LECTURE 8. COLLIGATIVE PROPERTIES OF SOLUTIONS
A solution is a homogeneous mixture of two or more chemically non-reacting substances whose composition can be varied within certain limits.
Solutions can be classified on the basis of their state: solid, liquid, or gas.
The substances making up the solutions are called components. The components of a binary solution are solute and solvent. Solvent is a component which is present in excess, in other words a solvent is a substance in which dissolution takes place. Solvent doesn’t change its physical state during reaction of dissolution. Solute is a component which is present in lesser quantity. Or solute is a substance that dissolves. In a solution, the particles are of molecular size (about 1000 pm) and the different components cannot be separated by any of the physical methods such as filtration, setting ,centrifugation, etc.)
Colligative properties of solutions.
What are Colligative Properties?
Colligative properties are properties of solutions that depend on the number of molecules in a given volume of solvent and not on the properties (e.g. size or mass) of the molecules. Colligative properties include: lowering of vapor pressure; elevation of boiling point; depression of freezing point and osmotic pressure. Measurements of these properties for a dilute aqueous solution of a non-ionized solute such as urea or glucose can lead to accurate determinations of relative molecular masses. Alternatively, measurements for ionized solutes can lead to an estimation of the percentage of ionization taking place.
Henry’s Law: The solubility of a gas dissolved in a liquid is proportional to the partial pressure of the gas above the liquid.
This is a statement of Henry’s law, which can be written
X = KP
where X is the equilibrium mole fraction of the gas in solution (its solubility), P is its partial pressure in the gas phase, and K is a constant of proportionality, usually called the Henry’s-law constant.
The partial pressure is a part of common pressure, which one is a share of each gas in gas mixture.
Henry’s law applies only when the concentration of the solute and its partial pressure above the solution are both low, that is, when the gas and its solution are both essentially ideal, and when the solute does not interact
On no-bottoms, where the external pressure increases, the dissolubility of gases in a blood is augmented. At fast ascent from depth the dissolubility sharply decreases, they are excreted by the way is bubble and seal vessels – aeroembolism.
Properties of a solution which depend only on the concentration of the solute and not upon its identity are known as colligative properties. These include vapor-pressure lowering, boiling-point elevation, freezing-point depression, and osmotic pressure. Each of these properties is a consequence of a decrease in the escaping tendency of solvent molecules brought about by the presence of solute particles. Escaping tendency is the tendency shown by molecules to escape from the phase in which they exist.
Raoult’s Law and Vapor Pressure Lowering
When a nonvolatile solute is added to a liquid to form a solution, the vapor pressure above that solution decreases. To understand why that might occur, let’s analyze the vaporization process of the pure solvent then do the same for a solution. Liquid molecules at the surface of a liquid can escape to the gas phase when they have a sufficient amount of energy to break free of the liquid’s intermolecular forces. That vaporization process is reversible. Gaseous molecules coming into contact with the surface of a liquid can be trapped by intermolecular forces in the liquid. Eventually the rate of escape will equal the rate of capture to establish a constant, equilibrium vapor pressure above the pure liquid.
If we add a nonvolatile solute to that liquid, the amount of surface area available for the escaping solvent molecules is reduced because some of that area is occupied by solute particles. Therefore, the solvent molecules will have a lower probability to escape the solution than the pure solvent. That fact is reflected in the lower vapor pressure for a solution relative to the pure solvent. That statement is only true if the solvent is nonvolatile. If the solute has its own vapor pressure, then the vapor pressure of the solution may be greater than the vapor pressure of the solvent.
Note that we did not need to identify the nature of the solvent or the solute (except for its lack of volatility) to derive that the vapor pressure should be lower for a solution relative to the pure solvent. That is what makes vapor pressure lowering a colligative property–it only depends on the number of dissolved solute particles.
summarizes our discussion so far. On the surface of the pure solvent (shown on the left) there are more solvent molecules at the surface than in the right-hand solution flask. Therefore, it is more likely that solvent molecules escape into the gas phase on the left than on the right. Therefore, the solution should have a lower vapor pressure than the pure solvent.
Figure : The Vapor Pressure of a Solution is Lower than that of the Pure Solvent
The French chemist Francois Raoult discovered the law that mathematically describes the vapor pressure lowering phenomenon. Raoult’s law is given in :
Figure %: Raoult’s Law Describes the Mathematics of Vapor Pressure Lowering
Raoult’s law states that the vapor pressure of a solution, P, equals the mole fraction of the solvent, c solvent, multiplied by the vapor pressure of the pure solvent, Po. While that “law” is approximately obeyed by most solutions, some show deviations from the expected behavior. Deviations from Raoult’s law can either be positive or negative. A positive deviation means that there is a higher than expected vapor pressure above the solution. A negative deviation, conversely, means that we find a lower than expected vapor pressure for the solution. The reason for the deviation stems from a flaw in our consideration of the vapor pressure lowering event–we assumed that the solute did not interact with the solvent at all. That, of course, is not true most of the time. If the solute is strongly held by the solvent, then the solution will show a negative deviation from Raoult’s law because the solvent will find it more difficult to escape from solution. If the solute and solvent are not as tightly bound to each other as they are to themselves, then the solution will show a positive deviation from Raoult’s law because the solvent molecules will find it easier to escape from solution into the gas phase.
Solutions that obey Raoult’s law are called ideal solutions because they behave exactly as we would predict. Solutions that show a deviation from Raoult’s law are called non-ideal solutions because they deviate from the expected behavior. Very few solutions actually approach ideality, but Raoult’s law for the ideal solution is a good enough approximation for the non- ideal solutions that we will continue to use Raoult’s law. Raoult’s law is the starting point for most of our discussions about the rest of the colligative properties, as we shall see in the following section.
Boiling Point Elevation
One consequence of Raoult’s law is that the boiling point of a solution made of a liquid solvent with a nonvolatile solute is greater than the boiling point of the pure solvent. The boiling point of a liquid or is defined as the temperature at which the vapor pressure of that liquid equals the atmospheric pressure. For a solution, the vapor pressure of the solvent is lower at any given temperature. Therefore, a higher temperature is required to boil the solution than the pure solvent. is a phase diagram for both a pure solvent and a solution of that solvent and a nonvolatile solute that explains that point graphically.
Figure: Phase Diagram for a Solvent and its Solution with a Nonvolatile Solute
As you can see in the the vapor pressure of the solution is lower than that of the pure solvent. Because both pure solvent and solutioeed to reach the same pressure to boil, the solution requires a higher temperature to boil. If we represent the difference in boiling point between the pure solvent and a solution as ΔTb, we can calculate that change in boiling point from the :
In the we use the units molality, m, for the concentration, m, because molality is temperature independent. The term Kb is a boiling point elevation constant that depends on the particular solvent being used. The term i in the above equation is called the van’t Hoff factor and represents the number of dissociated moles of particles per mole of solute. The van’t Hoff factor is 1 for all non-electrolyte solutes and equals the total number of ions released for electrolytes. Therefore, the value of i for Na2SO4 is 3 because that salt releases three moles of ions per mole of the salt.
Freezing Point Depression
As you may have noticed when we looked at the , the freezing point is depressed due to the vapor pressure lowering phenomenon. The points out that fact:
Figure: Phase Diagram for a Solution and the Pure Solvent Indicating the Freezing Point Depression
In analogy to the boiling point elevation, we can calculate the amount of the freezing point depression with the :
Note that the sign of the change in freezing point is negative because the freezing point of the solution is less than that of the pure solvent. Just as we did for boiling point elevation, we use molality to measure the concentration of the solute because it is temperature independent. Do not forget about the van’t Hoff factor, i, in your freezing point calculations.
One way to rationalize the freezing point depression phenomenon without talking about Raoult’s law is to consider the freezing process. In order for a liquid to freeze it must achieve a very ordered state that results in the formation of a crystal. If there are impurities in the liquid, i.e. solutes, the liquid is inherently less ordered. Therefore, a solution is more difficult to freeze than the pure solvent so a lower temperature is required to freeze the liquid.
Osmotic Pressure
Osmosis refers to the flow of solvent molecules past a semipermeable membrane that stops the flow of solute molecules only. When a solution and the pure solvent used in making that solution are placed on either side of a semipermeable membrane, it is found that more solvent molecules flow out of the pure solvent side of the membrane than solvent flows into the pure solvent from the solution side of the membrane. That flow of solvent from the pure solvent side makes the volume of the solution rise. When the height difference between the two sides becomes large enough, the net flow through the membrane ceases due to the extra pressure exerted by the excess height of the solution chamber. Converting that height of solvent into units of pressure (by using the ) gives a measure of the osmotic pressure exerted on the solution by the pure solvent. P stands for pressure, r is the density of the solution, and h is the height of the solution.
shows a typical setup for measuring the osmotic pressure of a solution.
Figure Setup for Measuring the Osmotic Pressure of a Solution
You can understand why more molecules flow from the solvent chamber to the solution chamber in analogy to our discussion of Raoult’s law. More solvent molecules are at the membrane interface on the solvent side of the membrane than on the solution side. Therefore, it is more likely that a solvent molecule will pass from the solvent side to the solution side than vice versa. That difference in flow rate causes the solution volume to rise. As the solution rises, by the pressure depth equation, it exerts a larger pressure on the membrane’s surface. As that pressure rises, it forces more solvent molecules to flow from the solution side to the solvent side. When the flow from both sides of the membrane are equal, the solution height stops rising and remains at a height reflecting the osmotic pressure of the solution.
The equation relating the osmotic pressure of a solution to its concentration has a form quite similar to the ideal gas law:
Although the above equation may be more simple to remember, the is more useful. This form of the equation has been derived by realizing that/ V gives the concentration of the solute in units of molarity, M.
The water – main component of human organisms and also is part of medium, in which lives the people. The main water property is solved a lot of matters with formatted solutions.
The water in organisms of the person, animal, plant is by its constituent (in a yumrn’s organism about 70 -80 % of water), solvent, and also participates in exchange reactions of matters (hydrolysis, hydration, swelling (turgescence), digestion). It executes a role of a transport system in processes of a feeding, carry of enzymes, products of a metabolism, gases, antibodies. The water is supported a condition to a homeostasis in an organism of the person (acid – alkaline, osmotic, hemodinamicil, thermal equilibrium). The water is indispensable for secrets iones, maintenance of a turgor of cages.
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The colligative properties that we will consider in this and the next unit apply to to solutions in which the solute is non-volatile; that is, it does not make a significant contribution to the overall vapor pressure of the solution. Solutions of salt or sugar in water fulfill this condition exactly. Other solutes that have very small vapor pressures, such as iodine or ethylene glycol antifreeze, can often be considered nonvolatile in comparison to the solvent at the same temperature. Solutions in which both components possess significant vapor pressures, such as alcohol in water, will be treated in another section farther on.
Difference between Diffusion and Osmosis. The main points of difference between diffusion and osmosis nау be summed up as given below:
Osmosis:
1. In osmosis, а semi-permeable membrane is used.
2. In this process, there is only flow of solvent molecules and that too through the semi-permeable membrane.
3. It takes place from lower concentration to higher concentration.
4. It applies to solutions only.
5. It can be stopped or reversed by applying pressure on the solution with higher concentration.
Difference:
1. In diffusion, no semi-permeable membrane is used.
2. In this process, the solvent as well as the solute molecules move directly into each other.
3. It takes place from higher concentration со lower concentration.
4. It takes place in gases as well as solutions.
5. It cannot be stopped or reversed.
Semi-permeable membranes. The semi-permeable membranes (as defined above) are of two types:
(i) Natural semi-permeable membranes е.g vegetable membranes and animal membranes which are found just under the outer skin of the animals and plants. The pig’ s bladder is the most common animal membrane used.
(ii) Artificial semi-permeable membranes. The well known examples of the artificial semi-permeable membranes are parchment paper, cellophane and certain freshly precipitated inorganic substances е.у. copper ferrocyanide, silicates, of iron, cobalt, nickel etc. The precipitated substances have to be supported on some material and this is achieved by preparing the precipitate in the walls of а porous pot.
Measurement of osmotic pressure
The osmotic pressure of а solution can be measured by many methods, but only two methods will be described.
1. Pfeffer’s method – А very simple apparatus was used by Pfeffer for this purpose. А battery pot with, а semipermeable membrane deposited in its wall is cemented to, а wide glass tube which ends in а thin tube at the top and carries а manometer in the side. The manometer is closed at its upper end and is filled with Hg and N2. The solution under investigation is introduced into the pot through this tube, The apparatus is then made airtight by sealing off the tube at the top. А portion of the pot is immersed in distilled water kept at. а constant temperature. In the course of а few days, the manometer registers the maximum pressure, which is the osmotic pressure of the solution.
2. Freezing point determination method – It ha been found that there is а decrease of 1.8б0С in the freezing point of а solution when its osmotic pressure is0equal to one osmole. This method is much more rapid and accurate than Pfeffer’s method. A special apparatus is used to determine the freezing point о f the solution under investigation which is then compared with freezing point of the pure. solvent.
The decrease in the freezing point of the solution is one of the colligative properties of colloidal solutions. The other colligative properties e.g. elevation of boiling point and depression of the vapor density can also be used in the determination of the osmotic pressure of а solution.
Laws of osmotic pressure – These are the same as gas laws and apply to dilute solutions which occur in the living body.
1. The osmotic pressure is directly proportional to the concentration of the solute. For example, 1 % NaCI solution will have double the osmotic pressure of 0.5 % NaCl solution. The osmotic pressure of а solution is dependent upon only the number of dissolved or dissociated particles per unit volume and is independent of chemcial nature of particles. Thus а sodium ion (at. wt. 23), а molecule of glucose (mol. wt. 180) and а molecule of serum albumin (mol. wt. about 70,000) will exert equal osmotic pressure.
2. The osmotic pressure of а solution is directly proportional to the absolute temperature. For dilute solutions, the osmotic pressure is equal to the value CRT, where С = molar concentration, R = gas constant (0.082 liter atmosphere per degree per mole) and T = absolute temperature. А molar solution of a non-electrolyte (е.g. glucose or urea) at 00С (or 2730 absolute) will exert an osmotic pressure equal to CRT = 1 х 0.082 x 273 = 22.4 atmospheres = 22.4 х 760 mm Hg = 17024 mm Hg.
VALUE OSMOS IN BIOLOGICAL PROCESSES
The blood, lymph and also all intercellular lymphs alive organisms is by aqueous solutions of moleculas and ions of many matters – organic and mineral. These solutions have definite osmotic pressure. So, the osmotic pressure of a blood of the person is value a constant and equally 7,4 105 – 7,8 105 Pa(pascal). Such high value osmotic pressure in a blood is conditioned by availability in her of a plenty of ions. High-molecular connection, mainly, proteins (albumines, the globulins), introduce 0,5 % common osmotic pressure of a blood. This part of osmotic pressure of a blood call oncotical as pressure, the value which one is equal 3,5-3,9 kPa. Oncotical pressure has large value for alive organisms. At a decrease oncotical of pressure the water goes in the party by high pressure – in a tissue, producing so called oncotical edemas of a hypodermic fat.
The osmotic pressure of a blood of the person is responded osmomolar concentration Dissoluble in plasma of matters, which one equal 0,287 – 0,0303 mol / liter.
Solutions with osmotic pressure, which is equal osmotic pressure of standard solution, is called isotonic. The solutions with osmotic by pressure are called as maximum for standard, hypertonic, and solutions with the lowest osmotic pressure hypotonic. In medical practice isotonic call solutions with osmotic pressure equal to osmotic pressure of a blood plasma. Such solution is 0, 89 % solution of sodium salt, and also 4,5 -5 % solution of a glucose. The isoosmotic solutions can be entered into an organism of the person in plenties. The hypertonic salt solutions enter in an person’s organism only in small amounts. At the introducing of a plenty hypertonic of solution the erythrocytes owing to loss of water decrease in volument and shrivel. Such phenomenon is called as a plasmolysis.
Importance of osmotic pressure of plasma proteins – The plasma proteins form а colloidal solution and are the chief colloid of the plasma. The osmotic pressure of plasma proteins which is called oncotic pressure though negligible (25 to 30 mm Hg) as compared to that of plasma crystalloids (about 5,000 mm Hg) is the main force which tends to keep the plasma water within the blood vessels. This is so because the concentration of non-colloids is almost the same in plasma and in the extracellular fluid and the osmotic pressure exerted by the non-colloids on the outside and the inside of capillaries are therefore balanced by each other. If the concentration of plasma proteins decreases markedly, water leaks into tissue spaces and the pathological condition called edema results.
Isosmotic, isotonic, hyposmotic, hypotonic, hyperosmot1c and hypertonic solutions – Isosmotic solutions are those which have the same osmotic pressure; 0.15 molar or about 0.90 % NaC1 solution in water is isosmotic with the human blood plasma. If human red blood cells are placed in this solution, they remain intact and retain their original shape and volume. 0.90% NaC1 solution is therefore isosmotic as well as isotonic with red blood cells because in this case0the amount of water entering the cells is equal to that leaving them; thus there is no net gain or loss of water by the cells.
Osmotic pressure creates some critical problems for living organisms. Cells typically contain fairly high concentrations of solutes, that is, small organic molecules and ionic salts, as well as lower concentrations of macromolecules. If cells are placed in а solution that has an equal concentration of solute, there will be no net movement of water in either direction. Such solutions are called isotonic. For example, red blood cells are isotonic to а 0.9% NaCI solution. When cells are placed in а solution with а lower solute concentration (i.е., а hypotonic solution), water will move into the cells. Red blood cells, for example, will swell and rupture in а process called hemolysis when they are immersed in pure water. In hypertonic solutions, those with higher solute concentrations, cells shrivel because there is а net movement of water out of the cell. The shrinkage of red blood cells in hypertonic solution (е.g., а 3% NaCI solution) is referred to as crenation.
Fig. Osmotic Pressure and Plant Cells. (а) Isotonic solutions cause no changes in cell volume. (b) Plant cells typically exist in а hypotonic environment. Water enters these cells, and they become swollen. Cell bursting is prevented by the restraining force of rigid cell walls. (с) In а hypertonic environment the cell membrane pulls away from the cell wall because of water loss. This is the reason plants wilt when they receive insufficient water.
a) b) c)
References:
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Prepared by PhD Falfushynska H.
