Osmosis and Osmotic Pressure

Osmosis is a physical phenomenon in which molecules of solvent move from the region of low solute concentration to the region of high solute concentration through a semipermeable membrane. The study of osmosis is of interest in chemistry as well as in biology since many physical and biological processes are governed by it. We will mention some real-life examples later in this article.

Osmosis

Consider a U-shaped tube filled with pure solvent, as illustrated below. Both arms of the tube are separated by a semipermeable membrane placed at the center of the tube. Initially, solvent flows from the left to right arm, and vice versa in the absence of solute. So, the net flow is zero. And we describe the condition as initial equilibrium between both arms.

The situation changes when a solute is added to one of the arms, says left. Now, we have two solutions: one on the left with a higher concentration of solute and one on the right with a lower concentration of solute. As a consequence, our initial equilibrium is disturbed, and we see the net positive flow of solvent from the right (low solute concentration) to the left (high solute concentration).

One thing to note here is that only solvent molecules are allowed to diffuse from right to left; no solute is transferred to either arm. The movement of solute is restricted by the semipermeable membrane separating the two solutions.

As the flow of solvent continues, the liquid level rises on the left arm and falls on the other. After a while, we reach a new equilibrium, and the net flow becomes zero.

For some readers, osmosis may seem counter-intuitive. In many of our natural phenomena that we have experienced and have studied in other science subjects, the flow of a quantity is governed by potential difference, i.e—the flow occurs from higher to lower potential. For example, water flows from higher ground to lower, wind flows from higher pressure to lower, heat flows from higher temperature to lower, or particles diffuse from higher concentration to lower.

However, in osmosis, the flow is opposite of what we might expect: the flow is from lower concentration to higher. What if I tell you our intuition is right. The flow is from higher concentration to lower. But how so? You can think of it this way: in osmosis, solvent moves from the region of higher solvent concentration (equivalent to lower solute concentration) to lower solvent concentration (equivalent to higher solute concentration). If we redefine osmosis based on solvent concentration instead of solute concentration, our intuition holds true.

In the above-discussed case, the left arm has higher solute concentration, which is equivalent to saying the left arm has lower solvent concentration. And the right arm has lower solute concentration (or the right arm has higher solvent concentration). So, the solvent diffuses from right to left because the right arm is at higher potential (higher solvent concentration) compared to the left.

The true understanding of how osmosis works and what decides the flow occurs involves thermodynamics and determining thermodynamic potential, which is simply not in the scope of this article. Though the above explanation might be invalid in some cases, it acts as a good explanation for beginners to grasp the concept of osmosis.

Osmotic pressure

Now, we have some idea of osmosis, so we can move further and introduce ourselves to osmotic pressure. Osmosis is a tendency of solvent molecules to permeate themselves through a semipermeable membrane and diffuse into a higher-solute-concentration solution. So, it is a kind of pressure that forces solvent molecules to diffuse. We can stop osmosis by applying counter pressure. This applied counter pressure that nullifies osmosis is a measure of osmotic pressure.

Osmotic pressure is defined as the amount of pressure applied to stop the flow of solvent from entering into a higher solute concentration solution separated by a semipermeable membrane.

In the above diagram, we establish the original state of the system by applying counter pressure. At equilibrium, the net solvent flow is zero, and counter pressure balances osmotic pressure. Note that counter pressure arises from external factors, for example, created by us to stop osmosis, while osmotic pressure originates because of intrinsic reasons, which here is concentration difference. Osmotic pressure exists even without an external counter pressure since it is the property of the system.

Osmotic pressure equation

For a very dilute solution, i.e., the concentration of solute is very low, we can treat it as an ideal solution. Osmotic pressure for an ideal solution is expressed by the van’t Hoff equation.

∏ = icRT

Here, ∏ is the osmotic pressure, i is van’t Hoff index, c is the molar concentration of solute, R is the ideal gas constant, and T is the absolute temperature in the kelvin.

The van’t Hoff index is proportional to the degree of dissociation, which is the ability of a solute to dissociate when dissolved in the solution. As an example, NaCl dissociates into Na+ ions and Cl- ions when dissolved in a polar solvent like water. For nondissociative solutions, like nonelectrolytes in water, the van’t Hoff index is almost unity (i=1).

If you have noticed, the equation resembles the ideal gas equation, PV=nRT. If you rearrange the ideal gas equation, we have P = (n/V) RT. We can make an analogy between the two equations.

P, the pressure of gas ≍ ∏, the osmotic pressure

n/V, the molar density ≍ c, the concentration of solute

T, the temperature of gas ≍ T, the temperature of solution.

(≍ is equivalent to symbol)

Reverse osmosis

What if we apply the pressure in excess of the osmotic pressure in the previous explanation? It results in the movement of solvent molecules from a higher concentration of solute to a lower—from the left arm to the right. Thus, we are reversing the natural direction of the osmotic flow, so we named it reverse osmosis.

At equilibrium, the counter pressure equals the osmotic pressure, as illustrated above. As counter pressure exceeds osmotic pressure, solvent from higher concentration diffuses to the lower concentration arm. Consequently, the solute concentration increases in the left arm. It, in turn, increases osmotic pressure since the osmotic pressure is directly proportional solute concentration according to the van’t Hoff equation. So, now, we have to apply even more pressure. And this goes on, which makes it almost impossible after a certain point to separate complete solvent from the solution.

Reverse osmosis is a separation process and is widely used on an industrial scale for water purification.

Examples and applications

Some examples and applications of osmosis and reverse osmosis are mentioned below.

Wilting of leaves

When gardening plants are not watered, their leaves become wilted. Why? And when you start watering them, they quickly return back to their original shape and get inflated. The answer to it lies in osmosis.

A plant cell contains some amount of water and salt. Salt plays the role of solute, and water acts as a solvent. Under normal conditions, the exchange of water across the plant cell wall and the environmental fluid is equal, and we say the environment is isotonic with cells. When a plant cell is placed in an environment that has a relatively higher solute concentration (hypertonicity), there is a net outflow of water from the cell to the outside. As a result, the cell shrinks and might get plasmolyzed. In contrast, when a cell comes in contact with a hypotonic solution, the cell experiences a net inflow of water and swells.

Transport of water from soil to plant

Ever wonder how plants can suck water from the soil? Osmosis plays a very vital role here. The roots of plants are rich in salts. So, the water in soil has a relatively low concentration of salts, which creates potential. And water slips into the roots.

Blood cells

Animal cells are no exception; osmosis is also crucial for the maintenance and transport of constituents in animal and human cells.

(Public Domain: Mariana Ruiz Villareal)

The diagram above depicts the nature of blood cells under different conditions. Under isotonic, blood cells are in their regular shape. In a water-enriched environment, blood cells inflate and could burst in extreme conditions, which is highly undesirable. In the same way, in a water-deprived environment, cells collapse because of hypertonicity.

Marine and freshwater life

Osmosis can also have detrimental effects on marine and freshwater organisms. For example, marine fish are adapted to saltwater over the years of evolution. Their bodies exist in natural equilibrium with their natural environment. When they are suddenly introduced to freshwater, their cells will keep inflating. And eventually, they will burst to the degree that the creature will die.

Water purification

Water purification is the biggest application of reverse osmosis. The water in oceans and lakes contains undesirable solute particles, which covers heavy salts, toxics, pathogens, and impurities like dirt and sand grains. Many of these can be separated by physical and chemical processes, such as coagulation and flocculation. Pathogens can be killed by chlorination and radiation. But economically separating water from salt plus water is a big challenge (and still remains). That is where reverse osmosis comes in. As we have explained earlier, if applied an adequate amount of force, we can separate solvent from solution.

Reverse osmosis, abbreviated as RO, plants could be at an industrial scale or portable like the one in the below image. You can also install an RO unit in your backyard.

(Public domain: US Army)

A portable reverse osmosis unit purifies water from a lake and stores it in bladders.