Boiling point elevation The increase

Boiling point elevation The increase in the boiling point of a solvent caused by dissolution of a nonvolatile solute. 

Boiling point elevation (ATb) Increase in the boiling point caused by addition of a nonvolatile solute, 269-271 Bomb calorimeter Device used to measure heat flow, in which a reaction is carried out within a sealed metal container, 202-203... 

Because the presence of a nonvolatile solute lowers the vapor pressure of the solvent, the boiling point of the solvent rises. This increase is called boiling-point elevation. The elevation of the boiling point has the same origin as vapor-pressure lowering and is also due to the effect of the solute on the entropy of the solvent. 

Many molecular parameters, such as ionization, molecular and electronic structure, size, and stereochemistry, will influence the basic interaction between a solute and a solvent. The addition of any substance to water results in altered properties for this substance and for water itself. Solutes cause a change in water properties because the hydrate envelopes that are formed around dissolved molecules are more organized and therefore more stable than the flickering clusters of free water. The properties of solutions that depend on solute and its concentration are different from those of pure water. The differences can be seen in such phenomena as the freezing point depression, boiling point elevation, and increased osmotic pressure of solutions. 

The coUigative properties of antifreeze chemicals may also result in boiling point elevation. As the chemical is added to water, the boiling point of the mixture increases. Unlike the freeze depression, the boiling elevation does not experience a maximum the boiling point versus concentration curve is a smooth curve that achieves its maximum at the 100% antifreeze level. The boiling point elevation can be another important characteristic for antifreeze fluids in certain heat-transfer appHcations. 

The vapor pressure of a liquid dictates when a substance will boil. In fact, the boiling point of a substance is defined as the temperature at which the vapor pressure equals the external pressure. Typically, the external pressure is equal to atmospheric pressure, and we define the normal boiling point as the temperature when the vapor pressure equals 1 atmosphere. If we consider water heated on a stove, the bubbles that develop in the liquid contain water vapor that exerts a pressure at the specific vapor pressure of water at that temperature. For example, when water reaches 60°C, any bubbles that form will contain vapor at 149 mm Hg (see Table 9.4). At this pressure, and any other pressure below 760 mm Hg (1 atmosphere), the external pressure of 1 atmosphere causes the bubbles to immediately collapse. As the temperature of the water rises, the vapor pressure continually increases. At 100°C, the vapor pressure inside the bubbles finally reaches 760 mm Hg. The vapor pressure is now sufficient to allow the bubbles to rise to the surface without collapsing. At higher elevations where the external pressure is lower, liquids boil at a lower temperature. At the top of a 15,000-foot peak, water boils at approximately 85°C rather than 100°C. This increases the cooking time for items, as noted in the directions of many packaged food. If the external pressure is increased, the boiling temperature also increases. This is the concept behind a pressure cooker. The sealed cooker allows pressure to build up inside it... 

Boiling Point Elevation increase in normal boiling point of a pure liquid due to the presence of solute added to the liquid... 

The elevation of boiling point in this case is much greater than for any of the other systems. The assumptions that the heat term in the Gibbs-Duhem equation can be neglected and that the effect of the salt can be expressed in terms of its effect on the vapor pressure of each solvent independently become less viable as the boiling point elevation increases. 

Even with this unequal distribution there may be little effect on yield of distillate from a substantially fresh water feed hence the high output of the still from distilled water feed. With sea water, 3 to 4% NaCl equivalent, the average or effective boiling point elevation becomes unequal on the two rotors. Thus if a 50% cut is secured and the lower rotor receives twice the feed of the upper, the average residue concentrate of 7% brine from 3.5% feed could be an actual 10% from the upper periphery and 5% from the lower, supposing equal rates of distillation. Actually because of -the different elevations of boiling point (1.1° and 1.8° F.) the rate of evaporation from the upper rotor decreases while that from the lower rotor increases but less than proportionally because of the added thickness of the feed layer. Later experiments at Columbus on the No. 4 machine suggest that this situation existed in the No. 5 still. 

The increased solubility of quartz in basic organic solvent systems appears to be caused by aqueous potassium hydroxide reaction at temperatures above the boiling point of the aqueous system alone. The organic solvent fraction serves as a substrate which permits attainment of elevated temperatures. Increasing the pressure at which basic aqueous reactions are performed would serve as an alternative method which would eliminate the need for addition of organic solvents. This prospect is especially attractive for in situ removal of silicates from oil shale since geothermal gradient and overburden may provide the elevated temperature and pressure necessary for efficient silicate removal. 

Vapor-pressure lowering, boiling-point elevation, and freezing-point depression are very similar thermodynamically. For example, the increase in boiling point ATh is interpreted thermodynamically by using the boiling-point elevation constant Kb to obtain the molality of the solution, as stated in the equation... 

The final colligative property, osmotic pressure,24-29 is different from the others and is illustrated in Figure 2.2. In the case of vapor-pressure lowering and boiling-point elevation, a natural boundary separates the liquid and gas phases that are in equilibrium. A similar boundary exists between the solid and liquid phases in equilibrium with each other in melting-point-depression measurements. However, to establish a similar equilibrium between a solution and the pure solvent requires their separation by a semi-permeable membrane, as illustrated in the figure. Such membranes, typically cellulosic, permit transport of solvent but not solute. Furthermore, the flow of solvent is from the solvent compartment into the solution compartment. The simplest explanation of this is the increased entropy or disorder that accompanies the mixing of the transported solvent molecules with the polymer on the solution side of the membrane. Flow of liquid up the capillary on the left causes the solution to be at a hydrostatic pressure... 

---