# Osmotic pressure

Osmosis is defined as the fluid’s diffusion via semipermeable membrane. When the semipermeable membrane (Eg like fruits’ and vegetables’ skin, animal bladders) divides a solution from the solvent, only then, the solvent molecules are in the driver’s seat to go through that membrane.

Solution’s osmotic pressure is referred to the difference in pressure which is required to stop the solvent’s flow all over the semipermeable membrane. Solution’s osmotic pressure is proportional to solute particles’ molar concentration in the solution.

Π = nRT / V= MMRT

where,

• Π = osmotic pressure,
• R = ideal gas constant (with the value – 0.0821 L atm/mol K),
• T = temperature in Kelvin,
• n = number of moles of the solute exist,
• V = volume of a solution (then, n / V = molar concentration of a solute, and
• MM = molar mass of a solute.

Remember that the n/V co-relates with solution’s molarity of the solute non-dissociative in nature, or twice the molarity of entirely-dissociated solute like NaCl. Regarding this, molarity is defined as the total sum of concentrations of the whole of solute varieties.

Summoning that Π is the Greek equivalent of P, the re-arrangement from ΠV = nRT from the above equation should be visible simply and recognizable. So much extra effort was needed by the end of the 19th century to describe the similarity between ideal gas law and this above relation. Although, the equation of Van’t Hoff yields a much rough estimation of actual law of osmotic pressure, which is significantly much more complex and had the derivation after the formulation of Van’t Hoff. By its very nature, this equation gives the logical results for extremely dilute (i.e. ideal) solutions only.

As per the Van’t Hoff equation, the ideal solution which contains one mole of dissolved particles/litre of solvent at zero degrees temperature and will have the osmotic pressure at 22.4 atm.

Introduction

What do you mean by osmotic pressure? The Semipermeable membrane doesn’t allow the solvent to pass through. (let’s take sugar example).

The solvent will go towards that which is more focused on trying to make each side more alike! Since solvents’ flow is there, the height of each of the side varies, which is the osmotic pressure. When you’re working with aqueous solutions, people use the mm of the water molecule (H2O) to illustrate the difference.

The osmotic pressure is very much related to a few of the alternative properties of solutions like the colligative properties. These involve elevation in boiling point, depression in freezing point, and depression in vapor pressure, all leading up to dissolving solutes in the solution.

Oftentimes, the osmolarity is found out from depression in vapor pressure or the depression in freezing point, instead of, from direct measurements of osmosis pressure. Osmolarity refers to the essential concentration to witness these phenomena.

Osmotic equilibrium and osmotic pressure

A single method to stop is raising the hydrostatic pressure on the side of the solution of a membrane. This osmotic pressure compresses the molecules of the solvent very close together, to raise the escaping tendency of theirs from the phase. If you apply extra pressure or allow the pressure building up by the osmotic liquid’s flow into the enclosed area, the tendency to escape will ultimately rise comparatively to pure solvent’s molecules and the osmotic flow.

The pressure needed for achieving the osmotic equilibrium state is called the osmotic pressure.

Keep in mind, the osmotic pressure is that pressure which is needed to restrict osmosis, and not to help it further and maintain it.