A copolymer is a polymer derived from more than one species of monomer. The polymerization of monomers into copolymers is called copolymerization. Copolymers obtained by copolymerization of two monomer species are sometimes called bipolymers; those obtained from three and four monomers are called quaterpolymers, respectively. There are many commercially relevant copolymers; some examples include acrylonitrile butadiene styrene, styrene/butadiene co-polymer, nitrile rubber, styrene-acrylonitrile, styrene-isoprene-styrene and ethylene-vinyl acetate, all formed by chain-growth polymerization. Another production mechanism is step-growth polymerization, used to produce the nylon-12/6/66 copolymer of nylon 12, nylon 6 and nylon 66, as well as the copolyester family. Since a copolymer consists of at least two types of constituent units, copolymers can be classified based on how these units are arranged along the chain. Linear copolymers consist of a single main chain, include alternating copolymers, statistical copolymers and block copolymers.
Branched copolymers consist of a single main chain with one or more polymeric side chains, can be grafted, star shaped or have other architectures. The reactivity ratio of a growing copolymer chain terminating in a given monomer is the ratio of the reaction rate constant for addition of the same monomer and the rate constant for addition of the other monomer; that is, r 1 = k 11 k 12 and r 2 = k 22 k 21, where for example k 12 is the rate constant for propagation of a polymer chain ending in monomer 1 by addition of monomer 2. The composition and structural type of the copolymer depend on these reactivity ratios r1 and r2 according to the Mayo–Lewis equation called the copolymerization equation or copolymer equation, for the relative instantaneous rates of incorporation of the two monomers. D d = Block copolymers comprise two or more homopolymer subunits linked by covalent bonds; the union of the homopolymer subunits may require an intermediate non-repeating subunit, known as a junction block.
Block copolymers with two or three distinct blocks are called diblock copolymers and triblock copolymers, respectively. Technically, a block is a portion of a macromolecule, comprising many constitutional units, that has at least one feature, not present in the adjacent portions. A possible sequence of repeat units A and B in a triblock copolymer might be ~A-A-A-A-A-A-A-B-B-B-B-B-B-B-A-A-A-A-A~. Block copolymers are made up of blocks of different polymerized monomers. For example, polystyrene-b-poly or PS-b-PMMA is made by first polymerizing styrene, subsequently polymerizing methyl methacrylate from the reactive end of the polystyrene chains; this polymer is a "diblock copolymer". Triblocks, multiblocks, etc. can be made. Diblock copolymers are made using living polymerization techniques, such as atom transfer free radical polymerization, reversible addition fragmentation chain transfer, ring-opening metathesis polymerization, living cationic or living anionic polymerizations. An emerging technique is chain shuttling polymerization.
The synthesis of block copolymers requires that both reactivity ratios are much larger than unity under the reaction conditions, so that the terminal monomer unit of a growing chain tends to add a similar unit most of the time. The "blockiness" of a copolymer is a measure of the adjacency of comonomers vs their statistical distribution. Many or most synthetic polymers are in fact copolymers, containing about 1-20% of a minority monomer. In such cases, blockiness is undesirable. A block index has been proposed as a quantitative measure of blockiness or deviation from random monomer composition. An alternating copolymer has regular alternating A and B units, is described by the formula: -A-B-A-B-A-B-A-B-A-B-, or -n-; the molar ratio of each monomer in the polymer is close to one, which happens when the reactivity ratios r1 and r2 are close to zero, as can be seen from the Mayo–Lewis equation. For example, in the free-radical copolymerization of
Simplified molecular-input line-entry system
The simplified molecular-input line-entry system is a specification in the form of a line notation for describing the structure of chemical species using short ASCII strings. SMILES strings can be imported by most molecule editors for conversion back into two-dimensional drawings or three-dimensional models of the molecules; the original SMILES specification was initiated in the 1980s. It has since been extended. In 2007, an open standard called. Other linear notations include the Wiswesser line notation, ROSDAL, SYBYL Line Notation; the original SMILES specification was initiated by David Weininger at the USEPA Mid-Continent Ecology Division Laboratory in Duluth in the 1980s. Acknowledged for their parts in the early development were "Gilman Veith and Rose Russo and Albert Leo and Corwin Hansch for supporting the work, Arthur Weininger and Jeremy Scofield for assistance in programming the system." The Environmental Protection Agency funded the initial project to develop SMILES. It has since been modified and extended by others, most notably by Daylight Chemical Information Systems.
In 2007, an open standard called "OpenSMILES" was developed by the Blue Obelisk open-source chemistry community. Other'linear' notations include the Wiswesser Line Notation, ROSDAL and SLN. In July 2006, the IUPAC introduced the InChI as a standard for formula representation. SMILES is considered to have the advantage of being more human-readable than InChI; the term SMILES refers to a line notation for encoding molecular structures and specific instances should be called SMILES strings. However, the term SMILES is commonly used to refer to both a single SMILES string and a number of SMILES strings; the terms "canonical" and "isomeric" can lead to some confusion when applied to SMILES. The terms are not mutually exclusive. A number of valid SMILES strings can be written for a molecule. For example, CCO, OCC and CC all specify the structure of ethanol. Algorithms have been developed to generate the same SMILES string for a given molecule; this SMILES is unique for each structure, although dependent on the canonicalization algorithm used to generate it, is termed the canonical SMILES.
These algorithms first convert the SMILES to an internal representation of the molecular structure. Various algorithms for generating canonical SMILES have been developed and include those by Daylight Chemical Information Systems, OpenEye Scientific Software, MEDIT, Chemical Computing Group, MolSoft LLC, the Chemistry Development Kit. A common application of canonical SMILES is indexing and ensuring uniqueness of molecules in a database; the original paper that described the CANGEN algorithm claimed to generate unique SMILES strings for graphs representing molecules, but the algorithm fails for a number of simple cases and cannot be considered a correct method for representing a graph canonically. There is no systematic comparison across commercial software to test if such flaws exist in those packages. SMILES notation allows the specification of configuration at tetrahedral centers, double bond geometry; these are structural features that cannot be specified by connectivity alone and SMILES which encode this information are termed isomeric SMILES.
A notable feature of these rules is. The term isomeric SMILES is applied to SMILES in which isotopes are specified. In terms of a graph-based computational procedure, SMILES is a string obtained by printing the symbol nodes encountered in a depth-first tree traversal of a chemical graph; the chemical graph is first trimmed to remove hydrogen atoms and cycles are broken to turn it into a spanning tree. Where cycles have been broken, numeric suffix labels are included to indicate the connected nodes. Parentheses are used to indicate points of branching on the tree; the resultant SMILES form depends on the choices: of the bonds chosen to break cycles, of the starting atom used for the depth-first traversal, of the order in which branches are listed when encountered. Atoms are represented by the standard abbreviation of the chemical elements, in square brackets, such as for gold. Brackets may be omitted in the common case of atoms which: are in the "organic subset" of B, C, N, O, P, S, F, Cl, Br, or I, have no formal charge, have the number of hydrogens attached implied by the SMILES valence model, are the normal isotopes, are not chiral centers.
All other elements must be enclosed in brackets, have charges and hydrogens shown explicitly. For instance, the SMILES for water may be written as either O or. Hydrogen may be written as a separate atom; when brackets are used, the symbol H is added if the atom in brackets is bonded to one or more hydrogen, followed by the number of hydrogen atoms if greater than 1 by the sign + for a positive charge or by - for a negative charge. For example, for ammonium. If there is more than one charge, it is written as digit.
The melting point of a substance is the temperature at which it changes state from solid to liquid. At the melting point the solid and liquid phase exist in equilibrium; the melting point of a substance depends on pressure and is specified at a standard pressure such as 1 atmosphere or 100 kPa. When considered as the temperature of the reverse change from liquid to solid, it is referred to as the freezing point or crystallization point; because of the ability of some substances to supercool, the freezing point is not considered as a characteristic property of a substance. When the "characteristic freezing point" of a substance is determined, in fact the actual methodology is always "the principle of observing the disappearance rather than the formation of ice", that is, the melting point. For most substances and freezing points are equal. For example, the melting point and freezing point of mercury is 234.32 kelvins. However, certain substances possess differing solid-liquid transition temperatures.
For example, agar melts at 85 °C and solidifies from 31 °C. The melting point of ice at 1 atmosphere of pressure is close to 0 °C. In the presence of nucleating substances, the freezing point of water is not always the same as the melting point. In the absence of nucleators water can exist as a supercooled liquid down to −48.3 °C before freezing. The chemical element with the highest melting point is tungsten, at 3,414 °C; the often-cited carbon does not melt at ambient pressure but sublimes at about 3,726.85 °C. Tantalum hafnium carbide is a refractory compound with a high melting point of 4215 K. At the other end of the scale, helium does not freeze at all at normal pressure at temperatures arbitrarily close to absolute zero. Many laboratory techniques exist for the determination of melting points. A Kofler bench is a metal strip with a temperature gradient. Any substance can be placed on a section of the strip, revealing its thermal behaviour at the temperature at that point. Differential scanning calorimetry gives information on melting point together with its enthalpy of fusion.
A basic melting point apparatus for the analysis of crystalline solids consists of an oil bath with a transparent window and a simple magnifier. The several grains of a solid are placed in a thin glass tube and immersed in the oil bath; the oil bath is heated and with the aid of the magnifier melting of the individual crystals at a certain temperature can be observed. In large/small devices, the sample is placed in a heating block, optical detection is automated; the measurement can be made continuously with an operating process. For instance, oil refineries measure the freeze point of diesel fuel online, meaning that the sample is taken from the process and measured automatically; this allows for more frequent measurements as the sample does not have to be manually collected and taken to a remote laboratory. For refractory materials the high melting point may be determined by heating the material in a black body furnace and measuring the black-body temperature with an optical pyrometer. For the highest melting materials, this may require extrapolation by several hundred degrees.
The spectral radiance from an incandescent body is known to be a function of its temperature. An optical pyrometer matches the radiance of a body under study to the radiance of a source, calibrated as a function of temperature. In this way, the measurement of the absolute magnitude of the intensity of radiation is unnecessary. However, known temperatures must be used to determine the calibration of the pyrometer. For temperatures above the calibration range of the source, an extrapolation technique must be employed; this extrapolation is accomplished by using Planck's law of radiation. The constants in this equation are not known with sufficient accuracy, causing errors in the extrapolation to become larger at higher temperatures. However, standard techniques have been developed to perform this extrapolation. Consider the case of using gold as the source. In this technique, the current through the filament of the pyrometer is adjusted until the light intensity of the filament matches that of a black-body at the melting point of gold.
This establishes the primary calibration temperature and can be expressed in terms of current through the pyrometer lamp. With the same current setting, the pyrometer is sighted on another black-body at a higher temperature. An absorbing medium of known transmission is inserted between this black-body; the temperature of the black-body is adjusted until a match exists between its intensity and that of the pyrometer filament. The true higher temperature of the black-body is determined from Planck's Law; the absorbing medium is removed and the current through the filament is adjusted to match the filament intensity to that of the black-body. This establishes a second calibration point for the pyrometer; this step is repeated to carry the calibration to hi
The flash point of a volatile material is the lowest temperature at which vapours of the material will ignite, when given an ignition source. The flash point is sometimes confused with the autoignition temperature, the temperature that results in spontaneous autoignition; the fire point is the lowest temperature at which vapors of the material will keep burning after the ignition source is removed. The fire point is higher than the flash point, because at the flash point more vapor may not be produced enough to sustain combustion. Neither flash point nor fire point depends directly on the ignition source temperature, but ignition source temperature is far higher than either the flash or fire point; the flash point is a descriptive characteristic, used to distinguish between flammable fuels, such as petrol, combustible fuels, such as diesel. It is used to characterize the fire hazards of fuels. Fuels which have a flash point less than 37.8 °C are called flammable, whereas fuels having a flash point above that temperature are called combustible.
All liquids have a specific vapor pressure, a function of that liquid's temperature and is subject to Boyle's Law. As temperature increases, vapor pressure increases; as vapor pressure increases, the concentration of vapor of a flammable or combustible liquid in the air increases. Hence, temperature determines the concentration of vapor of the flammable liquid in the air. A certain concentration of a flammable or combustible vapor is necessary to sustain combustion in air, the lower flammable limit, that concentration is different and is specific to each flammable or combustible liquid; the flash point is the lowest temperature at which there will be enough flammable vapor to induce ignition when an ignition source is applied There are two basic types of flash point measurement: open cup and closed cup. In open cup devices, the sample is contained in an open cup, heated and, at intervals, a flame brought over the surface; the measured flash point will vary with the height of the flame above the liquid surface and, at sufficient height, the measured flash point temperature will coincide with the fire point.
The best-known example is the Cleveland open cup. There are two types of closed cup testers: non-equilibrial, such as Pensky-Martens, where the vapours above the liquid are not in temperature equilibrium with the liquid, equilibrial, such as Small Scale, where the vapours are deemed to be in temperature equilibrium with the liquid. In both these types, the cups are sealed with a lid through which the ignition source can be introduced. Closed cup testers give lower values for the flash point than open cup and are a better approximation to the temperature at which the vapour pressure reaches the lower flammable limit; the flash point is an empirical measurement rather than a fundamental physical parameter. The measured value will vary with equipment and test protocol variations, including temperature ramp rate, time allowed for the sample to equilibrate, sample volume and whether the sample is stirred. Methods for determining the flash point of a liquid are specified in many standards. For example, testing by the Pensky-Martens closed cup method is detailed in ASTM D93, IP34, ISO 2719, DIN 51758, JIS K2265 and AFNOR M07-019.
Determination of flash point by the Small Scale closed cup method is detailed in ASTM D3828 and D3278, EN ISO 3679 and 3680, IP 523 and 524. CEN/TR 15138 Guide to Flash Point Testing and ISO TR 29662 Guidance for Flash Point Testing cover the key aspects of flash point testing. Gasoline is a fuel used in a spark-ignition engine; the fuel is mixed with air within its flammable limits and heated by compression and subject to Boyle's Law above its flash point ignited by the spark plug. To ignite, the fuel must have a low flash point, but in order to avoid preignition caused by residual heat in a hot combustion chamber, the fuel must have a high autoignition temperature. Diesel fuel flash points vary between 52 and 96 °C. Diesel is suitable for use in a compression-ignition engine. Air is compressed until it has been heated above the autoignition temperature of the fuel, injected as a high-pressure spray, keeping the fuel–air mix within flammable limits. In a diesel-fueled engine, there is no ignition source.
Diesel fuel must have a high flash point and a low autoignition temperature. Jet fuel flash points vary with the composition of the fuel. Both Jet A and Jet A-1 have flash points between 38 and 66 °C, close to that of off-the-shelf kerosene, yet both Jet B and JP-4 have flash points between −23 and −1 °C. Flash points of substances are measured according to standard test methods described and defined in a 1938 publication by T. L. Ainsley of South Shields entitled "Sea Transport of Petroleum"; the test methodology defines the apparatus required to carry out the measurement, key test parameters, the procedure for the operator or automated apparatus to follow, the precision of the test method. Standard test methods are written and controlled by a number of national and international committees and organizations; the three main bodies are the CEN / ISO Joint Working Group on Flash Point, ASTM D02.8B Flammability Section and the Energy Institute's TMS SC-B-4 Flammability Panel. Autoignition temperature Fire point Safety data sheet
In chemistry, a radical is an atom, molecule, or ion that has an unpaired valence electron. With some exceptions, these unpaired electrons make radicals chemically reactive. Many radicals spontaneously dimerize. Most organic radicals have short lifetimes. A notable example of a radical is the hydroxyl radical, a molecule that has one unpaired electron on the oxygen atom. Two other examples are triplet triplet carbene which have two unpaired electrons. Radicals may be generated in a number of ways. Ionizing radiation, electrical discharges, electrolysis are known to produce radicals. Radicals are intermediates in many chemical reactions, more so than is apparent from the balanced equations. Radicals are important in combustion, atmospheric chemistry, plasma chemistry and many other chemical processes. A large fraction of natural products is generated by radical-generating enzymes. In living organisms, the radicals superoxide and nitric oxide and their reaction products regulate many processes, such as control of vascular tone and thus blood pressure.
They play a key role in the intermediary metabolism of various biological compounds. Such radicals can be messengers in a process dubbed redox signaling. A radical may be otherwise bound. In chemical equations, radicals are denoted by a dot placed to the right of the atomic symbol or molecular formula as follows: C l 2 → U V 2 C l ⋅ Radical reaction mechanisms use single-headed arrows to depict the movement of single electrons: The homolytic cleavage of the breaking bond is drawn with a'fish-hook' arrow to distinguish from the usual movement of two electrons depicted by a standard curly arrow; the second electron of the breaking bond moves to pair up with the attacking radical electron. Radicals take part in radical addition and radical substitution as reactive intermediates. Chain reactions involving radicals can be divided into three distinct processes; these are initiation and termination. Initiation reactions are those, they may involve the formation of radicals from stable species as in Reaction 1 above or they may involve reactions of radicals with stable species to form more radicals.
Propagation reactions are those reactions involving radicals in which the total number of radicals remains the same. Termination reactions are those reactions resulting in a net decrease in the number of radicals. Two radicals combine to form a more stable species, for example: 2Cl·→ Cl2 Radicals can form by breaking of covalent bonds by homolysis; the homolytic bond dissociation energies abbreviated as "ΔH °" are a measure of bond strength. Splitting H2 into 2H•, for example, requires a ΔH ° of +435 kJ·mol-1, while splitting Cl2 into 2Cl• requires a ΔH ° of +243 kJ·mol-1. For weak bonds, homolysis can be induced thermally. Strong bonds require high energy photons or flames to induce homolysis. Radicals or charged species add to non-radicals to give new radicals; this process is the basis of the radical chain reaction. Being prevalent and a diradical, O2 reacts with many organic compounds to generate radicals together with the hydroperoxide radical; this process is related to rancidification of unsaturated fats.
Radicals may be formed by single-electron oxidation or reduction of an atom or molecule. These redox reactions occur in electrochemical cells and in ionization chambers of mass spectrometers. Although radicals are short-lived due to their reactivity, there are long-lived radicals; these are categorized as follows: The prime example of a stable radical is molecular dioxygen. Another common example is nitric oxide. Organic radicals can be long lived if they occur in a conjugated π system, such as the radical derived from α-tocopherol. There are hundreds of examples of thiazyl radicals, which show low reactivity and remarkable thermodynamic stability with only a limited extent of π resonance stabilization. Persistent radical compounds are those whose longevity is due to steric crowding around the radical center, which makes it physically difficult for the radical to react with another molecule. Examples of these include Gomberg's triphenylmethyl radical, Fremy's salt, such as TEMPO, TEMPOL, nitronyl nitroxides, azephenylenyls and radicals derived from PTM and TTM.
Persistent radicals are generated in great quantity during combustion, "may be responsible for the oxidative stress resulting in cardiopulmonary disease and cancer, attributed to exposure to airborne fine particles". Gomberg's free radical can be generated by following reaction in lab - 3C-Cl + Ag === 3C• + AgCl The reason for persistivity of free radicals is either the delocalisation of unpaired electron or the unavailability of unpaired electron to other species due to the screening of neighbouring atoms/groups. Diradicals are molecules containing two radical centers. Multiple radical centers can exist in a molecule. Atmospheric oxygen exists as a diradical in its ground state as triplet oxygen; the low reactivity of atmospheric oxygen is due to its diradical state. Non-radical states of dioxygen are less stable tha
In organic chemistry, an alkene is an unsaturated hydrocarbon that contains at least one carbon–carbon double bond. The words alkene and olefin are used interchangeably. Acyclic alkenes, with only one double bond and no other functional groups, known as mono-enes, form a homologous series of hydrocarbons with the general formula CnH2n. Alkenes have two hydrogen atoms fewer than the corresponding alkane; the simplest alkene, with the International Union of Pure and Applied Chemistry name ethene, is the organic compound produced on the largest scale industrially. Aromatic compounds are drawn as cyclic alkenes, but their structure and properties are different and they are not considered to be alkenes. Like a single covalent bond, double bonds can be described in terms of overlapping atomic orbitals, except that, unlike a single bond, a carbon–carbon double bond consists of one sigma bond and one pi bond; this double bond is stronger than a single covalent bond and shorter, with an average bond length of 1.33 ångströms.
Each carbon of the double bond uses its three sp2 hybrid orbitals to form sigma bonds to three atoms. The unhybridized 2p atomic orbitals, which lie perpendicular to the plane created by the axes of the three sp² hybrid orbitals, combine to form the pi bond; this bond lies outside the main C–C axis, with half of the bond on one side of the molecule and half on the other. With a strength of 65 kcal/mol, the pi bond is weaker than the sigma bond. Rotation about the carbon–carbon double bond is restricted because it incurs an energetic cost to break the alignment of the p orbitals on the two carbon atoms; as a consequence, substituted alkenes may exist as one of called cis or trans isomers. More complex alkenes may be named with the E–Z notation for molecules with three or four different substituents. For example, of the isomers of butene, the two methyl groups of -but-2-ene appear on the same side of the double bond, in -but-2-ene the methyl groups appear on opposite sides; these two isomers of butene are different in their chemical and physical properties.
Twisting to a 90° dihedral angle between two of the groups on the carbons requires less energy than the strength of a pi bond, the bond still holds. The carbons of the double bond become pyramidal, which allows preserving some p orbital alignment—and hence pi bonding; the other two attached. This contradicts a common textbook assertion that the two carbons retain their planar nature when twisting, in which case the p orbitals would rotate enough away from each other to be unable to sustain a pi bond. In a 90°-twisted alkene, the p orbitals are only misaligned by 42° and the strain energy is only around 40 kcal/mol. In contrast, a broken pi bond has an energetic cost of around 65 kcal/mol; some pyramidal alkenes are stable. For example, trans-cyclooctene is a stable strained alkene and the orbital misalignment is only 19°, despite having a significant dihedral angle of 137° and a degree of pyramidalization of 18°. Trans-cycloheptene is stable at low temperatures; as predicted by the VSEPR model of electron pair repulsion, the molecular geometry of alkenes includes bond angles about each carbon in a double bond of about 120°.
The angle may vary because of steric strain introduced by nonbonded interactions between functional groups attached to the carbons of the double bond. For example, the C–C–C bond angle in propylene is 123.9°. For bridged alkenes, Bredt's rule states that a double bond cannot occur at the bridgehead of a bridged ring system unless the rings are large enough. Following Fawcett and defining S as the total number of non-bridgehead atoms in the rings, bicyclic systems require S ≥ 7 for stability and tricyclic systems require S ≥ 11. Many of the physical properties of alkenes and alkanes are similar: they are colourless and combustable; the physical state depends on molecular mass: like the corresponding saturated hydrocarbons, the simplest alkenes, ethene and butene are gases at room temperature. Linear alkenes of five to sixteen carbons are liquids, higher alkenes are waxy solids; the melting point of the solids increases with increase in molecular mass. Alkenes have stronger smells than the corresponding alkane.
Ethylene is described to have a "sweet" odor, for example. The binding of cupric ion to the olefin in the mammalian olfactory receptor MOR244-3 is implicated in the smell of alkenes. Strained alkenes, in particular, like norbornene and trans-cyclooctene are known to have strong, unpleasant odors, a fact consistent with the stronger π complexes they form with metal ions including copper. Alkenes are stable compounds, but are more reactive than alkanes, either because of the reactivity of the carbon–carbon pi-bond or the presence of allylic CH centers. Most reactions of alkenes involve additions to this pi bond. Alkenes serve as a feedstock for the petrochemical industry because they can participate in a wide variety of reactions, prominently polymerization and alkylation. Alkenes react in many addition reactions. Most of these addition reactions follow the mechanism of electrophilic addition. Examples are hydrohalogenation, halohydrin formation, hydroboration, dichlorocarbene addition, Simmons–Smith reaction, catalytic hydrogenation, epox
The boiling point of a substance is the temperature at which the vapor pressure of a liquid equals the pressure surrounding the liquid 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 at 93.4 °C at 1,905 metres altitude. For a given pressure, different liquids will boil at different temperatures; the normal 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. 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.
The heat of vaporization is the energy required to transform a given quantity of a substance from a liquid into a gas at a given 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. A saturated liquid contains as much thermal energy. 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, a vapor at saturation temperature will begin to condense into its liquid phase as thermal energy is removed.
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 the IUPAC standard pressure of 100.000 kPa. At higher elevations, where the atmospheric pressure is much lower, the boiling point is 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; 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: T B = − 1, where: T B is the boiling point at the pressure of interest, R is the ideal gas constant, P is the vapour pressure of the liquid at the pressure of interest, P 0 is some pressure where the corresponding T 0 is known, Δ H vap is the heat of vaporization of the liquid, T 0 is the boiling temperature, ln 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, vapor at saturation pressure and temperature will begin to condense into its liquid phase as the system pressure is increased. 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 at a pressure of 1 atm. The IUPAC recommended standard boiling point of water at a standard pressure of 100 kPa is 99.61 °C. For comparison, on top of Mount Everest, at 8,848 m elevation, the pressure is about 34 kPa and the boiling point of water is 71 °C; the Celsius temperature scale was defined until 1954 by two points: 0 °C being defined by the wate