Limitations of methods of osmometry: measuring the osmolality of biological fluids TERRENCE E SWEENEY AND CAROL A. BEUCHAT Department of Biology, University of Scranton, Scranton, Pennsylvania 1851 O-4625; Department of Biology, San Diego State University, San Diego 92182; and Hubbs-Sea World Research Institute, San Diego, California 92109 Sweeney, Terrence E., and Carol A. Beuchat. Limitations of meth- ods of osmometry: measuringthe osmolality of biological fluids. Am. J. Physiol. 264 (Regulatory Integrative Camp. Physiol. 33): R469-R480, 1993. -0smometry is an important tool in the investigation of biological phe- nomena, and commercially available instruments for freezing point and vapor pressure osmometrycan determinethe osmolality of solutions quickly and inexpensively. However, accurate measurements of osmolality using these techniques require that the solutions have specific characteristics, and that measurements do not exceed the limitations inherent to each method or instrument. The thermodynamic principles underlying osmometry con- strain the rangeand accuracyof eachmeasurement method, and thesemust be considered in establishing the usefulness of each technique. This paper addresses the principles and limitations of routine osmometry techniques. We begin by discussing definitions of osmolality and the thermodynamic concepts of solute-solvent systemsthat are central to understanding os- mometry of biological (i.e., aqueous) solutions. We then explore the appli- cation of various methods of measuringosmolality, the nature of errors introduced by overextension or misapplication of osmometry techniques, and the interpretation of data in the literature acquiredby various methods and protocols. concentration; water; chemical potential; activity; osmotic pressure; vapor pressure; colligative properties; freezing point depression OSMOLALITY describes the osmotic pressure of a solu- tion when a solute, such as a salt or a sugar, is added to a pure solvent such as water. Because osmotic pressure reflects fundamental aspects of the behavior of a solu- tion (i.e., its colligative properties), the measurement of osmolality has broad applicability in the biological and health sciences. Osmolality affects physiological pro- cesses that involve the movement of solutes or water across a membrane. The change in volume of red blood cells in a hypo- or hyperosmotic medium, the concen- tration of urine by the renal medulla, and changes in transvascular exchange during dehydration are typical examples. Consideration of osmolality is important in situations in which transmembrane differences in osmo- lality may produce unwanted or deleterious movements of water in response to an osmotic gradient, as in the preparation of cell culture media or solutions for paren- teral infusion, and in the determination of the proper composition of fixative solutions. Osmolality is critical to organisms that live in aqueous environments (e.g., fish, invertebrates, plants, bacteria) because these or- ganisms face the problem of water and solute exchange with a medium that can have an osmolality substantially higher or lower than their own. Osmolality was originally defined to provide an anal- ogy between the chemistry of solutions and that of gases as described by the ideal gas law. That is, just as molality is measured as moles of solute per kilogram of solution, osmolality would be measurable as osmoles of “osmoti- cally active” particles in the same manner. In biological texts, osmolality is now often defined as the total num- ber of dissolved solute particles in solution, where the nature of the particles, that is, their shape, weight, or charge, does not influence osmolality. However, the chemist van’t Hoff, who postulated the analogy between osmolality and the ideal gas law (19), realized that it was incorrect. Whereas the mole is mea- surable as a defined number of molecules, the osmole is not measurable as a defined number of particles. This is because the osmotic pressure of a solution that is not extremely dilute is not solely determined by the number of particles in the solvent; it is also a function of their size, shape, and charge, because these factors influence the solute’s behavior in solution. For example, a salt that dissociates into its ionic components when dissolved in solution creates a greater number of particles than the same molar amount of a nondissociating solute such as a sugar. However, adding more solute to a solvent does 0363-6119/93 $2.00 Copyright 0 1993 the American Physiological Society R469 at Ohio State University-HEA on November 6, 2012 http://ajpregu.physiology.org/ Downloaded from