Boiling

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Boiling water

Boiling is the process of rapidly converting a liquid to its gaseous (vapor) state, typically by heating the liquid to a temperature called its boiling point. The boiling of a substance is known as a phase change or phase transition. Chemically, the substance remains the same, but its physical state (or "phase") changes.

The boiling point of a substance is the temperature at which it can change its state from liquid to gas throughout the bulk of the liquid at a given pressure. It should be noted that the boiling point of a substance is sensitive to the pressure of the surroundings. Thus, for example, the boiling point of water is lower at a high altitude than it is at sea level, because the air pressure at high altitudes is lower than that at sea level. Based on this understanding, the boiling point of a substance can be defined as the temperature at which the vapor pressure of the liquid substance is equal to the pressure of the surrounding gases.

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Comparing boiling with evaporation

Any change of state from a liquid to a gas is considered vaporization. If this change of state takes place below the boiling point of the liquid, it is called evaporation. Evaporation, however, is a surface phenomenon in which only those molecules located near the gas/liquid interface can evaporate. Boiling, on the other hand, is a bulk process, which means that at the boiling point, molecules anywhere in the liquid may be vaporized, resulting in the formation of bubbles of vapor throughout the liquid.[1]

The 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.

Latent heat of vaporization

When a liquid is heated, its temperature will rise until it reaches the boiling point of the liquid. If more heat is supplied, it goes toward the phase change from liquid to gas, while the temperature remains constant. The heat required to change 1 gram of the liquid to the gaseous phase (at a particular pressure) is called the latent heat of vaporization. The word "latent" is derived from a Latin word that means "hidden," implying that at the boiling point, the heat added to the liquid seems to disappear, without raising the temperature of the liquid.

Understanding boiling on a molecular level

The molecules within a liquid interact with one another with various attractive forces, including what are called hydrogen bonds and dipole-dipole attractions. The boiling point represents the temperature at which the liquid molecules possess enough heat energy to overcome the various intermolecular attractions that bind the molecules into the liquid. Therefore the boiling point is also an indicator of the strength of these attractive forces.

Boiling points of water and some elements

The boiling point of water is 100 °C (212 °F) at standard pressure. Strictly speaking, 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 °C.

The boiling point of water (or other liquid) can be reduced by lowering the pressure of the surrounding gases, such as by using a vacuum pump or by going to high altitudes. On top of Mount Everest, for example, the pressure is about 260 mbar (26 kPa), so the boiling point of water is 69 °C. Conversely, the boiling of water is higher in a pressure cooker because there is greater pressure within the cooker.

The element with the lowest boiling point is helium. Both the boiling points of rhenium and tungsten exceed 5,000 Kelvin (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.

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.

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 chance 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 coffee to the superheated water. This sometimes results in an "explosion" of bubbles. There are some ways to prevent superheating in your microwave, such as putting 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 stabilized 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 of boiling:

  • Older, tougher, cheaper joints of meat and poultry can be made digestible
  • It is appropriate for large-scale cookery
  • Nutritious, well-flavored stock is produced
  • It is safe and simple
  • Maximum color 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


Notes

  1. The bubbles that precede real boiling in the pot on the stove are either previously dissolved gas or water vapor forming on the very hot bottom of the pot that will condense before reaching the top of the liquid.

References

  • DeVoe, Howard. Thermodynamics and Chemistry. Upper Saddle River, NJ: Prentice-Hall, 2001. ISBN 0023287411

External links

All links retrieved February 15, 2013.

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