Boiling

Boiling water.

Boiling, a type of phase transition, is the rapid vaporization of a liquid, which typically occurs when a liquid is heated to its boiling point, the temperature such that its vapor pressure is above that of the surroundings, such as atmospheric pressure. Thus, a liquid may also boil when the pressure of the surrounding atmosphere is sufficiently reduced, such as the use of a vacuum pump or at high altitudes. Boiling occurs in three characteristic stages, which are nucleate, transition and film boiling. These stages generally take place from low to high surface temperatures, respectively.

Nucleate boiling is characterized by the incipience and growth of bubbles on a heated surface, which rise from discrete points on a surface, whose temperature is only slightly above the liquid’s saturation temperature. In general, the number of nucleation sites are increased by an increasing surface temperature. An irregular surface of the boiling vessel (i.e. increased surface roughness) can create additional nucleation sites, while an exceptionally smooth surface, such as glass, lends itself to superheating. Under special conditions, a heated liquid may show boiling delay when heated over its boiling point, by starting to boil suddenly and violently.

When the surface temperature reaches a maximum value, the critical superheat, vapor begins to form faster than liquid can reach the surface. Thus, the heated surface suddenly becomes covered with a vapor layer. Because of the vapor layer’s lower thermal conductivity, this vapor layer insulates the surface. This condition of a vapor film insulating the surface from the liquid characterizes film boiling.

Transition boiling may be defined as the unstable boiling, which occurs at surface temperatures between the maximum attainable in nucleate and the minimum attainable in film boiling.

Boiling point

The boiling point of a substance is the temperature at which it can change its state from a liquid to a gas throughout the bulk of the liquid at a given pressure. A liquid may change to a gas at temperatures below the boiling point through the process of evaporation. Any change of state from a liquid to a gas at boiling point is considered vaporization. However, evaporation is a surface phenomenon, in which only molecules located near the gas/liquid surface could evaporate. Boiling on the other hand is a bulk process, so at the boiling point molecules anywhere in the liquid may be vaporized, resulting in the formation of vapor bubbles.

A somewhat clearer (and perhaps more useful) definition of boiling point is "the temperature at which the vapor pressure of the liquid equals the pressure of the surroundings."

The reaction

Something that should not emembered is that boiling is evidenced by the appearance of bubbles containing vapor from the liquid. [Note: The bubbles that precede real boiling in the pot on the stove are either (formerly) dissolved gas or water vapor forming on the very hot bottom of the pot that will be condensed before it can get to the top of the liquid.] Production of vapor requires energy and thus does not occur without some source of energy. This source can be a hot surface or even the liquid itself. Hot liquid will boil as it rises through the bulk liquid if the pressure of the environment drops to the vapor pressure of the liquid at its temperature. This production of vapor will not quickly stop because the temperature of the liquid will not be reduced by the vaporization thus reducing the vapor pressure.

Saturation temperature and pressure

A saturated liquid or saturated vapor contains as much thermal energy as it can without boiling or condensing.

Saturation temperature means boiling point. The saturation temperature is the temperature for a corresponding saturation pressure at which a liquid boils into its vapor phase. The liquid can be said to be saturated with thermal energy. Any addition of thermal energy results in a phase change.

If the pressure in a system remains constant (isobaric), a vapor at saturation temperature will begin to condense into its liquid phase as thermal energy (heat) is removed. Similarly, a liquid at saturation temperature and pressure will boil into its vapor phase as additional thermal energy is applied.

The boiling point corresponds to the temperature at which the vapor pressure of the substance equals the ambient pressure. Thus the boiling point is dependent on the pressure. Usually, boiling points are published with respect to standard pressure (101.325 kilopascals or 1 atm). At higher elevations, where the atmospheric pressure is much lower, the boiling point is also lower. The boiling point increases with increased ambient pressure up to the critical point, where the gas and liquid properties become identical. The boiling point cannot be increased beyond the critical point. Likewise, the boiling point decreases with decreasing ambient pressure until the triple point is reached. The boiling point cannot be reduced below the triple point.

Saturation Pressure, or vapor point, is the pressure for a corresponding saturation temperature at which a liquid boils into its vapor phase. Saturation pressure and saturation temperature have a direct relationship: as saturation pressure is increased so is saturation temperature.

If the temperature in a system remains constant (an isothermal system), vapor at saturation pressure and temperature will begin to condense into its liquid phase as the system pressure is increased. Similarly, a liquid at saturation pressure and temperature will tend to flash into its vapor phase as system pressure is decreased.

Latent heat

The process of changing from a liquid to a gas requires an amount of heat called the latent heat of vaporization. As heat is added to a liquid at its boiling point, all of this heat goes toward the phase change from liquid to gas, thus the temperature of the substance remains constant even though heat has been added. The word latent, which comes from Latin and means hidden, is used to describe this "disappearing" heat that is added, but doesn't result in an increase in temperature. Since heat is added with no corresponding change in temperature, the heat capacity of the liquid is essentially infinite at the boiling point.

Intermolecular interactions

In terms of intermolecular interactions, the boiling point represents the point at which the liquid molecules possess enough heat energy to overcome the various intermolecular attractions binding the molecules into the liquid (eg. dipole-dipole attraction, instantaneous-dipole induced-dipole attractions, and hydrogen bonds). Therefore the boiling point is also an indicator of the strength of these attractive forces.

The boiling point of water is 100 °C (212 °F) at standard pressure. On top of Mount Everest the pressure is about 260 mbar (26 kPa) so the boiling point of water is 69 °C.

For purists with a knowledge of thermodynamics, the normal boiling point of water is 99.97 degrees Celsius (at a pressure of 1 atm, i.e. 101.325 kPa). Until 1982 this was also the standard boiling point of water, but the IUPAC now recommends a standard pressure of 1 bar (100 kPa). At this slightly reduced pressure, the standard boiling point of water is 99.61 degrees Celsius.

Properties of other elements

The element with the lowest boiling point is helium. Both the boiling points of rhenium and tungsten exceed 5000 K at standard pressure. Due to the experimental difficulty of precisely measuring extreme temperatures without bias, there is some discrepancy in the literature as to whether tungsten or rhenium has the higher boiling point.

Superheating (boiling delay)

In physics, superheating (sometimes referred to as boiling retardation, boiling delay, or defervescence) is the phenomenon in which a liquid is heated to a temperature higher than its standard boiling point, without actually boiling. This can be caused by rapidly heating a homogeneous substance while leaving it undisturbed (so as to avoid the introduction of bubbles at nucleation sites).

Because a superheated fluid is the result of artificial circumstances, it is metastable, and is disrupted as soon as the circumstances abate, leading to the liquid boiling very suddenly and violently (a steam explosion). Superheating is sometimes a concern with microwave ovens, some of which can quickly heat water without physical disturbance. A person agitating a container full of superheated water by attempting to remove it from a microwave could easily be scalded.

Superheating is common when a person puts an undisturbed cup of water into the microwave and heats it. Once finished, the water appears to have not come to a boil. Once the water is disturbed, it violently comes to a boil. This can be simply from contact with the cup, or the addition of substances like instant coffee or sugar, which could result in hot scalding water shooting out. The chances of superheating is more common with smooth containers, such as brand-new glassware that lacks any scratches (scratches can house small pockets of air, which can serve as a nucleation point.

Rotating dishes in modern microwave ovens can also provide enough perturbation to prevent superheating.

There have been some injuries by superheating water, such as when a person makes instant coffee and adds the the coffee to the superhated water. This sometimes results in an "explosion" of bubbles. There are some ways to prevent superheating in your microwave, such as ptting a popsicle stick in the glass, or having a scratched container to cook the water in.

Boiling point elevation

Boiling-point elevation is a colligative property that states that a solution will have a higher boiling point than that of a pure solvent. Based on this knowledge, it is often thought that the addition of salt to water when cooking food will significantly elevate the boiling point of the water. That view, however, is mistaken. The amount of salt added when cooking is generally not enough to raise the temperature by a single degree. The salt is added simply to season the food and prevent pasta from sticking.

Scope restriction

Milk and water with starch content does not boil over because of superheating, but because of extreme foam build up. This foam is stabilised by special substances in the liquids and therefore does not burst.

Boiling in cookery

In cookery, boiling is cooking food in boiling water, or other water-based liquid such as stock or milk. Simmering is gentle boiling, while in poaching the cooking liquid moves but scarcely bubbles.

In places where the available water supply is contaminated with disease-causing bacteria, boiling water and allowing it to cool before drinking it is a valuable health measure. Boiling water for a few minutes kills most bacteria, amoebas, and other microbial pathogens. It thus can help prevent cholera, dysentery, and other diseases caused by microorganisms.

The temperature of a substance is constant as it undergoes a phase transition. Therefore, increasing the temperature of a liquid already boiling by increasing the rate of heat transfer is impossible, it will just boil more quickly. Once it has turned into steam, water will increase in temperature as heat is applied to it. Pressure and a change in composition of the liquid may alter the boiling point of the liquid. For this reason, high elevation cooking generally takes longer since boiling point is a function of atmospheric pressure. In Denver, Colorado, which is at an elevation of about one mile, water boils at approximately 95 C. [1] Depending on the type of food and the elevation, the boiling water may not be hot enough to cook the food properly. The boiling point is defined as the temperature at which the vapor pressure of the substance equals the pressure above the substance. Increasing the pressure as in a pressure cooker raises the temperature of the contents above the open air boiling point. Adding a water soluble substance, such as salt or sugar also increases the boiling point. This is called boiling-point elevation. However, the effect is very small, and the boiling point will be increased by an insignificant amount. On the other hand, salt or ethylene glycol can cause significant freezing point depression. Due to variations in composition and pressure, the boiling point of water is almost never exactly 212 F / 100 C, but rather close enough for cooking.

Foods suitable for boiling include:

• Fish
• Vegetables
• Farinaceous foods such as pasta
• Eggs
• Meats
• Sauces
• Stocks and soups

Advantages:

• Older, tougher, cheaper joints of meat and poultry can be made digestible
• It is appropriate for large-scale cookery
• Nutritious, well flavoured stock is produced
• It is safe and simple
• Maximum colour and nutritive value is retained when cooking green vegetables, provided boiling time is kept to the minimum

Disadvantages:

• There is a loss of soluble vitamins in the water
• It can be a slow method
• Foods can look unattractive

Boiling can be done in two ways: The food can be placed into already rapidly boiling water and left to cook, the heat can be turned down and the food can be simmered; or the food can also be placed into the pot, and cold water may be added to the pot. This may then be boiled until the food is satisfactory.

See also

• Heat
• Melting
• Phase (matter)
• Temperature
• List of elements by boiling point
• Leidenfrost effect
• critical temperature
• triple point
• Supersaturation
• Critical point
• Microwaving
• Autoclave

References

ISBN links support NWE through referral fees

DeVoe, Howard, Thermodynamics and Chemistry. Prentice-Hall, 2001

External links

• Explanation of concepts of saturated vapor pressure, evaporation, boiling, boiling point.
• WebElements.com - Boiling point
• Boiling point of water
• Video of superheated water in a microwave explosively flash boiling, why it happens, and why it's dangerous.
• A series of superheated water with oil film experiments done in the microwave by Louis A. Bloomfield, physics professor at the University of Virginia. Experiment #13 proceeds with surprising violence.

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