Introduction:

Solutions are the homogeneous mixtures which include two or more substances. The main component is termed as the solvent and the minor component(s) are termed as the solute(s). As the solution is mainly comprised of solvent, the physical properties of the solution look like those of the pure solvent. Though, some of these physical properties, termed as Colligative properties, are independent of the nature of the solute and based only on the concentration of solute particles. Illustrations of Colligative properties comprise vapor pressure reduction, boiling point elevation, freezing point depression and osmotic pressure.

The Colligative properties are those properties of solutions which based on the number of dissolved particles in solution, however not on the identities of the solutes. For illustration, the freezing point of salt water is lower than that of pure water, due to the presence of salt dissolved in the water. To a fine approximation, it doesn't matter whether the salt dissolved in water is sodium chloride or potassium nitrate; if the molar amounts of solute are similar and the number of ions is similar, the freezing points will be the same. For illustration: AlCl₃ and K₃PO₄ would show essentially the similar Colligative properties, as each compound dissolves to generate four ions per formula unit.

To illustrate the difference between the two sets (that is, Colligative and non-Colligative) of solution properties, we will compare the properties of a 1.0 M aqueous sugar solution to a 0.5 M solution of table salt (NaCl) in water. In spite of the concentration of sodium chloride being half of the sucrose concentration, both the solutions have exactly the similar number of dissolved particles as each sodium chloride unit makes two particles on dissolution-a sodium ion, Na⁺, and a chloride ion, Cl^-. Thus, any difference in the properties of such two solutions is due to the non-Colligative property. Both the solutions encompass the similar freezing point, boiling point, vapor pressure and osmotic pressure as those Colligative properties of a solution based only on the number of dissolved particles. The taste of two solutions, though, is markedly different. The sugar solution is sweet and the salt solution tastes salty. Thus, the taste of the solution is not a Colligative property. The other non-Colligative property is the color of a solution. A 0.5 M solution of CuSO₄ is bright blue in contrast to the colorless salt and sugar solutions. The other non-Colligative properties comprise surface tension, viscosity and solubility.

The Colligative properties don't based on the identity of either the solvent or the solute(s) particles (type, size or charge) in the solution.

Colligative properties are:

1) Vapor pressure lowering

2) Boiling point elevation

3) Freezing point Depression

4) Osmotic pressure.

As such properties yield information on the number of solute particles in solution; one can utilize them to get the molecular weight of the solute.

Vapor Pressure Lowering:

The vapor-pressure of a liquid is the equilibrium pressure of the gas molecules from that liquid (that is, the results of evaporation) above the liquid itself. The glass of water put in an open room will evaporate fully (and therefore never reach equilibrium); though, if a cover is put on the glass, the space above the liquid will ultimately include a constant amount of water vapor. How much water vapor is present based on the temperature, however not on the amount of liquid which is present at equilibrium (given some liquid is present at the equilibrium). (At room temperature, the vapor-pressure of pure water is around 20 Torr, which is around one-fourth of the net atmospheric pressure on a 'normal' day at sea level.)

If, rather than pure water, an aqueous solution is positioned in the glass, the equilibrium pressure will be lower than it would be for pure water. Raoult's law defines that the vapor pressure of the solvent over the solution is proportional to the fraction of solvent molecules in the solution; that is, if two thirds of the molecules are solvent molecules, the vapor pressure because of the solvent is around two-thirds of what it would be for pure solvent at that temperature. If the solute consists of a vapor pressure of its own, then the net vapor pressure over the solution would be:

P_vap = 2/3 (Pure solvent vapor pressure) + 1/2 (Pure solute vapor pressure)

Usually, one anticipates that the solutes which are liquids in their pure form (like ethyl alcohol) will encompass some vapor pressure of their own, while ionic compounds (like sodium chloride) will not contribute to the net vapor pressure over the solution.

One effect of this lowering in vapor pressure might be noticed in a spilled can of soda. As the water evaporates, the soda becomes more sugar and less water, till the vapor pressure of the water is so low that it hardly evaporates. As an outcome, the spilled soda remains sticky for a long time. Contrast this behavior with that of the water spill.

Raoult's Law:

In the mid-1800s, it was invented that the vapor pressure of a solution was lowered and that the amount was more-or-less proportional to the amount of solute. In the year 1880, Francios Marie Raoult was capable to find out the equation that regulates this property:

P_solution = (χ_solvent) (P°_solvent)

Here,

P_solvent is the vapor-pressure of the solvent over the solution

χ_solvent is the mole fraction of the solvent in the solution

P^o_solvent is the vapor pressure of the pure solvent.

=> Note that:

If χ_solvent = 1 (pure solvent), P_solvent = P^o_solvent, and

If χ_solvent < 1 (solute(s) present), P_solvent < P^o_solvent (that is, the vapor pressure of the solvent above the solution is lower than that of the vapor pressure above the pure solvent).

Boiling point elevation:

The boiling point of solution is higher than that of the pure solvent. Accordingly, the use of a solution, instead of a pure liquid, in antifreeze serves to maintain the mixture from boiling in a hot automobile engine. As by freezing point depression, the effect based on the number of solute particles present in a specific amount of solvent, however not the identity of such particles. If 10 grams (that is, 0.35 ounces) of sodium chloride are dissolved in 100 grams (that is, 3.5 ounces) of water, the boiling point of the solution is 101.7°C (215.1°F; which is 1.7°C (3.1°F) higher than that of the boiling point of pure water). The formula employed to compute the change in boiling point (ΔT_b) relative to the pure solvent is identical to that employed for freezing point depression:

ΔT_b = i K_b m,

Here, K_b is the boiling point elevation constant for the solvent (that is, 0.52°C·kg/mol for water), and 'm' and 'i' encompass the similar meanings as in the freezing point depression formula. Note that ΔT_b symbolizes an increase in the boiling point, while ΔT_f symbolizes a decrease in the freezing point. As by the freezing point depression formula, this one is most precise at low solute concentrations.

Freezing point Depression:

The presence of a solute lowers the freezing-point of a solution relative to that of the pure solvent. For illustration, pure water freezes at 0°C (32°F); if one dissolves 10 grams (0.35 ounces) of sodium chloride (that is, table salt) in 100 grams (that is, 3.53 ounces) of water, the freezing point goes down to - 5.9°C (21.4°F). If one employs sucrose (that is, table sugar) rather than sodium chloride, 10 grams (that is, 0.35 ounces) in 100 grams (3.53 ounces) of water makes a solution having a freezing point of - 0.56°C (31°F). The main reason that the salt solution consists of a lower freezing point than the sugar solution is that there are more particles in 10 grams (0.35 ounces) of the sodium chloride than in 10 grams (0.35 ounces) of sucrose. As sucrose, C₁₂H₂₂O₁₁ consists of a molecular weight of 342.3 grams (12.1 ounces) per mole and sodium chloride consists of a molecular weight of 58.44 grams (2.06 ounces) per mole, 1 gram (0.035 ounces) of sodium chloride consists of almost six times as most of the sodium chloride units as there are sucrose units in a gram of sucrose. Moreover, each and every sodium chloride unit comes apart into two ions (that is, a sodium cation and a chloride anion) whenever dissolved in water. Sucrose is a non-electrolyte that implies that the solution includes whole C₁₂H₂₂O₁₁ molecules. In expecting the expected freezing point of a solution, one should consider not only the number of formula units present, however as well the number of ions which result from each formula unit, in the case of ionic compounds. One can compute the change in freezing point (Δ T_f) relative to the pure solvent by employing the equation:

ΔT_f = i K_f m

Here, 'K_f' is the freezing point depression constant for the solvent (1.86°C·kg/mol for water), 'm' is the number of moles of solute in solution per kilogram of solvent and 'i' is the number of ions present per formula unit (example: i = 2 for NaCl). This formula is estimated; however it works well for the low solute concentrations.

As the presence of a solute lowers the freezing point, most of the communities place salt on their roads after a snowfall, to keep the melted snow from refreezing. As well, the antifreeze employed in the automobile heating and cooling systems is the solution of water and ethylene glycol (or propylene glycol); this solution consists of a lower freezing point than either pure water or pure ethylene glycol.

Osmotic pressure:

Osmosis is a method whereby liquids pass via semi-permeable membranes, spontaneously. This method is driven via changes in entropy. Let consider two containers of liquid, joined by a semi-permeable membrane. The semi-permeable membrane can be simply a device that holes small adequate to allow the solvent to pass via however not the solute. In one container is a pure solvent and in the other is similar solvent however by some solute dissolved in it. In this case, there will be solvent passing via in both the directions arbitrarily. The rate of the two methods based on the relative entropy of the two sides and on the relative pressure exerted to the solutions.

Fig: osmotic pressure

As the solution will encompass higher entropy than the solute, it is reasonable to anticipate that the spontaneous method is the one where the solvent passes via the membrane from the pure solvent side to the solution side where the entropy is higher. This will take place till the difference in the height of the two columns of liquid (h) is large adequate that the pressure caused through this column of liquid precisely stops the total flow of solvent. This pressure is equivalent to the osmotic pressure.

Osmotic pressure is represented by the symbol π (Greek equivalent of P) and the equation associating the osmotic pressure to the amount of solute is:

π V = nRT

Note that n/V is the concentration in mol/liter so we can as well write:

π = C_M,solute RT

π can be computed at room temperature and is thus, more helpful for measuring the solutes that might decompose at higher (boiling point elevation) temperature measurements. It is as well extremely sensitive to small amounts of solute and is thus useful for evaluating very big molar mass values.

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