A metal is a material that, when freshly prepared, polished, or fractured, shows a lustrous appearance, conducts electricity and heat well. Metals are malleable or ductile. A metal may be an alloy such as stainless steel. In physics, a metal is regarded as any substance capable of conducting electricity at a temperature of absolute zero. Many elements and compounds that are not classified as metals become metallic under high pressures. For example, the nonmetal iodine becomes a metal at a pressure of between 40 and 170 thousand times atmospheric pressure; some materials regarded as metals can become nonmetals. Sodium, for example, becomes a nonmetal at pressure of just under two million times atmospheric pressure. In chemistry, two elements that would otherwise qualify as brittle metals—arsenic and antimony—are instead recognised as metalloids, on account of their predominately non-metallic chemistry. Around 95 of the 118 elements in the periodic table are metals; the number is inexact as the boundaries between metals and metalloids fluctuate due to a lack of universally accepted definitions of the categories involved.
In astrophysics the term "metal" is cast more to refer to all chemical elements in a star that are heavier than the lightest two and helium, not just traditional metals. A star fuses lighter atoms hydrogen and helium, into heavier atoms over its lifetime. Used in that sense, the metallicity of an astronomical object is the proportion of its matter made up of the heavier chemical elements. Metals are present in many aspects of modern life; the strength and resilience of some metals has led to their frequent use in, for example, high-rise building and bridge construction, as well as most vehicles, many home appliances, tools and railroad tracks. Precious metals were used as coinage, but in the modern era, coinage metals have extended to at least 23 of the chemical elements; the history of metals is thought to begin with the use of copper about 11,000 years ago. Gold, iron and brass were in use before the first known appearance of bronze in the 5th millennium BCE. Subsequent developments include the production of early forms of steel.
Metals are lustrous, at least when freshly prepared, polished, or fractured. Sheets of metal thicker than a few micrometres appear opaque; the solid or liquid state of metals originates in the capacity of the metal atoms involved to lose their outer shell electrons. Broadly, the forces holding an individual atom’s outer shell electrons in place are weaker than the attractive forces on the same electrons arising from interactions between the atoms in the solid or liquid metal; the electrons involved become delocalised and the atomic structure of a metal can be visualised as a collection of atoms embedded in a cloud of mobile electrons. This type of interaction is called a metallic bond; the strength of metallic bonds for different elemental metals reaches a maximum around the center of the transition metal series, as these elements have large numbers of delocalized electrons. Although most elemental metals have higher densities than most nonmetals, there is a wide variation in their densities, lithium being the least dense and osmium the most dense.
Magnesium and titanium are light metals of significant commercial importance. Their respective densities of 1.7, 2.7 and 4.5 g/cm3 can be compared to those of the older structural metals, like iron at 7.9 and copper at 8.9 g/cm3. An iron ball would thus weigh about as much as three aluminium balls. Metals are malleable and ductile, deforming under stress without cleaving; the nondirectional nature of metallic bonding is thought to contribute to the ductility of most metallic solids. In contrast, in an ionic compound like table salt, when the planes of an ionic bond slide past one another, the resultant change in location shifts ions of the same charge into close proximity, resulting in the cleavage of the crystal; such a shift is not observed in a covalently bonded crystal, such as a diamond, where fracture and crystal fragmentation occurs. Reversible elastic deformation in metals can be described by Hooke's Law for restoring forces, where the stress is linearly proportional to the strain. Heat or forces larger than a metal's elastic limit may cause a permanent deformation, known as plastic deformation or plasticity.
An applied force may be a compressive force, or a shear, bending or torsion force. A temperature change may affect the movement or displacement of structural defects in the metal such as grain boundaries, point vacancies and screw dislocations, stacking faults and twins in both crystalline and non-crystalline metals. Internal slip and metal fatigue may ensue; the atoms of metallic substances are arranged in one of three common crystal structures, namely body-centered cubic, face-centered cubic, hexagonal close-packed. In bcc, each atom is positioned at the center of a cube of eight others. In fcc and hcp, each atom is surrounded by twelve others; some metals adopt different structures depending on the temperature. The
Citric acid is a weak organic acid that has the chemical formula C6H8O7. It occurs in citrus fruits. In biochemistry, it is an intermediate in the citric acid cycle, which occurs in the metabolism of all aerobic organisms. More than a million tons of citric acid are manufactured every year, it is used as an acidifier, as a flavoring and chelating agent. A citrate is a derivative of citric acid. An example of the former, a salt is trisodium citrate; when part of a salt, the formula of the citrate ion is written as C6H5O3−7 or C3H5O3−3. Citric acid exists in greater than trace amounts in a variety of fruits and vegetables, most notably citrus fruits. Lemons and limes have high concentrations of the acid; the concentrations of citric acid in citrus fruits range from 0.005 mol/L for oranges and grapefruits to 0.30 mol/L in lemons and limes. Within species, these values vary depending on the cultivar and the circumstances in which the fruit was grown. Industrial-scale citric acid production first began in 1890 based on the Italian citrus fruit industry, where the juice was treated with hydrated lime to precipitate calcium citrate, isolated and converted back to the acid using diluted sulfuric acid.
In 1893, C. Wehmer discovered. However, microbial production of citric acid did not become industrially important until World War I disrupted Italian citrus exports. In 1917, American food chemist James Currie discovered certain strains of the mold Aspergillus niger could be efficient citric acid producers, the pharmaceutical company Pfizer began industrial-level production using this technique two years followed by Citrique Belge in 1929. In this production technique, still the major industrial route to citric acid used today, cultures of A. niger are fed on a sucrose or glucose-containing medium to produce citric acid. The source of sugar is corn steep liquor, hydrolyzed corn starch or other inexpensive sugary solutions. After the mold is filtered out of the resulting solution, citric acid is isolated by precipitating it with calcium hydroxide to yield calcium citrate salt, from which citric acid is regenerated by treatment with sulfuric acid, as in the direct extraction from citrus fruit juice.
In 1977, a patent was granted to Lever Brothers for the chemical synthesis of citric acid starting either from aconitic or isocitrate/alloisocitrate calcium salts under high pressure conditions. This produced citric acid in near quantitative conversion under what appeared to be a reverse non-enzymatic Krebs cycle reaction. In 2007, worldwide annual production stood at 1,600,000 tons. More than 50% of this volume was produced in China. More than 50% was used as an acidity regulator in beverages, some 20% in other food applications, 20% for detergent applications and 10% for related applications other than food, such as cosmetics, pharmaceutics and in the chemical industry. Citric acid was first isolated in 1784 by the chemist Carl Wilhelm Scheele, who crystallized it from lemon juice, it can exist either as a monohydrate. The anhydrous form crystallizes from hot water, while the monohydrate forms when citric acid is crystallized from cold water; the monohydrate can be converted to the anhydrous form at about 78 °C.
Citric acid dissolves in absolute ethanol at 15 °C. It decomposes with loss of carbon dioxide above about 175 °C. Citric acid is considered to be a tribasic acid, with pKa values, extrapolated to zero ionic strength, of 5.21, 4.28 and 2.92 at 25 °C. The pKa of the hydroxyl group has been found, by means of 13C NMR spectroscopy, to be 14.4. The speciation diagram shows that solutions of citric acid are buffer solutions between about pH 2 and pH 8. In biological systems around pH 7, the two species present are the citrate ion and mono-hydrogen citrate ion; the SSC 20X hybridization buffer is an example in common use. Tables compiled for biochemical studies are available. On the other hand, the pH of a 1 mM solution of citric acid will be about 3.2. The pH of fruit juices from citrus fruits like oranges and lemons depends on the citric acid concentration, being lower for higher acid concentration and conversely. Acid salts of citric acid can be prepared by careful adjustment of the pH before crystallizing the compound.
See, for example, sodium citrate. The citrate ion forms complexes with metallic cations; the stability constants for the formation of these complexes are quite large because of the chelate effect. It forms complexes with alkali metal cations. However, when a chelate complex is formed using all three carboxylate groups, the chelate rings have 7 and 8 members, which are less stable thermodynamically than smaller chelate rings. In consequence, the hydroxyl group can be deprotonated, forming part of a more stable 5-membered ring, as in ammonium ferric citrate, 5Fe2·2H2O. Citric acid can be esterified at one or more of the carboxylic acid functional groups on the molecule, to form any of a variety of mono-, di-, tri-, mixed esters. Citrate is an intermediate in the TCA cycle, a central metabolic pathway for animals and bacteria. Citrate synthase catalyzes the condensation of oxaloacetate with acetyl CoA to form citrate. Citrate acts as the substrate for aconitase and is converted into aconitic acid.
The cycle ends with regeneration of oxaloacetate. This series
A leachate is any liquid that, in the course of passing through matter, extracts soluble or suspended solids, or any other component of the material through which it has passed. Leachate is a used term in the environmental sciences where it has the specific meaning of a liquid that has dissolved or entrained environmentally harmful substances that may enter the environment, it is most used in the context of land-filling of putrescible or industrial waste. In the narrow environmental context leachate is therefore any liquid material that drains from land or stockpiled material and contains elevated concentrations of undesirable material derived from the material that it has passed through. Leachate from a landfill varies in composition depending on the age of the landfill and the type of waste that it contains, it contains both dissolved and suspended material. The generation of leachate is caused principally by precipitation percolating through waste deposited in a landfill. Once in contact with decomposing solid waste, the percolating water becomes contaminated, if it flows out of the waste material it is termed leachate.
Additional leachate volume is produced during this decomposition of carbonaceous material producing a wide range of other materials including methane, carbon dioxide and a complex mixture of organic acids, aldehydes and simple sugars. The risks of leachate generation can be mitigated by properly designed and engineered landfill sites, such as those that are constructed on geologically impermeable materials or sites that use impermeable liners made of geomembranes or engineered clay; the use of linings is now mandatory within the United States and the European Union except where the waste is deemed inert. In addition, most toxic and difficult materials are now excluded from landfilling. However, despite much stricter statutory controls, leachates from modern sites are found to contain a range of contaminants stemming from illegal activity or discarded household and domestic products; when water percolates through waste, it promotes and assists the process of decomposition by bacteria and fungi.
These processes in turn release by-products of decomposition and use up any available oxygen, creating an anoxic environment. In decomposing waste, the temperature rises and the pH falls with the result that many metal ions that are insoluble at neutral pH become dissolved in the developing leachate; the decomposition processes themselves release more water. Leachate reacts with materials that are not prone to decomposition themselves, such as fire ash, cement-based building materials and gypsum-based materials changing the chemical composition. In sites with large volumes of building waste those containing gypsum plaster, the reaction of leachate with the gypsum can generate large volumes of hydrogen sulfide, which may be released in the leachate and may form a large component of the landfill gas. In a landfill that receives a mixture of municipal and mixed industrial waste but excludes significant amounts of concentrated chemical waste, landfill leachate may be characterized as a water-based solution of four groups of contaminants: dissolved organic matter, inorganic macro components, heavy metals, xenobiotic organic compounds such as halogenated organics.
The physical appearance of leachate when it emerges from a typical landfill site is a odoured black-, yellow- or orange-coloured cloudy liquid. The smell is acidic and offensive and may be pervasive because of hydrogen-, nitrogen- and sulfur-rich organic species such as mercaptans. In older landfills and those with no membrane between the waste and the underlying geology, leachate is free to leave the waste and flow directly into the groundwater. In such cases, high concentrations of leachate are found in nearby springs and flushes; as leachate first emerges it can be black in colour and effervescent, with dissolved and entrained gases. As it becomes oxygenated it tends to turn brown or yellow because of the presence of iron salts in solution and in suspension, it quickly develops a bacterial flora comprising substantial growths of Sphaerotilus natans. In the UK, in the late 1960s, central Government policy was to ensure new landfill sites were being chosen with permeable underlying geological strata to avoid the build-up of leachate.
This policy was dubbed "dilute and disperse". However, following a number of cases where this policy was seen to be failing, an exposee in The Sunday Times of serious environmental damage being caused by inappropriate disposal of industrial wastes, both policy and the law were changed; the Deposit of Poisonous Wastes Act 1972, together with The 1974 Local Government Act, made local government responsible for waste disposal and for the enforcement of environmental standards regarding waste disposal. Proposed landfill locations had to be justified not only by geography but scientifically. Many European countries decided to select landfill sites in groundwater-free clay geological conditions or to require that the site have an engineered lining. In the wake of European advancements, the United States increased its development of leachate retaining and collection systems; this led from lining in principle to the use of multiple lining layers in all landfills. The primary criterion for design of the leachate system is that all leachate be collected and removed from the landfill at a rate sufficient to pr
Cobalt is a chemical element with symbol Co and atomic number 27. Like nickel, cobalt is found in the Earth's crust only in chemically combined form, save for small deposits found in alloys of natural meteoric iron; the free element, produced by reductive smelting, is a hard, silver-gray metal. Cobalt-based blue pigments have been used since ancient times for jewelry and paints, to impart a distinctive blue tint to glass, but the color was thought by alchemists to be due to the known metal bismuth. Miners had long used the name kobold ore for some of the blue-pigment producing minerals. In 1735, such ores were found to be reducible to a new metal, this was named for the kobold. Today, some cobalt is produced from one of a number of metallic-lustered ores, such as for example cobaltite; the element is however more produced as a by-product of copper and nickel mining. The copper belt in the Democratic Republic of the Congo and Zambia yields most of the global cobalt production; the DRC alone accounted for more than 50% of world production in 2016, according to Natural Resources Canada.
Cobalt is used in the manufacture of magnetic, wear-resistant and high-strength alloys. The compounds cobalt silicate and cobalt aluminate give a distinctive deep blue color to glass, inks and varnishes. Cobalt occurs as only one stable isotope, cobalt-59. Cobalt-60 is a commercially important radioisotope, used as a radioactive tracer and for the production of high energy gamma rays. Cobalt is the active center of a group of coenzymes called cobalamins. Vitamin B12, the best-known example of the type, is an essential vitamin for all animals. Cobalt in inorganic form is a micronutrient for bacteria and fungi. Cobalt is a ferromagnetic metal with a specific gravity of 8.9. The Curie temperature is 1,115 °C and the magnetic moment is 1.6–1.7 Bohr magnetons per atom. Cobalt has a relative permeability two-thirds. Metallic cobalt occurs as two crystallographic structures: fcc; the ideal transition temperature between the hcp and fcc structures is 450 °C, but in practice the energy difference between them is so small that random intergrowth of the two is common.
Cobalt is a weakly reducing metal, protected from oxidation by a passivating oxide film. It is attacked by halogens and sulfur. Heating in oxygen produces Co3O4 which loses oxygen at 900 °C to give the monoxide CoO; the metal reacts with fluorine at 520 K to give CoF3. It does not react with hydrogen gas or nitrogen gas when heated, but it does react with boron, phosphorus and sulfur. At ordinary temperatures, it reacts with mineral acids, slowly with moist, but not with dry, air. Common oxidation states of cobalt include +2 and +3, although compounds with oxidation states ranging from −3 to +5 are known. A common oxidation state for simple compounds is +2; these salts form the pink-colored metal aquo complex 2+ in water. Addition of chloride gives the intensely blue 2−. In a borax bead flame test, cobalt shows deep blue in both reducing flames. Several oxides of cobalt are known. Green cobalt oxide has rocksalt structure, it is oxidized with water and oxygen to brown cobalt hydroxide. At temperatures of 600 -- 700 °C, CoO oxidizes to the blue cobalt oxide.
Black cobalt oxide is known. Cobalt oxides are antiferromagnetic at low temperature: CoO and Co3O4, analogous to magnetite, with a mixture of +2 and +3 oxidation states; the principal chalcogenides of cobalt include the black cobalt sulfides, CoS2, which adopts a pyrite-like structure, cobalt sulfide. Four dihalides of cobalt are known: cobalt fluoride, cobalt chloride, cobalt bromide, cobalt iodide; these halides exist in hydrated forms. Whereas the anhydrous dichloride is blue, the hydrate is red; the reduction potential for the reaction Co3+ + e− → Co2+ is +1.92 V, beyond that for chlorine to chloride, +1.36 V. Consequently and chloride would result in the cobalt being reduced to cobalt; because the reduction potential for fluorine to fluoride is so high, +2.87 V, cobalt fluoride is one of the few simple stable cobalt compounds. Cobalt fluoride, used in some fluorination reactions, reacts vigorously with water; as for all metals, molecular compounds and polyatomic ions of cobalt are classified as coordination complexes, that is, molecules or ions that contain cobalt linked to several ligands.
The principles of electronegativity and hardness–softness of a series of ligands can be used to explain the usual oxidation state of cobalt. For example, Co+3 complexes tend to have ammine ligands; because phosphorus is softer than nitrogen, phosphine ligands tend to feature the softer Co2+ and Co+, an example being triscobalt chloride. The more electronegative oxide and fluoride can stabilize Co4+ and Co5+ derivatives, e.g. caesium hexafluorocobaltate and potassium percobaltate. Alfred Werner, a Nobel-prize winning pioneer in coordination chemistry, worked with compounds of empirical formula 3+. One of the isomers
An acid is a molecule or ion capable of donating a hydron, or, capable of forming a covalent bond with an electron pair. The first category of acids is the proton donors or Brønsted acids. In the special case of aqueous solutions, proton donors form the hydronium ion H3O+ and are known as Arrhenius acids. Brønsted and Lowry generalized the Arrhenius theory to include non-aqueous solvents. A Brønsted or Arrhenius acid contains a hydrogen atom bonded to a chemical structure, still energetically favorable after loss of H+. Aqueous Arrhenius acids have characteristic properties which provide a practical description of an acid. Acids form aqueous solutions with a sour taste, can turn blue litmus red, react with bases and certain metals to form salts; the word acid is derived from the Latin acidus/acēre meaning sour. An aqueous solution of an acid has a pH less than 7 and is colloquially referred to as'acid', while the strict definition refers only to the solute. A lower pH means a higher acidity, thus a higher concentration of positive hydrogen ions in the solution.
Chemicals or substances having the property of an acid are said to be acidic. Common aqueous acids include hydrochloric acid, acetic acid, sulfuric acid, citric acid; as these examples show, acids can be solutions or pure substances, can be derived from acids that are solids, liquids, or gases. Strong acids and some concentrated weak acids are corrosive, but there are exceptions such as carboranes and boric acid; the second category of acids are Lewis acids. An example is boron trifluoride, whose boron atom has a vacant orbital which can form a covalent bond by sharing a lone pair of electrons on an atom in a base, for example the nitrogen atom in ammonia. Lewis considered this as a generalization of the Brønsted definition, so that an acid is a chemical species that accepts electron pairs either directly or by releasing protons into the solution, which accept electron pairs. However, hydrogen chloride, acetic acid, most other Brønsted-Lowry acids cannot form a covalent bond with an electron pair and are therefore not Lewis acids.
Conversely, many Lewis acids are not Brønsted-Lowry acids. In modern terminology, an acid is implicitly a Brønsted acid and not a Lewis acid, since chemists always refer to a Lewis acid explicitly as a Lewis acid. Modern definitions are concerned with the fundamental chemical reactions common to all acids. Most acids encountered in everyday life are aqueous solutions, or can be dissolved in water, so the Arrhenius and Brønsted-Lowry definitions are the most relevant; the Brønsted-Lowry definition is the most used definition. Hydronium ions are acids according to all three definitions. Although alcohols and amines can be Brønsted-Lowry acids, they can function as Lewis bases due to the lone pairs of electrons on their oxygen and nitrogen atoms; the Swedish chemist Svante Arrhenius attributed the properties of acidity to hydrogen ions or protons in 1884. An Arrhenius acid is a substance that, when added to water, increases the concentration of H+ ions in the water. Note that chemists write H+ and refer to the hydrogen ion when describing acid-base reactions but the free hydrogen nucleus, a proton, does not exist alone in water, it exists as the hydronium ion, H3O+.
Thus, an Arrhenius acid can be described as a substance that increases the concentration of hydronium ions when added to water. Examples include molecular substances such as acetic acid. An Arrhenius base, on the other hand, is a substance which increases the concentration of hydroxide ions when dissolved in water; this decreases the concentration of hydronium because the ions react to form H2O molecules: H3O+ + OH− ⇌ H2O + H2ODue to this equilibrium, any increase in the concentration of hydronium is accompanied by a decrease in the concentration of hydroxide. Thus, an Arrhenius acid could be said to be one that decreases hydroxide concentration, while an Arrhenius base increases it. In an acidic solution, the concentration of hydronium ions is greater than 10−7 moles per liter. Since pH is defined as the negative logarithm of the concentration of hydronium ions, acidic solutions thus have a pH of less than 7. While the Arrhenius concept is useful for describing many reactions, it is quite limited in its scope.
In 1923 chemists Johannes Nicolaus Brønsted and Thomas Martin Lowry independently recognized that acid-base reactions involve the transfer of a proton. A Brønsted-Lowry acid is a species. Brønsted-Lowry acid-base theory has several advantages over Arrhenius theory. Consider the following reactions of acetic acid, the organic acid that gives vinegar its characteristic taste: CH3COOH + H2O ⇌ CH3COO− + H3O+ CH3COOH + NH3 ⇌ CH3COO− + NH+4Both theories describe the first reaction: CH3COOH acts as an Arrhenius acid because it acts as a source of H3O+ when dissolved in water, it acts as a Brønsted acid by donating a proton to water. In the second example CH3COOH undergoes the same transformation, in this case donating a proton to ammonia, but does not relate to the Arrhenius definition of an acid because the reaction does not produce hydronium. CH3COOH is
Redox is a chemical reaction in which the oxidation states of atoms are changed. Any such reaction involves both a reduction process and a complementary oxidation process, two key concepts involved with electron transfer processes. Redox reactions include all chemical reactions; the chemical species from which the electron is stripped is said to have been oxidized, while the chemical species to which the electron is added is said to have been reduced. It can be explained in simple terms: Oxidation is the loss of electrons or an increase in oxidation state by a molecule, atom, or ion. Reduction is a decrease in oxidation state by a molecule, atom, or ion; as an example, during the combustion of wood, oxygen from the air is reduced, gaining electrons from carbon, oxidized. Although oxidation reactions are associated with the formation of oxides from oxygen molecules, oxygen is not included in such reactions, as other chemical species can serve the same function; the reaction can occur slowly, as with the formation of rust, or more in the case of fire.
There are simple redox processes, such as the oxidation of carbon to yield carbon dioxide or the reduction of carbon by hydrogen to yield methane, more complex processes such as the oxidation of glucose in the human body. "Redox" is a portmanteau of the words "reduction" and "oxidation". The word oxidation implied reaction with oxygen to form an oxide, since dioxygen was the first recognized oxidizing agent; the term was expanded to encompass oxygen-like substances that accomplished parallel chemical reactions. The meaning was generalized to include all processes involving loss of electrons; the word reduction referred to the loss in weight upon heating a metallic ore such as a metal oxide to extract the metal. In other words, ore was "reduced" to metal. Antoine Lavoisier showed. Scientists realized that the metal atom gains electrons in this process; the meaning of reduction became generalized to include all processes involving a gain of electrons. Though "reduction" seems counter-intuitive when speaking of the gain of electrons, it might help to think of reduction as the loss of oxygen, its historical meaning.
Since electrons are negatively charged, it is helpful to think of this as reduction in electrical charge. The electrochemist John Bockris has used the words electronation and deelectronation to describe reduction and oxidation processes when they occur at electrodes; these words are analogous to protonation and deprotonation, but they have not been adopted by chemists worldwide. The term "hydrogenation" could be used instead of reduction, since hydrogen is the reducing agent in a large number of reactions in organic chemistry and biochemistry. But, unlike oxidation, generalized beyond its root element, hydrogenation has maintained its specific connection to reactions that add hydrogen to another substance; the word "redox" was first used in 1928. The processes of oxidation and reduction occur and cannot happen independently of one another, similar to the acid–base reaction; the oxidation alone and the reduction alone are each called a half-reaction, because two half-reactions always occur together to form a whole reaction.
When writing half-reactions, the gained or lost electrons are included explicitly in order that the half-reaction be balanced with respect to electric charge. Though sufficient for many purposes, these general descriptions are not correct. Although oxidation and reduction properly refer to a change in oxidation state — the actual transfer of electrons may never occur; the oxidation state of an atom is the fictitious charge that an atom would have if all bonds between atoms of different elements were 100% ionic. Thus, oxidation is best defined as an increase in oxidation state, reduction as a decrease in oxidation state. In practice, the transfer of electrons will always cause a change in oxidation state, but there are many reactions that are classed as "redox" though no electron transfer occurs. In redox processes, the reductant transfers electrons to the oxidant. Thus, in the reaction, the reductant or reducing agent loses electrons and is oxidized, the oxidant or oxidizing agent gains electrons and is reduced.
The pair of an oxidizing and reducing agent that are involved in a particular reaction is called a redox pair. A redox couple is a reducing species and its corresponding oxidizing form, e.g. Fe2+/Fe3+ Substances that have the ability to oxidize other substances are said to be oxidative or oxidizing and are known as oxidizing agents, oxidants, or oxidizers; that is, the oxidant removes electrons from another substance, is thus itself reduced. And, because it "accepts" electrons, the oxidizing agent is called an electron acceptor. Oxygen is the quintessential oxidizer. Oxidants are chemical substances with elements in high oxidation states, or else electronegative elements that can gain extra electrons by oxidizing another substance. Substances that have the ability to reduce other substances are said to be reductive or reducing and are known as
Drying is a mass transfer process consisting of the removal of water or another solvent by evaporation from a solid, semi-solid or liquid. This process is used as a final production step before selling or packaging products. To be considered "dried", the final product must be solid, in the form of a continuous sheet, long pieces, particles or powder. A source of heat and an agent to remove the vapor produced by the process are involved. In bioproducts like food and pharmaceuticals like vaccines, the solvent to be removed is invariably water. Desiccation considered an extreme form of drying. In the most common case, a gas stream, e.g. air, applies the heat by convection and carries away the vapor as humidity. Other possibilities are vacuum drying, where heat is supplied by conduction or radiation, while the vapor thus produced is removed by the vacuum system. Another indirect technique is drum drying, where a heated surface is used to provide the energy, aspirators draw the vapor outside the room.
In contrast, the mechanical extraction of the solvent, e.g. water, by filtration or centrifugation, is not considered "drying" but rather "draining". In some products having a high initial moisture content, an initial linear reduction of the average product moisture content as a function of time may be observed for a limited time known as a "constant drying rate period". In this period, it is surface moisture outside individual particles, being removed; the drying rate during this period is dependent on the rate of heat transfer to the material being dried. Therefore, the maximum achievable drying rate is considered to be heat-transfer limited. If drying is continued, the slope of the curve, the drying rate, becomes less steep and tends to nearly horizontal at long times; the product moisture content is constant at the "equilibrium moisture content", where it is, in practice, in equilibrium with the dehydrating medium. In the falling-rate period, water migration from the product interior to the surface is by molecular diffusion, i,e. the water flux is proportional to the moisture content gradient.
This means that water moves from zones with higher moisture content to zones with lower values, a phenomenon explained by the second law of thermodynamics. If water removal is considerable, the products undergo shrinkage and deformation, except in a well-designed freeze-drying process; the drying rate in the falling-rate period is controlled by the rate of removal of moisture or solvent from the interior of the solid being dried and is referred to as being "mass-transfer limited". This is noticed in hygroscopic products such as fruits and vegetables, where drying occurs in the falling rate period with the constant drying rate period said to be negligible; the following are some general methods of drying: Application of hot air. Air heating accelerates drying, it reduces air relative humidity, further increasing the driving force for drying. In the falling rate period, as moisture content falls, the solids heat up and the higher temperatures speed up diffusion of water from the interior of the solid to the surface.
However, product quality considerations limit the applicable rise to air temperature. Excessively hot air can completely dehydrate the solid surface, so that its pores shrink and close, leading to crust formation or "case hardening", undesirable. For instance in wood drying, air is heated though some steam is added to it in order to avoid excessive surface dehydration and product deformation owing to high moisture gradients across timber thickness. Spray drying belongs in this category. Indirect or contact drying, as drum drying, vacuum drying. Again, higher wall temperatures will speed up drying but this is limited by product degradation or case-hardening. Drum drying belongs in this category. Dielectric drying is the focus of intense research nowadays, it may be used to assist air vacuum drying. Researchers have found that microwave finish drying speeds up the otherwise low drying rate at the end of the classical drying methods. Freeze drying or lyophilization is a drying method where the solvent is frozen prior to drying and is sublimed, i.e. passed to the gas phase directly from the solid phase, below the melting point of the solvent.
It is applied to dry foods, beyond its classical pharmaceutical or medical applications. It keeps biological properties of proteins, retains vitamins and bioactive compounds. Pressure can be reduced by a high vacuum pump. If using a vacuum pump, the vapor produced by sublimation is removed from the system by converting it into ice in a condenser, operating at low temperatures, outside the freeze drying chamber. Supercritical drying involves steam drying of products containing water; this process is feasible because water in the product is boiled off, joined with the drying medium, increasing its flow. It is employed in closed circuit and allows a proportion of latent heat to be recovered by recompression, a feature, not possible with conventional air drying, for instance; the process has potential for use in foods if carried out at reduced pressure