Kelvin) is doubled, the volume must also be doubled. This part of the ideal gas law is called Charles's Law.
The gas laws are great, but in biology, gases such as O2 and CO2 aren't always in gas form. The most important properties of gases happen when they are dissolved in aqueous solutions.
This page updated September 20, 2011