Group 2A — The Alkaline Earth Metals

1A 2A 3A 4A 5A 6A 7A 8A
(1) (2) (13) (14) (15) (16) (17) (18)
3B 4B 5B 6B 7B 8B 1B 2B
(3) (4) (5) (6) (7) (8) (9) (10) (11) (12)
1 H He
2 Li Be B C N O F Ne
3 Na Mg Al Si P S Cl Ar
4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
6 Cs Ba La   Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
7 Fr Ra Ac   Rf Db Sg Bh Hs Mt Ds Rg Uub Uuq
6   Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
7   Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

 

Group 2A (or IIA) of the periodic table are the alkaline earth metals:  beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra).  They are harder and less reactive than the alkali metals of Group 1A.  The name comes from the fact that the oxides of these metals produced basic solutions when dissolved in water, and they remained solids at the temperatures available to the ancient alchemists.  Like the Group 1A elements, the alkaline earth metals are too reactive to be found in nature in their elemental form.

The alkaline earth metals have two valence electrons in their highest-energy orbitals (ns2). They are smaller than the alkali metals of the same period, and therefore have higher ionization energies.  In most cases, the alkaline earth metals are ionized to form a 2+ charge.

The alkaline earth metals have much higher melting points than the alkali metals: beryllium melts at 1287C, magnesium at 649C, calcium at 839C, strontium at 768C, barium at 727C, and radium at 700C.  They are harder metals than the Group 1A elements, but are soft and lightweight compared to many of the transition metals.

Salts of the Group 2A metals are less soluble in water than those of Group 1A because of the higher charge densities on the 2+ cations; nevertheless, many Group 2A salts are at least moderately soluble.   Some Group 2A salts bond strongly to water molecules, and crystallize as hydrates; among these are Epsom salt, MgSO47H2O, and gypsum, CaSO42H2O.

 

Beryllium (Be, Z=4).

Beryllium is a silver-white, soft metal.  Its name is derived from the Greek word for the mineral beryl, beryllo.  It is found in the Earth's crust at a concentration of 2.6 ppm, making it the 47th most abundant element.  The primary ores of beryllium are beryl [beryllium aluminium silicate, Be3Al2(SiO3)6] and bertrandite [beryllium silicate hydroxide, Be4Si2O7(OH)2].  Gemstone-quality beryls include emeralds and aquamarine; the green color of these gems comes from trace amounts of chromium.

Because of its small size and high charge density, beryllium bonds through covalent bonding instead of ionic bonding.  Elemental beryllium is very unreactive towards air and water, even at high temperatures.

Beryllium is used to make windows for X-ray tubes (it is transparent to X-rays), and is used in alloys with other metals, such as copper and nickel, to make spark-proof tools and watch springs.  Beryllium is also using in casings for nuclear weapons and in nuclear power plants because of its ability to reflect neutrons.  Beryllium accumulates in bones; long-term exposure to beryllium results in inflammation in the lungs and shortness of breath (a condition called berylliosis).

 

Magnesium (Mg, Z=12).

Magnesium is a silver-white, relatively soft metal.  The name of the element is derived from Magnesia, a district in Thessaly in central Greece.  It is found in the Earth's crust at a concentration of 2.3%, making it the 7th most abundant element.  Large amounts of magnesium are also present in the minerals in the Earth's mantle.  It is obtained from seawater, carnalite [MgKCl36H2O], dolomite [a mixture of calcium and magnesium carbonate, CaMg(CO3)2], and magnesite [magnesium carbonate, MgCO3].

Magnesium alloyed with aluminum and traces of other metals is used in car and aircraft construction; magnesium alloys are also used in other lightweight devices, such ladders, cameras, bicycle frames, hard disk drives, etc.  Magnesium is more easily oxidized than iron, and is used in sacrificial anodes to protect iron pipes and other structures that corrode easily.

Magnesium burns in air with a brilliant white flame, and is used in fireworks and incendiary bombs.  (It used to be used in disposable flashbulbs, but this use has been supplanted by other types of illumination.)  Magnesium fires are very difficult to put out, since even in the absence of air, burning magnesium reacts with nitrogen to form magnesium nitride (Mg3N2), and with water to produce magnesium hydroxide and hydrogen gas.

Magnesium is found in a number of familiar compounds.  Magnesium oxide, MgO, is used in refractory bricks that are capable of withstanding the high temperatures in fireplaces and furnaces (magnesium oxide melts at 2800 C).  Magnesium sulfate heptahydrate, MgSO47H2O, better known as Epsom salt, is a muscle relaxant and a mild laxative.  Magnesium hydroxide, Mg(OH)2, also known as milk of magnesia, is a laxative and antacid.  (The "milk" in "milk of magnesia" refers to the fact that since magnesium hydroxide is not very soluble in water, it tends to form a chalky, white suspension that looks like milk — but with considerably different physiological effects.)

Green plants contain a molecule called chlorophyll, which consists of a flat ring of carbon and nitrogen atoms with a large open space in the middle, in which a magnesium ion is bound, held in place by the nitrogen atoms.  The chlorophyll molecule absorbs light from the sun, and in the process of photosynthesis, the energy from the light is converted into chemical energy that the plant can use to power a multitude of processes.

In organic chemistry, magnesium reacts with bromoalkanes (hydrocarbons containing carbon-bromine bonds) to form organomagnesium compounds known as Grignard reagents (after their discoverer, Victor Grignard, who won a Nobel Prize in Chemistry, 1912).  These compounds are extremely useful in forming new carbon-carbon bonds, and are often used in the synthesis of organic compounds.  Grignard reagents are notoriously sensitive to water, and care must be taken to ensure that the apparatus in which the reaction is being carried out is extremely dry.

 

Calcium (Ca, Z=20).

Calcium is a silver-colored, relatively soft metal.  The name of the element is derived from the Latin word for lime, calx.  It is found in the Earth's crust at a concentration of 4.1%, making it the 5th most abundant element.  The major sources of calcium are calcite and limestone [calcium carbonate, CaCO3], anhydrite [calcium sulfate, CaSO4], gypsum [calcium sulfate dihydrate, CaSO42H2O], and dolomite [a mixture of calcium and magnesium carbonate, CaMg(CO3)2].

Calcium salts forms the hard parts of the bodies of most living creatures, from the shells of marine organisms and the coral of coral reefs (in the form of calcium carbonate, CaCO3) to the bones and teeth of land-dwelling creatures (in the form of hydroxyapatite crystals, Ca3(PO4)2]3Ca(OH)2).

Since calcium forms such hard minerals, it is useful in building materials, such as plaster, mortar, and cement.  Mortar is made from calcium oxide, CaO, also known as lime, or quicklime.  When calcium oxide is treated with water it forms calcium hydroxide, Ca(OH)2, or slaked lime, which absorbs carbon dioxide from the air and gradually forms calcium carbonate, CaCO3.  Lime, heated by hydrogen burning in oxygen, burns with a brilliant white light, which can be focused into a narrow beam visible over great distances.  This kind of lighting was used in lighthouses, in surveying, and in theaters to produce spotlights (leaving the actor "in the limelight").

Calcium chloride is a deliquescent (it absorbs enough water from the air that it dissolves in the solution), and is used to remove moisture from the air in damp basements.  (It would take a stronger person that me to resist calling a freshly opened box of calcium chloride that wasn't behaving properly as a "juvenile deliquescent.")

"Hard water" contains dissolved minerals having 2+ or 3+ charges, such as calcium and magnesium; these salts cause some soaps and detergents to precipitate out as "soap scum"; these minerals precipitate out over time to form "scale" in water heaters and pots.  The calcium can be removed by water softeners, which exchange the calcium ions for sodium ions, which have 1+ charges, and do not readily precipitate out.

 

Strontium (Sr, Z=38).

Strontium is a shiny, relatively soft metal.  The name of the element is derived from Strontian, a town in Scotland where the mineral strontianite was discovered, from which strontium was first isolated.  It is found in the Earth's crust at a concentration of 370 ppm, making it the 16th most abundant element.  It is found in the ores celestite [strontium sulfate, SrSO4] and strontianite [strontium carbonate, SrCO3].

Strontium salts produce brilliant red colors when heated, and are used in fireworks and flares for this reason.  Radioactive strontium-90 (a beta-emitter) is produced in nuclear explosions; since it is chemically similar to calcium, it becomes incorporated into bone in people who are exposed to it.  Strontium-90 is a beta-emitter, and interferes with the production of red blood cells.

 

Barium (Ba, Z=56).

Barium is a shiny, soft metal.  The name of the element is derived from the Greek word barys, which means "heavy," in reference to the high density of some barium minerals.  It is found in the Earth's crust at a concentration of 500 ppm, making it the 14th most abundant element.  It is found in the ores barite [barium sulfate, BaSO4] and witherite [barium carbonate, BaCO3].

Barium was discovered in the 1500s in the form of the "Bologna stones" (now known to be barium sulfate, BaSO4) discovered near Bologna, Italy .  These stones glowed in the presence of light, and also when heated.  Barium salts give off a green color when heated, and are used in fireworks (in the form of barium nitrate, Ba(NO3)2).

Barium sulfate, BaSO4, is poisonous, but it is so insoluble that that it passes through the body before any absorption of barium can take place.  It is used in the diagnosis of some intestinal problems in the form of "barium enemas":  barium sulfate is opaque to X-rays, and can be used to take X-rays of the digestive tract.

 

Radium (Ra, Z=88).

Radium is a soft, shiny, radioactive metal.  The name of the element was derived from the Latin word for "ray," radius, because of its ability to glow in the dark with a faint blue light.  It is found in the Earth's crust at a concentration of 0.6 ppt (parts per trillion), making it the 84th most abundant element.  It is found in trace amounts in uranium ores, but commercially used radium is more easily obtained from spent nuclear fuel.

Radium was discovered by Pierre and Marie Curie in 1898; they extracted a milligram of radium from three tons of uranium ore.  Radium is produced in the radioactive decay of uranium-235, uranium-238, thorium-232, and plutonium-241.  After its discovery, and before the dangers of radiation were understood, radium was used in a lot of quack cures and patent medicines.  Radium was used to make glow-in-the-dark clock faces in the early 1900's; the alpha particles emitted by the radium struck particles of zinc sulfide, causing them to glow, but were stopped by the clock's casing by by the glass in the clock face.  Many of the workers who painted these clock faces became ill, or died of radiation sickness.

 

 

References

John Emsley, The Elements, 3rd edition.  Oxford:  Clarendon Press, 1998.

John Emsley, Nature's Building Blocks:  An A-Z Guide to the Elements.  Oxford:  Oxford University Press, 2001.

David L. Heiserman, Exploring Chemical Elements and their Compounds.  New York:  TAB Books, 1992.