Chemistry for Kids: Elements - Magnesium Magnesium 1.738 grams per cm cubed Melting Point: Discovered by: Joseph Black in 1755. Isolated by Sir Humphry Davy in 1808. Magnesium is an alkaline earth metal and is the second element located in the second row of the periodic table. It is the eighth most abundant element on Earth. Magnesium atoms have 12 electrons and 12 protons. There are two valence electrons in the outer shell. Characteristics and Properties In standard conditions magnesium is a light metal with a silvery-white color. When exposed to air, magnesium will tarnish and become protected by a thin layer of oxide. When coming into contact with water, magnesium will react and produce hydrogen gas. If submerged in water, you will see gas bubbles start to form. Magnesium burns with a very bright white light. At one time magnesium powder was used to produce a bright flash for photography. Where is magnesium found on Earth? Magnesium is fairly abundant on Earth in compounds and is found in over 60 different minerals in the Earth's crust.

Some of the most important minerals include dolomite, magnesite, talc, and carnallite. The compound magnesium oxide (MgO) is the second most abundant compound in the Earth's crust making up around 35% of the crust by weight. A significant amount of Magnesium is also found dissolved in ocean water. In ocean water it takes the form of the cation Mg2+. A lot of commercial magnesium used in the United States comes from a process using electrolysis to extract it from sea water. How is magnesium used today? One of the main uses of magnesium metal is in metal alloys. This is because it is both strong and light. It is often mixed with aluminum, zinc, manganese, silicon, and copper to make strong and light alloys for use as automobile parts, aircraft components, and missiles. Magnesium metal is also used in electronic components. Its light weight and good electrical properties make it a good element for use in cameras, mobile phones, laptop computers, and other handheld electronic components. Another application of magnesium is in various compounds.

Some compounds are used as medicines such as magnesium hydroxide which used to help indigestion (Milk of Magnesia) and magnesium sulfate (Epsom salts) which is used in baths to soothe sore muscles. The human body needs magnesium for good health. It is used to make proteins, strong bones, and to regulate the body's temperature. How was it discovered? Scottish chemist Joseph Black first demonstrated in 1755 that the substance magnesia alba was a compound of different elements, one of them being magnesium. The element was first isolated by English chemist Sir Humphry Davy in 1808. Where did magnesium get its name? Magnesium gets its name from the district of Magnesia in Greece where the compound magnesium carbonate was first found. Isotopes Magnesium has three stable isotopes including magnesium-24, magnesium-25, and magnesium-26. Interesting Facts about Magnesium For many years magnesium was thought to be the same element as calcium. A magnesium fire is very difficult to put out as it can burn in nitrogen, carbon dioxide, and water. It is used in flares and fireworks because of its bright white light when it burns.

If you pour water on a magnesium fire, it will only make the fire worse. Magnesium is sometimes used to help shorten the length of migraine headaches. More on the Elements and the Periodic Table It's Elemental - The Element Magnesium It's Elemental Melting Point: 923 K (650°C or 1202°F) Boiling Point: 1363 K (1090°C or 1994°F) Density: 1.74 grams per cubic centimeter Phase at Room Temperature: Solid Element Classification: Metal Period Number: 3 Group Number: 2 Group Name: Alkaline Earth Metal What's in a name? For Magnesia, a district in the region of Thessaly, Greece. Say what? Magnesium is pronounced as mag-NEE-zhi-em. History and Uses: Although it is the eighth most abundant element in the universe and the seventh most abundant element in the earth's crust , magnesium is never found free in nature.

Magnesium was first isolated by Sir Humphry Davy, an English chemist, through the electrolysis of a mixture of magnesium oxide (MgO) and mercuric oxide (HgO) in 1808. Today, magnesium can be extracted from the minerals dolomite (CaCO3·MgCO3) and carnallite (KCl·MgCl2·6H2O), but is most often obtained from seawater. Every cubic kilometer of seawater contains about 1.3 billion kilograms of magnesium (12 billion pounds per cubic mile). Magnesium burns with a brilliant white light and is used in pyrotechnics, flares and photographic flashbulbs. Magnesium is the lightest metal that can be used to build things, although its use as a structural material is limited since it burns at relatively low temperatures. Magnesium is frequently alloyed with aluminum , which makes aluminum easier to roll, extrude and weld. Magnesium-aluminum alloys are used where strong, lightweight materials are required, such as in airplanes, missiles and rockets. Cameras, horseshoes, baseball catchers' masks and snowshoes are other items that are made from magnesium alloys.

Magnesium oxide (MgO), also known as magnesia, is the second most abundant compound in the earth's crust . Magnesium oxide is used in some antacids, in making crucibles and insulating materials, in refining some metals from their ores and in some types of cements. When combined with water (H2O), magnesia forms magnesium hydroxide (Mg(OH)2), better known as milk of magnesia, which is commonly used as an antacid and as a laxative. Hydrated magnesium sulphate (MgSO4·7H2O), better known as Epsom salt, was discovered in 1618 by a farmer in Epsom, England, when his cows refused to drink the water from a certain mineral well. He tasted the water and found that it tasted very bitter. He also noticed that it helped heal scratches and rashes on his skin. Epsom salt is still used today to treat minor skin abrasions. Other magnesium compounds include magnesium carbonate (MgCO3) and magnesium fluoride (MgF2). Magnesium carbonate is used to make some types of paints and inks and is added to table salt to prevent caking.

A thin film of magnesium fluoride is applied to optical lenses to help reduce glare and reflections. Estimated Crustal Abundance: 2.33×104 milligrams per kilogram Estimated Oceanic Abundance: 1.29×103 milligrams per liter Number of Stable Isotopes: 3 ( View all isotope data ) Ionization Energy: 7.646 eV alkali metal | chemical element | Britannica.com chemical element Alternative Titles: Group 1 element, Group Ia element Related Topics chemical element Alkali metal, any of the six chemical elements that make up Group 1 (Ia) of the periodic table —namely, lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). The alkali metals are so called because reaction with water forms alkalies (i.e., strong bases capable of neutralizing acids ). Sodium and potassium are the sixth and seventh most abundant of the elements, constituting , respectively, 2.6 and 2.4 percent of Earth’s crust .

The other alkali metals are considerably more rare, with rubidium, lithium, and cesium, respectively, forming 0.03, 0.007, and 0.0007 percent of Earth’s crust. Francium, a natural radioactive isotope , is very rare and was not discovered until 1939. Modern version of the periodic table of the elements. Encyclopædia Britannica, Inc. The alkali metals are so reactive that they are generally found in nature combined with other elements. Simple minerals , such as halite (sodium chloride, NaCl), sylvite (potassium chloride, KCl), and carnallite (a potassium-magnesium chloride, KCl · MgCl2· 6H2O), are soluble in water and therefore are easily extracted and purified. More complex, water-insoluble minerals are, however, far more abundant in Earth’s crust.

A very dilute gas of atomic sodium (about 1,000 atoms per cubic cm [about 16,000 atoms per cubic inch]) is produced in Earth’s mesosphere (altitude about 90 km [60 miles]) by ablation of meteors . Subsequent reaction of sodium with ozone and atomic oxygen produces excited sodium atoms that emit the light we see as the “tail” of a meteor as well as the more diffuse atmospheric nightglow. Smaller amounts of lithium and potassium are also present. The alkali metals have the silver-like lustre, high ductility, and excellent conductivity of electricity and heat generally associated with metals. Lithium is the lightest metallic element. The alkali metals have low melting points, ranging from a high of 179 °C (354 °F) for lithium to a low of 28.5 °C (83.3 °F) for cesium. Alloys of alkali metals exist that melt as low as −78 °C (−109 °F). Similar Topics beryllium (Be) The alkali metals react readily with atmospheric oxygen and water vapour.

(Lithium also reacts with nitrogen .) They react vigorously, and often violently, with water to release hydrogen and form strong caustic solutions. Most common nonmetallic substances such as halogens , halogen acids, sulfur , and phosphorus react with the alkali metals. The alkali metals themselves react with many organic compounds , particularly those containing a halogen or a readily replaceable hydrogen atom. Sodium is by far the most important alkali metal in terms of industrial use. The metal is employed in the reduction of organic compounds and in the preparation of many commercial compounds. As a free metal, it is used as a heat-transfer fluid in some nuclear reactors . Hundreds of thousands of tons of commercial compounds that contain sodium are used annually, including common salt (NaCl), baking soda (NaHCO3), sodium carbonate (Na2CO3), and caustic soda (NaOH). Potassium has considerably less use than sodium as a free metal. Potassium salts, however, are consumed in considerable tonnages in the manufacture of fertilizers . Lithium metal is used in certain light-metal alloys and as a reactant in organic syntheses. An important use of lithium is in the construction of lightweight batteries .

Primary lithium batteries (not rechargeable) are widely used in many devices such as cameras, cellular telephones, and pacemakers. Rechargeable lithium storage batteries that could be suitable for vehicle propulsion or energy storage are the subject of intensive research. Rubidium and cesium and their compounds have limited use, but cesium metal vapour is used in atomic clocks , which are so accurate that they are used as time standards. History Alkali metal salts were known to the ancients. The Old Testament refers to a salt called neter (sodium carbonate), which was extracted from the ash of vegetable matter. Saltpetre (potassium nitrate) was used in gunpowder , which was invented in China about the 9th century ad and had been introduced into Europe by the 13th century. Periodic Table of the Elements In October 1807 the English chemist Sir Humphry Davy isolated potassium and then sodium . The name sodium is derived from the Italian soda, a term applied in the Middle Ages to all alkalies; potassium comes from the French potasse, a name used for the residue left in the evaporation of aqueous solutions derived from wood ashes. Britannica Stories Ringling Bros.

Folds Its Tent Lithium was discovered by the Swedish chemist Johan August Arfwedson in 1817 while analyzing the mineral petalite. The name lithium is derived from lithos, the Greek word for “stony.” The element was not isolated in pure form until Davy produced a minute quantity by the electrolysis of lithium chloride. While the German chemists Robert Bunsen and Gustav Kirchhoff were investigating the mineral waters in the Palatinate in 1860, they obtained a filtrate that was characterized by two lines in the blue region of its spectrum (the light emitted when the sample was inserted into a flame ). They suggested the presence of a new alkali element and called it cesium , derived from the Latin caesius, used to designate the blue of the sky. The same researchers, on extracting the alkalies from the mineral lepidolite , separated another solution, which yielded two spectral lines of red colour. They proposed the name rubidium for the element in this solution from the Latin rubidus, which was used for the darkest red colour. Francium was not discovered until 1939 by Marguerite Perey of the Radium Institute in Paris.

In the 19th century the only use for the alkali metals was the employment of sodium as a reagent in the manufacture of aluminum . When the electrolytic process for aluminum purification was established, it appeared that large-scale use of sodium would cease. Subsequent improvements in the electrolytic production of sodium, however, reduced the cost of this element to such an extent that it can be employed economically to manufacture gasoline additives, reagents for chemical industry, herbicides , insecticides , nylon , pharmaceuticals, and reagents for metal refining. The continuous electrolysis of sodium hydroxide , a technique called the Castner process , was replaced in 1926 by the Downs cell process . This process, in which a molten sodium chloride–calcium chloride mixture (to reduce the melting point ) is electrolyzed, produces both sodium metal and chlorine . General properties of the group Connect with Britannica Eyjafjallajökull volcano The alkali metals have played an important role in quantum physics. Some alkali metal isotopes, such as rubidium-87, are bosons .

Dilute atomic gases of such alkali metal isotopes, confined by magnetic fields or “ laser mirrors” and cooled to temperatures near absolute zero , form Bose-Einstein condensates. In this state, the cluster of atoms is in a single quantum state and exhibits macroscopic behaviours normally seen only with atomic-sized particles. These include interference effects and coherent motion of the entire “cloud” of atoms. Chemical properties Since the alkali metals are the most electropositive (the least electronegative) of elements, they react with a great variety of nonmetals. In its chemical reactivity, lithium more closely resembles Group 2 (IIa) of the periodic table than it does the other metals of its own group. It is less reactive than the other alkali metals with water , oxygen , and halogens and more reactive with nitrogen , carbon , and hydrogen . Reactions with oxygen The alkali metals tend to form ionic solids in which the alkali metal has an oxidation number of +1. Therefore, neutral compounds with oxygen can be readily classified according to the nature of the oxygen species involved.

Ionic oxygen species include the oxide, O2-, peroxide , O22-, superoxide, O2-, and ozonide O3-. Compounds that can be prepared that contain an alkali metal, M, and oxygen are therefore the monoxide, M2O, peroxide, M2O2, superoxide, MO2, and ozonide, MO3. Rubidium and cesium and, possibly, potassium also form the sesquioxide, M4O6, which contains two peroxide anions and one superoxide anion per formula unit. Lithium forms only the monoxide and the peroxide. All the alkali metals react directly with oxygen; lithium and sodium form monoxides, Li2O and Na2O, and the heavier alkali metals form superoxides, MO2. The rate of reaction with oxygen, or with air, depends upon whether the metals are in the solid or liquid state , as well as upon the degree of mixing of the metals with the oxygen or air. In the liquid state, alkali metals can be ignited in air with ease, generating copious quantities of heat and a dense choking smoke of the oxide .

The free energy of formation (a measure of stability) of the alkali metal oxides at 25 °C (77 °F) varies widely from a high of −133 kcal/mole for lithium oxide to −63 kcal/mole for cesium oxide. The close approach of the small lithium ion to the oxygen atom results in the unusually high free energy of formation of the oxide. The peroxides (Li2O2and Na2O2) can be made by passing oxygen through a liquid- ammonia solution of the alkali metal, although sodium peroxide is made commercially by oxidation of sodium monoxide with oxygen. Sodium superoxide (NaO2) can be prepared with high oxygen pressures, whereas the superoxides of rubidium, potassium, and cesium can be prepared directly by combustion in air. By contrast, no superoxides have been isolated in pure form in the case of lithium or the alkaline-earth metals, although the heavier members of that group can be oxidized to the peroxide state. The cyanides of potassium, rubidium, and cesium, which are less stable than the lower oxides, can be prepared by the reaction of the superoxides with ozone.

Reactions with water The alkali metals all react violently with water according to M + H2O → MOH + 1/2 H2. The rate of the reaction depends on the degree of metal surface presented to the liquid. With small metal droplets or thin films of alkali metal, the reaction can be explosive. The rate of the reaction of water with the alkali metals increases with increasing atomic weight of the metal. With the heavier alkali metals, the hydroxides are highly soluble; thus, they are removed readily from the reacting surface, and the reaction can proceed with unabated vigour. The reaction involves equimolar mixtures (that is, equal numbers of atoms or molecules ) of the alkali metal and water to form a mole (an amount equal to that of the reactants) of alkali metal hydroxide and half a mole of hydrogen gas. These reactions are highly exothermic (give off heat), and the hydrogen that is generated can react with oxygen to increase further the heat that is generated. Reactions with nonmetals Of the alkali metals, only lithium reacts with nitrogen, and it forms a nitride (Li3N).

In this respect it is more similar to the alkaline-earth metals than to the Group 1 metals. Lithium also forms a relatively stable hydride , whereas the other alkali metals form hydrides that are more reactive. Lithium forms a carbide (Li2C2) similar to that of calcium . The other alkali metals do not form stable carbides, although they do react with the graphite form of carbon to give intercalation compounds (substances in which the metal atoms are inserted between layers of carbon atoms in the graphite structure). The alkali metals can be burned in atmospheres of the various halogens to form the corresponding halides . The reactions are highly exothermic , producing up to 235 kcal/mole for lithium fluoride. The alkali metals react with nonmetals in Groups 15 and 16 (Va and VIa) of the periodic table. Sulfides can be formed by the direct reaction of the alkali metals with elemental sulfur , furnishing a variety of sulfides. Phosphorus combines with the alkali metals to form phosphides with the general formula M3P.

Formation of alloys The characteristics of alloy behaviour in alkali metals can be evaluated in terms of the similarity of the elements participating in the alloy. Elements with similar atomic volumes form solid solutions (that is, mix completely in all proportions); some dissimilarity in atomic volumes results in eutectic-type systems (solutions formed over limited concentration ranges), and further dissimilarity results in totally immiscible systems. The high-pressure transition in potassium, rubidium, and cesium that converts these s-type metals to more transition metal-like d-type metals yields atomic volumes that are similar to those of many transition metals at the same pressure. This permits alloys or compounds to form between these alkali metals and such transition metals as nickel or iron. The elements potassium, rubidium, and cesium, which have rather similar atomic volumes and ionization energies, form complete solid solutions and mixed crystals. Sodium, which is a significantly smaller atom than potassium and has a higher ionization energy , tends to form eutectic systems with potassium, rubidium, and cesium. Even greater dissimilarity exists in the atomic volumes of sodium and lithium, resulting in insolubilities of the liquid phases.

The consolute temperature (the temperature at which the two liquids become completely miscible) increases on going from the lithium-sodium alloy system to the lithium-cesium system. Lithium and cesium can coexist as two separate liquid phases at temperatures up to at least 1,100 °C (2,000 °F). There is only one example of solid miscibility in alkali–alkaline-earth-metal binaries—the lithium-magnesium system, in which the two elements are very similar. Sodium forms compounds only with barium in the alkaline-earth-metal series. The heavier alkali metals all tend to form immiscible liquid phases with the alkaline earth metals. Several elements in Group 12 (IIb) of the periodic table ( zinc , cadmium , and mercury ) react with the alkali metals to form compounds. Mercury forms at least six compounds, commonly termed amalgams, with each of the five alkali metals, and with the exception of the amalgam with lithium, the highest melting point compound in each series has the formula MHg2. Lithium and sodium also form compounds with cadmium and zinc.

Formation of complexes Until the late 1960s there were few complexes of the alkali metal cations with organic molecules. Specialized biological molecules such as valinomycin were known to complex selectively the potassium cation K+ for transport across cell membranes , but synthetic ionophores (molecules that can form complexes with ions) were rare. All the alkali cations have a charge of +1 and, except for lithium , are chemically similar and rather inert. The only significant difference between one alkali cation and another is the size. The synthesis of crown ethers by American chemist Charles J. Pedersen in 1967 provided size-selective cyclic molecules consisting of ether oxygens forming a ring or “crown” that could complex a cation of the right size to fit into the hole in the centre of the molecule. In some cases two crown ether molecules can encapsulate a cation in a “sandwich” fashion.

For example, K+ just fits into the centre of an 18-crown-6 ring (18 atoms in the ring, 12 of which are carbon atoms and 6 are ether oxygen atoms) to form a 1:1 complex (that is, 1 cation:1 crown ether), K+(18C6). Cs+ is too large to fit into the ring but can be complexed on one side to form the Cs+(18C6) complex or can be sandwiched between two 18-crown-6 molecules to form the 1:2 complex, Cs+(18C6)2. Thus, the selectivity of a crown ether for a particular cation depends on the ring size. Common crown ethers are 12-crown-4, 15-crown-5, and 18-crown-6. These molecules are selective for Li+, Na+, and K+, respectively. Even greater affinity for alkali cations was achieved by the synthesis of cryptands by French chemist Jean-Marie Lehn in 1968 and spherands by American chemist Donald Cram in 1979.

These are three-dimensional molecules with an internal cavity or crypt that can completely encapsulate the alkali cation. By synthesizing molecules with different cavity sizes, the selectivity for particular cations over those of the “wrong” size to fit in the cavity can be controlled. It should be noted, however, that these molecules are not rigid and that flexibility of the framework can alter the cavity size to accommodate alkali cations of different sizes, although with differences in the strength of complexation. Since the initial syntheses of crown ethers and cryptands, thousands of complexants for cations of various sizes, charges, and geometries have been synthesized. This has led to an entirely new branch of chemistry called supramolecular chemistry. Analytical chemistry of the alkali metals Classical methods of separation and analysis of alkali metals are rather difficult and time consuming. For lithium they include such procedures as selective extraction of lithium chloride into organic solvents and the detection of lithium with azo dyes that give highly sensitive colour reactions in alkaline solutions. A modification of the uranyl acetate test (the precipitation of an insoluble sodium salt with uranyl acetate) has been used as a standard test for the presence of sodium.

The use of a cobaltinitrite solution permits separation of potassium from sodium by precipitation of the insoluble potassium salt. There are essentially no satisfactory analytical methods for rubidium and cesium based on the use of reagents in solution. Classical methods of separation of the alkali metals have been largely supplanted by chromatographic elution. Strongly acidic cation -exchange resins and aqueous acidic solutions are used. Generally the affinity increases with atomic weight so that the ions are eluted in the order Fr+ > Cs+ > Rb+ > K+ > Na+ > Li+, which is the order of decreasing size of the hydrated ions. Ion-exchange resins that are specific for lithium have been developed. Macrocyclic compounds such as crown ethers and cryptands that are selective for particular alkali metal ions have been synthesized. They form cationic complexes that can be dissolved in organic solvents such as chloroform (CHCl3) with counterions such as picrate (C6H2[NO2]3O-).

The characteristic flame colours of the alkali metals (red, yellow, violet, red, and blue for Li, Na, K, Rb, and Cs, respectively) are qualitative indicators of the modern analytical methods used to determine the concentrations of alkali-metal salts in aqueous solution. The intensities of the characteristic spectral lines in emission after excitation by a flame or ICP (inductively coupled plasma) give quantitative measures of the individual alkali metal concentration in the parts per million range or lower. Determination of one alkali metal in the presence of another, however, can result in interference, which can be reduced by using specially prepared standard solutions that contain known amounts of the interfering metals. The analysis of the alkali-metal samples for the presence of nonmetallic elements, such as oxygen , carbon , hydrogen , and nitrogen , requires specialized techniques. The oxygen content of sodium and potassium samples can be determined by extraction of the free alkali metal with mercury , leaving behind mercury-insoluble oxides and carbonates , which can subsequently be analyzed by means of solution methods. The oxygen content of rubidium and cesium can be accurately determined by precise measurement of the freezing point of these two elements.

The carbon content of alkali metals can be analyzed by oxidation of the alkali metal in pure oxygen, followed by infrared measurement of the carbon dioxide generated during combustion. For the analysis of nitride in lithium, the nitride commonly is converted to ammonia , and the ammonia is measured by colorimetric analysis. Chemistry for Kids: Elements - Sodium Sodium 0.968 grams per cm cubed Melting Point: Discovered by: Sir Humphry Davy in 1807 Sodium is an alkali metal located in the first group or column of the periodic table. The sodium atom has 11 electrons and 11 protons with one valence electron in the outer shell. Characteristics and Properties Sodium in its pure form is very reactive. It is a very soft metal that can be easily cut with a knife. It is silvery-white in color and burns with a yellow flame. Sodium will float on water, but it will also react violently when coming into contact with water. When sodium reacts with water it produces sodium hydroxide and hydrogen gas. Sodium is most famous for its many useful compounds such as table salt (NaCl), sodium nitrate (Na2CO3), and baking soda (NaHCO3).

Many of the compounds that sodium forms are water-soluble, meaning they dissolve in water. Where is sodium found on Earth? Sodium is the sixth most abundant element on Earth. It is never found in its pure form because it is so reactive. It is only found in compounds such as sodium chloride (NaCL) or table salt. Sodium chloride is found in ocean water (salt water), salt lakes, and underground deposits. Pure sodium can be recovered from sodium chloride through electrolysis. How is sodium used today? Sodium is primarily used in the form of compounds with other elements. The average person uses sodium every day in the form of table salt in their food. Table salt is the compound sodium chloride (NaCl). Table salt is needed for animals to survive, but most people use it for adding flavor to their food. Another popular use of sodium is in baking soda which is the chemical compound sodium bicarbonate. Baking soda is used as a leavening agent in cooking foods like pancakes, cakes, and breads. A lot of soaps are forms of sodium salts. Sodium hydroxide is a key ingredient when making soaps. Other applications include de-icing, medicine, organic chemistry, street lights, and cooling nuclear reactors.

How was it discovered? Sodium was discovered by English chemist Sir Humphry Davy in 1807. He isolated sodium by applying electrolysis to caustic soda. Where did sodium get its name? Sodium gets its name from the English word soda. This is because Sir Humphry Davy used caustic soda when isolating the element. The symbol Na comes from the Latin word natrium. Isotopes Only one of the 20 known isotopes of sodium is stable, sodium-23. Interesting Facts about Sodium Sir Humphry Davy discovered sodium only a few days after he discovered potassium. Sodium comprises around 2.6% of the Earth's crust. It helps to maintain the proper fluid balance in the body's cells and also helps us to digest our food. Our bodies lose sodium when we sweat. However, most people eat far more sodium than their bodies actually need. If the body runs low on sodium, it can cause the muscles to cramp. Sodium is considered non-toxic, but too much of it can cause high blood pressure.

More on the Elements and the Periodic Table Sodium - English to English Translation Sodium Hepatitis Central (TM) Liver Disease Medical Glossary Essential mineral that our bodies regulate and conserve. Excess sodium retention increases the fluid volume (edema) and low sodium leads to less fluid and relative dehydration. The adult body averages a total content of over 100 grams of sodium, of which one-third is in bone. A small amount of sodium does get into cell interiors, but this represents only about 10% of the body content. The remaining 57 % or so of the body sodium content is in the fluid immediately surrounding the cells, where it is the major cation (positive ion). The role of sodium in the extracellular fluid is maintaining osmotic equilibrium (the proper difference in ions dissolved in the fluids inside and outside the cell) and extracellular fluid volume. Also involved in nerve impulse transmission, muscle tone and nutrient transport. Interelated to potassim.

Increase in serum sodium is seen in conditions with water loss in excess of salt loss, as in profuse sweating, severe diarrhea or vomiting, polyuria (as in diabetes mellitus or insipidus), hypergluco ormineralocorticoidism, and inadequate water intake. Drugs causing elevated sodium include steroids with mineral ocorticoid activity, carbenoxolone, diazoxide, guanethidine, licorice, methyldopa, oxyphenbutazone, sodium bicarbonate, methoxyflurane, and reserpine. Decrease in sodium is seen in states characterized by intake of freewater or hypotonic solutions, as may occur in fluid replacement following sweating, diarrhea, vomiting, and diuretic abuse. Dilutional hyponatremia may occur in cardiac failure, liver failure, nephrotic syndrome, malnutrition, and SIADH. There are many other causes of hyponatremia, mostly related to corticosteroid metabolic defects or renal tubular abnormalities.

Drugs other than diuretics may cause hyponatremia, including ammonium chloride, chlorpropamide, heparin, amino glutethimide, vasopressin, cyclophosphamide, and vincristine hEnglish - advanced version sodium \so"di*um\ (?), n. [nl., fr.e. soda.] (chem.) a common metallic element of the alkali group, in nature always occuring combined, as in common salt, in albite, etc. it is isolated as a soft, waxy, white, unstable metal, so readily oxidized that it combines violently with water, and to be preserved must be kept under petroleum or some similar liquid. sodium is used combined in many salts, in the free state as a reducer, and as a means of obtaining other metals (as magnesium and aluminium) is an important commercial product. symbol na (natrium). atomic weight 23. specific gravity 0.97. sodium amalgam, an alloy of sodium and mercury, usually produced as a gray metallic crystalline substance, which is used as a reducing agent, and otherwise.

sodium bicarbonate, a white crystalline substance, hnaco3, with a slight alkaline taste resembling that of sodium carbonate. it is found in many mineral springs and also produced artificially,. it is used in cookery, in baking powders, and as a source of carbonic acid gas (carbon dioxide) for soda water. called also cooking soda , saleratus, and technically, acid sodium carbonate , primary sodium carbonate, sodium dicarbonate , etc. sodium carbonate, a white crystalline substance, na2co3.10h2o , having a cooling alkaline taste, found in the ashes of many plants, and produced artifically in large quantities from common salt. it is used in making soap, glass, paper, etc., and as alkaline agent in many chemical industries. called also sal soda, washing soda , or soda. cf. sodium bicarbonate, above and trona. sodium chloride, common, or table, salt, nacl. sodium hydroxide, a white opaque brittle solid, naoh, having a fibrous structure, produced by the action of quicklime, or of calcium hydrate (milk of lime), on sodium carbonate.

it is a strong alkali, and is used in the manufacture of soap, in making wood pulp for paper, etc. called also sodium hydrate, and caustic soda. by extension, a solution of sodium hydroxide. sodium n : a silvery soft waxy metallic element of the alkali metal group; occurs abundantly in natural compounds (especially in salt water); burns with a yellow flame and reacts violently in water; occurs in sea water and in the mineral halite (rock salt) [syn: na, atomic number 11] sodium symbol: na atomic number: 11 atomic weight: 22.9898 soft silvery reactive element belonging to group 1 of the periodic table (alkali metals). it is highly reactive, oxidizing in air and reacting violently with water, forcing it to be kept under oil. it was first isolated by humphrey davy in 1807. -- elements Magnesium, Chemical Element - water, uses, elements, examples, metal, gas, number, name PRONUNCIATION mag-NEE-zee-um Large amounts of magnesium are used to make alloys. An alloy is made by melting or mixing two or more metals.

The mixture has properties different from those of the individual metals. Magnesium alloys are quite light, yet very strong. This property makes them useful in the construction of airplanes and space-craft. About 70 percent of the magnesium compounds produced in the United States are used in the manufacture of refractory materials. A refractory material is one that can withstand very high temperatures by reflecting heat. Refractory materials are used to line the ovens that maintain high temperatures. The remaining 30 percent of magnesium compounds are used in agriculture, construction, industrial, and chemical operations. Discovery and naming Compounds of magnesium are very abundant in the Earth. Dolomite, or calcium magnesium carbonate (CaMg(CO 3 2 , is an example. Dolomite has been used as a building material for centuries. Another well-known magnesium compound is Epsom salts, or magnesium sulfate (MgSO 4 . Epsom salts are known for their soothing qualities, most notably when added to a bath. (See accompanying sidebar.) The allure of Epsom salts P erhaps the best know magnesium compound is magnesium sulfate (MgSO 4 ). It is popularly known as Epson salts.

One of the earliest stories about Epsom salts dates back to 1618. The town of Epsom, in Surrey, England, was suffering from a severe drought. A farmer named Henry Wicker brought his cattle to drink from a water hole on the town commons (central park). But the cattle would not drink the water. Wicker was surprised because he knew they were very thirsty. He tasted the water himself and found that it was very bitter. The bitterness was due to magnesium sulfate in the water. This compound became known as Epsom salts. People soon learned that soaking in the natural waters that contained Epsom salts made them feel better. The salts seemed to have properties that soothed the body. Before long, soaking in these waters became very popular. Today, Epsom salts are used in bath water. They relax sore muscles and remove rough skin. Many people believe the salts have the same relaxing effect as hot springs. Some gardeners even believe that sprinkling Epsom salts in the garden helps flowers and vegetables grow! Careful studies of magnesium and its compounds began in the middle 1700s.

Scottish physician and chemist Joseph Black (1728-99) carried out some of the earliest experiments on magnesium compounds. He reported on his research in an article that became famous. Black is sometimes given credit for "discovering" magnesium because of his work with the element. By 1800, chemists knew that magnesium was an element. But no one had been able to prepare pure magnesium metal. Magnesium holds very tightly to other elements in its compounds. No one had found a way to break the bonds between magnesium and these other elements. In 1808, English chemist Humphry Davy (1778-1829) solved the problem by passing an electric current through molten (melted) magnesium oxide (MgO). The current caused the compound to break apart, forming magnesium metal and oxygen gas: Davy used this method to discover a number of other elements. (See sidebar on Davy in the calcium entry in Volume 1.) Like magnesium, these elements form compounds that are very difficult to break apart. An electric current provides the energy to break these compounds down into their elements. The name magnesium goes back many centuries.

It was selected in honor of a region in Greece known as Magnesia. The region contains large supplies of magnesium compounds. Physical properties Magnesium is a moderately hard, silvery-white metal. It is the lightest of all structural metals. These metals are strong enough to be used to build buildings, bridges, automobiles, and airplanes. Magnesium is easily fabricated. Fabrication means shaping, molding, bending, cutting, and working with a metal. Metals must be fabricated before they can be turned into useful products. Metals that are strong, tough, or hard are not easily fabricated. They must be converted to an alloy. A metal that is more easily fabricated (such as magnesium) is combined with them. The melting point of magnesium is 651°C (1,200°F) and its boiling point is 1,100°C (2,000°F). Its density is 1.738 grams per cubic centimeter. At high temperatures, magnesium bums with a blinding white light. Chemical properties Magnesium is a fairly active metal. It reacts slowly with cold water and more rapidly with hot water. It combines with oxygen at room temperature to form a thin skin of magnesium oxide.

It burns with a blinding white light at higher temperatures. Magnesium reacts with most acids and with some alkalis. An alkali is a chemical with properties opposite those of an acid. Sodium hydroxide (common lye such as Drano) and lime-water are examples of alkalis. Magnesium also combines easily with many non-metals, including nitrogen, sulfur, phosphorus, chlorine, fluorine, bromine, and iodine. It also reacts readily with a number of compounds, such as carbon monoxide (CO), carbon dioxide (CO 2 , sulfur dioxide (SO 2 , and nitric oxide (NO). Occurrence in nature The abundance of magnesium in the Earth's crust is estimated to be about 2.1 percent. That makes it the sixth most common element in the earth. It also occurs in seawater. A cubic mile of seawater is estimated to contain up to six million tons of magnesium. There are many naturally occurring minerals of magnesium. Some of the most important are dolomite; magnesite, or magnesium carbonate (MgCO 3 ; carnallite, or potassium magnesium chloride (KMgCl 3 ; and epsomite, or magnesium sulfate (MgSO 4 .

The largest producer of magnesium ores is Turkey. Other large producers are North Korea, China, Slovakia, Austria, and Russia. The amount of magnesium produced in the United States is not reported in order to protect trade secrets. Magnesium produced in the United States comes from three sources: seawater, brine, and mines. Seawater is processed to obtain magnesium by companies in California, Delaware, Florida, and Texas. Magnesium is obtained from brine in Michigan and Utah. Brine is water that is even saltier than seawater. Finally, some magnesium compounds are taken from mines in Nevada, North Carolina, and Washington. Isotopes There are three naturally occurring isotopes of magnesium: magnesium-24, magnesium-25, and magnesium-26. Isotopes are two or more forms of an element. Isotopes differ from each other according to their mass number. The number written to the right of the element's name is the mass number. The mass number represents the number of protons plus neutrons in the nucleus of an atom of the element. The number of protons determines the element, but the number of neutrons in the atom of any one element can vary. Each variation is an isotope.

One radioactive isotope of magnesium, magnesium-28, also exists. A radioactive isotope is one that breaks apart and gives off some form of radiation. Radioactive isotopes are produced when very small particles are fired at atoms. These particles stick in the atoms and make them radioactive. Magnesium-28 has no important commercial uses. Extraction Magnesium is prepared by one of two methods. The first method is similar to the method used by Davy in 1808. An electric current is passed through molten (melted) magnesium chloride: Magnesium alloys are often used in the production of skis. The second method involves reacting magnesium oxide with ferrosilicon. Ferrosilicon is an alloy of iron and silicon. When magnesium oxide and ferrosilicon react, free magnesium metal is formed. Uses Although most cameras now use electronic flashes, magnesium metal is often contained in cameras that use flash bulbs. A thin strip of magnesium metal is inside the bulb. When the Magnesium alloys are used in power tools, such as the reciprocating saw shown here. flash is ignited, the magnesium strip catches fire. It burns with a very bright white light. The light from the bulb illuminates a scene for the photograph.

A common use of magnesium metal is in fireworks. Most firework displays include some brilliant flashes of very white light. Those flashes are produced by the burning of magnesium metal. Magnesium is commonly alloyed with other metals. Magnesium and aluminum, for instance, are two metals that combine to form alloys that are very strong and resistant to corrosion (rust). But they weigh much less than steel alloys with similar properties. Strength and low density are important properties in the manufacture of airplanes, automobiles, metal luggage, ladders, shovels and other gardening equipment, racing bikes, skis, race cars, cameras, and power tools. A typical magnesium alloy contains about 90 percent magnesium, 2 to 9 percent aluminum, and small amounts of zinc and manganese. Compounds The largest single use of magnesium compounds is in refractories.

Other magnesium compounds are used in the following categories: medicine: pain killer and fever reducer (magnesium acetylsalicylate); antacid to neutralize stomach acid (magnesium hydroxide; magnesium phosphate; magnesium silicate); laxative to loosen the bowels (magnesium carbonate; magnesium chloride; magnesium citrate; magnesium hydroxide; magnesium lactate; magnesium phosphate); antiseptic to kill germs (magnesium borate; magnesium salicylate; magnesium sulfate); sedative to help one get sleep (magnesium bromide) production of glass and ceramics: magnesium fluoride; magnesium oxide mothproofing of textiles: magnesium hexafluorosilicate fireproofing wood for construction: magnesium phosphate manufacture of paper: magnesium sulfite Health effects Magnesium is essential for good health in both plants and animals. It forms part of the chlorophyll molecule found in all green plants. Chlorophyll is the molecule in green plants that controls the conversion of carbon dioxide and water to carbohydrates, such as starch and sugar. Plants that do not get enough magnesium cannot make enough chlorophyll. Their leaves develop yellowish blotches as a result. Magnesium is found in many enzymes in both plants and animals.

An enzyme is a catalyst in a living organism. It speeds up the rate at which certain changes take place in the body. Enzymes are essential in order for living cells to function properly. It is difficult not to get enough magnesium in one's daily diet. It is found in nuts, cereals, seafoods, and green vegetables. Most people have no problem getting the 300 to 400 milligrams of magnesium recommended in the daily diet. Magnesium forms part of the chlorophyll molecule found in all green plants. A lack of magnesium can occur, however. For example, alcoholics and children in poor countries sometimes develop a magnesium deficiency. In such cases, magnesium deficiency may cause a person to become easily upset or overly aggressive. On the other hand, it is also possible to be exposed to too much magnesium. For example, inhaling magnesium powder can produce irritation of the throat and eyes, resulting in a fever. In large doses, magnesium can cause damage to muscles and nerves. It can eventually result in loss of feeling and paralysis (inability to move parts of the body). Such conditions are rare. They are likely to occur only among people who have to work with magnesium metal on a regular basis.

Also read article about Magnesium from Wikipedia User Contributions:

Question: Which English chemist isolated the metals sodium and magnesium, etc?
Answer: Sir Humphry Davy