mentioj any two ways for preparation of oxide
Answers
Answer:
Explanation:
Oxides can be prepared in the form of single crystals or polycrystalline or amorphous samples. For studies where surface atomic arrangements need to be known, single crystal samples are used from which well-ordered crystallographic surface planes can be prepared. There are many well-established methods to grow single crystals. The most common method is growth from the melt. In this method, very high temperatures are required, and the oxygen partial pressure above the melt must be carefully controlled to ensure that the crystal attains the desired stoichiometry. When the melting point of the oxide is so high that the oxide decomposes, other lower temperature methods may be employed, such as vapor transport. For most practical purposes where a large specific surface area is essential, polycrystalline or amorphous samples are desirable. Experience has shown that, generally, low temperature processes are necessary to obtain oxides of large surface areas or small particle sizes. However, methods can be developed, such as in the preparation of aerogels, in which high temperatures are used. A low temperature treatment does not necessarily lead to the formation of the thermodynamically most stable phase. Instead, depending on the details of the preparation procedure, metastable phases and/or amorphous samples may be obtained. A low temperature treatment may also lead to the formation of hydrous oxides from the precursors obtained by aqueous precipitation.
Answer:
Due to its electronegativity, oxygen forms stable chemical bonds with almost all elements to give the corresponding oxides. Noble metals (such as gold or platinum) are prized because they resist direct chemical combination with oxygen, and substances like gold(III) oxide must be generated by indirect routes.
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 (or simply air), some elements— sodium—react rapidly, to give the hydroxides. In part, for this reason, alkali and alkaline earth metals are not found in nature in their metallic, i.e., native, form. Cesium is so reactive with oxygen that it is used as a getter in vacuum tubes, and solutions of potassium and sodium, so-called NaK are used to deoxygenate and dehydrate some organic solvents. 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 thin film of aluminium oxide that passivates the metal, slowing further corrosion. The aluminum oxide layer can be built to greater thickness by the process of electrolytic anodizing. Though solid magnesium and aluminum 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 dangerously explosive in air. Consequently, they are often used in solid-fuel rockets.
Oxides, such as iron(III) oxide or rust, which consists of hydrated iron(III) oxides Fe2O3·nH2O and iron(III) oxide-hydroxide (FeO(OH), Fe(OH)3), form when oxygen combines with other elements
In dry oxygen, iron readily forms iron(II) oxide, but the formation of the hydrated ferric oxides, Fe2O3−x(OH)2x, that mainly comprise rust, typically requires oxygen and water. Free oxygen production by photosynthetic bacteria some 3.5 billion years ago precipitated iron out of solution in the oceans as Fe2O3 in the economically important iron ore hematite.
Structure
Oxides have a range of different structures, from individual molecules to polymeric and crystalline structures. At standard conditions, oxides may range from solids to gases.
Oxides of metals
Oxides of most metals adopt polymeric structures.[3] The oxide typically links three metal atoms (e.g., rutile structure) or six metal atoms (carborundum or rock salt structures). Because the M-O bonds are typically strong and these compounds are crosslinked polymers, the solids tend to be insoluble in solvents, though they are attacked by acids and bases. The formulas are often deceptively simple where many are nonstoichiometric compounds.[2]
Molecular oxides
Some important gaseous oxides
Carbon dioxide is the main product of fossil fuel combustion.
Carbon monoxide is the product of the incomplete combustion of carbon-based fuels and a precursor to many useful chemicals.
Nitrogen dioxide is a problematic pollutant from internal combustion engines.
Sulfur dioxide, the principal oxide of sulfur, is emitted from volcanoes.
Nitrous oxide ("laughing gas") is a potent greenhouse gas produced by soil bacteria.
Although most metal oxides are polymeric, some oxides are molecules. Examples of molecular oxides are carbon dioxide and carbon monoxide. All simple oxides of nitrogen are molecular, e.g., NO, N2O, NO2 and N2O4. Phosphorus pentoxide is a more complex molecular oxide with a deceptive name, the real formula being P4O10. Some polymeric oxides depolymerize when heated to give molecules, examples being selenium dioxide and sulfur trioxide. Tetroxides are rare. The more common examples: ruthenium tetroxide, osmium tetroxide, and xenon tetroxide.
Many oxyanions are known, such as polyphosphates and polyoxometalates. Oxycations are rarer, some examples being nitrosonium (NO+), vanadyl (VO2+), and uranyl (UO2+
2). Of course many compounds are known with both oxides and other groups. In organic chemistry, these include ketones and many related carbonyl compounds. For the transition metals, many oxo complexes are known as well as oxyhalides.