The properties of a solvent containing a solute are different from those of a pure solvent. The properties that depend on the number of particles dissolved and not on their chemical identity and nature (kind, size, and charge) are known as colligative properties. Some of these colligative properties are lowering the vapor pressure, depression of the freezing point, the elevation of the boiling point, and creating osmotic pressure.
Considering that osmosis and osmotic pressure are the most important for medical students, special emphasis is placed on this colligative property.
1. Osmosis and Osmotic Pressure of Solutions
When a solution and its pure solvent are separated by a semipermeable membrane (one through which the solvent can pass but not the solute), the pure solvent will diffuse through the membrane and dilute the solution. This process of transferring water molecules through a semipermeable membrane is called osmosis. The pressure required to stop the osmosis from a pure solvent into a solution is called the osmotic pressure, π, of the solution. The osmotic pressure of a dilute solution can be calculated from the expression:
πosm = CM · R · T · 1000, (2.1)
where CM is the molar concentration of the solute, mol/L; R is the gas constant, R = 8.31 J/mol·K;
T is the temperature of the solution on the Kelvin scale, K;
T = 273 + t°.
1000 we use to get [SI] units of pressure, Pa.
Note. Notice that CM units [mol/L] = [mol/dm3], and we need [mol/m3] instead, that is why 103 is included in formula (2.1).
This means that the relationship between osmotic pressure, π, and the molarity of the solute particles is similar to the ideal gas law:
πosm = C · R · T = (ν/V)RT => πV = ν · R · T (Van’t Hoff’s law).
If the solute is an electrolyte, we have to remember, that the number of particles may be doubled (e.g., NaCl ↔ Na++ Cl–) or tripled (CaCl2 ↔ Ca2+ + 2 Cl¯),… so it is necessary to use coefficient “i” (i ≈ 2,3,…):
πosm = i · CM · R · T · 1000. (2.2)
Solute concentrations are particularly important when solutions are injected into the body. The body cell fluids and blood serum have the following normal value of osmotic pressure:
7.7 atmospheres ≈ 7.7 · 105 Pa =770 kPa
Solutions injected into the body must have the same osmotic pressure as blood serum; that is, they are supposed to be isotonic with blood serum. (e.g., 0.9% sol. NaCl or 5% dextrose = glucose).
If a less concentrated solution, a hypotonic solution (like 0.1% NaCl or 2% dextrose), is injected in sufficient quantity to dilute the blood serum, water from the diluted serum will pass into the blood cells by osmosis, causing the cells to expand and rupture (hemolysis). We never use such solutions intravenously.
If a more concentrated solution, a hypertonic solution (like 10% NaCl or 20% dextrose), is injected, the cells will lose water to the more concentrated solution, shrivel, and possibly die (plasmolysis). These solutions may be used intravenously in a very small amount or they are diluted by isotonic ones listed above. (For example, 25% MgSO4 will be dissolved in 200 mL of 0.9% NaCl for intravenous infusion, or 5 mL of this solution will be injected). Most of the solutions, which ambulance doctors use, are hypertonic. But it is known, that plasmolysis is a reversible process, and it is not that dangerous compared with irreversible hemolysis. Some cells will change their shape because of crenation, but recover later.
Physiologists use the osmole and a related term osmolarity (Cosm) to calculate the number of osmotic active particles. For example, Cosm of the physiological solution that is 0,15 M (0,9%) NaCl will be twice the molarity since 1 mole of NaCl yields 2 moles of ions in solution (1NaCl ↔ 1Na++ 1Cl¯; i = 2):
0.15 mol/L (NaCl) = (0.15 mol/L × 2) Osm/L = 0.30 Osm/L = 300 mosm/L.
The normal value of osmolarity for blood: 290–310 mosm/L (≈ 0.3 Osm/L).
It is rather convenient to use osmolarity instead of osmotic pressure because we can predict the state of the cells in such a solution so easily.
Сosm = СM · i (2.3)
[Osm/L, mosm/L].
The effects of osmosis are particularly evident in biological systems, since the cells are surrounded by semipermeable membranes. Carrots and celery that have become limp due to loss of water to the atmosphere can be made crisp again by placing them in water. Water moves into the carrot or celery cells by osmosis. And vice versa, a cucumber placed in a concentrated salt solution loses water by osmosis and becomes a pickle due to diffusion of salt, also most microorganisms die because their cells lose water to the more concentrated salt solution (plasmolysis). A similar occurrence happens when you prepare a jam using a certain amount of sugar.
We use some hypertonic solutions as laxatives (MgSO4·7H2O, Na2SO4·10H2O), for hypertonic bandages (10% NaCl), in enemas, and so on.
The laxative action is based on the rules of osmosis: these salts are not digestible and create hyper-concentration in the intestine; as a result, water moves inside, into the bowels, passes through the intestinal walls in order to dilute this concentrated solution, stimulating laxative effect.
To clean the festering wounds the surgeons recommend some hypertonic bandages: all the rot products (the pus) move outside, trying to dilute the hypertonic solution, which was put on a gauze bandage.