Elevation of boiling point.

Introduction

The increase in the boiling point of a solvent caused by the addition of a solute is referred to as boiling point elevation. When a solute and a non-volatile solvent are combined, the resulting solution has a higher boiling point than the pure solvent. Boiling point elevation is a colligative property of matter, which means it is affected by the solute-to-solvent ratio but not by the identity of the solute. This means that the amount of solute added to a solution causes the boiling point to rise. The higher the solute concentration in the solution, the higher the boiling point elevation.

Why Does Boiling Point Elevation Occur?

A liquid’s boiling point is the temperature at which its vapour pressure equals the pressure in its surrounding environment. Non-volatile substances do not evaporate easily and have very low vapour pressures (assumed to be zero).

As a result, more heat must be applied to the solution for it to boil. The boiling point elevation is the increase in the boiling point of the solution. An increase in the concentration of added solute results in a further decrease in the solution’s vapour pressure and an increase in the solution’s boiling point.

Boiling point elevation formula

The boiling purpose of a non-volatile substance answer is expressed as follows:

The boiling point of a solution is equal to the boiling point of the pure solvent plus the boiling point elevation (ΔTb).

The increase in boiling purpose (ΔTb) is proportional to the concentration of the substance within the solution.

The equations are.

ΔTb = i*Kb*m

Where,

• Kb is the ebullioscopic constant
• m is the molality of the solute
• i is the Van’t Hoff factor

It is important to note that when the concentration of the solute is very high, this formula becomes less precise. Furthermore, this formula does not apply to volatile solvents.

The ebullioscopic constant (Kb) is frequently expressed in oC/molal or oC.kg.mol-1. The values of Kb for some common solvents are shown in the table below.

The boiling point elevation formula can be used to calculate the degree of dissociation of the solute as well as the molar mass of the solute.

Elevation in Boiling Point Derivation

A liquid’s boiling point is an important physical property. A liquid’s boiling point is the temperature at which its vapour pressure equals atmospheric pressure (1 atm). When a nonvolatile solute is added to a pure solvent at its boiling point, the solution’s vapour pressure falls below 1 atm. The temperature of the solution must be raised to restore the vapour pressure to 1 atm. As a result, the solution boils at a higher temperature (Tb) than the pure solvent’s boiling point (Tb°). This rise in the boiling point is referred to as the boiling point elevation. 

Suppose the vapour pressure of the solution increases as the temperature rises. The violet-colored curve depicts the variation of vapour pressure with temperature for pure water. Water has a vapour pressure of 1 atm at 100 0C. As a result, the boiling point of water is 1000 C (Tb°). When a solute is added to water, the resulting solution’s vapour pressure decreases. The green curve represents the variation of vapour pressure concerning temperature for the solution. The graph shows that the solution’s vapour pressure is equal to 1 atm pressure at Tb temperatures greater than Tb°. The difference between these two temperatures ( Tb–Tb°) gives the boiling point elevation.

The elevation of the boiling point (ΔTb)= Tb – Tb°

The increase in boiling point is proportional to the concentration of solute particles.

ΔTb α m (9.23)

m denotes the concentration of a solution in molality.

ΔTb = Kb m (9.23)

Where,

Kb= is the molal boiling point elevation constant, also known as the Ebullioscopic constant.

If m=1, then ΔTb=Kb;

As a result, the Kb is equal to the increase in boiling point for one molal solution. The following expression is used to calculate Kb.

Boiling Point Elevation Definition

The phenomenon of boiling point elevation occurs when the boiling point of a liquid (a solvent) is increased by the addition of another compound, resulting in a solution with a higher boiling point than the pure solvent. When a non-volatile solute is added to a pure solvent, the boiling point rises.

While the number of dissolved particles in a solution affects boiling point elevation, their identity is unimportant. Interactions between solvents and solutions do not affect boiling point elevation.

An ebullioscope is used to accurately measure boiling point and thus detect whether or not boiling point elevation has occurred, as well as how much the boiling point has changed.

Examples Of Boiling Point Elevation

Calculate the boiling point of a 3.5% solution (by weight) of sodium chloride in water.

1 kg of the given solution contains 0.035kg of NaCl and 0.965kg of H2O. Since the molar mass of NaCl is 58.5, the number of moles of NaCl in 1 kg of the solution is:

(35g)/(58.5g.mol-1) = 0.598 moles

The molality of NaCl in 1kg of the solution can be calculated as:

m = (0.598mol)/(0.965 kg) = 0.619 molal

The boiling point elevation constant of water is 0.512 oC.kg/molal. Since NaCl dissociates into 2 ions, the Van’t Hoff factor for this compound is 2. Therefore, the boiling point elevation (ΔTb) can be calculated as follows:

ΔTb = 2*(0.52oC/molal)*(0.619 molal) = 0.643oC

Boiling point of the solution = boiling point of pure solvent + boiling point elevation

= 100oC + 0.643oC = 100.643oC

Therefore, the boiling point of the 3.5% NaCl solution is 100.643oC.

Conclusion

 The addition of heat causes the transformation of the liquid into its vapour without raising the temperature.  At any temperature, a liquid partially vapourises in the space above it until the pressure exerted by the vapour reaches a characteristic value called the vapour pressure of the liquid at that temperature. As the temperature increases, the vapour pressure increases; at the boiling point, vapour bubbles form inside the liquid and rise to the surface. The boiling point of a liquid varies according to the pressure applied; the normal boiling point is the temperature at which the vapour pressure is equal to the normal atmospheric pressure at sea level (760 mm [29.92 in] of mercury). At sea level, the water boils at 100 ° C (212 ° F). At higher altitudes, the boiling point temperature is lower.