Preparation of the Alkali Metals Because the alkali metals are among the most potent reductants known, obtaining them in pure form requires a considerable input of energy. The key steps are acid leaching, basic precipitation of aluminum hydroxide, selective precipitation of insoluble lithium carbonate, conversion to lithium chloride, evaporation, and electrolysis. The other alkali metals and the alkaline earth metals are recovered from their ores by similar processes.
A crystal of spodumene LiAlSi 2 O 6. This mineral is one of the most important lithium ores. Reacting graphite with alkali metals such as K, Rb, and Cs results in partial reduction of the graphite and insertion of layers of alkali metal cations between sets of n layers of carbon atoms.
This schematic diagram illustrates the most common structures that have been observed. This trend, which is not consistent with the relative magnitudes of the reduction potentials of the elements, serves as another example of the complex interplay of different forces and phenomena—in this case, kinetics and thermodynamics.
Although the driving force for the reaction is greatest for lithium, the heavier metals have lower melting points. The heat liberated by the reaction causes them to melt, and the larger surface area of the liquid metal in contact with water greatly accelerates the reaction rate. Liquid Ammonia Solutions A remarkable feature of the alkali metals is their ability to dissolve reversibly in liquid ammonia.
Solvated electrons. Most metals are insoluble in virtually all solvents, but the alkali metals and the heavier alkaline earth metals dissolve readily in liquid ammonia to form solvated metal cations and solvated electrons, which give the solution a deep blue color. Image copyrighted by the Klein research group.
Organometallic Compounds of the Group 1 Elements Compounds that contain a metal covalently bonded to a carbon atom of an organic species are called organometallic compounds. The structure consists of a tetrahedral arrangement of four lithium atoms, with the carbon atom of a methyl group located above the middle of each triangular face of the tetrahedron.
The carbon atoms thus bridge three lithium atoms to form four-center, two-electron bonds. Organosodium and organopotassium compounds are more ionic than organolithium compounds. Uses of the Alkali Metals Because sodium remains liquid over a wide temperature range To extinguish a fire caused by burning lithium metal, would you use water, CO 2 , N 2 gas, or sand SiO 2?
Given: application and selected alkali metals Asked for: appropriate metal for each application Strategy: Use the properties and reactivities discussed in this section to determine which alkali metal is most suitable for the indicated application.
Thus LiOH is the better choice. Lithium is a potent reductant that reacts with water to form LiOH and H 2 gas, so adding a source of hydrogen such as water to a lithium fire is likely to produce an explosion. Lithium also reacts with oxygen and nitrogen in the air to form Li 2 O and Li 3 N, respectively, so we would not expect nitrogen to extinguish a lithium fire.
Thus water, N 2 , and CO 2 are all unsuitable choices for extinguishing a lithium fire. In contrast, sand is primarily SiO 2 , which is a network solid that is not readily reduced. Smothering a lithium fire with sand is therefore the best choice. The salt with the smaller cation has the higher lattice energy, and high lattice energies tend to decrease the solubility of a salt. However, the solvation energy of the cation is also important in determining solubility, and small cations tend to have higher solvation energies.
Recall that high solvation energies tend to increase the solubility of ionic substances. Thus CsI should be the least soluble of the alkali metal iodides, and LiI the most soluble. Consequently, CsNO 3 is the better choice.
B If a reaction is predicted to occur, balance the chemical equation. Solution: A Sodium is a reductant, and oxygen is an oxidant, so a redox reaction is most likely.
Under normal reaction conditions, the product of the reaction of an alkali metal with oxygen depends on the identity of the metal. The other reactant, water, is both a weak acid and a weak base, so we can predict that an acid—base reaction will occur. A Potassium is a reductant, whereas methanol is both a weak acid and a weak base similar to water. This reaction, therefore, is an acid dissociation that is driven to completion by a reduction of the protons as they are released.
A One of the reactants is an alkali metal, a potent reductant, and the other is an alkyl halide. Any compound that contains a carbon—halogen bond can, in principle, be reduced, releasing a halide ion and forming an organometallic compound. It undergoes electron capture to produce argon; a comparison of the ratio of potassium to argon in rocks can be used to determine the age of the rock potassium-argon dating. Trace amounts of potassium are found in all sources of potassium; in a typical human, about , atoms of potassium decay every second.
The energy released by the decay of potassium is partially responsible for the interior heat of the Earth, along with the decays of thorium and uranium. There are a number of widely-used compounds of potassium. Potassium chloride, KCl, is used in salt substitutes mixed with sodium chloride to improve its flavor , and in fertilizers; massive amounts of potassium chloride are used in lethal injections to cause rapid death by cardiac arrest. Potassium carbonate, K 2 CO 3 , also known as potash, is used in the manufacture of glass.
Potassium hydroxide, KOH, also known as caustic potash, is used in making soaps and detergents. Potassium nitrate, KNO 3 , also known as saltpeter, is a powerful oxidizer, and is one of the ingredients of gunpowder.
Potassium chlorate, KClO 3 , is a very powerful oxidizer, and is used in match heads and fireworks. Potassium superoxide, KO 2 , reacts with carbon dioxide to produce potassium carbonate and oxygen gas; it is used in rebreathers and respiration equipment to generate oxygen, and is also used in mines, submarines, and spacecraft.
Rubidium is a soft, white metal; it is similar to sodium and potassium in its reaction with water, but the reaction is even more violently exothermic. Its name is derived from the Latin word for deep red ruby , rubidius. It is found in the Earth's crust at a concentration of 90 ppm, making it the 22nd most abundant element.
It is not found in any unique minerals, but is present in trace amounts in lepidolite, pollucite, carnallite, zinnwaldite, and leucite. Metallic rubidium spontaneously combusts in air. In flame tests, rubidium salts produce a reddish-violet color, and are sometimes used in fireworks. Rubidium is used in the manufacture of vacuum tubes and cathode ray tubes CRTs , and is used in some atomic clocks. Cesium undergoes the same reaction in water as lithium, sodium, and potassium, but even more violently; because cesium is a very large atom, the outermost electron is lost very easily, and the process is extremely exothermic.
The name is derived from the Latin word caesius , which means "sky blue," because salts of cesium produce a blue color when heated. Cesium is found in the Earth's crust at a concentration of 3 ppm, making it the 46th most abundant element. The main ore of cesium is pollucite [CsAlSi 2 O 6 ]; the refining of pure cesium is made even more difficult by the presence of trace amounts of rubidium in the ore, which is chemically very similar to cesium and thus difficult to separate.
Because cesium is so reactive, it is used as a "getter" to remove all traces of other gases from vacuum chambers, cathode ray tubes, and vacuum tubes. Some cesium salts give off light when exposed to X-rays and gamma rays; they are also used in photoelectric cells. Cesium is used in atomic clocks. In the SI system, a second is defined as 9,,, cycles of the radiation corresponding to the energy difference between the ground state and one of the excited states of the cesium atom.
Radioactive cesium is produced in the testing of nuclear weapons, and in nuclear power plants; the explosion at the Chernobyl power plant in released large amounts of cesium, which contaminated a great deal of Western Europe. Cesium has a half-life of 30 years, and undergoes beta-decay to produce bariumm, a metastable isotope of barium with a half-life of 2.
Since cesium ions are so heavy, research on the use of cesium in ion propulsion drives aboard spacecraft and satellites is being conducted. Francium is an extremely rare, radioactive metal. Its is named for France, the country in which it was first isolated. It is found in the Earth's crust only in trace amounts, and is one of the least abundant elements on the Earth.
Traces of it are found in uranium ores, where it is produced in the decay series of uranium; there is probably only about 20 to 30 grams of naturally-occurring francium in the entire Earth.
All of the isotopes of francium are radioactive, and most have half-lives of less than five minutes; the longest-lived isotope francium has a half-life of The possible existence of francium was predicted by Mendeleev from a gap in his periodic table, but the element wasn't discovered until , by Marguerite Perey, an assistant to Marie Curie at the Radium Institute in Paris. John Emsley, The Elements , 3rd edition. Oxford: Clarendon Press, Oxford: Oxford University Press, David L.
Heiserman, Exploring Chemical Elements and their Compounds. Resources, Conservation and Recycling , , Journal of Chemical Education , 92 1 , Verbindungen der seltenen Alkalimetalle.
Pair your accounts. Your Mendeley pairing has expired. Please reconnect. Finally, all the alkali metals are also incredibly useful teaching tools in the field of chemistry. Teachers love demonstrating the principle of reactivity by dropping an alkali metal in water only for the class to watch in awe as it spews fire and explodes. Francium is the rarest of the alkali metals and the second rarest element in the Earth's crust only grams or about 1 pound is estimated to be in the Earth's crust.
It also happens to be highly radioactive and has a maximum life of only 22 minutes. Francium has never been dropped in water, because it's so rare and so expensive, but scientists do expect it would have the highest reaction of any alkali metal.
Sign up for our Newsletter! Mobile Newsletter banner close. Mobile Newsletter chat close. Mobile Newsletter chat dots. Mobile Newsletter chat avatar. Mobile Newsletter chat subscribe. Physical Science. Chemical Elements. The alkali metals are on the left column of the periodic table highlighted in hot pink. AA alkaline batteries line up in rows. These are made with lithium, one of the alkali metals on the periodic tables.
This illustration of a cesium atomic clock shows the cesium beam tube. Cesium atomic clocks are extremely accurate.
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