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Vapor Pressure
Vapor
Vapor
pressure 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
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Thermodynamic Equilibrium
Thermodynamic equilibrium
Thermodynamic equilibrium
is an axiomatic concept of thermodynamics. It is an internal state of a single thermodynamic system, or a relation between several thermodynamic systems connected by more or less permeable or impermeable walls. In thermodynamic equilibrium there are no net macroscopic flows of matter or of energy, either within a system or between systems. In a system in its own state of internal thermodynamic equilibrium, no macroscopic change occurs. Systems in mutual thermodynamic equilibrium are simultaneously in mutual thermal, mechanical, chemical, and radiative equilibria. Systems can be in one kind of mutual equilibrium, though not in others
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Thermodynamic Activity
In chemical thermodynamics, activity (symbol a) is a measure of the "effective concentration" of a species in a mixture, in the sense that the species' chemical potential depends on the activity of a real solution in the same way that it would depend on concentration for an ideal solution. By convention, activity is treated as a dimensionless quantity, although its value depends on customary choices of standard state for the species. The activity of pure substances in condensed phases (solid or liquids) is normally taken as unity (the number 1). Activity depends on temperature, pressure and composition of the mixture, among other things. For gases, the activity is the effective partial pressure, and is usually referred to as fugacity. The difference between activity and other measures of composition arises because molecules in non-ideal gases or solutions interact with each other, either to attract or to repel each other
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Knudsen Effusion Cell
In crystal growth, a Knudsen cell is an effusion evaporator source for relatively low partial pressure elementary sources (e.g. Ga, Al, Hg, As). Because it is easy to control the temperature of the evaporating material in Knudsen cells, they are commonly used in molecular-beam epitaxy. Development[edit] The Knudsen effusion cell was developed by Martin Knudsen (1871-1949). A typical Knudsen cell contains a crucible (made of pyrolytic boron nitride, quartz, tungsten or graphite), heating filaments (often made of metal tantalum), water cooling system, heat shields, and an orifice shutter. Vapor pressure
Vapor pressure
measurement[edit] The Knudsen cell is used to measure the vapor pressures of a solid with very low vapor pressure. Such a solid forms a vapor at low pressure by sublimation
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Millimeter Of Mercury
A millimeter of mercury is a manometric unit of pressure, formerly defined as the extra pressure generated by a column of mercury one millimetre high and now defined as precisely 7002133322387415000♠133.322387415 pascals.[1] It is denoted by the symbol mmHg[2] or mm Hg.[3] Although not an SI unit, the millimeter of mercury is still routinely used in medicine, meteorology, aviation, and many other scientific fields. One millimeter of mercury is approximately 1 Torr, which is 1/760 of standard atmospheric pressure (7005101325000000000♠101325/760 = 7002133322368421053♠133.322368421053 pascals)
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Volatile Anesthetics
An inhalational anaesthetic is a chemical compound possessing general anaesthetic properties that can be delivered via inhalation. They are administered by anaesthetists (a term which includes anaesthesiologists, nurse anaesthetists, and anaesthesiologist assistants) through an anaesthesia mask, laryngeal mask airway or tracheal tube connected to an anaesthetic vaporiser and an anaesthetic delivery system
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Methyl Chloride
Chloromethane, also called methyl chloride, Refrigerant-40, R-40 or HCC 40, is a chemical compound of the group of organic compounds called haloalkanes. It was once widely used as a refrigerant. It is a colorless extremely flammable gas with a mildly sweet odor. Due to concerns about its toxicity, it is no longer present in consumer products. Chloromethane
Chloromethane
was first synthesized by the French chemists Jean-Baptiste Dumas
Jean-Baptiste Dumas
and Eugene Peligot
Eugene Peligot
in 1835 by boiling a mixture of methanol, sulfuric acid, and sodium chloride. This method is similar to that used today.Contents1 Occurrence1.1 Marine 1.2 Biogenesis2 Production 3 Uses 4 Safety 5 References 6 External linksOccurrence[edit] Chloromethane
Chloromethane
is the most abundant organohalogen, anthropogenic or natural, in the atmosphere
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Raoult's Law
Raoult's law
Raoult's law
(/ˈrɑːuːlz/ law) is a law of thermodynamics established by French chemist François-Marie Raoult
François-Marie Raoult
in 1887. [1] It states that the partial vapor pressure of each component of an ideal mixture of liquids is equal to the vapour pressure of the pure component multiplied by its mole fraction in the mixture
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Fugacity
In chemical thermodynamics, the fugacity of a real gas is an effective partial pressure which replaces the mechanical partial pressure in an accurate computation of the chemical equilibrium constant. It is equal to the pressure of an ideal gas which has the same chemical potential as the real gas. For example, nitrogen gas (N2) at 0 °C and a pressure of P = 100 atm has a fugacity of f = 97.03 atm.[1] This means that the chemical potential of real nitrogen at a pressure of 100 atm is less than if nitrogen were an ideal gas at 100 atm. Ideal-gas nitrogen at 97.03 atm would have the same chemical potential as real nitrogen at 100 atm. Fugacities are determined experimentally or estimated from various models such as a Van der Waals gas
Van der Waals gas
that are closer to reality than an ideal gas
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Mole Fraction
In chemistry, the mole fraction or molar fraction (xi) is defined as the amount of a constituent (expressed in moles), ni, divided by the total amount of all constituents in a mixture (also expressed in moles), ntot:[1] x i = n i n t o t displaystyle x_ i = frac n_ i n_ mathrm tot The sum of all the mole fractions is equal to 1: ∑ i = 1 N n i = n t o t ; ∑ i = 1 N x i = 1 displaystyle sum _ i=1 ^ N n_ i =n_ mathrm tot ;;sum _ i=1 ^ N x_
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London Force
London dispersion forces (LDF, also known as dispersion forces, London forces, instantaneous dipole–induced dipole forces, or loosely van der Waals forces) are a type of force acting between atoms and molecules.[1] They are part of the van der Waals forces. The LDF is named after the German-American physicist Fritz London. The LDF is a weak intermolecular force arising from quantum-induced instantaneous polarization multipoles in molecules. They can therefore act between molecules without permanent multipole moments.Contents1 Introduction 2 Quantum mechanical theory of dispersion forces 3 Relative magnitude 4 ReferencesIntroduction[edit] London forces are exhibited by non-polar molecules because of the presence of correlated movements of the electrons in interacting molecules
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Azeotrope
An azeotrope (gK /əˈziːəˌtrəʊp/, US /əˈziəˌtroʊp/)[1] or a constant boiling point mixture is a mixture of two or more liquids whose proportions cannot be altered or changed by simple distillation.[2] 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,[3] and many azeotropes of three or more compounds are also known.[4] 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
Raoult's law
forms a minimum boiling azeotrope at a specific composition
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Crystal
A crystal or crystalline solid is a solid material whose constituents (such as atoms, molecules, or ions) are arranged in a highly ordered microscopic structure, forming a crystal lattice that extends in all directions.[1][2] In addition, macroscopic single crystals are usually identifiable by their geometrical shape, consisting of flat faces with specific, characteristic orientations. The scientific study of crystals and crystal formation is known as crystallography. The process of crystal formation via mechanisms of crystal growth is called crystallization or solidification. The word crystal derives from the Ancient Greek
Ancient Greek
word κρύσταλλος (krustallos), meaning both "ice" and "rock crystal",[3] from κρύος (kruos), "icy cold, frost".[4][5] Examples of large crystals include snowflakes, diamonds, and table salt. Most inorganic solids are not crystals but polycrystals, i.e. many microscopic crystals fused together into a single solid
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Square Meter
The square metre (International spelling as used by the International Bureau of Weights and Measures) or square meter (American spelling) is the SI derived unit
SI derived unit
of area, with symbol m2 (33A1 in Unicode[1]). It is the area of a square whose sides measure exactly one metre. The square metre is derived from the SI base unit
SI base unit
of the metre, which itself is defined as the length of the path travelled by light in absolute vacuum during a time interval of 1/299 792 458 of a second. Adding and subtracting SI prefixes creates multiples and submultiples; however, as the unit is squared, the order of magnitude difference between units doubles from their comparable linear units
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Sublimation (physics)
Sublimation is the phase transition of a substance directly from the solid to the gas phase without passing through the intermediate liquid phase.[1] Sublimation is an endothermic process that occurs at temperatures and pressures below a substance's triple point in its phase diagram, which corresponds to the lowest pressure at which the substance can exist as a liquid. The reverse process of sublimation is deposition or desublimation, in which a substance passes directly from a gas to a solid phase.[2] Sublimation has also been used as a generic term to describe a solid-to-gas transition (sublimation) followed by a gas-to-solid transition (deposition).[3] At normal pressures, most chemical compounds and elements possess three different states at different temperatures. In these cases, the transition from the solid to the gaseous state requires an intermediate liquid state. The pressure referred to is the partial pressure of the substance, not the total (e.g
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Naphthalene
Naphthalene
Naphthalene
is an organic compound with formula C 10H 8. It is the simplest polycyclic aromatic hydrocarbon, and is a white crystalline solid with a characteristic odor that is detectable at concentrations as low as 0.08 ppm by mass.[13] As an aromatic hydrocarbon, naphthalene's structure consists of a fused pair of benzene rings
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