Osmosis and osmotic pressureDiffusion through a semipermeable membrane

Osmotic pressure is the fourth member of the quartet of colligative properties that arise from the dilution of a solvent by non-volatile solutes. Because of its great importance, we are devoting a separate section to this topic with special emphasis on some of its many practical applications.

 

 

 

 

Semi permeable membranes and osmotic flow

Osmosis is the process in which a liquid passes through a membrane whose pores permit the passage of solvent molecules but are too small for the larger solute molecules to pass through.

The figure shows a simple osmotic cell. Both compartments contain water, but the one on the left also contains a solute whose molecules (represented by blue circles) are too large to pass through the membrane. Many artificial and natural substances are capable of acting as semi-permeable membranes. The walls of most plant and animal cells fall into this category.

If the cell is set up so that the liquid level is initially the same in both compartments, you will soon notice that the liquid rises in the left compartment and falls in the right side, indicating that water molecules from the right compartment are migrating through the semipermeable membrane and into the left compartment. This migration of the solvent is known as osmotic flow, or simply osmosis .

The escaping tendency of a substance from a phase increases with its concentration in the phase

What is the force that drives the molecules through the membrane? This is a misleading question, because there is no real “force” in the physical sense other than the thermal energies all molecules possess. Osmosis is a consequence of simple statistics: the randomly directed motions of a collection of molecules will cause more

to leave a region of high concentration than return to it; the escaping tendency of a substance from a phase increases with its concentration in the phase.

Diffusion and osmotic flow

Suppose you drop a lump of sugar into a cup of tea, without stirring. Initially there will be a very high concentration of dissolved sugar at the bottom of the cup, and a very low concentration near the top. Since the molecules are in random motion, there will be more sugar molecules moving from the high concentration region to the low concentration region than in the opposite direction. The motion of a substance from a region of high concentration to one of low concentration is known as diffusion. Diffusion is a consequence of a concentration gradient (which is a measure of the difference in escaping tendency of the substance in different regions of the solution.

You must clearly understand that there is really no special force on the individual molecules; diffusion is purely a consequence of statistics.

Osmotic flow is simply diffusion of a solvent through a membrane impermeable to solute molecules

Now take two solutions of differing solvent concentration, and separate them by a semipermeable membrane. Being semipermeable, the membrane is essentially invisible to the solvent molecules, so they diffuse from the high concentration region to the low concentration region just as before. This flow of solvent constitutes osmotic flow, or osmosis.

This illustration shows water molecules (blue) passing freely in both directions through the semipermeable membrane, while the larger solute molecules remain trapped in the left compartment, diluting the water and reducing its escaping tendency from this cell, compared to the water in the right side. This results in a net osmotic flow of water from the right side which continues until the increased hydrostatic pressure on the left side raises the escaping tendency of the diluted water to that of the pure water at

1 atm, at which point osmotic equilibrium is achieved.

 

In the absence of the semipermeable membrane, diffusion would continue until the concentrations of all substances are uniform throughout the liquid phase. With the semipermeable membrane in place, and if one compartment contains the pure solvent, this can never happen; no matter how much liquid flows through the membrane, the solvent in the right side will always be more concentrated than that in the left side. Osmosis will continue indefinitely until we run out of solvent, or something else stops it.

Osmotic equilibrium and osmotic pressure

Osmotic pressure is the pressure required to stop osmotic flowCaution! It is common usage to say that a solution “has” an osmotic pressure of "x atmospheres". It is important to understand that this means nothing more than that a pressure of this value must be applied to the solution in order to prevent flow of pure solvent into this solution through a semipermeable membrane separating the two liquids.

Osmotic pressure and solute concentration

The osmotic pressure Π (Pi) of a solution containing n moles of solute particles in a solution of volume V is given by the van't Hoff equation:

In contrast to the need to employ solute molality to calculate the effects of a non-volatile solute on changes in the freezing and boiling points of a solution, we can use solute molarity to calculate osmotic pressures.

Π = nRT / V

in which R is the gas constant (0.0821 L atm mol–1 K–1) and T is the absolute temperature.

Note that the fraction n/V corresponds to the molarity of a solution of a non-dissociating solute, or to twice the molarity of a totally-dissociated solute such as NaCl. In this context, molarity refers to the summed total of the concentrations of all solute species.

Recalling that Π is the Greek equivalent of P, the re-arranged form ΠV = nRT of the above equation should look familiar. Much effort was expended around the end of the 19th century to explain the similarity between this relation and the ideal gas law, but in fact, the Van’t Hoff equation turns out to be only a very rough approximation of the real osmotic pressure law, which is considerably more complicated and was derived

after van't Hoff's formulation. As such, this equation gives valid results only for extremely dilute ("ideal") solutions.

According to the Van't Hoff equation, an ideal solution containing 1 mole of dissolved particles per liter of solvent at 0° C will have an osmotic pressure of 22.4 atm.

Problem Example 1

Sea water contains dissolved salts at a total ionic concentration of about 1.13 mol L–1. What pressure must be applied to prevent osmotic flow of pure water into sea water through a membrane permeable only to water molecules?

Solution:

Π = MRT = (1.13 mol L–1)(0.0821 L atm mol–1 K–1)(298 K) = 27.6 atm

2  Some practical applications of osmosis

 

Reverse osmosis

If it takes a pressure of Π atm to bring about osmotic equilibrium, then it follows that applying a hydrostatic pressure greater than this to the high-solute side of an osmotic cell will force water to flow back into the fresh-water side. This process, known as reverse osmosis, is now the major technology employed to desalinate ocean water and to reclaim "used" water from power plants, runoff, and even from sewage. It is also widely used to

deionize ordinary water and to purify it for for industrial uses (especially beverage and food manufacture) and drinking purposes.

 

Pre-treatment commonly employs activated-carbon filtration to remove organics and chlorine (which tends to damage RO membranes). Although bacteria are unable to pass through semipermeable membranes, the latter can develop pinhole leaks, so some form of disinfection is often advised.

Membranes for reverse osmosis

The efficiency and cost or RO is critically dependent on the properties of the semipermeable membrane.

 

 

 

 

 

Using "normal saline solution" to prevent osmotic disruption of cells

The interiors of cells contain salts and other solutes that dilute the intracellular water. If the cell membrane is permeable to water, placing the cell in contact with pure water will draw water into the cell, tending to rupture it.

This is easily and dramatically seen if red blood cells are placed in a drop of water and observed through a microscope as they burst. This is the reason that "normal saline solution", rather than pure water, is administered in order to maintain blood volume or to infuse therapeutic agents during medical procedures.

In order to prevent irritation of sensitive membranes, one should always add some salt to water used to irrigate the eyes, nose, throat or bowel.

Normal saline contains 0.91% w/v of sodium chloride, corresponding to 0.154 M, making its osmotic pressure close to that of blood.

Osmotic pressure and food preservation

The drying of fruit, the use of sugar to preserve jams and jellies, and the use of salt to preserve certain meats, are age-old methods of preserving food. The idea is to reduce the water concentration to a level below that in living organisms. Any bacterial cell that wanders into such a medium will have water osmotically drawn out of it, and will die of dehydration. A similar effect is noticed by anyone who holds a hard sugar candy against the inner wall of the mouth for an extended time; the affected surface becomes dehydrated and noticeably rough when touched by the tongue.

In the food industry, what is known as water activity is measured on a scale of 0 to 1, where 0 indicates no water and 1 indicates all water. Food spoilage micro-organisms, in general, are inhibited in food where the water activity is below 0.6. However, if the pH of the food is less than 4.6, micro-organisms are inhibited (but not immediately killed] when the water activity is below 0.85.

raisinssalted codfish

Osmosis and diarrhea

The presence of excessive solutes in the bowel draws water from the intestinal walls, giving rise to diarrhea. This can occur when a food is eaten that cannot be properly digested (as, for example, milk in lactose-intolerant people). The undigested material contributes to the solute concentration, raising its osmotic pressure. The situation is made even worse if the material undergoes bacterial fermentation which results in the formation of methane and carbon dioxide, producing a frothy discharge. [image]

Water transport in plants: osmosis pushes, hydrogen-bonding pulls

Osmotic flow plays an important role in the transport of water from its source in the soil to its release by transpiration from the leaves, it is helped along by hydrogen-bonding forces between the water molecules. Capillary rise is not believed to be a significant factor.

isotonic solutions isotonic solutions are solutions having the same osmotic pressure as that

of the body fluids when separated by a biological membrane. Biological fluids including blood and lachrymal fluid normally have an osmotic pressure corresponding to that of 0.9% w/v solution of sodium chloride. Thus 0.9% solution of sodium chloride is said to be isotonic with the

physiological fluids .

Isotonic solutions cause no swelling or contraction of the tissues with which they come in contact and produce no discomfort when instilled in the eye, nasal tract, blood, or other body tissues. Isotonic sodium chloride is a familiar pharmaceutical example of such a preparation The need to achieve isotonic conditions with solutions to be applied to delicate membranes is dramatically illustrated by mixing a small quantity of blood with aqueous sodium chloride solutions of varying tonicity.

hypertonic If the red blood cells are suspended in a 2.0% NaCl solution, the water within the cells passes through the cell membrane in an attempt to dilute the surrounding salt solution. This outward passage of water causes the cells to shrink and become wrinkled. hypotonic if the blood is mixed with 0.2% NaCl solution or with distilled water, water enters the blood cells, causing them to swell and finally burst, with the liberation of hemoglobin .The salt solution in this instance is said to be with respect to the blood cell contents. Finally, This

phenomenon is known as hemolysis. Isosmotic solutionsThe red blood cell membrane is not impermeable to all drugs; that is, it is not a perfect semipermeable membrane. Thus, it will permit the passage of not only water molecules but also solutes such as urea, ammonium chloride, alcohol, and boric Acid .These solutes are regarded as solvent and they do not exert an osmotic pressure on the membrane (the solutions are isosmotic but not isotonic) A 2.0% solution of boric acid has the same osmotic pressure as the blood cell contents when determined by the freezing point method and is therefore said to be isosmotic with blood. The molecules of boric

acid pass freely through the erythrocyte membrane, however, regardless of concentration. As a result, this solution acts essentially as water when in contact with blood cells. Because it is extremely hypotonic with respect to the blood, boric acid solution brings about rapid hemolysis. Therefore, a solution containing a quantity of drug calculated to be isosmotic with blood is isotonic .Accordingly, a 2.0% boric acid solution serves as an isotonic ophthalmic preparation.

An isotonic solution is one that has the same osmolarity, or solute concentration, as another solution. If these two solutions are separated by a semipermeable membrane, water will flow in equal parts out of each solution and into the other. The effect is zero water flow between the two solutions, although water is moving both ways. In biology, some cells must be maintained in an isotonic solution to support cellular functions. Many animal cells, which lack a cell wall to provide support against the effects of water pressure, rely on the stability of the external environment to maintain their shape. Most animals maintain the pH and osmolarity of the fluids inside of their bodies to create isotonic solutions to bathe their cells in. This solution can carry nutrients and water, but only in proportions equal to that inside the cell.

A depiction of a cell in an isotonic solution can be seen above. Note that because there is the same concentration of solute molecules inside and outside of the cell that water molecules are simply exchanged through the cell membrane. This can be contrasted to the effects of a hypertonic solution, in which water molecules leave the cell, or a hypotonic solution in which water enters the cell.

Examples of Isotonic Solution

Blood Cells

When the plasma surrounding blood cells is an isotonic solution, compared to the solution inside the blood cells, the cells function normally. The isotonic solution allow the cells to move water and nutrients in and out of the cells. This is necessary for blood cells to perform their function of delivering oxygen and other nutrients to other parts of the body. If the cells are in a hypertonic environment, they will become plasmolyzed and will not contain enough water to perform cellular functions. If the cells exist in a hypotonic environment, they will lyse, spilling their contents into the bloodstream. This can cause dangerous side effects, as well as the loss of many blood cells. These events can be seen in the graphic below.

To avoid either of the negative situations from happening during the transfusion of nutrients and medicine, the solution that carries the medicine must be an isotonic solution, compared to the patient’s blood. The osmolarity of the IV fluid can be adjusted using special salts and sugars that act simply as solutes to dilute or strengthen a substance. Once a medicine is an isotonic solution compared to the blood, it can be added through an IV and no damage will occur to blood cells.

Hypotonic – When a solution has comparatively more water and less solute.

Hypertonic – A solution with less water and more solute than another solution.

Osmolarity – The overall solute concentration of a solution.

Precipitations Analysis

Chemical precipitation, formation of a separable solid substance from a solution, either by converting the substance into an insoluble form or by changing the composition of the solvent to diminish the solubility of the substance in it. The distinction between precipitation and crystallization lies largely in whether emphasis is placed on the process by which the solubility is reduced or on that by which the structure of the solid substance becomes organized.

Precipitation often is used to remove metal ions from aqueous solutions: silver ions present in a solution of a soluble salt, such as silver nitrate, are precipitated by addition of chloride ions, provided, for example, by a solution of sodium chloride; the chloride ions and the silver ions combine to form silver chloride, a compound that is not soluble in water. Similarly, barium ions are precipitated by sulfate ions, and calcium by oxalate . In many cases it is possible to select conditions under which a substance precipitates in highly pure and easily separable form. Isolation of such precipitates and determination of their weights constitute accurate methods for determining the amounts of various compounds. In attempts to precipitate a single substance from a solution containing several components, undesired constituents often are incorporated in the crystals, reducing their purity and impairing the accuracy of the analysis. Such contamination can be reduced by carrying out the operations with dilute solutions and by adding the precipitating agent slowly; an effective technique is that called homogeneous precipitation, in which the precipitating agent is synthesized in the solution rather than added

mechanically. In difficult cases it may be necessary to isolate an impure precipitate, re dissolve it, and precipitate it; most of the interfering substances are removed in the original solution, and the second precipitation is performed in their absence. Precipitation reactions occur when cations and anions in aqueous solution combine to form an insoluble ionic solid called a precipitate. Whether or not such a reaction occurs can be determined by using the solubility rules for common ionic solids. Because not all aqueous reactions form precipitates, one must consult the solubility rules before determining the state of the products and writing a net ionic equation. The ability to predict these reactions allows scientists to determine which ions are present in a solution, and allows industries to form chemicals by extracting components from these reactions.

Properties of Precipitates

Precipitates are insoluble ionic solid products of a reaction, formed when certain cations and anions combine in an aqueous solution. The determining factors of the formation of a precipitate can vary. Some reactions depend on temperature, such as solutions used for buffers, whereas others are dependent only on solution concentration. The solids produced in precipitate reactions are crystalline solids, and can be suspended throughout the liquid or fall to the bottom of the solution. The remaining fluid is called supernatant liquid. The two components of the mixture (precipitate and supernate) can be separated by various methods, such as filtration, centrifuging, or decanting.

 

Figure 1: Above is a diagram of the formation of a precipitate in solution.