1.
Solvable group
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In mathematics, more specifically in the field of group theory, a solvable group or soluble group is a group that can be constructed from abelian groups using extensions. Equivalently, a group is a group whose derived series terminates in the trivial subgroup. Historically, the word solvable arose from Galois theory and the proof of the unsolvability of quintic equation. Specifically, an equation is solvable by radicals if and only if the corresponding Galois group is solvable. Or equivalently, if its derived series, the normal series G ▹ G ▹ G ▹ ⋯. These two definitions are equivalent, since for every group H and every normal subgroup N of H, the least n such that G = is called the derived length of the solvable group G. For finite groups, an equivalent definition is that a group is a group with a composition series all of whose factors are cyclic groups of prime order. This is equivalent because a group has finite composition length. The Jordan–Hölder theorem guarantees that if one composition series has this property, for the Galois group of a polynomial, these cyclic groups correspond to nth roots over some field. All abelian groups are trivially solvable – a subnormal series being given by just the group itself, but non-abelian groups may or may not be solvable. More generally, all nilpotent groups are solvable, in particular, finite p-groups are solvable, as all finite p-groups are nilpotent. A small example of a solvable, non-nilpotent group is the symmetric group S3, in fact, as the smallest simple non-abelian group is A5, it follows that every group with order less than 60 is solvable. The group S5 is not solvable — it has a series, giving factor groups isomorphic to A5 and C2. Generalizing this argument, coupled with the fact that An is a normal, maximal, non-abelian simple subgroup of Sn for n >4, we see that Sn is not solvable for n >4. This is a key step in the proof that for every n >4 there are polynomials of n which are not solvable by radicals. This property is used in complexity theory in the proof of Barringtons theorem. The celebrated Feit–Thompson theorem states that every group of odd order is solvable. In particular this implies that if a group is simple
2.
Solid
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Solid is one of the four fundamental states of matter. It is characterized by structural rigidity and resistance to changes of shape or volume, unlike a liquid, a solid object does not flow to take on the shape of its container, nor does it expand to fill the entire volume available to it like a gas does. The atoms in a solid are tightly bound to other, either in a regular geometric lattice or irregularly. The branch of physics deals with solids is called solid-state physics. Materials science is concerned with the physical and chemical properties of solids. Solid-state chemistry is concerned with the synthesis of novel materials, as well as the science of identification. The atoms, molecules or ions which make up solids may be arranged in a repeating pattern. Materials whose constituents are arranged in a regular pattern are known as crystals, in some cases, the regular ordering can continue unbroken over a large scale, for example diamonds, where each diamond is a single crystal. Almost all common metals, and many ceramics, are polycrystalline, in other materials, there is no long-range order in the position of the atoms. These solids are known as amorphous solids, examples include polystyrene, whether a solid is crystalline or amorphous depends on the material involved, and the conditions in which it was formed. Solids which are formed by slow cooling will tend to be crystalline, likewise, the specific crystal structure adopted by a crystalline solid depends on the material involved and on how it was formed. While many common objects, such as an ice cube or a coin, are chemically identical throughout, for example, a typical rock is an aggregate of several different minerals and mineraloids, with no specific chemical composition. Wood is an organic material consisting primarily of cellulose fibers embedded in a matrix of organic lignin. In materials science, composites of more than one constituent material can be designed to have desired properties, the forces between the atoms in a solid can take a variety of forms. For example, a crystal of sodium chloride is made up of sodium and chlorine. In diamond or silicon, the atoms share electrons and form covalent bonds, in metals, electrons are shared in metallic bonding. Some solids, particularly most organic compounds, are together with van der Waals forces resulting from the polarization of the electronic charge cloud on each molecule. The dissimilarities between the types of solid result from the differences between their bonding, metals typically are strong, dense, and good conductors of both electricity and heat
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Liquid
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A liquid is a nearly incompressible fluid that conforms to the shape of its container but retains a constant volume independent of pressure. As such, it is one of the four states of matter. A liquid is made up of tiny vibrating particles of matter, such as atoms, water is, by far, the most common liquid on Earth. Like a gas, a liquid is able to flow and take the shape of a container, most liquids resist compression, although others can be compressed. Unlike a gas, a liquid does not disperse to fill every space of a container, a distinctive property of the liquid state is surface tension, leading to wetting phenomena. The density of a liquid is usually close to that of a solid, therefore, liquid and solid are both termed condensed matter. On the other hand, as liquids and gases share the ability to flow, although liquid water is abundant on Earth, this state of matter is actually the least common in the known universe, because liquids require a relatively narrow temperature/pressure range to exist. Most known matter in the universe is in form as interstellar clouds or in plasma form within stars. Liquid is one of the four states of matter, with the others being solid, gas. Unlike a solid, the molecules in a liquid have a greater freedom to move. The forces that bind the molecules together in a solid are only temporary in a liquid, a liquid, like a gas, displays the properties of a fluid. A liquid can flow, assume the shape of a container, if liquid is placed in a bag, it can be squeezed into any shape. These properties make a suitable for applications such as hydraulics. Liquid particles are bound firmly but not rigidly and they are able to move around one another freely, resulting in a limited degree of particle mobility. As the temperature increases, the vibrations of the molecules causes distances between the molecules to increase. When a liquid reaches its point, the cohesive forces that bind the molecules closely together break. If the temperature is decreased, the distances between the molecules become smaller, only two elements are liquid at standard conditions for temperature and pressure, mercury and bromine. Four more elements have melting points slightly above room temperature, francium, caesium, gallium and rubidium, metal alloys that are liquid at room temperature include NaK, a sodium-potassium metal alloy, galinstan, a fusible alloy liquid, and some amalgams
4.
Gas
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Gas is one of the four fundamental states of matter. A pure gas may be made up of atoms, elemental molecules made from one type of atom. A gas mixture would contain a variety of pure gases much like the air, what distinguishes a gas from liquids and solids is the vast separation of the individual gas particles. This separation usually makes a colorless gas invisible to the human observer, the interaction of gas particles in the presence of electric and gravitational fields are considered negligible as indicated by the constant velocity vectors in the image. One type of commonly known gas is steam, the gaseous state of matter is found between the liquid and plasma states, the latter of which provides the upper temperature boundary for gases. Bounding the lower end of the temperature scale lie degenerative quantum gases which are gaining increasing attention, high-density atomic gases super cooled to incredibly low temperatures are classified by their statistical behavior as either a Bose gas or a Fermi gas. For a comprehensive listing of these states of matter see list of states of matter. The only chemical elements which are stable multi atom homonuclear molecules at temperature and pressure, are hydrogen, nitrogen and oxygen. These gases, when grouped together with the noble gases. Alternatively they are known as molecular gases to distinguish them from molecules that are also chemical compounds. The word gas is a neologism first used by the early 17th-century Flemish chemist J. B. van Helmont, according to Paracelsuss terminology, chaos meant something like ultra-rarefied water. An alternative story is that Van Helmonts word is corrupted from gahst and these four characteristics were repeatedly observed by scientists such as Robert Boyle, Jacques Charles, John Dalton, Joseph Gay-Lussac and Amedeo Avogadro for a variety of gases in various settings. Their detailed studies ultimately led to a relationship among these properties expressed by the ideal gas law. Gas particles are separated from one another, and consequently have weaker intermolecular bonds than liquids or solids. These intermolecular forces result from interactions between gas particles. Like-charged areas of different gas particles repel, while oppositely charged regions of different gas particles attract one another, transient, randomly induced charges exist across non-polar covalent bonds of molecules and electrostatic interactions caused by them are referred to as Van der Waals forces. The interaction of these forces varies within a substance which determines many of the physical properties unique to each gas. A comparison of boiling points for compounds formed by ionic and covalent bonds leads us to this conclusion, the drifting smoke particles in the image provides some insight into low pressure gas behavior
5.
Chemical substance
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A chemical substance is a form of matter that has constant chemical composition and characteristic properties. It cannot be separated into components by physical methods, i. e. without breaking chemical bonds. Chemical substances can be chemical elements, chemical compounds, ions or alloys, Chemical substances are often called pure to set them apart from mixtures. A common example of a substance is pure water, it has the same properties. Other chemical substances commonly encountered in pure form are diamond, gold, table salt, however, in practice, no substance is entirely pure, and chemical purity is specified according to the intended use of the chemical. Chemical substances exist as solids, liquids, gases, or plasma, Chemical substances may be combined or converted to others by means of chemical reactions. Forms of energy, such as light and heat, are not matter, a chemical substance may well be defined as any material with a definite chemical composition in an introductory general chemistry textbook. According to this definition a chemical substance can either be a chemical element or a pure chemical compound. But, there are exceptions to this definition, a substance can also be defined as a form of matter that has both definite composition and distinct properties. The chemical substance index published by CAS also includes several alloys of uncertain composition, in geology, substances of uniform composition are called minerals, while physical mixtures of several minerals are defined as rocks. Many minerals, however, mutually dissolve into solid solutions, such that a rock is a uniform substance despite being a mixture in stoichiometric terms. Feldspars are an example, anorthoclase is an alkali aluminium silicate. In law, chemical substances may include both pure substances and mixtures with a composition or manufacturing process. For example, the EU regulation REACH defines monoconstituent substances, multiconstituent substances and substances of unknown or variable composition, the latter two consist of multiple chemical substances, however, their identity can be established either by direct chemical analysis or reference to a single manufacturing process. For example, charcoal is a complex, partially polymeric mixture that can be defined by its manufacturing process. Therefore, although the chemical identity is unknown, identification can be made to a sufficient accuracy. The CAS index also includes mixtures, polymers almost always appear as mixtures of molecules of multiple molar masses, each of which could be considered a separate chemical substance. However, the polymer may be defined by a precursor or reaction
6.
Solution
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In chemistry, a solution is a homogeneous mixture composed of two or more substances. In such a mixture, a solute is a substance dissolved in another substance, the mixing process of a solution happens at a scale where the effects of chemical polarity are involved, resulting in interactions that are specific to solvation. The solution assumes the characteristics of the solvent when the solvent is the fraction of the mixture. The concentration of a solute in a solution is the mass of that solute expressed as a percentage of the mass of the whole solution, a solution is a homogeneous mixture of two or more substances. The particles of solute in a solution cannot be seen by the naked eye, a solution does not allow beams of light to scatter. The solute from a solution cannot be separated by filtration and it is composed of only one phase. Homogeneous means that the components of the form a single phase. Heterogeneous means that the components of the mixture are of different phase, the properties of the mixture can be uniformly distributed through the volume but only in absence of diffusion phenomena or after their completion. Usually, the present in the greatest amount is considered the solvent. Solvents can be gases, liquids or solids, one or more components present in the solution other than the solvent are called solutes. The solution has the physical state as the solvent. If the solvent is a gas, only gases are dissolved under a set of conditions. An example of a solution is air. Since interactions between molecules play almost no role, dilute gases form rather trivial solutions, in part of the literature, they are not even classified as solutions, but addressed as mixtures. If the solvent is a liquid, then almost all gases, liquids, here are some examples, Gas in liquid, Oxygen in water Carbon dioxide in water – a less simple example, because the solution is accompanied by a chemical reaction. Liquid in liquid, The mixing of two or more substances of the same chemistry but different concentrations to form a constant, alcoholic beverages are basically solutions of ethanol in water. Solid in liquid, Sucrose in water Sodium chloride or any other salt in water, solutions in water are especially common. Counterexamples are provided by liquid mixtures that are not homogeneous, colloids, body fluids are examples for complex liquid solutions, containing many solutes
7.
Dissolution (chemistry)
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Solid solutions are the result of dissolution of one solid into another, and occur, e. g. in metal alloys, where their formation is governed and described by the relevant phase diagram. In the case of a crystalline solid dissolving in a liquid, for liquids and gases, the molecules must be able to form non-covalent intermolecular interactions with those of the solvent for a solution to form. Some distinctions can be made between solvation, dissolution, and solubility, gaseous elements and compounds will dissolve in liquids dependent on the interaction of their bonds with the liquid solvent. Gaseous elements and compounds may dissolve in another liquid depending on the compatibility of the chemical and physical bonds in the substance with those of the solvent. Hydrogen bonds play an important role in aqueous dissolution, for ionic compounds, dissolution takes place when the ionic lattice breaks up and the separate ions are then solvated. e. NaCl ⇌ Na+ + Cl− The solubility of salts in water is generally determined by the degree of solvation of the ions by water molecules. Such coordination complexes occur by water donating spare electrons on the atom to the ion. This in turn results in the remaining particle possessing either a net surface charge. The dissolution of minerals such as silicates occurs by several mechanisms which depend on the composition of the mineral. Dissolution rates partially depend on solution pH, adsorbed protons or hydroxides polarize the mineral surface and weaken cation-oxygen bonds, accelerating dissolution. Silicate minerals containing metal cations undergo incongruent dissolution as the cations out of the mineral faster than the silica lattice degrades. Incongruent dissolution results in a layer with different composition than the bulk. The reductive dissolution of a metal oxide can occur when a redox event in solution reduces a cation. Dissolution occurs when the cation is unstable in the solid material. In minerals such as oxides, reduction may be caused by electron transfer from organic molecules or bacteria in anoxic waters or soils. Charge carriers responsible for reductive dissolution may also be introduced by photoexcitation or by electrochemical poising at negative potentials, reductive dissolution is integral to natural geochemical phenomena such as the iron cycle. However, Fe 2 + is unstable in the oxide lattice relative to the solution and is subsequently solvated, reductants causing reductive dissolution include natural electron donors such as ascorbic acid and Fe 2 +. Chelating species such as oxalate accelerate the process by detaching surface-bound Fe 2 +, reductive dissolution is also promoted by light
8.
Saturation (chemistry)
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In chemistry, saturation has diverse meanings, all based on the idea of reaching a maximum capacity. In organic chemistry, a compound is a hydrocarbon which has no double or triple bonds. For example, consider three similar organic compounds that are less saturated, ethane, ethylene, and ethyne. Ethane, a saturated compound, has only single bonds between its hydrogen and carbon atoms. Ethylene, a compound of ethane, has a carbon double bond. Finally ethyne, the completely unsaturated compound of ethane, has a triple bond. The concept of saturation can be described using various naming systems, formulas, for instance, IUPAC nomenclature is a system of naming conventions used to describe the type and location of unsaturation within organic compounds. The degree of unsaturation is a used to summarize and diagram the amount of hydrogen that a compound can bind. Unsaturation can be determined by NMR, mass spectrometry and IR spectroscopy, tallow consists mainly of triglycerides, whose major constituents are derived from the saturated stearic and monounsaturated oleic acids. Many vegetable oils contain fatty acids with one or more double bonds in them, in organometallic chemistry, an unsaturated complex has fewer than 18 valence electrons and thus is susceptible to oxidative addition or coordination of an additional ligand. Unsaturation is characteristic of many catalysts because it is usually a requirement for substrate activation, in contrast, a coordinatively saturated complex resists undergoing substitution and oxidative addition reactions. In physical chemistry, saturation is the point at which the solute of a substance can dissolve no more of that substance and additional amounts of it will appear as a separate phase. This point of concentration, the saturation point, depends on the temperature and pressure of the solution as well as the chemical nature of the substances involved. Impurities, being present in lower concentration, do not saturate the solvent. If a change in conditions means that the concentration is higher than the saturation point. In physical chemistry, when referring to processes, saturation denotes the degree at which a binding site is fully occupied. For example, base refers to the fraction of exchangeable cations that are base cations. In biochemistry, the saturation refers to the fraction of total protein binding sites that are occupied at any given time
9.
Solid solution
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A solid solution is a solid-state solution of one or more solutes in a solvent. The solid solution needs to be distinguished from a mechanical mixtures of powdered solids like two salts, sugar and salt, etc, the mechanical mixtures have total or partial miscibility gap in solid state. Examples of solid solutions include crystallized salts from their liquid mixture, metal alloys, in the case of metal alloys intermetallic compounds occur frequently. The solute may incorporate into the solvent crystal lattice substitutionally, by replacing a solvent particle in the lattice, or interstitially, by fitting into the space between solvent particles. Both of these types of solid solution affect the properties of the material by distorting the crystal lattice, some mixtures will readily form solid solutions over a range of concentrations, while other mixtures will not form solid solutions at all. The propensity for any two substances to form a solution is a complicated matter involving the chemical, crystallographic. 1 displays an alloy of two metals which forms a solution at all relative concentrations of the two species. Solid solutions have important commercial and industrial applications, as such mixtures often have superior properties to pure materials, many metal alloys are solid solutions. Even small amounts of solute can affect the electrical and physical properties of the solvent, the binary phase diagram in Fig.2 shows the phases of a mixture of two substances in varying concentrations, A and B. The region labeled α is a solution, with B acting as the solute in a matrix of A. On the other end of the scale, the region labeled β is also a solid solution. The large solid region in between the α and β solid solutions, labeled α + β, is not a solid solution, at other proportions, the material will enter a mushy or pasty phase until it warms up to being completely melted. The mixture at the dip point of the diagram is called a eutectic alloy, lead-tin mixtures formulated at that point are useful when soldering electronic components, particularly if done manually, since the solid phase is quickly entered as the solder cools. In contrast, when lead-tin mixtures were used to solder seams in automobile bodies a pasty state enabled a shape to be formed with a paddle or tool. When a solid solution becomes unstable — due to a temperature, for example — exsolution occurs. This is mainly caused by difference in cation size, cations which have a large difference in radii are not likely to readily substitute. Take the alkali feldspar minerals for example, whose end members are albite, NaAlSi3O8 and microcline and this leads to exsolution where they will separate into two separate phases. In the case of the feldspar minerals, thin white albite layers will alternate between typically pink microcline
10.
Solubility equilibrium
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Solubility equilibrium is a type of dynamic equilibrium that exists when a chemical compound in the solid state is in chemical equilibrium with a solution of that compound. The solid may dissolve unchanged, with dissociation or with chemical reaction with another constituent of the solvent, each type of equilibrium is characterized by a temperature-dependent equilibrium constant. Solubility equilibria are important in pharmaceutical, environmental and many other scenarios, a solubility equilibrium exists when a chemical compound in the solid state is in chemical equilibrium with a solution of that compound. When equilibrium is established, the solution is said to be saturated, the concentration of the solute in a saturated solution is known as the solubility. Units of solubility may be molar or expressed as mass per unit volume, a solution containing a higher concentration of solute than the solubility is said to be supersaturated. A supersaturated solution may be induced to come to equilibrium by the addition of a seed which may be a crystal of the solute, or a tiny solid particle. There are three types of solubility equilibria. The equilibrium constant is known in this case as a solubility product and this is characteristic of the dissolution of weak acids or weak bases in aqueous media of varying pH. In each case a constant can be specified as a quotient of activities. This equilibrium constant is dimensionless as activity is a dimensionless quantity, however, use of activities is very inconvenient, so the equilibrium constant is usually divided by the quotient of activity coefficients, to become a quotient of concentrations. See equilibrium chemistry#Equilibrium constant for details, moreover, the concentration of solvent is usually taken to be constant and so is also subsumed into the equilibrium constant. For these reasons, the constant for a solubility equilibrium has dimensions related to the scale on which concentrations are measured, equilibria are defined for specific crystal phases. Therefore, the solubility product is expected to be different depending on the phase of the solid, for example, aragonite and calcite will have different solubility products even though they have both the same chemical identity. Under any given conditions one phase will be more stable than the other, therefore. However, kinetic factors may favor the formation the unfavorable precipitate, the thermodynamic solubility constant is defined for large monocrystals. Solubility will increase with decreasing size of solute particle because of the surface energy. This effect is small unless particles become very small, typically smaller than 1 μm. The effect of the size on solubility constant can be quantified as follows
11.
Supersaturation
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Supersaturation is a state of a solution that contains more of the dissolved material than could be dissolved by the solvent under normal circumstances. It can also refer to a vapor of a compound that has a higher pressure than the pressure of that compound. Special conditions need to be met in order to generate a supersaturated solution, one of the easiest ways to do this relies on the temperature dependence of solubility. As a general rule, the heat is added to a system. Therefore, at high temperatures, more solute can be dissolved than at room temperature, the precipitation or crystallization of the solute takes longer than the actual cooling time because the molecules need to meet up and form the precipitate without being knocked apart by water. Thus, the larger the molecule, the longer it will take to crystallize due to the principles of Brownian motion, the condition of supersaturation does not necessarily have to be reached through the manipulation of heat. The ideal gas law suggests that pressure and volume can also be changed to force a system into a supersaturated state, if the volume of solvent is decreased, the concentration of the solute can be above the saturation point and thus create a supersaturated solution. The decrease in volume is most commonly generated through evaporation, similarly, an increase in pressure can drive a solution to a supersaturated state. All three of these rely on the fact that the conditions of the solution can be changed quicker than the solute can precipitate or crystallize out. Supersaturated solutions will also undergo crystallization under specific conditions, in a normal solution, once the maximum amount of solute is dissolved, adding more solute would either cause the dissolved solute to precipitate out and/or for the solute to not dissolve at all. Similarly, there are cases wherein solubility of a solution is decreased by manipulating temperature, pressure, or volume. In these cases, the solute will simply precipitate out and this is due to the fact that a supersaturated solution is in a higher energy state than a saturated solution. Crystallization will occur to allow the solution to reach an energy state. The activation energy comes in the form of a crystal being added to the liquid solution. This nuclei can be added from another source, which is known as seeding, or can spontaneously form within the solution due partly to ion. This process is known as primary nucleation and it is necessary for the nuclei to be identical to the solute that is crystallizing. This will allow for the ions to build up on the nuclei. There are a multitude of factors that affect the rate and order of magnitude with which crystallization proceeds as well as the difference in formation of crystallites
12.
Metastability
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In physics, metastability denotes the phenomenon when a dynamical system spends an extended time in a configuration other than the systems state of least energy. During a metastable state of finite lifetime, all state-describing parameters reach, a conceptual analogy may be drawn with a ball resting in a hollow on a slope. Small perturbations will result in the ball settling back into its current hollow, isomerisation is another common example of metastability. Higher energy isomers are long lived as they are prevented from rearranging to their ground state by barriers in the potential energy. The metastability concept originated in the physics of first-order phase transitions and it then acquired new meaning in the study of aggregated subatomic particles or in molecules, macromolecules or clusters of atoms and molecules. Later, it was borrowed for the study of decision-making and information transmission systems, many complex natural and man-made systems can demonstrate metastability. Metastability is common in physics and chemistry – from an atom to statistical ensembles of molecules at molecular levels or as a whole, the abundance of states is more prevalent as the systems grow larger and/or if the forces of their mutual interaction are spatially less uniform or more diverse. In these systems, the equivalent of thermal fluctuations in molecular systems is the noise that affects signal propagation. Non-equilibrium thermodynamics is a branch of physics that studies the dynamics of statistical ensembles of molecules via unstable states, being stuck in a thermodynamic trough without being at the lowest energy state is known as having kinetic stability or being kinetically persistent. The particular motion or kinetics of the atoms involved has resulted in getting stuck, metastable states of matter range from melting solids, boiling liquids and sublimating solids to supercooled liquids or superheated liquid-gas mixtures. Extremely pure, supercooled water stays liquid below 0 °C and remains so until applied vibrations or condensing seed doping initiates crystallization centers and this is a common situation for the droplets of atmospheric clouds. Metastable phases are common in condensed matter, for example, diamond is a metastable form of carbon at standard temperature and pressure. It can be converted to graphite, but only after overcoming an activation energy – an intervening hill, martensite is a metastable phase used to control the hardness of most steel. The bonds between the blocks of polymers such as DNA, RNA and proteins are also metastable. Metastable polymorphs of silica are commonly observed, in some cases, such as in the allotropes of solid boron, acquiring a sample of the stable phase is difficult. Generally speaking, emulsions/colloidal systems and glasses are metastable, sandpiles are one system which can exhibit metastability if a steep slope or tunnel is present. Sand grains form a pile due to friction and it is possible for an entire large sand pile to reach a point where it is stable, but the addition of a single grain causes large parts of it to collapse. The avalanche is a problem with large piles of snow
13.
Effervescence
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Effervescence is the escape of gas from an aqueous solution and the foaming or fizzing that results from that release. The word effervescence is derived from the Latin verb fervere, preceded by the adverb ex and it has the same linguistic root as the word fermentation. Effervescence can also be observed when opening a bottle of champagne, the visible bubbles are produced by the escape from solution of the dissolved gas. Although CO2 is most common for beverages, nitrogen gas is sometimes added to certain draught beers. The smaller bubble size creates a beer head. Due to the solubility of nitrogen in beer, kegs or widgets are used for this. In the laboratory, an example of effervescence is seen if hydrochloric acid is added to a block of limestone. If a few pieces of marble or a tablet are put in hydrochloric acid in a test tube fitted with a bung
14.
Chemical reaction
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A chemical reaction is a process that leads to the transformation of one set of chemical substances to another. Nuclear chemistry is a sub-discipline of chemistry that involves the reactions of unstable. The substance initially involved in a reaction are called reactants or reagents. Chemical reactions are characterized by a chemical change, and they yield one or more products. Reactions often consist of a sequence of individual sub-steps, the elementary reactions. Chemical reactions are described with chemical equations, which present the starting materials, end products. Chemical reactions happen at a characteristic reaction rate at a given temperature, typically, reaction rates increase with increasing temperature because there is more thermal energy available to reach the activation energy necessary for breaking bonds between atoms. Reactions may proceed in the forward or reverse direction until they go to completion or reach equilibrium, Reactions that proceed in the forward direction to approach equilibrium are often described as spontaneous, requiring no input of free energy to go forward. Non-spontaneous reactions require input of energy to go forward. Different chemical reactions are used in combinations during chemical synthesis in order to obtain a desired product, in biochemistry, a consecutive series of chemical reactions form metabolic pathways. These reactions are catalyzed by protein enzymes. Chemical reactions such as combustion in fire, fermentation and the reduction of ores to metals were known since antiquity, in the Middle Ages, chemical transformations were studied by Alchemists. They attempted, in particular, to lead into gold, for which purpose they used reactions of lead. The process involved heating of sulfate and nitrate minerals such as sulfate, alum. In the 17th century, Johann Rudolph Glauber produced hydrochloric acid and sodium sulfate by reacting sulfuric acid, further optimization of sulfuric acid technology resulted in the contact process in the 1880s, and the Haber process was developed in 1909–1910 for ammonia synthesis. From the 16th century, researchers including Jan Baptist van Helmont, Robert Boyle, the phlogiston theory was proposed in 1667 by Johann Joachim Becher. It postulated the existence of an element called phlogiston, which was contained within combustible bodies. This proved to be false in 1785 by Antoine Lavoisier who found the explanation of the combustion as reaction with oxygen from the air
15.
Water
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Water is a transparent and nearly colorless chemical substance that is the main constituent of Earths streams, lakes, and oceans, and the fluids of most living organisms. Its chemical formula is H2O, meaning that its molecule contains one oxygen, Water strictly refers to the liquid state of that substance, that prevails at standard ambient temperature and pressure, but it often refers also to its solid state or its gaseous state. It also occurs in nature as snow, glaciers, ice packs and icebergs, clouds, fog, dew, aquifers, Water covers 71% of the Earths surface. It is vital for all forms of life. Only 2. 5% of this water is freshwater, and 98. 8% of that water is in ice and groundwater. Less than 0. 3% of all freshwater is in rivers, lakes, and the atmosphere, a greater quantity of water is found in the earths interior. Water on Earth moves continually through the cycle of evaporation and transpiration, condensation, precipitation. Evaporation and transpiration contribute to the precipitation over land, large amounts of water are also chemically combined or adsorbed in hydrated minerals. Safe drinking water is essential to humans and other even though it provides no calories or organic nutrients. There is a correlation between access to safe water and gross domestic product per capita. However, some observers have estimated that by 2025 more than half of the population will be facing water-based vulnerability. A report, issued in November 2009, suggests that by 2030, in developing regions of the world. Water plays an important role in the world economy, approximately 70% of the freshwater used by humans goes to agriculture. Fishing in salt and fresh water bodies is a source of food for many parts of the world. Much of long-distance trade of commodities and manufactured products is transported by boats through seas, rivers, lakes, large quantities of water, ice, and steam are used for cooling and heating, in industry and homes. Water is an excellent solvent for a variety of chemical substances, as such it is widely used in industrial processes. Water is also central to many sports and other forms of entertainment, such as swimming, pleasure boating, boat racing, surfing, sport fishing, Water is a liquid at the temperatures and pressures that are most adequate for life. Specifically, at atmospheric pressure of 1 bar, water is a liquid between the temperatures of 273.15 K and 373.15 K
16.
International Union of Pure and Applied Chemistry
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The International Union of Pure and Applied Chemistry /ˈaɪjuːpæk/ or /ˈjuːpæk/ is an international federation of National Adhering Organizations that represents chemists in individual countries. It is a member of the International Council for Science, IUPAC is registered in Zürich, Switzerland, and the administrative office, known as the IUPAC Secretariat, is in Research Triangle Park, North Carolina, United States. This administrative office is headed by IUPACs executive director, currently Lynn Soby, IUPAC was established in 1919 as the successor of the International Congress of Applied Chemistry for the advancement of chemistry. Its members, the National Adhering Organizations, can be national chemistry societies, national academies of sciences, there are fifty-four National Adhering Organizations and three Associate National Adhering Organizations. IUPACs Inter-divisional Committee on Nomenclature and Symbols is the world authority in developing standards for the naming of the chemical elements. Since its creation, IUPAC has been run by different committees with different responsibilities. These committees run different projects which include standardizing nomenclature, finding ways to bring chemistry to the world, IUPAC is best known for its works standardizing nomenclature in chemistry and other fields of science, but IUPAC has publications in many fields including chemistry, biology and physics. IUPAC is also known for standardizing the atomic weights of the elements through one of its oldest standing committees, the need for an international standard for chemistry was first addressed in 1860 by a committee headed by German scientist Friedrich August Kekulé von Stradonitz. This committee was the first international conference to create an international naming system for organic compounds, the ideas that were formulated in that conference evolved into the official IUPAC nomenclature of organic chemistry. IUPAC stands as a legacy of this meeting, making it one of the most important historical international collaborations of chemistry societies, since this time, IUPAC has been the official organization held with the responsibility of updating and maintaining official organic nomenclature. IUPAC as such was established in 1919, one notable country excluded from this early IUPAC is Germany. Germanys exclusion was a result of prejudice towards Germans by the Allied powers after World War I, Germany was finally admitted into IUPAC during 1929. However, Nazi Germany was removed from IUPAC during World War II, during World War II, IUPAC was affiliated with the Allied powers, but had little involvement during the war effort itself. After the war, East and West Germany were eventually readmitted to IUPAC, since World War II, IUPAC has been focused on standardizing nomenclature and methods in science without interruption. In 2016, IUPAC denounced the use of chlorine as a chemical weapon, the letter stated, Our organizations deplore the use of chlorine in this manner. According to the CWC, the use, stockpiling, distribution, IUPAC is governed by several committees that all have different responsibilities. Each committee is made up of members of different National Adhering Organizations from different countries, the steering committee hierarchy for IUPAC is as follows, All committees have an allotted budget to which they must adhere. Any committee may start a project, if a projects spending becomes too much for a committee to continue funding, it must take the issue to the Project Committee
17.
Miscibility
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Miscibility /mɪsᵻˈbɪlᵻti/ is the property of substances to mix in all proportions, forming a homogeneous solution. The term is most often applied to liquids, but applies also to solids, water and ethanol, for example, are miscible because they mix in all proportions. By contrast, substances are said to be if a significant proportion does not form a solution. Otherwise, the substances are considered miscible, for example, butanone is significantly soluble in water, but these two solvents are not miscible because they are not soluble in all proportions. In organic compounds, the percent of hydrocarbon chain often determines the compounds miscibility with water. For example, among the alcohols, ethanol has two atoms and is miscible with water, whereas 1-octanol with eight carbons is not. Octanols immiscibility leads it to be used as a standard for partition equilibria and this is also the case with lipids, the very long carbon chains of lipids cause them almost always to be immiscible with water. Analogous situations occur for other functional groups, acetic acid is miscible with water, whereas valeric acid is not. Immiscible metals are unable to form alloys with each other, typically, a mixture will be possible in the molten state, but upon freezing the metals separate into layers. This property allows solid precipitates to be formed by freezing a molten mixture of immiscible metals. One example of immiscibility in metals is copper and cobalt, where rapid freezing to form solid precipitates has been used to create granular GMR materials, there also exist metals that are immiscible in the liquid state. One with industrial importance is that liquid zinc and liquid silver are immiscible in liquid lead and this leads to the Parkes process, an example of liquid-liquid extraction, whereby lead containing any amount of silver is melted with zinc. The silver migrates to the zinc, which is skimmed off the top of the liquid. Substances with extremely low configurational entropy, especially polymers, are likely to be immiscible in one even in the liquid state. Miscibility of two materials is often determined optically, when the two miscible liquids are combined, the resulting liquid is clear. If the mixture is cloudy the two materials are immiscible, care must be taken with this determination. If the indices of refraction of the two materials are similar, an immiscible mixture may be clear and give an incorrect determination that the two liquids are miscible, miscibility gap Emulsion Heteroazeotrope ITIES Multiphasic liquid
18.
Ethanol
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Ethanol, also called alcohol, ethyl alcohol, and drinking alcohol, is the principal type of alcohol found in alcoholic beverages. It is a volatile, flammable, colorless liquid with a characteristic odor. Its chemical formula is C 2H 6O, which can be written also as CH 3-CH 2-OH or C 2H 5-OH, ethanol is mostly produced by the fermentation of sugars by yeasts, or by petrochemical processes. It is a psychoactive drug, causing a characteristic intoxication. It is widely used as a solvent, as fuel, and as a feedstock for synthesis of other chemicals, the eth- prefix and the qualifier ethyl in ethyl alcohol originally come from the name ethyl assigned in 1834 to the group C 2H 5- by Justus Liebig. He coined the word from the German name Aether of the compound C 2H 5-O-C 2H5, according to the Oxford English Dictionary, Ethyl is a contraction of the Ancient Greek αἰθήρ and the Greek word ύλη. The name ethanol was coined as a result of a resolution that was adopted at the International Conference on Chemical Nomenclature that was held in April 1892 in Geneva, Switzerland. The term alcohol now refers to a class of substances in chemistry nomenclature. The Oxford English Dictionary claims that it is a loan from Arabic al-kuḥl, a powdered ore of antimony used since aniquity as a cosmetic. The use of alcohol for ethanol is modern, first recorded 1753, the systematic use in chemistry dates to 1850. Ethanol is used in medical wipes and most common antibacterial hand sanitizer gels as an antiseptic, ethanol kills organisms by denaturing their proteins and dissolving their lipids and is effective against most bacteria and fungi, and many viruses. However, ethanol is ineffective against bacterial spores, ethanol may be administered as an antidote to methanol and ethylene glycol poisoning. Ethanol, often in high concentrations, is used to dissolve many water-insoluble medications, as a central nervous system depressant, ethanol is one of the most commonly consumed psychoactive drugs. The amount of ethanol in the body is typically quantified by blood alcohol content, small doses of ethanol, in general, produce euphoria and relaxation, people experiencing these symptoms tend to become talkative and less inhibited, and may exhibit poor judgment. Ethanol is commonly consumed as a drug, especially while socializing. The largest single use of ethanol is as a fuel and fuel additive. Brazil in particular relies heavily upon the use of ethanol as an engine fuel, gasoline sold in Brazil contains at least 25% anhydrous ethanol. Hydrous ethanol can be used as fuel in more than 90% of new gasoline fueled cars sold in the country, Brazilian ethanol is produced from sugar cane and noted for high carbon sequestration
19.
Silver chloride
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Silver chloride is a chemical compound with the chemical formula AgCl. This white crystalline solid is known for its low solubility in water. Upon illumination or heating, silver chloride converts to silver, which is signaled by greyish or purplish coloration to some samples, AgCl occurs naturally as a mineral chlorargyrite. Silver chloride is synthesized by combining aqueous solutions of silver nitrate. AgNO3 + NaCl ⟶ AgCl ↓ + NaNO3 The solid adopts the fcc NaCl structure, however, the crystallography depends on the condition of crystallization, primarily free silver ion concentration, as is shown on the pictures left. AgCl dissolves in solutions containing ligands such as chloride, cyanide, triphenylphosphine, thiosulfate, thiocyanate, most complexes derived from AgCl are two-, three-, and, in rare cases, four-coordinate, adopting linear, trigonal planar, and tetrahedral coordination geometries, respectively. Due to its conspicuousness it is used in titration, which gives the typical case of argentometry. The solubility product, Ksp, for AgCl in water is 1. 77×10−10 at room temperature, ions that interfere with this test are bromide and iodide, as well as a variety of ligands. For AgBr and AgI, the Ksp values are 5.2 x 10−13 and 8.3 x 10−17, silver bromide and silver iodide are also significantly more photosensitive than is AgCl. AgCl quickly darkens on exposure to light by disintegrating into elemental chlorine and this reaction is used in photography and film. The silver chloride electrode is a reference electrode in electrochemistry. Silver chlorides low solubility makes it an addition to pottery glazes for the production of Inglaze lustre. Silver chloride has been used as an antidote for poisoning, assisting in the elimination of mercury
20.
Solvation
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Solvation, also sometimes called dissolution, is the attraction and association of molecules of a solvent with molecules or ions of a solute. When a solute is soluble in a solvent, the solute’s molecules or ions will spread out. A molecule or ion of solute surrounded by solvent is known as a solvation complex, solvation is the process of reorganizing solvent and solute molecules into solvation complexes until the solute is distributed evenly within the solvent. Solvation depends on such as hydrogen bonding and van der Waals forces. Insoluble solutes prefer to maintain interactions among solute molecules rather than break apart, solvation of a solute by water is called hydration. By an IUPAC definition, solvation is an interaction of a solute with the solvent, in the solvated state, an ion in a solution is surrounded or complexed by solvent molecules. Solvated species can be described by number and the complex stability constants. The concept of the interaction can also be applied to an insoluble material, for example. Solvation is, in concept, distinct from solubility, solvation or dissolution is a kinetic process and is quantified by its rate. Solubility quantifies the dynamic equilibrium state achieved when the rate of dissolution equals the rate of precipitation, the consideration of the units makes the distinction clearer. The typical unit for dissolution rate is mol/s, the units for solubility express a concentration, mass per volume, molarity, etc. Solvation involves different types of interactions, hydrogen bonding, ion-dipole interactions. Which of these forces are at play depends on the structure and properties of the solvent. The similarity or complementary character of these properties between solvent and solute determines how well a solute can be solvated by a particular solvent, Solvent polarity is the most important factor in determining how well it solvates a particular solute. Polar solvents have molecular dipoles, meaning part of the solvent molecule has more electron density than another part of the molecule. The part with more electron density will experience a negative charge while the part with less electron density will experience a partial positive charge. Polar solvent molecules can solvate polar solutes and ions because they can orient the appropriate partially charged portion of the molecule towards the solute through electrostatic attraction and this stabilizes the system and creates a solvation shell around each particle of solute. Water is the most common and well-studied polar solvent, but others exist, such as ethanol, methanol, acetone, acetonitrile, polar solvents are often found to have a high dielectric constant, although other solvent scales are also used to classify solvent polarity
21.
Precipitation (chemistry)
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Precipitation is the creation of a solid from a solution. When the reaction occurs in a solution, the solid formed is called the precipitate. The chemical that causes the solid to form is called the precipitant, without sufficient force of gravity to bring the solid particles together, the precipitate remains in suspension. After sedimentation, especially when using a centrifuge to press it into a compact mass, Precipitation can be used as a medium. The precipitate-free liquid remaining above the solid is called the supernate or supernatant, powders derived from precipitation have also historically been known as flowers. When the solid appears in the form of fibers which have been through chemical processing. Sometimes the formation of a precipitate indicates the occurrence of a chemical reaction, if silver nitrate solution is poured into a solution of sodium chloride, a chemical reaction occurs forming a white precipitate of silver chloride. When potassium iodide solution reacts with lead nitrate solution, a precipitate of lead iodide is formed. Precipitation may occur if the concentration of a compound exceeds its solubility, Precipitation may occur rapidly from a supersaturated solution. Precipitation in solids is routinely used to synthesize nanoclusters, an important stage of the precipitation process is the onset of nucleation. The creation of a solid particle includes the formation of an interface, which requires some energy based on the relative surface energy of the solid. If this energy is not available, and no suitable nucleation surface is available, Precipitation reactions can be used for making pigments, removing salts from water in water treatment, and in classical qualitative inorganic analysis. Precipitation is also useful to isolate the products of a reaction during workup, ideally, the product of the reaction is insoluble in the reaction solvent. Thus, it precipitates as it is formed, preferably forming pure crystals, an example of this would be the synthesis of porphyrins in refluxing propionic acid. Thereafter, the precipitate may easily be separated by filtration, decanting, an example would be the synthesis of chromic tetraphenylporphyrin chloride, water is added to the DMF reaction solution, and the product precipitates. Precipitation is also useful in purifying products, crude bmim-Cl is taken up in acetonitrile, and dropped into ethyl acetate, another important application of an antisolvent is in ethanol precipitation of DNA. In metallurgy, precipitation from a solution is also a useful way to strengthen alloys. An example of a reaction, Aqueous silver nitrate is added to a solution containing potassium chloride
22.
Oxide
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An oxide /ˈɒksaɪd/ is a chemical compound that contains at least one oxygen atom and one other element in its chemical formula. Oxide itself is the dianion of oxygen, an O2– atom, Metal oxides thus typically contain an anion of oxygen in the oxidation state of −2. Most of the Earths crust consists of oxides, the result of elements being oxidized by the oxygen in air or in water. Hydrocarbon combustion affords the two principal carbon oxides, carbon monoxide and carbon dioxide, even materials considered pure elements often develop an oxide coating. For example, aluminium foil develops a skin of Al2O3 that protects the foil from further corrosion. Individual elements can form multiple oxides, each containing different amounts of the element. Certain elements can form many different oxides, such those of nitrogen, due to its electronegativity, oxygen forms stable chemical bonds with almost all elements to give the corresponding oxides. Noble metals are prized because they resist direct chemical combination with oxygen, two independent pathways for corrosion of elements are hydrolysis and oxidation by oxygen. The combination of water and oxygen is even more corrosive, virtually all elements burn in an atmosphere of oxygen, or an oxygen rich environment. In the presence of water and oxygen, some elements— sodium—react rapidly, even dangerously, in part for this reason, alkali and alkaline earth metals are not found in nature in their metallic, i. e. native, form. Caesium is so reactive with oxygen that it is used as a getter in vacuum tubes, the surface of most metals consists of oxides and hydroxides in the presence of air. A well-known example is aluminium foil, which is coated with a film of aluminium oxide that passivates the metal. The aluminium oxide layer can be built to greater thickness by the process of electrolytic anodising, though solid magnesium and aluminium react slowly with oxygen at STP—they, like most metals, burn in air, generating very high temperatures. Finely grained powders of most metals can be explosive in air. Consequently, they are used in Solid-fuel rockets. In dry oxygen, iron forms iron oxide, but the formation of the hydrated ferric oxides, Fe2O3−x2x. Free oxygen production by photosynthetic bacteria some 3.5 billion years ago precipitated iron out of solution in the oceans as Fe2O3 in the important iron ore hematite. Oxides have a range of different structures, from molecules to polymeric
23.
Aqueous solution
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An aqueous solution is a solution in which the solvent is water. It is usually shown in chemical equations by appending to the relevant chemical formula, for example, a solution of table salt, or sodium chloride, in water would be represented as Na+ + Cl−. The word aqueous means pertaining to, related to, similar to, as water is an excellent solvent and is also naturally abundant, it is a ubiquitous solvent in chemistry. Substances that are hydrophobic often do not dissolve well in water, an example of a hydrophilic substance is sodium chloride. Acids and bases are aqueous solutions, as part of their Arrhenius definitions, the ability of a substance to dissolve in water is determined by whether the substance can match or exceed the strong attractive forces that water molecules generate between themselves. If the substance lacks the ability to dissolve in water the molecules form a precipitate, reactions in aqueous solutions are usually metathesis reactions. Metathesis reactions are another term for double-displacement, that is, when a cation displaces to form a bond with the other anion. The cation bonded with the latter anion will dissociate and bond with the other anion, aqueous solutions that conduct electric current efficiently contain strong electrolytes, while ones that conduct poorly are considered to have weak electrolytes. Those strong electrolytes are substances that are ionized in water. Nonelectrolytes are substances that dissolve in water yet maintain their molecular integrity, examples include sugar, urea, glycerol, and methylsulfonylmethane. When writing the equations of reactions, it is essential to determine the precipitate. To determine the precipitate, one must consult a chart of solubility, soluble compounds are aqueous, while insoluble compounds are the precipitate. Remember that there may not always be a precipitate, when performing calculations regarding the reacting of one or more aqueous solutions, in general one must know the concentration, or molarity, of the aqueous solutions. Solution concentration is given in terms of the form of the prior to it dissolving. Metal ions in aqueous solution Solubility Dissociation Acid-base reaction theories Properties of water Zumdahl S.1997, 4th ed. Boston, Houghton Mifflin Company
24.
Iron(II) hydroxide
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Iron hydroxide or ferrous hydroxide is a inorganic compound with the formula Fe2. It is produced when iron salts, from a compound such as sulfate, are treated with hydroxide ions. Iron hydroxide is a solid, but even traces of oxygen impart a greenish tinge. The air-oxidized solid is known as green rust. Iron hydroxide is poorly soluble, or 10−14 mol/L, iron ions are easily substituted by iron ions produced by its progressive oxidation. It is also formed as an by-product of other reactions, a. o. in the synthesis of siderite. Fe2 is a layer double hydroxide, the mineralogical form of green rust is a recently discovered. All forms of green rust are more complex and variable than the iron hydroxide compound. The natural analogue of iron compound is a very rare mineral amakinite,2. Under anaerobic conditions, the hydroxide can be oxidized by the protons of water to form magnetite. The resulting products are poorly soluble, iron hydroxide has also been investigated as an agent for the removal of toxic selenate and selenite ions from water systems such as wetlands. The iron hydroxide reduces these ions to elemental selenium, which is insoluble in water, in a basic solution iron hydroxide is the electrochemically active material of the negative electrode of the nickel-iron battery
25.
PH
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In chemistry, pH is a numeric scale used to specify the acidity or basicity of an aqueous solution. It is approximately the negative of the base 10 logarithm of the concentration, measured in units of moles per liter. More precisely it is the negative of the logarithm to base 10 of the activity of the hydrogen ion, solutions with a pH less than 7 are acidic and solutions with a pH greater than 7 are basic. Pure water is neutral, at pH7, being neither an acid nor a base, contrary to popular belief, the pH value can be less than 0 or greater than 14 for very strong acids and bases respectively. The pH scale is traceable to a set of standard solutions whose pH is established by international agreement, the pH of aqueous solutions can be measured with a glass electrode and a pH meter, or an indicator. In the first papers, the notation had the H as a subscript to the p, as so. The exact meaning of the p in pH is disputed, but according to the Carlsberg Foundation and it has also been suggested that the p stands for the German Potenz, others refer to French puissance. Another suggestion is that the p stands for the Latin terms pondus hydrogenii, potentia hydrogenii and it is also suggested that Sørensen used the letters p and q simply to label the test solution and the reference solution. Currently in chemistry, the p stands for decimal cologarithm of, PH is defined as the decimal logarithm of the reciprocal of the hydrogen ion activity, aH+, in a solution. P H = − log 10 = log 10 For example and this definition was adopted because ion-selective electrodes, which are used to measure pH, respond to activity. For H+ number of electrons transferred is one and it follows that electrode potential is proportional to pH when pH is defined in terms of activity. The reference electrode may be a silver chloride electrode or a calomel electrode, the hydrogen-ion selective electrode is a standard hydrogen electrode. Reference electrode | concentrated solution of KCl || test solution | H2 | Pt Firstly, the cell is filled with a solution of hydrogen ion activity. Then the emf, EX, of the cell containing the solution of unknown pH is measured. PH = pH + E S − E X z The difference between the two measured emf values is proportional to pH and this method of calibration avoids the need to know the standard electrode potential. The proportionality constant, 1/z is ideally equal to 12.303 R T / F the Nernstian slope, to apply this process in practice, a glass electrode is used rather than the cumbersome hydrogen electrode. A combined glass electrode has a reference electrode. It is calibrated against buffer solutions of hydrogen ion activity
26.
Phase (matter)
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In the physical sciences, a phase is a region of space, throughout which all physical properties of a material are essentially uniform. Examples of physical properties include density, index of refraction, magnetization, a simple description is that a phase is a region of material that is chemically uniform, physically distinct, and mechanically separable. In a system consisting of ice and water in a jar, the ice cubes are one phase, the water is a second phase. The glass of the jar is another separate phase, the term phase is sometimes used as a synonym for state of matter, but there can be several immiscible phases of the same state of matter. Also, the phase is sometimes used to refer to a set of equilibrium states demarcated in terms of state variables such as pressure and temperature by a phase boundary on a phase diagram. Distinct phases may be described as different states of such as gas, liquid, solid. Useful mesophases between solid and liquid form other states of matter, distinct phases may also exist within a given state of matter. As shown in the diagram for iron alloys, several phases exist for both the solid and liquid states, phases may also be differentiated based on solubility as in polar or non-polar. A mixture of water and oil will separate into two phases. Water has a low solubility in oil, and oil has a low solubility in water. Solubility is the amount of a solute that can dissolve in a solvent before the solute ceases to dissolve. A mixture can separate into more than two phases and the concept of phase separation extends to solids, i. e. solids can form solid solutions or crystallize into distinct crystal phases. Metal pairs that are mutually soluble can form alloys, whereas metal pairs that are mutually insoluble cannot, as many as eight immiscible liquid phases have been observed. Mutually immiscible liquid phases are formed from water, hydrophobic solvents, perfluorocarbons, silicones, several different metals. Not all organic solvents are completely miscible, e. g. a mixture of ethylene glycol, phases do not need to macroscopically separate spontaneously. Emulsions and colloids are examples of immiscible phase pair combinations that do not physically separate, left to equilibration, many compositions will form a uniform single phase, but depending on the temperature and pressure even a single substance may separate into two or more distinct phases. Within each phase, the properties are uniform but between the two phases properties differ, water in a closed jar with an air space over it forms a two phase system. Most of the water is in the phase, where it is held by the mutual attraction of water molecules
27.
Aragonite
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Aragonite is a carbonate mineral, one of the two most common, naturally occurring, crystal forms of calcium carbonate, CaCO3. It is formed by biological and physical processes, including precipitation from marine, aragonites crystal lattice differs from that of calcite, resulting in a different crystal shape, an orthorhombic crystal system with acicular crystal. Repeated twinning results in pseudo-hexagonal forms, Aragonite may be columnar or fibrous, occasionally in branching stalactitic forms called flos-ferri from their association with the ores at the Carinthian iron mines. The type location for aragonite is Molina de Aragón,25 km from Aragon for which it was named in 1797, an aragonite cave, the Ochtinská Aragonite Cave, is situated in Slovakia. In the US, aragonite in the form of stalactites and cave flowers is known from Carlsbad Caverns, massive deposits of oolitic aragonite sand are found on the seabed in the Bahamas. Aragonite is the high pressure polymorph of calcium carbonate, as such, it occurs in high pressure metamorphic rocks such as those formed at subduction zones. Aragonite forms naturally in almost all mollusk shells, and as the calcareous endoskeleton of warm-, because the mineral deposition in mollusk shells is strongly biologically controlled, some crystal forms are distinctively different from those of inorganic aragonite. In some mollusks, the shell is aragonite, in others. The nacreous layer of the fossil shells of some extinct ammonites forms an iridescent material called ammolite. Aragonite also forms in the ocean and in caves as inorganic precipitates called marine cements and speleothems, Aragonite is not uncommon in serpentinites where high Mg in pore solutions apparently inhibits calcite growth and promotes aragonite precipitation. Aragonite is metastable at the low pressures near the Earths surface and is commonly replaced by calcite in fossils. Aragonite older than the Carboniferous is essentially unknown and it can also be synthesized by adding a calcium chloride solution to a sodium carbonate solution at temperatures above 60 °C or in water-ethanol mixtures at ambient temperatures. Aragonite is thermodynamically unstable at temperature and pressure, and tends to alter to calcite on scales of 107 to 108 years. The mineral vaterite, also known as μ-CaCO3, is another phase of calcium carbonate that is metastable at ambient conditions typical of Earths surface, in aquaria, aragonite is considered essential for the replication of reef conditions. Aragonite provides the necessary for much sea life and also keeps the pH of the water close to its natural level. Aragonite has been tested for the removal of pollutants like zinc, cobalt. Aragonite sea Ikaite, CaCO3·6H2O Monohydrocalcite, CaCO3·H2O Nacre, otherwise known as Mother-of-Pearl Oolitic aragonite sand The Ochtinska aragonite cave Kosovo Caves Aragonite Formations
28.
Calcite
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Calcite is a carbonate mineral and the most stable polymorph of calcium carbonate. The Mohs scale of hardness, based on scratch hardness comparison. Other polymorphs of calcium carbonate are the minerals aragonite and vaterite, aragonite will change to calcite at 380–470 °C, and vaterite is even less stable. Calcite is derived from the German Calcit, a term coined in the 19th century from the Latin word for lime and it is thus etymologically related to chalk. Calcite crystals are trigonal-rhombohedral, though actual calcite rhombohedra are rare as natural crystals, however, they show a remarkable variety of habits including acute to obtuse rhombohedra, tabular forms, prisms, or various scalenohedra. Calcite exhibits several twinning types adding to the variety of observed forms and it may occur as fibrous, granular, lamellar, or compact. Cleavage is usually in three directions parallel to the rhombohedron form and its fracture is conchoidal, but difficult to obtain. It has a defining Mohs hardness of 3, a gravity of 2.71. Color is white or none, though shades of gray, red, orange, yellow, green, blue, violet, brown, calcite is transparent to opaque and may occasionally show phosphorescence or fluorescence. A transparent variety called Iceland spar is used for optical purposes, acute scalenohedral crystals are sometimes referred to as dogtooth spar while the rhombohedral form is sometimes referred to as nailhead spar. Single calcite crystals display an optical property called birefringence and this strong birefringence causes objects viewed through a clear piece of calcite to appear doubled. The birefringent effect was first described by the Danish scientist Rasmus Bartholin in 1669, at a wavelength of ~590 nm calcite has ordinary and extraordinary refractive indices of 1.658 and 1.486, respectively. Between 190 and 1700 nm, the refractive index varies roughly between 1.9 and 1.5, while the extraordinary refractive index varies between 1.6 and 1.4. Calcite, like most carbonates, will dissolve with most forms of acid, calcite can be either dissolved by groundwater or precipitated by groundwater, depending on several factors including the water temperature, pH, and dissolved ion concentrations. Although calcite is fairly insoluble in water, acidity can cause dissolution of calcite. Ambient carbon dioxide, due to its acidity, has a slight solubilizing effect on calcite, calcite exhibits an unusual characteristic called retrograde solubility in which it becomes less soluble in water as the temperature increases. When conditions are right for precipitation, calcite forms mineral coatings that cement the existing rock grains together or it can fill fractures. On a landscape scale, continued dissolution of calcium carbonate-rich rocks can lead to the expansion and eventual collapse of cave systems, high-grade optical calcite was used in World War II for gun sights, specifically in bomb sights and anti-aircraft weaponry
29.
Polymorphism (materials science)
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In materials science, polymorphism is the ability of a solid material to exist in more than one form or crystal structure. Polymorphism can potentially be found in any crystalline material including polymers, minerals, and metals, and is related to allotropy, the complete morphology of a material is described by polymorphism and other variables such as crystal habit, amorphous fraction or crystallographic defects. Polymorphism is relevant to the fields of pharmaceuticals, agrochemicals, pigments, dyestuffs, foods, when polymorphism exists as a result of a difference in crystal packing, it is called packing polymorphism. Polymorphism can also result from the existence of different conformers of the molecule in conformational polymorphism. In pseudopolymorphism the different crystal types are the result of hydration or solvation and this is more correctly referred to as solvomorphism as different solvates have different chemical formulae. An example of an organic polymorph is glycine, which is able to form monoclinic, silica is known to form many polymorphs, the most important of which are, α-quartz, β-quartz, tridymite, cristobalite, coesite, and stishovite. A classical example is the pair of minerals, calcite and aragonite, an analogous phenomenon for amorphous materials is polyamorphism, when a substance can take on several different amorphous modifications. In terms of thermodynamics, there are two types of polymorphic behaviour, for an enantiotropic system, a plot of the free energy against temperature shows a crossing point threshold before the various melting points. It may also be possible to revert interchangeably between the two polymorphs by heating or cooling, or through contact with a lower energy polymorph. Present-day analysis identifies three polymorphs for benzamide, the least stable one, formed by flash cooling is the orthorhombic form II and this type is followed by the monoclinic form III. The most stable form is monoclinic form I, the hydrogen bonding mechanisms are the same for all three phases, however, they differ strongly in their pi-pi interactions. Polymorphs have different stabilities and may spontaneously convert from a form to the stable form at a particular temperature. Most polymorphs of organic molecules only differ by a few kJ/mol in lattice energy, approximately 50% of known polymorph pairs differ by less than 2 kJ/mol and stability differences of more than 10 kJ/mol are rare. They also exhibit different melting points, solubilities, X-ray crystal, various conditions in the crystallisation process is the main reason responsible for the development of different polymorphic forms. In 2006 a new form of maleic acid was discovered 124 years after the first crystal form was studied. Maleic acid is a manufactured on a very large scale in the chemical industry and is a salt forming component in medicine. The new crystal type is produced when a co-crystal of caffeine and maleic acid is dissolved in chloroform, whereas form I has monoclinic space group P21/c, the new form has space group Pc. 1,3, 5-Trinitrobenzene is more than 125 years old and was used as an explosive before the arrival of the safer 2,4, only one crystal form of 1,3, 5-trinitrobenzene was known in the space group Pbca
30.
Calcium carbonate
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Calcium carbonate is a chemical compound with the formula CaCO3. It is a substance found in rocks as the minerals calcite and aragonite and is the main component of pearls and the shells of marine organisms, snails. Calcium carbonate is the ingredient in agricultural lime and is created when calcium ions in hard water react with carbonate ions to create limescale. It is medicinally used as a supplement or as an antacid. Calcium carbonate shares the properties of other carbonates. CaCO3 + CO2 + H2O → Ca2 This reaction is important in the erosion of rock, forming caverns. An unusual form of calcium carbonate is the hexahydrate, ikaite, ikaite is stable only below 6 °C. The vast majority of calcium used in industry is extracted by mining or quarrying. Pure calcium carbonate, can be produced from a quarried source. Alternatively, calcium carbonate is prepared from calcium oxide, other forms can be prepared, the denser, orthorhombic λ-CaCO3 and μ-CaCO3, occurring as the mineral vaterite. The aragonite form can be prepared by precipitation at temperatures above 85 °C, calcite contains calcium atoms coordinated by 6 oxygen atoms, in aragonite they are coordinated by 9 oxygen atoms. The vaterite structure is not fully understood, magnesium carbonate MgCO3 has the calcite structure, whereas strontium and barium carbonate adopt the aragonite structure, reflecting their larger ionic radii. Calcite, aragonite and vaterite are pure calcium carbonate minerals, industrially important source rocks which are predominantly calcium carbonate include limestone, chalk, marble and travertine. Eggshells, snail shells and most seashells are predominantly calcium carbonate, oyster shells have enjoyed recent recognition as a source of dietary calcium, but are also a practical industrial source. While not practical as a source, dark green vegetables such as broccoli. Beyond Earth, strong evidence suggests the presence of Calcium carbonate on Mars, signs of Calcium Carbonate have been detected at more than one location. This provides some evidence for the past presence of liquid water, Carbonate is found frequently in geologic settings and constitute an enormous carbon reservoir. Calcium carbonate occurs as aragonite, calcite and dolomite, the carbonate minerals form the rock types, limestone, chalk, marble, travertine, tufa, and others
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Chemical formula
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These are limited to a single typographic line of symbols, which may include subscripts and superscripts. A chemical formula is not a name, and it contains no words. Although a chemical formula may imply certain simple chemical structures, it is not the same as a full chemical structural formula. Chemical formulas can fully specify the structure of only the simplest of molecules and chemical substances, the simplest types of chemical formulas are called empirical formulas, which use letters and numbers indicating the numerical proportions of atoms of each type. Molecular formulas indicate the numbers of each type of atom in a molecule. For example, the formula for glucose is CH2O, while its molecular formula is C6H12O6. This is possible if the relevant bonding is easy to show in one dimension, an example is the condensed molecular/chemical formula for ethanol, which is CH3-CH2-OH or CH3CH2OH. For reasons of structural complexity, there is no condensed chemical formula that specifies glucose, chemical formulas may be used in chemical equations to describe chemical reactions and other chemical transformations, such as the dissolving of ionic compounds into solution. A chemical formula identifies each constituent element by its chemical symbol, in empirical formulas, these proportions begin with a key element and then assign numbers of atoms of the other elements in the compound, as ratios to the key element. For molecular compounds, these numbers can all be expressed as whole numbers. For example, the formula of ethanol may be written C2H6O because the molecules of ethanol all contain two carbon atoms, six hydrogen atoms, and one oxygen atom. Some types of compounds, however, cannot be written with entirely whole-number empirical formulas. An example is boron carbide, whose formula of CBn is a variable non-whole number ratio with n ranging from over 4 to more than 6.5. When the chemical compound of the consists of simple molecules. These types of formulas are known as molecular formulas and condensed formulas. A molecular formula enumerates the number of atoms to reflect those in the molecule, so that the formula for glucose is C6H12O6 rather than the glucose empirical formula. However, except for very simple substances, molecular chemical formulas lack needed structural information, for simple molecules, a condensed formula is a type of chemical formula that may fully imply a correct structural formula. For example, ethanol may be represented by the chemical formula CH3CH2OH
32.
Intermolecular force
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Intermolecular forces are weak relative to intramolecular forces – the forces which hold a molecule together. For example, the covalent bond, involving sharing electron pairs between atoms, is stronger than the forces present between neighboring molecules. Both sets of forces are essential parts of force fields used in molecular mechanics. The investigation of intermolecular forces starts from macroscopic observations which indicate the existence and these observations include non-ideal-gas thermodynamic behavior reflected by virial coefficients, vapor pressure, viscosity, superficial tension, and absorption data. The first reference to the nature of microscopic forces is found in Alexis Clairauts work Theorie de la Figure de la Terre, other scientists who have contributed to the investigation of microscopic forces include, Laplace, Gauss, Maxwell and Boltzmann. The link to microscopic aspects is given by virial coefficients and Lennard-Jones potentials, dipole-dipole interactions are electrostatic interactions between permanent dipoles in molecules. These interactions tend to align the molecules to increase attraction, an example of a dipole-dipole interaction can be seen in hydrogen chloride, the positive end of a polar molecule will attract the negative end of the other molecule and influence its position. Polar molecules have a net attraction between them, examples of polar molecules include hydrogen chloride and chloroform. H δ + − Cl δ − ⋯ H δ + − Cl δ − Often molecules contain dipolar groups and this occurs if there is symmetry within the molecule that causes the dipoles to cancel each other out. This occurs in such as tetrachloromethane and carbon dioxide. The dipole-dipole interaction between two atoms is usually zero, since atoms rarely carry a permanent dipole. Ion-dipole and ion-induced dipole forces are similar to dipole-dipole and induced-dipole interactions but involve ions, ion-dipole and ion-induced dipole forces are stronger than dipole-dipole interactions because the charge of any ion is much greater than the charge of a dipole moment. Ion-dipole bonding is stronger than hydrogen bonding, an ion-dipole force consists of an ion and a polar molecule interacting. They align so that the positive and negative groups are next to one another, an ion-induced dipole force consists of an ion and a non-polar molecule interacting. Like a dipole-induced dipole force, the charge of the ion causes distortion of the cloud on the non-polar molecule. A hydrogen bond is the attraction between the pair of an electronegative atom and a hydrogen atom that is bonded to either nitrogen, oxygen. The hydrogen bond is described as a strong electrostatic dipole-dipole interaction. Intermolecular hydrogen bonding is responsible for the boiling point of water compared to the other group 16 hydrides
33.
Entropy
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In statistical thermodynamics, entropy is a measure of the number of microscopic configurations Ω that a thermodynamic system can have when in a state as specified by some macroscopic variables. Formally, S = k B ln Ω, for example, gas in a container with known volume, pressure, and temperature could have an enormous number of possible configurations of the collection of individual gas molecules. Each instantaneous configuration of the gas may be regarded as random, Entropy may be understood as a measure of disorder within a macroscopic system. The second law of thermodynamics states that an isolated systems entropy never decreases, such systems spontaneously evolve towards thermodynamic equilibrium, the state with maximum entropy. Non-isolated systems may lose entropy, provided their environments entropy increases by at least that amount, since entropy is a function of the state of the system, a change in entropy of a system is determined by its initial and final states. This applies whether the process is reversible or irreversible, however, irreversible processes increase the combined entropy of the system and its environment. The above definition is called the macroscopic definition of entropy because it can be used without regard to any microscopic description of the contents of a system. The concept of entropy has found to be generally useful and has several other formulations. Entropy was discovered when it was noticed to be a quantity that behaves as a function of state and it has the dimension of energy divided by temperature, which has a unit of joules per kelvin in the International System of Units. But the entropy of a substance is usually given as an intensive property—either entropy per unit mass or entropy per unit amount of substance. In statistical mechanics this reflects that the state of a system is generally non-degenerate. Understanding the role of entropy in various processes requires an understanding of how. It is often said that entropy is an expression of the disorder, or randomness of a system, the second law is now often seen as an expression of the fundamental postulate of statistical mechanics through the modern definition of entropy. In other words, in any natural process there exists an inherent tendency towards the dissipation of useful energy and he made the analogy with that of how water falls in a water wheel. This was an insight into the second law of thermodynamics. g. Clausius described entropy as the transformation-content, i. e. dissipative energy use and this was in contrast to earlier views, based on the theories of Isaac Newton, that heat was an indestructible particle that had mass. Later, scientists such as Ludwig Boltzmann, Josiah Willard Gibbs, henceforth, the essential problem in statistical thermodynamics, i. e. according to Erwin Schrödinger, has been to determine the distribution of a given amount of energy E over N identical systems. Carathéodory linked entropy with a definition of irreversibility, in terms of trajectories
34.
Complex (chemistry)
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Many metal-containing compounds, especially those of transition metals, are coordination complexes. A coordination complex whose centre is an atom is called a metal complex. Coordination complexes are so pervasive that their structures and reactions are described in many ways, the atom within a ligand that is bonded to the central metal atom or ion is called the donor atom. In a typical complex, an ion is bonded to several donor atoms. A polydentate ligand is a molecule or ion that bonds to the atom through several of the ligands atoms. These complexes are called chelate complexes, the formation of complexes is called chelation, complexation. The central atom or ion, together with all ligands comprise the coordination sphere, the central atoms or ion and the donor atoms comprise the first coordination sphere. Coordination refers to the covalent bonds between the ligands and the central atom. Originally, a complex implied a reversible association of molecules, atoms, as applied to coordination chemistry, this meaning has evolved. Some metal complexes are formed virtually irreversibly and many are bound together by bonds that are quite strong, the number of donor atoms attached to the central atom or ion is called the coordination number. The most common coordination numbers are 2,4 and especially 6, for example, the cobalt hexahydrate ion or the hexaaquacobalt ion 2+, is a hydrated-complex ion that consists of six water molecules attached to a metal ion Co. The oxidation state and the number reflect the number of bonds formed between the metal ion and the ligands in the complex ion. However the coordination number of Pt2+ 2is 4 since it has two ligands, which contain four donor atoms in total. Coordination complexes have been known since the beginning of modern chemistry, early well-known coordination complexes include dyes such as Prussian blue. Their properties were first well understood in the late 1800s, following the 1869 work of Christian Wilhelm Blomstrand, Blomstrand developed what has come to be known as the complex ion chain theory. The theory claimed that the reason coordination complexes form is because in solution and he compared this effect to the way that various carbohydrate chains form. Following this theory, Danish scientist Sophus Mads Jorgensen made improvements to it and it was not until 1893 that the most widely accepted version of the theory today was published by Alfred Werner. Werner’s work included two important changes to the Blomstrand theory, the first was that Werner described the two different ion possibilities in terms of location in the coordination sphere
35.
Ligand
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In coordination chemistry, a ligand is an ion or molecule that binds to a central metal atom to form a coordination complex. The bonding with the metal generally involves formal donation of one or more of the electron pairs. The nature of bonding can range from covalent to ionic. Furthermore, the bond order can range from one to three. Ligands are viewed as Lewis bases, although rare cases are known to involve Lewis acidic ligand, metals and metalloids are bound to ligands in virtually all circumstances, although gaseous naked metal ions can be generated in high vacuum. Ligands in a complex dictate the reactivity of the atom, including ligand substitution rates, the reactivity of the ligands themselves. Ligand selection is a consideration in many practical areas, including bioinorganic and medicinal chemistry, homogeneous catalysis. Ligands are classified in many ways, including, charge, size, the identity of the atom. The size of a ligand is indicated by its cone angle, the composition of coordination complexes have been known since the early 1800s, such as Prussian blue and copper vitriol. The key breakthrough occurred when Alfred Werner reconciled formulas and isomers and he showed, among other things, that the formulas of many cobalt and chromium compounds can be understood if the metal has six ligands in an octahedral geometry. The first to use the term ligand were Alfred Stock and Carl Somiesky, the theory allows one to understand the difference between coordinated and ionic chloride in the cobalt ammine chlorides and to explain many of the previously inexplicable isomers. He resolved the first coordination complex called hexol into optical isomers, in general, ligands are viewed as electron donors and the metals as electron acceptors. This is because the ligand and central metal are bonded to one another, bonding is often described using the formalisms of molecular orbital theory. The HOMO can be mainly of ligands or metal character, ligands and metal ions can be ordered in many ways, one ranking system focuses on ligand hardness. Metal ions preferentially bind certain ligands, in general, soft metal ions prefer weak field ligands, whereas hard metal ions prefer strong field ligands. According to the orbital theory, the HOMO of the ligand should have an energy that overlaps with the LUMO of the metal preferential. Metal ions bound to strong-field ligands follow the Aufbau principle, whereas complexes bound to weak-field ligands follow Hunds rule. Binding of the metal with the results in a set of molecular orbitals, where the metal can be identified with a new HOMO and LUMO
36.
Mineral redox buffer
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In geology, a redox buffer is an assemblage of minerals or compounds that constrains oxygen fugacity as a function of temperature. Knowledge of the conditions at which a rock forms and evolves can be important for interpreting the rock history. Iron, sulfur, and manganese are three of the relatively abundant elements in the Earths crust that occur in more than one oxidation state, for instance, iron, the fourth most abundant element in the crust, exists as native iron, ferrous iron, and ferric iron. The redox state of a rock affects the relative proportions of the states of these elements. If a rock contains pure minerals that constitute a redox buffer, for other rocks with suitable minerals, oxygen fugacities can be calculated, and the redox conditions can be obtained by comparison to the fugacity-temperature diagram. Common redox buffers Redox buffers were developed in part to control oxygen fugacities in laboratory experiments to investigate mineral stabilities, each of the curves plotted in the fugacity-temperature diagram is for an oxidation reaction occurring in a buffer. These redox buffers are listed here in order of decreasing oxygen fugacity at a given temperature—in other words, as long as all the pure minerals are present in a buffer assemblage, the oxidizing conditions are fixed on the curve for that buffer. Pressure has only a minor influence on these buffer curves for conditions in the Earths crust, within a rock of a given chemical composition, iron enters minerals based on the bulk chemical composition and the mineral phases which are stable at that temperature and pressure. For instance, at conditions more oxidizing than the MH buffer, at least much of the iron is likely to be present as Fe3+. Iron may only enter minerals such as if it is present as Fe2+. Elements in olivine such as magnesium, however, stabilize olivine containing Fe2+ to conditions more oxidizing than those required for fayalite stability, solid solution between magnetite and the titanium-bearing endmember, ulvospinel, enlarges the stability field of magnetite. Likewise, at conditions more reducing than the IW buffer, minerals such as pyroxene can still contain Fe3+, the redox buffers therefore are only approximate guides to the proportions of Fe2+ and Fe3+ in minerals and rocks. Terrestrial igneous rocks commonly record crystallization at oxygen fugacities more oxidizing than the WM buffer and their oxidizing conditions thus are not far from those of the FMQ redox buffer. Nonetheless, there are differences that correlate with tectonic setting. Igneous rock emplaced and erupted in island arcs typically record oxygen fugacities 1 or more log units more oxidizing than those of the NiNiO buffer. In contrast, basalt and gabbro in non-arc settings typically record oxygen fugacities from about those of the FMQ buffer to a log unit or so more reducing than that buffer, oxidizing conditions are common in some environments of deposition and diagenesis of sedimentary rocks. Other sedimentary environments, such as the environments for formation of black shale, are relatively reducing, oxygen fugacities during metamorphism extend to higher values than those in magmatic environments, because of the more oxidizing compositions inherited from some sedimentary rocks. Nearly pure hematite is present in some metamorphosed banded iron formations, in contrast, native nickel-iron is present in some serpentinites
37.
Specific surface area
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Specific surface area is a property of solids defined as the total surface area of a material per unit of mass, or solid or bulk volume. It is a derived scientific value that can be used to determine the type and it has a particular importance for adsorption, heterogeneous catalysis, and reactions on surfaces. Values obtained for specific surface area depend on the method of measurement, in adsorption based methods, the size of the adsorbate molecule, the exposed crystallographic planes at the surface and measurement temperature all affect the obtained specific surface area. The SSA can be calculated from a particle size distribution. This method, however, fails to account for surface associated with the texture of the particles. The SSA can be measured by using the BET isotherm. This has the advantage of measuring the surface of fine structures, however, the results can differ markedly depending on the substance adsorbed. This depends upon a relationship between the surface area and the resistance to gas-flow of a porous bed of powder. The method is simple and quick, and yields a result that often correlates well with the chemical reactivity of a powder, however, it fails to measure much of the deep surface texture
38.
Ostwald ripening
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Dissolution of small crystals or sol particles and the redeposition of the dissolved species on the surfaces of larger crystals or sol particles was first described by Wilhelm Ostwald in 1896. Ostwald ripening is generally found in water-in-oil emulsions, while flocculation is found in oil-in-water emulsions and this thermodynamically-driven spontaneous process occurs because larger particles are more energetically favored than smaller particles. This stems from the fact that molecules on the surface of a particle are energetically less stable than the ones in the interior. Large particles are more favorable since, continuing with this example, more atoms are bonded to 6 neighbors. As the system tries to lower its energy, molecules on the surface of a small particle will tend to detach from the particle, as per the Kelvin equation. When all small particles do this, it increases the concentration of molecules in solution. When the free molecules in solution are supersaturated, the molecules have a tendency to condense on the surface of larger particles. Therefore, all smaller particles shrink, while larger particles grow, as time tends to infinity, the entire population of particles becomes one large spherical particle to minimize the total surface area. The history of progress in quantitatively modeling Ostwald Ripening is long. In 1958, Lifshitz and Slyozov performed an investigation of Ostwald ripening in the case where diffusion of material is the slowest process. They began by stating how a single particle grows in a solution and this equation describes where the boundary is between small, shrinking particles and large, growing particles. In contexts where the value of ⟨R⟩0 is irrelevant, an approach that respects the meanings of all terms is to take the time derivative of the equation to eliminate ⟨R⟩0. Another such approach is to change the ⟨R⟩0 to ⟨R⟩i with the initial time i having a positive value, also contained in the Lifshitz and Slyozov derivation is an equation for the size distribution function f of particles. For convenience, the radius of particles is divided by the radius to form a new variable. F =49 ρ273113 exp , ρ <1, although his calculations and approach were different, Wagner came to the same conclusions as Lifshitz and Slyozov for slow-diffusion systems. This duplicate derivation went unnoticed for years because the two papers were published on opposite sides of the Iron Curtain in 1961. It was not until 1975 that Kahlweit addressed the fact that the theories were identical, many experiments and simulations have shown LSW theory to be robust and accurate. Even some systems that undergo spinodal decomposition have been shown to quantitatively obey LSW theory after initial stages of growth
39.
Critical point (thermodynamics)
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In thermodynamics, a critical point is the end point of a phase equilibrium curve. The most prominent example is the critical point, the end point of the pressure-temperature curve that designates conditions under which a liquid. At the critical point, defined by a critical temperature Tc, other examples include the liquid–liquid critical points in mixtures. For simplicity and clarity, the notion of critical point is best introduced by discussing a specific example. This was the first critical point to be discovered, and it is still the best known, the figure to the right shows the schematic PT diagram of a pure substance. The commonly known phases solid, liquid and vapor are separated by phase boundaries, at the triple point, all three phases can coexist. However, the liquid-vapor boundary terminates in an endpoint at some critical temperature Tc, in water, the critical point occurs at around 647 K and 22.064 MPa. In the vicinity of the point, the physical properties of the liquid. For instance, liquid water under normal conditions is nearly incompressible, has a low thermal expansion coefficient, has a dielectric constant. At the critical point, only one phase exists, the heat of vaporization is zero. There is a point in the constant-temperature line on a PV diagram. This means that at the point, T = T =0 Above the critical point there exists a state of matter that is continuously connected with both the liquid and the gaseous state. The existence of a point was first discovered by Charles Cagniard de la Tour in 1822 and named by Dmitri Mendeleev in 1860. Cagniard showed that CO2 could be liquefied at 31 °C at a pressure of 73 atm, but not at a slightly higher temperature, even under pressures as high as 3,000 atm. Solving the above condition T =0 for the van der Waals equation, however, the van der Waals equation, based on a mean field theory, does not hold near the critical point. In particular, it predicts wrong scaling laws, the principle of corresponding states indicates that substances at equal reduced pressures and temperatures have equal reduced volumes. This relationship is true for many substances, but becomes increasingly inaccurate for large values of pr. The liquid–liquid critical point of a solution, which occurs at the solution temperature
40.
Relative permittivity
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The relative permittivity of a material is its permittivity expressed as a ratio relative to the permittivity of vacuum. Permittivity is a property that affects the Coulomb force between two point charges in the material. Relative permittivity is the factor by which the field between the charges is decreased relative to vacuum. Likewise, relative permittivity is the ratio of the capacitance of a capacitor using that material as a dielectric, relative permittivity is also commonly known as dielectric constant, a term deprecated in physics and engineering as well as in chemistry. Relative permittivity is typically denoted as εr and is defined as ε r = ε ε0, where ε is the complex frequency-dependent absolute permittivity of the material, and ε0 is the vacuum permittivity. Relative permittivity is a number that is in general complex-valued, its real and imaginary parts are denoted as. The relative permittivity of a medium is related to its electric susceptibility, χe, in anisotropic media the relative permittivity is a second rank tensor. The relative permittivity of a material for a frequency of zero is known as its relative permittivity. The historical term for the relative permittivity is dielectric constant and it is still commonly used, but has been deprecated by standards organizations, because of its ambiguity, as some older authors used it for the absolute permittivity ε. The permittivity may be quoted either as a property or as a frequency-dependent variant. It has also used to refer to only the real component εr of the complex-valued relative permittivity. In the causal theory of waves, permittivity is a complex quantity, the imaginary part corresponds to a phase shift of the polarization P relative to E and leads to the attenuation of electromagnetic waves passing through the medium. By definition, the relative permittivity of vacuum is equal to 1. The relative static permittivity, εr, can be measured for static electric fields as follows, first the capacitance of a test capacitor, then, using the same capacitor and distance between its plates, the capacitance Cx with a dielectric between the plates is measured. The relative dielectric constant can be calculated as ε r = C x C0. For time-variant electromagnetic fields, this quantity becomes frequency-dependent, an indirect technique to calculate εr is conversion of radio frequency S-parameter measurement results. A description of frequently used S-parameter conversions for determination of the frequency-dependent εr of dielectrics can be found in this bibliographic source, alternatively, resonance based effects may be employed at fixed frequencies. The relative permittivity is a piece of information when designing capacitors
41.
Salt (chemistry)
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In chemistry, a salt is an ionic compound that results from the neutralization reaction of an acid and a base. Salts are composed of related numbers of cations and anions so that the product is electrically neutral and these component ions can be inorganic, such as chloride, or organic, such as acetate, and can be monatomic, such as fluoride, or polyatomic, such as sulfate. There are several varieties of salts, salts that hydrolyze to produce hydroxide ions when dissolved in water are alkali salts, whilst those that hydrolyze to produce hydronium ions in water are acidic salts. Neutral salts are those salts that are neither acidic nor basic, zwitterions contain an anionic centre and a cationic centre in the same molecule, but are not considered to be salts. Examples of zwitterions include amino acids, many metabolites, peptides, usually, non-dissolved salts at standard conditions for temperature and pressure are solid, but there are exceptions. Molten salts and solutions containing dissolved salts are called electrolytes, as they are able to conduct electricity. As observed in the cytoplasm of cells, in blood, urine, plant saps and mineral waters, therefore, their salt content is given for the respective ions. Salts can appear to be clear and transparent, opaque, and even metallic, in many cases, the apparent opacity or transparency are only related to the difference in size of the individual monocrystals. Since light reflects from the boundaries, larger crystals tend to be transparent. The color of the salt is due to the electronic structure in the d-orbitals of transition elements or in the conjugated organic dye framework. Different salts can elicit all five basic tastes, e. g. salty, sweet, sour, bitter, and umami or savory. Salts of strong acids and strong bases are non-volatile and odorless and that slow, partial decomposition is usually accelerated by the presence of water, since hydrolysis is the other half of the reversible reaction equation of formation of weak salts. Many ionic compounds can be dissolved in water or other similar solvents, the exact combination of ions involved makes each compound have a unique solubility in any solvent. The solubility is dependent on how well each ion interacts with the solvent, for example, all salts of sodium, potassium and ammonium are soluble in water, as are all nitrates and many sulfates – barium sulfate, calcium sulfate and lead sulfate are examples of exceptions. However, ions that bind tightly to each other and form highly stable lattices are less soluble, for example, most carbonate salts are not soluble in water, such as lead carbonate and barium carbonate. Some soluble carbonate salts are, sodium carbonate, potassium carbonate, solid salts do not conduct electricity. Moreover, solutions of salts also conduct electricity, the name of a salt starts with the name of the cation followed by the name of the anion. Salts are often referred to only by the name of the cation or by the name of the anion. g
42.
Celsius
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Celsius, also known as centigrade, is a metric scale and unit of measurement for temperature. As an SI derived unit, it is used by most countries in the world and it is named after the Swedish astronomer Anders Celsius, who developed a similar temperature scale. The degree Celsius can refer to a temperature on the Celsius scale as well as a unit to indicate a temperature interval. Before being renamed to honour Anders Celsius in 1948, the unit was called centigrade, from the Latin centum, which means 100, and gradus, which means steps. The scale is based on 0° for the point of water. This scale is widely taught in schools today, by international agreement the unit degree Celsius and the Celsius scale are currently defined by two different temperatures, absolute zero, and the triple point of VSMOW. This definition also precisely relates the Celsius scale to the Kelvin scale, absolute zero, the lowest temperature possible, is defined as being precisely 0 K and −273.15 °C. The temperature of the point of water is defined as precisely 273.16 K at 611.657 pascals pressure. This definition fixes the magnitude of both the degree Celsius and the kelvin as precisely 1 part in 273.16 of the difference between absolute zero and the point of water. Thus, it sets the magnitude of one degree Celsius and that of one kelvin as exactly the same, additionally, it establishes the difference between the two scales null points as being precisely 273.15 degrees. In his paper Observations of two persistent degrees on a thermometer, he recounted his experiments showing that the point of ice is essentially unaffected by pressure. He also determined with precision how the boiling point of water varied as a function of atmospheric pressure. He proposed that the point of his temperature scale, being the boiling point. This pressure is known as one standard atmosphere, the BIPMs 10th General Conference on Weights and Measures later defined one standard atmosphere to equal precisely 1013250dynes per square centimetre. On 19 May 1743 he published the design of a mercury thermometer, in 1744, coincident with the death of Anders Celsius, the Swedish botanist Carolus Linnaeus reversed Celsiuss scale. In it, Linnaeus recounted the temperatures inside the orangery at the University of Uppsala Botanical Garden, since the 19th century, the scientific and thermometry communities worldwide referred to this scale as the centigrade scale. Temperatures on the scale were often reported simply as degrees or. More properly, what was defined as centigrade then would now be hectograde.2 gradians, for scientific use, Celsius is the term usually used, with centigrade otherwise continuing to be in common but decreasing use, especially in informal contexts in English-speaking countries
. PMID 23795551.
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