Magnesium

12

sodium ← magnesium → aluminium


Be
↑
Mg
↓
Ca




periodic table




General


Name, Symbol, Number

magnesium, Mg, 12


Chemical series
alkaline earth metals


Group, Period, Block

2, 3, s


Appearance
silvery white 



Atomic mass
24.3050(6) g/mol


Electron configuration
[Ne] 3s2


Electrons per shell
2, 8, 2


Physical properties


Phase
solid


Density (near r.t.)
1.738 g/cm³


Liquid density at m.p.
1.584 g/cm³


Melting point
923 K
(650 °C, 1202 °F)


Boiling point
1363 K
(1090 °C, 1994 °F)


Heat of fusion
8.48 kJ/mol


Heat of vaporization
128 kJ/mol


Heat capacity
(25 °C) 24.869 J/(mol·K)




Vapor pressure


P/Pa
1
10
100
1 k
10 k
100 k


at T/K
701
773
861
971
1132
1361



Atomic properties


Crystal structure
hexagonal


Oxidation states
2
(strongly basic oxide)


Electronegativity
1.31 (Pauling scale)


Ionization energies
(more)

1st: 737.7 kJ/mol


2nd: 1450.7 kJ/mol


3rd: 7732.7 kJ/mol


Atomic radius
150 pm


Atomic radius (calc.)
145 pm


Covalent radius
130 pm


Van der Waals radius
173 pm


Miscellaneous


Magnetic ordering
paramagnetic


Electrical resistivity
(20 °C) 43.9 nΩ·m


Thermal conductivity
(300 K) 156 W/(m·K)


Thermal expansion
(25 °C) 24.8 µm/(m·K)


Speed of sound (thin rod)
(r.t.) (annealed)
4940 m/s


Speed of sound (thin rod)
(r.t.) 45 m/s


Shear modulus
17 GPa


Bulk modulus
45 GPa


Poisson ratio
0.29


Mohs hardness
2.5


Brinell hardness
260 MPa


CAS registry number
7439-95-4


Notable isotopes




Main article: Isotopes of magnesium


iso

NA

half-life

DM

DE (MeV)

DP


24Mg

78.99%

Mg is stable with 12 neutrons


25Mg

10%

Mg is stable with 13 neutrons


26Mg

11.01%

Mg is stable with 14 neutrons


Magnesium (chemical symbol Mg, atomic number 12) is the eighth most abundant chemical element in the Earth's crust, constituting about 2 percent of the crust by weight. It is the third most plentiful element dissolved in seawater. Silvery white in color, it is classified as an alkaline earth metal and is not found as a free metal in nature.

Contents

1 Discovery and occurrence
2 Notable characteristics of element and compounds

2.1 Isotopes


3 Health and nutrition

3.1 Food sources


4 Applications

4.1 Uses of the metal and its alloys
4.2 Uses of magnesium compounds


5 Precautions
6 References
7 External links
8 Credits



In its ionic form, magnesium is one of the most essential elements for all known living organisms. Its salts, therefore, are used as additives in foods and fertilizers. When alloyed with aluminum, it is widely used to make beverage cans, airplanes, and missiles. The brilliant white light produced by burning magnesium is used in flashlight photography and fireworks. Magnesium oxide is used for lining furnaces, the hydroxide (in the form of "milk of magnesia") is a laxative and antacid, its sulfate (Epsom salt) is used in medicine. Magnesium carbonate is used in brick, flooring, fireproofing, and fire-extinguishing compositions.

Discovery and occurrence
The name magnesium originates from the Greek word for a district called Magnesia in Thessaly. The element, which is always combined with other elements in nature, is found in deposits of over 60 minerals. Of these, only dolomite, magnesite, brucite, carnallite, talc, and olivine are of commercial importance.
In England in 1755 Joseph Black recognized magnesium as being an element. In 1808 Sir Humphry Davy electrolytically isolated pure magnesium metal from a mix of magnesia (magnesium oxide, MgO) and mercuric oxide (HgO). French chemist Antoine Bussy described another method of preparing it in 1831.
In the United States, this metal is principally obtained by electrolysis of fused magnesium chloride from brines, wells, and seawater:

cathode: Mg2+ + 2 e- → Mg
anode: 2 Cl- → Cl2 (gas) + 2 e-
The United States has traditionally been the major world supplier of this metal, supplying 45 percent of world production even as recently as 1995. Today, the U.S. market share is at 7 percent.[1] As of 2005, China has taken over as the dominant supplier, pegged at 60 percent world market share, which increased from 4 percent in 1995. To isolate the metal, China relies almost completely on what is called the Pidgeon process, in which magnesium oxide is reduced at high temperatures with silicon.

Notable characteristics of element and compounds
As a member of the series of alkaline earth metals, magnesium lies in group 2 (former group 2A) of the periodic table, between beryllium and calcium. In addition, it is placed in period 3, immediately following sodium.
Elemental magnesium is a fairly strong, lightweight metal (two thirds the density of aluminum). It slightly tarnishes when exposed to air, although unlike the alkaline metals, storage in an oxygen free environment is unnecessary because magnesium is protected by a thin layer of oxide that is fairly impermeable and hard to remove.
Like calcium, magnesium reacts with water at room temperature, although it reacts much more slowly than calcium. When it is submerged in water, hydrogen bubbles begin to form almost unnoticeably on the metal's surface, but the powdered metal reacts much more rapidly. At higher temperatures, the reaction will occur faster (see precautions).
Magnesium is a highly flammable metal that is easy to ignite when powdered or shaved into thin strips. It is, however, difficult to ignite in mass or bulk. Once ignited, it is difficult to extinguish, being able to burn in both nitrogen (forming magnesium nitride), and carbon dioxide (forming magnesium oxide and carbon). It burns at a temperature of approximately 2,500 K (2200 °C, 4000 °F). The autoignition temperature of magnesium is approximately 744 K (473 °C, 883 °F).
When it burns in air, magnesium produces a brilliant white light. This property was applied in the early days of photography, when magnesium powder (flash powder) was used as a source of illumination. Later, magnesium ribbon was used in electrically ignited flash bulbs.
When glowing white at high temperatures, magnesium has many chemical properties that it does not possess at lower temperatures. It also becomes more toxic, but the high temperature alone is extremely dangerous.
Magnesium compounds are typically white crystals. Most are soluble in water, providing the sour-tasting magnesium ion Mg2+. Small amounts of dissolved magnesium ion contributes to the tartness and taste of natural waters. In large amounts, magnesium ion is an ionic laxative. The so-called "milk of magnesia" is a water suspension of one of the few insoluble magnesium compounds, magnesium hydroxide, and it is named after the appearance of the undissolved particles. Milk of magnesia is a mild base.

Isotopes
Three isotopes of magnesium are found in nature: 24Mg, 25Mg, and 26Mg, at abundances of approximately 79%, 10%, and 11%, respectively. All three isotopes are stable.
26Mg has found application in isotopic geology, similar to that of aluminum. 26Mg is a radiogenic daughter product of 26Al, which has a half-life of 717,000 years. Large enrichments of stable 26Mg have been observed in the Ca-Al-rich inclusions of some carbonaceous chondrite meteorites. The anomalous abundance of 26Mg is attributed to the decay of its parent 26Al in the inclusions. Therefore, the meteorite must have formed in the solar nebula before the 26Al had decayed. Hence, these fragments are among the oldest objects in the solar system and have preserved information about its earliest history.

Health and nutrition
Magnesium ions are essential for the cells of all known living organisms. Many enzymes require the presence of magnesium ions for their catalytic action. Plants have an additional use for magnesium in that chlorophylls contain magnesium. The adult human daily nutritional requirement (which depends on various factors including sex and body size) is 300-400 milligrams (mg) per day.
Too much magnesium in the diet can make it difficult for the human body to absorb calcium. On the other hand, inadequate magnesium intake has been associated with muscle spasms, cardiovascular disease, diabetes, high blood pressure, anxiety disorders, and osteoporosis. Acute deficiency is rare, and is more common as a drug side effect (such as chronic alcohol or diuretic use) than from low food intake per se. The incidence of chronic deficiency, resulting in less than optimal health, is debated.
The Dietary Reference Intake (DRI) upper tolerated limit for supplemental magnesium is 350 mg/day (calculated as mg of Mg elemental in the salt). The most common symptom of excess oral magnesium intake is diarrhea. Given that the kidneys of adult humans excrete excess magnesium efficiently, oral magnesium poisoning in adults with normal renal function, is very rare. Infants have reduced ability to excrete excess magnesium even when healthy, so they should not be given magnesium supplements, except under a physician's care.

Food sources
Magnesium is present in many foods, but it usually occurs in small amounts. As with most nutrients, daily needs for magnesium cannot be met from a single food. Eating a wide variety of foods, including fruits, vegetables, and plenty of whole grains, helps ensure an adequate intake of magnesium. Green vegetables such as spinach provide magnesium because the center of the chlorophyll molecule contains magnesium. Nuts—especially almonds), seeds, and some whole grains—are also good sources of magnesium.
The magnesium content of refined foods is usually low. Whole-wheat bread, for example, has twice as much magnesium as white bread because the magnesium-rich germ and bran are removed when white flour is processed.
Water can provide magnesium, but the amount varies according to the water supply. "Hard" water contains more magnesium than "soft" water. Dietary surveys do not estimate magnesium intake from water, which may lead to underestimating total magnesium intake and its variability.
The following figures indicate the amount of magnesium in some foods:

spinach (1/2 cup): 80 mg
peanut butter (2 tablespoons): 50 mg
black-eyed peas (1/2 cup): 45 mg
milk: low fat (1 cup): 40 mg
Applications
Uses of the metal and its alloys
Magnesium is the third most commonly used structural metal, following steel and aluminum. In its purest form, magnesium can be compared to aluminum. Like aluminum, it is strong and light. Its principal use is as an alloying additive to aluminum—these aluminum-magnesium alloys, sometimes called magnalium or magnelium, are used mainly for beverage cans.
Magnesium alloy is also used in several high-volume parts manufacturing applications, including automotive and truck components. Specialty, high-grade car wheels of magnesium alloy are called "mag wheels." In 1957, a Corvette SS, designed for racing, was constructed with completely magnesium body panels. Volkswagen has used magnesium in its engine components for many years. For a long time, Porsche used magnesium alloy for its engine blocks, due to the weight advantage.
Recently, there appears to be renewed interest in magnesium engine blocks, as featured in the 2006 BMW 325i and 330i models. The award-winning BMW engine uses an aluminum alloy insert for the cylinder walls and cooling jackets surrounded by a high-temperature magnesium alloy AJ62A. The application of magnesium AE44 alloy in the 2006 Corvette Z06 engine cradle has advanced the technology of designing robust automotive parts in magnesium. Both these alloys are recent developments in high-temperature, low-creep magnesium alloys. New alloy development and lower costs, which are becoming competitive with aluminum, will increase the number of automotive applications.
Magnesium is also used in electronic devices. Given its low weight and good mechanical and electrical properties, it is widely used in the manufacture of mobile phones, laptops, cameras, and so forth.
There are several other uses of magnesium, some of which are listed below.

It is a reducing agent for the production of pure uranium and other metals from their salts. It is also used for the removal of sulfur from iron and steel.
The printing industry uses it for photoengraved plates.
As an alloying agent, it improves the mechanical, fabrication, and welding characteristics of aluminum.
Combined in alloys, it is essential for airplane and missile construction.
It is an additive agent for conventional propellants and used in producing nodular graphite in cast iron.
The extremely high temperature at which magnesium burns makes it a handy tool for starting emergency fires during outdoor recreation.
As burning magnesium produces a brilliant white light, it is used in flashlight photography, flares, pyrotechnics, sparklers, and incendiary bombs.
Uses of magnesium compounds
Magnesium compounds, primarily magnesium oxide, are used mainly as refractory materials in furnace linings for producing iron, steel, nonferrous metals, glass, and cement.
The hydroxide of magnesium, or milk of magnesia, is used as a laxative and antacid; its chloride and citrate are used as oral magnesium supplements; and its sulfate (Epsom salts) is used for various purposes in medicine.
Magnesium carbonate (MgCO3) is used in making bricks and in flooring, fireproofing, and fire-extinguishing compositions. Its powder is used by gymnasts and weightlifters to improve their grip on objects such as the lifting bar.
Magnesium stearate is a slightly flammable white powder with lubricative properties. The pharmaceutical industry uses it in the manufacture of tablets, to prevent the tablets from sticking to the equipment during the tablet compression process.
Given that magnesium (in the ionic form) is necessary for living organisms, magnesium salts are additives in foods, fertilizers, and tissue culture media.
Precautions
Pure magnesium metal and its alloys are highly flammable when in the molten state, powdered form, or ribbon form. Burning or molten magnesium metal reacts violently with water. Magnesium powder is an explosive hazard. One should wear safety glasses while working with magnesium. The bright white light (including ultraviolet) produced by burning magnesium can damage the eyes. Water should not be used to extinguish magnesium fires, because it can actually feed the fire, according to the following reaction:[2]

Mg (s) + 2 H2O (v) → Mg(OH)2 (aq) + H2 (g)
or in words:
Magnesium (solid) + steam → Magnesium hydroxide (aqueous) + Hydrogen (gas)
Carbon dioxide fire extinguishers should not be used either, because magnesium can burn in carbon dioxide (forming magnesium oxide, MgO, and carbon).[3] A dry chemical fire extinguisher (Class D) should be used if available, or else the fire should be covered with sand or magnesium foundry flux. An easy way to put out small metal fires is to place a polyethene bag filled with dry sand on top of the fire. The heat of the fire will melt the bag and the sand will flow out onto the fire.

ReferencesISBN links support NWE through referral fees

↑ [1] Vardi, Nathan, Man With Many Enemies Forbes.com, July 22, 2002. Accessed on June 26, 2006.

↑ Webelements.com. “Chemistry: Periodic Table: magnesium: chemical reaction data." Accessed on June 26, 2006.

↑ [2] "Demo Lab: Reaction Of Magnesium Metal With Carbon Dioxide" Accessed on June 26, 2006.


External links
All links retrieved November 5, 2022.

WebElements.com – Magnesium
The Magnesium Website

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