78 | The History and Use of Our Earth’s Chemical Elements and may cause death. Strontium-90 does have some utilitarian uses as the radiation source for
78 | The History and Use of Our Earth’s Chemical Elements and may cause death. Strontium-90 does have some utilitarian uses as the radiation source for
instruments that measure the thickness of materials during their production (e.g., cigarettes, building materials, and textiles).
Strontium hydroxide [Sr(OH) 2 ] is used to extract sugar from sugar beet molasses. It is also used in the manufacture of soaps, adhesives, plastics, glass, and lubricants that can be used in very high or low temperature environments.
Strontium iodide (SrI 2 ) is made by treating strontium carbonate with hydrochloric acid. It is used as a medicinal source of iodine. Strontium nitrate [Sr(NO 3 ) 2 ], because of the bright red flame it produces when burned, is used in fireworks, matches, marine signals, and so forth. Strontium carbonate (SrCO 3 ) is used to make radiation-resistant glass and TV picture tubes, as well as pyrotechnics. Strontium peroxide or strontium dioxide (SrO 2 ) can cause fires or explode when heated and in contact with organic substances. It is used as both a reducing agent and an oxidizing agent. Strontium sulfide (SrS) smells like rotten eggs. It is used as a depilatory to remove hair from skin and hides.
Hazards As a powder, strontium metal may spontaneously burst into flames. Both its metal and
some of its compounds will explode when heated. Some of the compounds will explode if struck with a hammer.
Both the metal and some compounds will react with water to produce strontium hydroxide [Sr(OH) 2 ] and release hydrogen gas. The heat from the exothermic reaction may cause the hydrogen to either burn or explode [Sr + 2H 2 O → Sr(OH) 2 +H 2 ↑]. Some compounds, such as strontium chromate and strontium fluoride, are carcinogens and toxic if ingested. Strontium-90 is particularly dangerous because it is a radioactive bone-seeker that replaces the calcium in bone tissue. Radiation poisoning and death may occur in people exposed to excessive doses of Sr-90. Strontium-90, as well as some other radioisotopes that are produced by explosions of nuclear weapons and then transported atmospherically, may be inhaled by plants and animals many miles from the source of the detonation. This and other factors led to the ban on atmospheric testing of nuclear and thermonuclear weapons.
BARIUM SYMBOL:Ba PERIOD:6 GROUP:2(IIA) ATOMICNO:56
ATOMICMASS:137.347amu VALENCE:2 OXIDATIONSTATE:+2 NATURALSTATE:Solid ORIGINOFNAME:ThenamebariumisderivedfromtheLatinwordbarys,whichmeans
“heavy.” ISOTOPES:Therearecurrently35knownisotopesofbarium,rangingfromBa-120toBa-
148.Sevenoftheseisotopesarestable.Thepercentagesofeachasfoundinnature areasfollows:Ba-130=0.106%,Ba-132=0.101%,Ba-134=2.147%,Ba-135= 6.592%,Ba-136=7.854%,Ba-137=11.23%,andBa-138=71.7%.
Guide to the Elements | 79
ELECTRONCONFIGURATION EnergyLevels/Shells/Electrons Orbitals/Electrons
1-K=2 s2
2-L=8
s2,p6
3-M=18
s2,p6,d10
4-N=18
s2,p6,d10
5-O=8
s2,p6
6-P=2
s2
Properties Barium is the fifth element in group 2 (IIA) of the alkali earth metals and has most of the
properties and characteristics of the other alkali earth metals in this group. For example, they all are called alkaline earths because, when first discovered, they exhibited both characteristics of alkaline (basic) substances and characteristics of the earth from which they came. Ancient humans did not know they were metals because their metallic forms do not exist in nature. Barium is a silvery metal that is somewhat malleable and machineable (can be worked on a lathe, stretched and pounded). Its melting point is 725°C, its boiling point is about 1640°C, and its density is 3.51 g/cm 3 . (The accurate figures for its properties are difficult to determine because of barium’s extreme activity—the pure metal will ignite when exposed to air, water, ammonia, oxygen, and the halogens.)
Characteristics When barium burns in air, it produces barium oxide (2Ba + O 2 → 2BaO). When metallic
barium burns in water, it forms barium hydroxide [Ba + 2H 2 O → Ba(OH) 2 +H 2 ↑]. Several barium compounds burn with a bright green flame, which make them useful for fireworks. Barium is more reactive with water than are calcium and strontium. This is a result of the valence electrons’ being further from the positive nucleus. Therefore, barium is more electro- negative than the alkali earth metals with smaller nuclei.
In powdered form, it will burst into a bright green flame at room temperature. AbundanceandSource Barium is the 17th most abundant element in the Earth’s crust, making up about 0.05%
of the crust. It is found in the minerals witherite, which is barium carbonate (BaCO 3 ), and barite, known as barium sulfate (BaSO 4 ). Pure barium metal does not exist on Earth—only as compounds or in minerals and ores. Barium ores are found in Missouri, Arkansas, Georgia, Kentucky, Nevada, California, Canada, and Mexico.
80 | The History and Use of Our Earth’s Chemical Elements
It is produced by the reduction of barium oxide (BaO), using aluminum or silicon in
a high-temperature vacuum. It is also commercially produced by the electrolysis of molten barium chloride (BaCl ) at about 950 o 2 C, wherein the barium metal is collected at the cathode and chlorine gas is emitted at the anode.
History Alchemists in the early Middle Ages knew about some barium minerals. Smooth round
pebble-like stones found in Bologna, Italy, were known as “Bologna stones.” When these odd stones were exposed to sunlight, or even a primitive reading lamp, they would continue to glow for several years. This characteristic made them attractive to witches as well as the alche-
mists. These stones are actually the mineral barite, barium sulfide (BaSO 4 ), which today is a major source of barium metal. Chemists did not discover the mineral witherite (BaCO 3 ) until the eighteenth century. Carl Wilhelm Scheele (1742–1786) discovered barium oxide in 1774, but he did not isolate or identify the element barium. It was not until 1808 that Sir Humphry Davy used molten barium compounds (baryta) as an electrolyte to separate, by electrolysis, the barium cations, which were deposited at the negative cathode as metallic barium. Therefore, Davy received the credit for barium’s discovery.
CommonUses Pure barium metal has few commercial uses because of it reactivity with air and water.
Nevertheless, this property makes it useful as a “getter” or scavenger to remove the last traces of gas from vacuum tubes. Barium metal is used to form alloys with other metals. One alloy is used to make sparkplugs that easily emit electrons when heated, thus improving the efficiency of internal combustion engines.
Its compounds have many practical uses. For example, when the mineral barite is ground up into a fine powder, it can be used as a filler and brightener for writing and computer paper. It is also used (along with zinc sulfide) as a pigment, called lithopone, for white paint. Barium compounds are also used in the manufacture of plastics, rubber, resins, ceramics, rocket fuel, fireworks, insecticides, and fungicides and to refine vegetable oils.
A major medical use is a solution of barium sulfide (with flavoring) that is ingested by patients undergoing stomach and intestinal X-ray and CT scan examinations. Barium sulfide is opaque to X-rays, and thus it blocks the transmission of the rays. The organs appear in con- trast against a background, which highlights any problems with the digestive system.
ExamplesofCompounds Barium acetate [Ba(C 2 H 3 O 2 ) 2 •H 2 O], a white crystal, is used as a dryer for paints and var-
nishes. It is produced by adding acetic acid to barium sulfate and recovering the crystals by evaporation. It is also used as a textile mordant and catalyst.
Barium bromate [Ba(BrO 3 ) 2 ] is used as a corrosion inhibitor to prevent rust and as an oxi- dizing agent and chemical reagent. Barium nitrate [Ba(NO 3 ) 2 ] burns with a bright green flame and is used in signal flares and pyrotechnics. It can be produced by treating barium carbonate with nitric acid. Barium chloride (BaCl 2 ) is used in the manufacture of paint pigments and dyeing textiles and as an additive in oils. It is also used as a water softener.
Guide to the Elements | 81 Barium peroxide (BaO 2 ) is a grayish-white dry powder that makes an excellent bleaching agent
that can be stored in paper packages. Its bleaching qualities are released when mixed with water. Barium sulfate (BaSO 4 ), as mentioned, is used in medicine as an opaque liquid medium to block X-rays when ingested, thus providing an image of ulcers and intestinal problems. It is also used in the manufacture of paints, rubber, and plastics.
Barium thiosulfate (BaS 2 O 3 ) is mainly used in explosives, matches, varnishes, and photog- raphy.
Hazards Barium metal, in powder form, is flammable at room temperature. It must be stored in an
oxygen-free atmosphere or in petroleum. Many of barium’s compounds are toxic, especially barium chloride, which affects the func- tioning of the heart, causing ventricular fibrillation, an erratic heartbeat that can lead to death. Several of barium’s compounds are explosive as well as toxic if ingested or inhaled. Care should
be used when working with barium and other alkali metals in the laboratory or in industry. RADIUM
SYMBOL:Ra PERIOD:7 GROUP:2(IIA) ATOMICNO:88 ATOMICMASS:226.03amu VALENCE:2 OXIDATIONSTATE:+2 NATURALSTATE:
Solid ORIGINOFNAME:Radium’snameisderivedfromtheLatinwordradius,whichmeans “ray.” ISOTOPES:Therearenostableisotopesofradium.Radiumhas25knownradioisotopes, rangingfromRa-206toRa-230.Theirhalf-livesrangefromafractionofasecondto hundredsofyears.Radium-226wasdiscoveredbytheCuriesandhasahalf-lifeof about1630years.Ra-226isthemostabundantisotope,andthus,Ra-226isusedto determineradium’satomicmass.
Variousradiumisotopesarederivedthroughaseriesofradioactivedecayprocesses.For example,Ra-223isderivedfromthedecayofactinium.Ra-228andRa-224aretheresult oftheseriesofthoriumdecays,andRa-226isaresultofthedecayoftheuraniumseries.
ELECTRONCONFIGURATION EnergyLevels/Shells/Electrons Orbitals/Electrons
s2,p6