Boiling point

The boiling point of a substance is the temperature at which the vapor pressure of the liquid equals the pressure surrounding the liquid[1][2] and the liquid changes into a vapor.

The boiling point of a liquid varies depending upon the surrounding environmental pressure. A liquid in a partial vacuum has a lower boiling point than when that liquid is at atmospheric pressure. A liquid at high pressure has a higher boiling point than when that liquid is at atmospheric pressure. For example, water boils at 100 °C (212 °F) at sea level, but at 93.4 °C (200.1 °F) at 1,905 metres (6,250 ft) [3] altitude. For a given pressure, different liquids will boil at different temperatures.

The normal boiling point (also called the atmospheric boiling point or the atmospheric pressure boiling point) of a liquid is the special case in which the vapor pressure of the liquid equals the defined atmospheric pressure at sea level, 1 atmosphere.[4][5] At that temperature, the vapor pressure of the liquid becomes sufficient to overcome atmospheric pressure and allow bubbles of vapor to form inside the bulk of the liquid. The standard boiling point has been defined by IUPAC since 1982 as the temperature at which boiling occurs under a pressure of 1 bar.[6]

The heat of vaporization is the energy required to transform a given quantity (a mol, kg, pound, etc.) of a substance from a liquid into a gas at a given pressure (often atmospheric pressure).

Liquids may change to a vapor at temperatures below their boiling points through the process of evaporation. Evaporation is a surface phenomenon in which molecules located near the liquid's edge, not contained by enough liquid pressure on that side, escape into the surroundings as vapor. On the other hand, boiling is a process in which molecules anywhere in the liquid escape, resulting in the formation of vapor bubbles within the liquid.

Kochendes wasser02
Boiling water

Saturation temperature and pressure

Demonstration of the lower boiling point of water at lower pressure, achieved by using a vacuum pump.

A saturated liquid contains as much thermal energy as it can without boiling (or conversely a saturated vapor contains as little thermal energy as it can without 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 transition.

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 liquid equals the surrounding environmental pressure. Thus, the boiling point is dependent on the pressure. Boiling points may be published with respect to the NIST, USA standard pressure of 101.325 kPa (or 1 atm), or the IUPAC standard pressure of 100.000 kPa. At higher elevations, where the atmospheric pressure is much lower, the boiling point is also lower. The boiling point increases with increased 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 pressure until the triple point is reached. The boiling point cannot be reduced below the triple point.

If the heat of vaporization and the vapor pressure of a liquid at a certain temperature are known, the boiling point can be calculated by using the Clausius–Clapeyron equation, thus:

where:

is the boiling point at the pressure of interest,
is the ideal gas constant,
is the vapour pressure of the liquid at the pressure of interest,
is some pressure where the corresponding is known (usually data available at 1 atm or 100 kPa),
is the heat of vaporization of the liquid,
is the boiling temperature,
is the natural logarithm.

Saturation pressure 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.

There are two conventions regarding the standard boiling point of water: The normal boiling point is 99.97 °C (211.9 °F) at a pressure of 1 atm (i.e., 101.325 kPa). The IUPAC recommended standard boiling point of water at a standard pressure of 100 kPa (1 bar)[7] is 99.61 °C (211.3 °F).[6][8] For comparison, on top of Mount Everest, at 8,848 m (29,029 ft) elevation, the pressure is about 34 kPa (255 Torr)[9] and the boiling point of water is 71 °C (160 °F). The Celsius temperature scale was defined until 1954 by two points: 0 °C being defined by the water freezing point and 100 °C being defined by the water boiling point at standard atmospheric pressure.

Relation between the normal boiling point and the vapor pressure of liquids

Vapor pressure chart
A log-lin vapor pressure chart for various liquids

The higher the vapor pressure of a liquid at a given temperature, the lower the normal boiling point (i.e., the boiling point at atmospheric pressure) of the liquid.

The vapor pressure chart to the right has graphs of the vapor pressures versus temperatures for a variety of liquids.[10] As can be seen in the chart, the liquids with the highest vapor pressures have the lowest normal boiling points.

For example, at any given temperature, methyl chloride has the highest vapor pressure of any of the liquids in the chart. It also has the lowest normal boiling point (−24.2 °C), which is where the vapor pressure curve of methyl chloride (the blue line) intersects the horizontal pressure line of one atmosphere (atm) of absolute vapor pressure.

The critical point of a liquid is the highest temperature (and pressure) it will actually boil at.

See also Vapour pressure of water.

Properties of the elements

The element with the lowest boiling point is helium. Both the boiling points of rhenium and tungsten exceed 5000 K at standard pressure; because it is difficult to measure extreme temperatures precisely without bias, both have been cited in the literature as having the higher boiling point.[11]

Boiling point as a reference property of a pure compound

As can be seen from the above plot of the logarithm of the vapor pressure vs. the temperature for any given pure chemical compound, its normal boiling point can serve as an indication of that compound's overall volatility. A given pure compound has only one normal boiling point, if any, and a compound's normal boiling point and melting point can serve as characteristic physical properties for that compound, listed in reference books. The higher a compound's normal boiling point, the less volatile that compound is overall, and conversely, the lower a compound's normal boiling point, the more volatile that compound is overall. Some compounds decompose at higher temperatures before reaching their normal boiling point, or sometimes even their melting point. For a stable compound, the boiling point ranges from its triple point to its critical point, depending on the external pressure. Beyond its triple point, a compound's normal boiling point, if any, is higher than its melting point. Beyond the critical point, a compound's liquid and vapor phases merge into one phase, which may be called a superheated gas. At any given temperature, if a compound's normal boiling point is lower, then that compound will generally exist as a gas at atmospheric external pressure. If the compound's normal boiling point is higher, then that compound can exist as a liquid or solid at that given temperature at atmospheric external pressure, and will so exist in equilibrium with its vapor (if volatile) if its vapors are contained. If a compound's vapors are not contained, then some volatile compounds can eventually evaporate away in spite of their higher boiling points.

Boiling point vs molar mass graph
Boiling points of alkanes, alkenes, ethers, halogenoalkanes, aldehydes, ketones, alcohols and carboxylic acids as a function of molar mass

In general, compounds with ionic bonds have high normal boiling points, if they do not decompose before reaching such high temperatures. Many metals have high boiling points, but not all. Very generally—with other factors being equal—in compounds with covalently bonded molecules, as the size of the molecule (or molecular mass) increases, the normal boiling point increases. When the molecular size becomes that of a macromolecule, polymer, or otherwise very large, the compound often decomposes at high temperature before the boiling point is reached. Another factor that affects the normal boiling point of a compound is the polarity of its molecules. As the polarity of a compound's molecules increases, its normal boiling point increases, other factors being equal. Closely related is the ability of a molecule to form hydrogen bonds (in the liquid state), which makes it harder for molecules to leave the liquid state and thus increases the normal boiling point of the compound. Simple carboxylic acids dimerize by forming hydrogen bonds between molecules. A minor factor affecting boiling points is the shape of a molecule. Making the shape of a molecule more compact tends to lower the normal boiling point slightly compared to an equivalent molecule with more surface area.

Comparison of butane isomer boiling points
Common name n-butane isobutane
IUPAC name butane 2-methylpropane
Molecular
form
Butane-3D-balls Isobutane-3D-balls
Boiling
point (°C)
−0.5 −11.7
Comparison of pentane isomer boiling points
Common name n-pentane isopentane neopentane
IUPAC name pentane 2-methylbutane 2,2-dimethylpropane
Molecular
form
Pentane-3D-balls Isopentane-3D-balls Neopentane-3D-balls
Boiling
point (°C)
36.0 27.7 9.5
Binary Boiling Point Diagram new
Binary boiling point diagram of two hypothetical only weakly interacting components without an azeotrope

Most volatile compounds (anywhere near ambient temperatures) go through an intermediate liquid phase while warming up from a solid phase to eventually transform to a vapor phase. By comparison to boiling, a sublimation is a physical transformation in which a solid turns directly into vapor, which happens in a few select cases such as with carbon dioxide at atmospheric pressure. For such compounds, a sublimation point is a temperature at which a solid turning directly into vapor has a vapor pressure equal to the external pressure.

Impurities and mixtures

In the preceding section, boiling points of pure compounds were covered. Vapor pressures and boiling points of substances can be affected by the presence of dissolved impurities (solutes) or other miscible compounds, the degree of effect depending on the concentration of the impurities or other compounds. The presence of non-volatile impurities such as salts or compounds of a volatility far lower than the main component compound decreases its mole fraction and the solution's volatility, and thus raises the normal boiling point in proportion to the concentration of the solutes. This effect is called boiling point elevation. As a common example, salt water boils at a higher temperature than pure water.

In other mixtures of miscible compounds (components), there may be two or more components of varying volatility, each having its own pure component boiling point at any given pressure. The presence of other volatile components in a mixture affects the vapor pressures and thus boiling points and dew points of all the components in the mixture. The dew point is a temperature at which a vapor condenses into a liquid. Furthermore, at any given temperature, the composition of the vapor is different from the composition of the liquid in most such cases. In order to illustrate these effects between the volatile components in a mixture, a boiling point diagram is commonly used. Distillation is a process of boiling and [usually] condensation which takes advantage of these differences in composition between liquid and vapor phases.

See also

References

  1. ^ Goldberg, David E. (1988). 3,000 Solved Problems in Chemistry (1st ed.). McGraw-Hill. section 17.43, p. 321. ISBN 0-07-023684-4.
  2. ^ Theodore, Louis; Dupont, R. Ryan; Ganesan, Kumar, eds. (1999). Pollution Prevention: The Waste Management Approach to the 21st Century. CRC Press. section 27, p. 15. ISBN 1-56670-495-2.
  3. ^ "Boiling Point of Water and Altitude". www.engineeringtoolbox.com.
  4. ^ General Chemistry Glossary Purdue University website page
  5. ^ Reel, Kevin R.; Fikar, R. M.; Dumas, P. E.; Templin, Jay M. & Van Arnum, Patricia (2006). AP Chemistry (REA) – The Best Test Prep for the Advanced Placement Exam (9th ed.). Research & Education Association. section 71, p. 224. ISBN 0-7386-0221-3.
  6. ^ a b Cox, J. D. (1982). "Notation for states and processes, significance of the word standard in chemical thermodynamics, and remarks on commonly tabulated forms of thermodynamic functions". Pure and Applied Chemistry. 54 (6): 1239. doi:10.1351/pac198254061239.
  7. ^ Standard Pressure IUPAC defines the "standard pressure" as being 105 Pa (which amounts to 1 bar).
  8. ^ Appendix 1: Property Tables and Charts (SI Units), Scroll down to Table A-5 and read the temperature value of 99.61 °C at a pressure of 100 kPa (1 bar). Obtained from McGraw-Hill's Higher Education website.
  9. ^ West, J. B. (1999). "Barometric pressures on Mt. Everest: New data and physiological significance". Journal of Applied Physiology. 86 (3): 1062–6. doi:10.1152/jappl.1999.86.3.1062. PMID 10066724.
  10. ^ Perry, R.H.; Green, D.W., eds. (1997). Perry's Chemical Engineers' Handbook (7th ed.). McGraw-Hill. ISBN 0-07-049841-5.
  11. ^ DeVoe, Howard (2000). Thermodynamics and Chemistry (1st ed.). Prentice-Hall. ISBN 0-02-328741-1.

External links

Azeotrope

An azeotrope (UK /əˈziːəˌtrəʊp/, US /əˈziəˌtroʊp/) or a constant boiling point mixture is a mixture of two or more liquids whose proportions cannot be altered or changed by simple distillation. This happens because when an azeotrope is boiled, the vapour has the same proportions of constituents as the unboiled mixture. Because their composition is unchanged by distillation, azeotropes are also called (especially in older texts) constant boiling point mixtures.

Many azeotropic mixtures of pairs of compounds are known, and many azeotropes of three or more compounds are also known. In such a case it is not possible to separate the components by fractional distillation. There are two types of azeotropes: minimum boiling azeotrope and maximum boiling azeotrope. A solution that shows greater positive deviation from Raoult's law forms a minimum boiling azeotrope at a specific composition. For example, an ethanol-water mixture (obtained by fermentation of sugars) on fractional distillation yields a solution containing approximately 95% by volume of ethanol. Once this composition has been achieved, the liquid and vapour have the same composition, and no further separation occurs.

A solution that shows large negative deviation from Raoult's law forms a maximum boiling azeotrope at a specific composition. Nitric acid and water is an example of this class of azeotrope. This azeotrope has an approximate composition of 68% nitric acid and 32% water by mass, with a boiling point of 393.5 K (120.4 °C).

Boiling

Boiling is the rapid vaporization of a liquid, which occurs when a liquid is heated to its boiling point, the temperature at which the vapour pressure of the liquid is equal to the pressure exerted on the liquid by the surrounding atmosphere. There are two main types of boiling: nucleate boiling where small bubbles of vapour form at discrete points, and critical heat flux boiling where the boiling surface is heated above a certain critical temperature and a film of vapor forms on the surface. Transition boiling is an intermediate, unstable form of boiling with elements of both types. The boiling point of water is 100 °C or 212 °F but is lower with the decreased atmospheric pressure found at higher altitudes.

Boiling water is used as a method of making it potable by killing microbes that may be present. The sensitivity of different micro-organisms to heat varies, but if water is held at 70 °C (158 °F) for ten minutes, many organisms are killed, but some are more resistant to heat and require one minute at the boiling point of water.

Boiling is also used in cooking. Foods suitable for boiling include vegetables, starchy foods such as rice, noodles and potatoes, eggs, meats, sauces, stocks, and soups. As a cooking method, it is simple and suitable for large-scale cookery. Tough meats or poultry can be given a long, slow cooking and a nutritious stock is produced. Disadvantages include loss of water-soluble vitamins and minerals. Commercially prepared foodstuffs are sometimes packed in polythene sachets and sold as "boil-in-the-bag" products.

Boiling-point elevation

Boiling-point elevation describes the phenomenon that the boiling point of a liquid (a solvent) will be higher when another compound is added, meaning that a solution has a higher boiling point than a pure solvent. This happens whenever a non-volatile solute, such as a salt, is added to a pure solvent, such as water. The boiling point can be measured accurately using an ebullioscope.

Boiling Point, California

Boiling Point is an unincorporated community located in the Antelope Valley of the Mojave Desert, in northern Los Angeles County, California.The settlement is located along the Sierra Highway, 12 mi (19 km), west of Palmdale.

Ritter Ranch Park, a multi-purpose recreational area, is located north of the settlement along Boiling Point Road.

Boiling Point (2012)

Boiling Point (2012) was a professional wrestling Internet pay-per-view (iPPV) event produced by Ring of Honor (ROH) which took place on August 11, 2012 at the Rhode Island Convention Center in Providence, Rhode Island.

Boiling Points

Boiling Points is a prank reality television show, much like the format used on Candid Camera. It was broadcast on MTV in the United States from 2004 to 2005. In each half-hour episode, annoying situations were set up and deliberately inflicted on one or more young adults who were unaware that they were being tested. Examples included poor or incompetent customer service in a store or restaurant, being accosted by a date's ex-love-interest while out together, and unprovoked rudeness from a total stranger.

While being watched via hidden camera, the subject of the show must refrain from displaying a temper or storming off for a predetermined length of time. If he or she passes, a prize of $100 cash is awarded to them on the spot.

Typically, the Boiling Point time on each segment is anywhere from 3 to 20 minutes. Most often, people lose because they use obscenities in their confusion and/or anger.

The cast included improvisers Alison Becker, Jonathan Blitt, Colton Dunn, Giselle Forte, Rebekka Johnson, Billy Merritt, Missy O'Reilly, Eric Wippo and Sauce.

Celsius

The Celsius scale, also known as the centigrade scale, is a temperature scale used by the International System of Units (SI). As an SI derived unit, it is used by all countries except the United States, the Bahamas, Belize, the Cayman Islands and Liberia. It is named after the Swedish astronomer Anders Celsius (1701–1744), who developed a similar temperature scale. The degree Celsius (°C) can refer to a specific temperature on the Celsius scale or a unit to indicate a difference between two temperatures or an uncertainty. Before being renamed to honor Anders Celsius in 1948, the unit was called centigrade, from the Latin centum, which means 100, and gradus, which means steps.

From 1743, the Celsius scale is based on 0 °C for the freezing point of water and 100 °C for the boiling point of water at 1 atm pressure. Prior to 1743, the scale was also based on the boiling and melting points of water, but the values were reversed (i.e. the boiling point was at 0 degrees and the melting point was at 100 degrees). The 1743 scale reversal was proposed by Jean-Pierre Christin.

By international agreement, since 1954 the unit degree Celsius and the Celsius scale are defined by absolute zero and the triple point of Vienna Standard Mean Ocean Water (VSMOW), a specially purified water. This definition also precisely relates the Celsius scale to the Kelvin scale, which defines the SI base unit of thermodynamic temperature with symbol K. Absolute zero, the lowest temperature possible, is defined as being exactly 0 K and −273.15 °C. The temperature of the triple point of water is defined as exactly 273.16 K (0.01 °C). This means that a temperature difference of one degree Celsius and that of one kelvin are exactly the same.On May 20, 2019, the degree Kelvin, and along with it the degree Celsius, will again be re-defined so that its value will be determined by definition of the Boltzmann constant.

Colligative properties

In chemistry, colligative properties are properties of solutions that depend on the ratio of the number of solute particles to the number of solvent molecules in a solution, and not on the nature of the chemical species present. The number ratio can be related to the various units for concentration of solutions. The assumption that solution properties are independent of nature of solute particles is only exact for ideal solutions, and is approximate for dilute real solutions. In other words, colligative properties are a set of solution properties that can be reasonably approximated by assuming that the solution is ideal.

Here we consider only properties which result from the dissolution of nonvolatile solute in a volatile liquid solvent. They are essentially solvent properties which are changed by the presence of the solute. The solute particles displace some solvent molecules in the liquid phase and therefore reduce the concentration of solvent, so that the colligative properties are independent of the nature of the solute. The word colligative is derived from the Latin colligatus meaning bound together.Colligative properties include:

Relative lowering of vapor pressure

Elevation of boiling point

Depression of freezing point

Osmotic pressureFor a given solute-solvent mass ratio, all colligative properties are inversely proportional to solute molar mass.

Measurement of colligative properties for a dilute solution of a non-ionized solute such as urea or glucose in water or another solvent can lead to determinations of relative molar masses, both for small molecules and for polymers which cannot be studied by other means. Alternatively, measurements for ionized solutes can lead to an estimation of the percentage of dissociation taking place.

Colligative properties are mostly studied for dilute solutions, whose behavior may often be approximated as that of an ideal solution.

Cryogenics

In physics, cryogenics is the production and behaviour of materials at very low temperatures. A person who studies elements that have been subjected to extremely cold temperatures is called a cryogenicist.

It is not well-defined at what point on the temperature scale refrigeration ends and cryogenics begins, but scientists assume a gas to be cryogenic if it can be liquefied at or below −150 °C (123 K; −238 °F). The U.S. National Institute of Standards and Technology has chosen to consider the field of cryogenics as that involving temperatures below −180 °C (93 K; −292 °F). This is a logical dividing line, since the normal boiling points of the so-called permanent gases (such as helium, hydrogen, neon, nitrogen, oxygen, and normal air) lie below −180 °C while the Freon refrigerants, hydrocarbons, and other common refrigerants have boiling points above −180 °C.Discovery of superconducting materials with critical temperatures significantly above the boiling point of liquid nitrogen has provided new interest in reliable, low cost methods of producing high temperature cryogenic refrigeration. The term "high temperature cryogenic" describes temperatures ranging from above the boiling point of liquid nitrogen, −195.79 °C (77.36 K; −320.42 °F), up to −50 °C (223 K; −58 °F), the generally defined upper limit of study referred to as cryogenics.Cryogenicists use the Kelvin or Rankine temperature scale, both of which measure from absolute zero, rather than more usual scales such as Celsius or Fahrenheit, with their zeroes at arbitrary temperatures.

Distillation

Distillation is the process of separating the components or substances from a liquid mixture by using selective boiling and condensation. Distillation may result in essentially complete separation (nearly pure components), or it may be a partial separation that increases the concentration of selected components in the mixture. In either case, the process exploits differences in the volatility of the mixture's components. In industrial chemistry, distillation is a unit operation of practically universal importance, but it is a physical separation process, not a chemical reaction.

Distillation has many applications. For example:

Distillation of fermented products produces distilled beverages with a high alcohol content or separates out other fermentation products of commercial value.

Distillation is an effective and traditional method of desalination.

In the fossil fuel industry, oil stabilization is a form of partial distillation that reduces vapor pressure of crude oil, thereby making it safe for storage and transport as well as reducing the atmospheric emissions of volatile hydrocarbons. In midstream operations at oil refineries, distillation is a major class of operation for transforming crude oil into fuels and chemical feed stocks.

Cryogenic distillation leads to the separation of air into its components – notably oxygen, nitrogen, and argon – for industrial use.

In the field of industrial chemistry, large amounts of crude liquid products of chemical synthesis are distilled to separate them, either from other products, from impurities, or from unreacted starting materials.An installation used for distillation, especially of distilled beverages, is called a distillery. The distillation equipment at a distillery is a still.

Liquid helium

At standard pressure, the chemical element helium exists in a liquid form only at the extremely low temperature of −270 °C (about 4 K or −452.2 °F). Its boiling point and critical point depend on which isotope of helium is present: the common isotope helium-4 or the rare isotope helium-3. These are the only two stable isotopes of helium. See the table below for the values of these physical quantities. The density of liquid helium-4 at its boiling point and a pressure of one atmosphere (101.3 kilopascals) is about 0.125 grams per cm3, or about 1/8th the density of liquid water.

Liquid nitrogen

Liquid nitrogen is nitrogen in a liquid state at an extremely low temperature. It is a colorless liquid with a density of 0.807 g/ml at its boiling point (−195.79 °C (77 K; −320 °F)) and a dielectric constant of 1.43. Nitrogen was first liquefied at the Jagiellonian University on 15 April 1883 by Polish physicists, Zygmunt Wróblewski and Karol Olszewski. It is produced industrially by fractional distillation of liquid air. Liquid nitrogen is often referred to by the abbreviation, LN2 or "LIN" or "LN" and has the UN number 1977. Liquid nitrogen is a diatomic liquid, which means that the diatomic character of the covalent N bonding in N2 gas is retained after liquefaction.Liquid nitrogen is a cryogenic fluid that can cause rapid freezing on contact with living tissue. When appropriately insulated from ambient heat, liquid nitrogen can be stored and transported, for example in vacuum flasks. The temperature is held constant at 77 K by slow boiling of the liquid, resulting in the evolution of nitrogen gas. Depending on the size and design, the holding time of vacuum flasks ranges from a few hours to a few weeks. The development of pressurised super-insulated vacuum vessels has enabled liquefied nitrogen to be stored and transported over longer time periods with losses reduced to 2% per day or less.The temperature of liquid nitrogen can readily be reduced to its freezing point 63 K (−210 °C; −346 °F) by placing it in a vacuum chamber pumped by a vacuum pump. Liquid nitrogen's efficiency as a coolant is limited by the fact that it boils immediately on contact with a warmer object, enveloping the object in insulating nitrogen gas. This effect, known as the Leidenfrost effect, applies to any liquid in contact with an object significantly hotter than its boiling point. Faster cooling may be obtained by plunging an object into a slush of liquid and solid nitrogen rather than liquid nitrogen alone.

List of chemical elements

This is a list of the 118 chemical elements which have been identified as of 2019. A chemical element, often simply called an element, is a species of atoms which all have the same number of protons in their atomic nuclei (i.e., the same atomic number, or Z).Perhaps the most popular visualization of all 118 elements is the periodic table of the elements, a convenient tabular arrangement of the elements by their chemical properties that uses abbreviated chemical symbols in place of full element names, but the simpler list format presented here may also be useful. Like the periodic table, the list below organizes the elements by the number of protons in their atoms; it can also be organized by other properties, such as atomic weight, density, and electronegativity. For more detailed information about the origins of element names, see List of chemical element name etymologies.

Solvent

A solvent (from the Latin solvō, "loosen, untie, solve") is a substance that dissolves a solute (a chemically distinct liquid, solid or gas), resulting in a solution. A solvent is usually a liquid but can also be a solid, a gas, or a supercritical fluid. The quantity of solute that can dissolve in a specific volume of solvent varies with temperature. Common uses for organic solvents are in dry cleaning (e.g. tetrachloroethylene), as paint thinners (e.g. toluene, turpentine), as nail polish removers and glue solvents (acetone, methyl acetate, ethyl acetate), in spot removers (e.g. hexane, petrol ether), in detergents (citrus terpenes) and in perfumes (ethanol). Water is a solvent for polar molecules and the most common solvent used by living things; all the ions and proteins in a cell are dissolved in water within a cell. Solvents find various applications in chemical, pharmaceutical, oil, and gas industries, including in chemical syntheses and purification processes.

UFC 54

UFC 54: Boiling Point was a mixed martial arts event held by the Ultimate Fighting Championship on August 20, 2005, at the MGM Grand Arena in Las Vegas, Nevada. The event was broadcast live on pay-per-view in the United States, and later released on DVD.

Vapor pressure

Vapor pressure (or vapour pressure in British spelling) or equilibrium vapor pressure is defined as the pressure exerted by a vapor in thermodynamic equilibrium with its condensed phases (solid or liquid) at a given temperature in a closed system. The equilibrium vapor pressure is an indication of a liquid's evaporation rate. It relates to the tendency of particles to escape from the liquid (or a solid). A substance with a high vapor pressure at normal temperatures is often referred to as volatile. The pressure exhibited by vapor present above a liquid surface is known as vapor pressure. As the temperature of a liquid increases, the kinetic energy of its molecules also increases. As the kinetic energy of the molecules increases, the number of molecules transitioning into a vapor also increases, thereby increasing the vapor pressure.

The vapor pressure of any substance increases non-linearly with temperature according to the Clausius–Clapeyron relation. The atmospheric pressure boiling point of a liquid (also known as the normal boiling point) is the temperature at which the vapor pressure equals the ambient atmospheric pressure. With any incremental increase in that temperature, the vapor pressure becomes sufficient to overcome atmospheric pressure and lift the liquid to form vapor bubbles inside the bulk of the substance. Bubble formation deeper in the liquid requires a higher temperature due to the higher fluid pressure, because fluid pressure increases above the atmospheric pressure as the depth increases. More important at shallow depths is the higher temperature required to start bubble formation. The surface tension of the bubble wall leads to an overpressure in the very small, initial bubbles. Thus, thermometer calibration should not rely on the temperature in boiling water.

The vapor pressure that a single component in a mixture contributes to the total pressure in the system is called partial pressure. For example, air at sea level, and saturated with water vapor at 20 °C, has partial pressures of about 2.3 kPa of water, 78 kPa of nitrogen, 21 kPa of oxygen and 0.9 kPa of argon, totaling 102.2 kPa, making the basis for standard atmospheric pressure.

Vapor–liquid equilibrium

In thermodynamics and chemical engineering, the vapor–liquid equilibrium (VLE) describes the distribution of a chemical species between the vapor phase and a liquid phase.

The concentration of a vapor in contact with its liquid, especially at equilibrium, is often expressed in terms of vapor pressure, which will be a partial pressure (a part of the total gas pressure) if any other gas(es) are present with the vapor. The equilibrium vapor pressure of a liquid is in general strongly dependent on temperature. At vapor–liquid equilibrium, a liquid with individual components in certain concentrations will have an equilibrium vapor in which the concentrations or partial pressures of the vapor components have certain values depending on all of the liquid component concentrations and the temperature. The converse is also true: if a vapor with components at certain concentrations or partial pressures is in vapor–liquid equilibrium with its liquid, then the component concentrations in the liquid will be determined dependent on the vapor concentrations and on the temperature. The equilibrium concentration of each component in the liquid phase is often different from its concentration (or vapor pressure) in the vapor phase, but there is a relationship. The VLE concentration data can be determined experimentally, or computed or approximated with the help of theories such as Raoult's law, Dalton's law, and Henry's law.

Such vapor–liquid equilibrium information is useful in designing columns for distillation, especially fractional distillation, which is a particular specialty of chemical engineers. Distillation is a process used to separate or partially separate components in a mixture by boiling (vaporization) followed by condensation. Distillation takes advantage of differences in concentrations of components in the liquid and vapor phases.

In mixtures containing two or more components, the concentrations of each component are often expressed as mole fractions. The mole fraction of a given component of a mixture in a particular phase (either the vapor or the liquid phase) is the number of moles of that component in that phase divided by the total number of moles of all components in that phase.

Binary mixtures are those having two components. Three-component mixtures are called ternary mixtures. There can be VLE data for mixtures with even more components, but such data is often hard to show graphically. VLE data is a function of the total pressure, such as 1 atm or whatever pressure the process is conducted at.

When a temperature is reached such that the sum of the equilibrium vapor pressures of the liquid components becomes equal to the total pressure of the system (it is otherwise smaller), then vapor bubbles generated from the liquid begin to displace the gas that was maintaining the overall pressure, and the mixture is said to boil. This temperature is called the boiling point of the liquid mixture at the given pressure. (It is assumed that the total pressure is held steady by adjusting the total volume of the system to accommodate the specific volume changes that accompany boiling.) The boiling point at an overall pressure of 1 atm is called the normal boiling point of the liquid mixture.

Volatility (chemistry)

In chemistry and physics, volatility is quantified by the tendency of a substance to vaporize. Volatility is directly related to a substance's vapor pressure. At a given temperature, a substance with higher vapor pressure vaporizes more readily than a substance with a lower vapor pressure.The term is primarily written to be applied to liquids; however, it may be used to describe the process of sublimation which is associated with solid substances, such as dry ice (solid carbon dioxide) and osmium tetroxide (OsO4), which can change directly from the solid state to a vapor, without becoming liquid.

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