In crystallography, crystal structure is a description of the ordered arrangement of atoms, ions or molecules in a crystalline material. Ordered structures occur from the intrinsic nature of the constituent particles to form symmetric patterns that repeat along the principal directions of three-dimensional space in matter; the smallest group of particles in the material that constitutes this repeating pattern is the unit cell of the structure. The unit cell reflects the symmetry and structure of the entire crystal, built up by repetitive translation of the unit cell along its principal axes; the translation vectors define the nodes of the Bravais lattice. The lengths of the principal axes, or edges, of the unit cell and the angles between them are the lattice constants called lattice parameters or cell parameters; the symmetry properties of the crystal are described by the concept of space groups. All possible symmetric arrangements of particles in three-dimensional space may be described by the 230 space groups.
The crystal structure and symmetry play a critical role in determining many physical properties, such as cleavage, electronic band structure, optical transparency. Crystal structure is described in terms of the geometry of arrangement of particles in the unit cell; the unit cell is defined as the smallest repeating unit having the full symmetry of the crystal structure. The geometry of the unit cell is defined as a parallelepiped, providing six lattice parameters taken as the lengths of the cell edges and the angles between them; the positions of particles inside the unit cell are described by the fractional coordinates along the cell edges, measured from a reference point. It is only necessary to report the coordinates of a smallest asymmetric subset of particles; this group of particles may be chosen so that it occupies the smallest physical space, which means that not all particles need to be physically located inside the boundaries given by the lattice parameters. All other particles of the unit cell are generated by the symmetry operations that characterize the symmetry of the unit cell.
The collection of symmetry operations of the unit cell is expressed formally as the space group of the crystal structure. Vectors and planes in a crystal lattice are described by the three-value Miller index notation; this syntax uses the indices ℓ, m, n as directional orthogonal parameters, which are separated by 90°. By definition, the syntax denotes a plane that intercepts the three points a1/ℓ, a2/m, a3/n, or some multiple thereof; that is, the Miller indices are proportional to the inverses of the intercepts of the plane with the unit cell. If one or more of the indices is zero, it means. A plane containing a coordinate axis is translated so that it no longer contains that axis before its Miller indices are determined; the Miller indices for a plane are integers with no common factors. Negative indices are indicated with horizontal bars, as in. In an orthogonal coordinate system for a cubic cell, the Miller indices of a plane are the Cartesian components of a vector normal to the plane. Considering only planes intersecting one or more lattice points, the distance d between adjacent lattice planes is related to the reciprocal lattice vector orthogonal to the planes by the formula d = 2 π | g ℓ m n | The crystallographic directions are geometric lines linking nodes of a crystal.
The crystallographic planes are geometric planes linking nodes. Some directions and planes have a higher density of nodes; these high density planes have an influence on the behavior of the crystal as follows: Optical properties: Refractive index is directly related to density. Adsorption and reactivity: Physical adsorption and chemical reactions occur at or near surface atoms or molecules; these phenomena are thus sensitive to the density of nodes. Surface tension: The condensation of a material means that the atoms, ions or molecules are more stable if they are surrounded by other similar species; the surface tension of an interface thus varies according to the density on the surface. Microstructural defects: Pores and crystallites tend to have straight grain boundaries following higher density planes. Cleavage: This occurs preferentially parallel to higher density planes. Plastic deformation: Dislocation glide occurs preferentially parallel to higher density planes; the perturbation carried by the dislocation is along a dense direction.
The shift of one node in a more dense direction requires a lesser distortion of the crystal lattice. Some directions and planes are defined by symmetry of the crystal system. In monoclinic, rhombohedral and trigonal/hexagonal systems there is one unique axis which has higher rotational symmetry than the other two axes; the basal plane is the plane perpendicular to the principal axis in these crystal systems. For triclinic and cubic crystal systems the axis designation is arbitrary and there is no principal axis. For the special case of simple cubic crystals, the lattice vectors are orthogonal and of equal length. So, in this common case, the Miller indices and both denote normals/directions in Cartesian coordinates. For cubic crystals with lattice constant a, the spacing d between adjacent l
Iron chloride called ferric chloride, is an industrial scale commodity chemical compound with iron in the +3 oxidation state. The compound exist as a hexahydrate with the formula trans-Cl · 2H2O written as FeCl3 · 6H2O; the anhydrous compound is a crytalline solid with a melting point of 307.6 °C. The color depends on the viewing angle: by reflected light the crystals appear dark green, but by transmitted light they appear purple-red; the hexahydrate appears orange-brown in color. In nature, iron chloride is known as the mineral molysite, but it is rare and found from some fumaroles, it is however an industrial scale commodity. Iron chloride dissolves in water, but undergoes partial hydrolysis in an exothermic reaction, result in a acidic solution; the resulting brown and corrosive solution is used as a flocculant in sewage treatment and drinking water production, as an etchant for copper-based metals in printed circuit boards. Anhydrous iron chloride is deliquescent, it is a strong Lewis acid, it is used as a catalyst in organic synthesis.
Anhydrous iron chloride has the BiI3 structure, with octahedral Fe centres interconnected by two-coordinate chloride ligands. Iron chloride has a low melting point and boils at around 315 °C; the vapour consists of the dimer Fe2Cl6 which dissociates into the monomeric FeCl3 at higher temperature, in competition with its reversible decomposition to give iron chloride and chlorine gas. The hexahydrate is a orange-brown solid given as the simplified formula FeCl3 · 6H2O; the compound contains a cationic coordination complex, is more properly written as trans-Cl · 2H2O with the systematic name tetraaquadichloroiron chloride dihydrate. The two H2O molecules are embedded within the monoclinic crystal structure. Anhydrous iron chloride may be prepared by union of the elements: 2 Fe + 3 Cl 2 ⟶ 2 FeCl 3 Solutions of iron chloride are produced industrially both from iron and from ore, in a closed-loop process. Dissolving iron ore in hydrochloric acid Fe 3 O 4 + 8 HCl ⟶ FeCl 2 + 2 FeCl 3 + 4 H 2 O Oxidation of iron chloride with chlorine 2 FeCl 2 + Cl 2 ⟶ 2 FeCl 3 Oxidation of iron chloride with oxygen 4 FeCl 2 + O 2 + 4 HCl ⟶ 4 FeCl 3 + 2 H 2 O Small amounts can be produced by reacting iron with hydrochloric acid with hydrogen peroxide.
The hydrogen peroxide is the oxidant in turning ferrous chloride into ferric chloride Anhydrous iron chloride cannot be obtained from the hydrate by heating. Instead, the solid decomposes into HCl and iron oxychloride; the conversion can be accomplished by treatment with thionyl chloride. Dehydration can be effected with trimethylsilyl chloride: FeCl 3 ⋅ 6 H 2 O + 12 Me 3 SiCl ⟶ FeCl 3 + 6 2 O + 12 HCl Iron chloride undergoes hydrolysis to give a acidic solution; when heated with iron oxide at 350 °C, iron chloride gives iron oxychloride, a layered solid and intercalation host. FeCl 3 + Fe 2 O 3 ⟶ 3 FeOCl The anhydrous salt is a moderately strong Lewis acid, forming adducts with Lewis bases such as triphenylphosphine oxide, it reacts with other chloride salts to give the yellow tetrahedral − ion. Salts of − in hydrochloric acid can be extract
Diethyl ether, or ether, is an organic compound in the ether class with the formula 2O, sometimes abbreviated as Et2O. It is a colorless volatile flammable liquid, it is used as a solvent in laboratories and as a starting fluid for some engines. It was used as a general anesthetic, until non-flammable drugs were developed, such as halothane, it has been used as a recreational drug to cause intoxication. Most diethyl ether is produced as a byproduct of the vapor-phase hydration of ethylene to make ethanol; this process uses solid-supported phosphoric acid catalysts and can be adjusted to make more ether if the need arises. Vapor-phase dehydration of ethanol over some alumina catalysts can give diethyl ether yields of up to 95%. Diethyl ether can be prepared both in laboratories and on an industrial scale by the acid ether synthesis. Ethanol is mixed with a strong acid sulfuric acid, H2SO4; the acid dissociates in the aqueous environment producing hydronium ions, H3O+. A hydrogen ion protonates the electronegative oxygen atom of the ethanol, giving the ethanol molecule a positive charge: CH3CH2OH + H3O+ → CH3CH2OH2+ + H2OA nucleophilic oxygen atom of unprotonated ethanol displaces a water molecule from the protonated ethanol molecule, producing water, a hydrogen ion and diethyl ether.
CH3CH2OH2+ + CH3CH2OH → H2O + H+ + CH3CH2OCH2CH3This reaction must be carried out at temperatures lower than 150 °C in order to ensure that an elimination product is not a product of the reaction. At higher temperatures, ethanol will dehydrate to form ethylene; the reaction to make diethyl ether is reversible, so an equilibrium between reactants and products is achieved. Getting a good yield of ether requires that ether be distilled out of the reaction mixture before it reverts to ethanol, taking advantage of Le Chatelier's principle. Another reaction that can be used for the preparation of ethers is the Williamson ether synthesis, in which an alkoxide performs a nucleophilic substitution upon an alkyl halide, it is important as a solvent in the production of cellulose plastics such as cellulose acetate. Diethyl ether has a high cetane number of 85–96 and is used as a starting fluid, in combination with petroleum distillates for gasoline and Diesel engines because of its high volatility and low flash point.
Ether starting fluid is sold and used in countries with cold climates, as it can help with cold starting an engine at sub-zero temperatures. For the same reason it is used as a component of the fuel mixture for carbureted compression ignition model engines. In this way diethyl ether is similar to one of its precursors, ethanol. Diethyl ether is a common laboratory aprotic solvent, it has limited solubility in water and dissolves 1.5 g/100 g water at 25 °C. This, coupled with its high volatility, makes it ideal for use as the non-polar solvent in liquid-liquid extraction; when used with an aqueous solution, the diethyl ether layer is on top as it has a lower density than the water. It is a common solvent for the Grignard reaction in addition to other reactions involving organometallic reagents. Due to its application in the manufacturing of illicit substances, it is listed in the Table II precursor under the United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances as well as substances such as acetone and sulfuric acid.
William T. G. Morton participated in a public demonstration of ether anesthesia on October 16, 1846 at the Ether Dome in Boston, Massachusetts. However, Crawford Williamson Long, is now known to have demonstrated its use as a general anesthetic in surgery to officials in Georgia, as early as March 30, 1842, Long publicly demonstrated ether's use as a surgical anesthetic on six occasions before the Boston demonstration. British doctors were aware of the anesthetic properties of ether as early as 1840 where it was prescribed in conjunction with opium. Diethyl ether supplanted the use of chloroform as a general anesthetic due to ether's more favorable therapeutic index, that is, a greater difference between an effective dose and a toxic dose. Diethyl ether increases tracheobronchial secretions. Diethyl ether could be mixed with other anesthetic agents such as chloroform to make C. E. mixture, or chloroform and alcohol to make A. C. E. Mixture. In the 21st century, ether is used; the use of flammable ether was displaced by nonflammable fluorinated hydrocarbon anesthetics.
Halothane was the first such anesthetic developed and other used inhaled anesthetics, such as isoflurane and sevoflurane, are halogenated ethers. Diethyl ether was found to have undesirable side effects, such as post-anesthetic nausea and vomiting. Modern anesthetic agents reduce these side effects. Prior to 2005 it was on the World Health Organization's List of Essential Medicines for use as an anesthetic. Ether was once used in pharmaceutical formulations. A mixture of alcohol and ether, one part of diethyl ether and three parts of ethanol, was known as "Spirit of ether", Hoffman's Anodyne or Hoffman's Drops. In the United States this concoction was removed from the Pharmacopeia at some point prior to June 1917, as a study published by William Procter, Jr. in the American Journal of Pharmacy as early as 1852 showed that there were differences in formulation to be found between commercial manufacturers, between international pharmacopoeia, from Hoffman's original recipe. The anesthetic and intoxicating effects of ether have made it a recreational drug.
Diethyl ether in anesthetic dosage is an inhalant which has a long history
Rhenium trioxide or rhenium oxide is an inorganic compound with the formula ReO3. It is a red solid with a metallic lustre copper in appearance, it is the only stable trioxide of the Group 7 elements. Rhenium trioxide can be formed by reducing rhenium oxide with carbon monoxide. Re2O7 + CO → 2 ReO3 + CO2Re2O7 can be reduced with dioxane. Rhenium oxide crystallizes with a primitive cubic unit cell, with a lattice parameter of 3.742 Å. The structure of ReO3 is similar to that of perovskite, without the large A cation at the centre of the unit cell; each rhenium center is surrounded by an octahedron defined by six oxygen centers. These octahedra share corners to form the 3-dimensional structure; the coordination number of O is 2. Upon heating to 400 °C under vacuum, it undergoes disproportionation: 3 ReO3 → Re2O7 + ReO2ReO3 is unusual for an oxide because it exhibits low resistivity, it behaves like a metal. At 300 K, its resistivity is 100.0 nΩ·m, whereas at 100 K, this decreases to 6.0 nΩ·m, 17 times less than at 300 K. Rhenium trioxide finds some use in organic synthesis as a catalyst for amide reduction
Solubility is the property of a solid, liquid or gaseous chemical substance called solute to dissolve in a solid, liquid or gaseous solvent. The solubility of a substance fundamentally depends on the physical and chemical properties of the solute and solvent as well as on temperature and presence of other chemicals of the solution; the extent of the solubility of a substance in a specific solvent is measured as the saturation concentration, where adding more solute does not increase the concentration of the solution and begins to precipitate the excess amount of solute. Insolubility is the inability to dissolve in a liquid or gaseous solvent. Most the solvent is a liquid, which can be a pure substance or a mixture. One may speak of solid solution, but of solution in a gas. Under certain conditions, the equilibrium solubility can be exceeded to give a so-called supersaturated solution, metastable. Metastability of crystals can lead to apparent differences in the amount of a chemical that dissolves depending on its crystalline form or particle size.
A supersaturated solution crystallises when'seed' crystals are introduced and rapid equilibration occurs. Phenylsalicylate is one such simple observable substance when melted and cooled below its fusion point. Solubility is not to be confused with the ability to'dissolve' a substance, because the solution might occur because of a chemical reaction. For example, zinc'dissolves' in hydrochloric acid as a result of a chemical reaction releasing hydrogen gas in a displacement reaction; the zinc ions are soluble in the acid. The solubility of a substance is an different property from the rate of solution, how fast it dissolves; the smaller a particle is, the faster it dissolves although there are many factors to add to this generalization. Crucially solubility applies to all areas of chemistry, inorganic, physical and biochemistry. In all cases it will depend on the physical conditions and the enthalpy and entropy directly relating to the solvents and solutes concerned. By far the most common solvent in chemistry is water, a solvent for most ionic compounds as well as a wide range of organic substances.
This is a crucial factor in much environmental and geochemical work. According to the IUPAC definition, solubility is the analytical composition of a saturated solution expressed as a proportion of a designated solute in a designated solvent. Solubility may be stated in various units of concentration such as molarity, mole fraction, mole ratio, mass per volume and other units; the extent of solubility ranges from infinitely soluble such as ethanol in water, to poorly soluble, such as silver chloride in water. The term insoluble is applied to poorly or poorly soluble compounds. A number of other descriptive terms are used to qualify the extent of solubility for a given application. For example, U. S. Pharmacopoeia gives the following terms: The thresholds to describe something as insoluble, or similar terms, may depend on the application. For example, one source states that substances are described as "insoluble" when their solubility is less than 0.1 g per 100 mL of solvent. Solubility occurs under dynamic equilibrium, which means that solubility results from the simultaneous and opposing processes of dissolution and phase joining.
The solubility equilibrium occurs. The term solubility is used in some fields where the solute is altered by solvolysis. For example, many metals and their oxides are said to be "soluble in hydrochloric acid", although in fact the aqueous acid irreversibly degrades the solid to give soluble products, it is true that most ionic solids are dissolved by polar solvents, but such processes are reversible. In those cases where the solute is not recovered upon evaporation of the solvent, the process is referred to as solvolysis; the thermodynamic concept of solubility does not apply straightforwardly to solvolysis. When a solute dissolves, it may form several species in the solution. For example, an aqueous suspension of ferrous hydroxide, Fe2, will contain the series + as well as other species. Furthermore, the solubility of ferrous hydroxide and the composition of its soluble components depend on pH. In general, solubility in the solvent phase can be given only for a specific solute, thermodynamically stable, the value of the solubility will include all the species in the solution.
Solubility is defined for specific phases. For example, the solubility of aragonite and calcite in water are expected to differ though they are both polymorphs of calcium carbonate and have the same chemical formula; the solubility of one substance in another is determined by the balance of intermolecular forces between the solvent and solute, the entropy change that accompanies the solvation. Factors such as temperature and pressure will alter this balance. Solubility may strongly depend on the presence of other species dissolved in the solvent, for example, complex-forming anions in liquids. Solubility will depend on the excess or deficiency of a common ion in the solution, a phenomenon known as the common-ion effect. To a lesser extent, solubility will depend on the ionic strength of solutions; the last two effects can be quantified using the equation for solubility equilibrium. For a solid that dissolves in a redox reaction, solubility is expe
Iron oxide or ferric oxide is the inorganic compound with the formula Fe2O3. It is one of the three main oxides of iron, the other two being iron oxide, rare; as the mineral known as hematite, Fe2O3 is the main source of iron for the steel industry. Fe2O3 is attacked by acids. Iron oxide is called rust, to some extent this label is useful, because rust shares several properties and has a similar composition. To a chemist, rust is considered an ill-defined material, described as hydrated ferric oxide. Fe2O3 can be obtained in various polymorphs. In the main ones, α and γ, iron adopts octahedral coordination geometry; that is, each Fe center is bound to six oxygen ligands. Α-Fe2O3 is the most common form. It occurs as the mineral hematite, mined as the main ore of iron, it is antiferromagnetic below ~260 K, exhibits weak ferromagnetism between 260 K and the Néel temperature, 950 K. It is easy to prepare using both thermal precipitation in the liquid phase, its magnetic properties are dependent on many factors, e.g. pressure, particle size, magnetic field intensity.
Γ-Fe2O3 has a cubic structure. It is metastable and converted from the alpha phase at high temperatures, it occurs as the mineral maghemite. It is ferromagnetic and finds application in recording tapes, although ultrafine particles smaller than 10 nanometers are superparamagnetic, it can be prepared by thermal dehydratation of gamma iron oxide-hydroxide. Another method involves the careful oxidation of iron oxide; the ultrafine particles can be prepared by thermal decomposition of iron oxalate. Several other phases have been claimed; the β-phase is cubic body-centered, at temperatures above 500 °C converts to alpha phase. It can be prepared by reduction of hematite by carbon, pyrolysis of iron chloride solution, or thermal decomposition of iron sulfate; the epsilon phase is rhombic, shows properties intermediate between alpha and gamma, may have useful magnetic properties. Preparation of the pure epsilon phase has proven challenging due to contamination with alpha and gamma phases. Material with a high proportion of epsilon phase can be prepared by thermal transformation of the gamma phase.
This phase is metastable, transforming to the alpha phase at between 500 and 750 °C. Can be prepared by oxidation of iron in an electric arc or by sol-gel precipitation from iron nitrate. Additionally at high pressure an amorphous form is claimed. Recent research has revealed epsilon iron oxide in ancient Chinese Jian ceramic glazes, which may provide insight into ways to produce that form in the lab. Several hydrates of Iron oxide exists; when alkali is added to solutions of soluble Fe salts, a red-brown gelatinous precipitate forms. This is not Fe3, but Fe2O3·H2O. Several forms of the hydrated oxide of Fe exist as well; the red lepidocrocite γ-FeOH, occurs on the outside of rusticles, the orange goethite, which occurs internally in rusticles. When Fe2O3·H2O is heated, it loses its water of hydration. Further heating at 1670 K converts Fe2O3 to black Fe3O4, known as the mineral magnetite. FeOH is soluble in acids, giving 3+. In concentrated aqueous alkali, Fe2O3 gives 3−; the most important reaction is its carbothermal reduction, which gives iron used in steel-making: Fe2O3 + 3 CO → 2 Fe + 3 CO2Another redox reaction is the exothermic thermite reaction with aluminium.
2 Al + Fe2O3 → 2 Fe + Al2O3This process is used to weld thick metals such as rails of train tracks by using a ceramic container to funnel the molten iron in between two sections of rail. Thermite is used in weapons and making small-scale cast-iron sculptures and tools. Partial reduction with hydrogen at about 400 °C produces magnetite, a black magnetic material that contains both Fe and Fe: 3 Fe2O3 + H2 → 2 Fe3O4 + H2OIron oxide is insoluble in water but dissolves in strong acid, e.g. hydrochloric and sulfuric acids. It dissolves well in solutions of chelating agents such as EDTA and oxalic acid. Heating iron oxides with other metal oxides or carbonates yields materials known as ferrates: ZnO + Fe2O3 → Zn2 Iron oxide is a product of the oxidation of iron, it can be prepared in the laboratory by electrolyzing a solution of sodium bicarbonate, an inert electrolyte, with an iron anode: 4 Fe + 3 O2 + 2 H2O → 4 FeOThe resulting hydrated iron oxide, written here as FeOH, dehydrates around 200 °C. 2 FeO → Fe2O3 + H2O The overwhelming application of iron oxide is as the feedstock of the steel and iron industries, e.g. the production of iron and many alloys.
A fine powder of ferric oxide is known as "jeweler's rouge", "red rouge", or rouge. It is used to put the final polish on metallic jewelry and lenses, as a cosmetic. Rouge cuts more than some modern polishes, such as cerium oxide, but is still used in optics fabrication and by jewelers for the superior finish it can produce; when polishing gold, the rouge stains the gold, which contributes to the appearance of the finished piece. Rouge is sold as a powder, laced on polishing cloths, or solid bar. Other polishing compounds are often called "rouge" when they do not contain iron oxide. Jewelers remove the residual rouge on jewelry by use of ultrasonic cleaning. Products sold as "stropping compound" are applied to a leather stro
Norman Neill Greenwood FRS CChem FRSC was an Australian-British chemist and Emeritus Professor at the University of Leeds. He is best known for the innovative textbook Chemistry of the Elements, co-authored with Alan Earnshaw, first published in 1984. After attending University High School, Greenwood read Chemistry at the University of Melbourne and graduated with a BSc in 1945 and an MSc in 1948. In 1948, he was awarded the Exhibition of 1851 Scholarship to enable him to read for a PhD at Sidney Sussex College, Cambridge under the supervision of Harry Julius Emeléus, he received the PhD in 1951. Greenwood was a senior research fellow at the Atomic Energy Research Establishment from 1951 until 1953 when he was appointed a lecturer at the University of Nottingham, his first PhD student at Nottingham was Kenneth Wade. Professor William Wynne-Jones, the Chairman of the School of Chemistry at Kings College, recruited Greenwood to the first established chair of inorganic chemistry in the country in 1961.
Greenwood was appointed professor and head of the Department of Inorganic and Structural Chemistry at the University of Leeds in 1971, a post which he held until his retirement in 1990 when he was given the title emeritus professor. Greenwood was elected a fellow of the Royal Society in 1987, his wide-ranging researches in inorganic and structural chemistry have made major advances in the chemistry of boron hydrides and other main-group element compounds. He pioneered the application of Mössbauer spectroscopy to problems in chemistry, he was a prolific writer and inspirational lecturer on chemical and educational themes, has held numerous visiting professorships throughout the world. He was appointed by NASA as principal investigator in the study of lunar rocks, he served as chairman of the IUPAC Commission on Atomic Weights from 1970 to 1975 and as president of the IUPAC Inorganic Chemistry Division. Greenwood, N. N.. Principles of Atomic Orbitals – Monograph for Teachers. Royal Society of Chemistry.
P. 48. ISBN 9780854040285. Greenwood, N. N.. Ionic crystals, lattice defects and nonstoichiometry. Butterworths. P. 194. Greenwood, N. N. C.. Mössbauer Spectroscopy. Chapman and Hall. P. 659. Greenwood, Norman N.. Chemistry of the Elements. Butterworth-Heinemann. P. 1340. ISBN 978-0-08-037941-8. Greenwood, N. N.. Recollections of a Scientist Volume 1. Boyhood and Youth in Australia. Xlibris Corporation. P. 288. ISBN 1-4691-7935-0. Greenwood, N. N.. Recollections of a Scientist, Volume 2: Expanding Horizons: England and Europe. Xlibris Corporation. P. 438. ISBN 978-1477151860. Editor: Spectroscopic Properties of Inorganic and Organometallic Compounds, Royal Society of Chemistry, Volume 1 to Volume 9 Norman Greenwood tells his life story at Web of Stories