Osmolarity and Tonicity

Osmolarity, Tonicity, and Equivalents Study Guide & Practice

1. Osmolarity

Recall that the osmotic pressure of a solution is created by solutes dissolved in a

solvent, and the dissolved solutes create a ‘pull’ for water to move INTO that

solution by osmosis.

Osmolarity allows

us to determine

the concentration

of multiple solutes

in a solution.

Osmolarity is simply another measure of concentration, like molarity. (See

Molarity Study Guide Handout for more on molarity.) However, unlike molarity

which tells us the concentration of one type of molecule, osmolarity allows us to

look at the concentration of multiple solutes dissolved in a solution. By now, you

can certainly appreciate that body fluids have many more than just one type of

particle dissolved in them. So when talking about body fluids it is much better to

use osmolarity.

The more solute dissolved in a solution, the more osmotic pressure that solution

has, and the higher the osmolarity of the solution.

Osmolarity

depends upon the

total NUMBER,

NOT THE TYPE,

of particles in a

solution.

The following is important to understand: The osmolarity of a solution depends

upon the total NUMBER, NOT THE TYPE, of particles dissolved in a solvent.

What do we mean by particles? In terms of osmolarity, particles can be atoms,

ions, molecules, or even cells in solution without regard to molecular weight or

electrical charge of the individual particles.

So, 1 particle (or mole or millimole) of glucose would be equivalent to 1 particle

(or mole, or millimole) of Na, or 1 particle of Cl, or 1 particle of Mg, or anything

else in solution. The type of particle does NOT matter; only the NUMBER of

particles in solution contributes to osmolarity.

Ionically bonded

molecules break

apart (dissociate)

in solution;

covalently bonded

molecules do not.

Covalently bonded molecules, e.g., glucose, proteins, lipids, do not break apart

when dissolved in aqueous solvent. (They are called...