balanced chemical equation for magnesium oxide
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Magnesium oxide, or magnesia, is a white solid mineral that occurs naturally as periclase and is a source of magnesium (see also oxide). It has an empirical formula of . It is formed by an ionic bond between one magnesium and one oxygen atom. Magnesium oxide is hygroscopic in nature and care must be taken to protect it from moisture. Magnesium hydroxide forms in the presence of water (MgO + H2O â†’ Mg(OH)2), but it can be reversed by heating it to separate moisture.
Magnesium oxide was historically known as magnesia alba (literally, the white mineral from Magnesia), to differentiate it from magnesia negra, a black mineral containing what is now known asmanganese.
A refractory material is one that is physically and chemically stable at high temperatures. "By far the largest consumer of magnesia worldwide is the refractory industry, which consumed about 56% of the magnesia in the United States in 2004, the remaining 44% being used in agricultural, chemical, construction, environmental, and other industrial applications."
MgO is one of the raw materials for making Portland cement in dry process plants. If too much MgO is added, the cement may become expansive. Production of MgO-based cement using serpentinite and waste CO2 (as opposed to conventional CaO-based cement using fossil fuels) may reduce anthropogenic emissions of CO2 .
MgO is an efficient moisture absorbent used by many libraries for preserving books.
In medicine, magnesium oxide is used for relief of heartburn and sore stomach, as an antacid, magnesium supplement, and as a short-term laxative. It is also used to improve symptoms of indigestion. Side effects of magnesium oxide may include nausea and cramping. In quantities sufficient to obtain a laxative effect, side effects of long-term use include enteroliths resulting in bowel obstruction.
- MgO is used as an insulator in industrial cables, as a basic refractory material for crucibles and as a principal fireproofing ingredient in construction materials. As a construction material, magnesium oxide wallboards have several attractive characteristics: fire resistance, moisture resistance, mold and mildew resistance, and strength.
- It is used as a reference white color in colorimetry, owing to its good diffusing and reflectivity properties. It may be smoked onto the surface of an opaque material to form an integrating sphere.
- It is used extensively in electrical heating as a component of "CalRod"-styled heating elements. There are several mesh sizes available and most commonly used ones are 40 and 80 mesh per the American Foundry Society. The extensive use is due to its high dielectric strength and average thermal conductivity. MgO is usually crushed and compacted with minimal airgaps or voids. The electrical heating industry also experimented with aluminium oxide, but it is not used anymore.
- Pressed MgO is used as an optical material. It is transparent from 0.3 to 7 Âµm. The refractive index is 1.72 at 1 Âµm and the Abbe number is 53.58. It is sometimes known by the Eastman Kodak trademarked name Irtran-5, although this designation is long since obsolete. Crystalline pure MgO is available commercially and has small use in infrared optics.
- It is packed around transuranic waste at the Waste Isolation Pilot Plant, to control the solubility of radionuclides.
- An aerosolized solution of MgO is used in library science and collections management for the deacidification of at-risk paper items. In this process, the alkalinity of MgO (and similar compounds) neutralizes the relatively high acidity characteristic of low-quality paper, thus slowing the rate of deterioration.
- It is also used as a protective coating in plasma displays.
Magnesium oxide is easily made by burning magnesium ribbon which oxidizes in a bright white light, resulting in a powder. However, the bright flame is very hard to extinguish and it emits a harmful intensity of UV light. Inhalation of magnesium oxide fumes can cause metal fume fever.
oxidation and reduction complementary chemical reactions characterized by the loss or gain, respectively, of one or more electrons by an atom or molecule. Originally the term oxidation was used to refer to a reaction in which oxygen combined with an element or compound, e.g., the reaction of magnesium with oxygen to form magnesium oxide or the combination of carbon monoxide with oxygen to form carbon dioxide. Similarly, reduction referred to a decrease in the amount of oxygen in a substance or its complete removal, e.g., the reaction of cupric oxide and hydrogen to form copper and water. When an atom or molecule combines with oxygen, it tends to give up electrons to the oxygen in forming a chemical bond . Similarly, when it loses oxygen, it tends to gain electrons. Such changes are now described in terms of changes in the oxidation number, or oxidation state, of the atom or molecule (see valence ). Thus oxidation has come to be defined as a loss of electrons or an increase in oxidation number, while reduction is defined as a gain of electrons or a decrease in oxidation number, whether or not oxygen itself is actually involved in the reaction. In the formation of magnesium oxide from magnesium and oxygen, the magnesium atoms have lost two electrons, or the oxidation number has increased from zero to +2. This is also true when magnesium reacts with chlorine to form magnesium chloride. In solution, ferrous iron (oxidation number +2) may be oxidized to ferric iron (oxidation number +3) by the loss of an electron. In the reduction of cupric oxide the oxidation number of copper has changed from +2 to zero by the gain of two electrons. The two processes, oxidation and reduction, occur simultaneously and in chemically equivalent quantities. In the formation of magnesium chloride, for every magnesium atom oxidized by a loss of two electrons, two chlorine atoms are reduced by a gain of one electron each. Oxidation-reduction reactions, called also redox reactions, are most simply balanced in the form of chemical equations by arranging the quantities of the substances involved so that the number of electrons lost by one substance is equaled by the number gained by another substance. In such reactions, the substance losing electrons (undergoing oxidation) is said to be an electron donor, or reductant, since its lost electrons are given to and reduce the other substance. Conversely, the substance that is gaining electrons (undergoing reduction) is said to be an electron acceptor, or oxidant. Common reductants (substances readily oxidized) are the active metals, hydrogen, hydrogen sulfide, carbon, carbon monoxide, and sulfurous acid. Common oxidants (substances readily reduced) include the halogens (especially fluorine and chlorine), oxygen, ozone, potassium permanganate, potassium dichromate, nitric acid, and concentrated sulfuric acid. Some substances are capable of acting either as reductants or as oxidants, e.g., hydrogen peroxide and nitrous acid. The corrosion of metals is a naturally occurring redox reaction. Industrially, many redox reactions are of great importance: combustion of fuels; electrolysis (oxidation occurs at the anode and reduction at the cathode); and metallurgical processes in which free metals are obtained from their ores.
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Answers:1. Alumin + Hydochloric Acid 2Al + 6HCl = 2AlCl3 + 3H2 [Single displacement] 2. Calcium hydroxide + nitric acid Ca(OH)2 + 2HNO3 = Ca(NO3)2 + 2H2O [Double displacement] 3. Potassium chlorate heated 2KClO3 = 2KCl + 3O2 [decomposition] 4. Magnesium + sulfur (S8) 4Mg + S = Mg4S [synthesis] 5. Ammonium phosphate + alumin Chloride (NH4)3PO4 + AlCl3 = AlPO4 + 3NH4Cl [double replacement] 6. Calcium oxide + water CaO + H2O = Ca(OH)2 [Synthesis] 7. Chromium + water Cr + 6H2O = Cr(H2O)6 [synthesis] 8. Tin + mercury(I) nitrate Sn + 2HgNO3 = Sn(NO3)2 + 2Hg [Single displacement] 9. Sodium Bromide + silver nitrate NaBr + AgNO3 = NaNO3 + AgBr [double replacement] 10. hydrogen + oxygen 2H2 + O2 = 2H2O [Synthesis] 11. hydogen peroxide H2O2 = H2 + O2 [decomposition] 12. fluorine + potasium bromide F + KBr = FBr + K [single displacement] 13. Carbon dioxide + water CO2 + H2O = H2CO3 [Synthesis] 14. calcium chloride + ammonium hydroxide CaCl2 + 2NH4OH = Ca(OH)2 + 2NH4Cl [double replacement] 15. Sodium + chlorine 2Na + Cl2 = 2NaCl [synthesis] 16. bromine + sodium chloride Br2 + 2NaCl = 2NaBr + Cl2 [single replacement] 17. mercury (II) oxide heated 2HgO = 2Hg + O2 [decomposition] 18. potassium + water 2K + 2H2O = 2KOH + H2 [SINGLE replacement] 19. strontium carbonate + nitric acid SrCO3 + 2HNO3 = Sr(NO3)2 + H2CO3 [double replacement] 20. potassium iodide + lead nitrate 2KI + Pb(NO3)2 = 2KNO3 + PbI2 [double replacement] :) whoo! that's a lot.
Answers:To balance an equation, you just need to play around with it. There is no trick. However, it is best to leave the single elements (like oxygen gas, O2) for last. Note that acids are aqueous. a. 2ZnS(s) + 3O2(g) --> 2ZnO(s) + 2SO2(g) b. 2HCl(aq) + Mg(OH)2(aq) --> MgCl2(aq) + 2H2O(l) c. 2HNO3(aq) + Ca(OH)2(aq) --> Ca(NO3)2(aq) + 2H2O(l) - Anonymous
Answers:The equation for the reactants is (this I am sure of): Mg(C2H3O2)2 + Li2CO3 I do not know for sure what the products are of this reaction, nor can I find it anywhere in reference. Unfortunately, all I can offer is my best guess, which is as follows: Mg(C2H3O2)2 + Li2CO3 --> H2O + CO2 + MgO + Li. If those were indeed the correct products, the balanced equation would be: 2Mg(C2H3O2) + 9Li2CO3 --> 3H2O + 13CO2 + 2MgO + 18Li Hope that is somewhat helpful!
Answers:3) CaO(s) + H2O(l) ==> Ca(OH)2(aq) This is a combination reaction. 4) CO2(g) + H2O(l) ==> H2CO3(aq) ...another combination reaction. 5) NiO(s) + 2HNO3(aq) ==> H2O(l) + Ni(NO3)2(aq) This is a double replacement reaction.