Metal Story: Lithium (Li) – The Lightest Metal
It was 150 years in 1967 since lithium, the first in the group of metals in the Mendeleyev Periodic Table, had been discovered. In all this time it has not lost its importance and still is “in the prime of life”. Enormous as its role in modem technology is, specialists do not believe that they know everything about this metal and predict a great future for it yet to come. But before going into that let us make a trip into the last century and look in at the quiet laboratory of the Swedish chemist Arvfedson. It is Sweden in the year 1817.
… For many days now the scientist has been busy analyzing the mineral petalite found at the Uto mine near Stockholm. He checks the results of his analysis again and again but each time the sum of all the components totals only 96 per cent What happens to the remaining 4 per cent? Could it be that…? Yes, this must definitely be it: the mineral contains a hitherto unknown element. Arvfedson tries again and again. Finally he is satisfied: he has discovered a new alkali metal. And since, unlike its “close relatives” potassium and sodium which were first discovered in organic products, the new element was found in a mineral, the scientist decides to gjve it the name lithium (from the Greek “lithos” for stone).
Soon Arvfedson detected this element in other min erals as well, and the Swedish chemist Berzelius identified it in the mineral waters of Karlsbad and Marienbad. As a matter of fact in our day too the mineral spring waters of Vichy in France enjoy widespread popularity owing precisely to the presence of lithium salts in the water giving it its excellent balneologic properties.
In 1855, the German chemist Bunsen, and independently of him, the British physicist Matheson, succeeded in isolating pure lithium by electrolysis of fused lithium chloride. It turned out to be a soft silvery white metal, barely half the weight of water. In lightness lithium has no rivals among metals: aluminium is five times heavier, iron, 15 and osmium 40 times heavier.
Even at room temperature lithium reacts vigorously with the nitrogen and oxygen of the air. Just try leaving a piece of lithium in a glass vessel with a ground glass plug. It will absorb all the air creating a vacuum in the vessel and atmospheric pressure will force the plug in so tightly that removing it will be a Herculean job. Therefore, storing lithium is a tough problem. While sodium can be successfully hidden in kerosene or gasoline, this will not do for lithium: it will immediately pop to the surface and flare up. To keep the reactive inclinations of lithium in check, lithium sticks are usually kept imbedded in vaseline or paraffin enveloping the sticks and playing the role of a protective coating.
Lithium combines even more readily with hydrogen. A small amount of it is capable of combining with staggering volumes of this gas: one kilogram of lithium hydrite contains 2 800 litres of hydrogen! During the Second World War American flyers were provided with lithium hydrite pellets as emergency portable hydrogen sources to be used if the plane crashed while flying above water: once in contact with water, the pellets would dissolve instantaneously releasing hydrogen to inflate rescue facilities, such as inflatable boats, life jackets, and signal balloon antennas.
The incredible ability of lithium compounds to absorb water has found extensive use as a means of purifying the air in submarines, respirators on aircraft, and air conditioners.
The first attempts to put lithium to industrial use date back to the beginning of this century. For almost a hundred years up to then, lithium had been mainly used in medicine to treat gout.
During the First World War Germany experienced a desperate need in tin for its industry. The country had no raw material to extract tin from and scientists had to make an urgent search for a replacement Lithium provided an excellent solution: a lead-lithium alloy (known as Bahnmetal) proved to be a fine antifriction material. Since then technology has constantly found use for lithium alloys, such as alloys with aluminium, beryllium, copper, zinc, and silver. There is a particularly promising future for the alloy with another lightweight metal — magnesium possessing, apart from other good properties, very valuable structural characteristics. A lithium-magnesium alloy containing not more than 50 per cent of magnesium is lighter than water. Some alloys of this compositional range have already been made. But unfortunately they are unstable, oxidizing readily in the air. At present scientists are working on the composition and production technology of an alloy which would be durable in service. A sample of a lithium-magnesium alloy which does not tarnish with time has already been displayed at the USSR Economic Exhibition in Moscow.
The high reactivity of lithium, its low melting point, and the light weight of its compounds make it an excellent degasser, deoxidizer and modifier in ferrous and nonferrous metallurgy.
In aluminium production lithium is used as a process catalyst. An addition of lithium compounds to the electrolyte bath increases the throughput of an aluminium electrolytic cell, while making it possible to lower the bath temperature and electricity consumption.
In the past the electrolyte of alkaline storage batteries consisted only of sodium hydroxide solutions. The introduction of several grams of lithium hydroxide increases the battery service life three times. The temperature range under which it remains serviceable is also considerably increased: it does not become discharged when the temperature rises to 40°C, nor does it freeze at temperatures down to 20°C below zero. A lithium-free electrolyte cannot withstand such temperatures. The new cell recently developed in Japan has one of its electrodes made of lithium. Its capacity is six to seven times greater than that of its zinc “predecessors”.
Some organic lithium compounds (stearate, palmitate and others) retain their physical characteristics within a wide temperature range, which makes it possible to use them in the production of lubricants for machines. The lithium-base lubricant enables cross-country vehicles in the Antarctic to make inland raids to regions where the air temperature is often as low as 60°C below zero. The lithium lubricant is very reliable in passenger cars as well. Owners of the Zhiguli car call it a “permanent lubricant”. It is enough to apply it once to some of the car’s rubbing parts to last them to the end of their service life.
One of the main characters in the Czechoslovak film Lemonade Joe, a parody on Hollywood hits, used to enjoy a “devil’s cocktail” which he followed up crunching a few glasses. According to eyewitness accounts, Indian Yogis do not mind partaking of the same “dish” now and then. They will chew a glass and swallow the splinters with such relish as though nothing can be more delicious. And what about you? Have you ever eaten glass? “What nonsense! Certainly not!” the reader is going to say at this point. But he will be mistaken: common glass dissolves in water. Not as readily as, say, sugar, but still it does. The most sensitive analytical balance indicates that we take about one ten-thousandth of a gram of glass with every glass of hot tea. But if salts of lanthanum, zirconium and lithium are added to the glass during manufacture, its solubility will reduce a hundred times. The glass will become resistant even to sulphuric acid.
The role of lithium in glassmaking is not exhausted with lowering the solubility of glass. Lithium-modified glasses exhibit valuable optical properties, excellent stability to heat, high electrical resistivity and low dielectric loss. Lithium, in particular, is used in the glasses from which television picture tubes are manufactured. When ordinary window glass is treated in a melt of lithium salts, the glass fonns a dense protective layer making it twice as strong and stable to elevated temperatures. Slight additions of lithium (0.5 to 1.5 per cent) lower the temperature of the glass melt considerably.
A dew drop has always been taken for a model of transparency. Yet glass transparent like a dew drop no longer satisfies modem technology which needs optical materials letting through not only visible rays of light but also invisible ones, say, ultraviolet Conventional telescopes are of little use to astronomers straining to catch the radiation of distant galaxies. Lithium fluoride possesses the highest transparency for the ultraviolet rays. Lenses made of lithium fluoride single crystals enable astronomers to penetrate deeper into the secrets of the universe.
Lithium is very valuable for making special glazes, enamels, paints, high-quality porcelain and faience. In the textile industry some lithium compounds are used for whitening and mordanting fabrics and others, for dyeing them.
Lithium salts make tracer bullets and projectiles leave a bright blue-green wake.
The following experiment demonstrates the pyro technical abilities of lithium. Try to set fire to a lump of sugar. Y ou will see that it will begin to melt but will not bum. But if the lump is first rubbed with tobacco ashes, it will readily bum with a pretty blue flame. This is to be explained by the fact that tobacco, just as many other plants, contains appreciable amounts of lithium. When tobacco leaves bum some of the lithium compounds remain deposited in the ash, making it possible to show this little chemical trick.
But everything we have discussed so far concerns only secondary, side jobs of lithium. Let us now get down to serious matters.
Scientists have found that the nucleus of the isotope lithium-6 can be easily disrupted by neutrons. On absorbing a neutron, the lithium nucleus becomes unstable md decays to form two new atoms: the light inert gas helium and the rarer superheavy hydrogen, known as tritium. At very high temperatures tritium atoms and 2 toms of hydrogen’s other heavy isotope, deuterium, combine releasing enormous quantities of energy, the thermonuclear energy.
Thermonuclear reactions are especially vigorous when neutrons bombard the compound formed by the isotope lithium-6 with deuterium, lithium deuteride. It serves as nuclear fuel in lithium reactors, which have some important advantages over uranium ones: lithium is more readily accessible and is cheaper than uranium, it does not form radioactive fission products and the process can be controlled more easily.
The considerable ability of lithium-6 to capture slow neutrons is utilized in the control of the rate of reactions occurring in uranium reactors as well. Thanks to this property, this isotope has. found application in radiation protective screens and in nuclear batteries of protracted service life. It is likely that in the near future lithium-6 will be used to absorb slow neutrons in nuclear propelled aircraft and spacecraft. Like other alkali metals, lithium is used as a coolant in nuclear installations. For this purpose the less scarce isotope lithium-7 is quite suitable (natural lithium contains about 93 per cent of it). Lithium-7 cannot serve as a raw material for the production of tritium like its lighter “brother”, hence has no significance to thermonuclear technology. But it is quite fit as a coolant. In this role it is made more efficient by its heat capacity and thermal conductivity, the large temperature range of its liquid phase (180°-1 336° Centigrade), its insignificant viscosity and low density.
Lately the rocket and aerospace industry has been making serious claims on lithium. A lot of power is required to overcome the pull of the earth’s gravity and break loose into outer space. The rocket which injected the space vehicle carrying the world’s first cosmonaut Yuri Gagarin into orbit had six engines with a total power output of 20 million horsepower! This is equivalent to the output of 20 power plants such as the Dnieper Hydroelectric Power Station.
Naturally the choice of a rocket propellant is a crucial issue. So far kerosene (yes, good old kerosene!) has been considered the most efficient fuel with liquid oxygen as the oxidizer. The calorific value of this propellant combination is 2 300 kcal per kilogram. (For comparison consider that only 1 480 kcal of heat are released when one kilogram of nitroglycerin, one of the most powerful chemical explosives known, is exploded.)
Application of metallic fuel has excellent prospects. The theory and technique of using metals as rocket propellants were first developed by the remarkable Soviet scientists Yu. V. Kondratyuk and F. A. Tsander several decades ago. Lithium is one of the metals best suited for this purpose. The combustion of one kilogram of it releases 10 270 kcal! Only beryllium has a greater calorific value. Patents on a solid rocket propellant containing 51-68 per cent of lithium metal have been published in the United States.
It is interesting to note that lithium acts against lithium during the operation of rocket engines. As a fuel component lithiium is capable of developing colossal temperatures. Lithium ceramic materials noted for their high heat resistance and refractoriness, (such as “stupalite”) are used in coatings for rocket nozzles and combustion chambers to protect them from the destructive effects of the lithium fuel.
In our time industry has at its disposal a large number of synthetic polymeric materials which can be used successfully as substitutes for steel, brass and glass. But production engineers frequently experience difficulties when polymers have to be joined with each other or with other materials. For example, the new fluorocarbon polymer known as teflon is an ideal anticorrosive coating, but until recently it could find no practical application because it could not be adhered to metal.
Recently Soviet scientists worked out an original technology for the nuclear “welding” of polymers to different materials. Small quantities of lithium or boron compounds are applied to the surfaces to be bonded as a sort of “nuclear glue”. When these surfaces are exposed to neutron bombardment, nuclear reactions develop releasing considerable energy, so that regions of the materials on a microscopic scale exhibit temperatures running to hundreds and even thousands of degrees in a very short time (less than a ten-thousand-millionth of a second). But even in those ultrashort intervals the molecules of the interface layers have time to become displaced, mix and sometimes form new chemical bonds between themselves — nuclear welding takes place.
As a rule, elements in the top left corner of the Mendeleyev Periodic Table are abundant in nature. But in contrast to most of its neighbours — sodium, potassium, magnesium, calcium and aluminium with which our planet is richly endowed — iithium is a comparatively rare element. It accounts for only 0.0065 per cent of the earth’s crust. About 20 minerals containing this valuable element are found in nature, the principal of them being spodumene (triphane), Crystals of this mineral, reminiscent in their shape of railroad cross-ties or tree trunks, sometimes have vast sizes: a crystal over 15 metres long, with a weight measured in tens of tons, was found in South Dakota (USA). Very beautiful emerald-green and pink-violet variants of spodumene, the semiprecious minerals hiddenite and kunzite have been found in American occurrences.
Granitic pegmatites, the reserves of which are practically inexhaustible, can serve as raw materials in lithium production. The estimate is that a cubic kilometre of granite contains 112 000 tons of lithium, 30 times more than all the lithium produced today in all the capitalist countries. Apart from lithium, granite deposits contain niobium, tantalum, zirconium, thorium, uranium, neodymium, cesium, cerium, praseodymium and other rare elements. But how can granite be made to share its wealth ?
Scientists are now at work searching for methods which, like the storybook words “Open, sesame!”, will open the granite storehouses of the earth. They will undoubtedly be successful in their efforts.
Before winding up our story about lithium, let us relate an amusing episode from the life of the American physicist Robert Wood in which lithium had a very important part to play. In 1891 Robert Wood, a Harvard graduate and the future celebrated scientist, arrived in Baltimore to study chemistry under the well-known Professor Remsen. He stayed at a boarding house near the University where he soon heard from other students there that the landlady often cooked their morning stew from dinner leftovers she collected from the plates the day before. But how was that to be proved?
Wood was famous for his ability to find original and simple solutions to problems. He was true to himself this time too. When a stake was served for dinner one day he left untouched several big hunks of meat which he had sprinkled with lithium chloride, an absolutely harmless substance looking and tasting much like common table salt Next day the students collected a few pieces of meat which had been served them for breakfast and examined them through a spectroscope. The red line of the spectrum produced by lithium put the dot over there. The thrifty landlady was exposed. Years later Wood recalled his “criminal investigation” with pleasure.
Source: Tales About Metals, S. Venetsky