Vapor pressure or equilibrium vapor pressure is defined as the pressure exerted by a vapor in thermodynamic equilibrium with its condensed phases 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. A substance with a high vapor pressure at normal temperatures is 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 increases; as the kinetic energy of the molecules increases, the number of molecules transitioning into a vapor 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 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 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, 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 pressure is measured in the standard units of pressure. The International System of Units recognizes pressure as a derived unit with the dimension of force per area and designates the pascal as its standard unit. One pascal is one newton per square meter. Experimental measurement of vapor pressure is a simple procedure for common pressures between 1 and 200 kPa. Most accurate results are obtained near the boiling point of substances and large errors result for measurements smaller than 1kPa. Procedures consist of purifying the test substance, isolating it in a container, evacuating any foreign gas measuring the equilibrium pressure of the gaseous phase of the substance in the container at different temperatures. Better accuracy is achieved when care is taken to ensure that the entire substance and its vapor are at the prescribed temperature; this is done, as with the use of an isoteniscope, by submerging the containment area in a liquid bath. Low vapor pressures of solids can be measured using the Knudsen effusion cell method.
In a medical context, vapor pressure is sometimes expressed in other units millimeters of mercury. This is important for volatile anesthetics, most of which are liquids at body temperature, but with a high vapor pressure. Anesthetics with a higher vapor pressure at body temperature will be excreted more as they are exhaled from the lungs; the Antoine equation is a mathematical expression of the relation between the vapor pressure and the temperature of pure liquid or solid substances. The basic form of the equation is: log P = A − B C + T and it can be transformed into this temperature-explicit form: T = B A − log P − C where: P is the absolute vapor pressure of a substance T is the temperature of the substance A, B and C are substance-specific coefficients log is either log 10 or log e A simpler form of the equation with only two coefficients is sometimes used: log P = A − B T which can be transformed to: T = B A − log P Sublimations and vaporizations of the same substance have separate sets of Antoine coefficients, as do components in mixtures.
Each parameter set for a specific compound is only applicable over a specified temperature range. Temperature ranges are chosen to maintain the equation's accuracy of a few up to 8–10 percent. For many volatile substances, several different sets of parameters are available and used for different temperature ranges; the Antoine equation has poor accuracy with any single parameter set when used from a compound's melting point to its critical temperature. Accuracy is usually poor when vapor pressure is under 10 Torr because of the limitations of the apparatus used to establish the Antoine parameter values; the Wagner equation gives "o
In the physical sciences, a partition coefficient or distribution coefficient is the ratio of concentrations of a compound in a mixture of two immiscible phases at equilibrium. This ratio is therefore a measure of the difference in solubility of the compound in these two phases; the partition coefficient refers to the concentration ratio of un-ionized species of compound, whereas the distribution coefficient refers to the concentration ratio of all species of the compound. In the chemical and pharmaceutical sciences, both phases are solvents. Most one of the solvents is water, while the second is hydrophobic, such as 1-octanol. Hence the partition coefficient measures how hydrophobic a chemical substance is. Partition coefficients are useful in estimating the distribution of drugs within the body. Hydrophobic drugs with high octanol/water partition coefficients are distributed to hydrophobic areas such as lipid bilayers of cells. Conversely, hydrophilic drugs are found in aqueous regions such as blood serum.
If one of the solvents is a gas and the other a liquid, a gas/liquid partition coefficient can be determined. For example, the blood/gas partition coefficient of a general anesthetic measures how the anesthetic passes from gas to blood. Partition coefficients can be defined when one of the phases is solid, for instance, when one phase is a molten metal and the second is a solid metal, or when both phases are solids; the partitioning of a substance into a solid results in a solid solution. Partition coefficients can be measured experimentally in various ways or estimated by calculation based on a variety of methods. Despite formal recommendation to the contrary, the term partition coefficient remains the predominantly used term in the scientific literature. In contrast, the IUPAC recommends that the title term no longer be used, that it be replaced with more specific terms. For example, partition constant, defined as where KD is the process equilibrium constant, represents the concentration of solute A being tested, "org" and "aq" refer to the organic and aqueous phases respectively.
The IUPAC further recommends "partition ratio" for cases where transfer activity coefficients can be determined, "distribution ratio" for the ratio of total analytical concentrations of a solute between phases, regardless of chemical form. The partition coefficient, abbreviated P, is defined as a particular ratio of the concentrations of a solute between the two solvents for un-ionized solutes, the logarithm of the ratio is thus log P; when one of the solvents is water and the other is a non-polar solvent the log P value is a measure of lipophilicity or hydrophobicity. The defined precedent is for the lipophilic and hydrophilic phase types to always be in the numerator and denominator respectively. To a first approximation, the non-polar phase in such experiments is dominated by the un-ionized form of the solute, electrically neutral, though this may not be true for the aqueous phase. To measure the partition coefficient of ionizable solutes, the pH of the aqueous phase is adjusted such that the predominant form of the compound in solution is the un-ionized, or its measurement at another pH of interest requires consideration of all species, un-ionized and ionized.
A corresponding partition coefficient for ionizable compounds, abbreviated log P I, is derived for cases where there are dominant ionized forms of the molecule, such that one must consider partition of all forms, ionized and un-ionized, between the two phases. M is used to indicate the number of ionized forms. For instance, for an octanol–water partition, it is log P oct/wat I = log . To distinguish between this and the standard, un-ionized, partition coefficient, the un-ionized is assigned the symbol log P0, such that the indexed log P oct/wat I expression for ionized solutes becomes an extension of this, into the range of values I > 0. The distribution co
Bernthsen acridine synthesis
The Bernthsen acridine synthesis is the chemical reaction of a diarylamine heated with a carboxylic acid and zinc chloride to form a 9-substituted acridine. Using zinc chloride, one must heat the reaction to 200-270 °C for 24hrs; the use of polyphosphoric acid will give acridine products at a lower temperature, but with decreased yields
Potassium dichromate, K2Cr2O7, is a common inorganic chemical reagent, most used as an oxidizing agent in various laboratory and industrial applications. As with all hexavalent chromium compounds, it is chronically harmful to health, it is a crystalline ionic solid with a bright, red-orange color. The salt is popular in the laboratory because it is not deliquescent, in contrast to the more industrially relevant salt sodium dichromate. Potassium dichromate is prepared by the reaction of potassium chloride on sodium dichromate. Alternatively, it can be obtained from potassium chromate by roasting chromite ore with potassium hydroxide, it is soluble in water and in the dissolution process it ionizes: K2Cr2O7 → 2 K+ + Cr2O72− Cr2O72− + H2O ⇌ 2 CrO42− + 2 H+ Potassium dichromate is an oxidising agent in organic chemistry, is milder than potassium permanganate. It is used to oxidize alcohols, it converts primary alcohols into aldehydes and, under more forcing conditions, into carboxylic acids. In contrast, potassium permanganate tends to give carboxylic acids as the sole products.
Secondary alcohols are converted into ketones. For example, menthone may be prepared by oxidation of menthol with acidified dichromate. Tertiary alcohols cannot be oxidized. In an aqueous solution the color change exhibited can be used to test for distinguishing aldehydes from ketones. Aldehydes reduce dichromate from the +6 to the +3 oxidation state, changing color from orange to green; this color change arises. A ketone will show no such change because it cannot be oxidized further, so the solution will remain orange; when heated it decomposes with the evolution of oxygen. 4K2Cr2O7 → 4K2CrO4 + 2Cr2O3 + 3O2When an alkali is added to an orange red solution containing dichromate ions, a yellow solution is obtained due to the formation of chromate ions. For example, potassium chromate is produced industrially using potash: K2Cr2O7 + K2CO3 → 2 K2CrO4 + CO2The reaction is reversible. Treatment with cold sulphuric acid gives red crystals of chromic anhydride: K2Cr2O7 + 2H2SO4 → 2CrO3 + 2 KHSO4 + H2OOn heating with concentrated acid, oxygen is evolved: 2 K2Cr2O7 + 8H2SO4 → 2 K2SO4 + 2 Cr23 + 8 H2O + 3O2 Potassium dichromate has few major applications, as the sodium salt is dominant industrially.
The main use is. Like other chromium compounds, potassium dichromate has been used to prepare "chromic acid" for cleaning glassware and etching materials; because of safety concerns associated with hexavalent chromium, this practice has been discontinued. It is used as an ingredient in cement in which it retards the setting of the mixture and improves its density and texture; this usage causes contact dermatitis in construction workers. Potassium dichromate has uses in photography and in photographic screen printing, where it is used as an oxidizing agent together with a strong mineral acid. In 1839, Mungo Ponton discovered that paper treated with a solution of potassium dichromate was visibly tanned by exposure to sunlight, the discoloration remaining after the potassium dichromate had been rinsed out. In 1852, Henry Fox Talbot discovered that exposure to ultraviolet light in the presence of potassium dichromate hardened organic colloids such as gelatin and gum arabic, making them less soluble.
These discoveries soon led to the carbon print, gum bichromate, other photographic printing processes based on differential hardening. After exposure, the unhardened portion was rinsed away with warm water, leaving a thin relief that either contained a pigment included during manufacture or was subsequently stained with a dye; some processes depended on the hardening only, in combination with the differential absorption of certain dyes by the hardened or unhardened areas. Because some of these processes allowed the use of stable dyes and pigments, such as carbon black, prints with an high degree of archival permanence and resistance to fading from prolonged exposure to light could be produced. Dichromated colloids were used as photoresists in various industrial applications, most in the creation of metal printing plates for use in photomechanical printing processes. Chromium intensification or Photochromos uses potassium dichromate together with equal parts of concentrated hydrochloric acid diluted down to 10% v/v to treat weak and thin negatives of black and white photograph roll.
This solution reconverts the elemental silver particles in the film to silver chloride. After thorough washing and exposure to actinic light, the film can be redeveloped to its end-point yielding a stronger negative, able to produce a more satisfactory print. A potassium dichromate solution in sulfuric acid can be used to produce a reversal negative; this is effected by developing a black and white film but allowing the development to proceed more or less to the end point. The development is stopped by copious washing and the film treated in the acid dichromate solution; this converts the silver metal to silver sulfate, a compound, insensitive to light. After thorough washing and exposure to actinic light, the film is developed again allowing the unexposed silver halide to be reduced to silver metal; the results obtained can be unpredictable, but sometimes excellent results are obtained producing images that would otherwise be unobtainable. This process can be coupled with solarisation so that the end product resembles a negative and is suitable for printing in the normal way.
CrVI compounds have the property of tanning an
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
Carbon tetrachloride known by many other names is an organic compound with the chemical formula CCl4. It is a colourless liquid with a "sweet" smell, it has no flammability at lower temperatures. It was widely used in fire extinguishers, as a precursor to refrigerants and as a cleaning agent, but has since been phased out because of toxicity and safety concerns. Exposure to high concentrations of carbon tetrachloride can affect the central nervous system, degenerate the liver and kidneys. Prolonged exposure can be fatal. Carbon tetrachloride was synthesized by the French chemist Henri Victor Regnault in 1839 by the reaction of chloroform with chlorine, but now it is produced from methane: CH4 + 4 Cl2 → CCl4 + 4 HClThe production utilizes by-products of other chlorination reactions, such as from the syntheses of dichloromethane and chloroform. Higher chlorocarbons are subjected to "chlorinolysis": C2Cl6 + Cl2 → 2 CCl4Prior to the 1950s, carbon tetrachloride was manufactured by the chlorination of carbon disulfide at 105 to 130 °C: CS2 + 3Cl2 → CCl4 + S2Cl2The production of carbon tetrachloride has steeply declined since the 1980s due to environmental concerns and the decreased demand for CFCs, which were derived from carbon tetrachloride.
In 1992, production in the U. S./Europe/Japan was estimated at 720,000 tonnes. In the carbon tetrachloride molecule, four chlorine atoms are positioned symmetrically as corners in a tetrahedral configuration joined to a central carbon atom by single covalent bonds; because of this symmetrical geometry, CCl4 is non-polar. Methane gas has the same structure, making carbon tetrachloride a halomethane; as a solvent, it is well suited to dissolving other non-polar compounds and oils. It can dissolve iodine, it is somewhat volatile, giving off vapors with a smell characteristic of other chlorinated solvents, somewhat similar to the tetrachloroethylene smell reminiscent of dry cleaners' shops. Solid tetrachloromethane has two polymorphs: crystalline II below −47.5 °C and crystalline I above −47.5 °C. At −47.3 °C it has monoclinic crystal structure with space group C2/c and lattice constants a = 20.3, b = 11.6, c = 19.9, β = 111°. With a specific gravity greater than 1, carbon tetrachloride will be present as a dense nonaqueous phase liquid if sufficient quantities are spilled in the environment.
In organic chemistry, carbon tetrachloride serves as a source of chlorine in the Appel reaction. One specialty use of carbon tetrachloride is in stamp collecting, to reveal watermarks on postage stamps without damaging them. A small amount of the liquid was placed on the back of a stamp, sitting in a black glass or obsidian tray; the letters or design of the watermark could be seen. Carbon tetrachloride was used as a dry cleaning solvent, as a refrigerant, in lava lamps. In case of the latter, carbon tetrachloride is a key ingredient that adds weight to the otherwise buoyant wax, it once was a popular solvent in organic chemistry, because of its adverse health effects, it is used today. It is sometimes useful as a solvent for infrared spectroscopy, because there are no significant absorption bands > 1600 cm−1. Because carbon tetrachloride does not have any hydrogen atoms, it was used in proton NMR spectroscopy. In addition to being toxic, its dissolving power is low, its use has been superseded by deuterated solvents.
Use of carbon tetrachloride in determination of oil has been replaced by various other solvents, such as tetrachloroethylene. Because it has no C-H bonds, carbon tetrachloride does not undergo free-radical reactions, it is a useful solvent for halogenations either by the elemental halogen or by a halogenation reagent such as N-bromosuccinimide. In 1910, the Pyrene Manufacturing Company of Delaware filed a patent to use carbon tetrachloride to extinguish fires; the liquid was vaporized by the heat of combustion and extinguished flames, an early form of gaseous fire suppression. At the time it was believed the gas displaced oxygen in the area near the fire, but research found that the gas inhibits the chemical chain reaction of the combustion process. In 1911, Pyrene patented a portable extinguisher that used the chemical; the extinguisher consisted of a brass bottle with an integrated handpump, used to expel a jet of liquid toward the fire. As the container was unpressurized, it could be refilled after use.
Carbon tetrachloride was suitable for liquid and electrical fires and the extinguishers were carried on aircraft or motor vehicles. In the first half of the 20th century, another common fire extinguisher was a single-use, sealed glass globe known as a "fire grenade," filled with either carbon tetrachloride or salt water; the bulb could be thrown at the base of the flames to quench the fire. The carbon tetrachloride type could be installed in a spring-loaded wall fixture with a solder-based restraint; when the solder melted by high heat, the spring would either break the globe or launch it out of the bracket, allowing the extinguishing agent to be automatically dispersed into the fire. A well-known brand was the "Red Comet,", variously manufactured with other fire-fighting equipment in the Denver, Colorado area by the Red Comet Manufacturing Company from its founding in 1919 until manufacturing operations were closed in the early 1980s. Prior to the Montreal Protocol, large quantities of carbon tetrachloride were used to produce the chlorofluorocarbon re
Pyridine is a basic heterocyclic organic compound with the chemical formula C5H5N. It is structurally related to benzene, with one methine group replaced by a nitrogen atom, it is a flammable, weakly alkaline, water-soluble liquid with a distinctive, unpleasant fish-like smell. Pyridine is colorless; the pyridine ring occurs in many important compounds, including agrochemicals and vitamins. Pyridine was produced from coal tar. Today it is synthesized on the scale of about 20,000 tonnes per year worldwide; the molecular electric dipole moment is 2.2 debyes. Pyridine is diamagnetic and has a diamagnetic susceptibility of −48.7 × 10−6 cm3·mol−1. The standard enthalpy of formation is 100.2 kJ·mol−1 in the liquid phase and 140.4 kJ·mol−1 in the gas phase. At 25 °C pyridine has a viscosity of 0.88 mPa/s and thermal conductivity of 0.166 W·m−1·K−1. The enthalpy of vaporization is 35.09 kJ · mol − 1 at normal pressure. The enthalpy of fusion is 8.28 kJ·mol−1 at the melting point. The critical parameters of pyridine are pressure 6.70 MPa, temperature 620 K and volume 229 cm3·mol−1.
In the temperature range 340–426 °C its vapor pressure p can be described with the Antoine equation log 10 p = A − B C + T where T is temperature, A = 4.16272, B = 1371.358 K and C = −58.496 K. Akin to benzene, pyridine ring forms a C5N hexagon. Electron localization in pyridine is reflected in the shorter C–N ring bond, whereas the carbon–carbon bonds in the pyridine ring have the same 139 pm length as in benzene; these bond lengths lie between the values for the single and double bonds and are typical of aromatic compounds. Pyridine crystallizes in an orthorhombic crystal system with space group Pna21 and lattice parameters a = 1752 pm, b = 897 pm, c = 1135 pm, 16 formula units per unit cell. For comparison, crystalline benzene is orthorhombic, with space group Pbca, a = 729.2 pm, b = 947.1 pm, c = 674.2 pm, but the number of molecules per cell is only 4. This difference is related to the lower symmetry of the individual pyridine molecule. A trihydrate is known; the optical absorption spectrum of pyridine in hexane contains three bands at the wavelengths of 195 nm, 251 nm and 270 nm.
The 1H nuclear magnetic resonance spectrum of pyridine contains three signals with the integral intensity ratio of 2:1:2 that correspond to the three chemically different protons in the molecule. These signals originate from γ-proton and β-protons; the carbon analog of pyridine, has only one proton signal at 7.27 ppm. The larger chemical shifts of the α- and γ-protons in comparison to benzene result from the lower electron density in the α- and γ-positions, which can be derived from the resonance structures; the situation is rather similar for the 13C NMR spectra of pyridine and benzene: pyridine shows a triplet at δ = 150 ppm, δ = 124 ppm and δ = 136 ppm, whereas benzene has a single line at 129 ppm. All shifts are quoted for the solvent-free substances. Pyridine is conventionally detected by mass spectrometry methods; because of the electronegative nitrogen in the pyridine ring, the molecule is electron deficient. It, enters less into electrophilic aromatic substitution reactions than benzene derivatives.
Correspondingly pyridine is more prone to nucleophilic substitution, as evidenced by the ease of metalation by strong organometallic bases. The reactivity of pyridine can be distinguished for three chemical groups. With electrophiles, electrophilic substitution takes place where pyridine expresses aromatic properties. With nucleophiles, pyridine reacts at positions 2 and 4 and thus behaves similar to imines and carbonyls; the reaction with many Lewis acids results in the addition to the nitrogen atom of pyridine, similar to the reactivity of tertiary amines. The ability of pyridine and its derivatives to oxidize, forming amine oxides, is a feature of tertiary amines; the nitrogen center of pyridine features a basic lone pair of electrons. This lone pair does not overlap with the aromatic π-system ring pyridine is a basic, having chemical properties similar to those of tertiary amines. Protonation gives pyridinium, C5H5NH+; the pKa of the conjugate acid is 5.25. The structures of pyridine and pyridinium are identical.
The pyridinium cation is isoelectronic with benzene. Pyridinium p-toluenesulfonate is an illustrative pyridinium salt. In addition to protonation, pyridine undergoes N-centered alkylation, N-oxidation. Pyridine has a conjugated system of six π electrons; the molecule is planar and, follows the Hückel criteria for aromatic systems. In contrast to benzene, the electron density is not evenly distributed over the ring, reflecting the negative inductive effect of the nitrogen atom. For this reason, pyridine has a dipole moment and a weaker resonant stabilization than benzene (re