68 | The History and Use of Our Earth’s Chemical Elements History
68 | The History and Use of Our Earth’s Chemical Elements History
From the days of the Egyptians, when emeralds were a particular favorite of kings, beryl has also been a favored gemstone. It was not until the late eighteenth century that Abbe René Just Haüy (1743–1822), the father of crystallography, studied the crystalline structures and densities of emeralds and beryl and determined that they were the same mineral. At about the same time, in 1798, Louis-Nicolas Vauquelin (1763–1829) discovered that both emeralds and beryl were composed of a new element with four protons in its nucleus. The element was named “glucina” because of its sweet taste. It was not until the nineteenth century that the
metal beryllium was extracted from beryllium chloride (BeCl 2 ) by chemical reactions. Late in the nineteenth century, P. Lebeau (dates unknown) separated the metal by the electrolytic process.
CommonUses In the mid-twentieth century, the determination that beryllium has a number of unique
properties led to the production of beryllium metal by electrolysis on a commercial scale. It proved valuable as an alloy metal to produce specialized, strong—but light—structural metals for use in satellites, aircraft, and spacecraft.
A 2% beryllium mixture with copper produces a unique alloy of bronze that is six times stronger than copper metal. This alloy does not give off sparks when struck with a hammer—a valuable characteristic when metals must be used in explosive gaseous environments. This alloy sometimes contains small amounts of other metals such as nickel or cobalt, which makes for excellent electrical conductivity for switching equipment, given the alloy’s simultaneous hard- ness and nonsparking qualities. Beryllium is also “transparent” to X-rays, which makes it ideal for windows for X-ray tubes.
In 1932 James Chadwick (1891–1974) bombarded beryllium with alpha particles (helium nuclei) that produced free neutrons. Since then, this nuclear process has made beryllium a reliable neutron emitter for laboratory nuclear research. Beryllium is not only an excellent moderator to slow down high-speed neutrons in nuclear reactors, but it also can act as a reflector of neutrons as well.
Beryllium is an excellent source of alpha particles, which are the nuclei of helium atoms. Alpha particles (radiation) are not very penetrating. These particles travel only a few inches in air and can be stopped by a sheet of cardboard. Alpha particles are produced in cyclotrons (atom smashers) and are used to bombard the nuclei of other elements to study their charac- teristics.
In the first part of the twentieth century, beryllium was used as coating inside fluorescent electric light tubes, but proved carcinogenic (causes cancer) when broken tubes produced beryllium dust that was inhaled. Because of this potential to cause cancer, since 1949 beryl- lium has no longer been used as the inside coating of fluorescent tubes. Beryllium is also used for computer parts, electrical instrument components, and solid propellant rocket fuels. Because it is one of the few metals that is transparent to X-rays, it is used to make special glass for X-ray equipment.
ExamplesofCompounds Beryllium carbide (Be 2 C) is used for the cores in nuclear reactors.
Guide to the Elements | 69 Beryllium chloride (BeCl 2 ) is used as a catalyst to accelerate many organic reactions, and
beryllium chloride is the electrolyte used along with NaCl in the electrolytic process to pro- duce beryllium metal.
Beryllium copper is not really a compound, but a very useful alloy that often contains other metals such as cobalt or nickel in small amounts. It is a hard, strong alloy with excellent elec- trical conductivity, which makes it very useful in electrical switching equipment owing to its nonsparking qualities. It makes excellent spot-welding electrodes, springs, and metal bushings, cams, and diaphragms.
Beryllium fluoride (BeF 2 ) is an example of beryllium that has an oxidation state of +2, combining with a negative anion element with an oxidation state of –1. Beryllium fluoride is also used along with magnesium metal in the chemical reduction process to produce beryl- lium metal.
Beryllium hydride (BeH 2 ) liberates hydrogen gas when mixed with water. It is used as a source of hydrogen in experimental rockets and fuel cells. Beryllium oxide (BeO) is a beryllium compound produced in significant commercial quan- tities. The chemical process starts with minerals containing aluminum silicate and silicon dioxide and undergoes a number of chemical reactions, some at high temperatures, to end up with BeO.
Hazards The elemental metallic form of beryllium is highly toxic, as are most of its compounds.
When inhaled, the fumes, dust, or particles of beryllium are highly carcinogenic. Some beryl- lium compounds are toxic when they penetrate cuts in the skin (e.g., when an old fluorescent tube breaks). Beryllium oxide when inhaled can result in a fatal disease known as berylliosis (similar to, but more toxic than, silicosis).
As with many other chemicals, beryllium has its positives and negatives. Although it is an important industrial chemical, the handling of beryllium is best left to experienced workers and laboratory personnel in proper facilities.
MAGNESIUM SYMBOL:Mg PERIOD:3 GROUP:2(IIA) ATOMICNO:12
ATOMICMASS:24.305amu VALENCE:2 OXIDATIONSTATE:+2 NATURALSTATE: Solid ORIGINOFNAME:MagnesiumisnamedafterMagnesia,anancientregionofThessaly, Greece,whereitwasmined.Magnesiumisoftenconfusedwithanotherelement,man- ganese.Onewaytoeliminatetheconfusionistothinkofmagnesium(Mg)as“12”and manganese(Mn)as“25”andtousethementaltrickofrememberingthat“g”comes before“n”inthealphabet,somagnesiumistheonewithloweratomicnumber.
ISOTOPES:Thereare15isotopesofmagnesium,rangingfromMg-20toMg-34.Three oftheseisotopesarestable:Mg-24makesup78.99%ofallmagnesiumfoundinthe Earth’scrust.Mg-25makesup10%,andMg-26constitutesmostoftherestat11%. Theother12isotopesareradioactiveandareproducedartificiallywithhalf-livesranging frommicrosecondstoafewhours.
70 | The History and Use of Our Earth’s Chemical Elements
ELECTRONCONFIGURATION EnergyLevels/Shells/Electrons Orbitals/Electrons
s2,p6
3-M=2
s2
Properties Magnesium is a lightweight, silvery-white, malleable alkali earth metal that is flammable.
It has a weak electronegativity (–1.31), which means it is highly reactive as it combines with some nonmetals. As with other alkali earth metals, magnesium is a good conductor of heat and electricity. Its melting point is 648.8°C, its boiling point is 1090°C, and its density is
1.74 g/cm 3 , making it about one-fifth the density of iron and only two-thirds as dense as aluminum.
Characteristics While in a thin solid form, magnesium ignites at 650°C, and it is more easily ignited in
a fine powder form. Burning magnesium produces a brilliant white light. It is also used as an oxidizer to displace several other metals from their compound minerals, salts, and ores. It is alloyed with other metals to make them lighter and more machinable, so that they can be rolled, pounded, formed into wires, and worked on a lathe.
The ground water in many regions of the United States contains relatively high percentages of magnesium, as well as some other minerals. A small amount improves the taste of water, but larger amounts result in “hard” water, which interferes with the chemical and physical action of soaps and detergents. The result is a scum-like precipitate that interferes with the cleansing action. The solution is the use of water softeners that treat hard water with either sodium chloride or potassium chloride, which displace the magnesium—making the water “soft,” resulting in a more effective cleansing action.
AbundanceandSource Magnesium is the eighth most abundant of the elements found in the entire universe, and
the seventh most abundant found in the Earth’s crust. Its oxide (MgO) is second in abun- dance to oxide of silicon (SiO 2 ), which is the most abundant oxide found in the Earth’s crust. Magnesium is found in great quantities in seawater and brines, which provide an endless sup- ply. Each cubic mile of seawater contains about 12 billion pounds of magnesium. Although magnesium metal cannot be extracted from seawater directly, it can be extracted by several
chemical processes through which magnesium chloride (MgCl 2 ) is produced. Electrolysis is then used with the magnesium chloride as the electrolyte at 714°C to produce metallic mag-
Guide to the Elements | 71 nesium and chlorine gas. Another method of securing magnesium is known as the Pigeon
process. This procedure uses the magnesium minerals dolomite or ferrosilicon. Dolomite (CaCO 3 ), which also contains MgCO 3 , is crushed and then heated to produce oxides of Ca and Mg. The oxides are heated to about 1200°C along with the ferrosilicon (an alloy of iron and silicon), and the silicon reduces the magnesium, producing a vapor of metallic magnesium that, as it cools, condenses to pure magnesium metal.
History Magnesium was known in ancient times, but it was not identified as an element until
much later. As the story goes, a seventeenth-century farmer’s cows refused to drink water from
a mineral well in Epsom, England. The farmer tasted the water and found it to be bitter but somewhat refreshing. He also noted that it seemed to heal a rash on his skin. After several years, the demand for his Epsom water was more than he could manage, so other sources were located and exploited. To no one’s surprise, the product became known as Epsom salt.
In the mid-eighteenth century, hundreds of chemists determined that the substance in Epsom salt (also known as magnesia alba) was a salty mixture of a sulfate and an oxide of what was believed to be a new element. In 1755 Joseph Black (1728–1799) separated magnesium oxide from lime. However, he was unable to separate the magnesium from the oxygen. Magnesium was not identified as an element until 1808, by Sir Humphry Davy (1778–1829). Davy used the electrolysis process with an electrolyte of a mixture of mag- nesium oxide (MgO = magnesia) and mercury oxide (HgO). A mixture of magnesium and mercury (an alloy-like amalgam) collected at the negative cathode. He then heated this amalgam, vaporizing the mercury with its lower melting point and leaving the magnesium metal behind.
Davy also discovered several other elements (potassium, barium, calcium, and strontium) by isolating the metals from their compounds through electrolysis. His work led to the devel- opment of electrochemistry, which is the use of electricity as the energy source to break up the oxides of these alkali and alkali earth elements.
CommonUses Small particles of powdered magnesium metal burn with a bright white flame that makes
the magnesium ideal for aerial flares dropped from airplanes that will light up ground areas. It is has also been used in aerial firebombs during wars to devastate a city by fire because water will not extinguish the flames—sand must be used. In the past decades, thin magnesium wire or foil was placed inside glass bulbs containing pure oxygen to form flash bulbs for photo- graphic purposes. When an electric charge ignites the magnesium, a brilliant light is produced. Today most flash cameras use a strobe light instead of flash bulbs.
Pure magnesium metal is lighter in weight than aluminum and, thus, would make an excellent construction metal were it not for its high reactivity and flammability at a rather low temperature when compared to other metals. It is an excellent metal to alloy with other metals for use in the aircraft, space, and automobile industries.
It is used for the production (thermal reduction) of other metals, such as zinc, iron, tita- nium, zirconium, and nickel. For instance, because of its strong electropositive nature, mag- nesium can “desulfurize” molten iron when it combines with the sulfur impurities in the iron to produce high-grade metallic iron plus MgS.
72 | The History and Use of Our Earth’s Chemical Elements
Milk of Magnesia is an alkaline (basic) water suspension and “creamy-like” suspended form
of magnesium hydroxide, Mg(OH) 2 . It is used as an antacid to neutralize excess stomach acid. Magnesium can also be used in the form of Epsom salts as a treatment for rashes and as a laxative. A more important commercial use of Epsom salts is in the tanning of leather, as well as in the dyeing of fabrics.
Magnesium is essential for proper nutrition in humans as well as other living organisms. It plays an important role in the process of photosynthesis in plant chlorophyll and is thus essential to green plants, which are, in turn, essential for most living organisms. Magnesium is also used as a dietary supplement for both humans and animals for maintaining proper enzyme levels.
Magnesium is an important element that acts as a catalyst in many life processes. In addi- tion to photosynthesis, it is also required for the oxidation in animal cells that produce energy and for the production of healthy red blood cells. Humans cannot live without magnesium— which we acquire mainly from various foods.
ExamplesofCompounds Magnesium acetate [Mg(C 2 H 3 O 2 ) 2 •4H 2 O] is used in the textile industry as a mordant
(“fixes” dyes so that they will not run). It is also used as a deodorant and antiseptic. Magnesium chloride ((MgCl 2 ) is mostly obtained from saltwater and has many uses, includ- ing as the source of magnesium metal during electrolysis, as a catalyst, and in the making of ceramics, lubricants, paper and textiles, and disinfectants.
Magnesium fluoride (MgF 2 ) is used to polarize corrective lenses of eyeglasses to reduce the glare of sunlight by selecting the orientation of the light waves passing through the lenses. MgF 2 is also used to polarize windows, sunglasses, and similar optical items. Magnesium chromate (MgCrO 4 ) has several medical uses, including as a dietary supplement and laxative. Magnesium oxide (MgO) is used as a lining for steel furnaces, as a component in ceramics, as food additives and pharmaceuticals, and to make strong window glass, fertilizers, paper, and rubber manufacturing.
Magnesium hydroxide [Mg(OH) 2 ] is a whitish solid and when suspended in water is used as an antacid known as milk of magnesia. Magnesium carbonate (MgCO 3 ) is found in a mixture of natural minerals. It can also be produced in several ways, including pumping carbon dioxide through magnesium oxide or magnesium hydroxide. It is used in pharmaceuticals such as magnesium citrate and as a des- iccant to keep hydroscopic products from caking (table salt) and to strengthen rubber and produce dyes, inks, and cosmetics.
Hazards Magnesium metal, particularly in the form of powder or small particles, can be ignited at
relatively low temperatures. The resulting fires are difficult to extinguish, requiring dry sand or dirt. Water will just accelerate the fire as hydrogen that will intensify the fire is released from the water.
Some magnesium compounds, whose molecules contain several atoms of oxygen— Mg(ClO 4 ), for example—are extremely explosive when in contact with moist organic sub- stance, such as your hands.
Guide to the Elements | 73 Although traces of magnesium are required for good nutrition and health, some com-
pounds of magnesium are poisonous when ingested. CALCIUM
SYMBOL:Ca PERIOD:4 GROUP:2(IIA) ATOMICNO:20 ATOMICMASS:40.078amu VALENCE:2 OXIDATIONSTATE:+2 NATURALSTATE:
Solid ORIGINOFNAME:Itsmodernnamewasderivedfromthewordcalcisorcalx,whichis Latinfor“lime.” ISOTOPES:Thereare20isotopesofcalcium,rangingfromCa-35toCa-54.Ofthesix
stableisotopes,Ca-40makesup96.941%ofthecalciumfoundintheEarth’scrust,and Ca-42=0.647%,Ca-43=0.135%,Ca-44=2.086%,Ca-46=0.004%,andCa-48= 0.187%foundonEarth.Alltheotherisotopesofcalciumareradioactiveandareartifi-
ciallyproducedwithhalf-livesrangingfromafewmicrosecondsto1×10 5 years.
RadioactiveCa-45emitsbetaparticles(high-speedelectrons)andhasahalf-lifeof about163days.Itisusedtodeterminethecalciumlevelsinbonesandinsoils.
ELECTRONCONFIGURATION EnergyLevels/Shells/Electrons Orbitals/Electrons
s2,p6
3-M=8