Diffusion is the random movement of molecules that leads to a net movement of those molecules from high to low concentration & an evenly mixed solution. Osmosis is just a special case diffusion when you have 2 “pools” of particles separated by a partition & only the solvent (for an aqueous solution, this is water), NOT what’s dissolved in it (the solute) can get through (we call the partition semipermeable). Since the solute can’t move, it’s up to the solvent to move to even things out. 

note: mix-matched & adapted from some super old posts I thought I’d revisit

We can measure the concentration of the dissolved “stuff” as osmolarity, which is # of particles (of any kind) in a given volume. Imbalances in osmolarity between sides of the partition create osmotic pressure. More on this in a minute. But I want to clear up some confusion (hopefully).

Osmolarity is often confused with tonicity. Tonicity is a comparative term that *only* takes into consideration concentration of nonpenetrating particles on either side of a semipermeable barrier. It “looks” at the equilibrium the mix will reach & tells you which direction you’ll have net movement of (or other solvent). 

The key point is that water moves FROM where there’s more nonpenetrating particles TO where there are fewer nonpenetrating particles, thereby balancing out the amount of “free water.” 

Why only nonpenetrating stuff? The penetrating particles can balance themselves out, so they’ll equally raise the osmolarity on both sides of barrier. BUT nonpenetrating particles will only raise the osmolarity on the side they’re on. This creates an imbalance, so water (or other solvent) has to move to balance out “effective” osmolarities. So, we can think of calculating the osmolarities of the solute molecules that are “osmotically active” (those that are nonpenetrating and thus contribute to the osmotic pressure. 

As I mentioned, tonicity is a comparative term so we use it to compare one solution to another:

  • hypertonic describes a solution that has higher concentrations of nonpenetrating solvents (you can remember hyper:over)
  • hypotonic describes a solution that has lower concentrations of nonpenetrating solvents (hypo:below)
  • isotonic describes a solution that has equal concentration of nonpenetrating solvents (iso:I so don’t care where I go because it’s all the same)

Solvent moves from a hypotonic solution to a hypertonic solution, counterbalanced by external pressure and all that water weight… So lets talk about that weight!

Imagine a U tube (no, not the one with videos, just a U-shaped tube) with a partition at the bottom that lets through solvent but not solute (i.e. the partition is permeable with respect to solvent but impermeable with respect to the solute; alternatively we can say the solvent is penetrating and the solute in nonpenetrating). Take that U tube & put pure solvent on one side & a solution on other side. What’s gonna happen?

Solvent will move from the solvent side to the solution side (volume of the solvent side decreases & solution side increases), diluting the solute. Thereby, the sides become more & more similar in terms of their composition. But, they never become identical…

You’d think the solvent side would dry itself out since the 2 sides can never balance in terms of composition (since there’s no solute on solvent side you can never have = solute concentrations on both sides). But it doesn’t. At some point, the solution side stops rising even though there’s still solvent on the other side. Why? Osmotic pressure (π)!

Because the liquid column’s taller on the solution side, there’s more hydrostatic pressure, weight of liquid pushing solution “back” towards solute side. This produces a kind of “back-pressure” making it harder to flow through in the solvent → solution direction

It’s important to remember that the solvent can (& does) flow in both directions, it’s just that more goes solvent→solution until equilibrium is reached, at which point it goes both directions at same rate). If you apply “artificial” pressure (like pushing down on it w/plunger) you can force the solvent to go the other way. In fact, this is how reverse osmosis purifies water.

So Osmotic pressure (π) is basically the pressure needed to overcome benefits that come with even-ing out the solutions.

π is a property of the solution. We can calculate osmotic pressure of any solution only knowing molar concentration (moles/L) of solute particles (M) & temperature (T)

π = MRT 

R is the gas constant (0.0821 L atm mol–1 K–1) & since it’s in Kelvin (K), we need to plug in temp in absolute temperature (temperature in K)) to match (you can easily convert: K = °C + 273.15)

More solute particles (M) leads to higher π, meaning you’d have to apply more pressure to stop solvent from moving in

Just concentration & temp? Not identity of the solute?! Yep – it’s another colligative property (joining vapor pressure depression, boiling point elevation, & freezing point depression to round out the set!) For more on those, see https://bit.ly/boilingpointvaporization & http://bit.ly/freezingpointdepression

Just like for those other colligative properties, you need to take into account number of particles not just “formula units”  e.g. if you have an electrolyte that splits into ions when you dissolve it, each of those ions counts as a separate particle (i.e. NaCl counts as 2 particles bc it dissociates into Na⁺ & 1 Cl⁻) More on this here: http://bit.ly/solutionconcentrations but here’s the gist

For non-electrolytes (like table sugar, sucrose), the number of dissolved particles is the same as the number of “formula units” (like what you’d buy it as – in this case, sucrose). so put in 6×10²³ sucroses (1 mole) & you get 6×10²³ dissolved particles. And if you dissolved it in 1L, you’d have a 1 Molar (1M) solution. It’s molarity is 1mol/L.

BUT for electrolytes (like table salt, NaCl) you get more dissolved particles than you put in because, unlike the strong covalent bonds holding together molecules like sucrose, the charge-based ionic bonds holding together these ionic compounds get broken & replaced w/water-solute interactions when you dissolve them. NaCl breaks into Na⁺ & Cl⁻ so if you put in 1 mol of NaCl you get 2 mol of dissolved particles.

We can use the term molarity(mol solute/L) to describe solute concentration, but clearly it can get confusing/ambiguous if you have electrolytes. By solute do you mean formula units or dissolved particles? (typically formula units & you multiply by the Van’t Hoff factor (which tells you how many pieces it breaks into upon dissolving) to get # of dissolved particles (n)). But you can also use molarity to describe the dissolved particles…

So if I put 1 mol of NaCl into 1 L of water I’d get a solution I could describe as…

  • 1M NaCl or
  • 1M Na+ & 1M Cl- or
  • 2M dissolved particles

Bring on the confusion & miscalculations?! No thanks… 

Instead, we can use a term called osmolarity. It’s similar to molarity except it clarifies unambiguously that we’re dealing with number of dissolved particles, not # of formula units & instead of moles, we can describe concentrations of these particles in osmoles. 1 osmole = 1 mole of dissolved particles.

This can help us greatly when considering colligative properties. But when it comes to osmosis, # of dissolved particles isn’t *really* all we usually care about. No, I didn’t lie to you before! Things are just normally not as simple as the classical definition assumes they are

Osmotic pressure really is a colligative property of a solution. It describes pressure needed to stop flow of pure solvent (often water) into a solution when solvent & solution are separated by a semipermeable partition that solvent, but not solute can get through. In this “normalized” definition, the solute is by definition nonpenetrating (it can’t get through the partition). 

So we can calculate a “normal” osmotic pressure for any solution. And if this is the case, differences in π will lead to osmosis so we can take any 2 solutions & predict which direction & how much solvent would need to flow between them to achieve equilibrium 

BUT… sometimes, instead of such an “optimized” situation we have solute(s) that can get through (they are penetrating). These particles still contribute to the osmotic pressure you calculate (based on that idealized situation) but they do affect osmosis…

So we need another term: tonicity. Tonicity is a comparative term that only takes into consideration concentration of nonpentrating solute & tells you which way you’ll have net movement to reach equilibrium

So just like permeability (which describes partition with respect to particles) & penetratability? (not sure if that’s a real word…) (which describes particles with respect to partition), tonicity is a relative term. It describes osmolarity of one part compared to osmolarity of the other part when you only consider nonpenetrating solute (so basically the “effective osmolarity”). Why’s this helpful?

Equilibrium is reached by movement of solvent from a hypotonic (hypO, belOW) solution to a hypertonic (hypER, over) solution until they are isotonic (I SO don’t care where I go because the osmolarity’s the same everywhere)

Biological cells have ways of preventing overfilling despite being packed with “stuff.” Here are a few examples:

  • Plants & bacteria have strong cell walls that can resist osmotic pressure. Plants even use osmotic pressure to their advantage… “excess” water that’s entered pushes agains the cell wall, (fruitlessly) trying to expand it. This creates hydrostatic pressure which provides stiffness so you get crispy vegetables
  • sugars are stored as polysaccharides (e.g. starch & glycogen), which are large chains of simple sugars (e.g. glucose) instead of as lots of simple sugars. This lets them store same amount of energy but with lower osmolarity
  • animals pump ions & other molecules into extracellular fluids to make those fluids isotonic 
  • serum albumin in the blood (a super abundant protein) prevents excess water from flowing into cells from blood


This also implications for medicine, in particular IV meds which can alter these osmolarities &/or tonicity. Consider 2 common IV solutions, normal saline (0.9% NaCl)(NS) & D-5-W (5% dextrose (glucose) in water). They’re isosmotic to each other (iso meaning same) & to cells BUT

  • NS *is* isotonic to cells because NaCl doesn’t enter the cells (it’s non penetrating so these molecules “count”) 
  • but D-5-W is hypotonic because glucose goes into cells (it’s a penetrating solute so it doesn’t “count” in determining tonicity)
    • have to think about how equilibrium will be reached… glucose can enter the cells -> lowers osmolarity outside -> makes the solution hypotonic ->  water flows in to help “dilute” it & balance out the concentrations of “free water”

So…

  • isosmotic solutions are NOT always isotonic. 
  • hyperosmotic solutions are NOT always hypertonic
  • BUT hyposmotic solutions ARE always hypotonic (hyposmotic tells you you have lower concentrations, so if that solute is penetrating, the difference in tonicity would be even greater)

Other lessons…

  • don’t drink salt water -> it’s hypertonic -> pulls water out of your cells -> dehydrates you (also other gross stuff in there…)
  • but, even if you are dehydrated, don’t give a pure water IV -> hypotonic -> burst cells
  • don’t let your dog pee on grass. it’s hypertonic so it pulls water out of grass, leading to a dead patch

Lab (as in laboratory not Labrador retrievers…) examples include


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