NJL

The Chemistry of Sodium and its Compounds

1. Discovery and Naming of Sodium

Sodium, in the form of salt, has been known by man for many, many centuries. Indeed, man's use of salt to season food was the basis for one of the earliest examples of international trade the journeying of caravans in biblical days to Sodom, Gomorrah and nearby cities to obtain this valuable commodity from the Romans who controlled the salt deposits near the Dead Sea.

Source: Comprehensive Inorganic Chemistry (J.C. Bailar et al.) Page 390 (WSU Library QD 451.2 C64 V.1)

The Egyptians called soda natron. Much later, the Romans used a similar name for the compound, natrium. These names explain the chemical symbol used for sodium, Na. The name sodium probably originated from an Arabic word suda, meaning "headache." Soda was sometimes used as a cure for headaches among early peoples. The word suda also carried Dyer into Latin to become sodanum, which also means "headache remedy."

Source: Chemical Elements (David E. Newton) Page 542 (WSU Reference QD 466.N464 1999V.3)

2. Occurrence and Extraction of Sodium

Sir Humphry Davy first isolated metallic sodium in 1807 by the electrolytic decomposition of sodium hydroxide. Later, the metal was produced experimentally by thermal reduction of the hydroxide with iron. In 1855, commercial production was started using the Deville process, in which sodium carbonate was reduced with carbon at 1100°C. In 1886 a process for the thermal reduction of sodium hydroxide with carbon was developed. Later sodium was made on a commercial scale by the electrolysis of sodium hydroxide. The process for the electrolytic decomposition of fused sodium chloride, patented in 1924, has been the preferred process since installation of the first electrolysis cells at Niagara Falls in 1925. Sodium chloride decomposition is widely used throughout the world.

Source: Encyclopedia of chemical technology fifth Edition volume 22 page 760 (WSU Reference TP9.K54 2004 V.22)

Sodium [is an] abundant element, its principal sources being rock salt (NaCI), natural brines and sea water. Rock salt is almost pure sodium chloride, and sodium chloride is the major

component of dissolved solids in the sea to a degree sufficient for it to be readily obtainable in warm climates by evaporation of sea water.

Source: Inorganic Chemistry (A. G. Sharpe) page 228 (WSU Library QD151.2.S48 1991)

3. General Properties of Sodium

Sodium is a soft, malleable solid readily cut with a knife or extruded as wire. It is commonly coated with a layer of white sodium monoxide, carbonate, or hydroxide, depending on the degree and kind of atmospheric exposure. In a strictly anhydrous inert atmosphere, the freshly cut surface has a faintly pink, bright metallic luster. Liquid sodium in such an atmosphere looks much like mercury. Both liquid and solid oxidize in air, but traces of moisture appear to be required for the reaction to proceed. Oxidation of the liquid is accelerated by an increase in temperature, or by increased velocity of sodium through an air or oxygen environment.

Source: Encyclopedia of chemical technology fifth Edition volume 15 page 761 (WSU Reference TP9.K54 2004 V.22)

4. Uses of Sodium Metal

Metallic sodium is used in the extraction of titanium, for the production of sodium cyanide and sodium peroxide and, used in conjunction with an alcohol, for the industrial reduction of fatty acids. Sodium is utilized as a heat transfer medium in some atomic power stations. Other applications of sodium and its alloys include the familiar yellow street lights, photoelectric cells, with potassium in high-temperature thermometers and with lead as an intermediary in the manufacture of tetraethyl lead (now discredited for health reasons).

Source: Comparative Inorganic Chemistry (Bernard Moody) Page 216 (WSU Library QD 151.2.M66 1991)

Sodium [metal] was first used commercially to make aluminum by reduction of sodium aluminum chloride. The principal application as of the mid-1990s is for the manufacture of tetraethyllead (TEL), the antiknock gasoline additive. However, TEL use is declining worldwide because of the recognized toxic effects of lead released to the environment Sodium use is growing for manufacture of sodium borohydride and agricultural crop protection chemicals Smaller amounts of sodium are used to produce sodium hydride, indigo dyes, tantalum metal powders, silicon, and sodium peroxide; in the preparation of many organic compounds, pharmaceuticals sodium azide, and copper; and in lead dross refining.

Source: Encyclopedia of chemical technology fifth Edition volume 22 page 760 (WSU Reference TP9.K54 2004 V.22)

Many former uses of sodium metal have been discontinued because of the development of cheaper competitive routes. For example, the first commercial use of sodium metal in the mid-nineteenth century involved its use in the reduction of aluminum chloride to produce aluminum metal; Hall's classic work on the cheaper electrolytic route to aluminum deprived sodium of its first industrial market. Similarly, its early use to produce NaCN by fusion with potassium ferrocyanide was displaced later by the reaction of sodium with ammonia and coke, and ultimately by the direct synthesis of hydrogen cyanide as a cyanide source. Sodium metal was also used formerly in large quantities to produce fatty alcohols for synthetic detergents by reduction of natural esters; however, the use of hydrogen reduction to produce alkylaryl sulphonates (as substitutes for the more expensive fatty alcohol sulphates) has almost completely eliminated this market for sodium.

Source: Comprehensive Inorganic Chemistry (J.C. Bailar et al.) Page 373 (WSU Library QD 451.2 C64 V.1)

5. Sodium Hydride

Sodium hydride is a useful condensation catalyst. It possesses considerable advantages over the other condensing agents, such as Na, NaOH, NaOR, NaNH2: (i) it acts more vigorously and more rapidly, (ii) no relatively large excess needs to be used, (iii) neither water nor alcohol is produced, (iv) there are few side reactions and reductions, (v) the hydrogen formed serves as a measure of the extent of reaction A suspension of NaH in oil is particularly suitable, since the particles are then protected against superficial oxidation. For condensations normally requiring days, hours (and frequently only minutes) are often sufficient when they are catalyzed with suspensions of NaH in oil, and generally, in addition, the reaction temperature can also be lowered.

Source: Hydrides (Wiberg and Amberger) page 33 (WSU Library QD 181.H1.W513)

6. Sodium Chloride

Common salt has played a major role in the history of civilization. Salt was one of the earliest commodities to be traded, and Roman soldiers were partially paid in salt (sal) hence our term for wages, salary. In central Europe during the Middle Ages, the Catholic Church controlled the salt mines, a source of wealth and power. Centuries later, the salt taxes in France were part of the cause of the French Revolution.

More sodium chloride is used for chemical manufacture than any other mineral, with world consumption exceeding 150 million tones per year.

Source: Descriptive Organic Chemistry second edition (Geoff Rayner-Canham) Page 200 (WSU Library QD 151.5.R39 1999)

Sodium chloride structure: the radius ratio is 0.52 and this suggests an octahedral arrangement. Each Na⁺ ion is surrounded by six CI^- ions at the corners of a regular octahedron and similarly each Cl^- ion is surrounded by six Na⁺ ions. The coordination is thus 6:6.

Source: Concise Inorganic Chemistry Fifth Edition (J. D. Lee) pages 47-48 (WSU Library QD 453.2.L44.1996)

Rock salt, solar salt, and to some degree, evaporated salt are used to maintain traffic safety and mobility during snow and ice conditions in Snowbelt regions worldwide. Sodium chloride melts ice at temperatures down to its eutectic point of -21.12°C. Most snowstorms occur when the temperature is near O°C, where salt is very effective. More than 40% of dry salt produced in the United States is used for highway deicing.

Source: Encyclopedia of chemical technology fifth Edition volume 22 page 817 (WSU Reference TP9.K54 2004 V.22)

7. Sodium Carbonate

Sodium carbonate Na2CO₃, also known as soda ash, is produced from both natural deposits and synthetic methods based on the Solvay process. Annual world production capacity is estimated at almost 44 x 106 metric tons. It is an essential ingredient in the production of glass, chemicals, soaps, detergents, pulp and paper.

Source: Encyclopedia of chemical technology fifth Edition volume 22 page 787 (WSU Reference TP9.K54 2004 V.22)

The Solvay process essentially is a process for converting salt and limestone to soda ash and calcium chloride by a type of metathesis, although one would scarcely recognize a metathetical [double displacement] reaction from the various chemical reactions involved. In essence, the Solvay process may be said to be a commercial way of combining the sodium ion from salt with the carbonate ion from limestone to produce sodium carbonate.

Source: Comprehensive Inorganic Chemistry (J.C. Bailar et al.) Page 453 (WSU Library QD 451.2 C64 V.1)

[I know this isn’t directly Na related, however I found it interesting and thought you might as well]

World production of Na₂CO₃ in 1993 was 31.5 million tones, and 49% of this was used in the glass industry. Smaller amounts were used to make various sodium phosphates and polyphosphates which are used for water softening (being added to various cleaning powders), and in wood pulp and paper making. The increased awareness of the effect of 'acid rain' on plants and buildings has led to a new use for Na₂CO₃ in treating the flue gases from coal- and oil-fired power stations, to remove SO₂ and H₂SO₄.

Source: Concise Inorganic Chemistry fifth edition (J.D. Lee) page 322 (WSU Library QD 453.2.L44 1996)

Sodium hydrogen carbonate [sodium bicarbonate] is less water soluble than sodium carbonate.

Thus it can be prepared by bubbling carbon dioxide through a saturated solution of the carbonate:

Na₂CO_3(aq) + CO_{2 (g)} + H₂0_(l) ↔ 2 NaHCO_3(s)

Heating sodium hydrogen carbonate causes it to decompose back to sodium carbonate:

2 NaHCO_3(s) ↔ Na₂CO_3(aq) + CO_{2 (g)} + H₂0_(g)

This reaction provides one application of sodium hydrogen carbonate, the major component in dry powder fire extinguishers. The sodium hydrogen carbonate powder itself smothers the fire, but, in addition, the solid decomposes to give carbon dioxide and water vapor, themselves fire-extinguishing gases.

The main use of sodium hydrogen carbonate is in the food industry, to cause bakery products to rise. It is commonly used as a mixture (baking powder) of sodium hydrogen carbonate and calcium dihydrogen phosphate, Ca(H₂PO4)₂ with some starch added as a filler. The calcium dihydrogen phosphate is acidic and, when moistened, reacts with the sodium hydrogen carbonate to generate carbon dioxide:

2 NaHCO_3(s) + Ca(H₂PO4)_2(S) -+ Na₂HPO_4(s) + CaHPO_4(s) + 2 CO_2(g) + 2 H₂O_(g)

Source: Descriptive Organic Chemistry second edition (Geoff Rayner-Canham) Page 200 (WSU Library QD 151.5.R39 1999)

8. Sodium in the Body

[Sodium is] among the 25 or so elements now thought to be essential for animal life. It performs electrophysiological functions based on the fact that the Na⁺:K⁺ concentration ratio is different in the fluids inside and outside cells. Due to these concentration gradients across cell membranes, a potential difference is set up and this is responsible for the transmission of nerve impulses.

Source: Encyclopedia of chemical technology fifth Edition volume 1 page 36 (WSU Reference TP9.K54 2004 V.1)

9. Sodium Ions

It is so easy to become locked into preconceptions. Everyone "knows" that the alkali metals "want" to lose an electron and form cations. In fact, this is not true. Left to itself, the alkali metal would prefer to complete its s orbital set by gaining an electron.

It was a Michigan State chemist, James Dye, who realized that the alkali metals have such positive electron affinities that it might just be possible to stabilize the alkali metal anion. After a number of attempts, he found a complex organic compound of formula C2oH3606 that

could just contain a sodium cation within its structure. He was hoping that, by adding this compound to a sample of sodium metal, some of the sodium atoms would pass their s electrons to neighboring sodium atoms, to produce sodium anions. This happened, as predicted:

2 Na_(s) + C₂₀H₃₆0₆ ^- [Na(C₂₀H₃₆0₆)]⁺•Na^-

The metallic-looking crystals were shown to contain the sodium anion, but the compound was found to be very reactive with almost everything. So, to the present day, this compound is no more than a laboratory curiosity. But its existence does remind us to question even the most commonly held beliefs.

Source: Descriptive Organic Chemistry second edition (Geoff Rayner-Canham) Page 183 (WSU Library QD 151.5.R39 1999)

[I included this section because the most common mention of Na in the literature was simply as a cation. That really seems to be all that most people, myself included, think about when they consider sodium]

10. Sodium Hydroxide

Caustic soda [sodium hydroxide] is the strongest alkali commonly manufactured. It is produced in various grades. Much of the total production of very pure alkali goes into the rayon industry and it is supplied to textile manufacturers for the mercerization of cotton, bleaching and dyeing processes and in paper-making; The hydrolysis (saponification) of natural oils and fats, by boiling with caustic soda solution, produces soap, as a mixture of sodium salts of certain organic acids (e.g. sodium stearate) and the trihydric alcohol, glycerol. Caustic soda is used in the removal of acidic compounds, such as phenol and the cresols, during the refining of coal tar. At least 400 products of the chemical industry require the use of caustic soda at some stage of their manufacture.

Source: Comparative Inorganic Chemistry (Bernard Moody) Page 208 (WSU Library QD 151.2.M66 1991)

Alkali and chlorine products are a group of commodity chemicals which include chlorine Cl₂; sodium hydroxide (caustic soda), NaOH; sodium carbonate (soda ash), Na2CO3; potassium hydroxide (caustic potash) KOH; and hydrochloric acid HCI. Chlorine and caustic soda are the two most important products in this group, ranking among the top ten chemicals in the United States.

Electrolysis of sodium chloride accounts for nearly all of today's installed capacity for sodium hydroxide.

2 NaCl + 2 H₂O --+ 2 NaOH + Cl₂ + H₂

As shown, chlorine is coproduced, so companies that are in the sodium hydroxide business are also usually involved in the chlorine business.

Source: Encyclopedia of chemical technology fifth Edition volume 22 page 760 (WSU Reference TP9.K54 2004 V.22)

Lithium and its Compounds

Lithium and its Compounds

1. Discovery and naming of Lithium

Lithium was discovered in 1817 by J.A. Arfvedson at Stockholm Sweden. Isolated in 1821 by W.T. Brande

Source: The Elements(John Emsley) Page 116; (WSU Reference QD 466.e48 1999)

The lithium containing materials petalite and spodumene, were discovered by Jose de Andrada between 1790 and 1800 on Utö island in Sweden. In 1817 J.A. Arfvedson discovered the new alkali metal lithium in petalite. He noted that lithium carbonate is sparingly soluble , that the hydroxide is much less soluble than the hydrogen oxides of other alkali metals and that lithium compounds are similar to those of other alkaline earth metals.

Source: Comprehensive Inorganic Chemistry (J.C. Bailar et al.) Page 331 (WSU Library QD 451.2 C64 V.1)

The Name Lithium was given to the new element in recognition of the fact that lithium was recovered first from a mineral whereas both sodium and potassium were first derived from plant matter.

Source: Comprehensive Inorganic Chemistry (J.C. Bailar et al.) Page 331 (WSU Library QD 451.2 C64 V.1)

[from the journal of J.A. Arfvedson] It’s solution could not be precipitated either by tartaric acid in excess or by platinum chloride. Consequently it could not be potassium. I mixed another portion of the solution with a few drops of pure potash but without it’s becoming cloudy. Therefore it contained no magnesia. … At last having studied more closely the sulfate in question I soon found that it contained a definite fixed alkali whose nature had not been previously known.

Source: The chemical elements (H.M. Davis and G. Seaborg) page 22 (WSU Library QD 466.d35 1961)

2. Occurrence and extraction of Lithium

The abundance of lithium in the earth’s crust is estimated to be about 0.005 percent. That places it among the top 15 elements found in the earth. The most common ores of lithium are spodumme, petalite, and lepidolite. Lithium is also obtained from salt water. … The worlds largest producer of lithium is the United States

Source: Chemical Elements (David e.Newton) Pages 309-310 (WSU Reference QD 466.N464 1999V.2)

Melt [lithium chloride] and then insert the electrodes so that they dip well into the solution. Shortly after you have immersed the electrodes a silver substance will form on the cathode. This is the metal lithium.

Source: Chemistry magic (Sweezey) page 72 (WSU Library QD 43 S8)

3. Uses of Lithium

The first commercial use of lithium occurred towards the end of World War I when small amounts were used in an aluminum zinc allow, Scleron. After the war lithium was used as a hardener in a lead bearing alloy-bearing material. Between world Wars I and II there was little production of lithium materials. Most lithium trade was as ores sold as additives of frit and glass formulators. ... LiH was used in military sea rescue equipment as a source of hydrogen upon reaction with water. the hydrogen inflated rescue balloons to carry radio antennas needed for the SOS signal broadcast. During this period all temperature greases using lithium stearate were produced for military applications.

Source: Encyclopedia of chemical technology fifth Edition volume 15 pages 120 to 121 (WSU Reference TP9.K54 2004 V.15)

4. General Properties of Lithium

The group has been much studied as it best illustrated the effects of increasing atomic

size and mass on chemical and physical properties.

Source: Encyclopedia of inorganic chemistry (R. Bruce King) Pages 35-36 (WSU Library QD 148.E53 1994 V.1)

Lithium is a very soft silvery metal. It has a melting point of 180.54°C and a boiling point of about 1335°C. Its density is 0.534 grams per cubic centimeter. … Lithium’s hardness is 0.6 on the Mohs scale. Lithium is an active element but not as active as the other alkali metals. It reacts slowly with water at room temperature. Lithium does not react with oxygen at room temperature. Under proper conditions lithium also combines with sulfur, hydrogen, nitrogen, and the halogens.

Source: Chemical Elements (David E. Newton) Page 308 (WSU Reference QD 466.N464 1999V.2)

One of the most interesting properties [of lithium] is the extremely low density 0.534 gm/cm3. At normal temperatures lithium has the lowest density of a non-gaseous element. The metal is harder and has a higher melting point than the other alkali metals.

Above -117 °C lithium metal has a body centered cubic crystal structure typical of the alkali metals. On cooling to -201°C the metal begins to transform to the face centered cubic structure.

Source: Comprehensive Inorganic Chemistry (J.C. Bailar et al.) Page 337(WSU Library QD 451.2 C64 V.1)

Just as the properties of lithium set it apart [physically] from the other alkali metals these same properties cause the compounds of lithium to differ from the other alkali metal compounds. ... Several general statements may be useful in the understanding of lithium compounds. The lattice energies of ionic lithium compounds are greater than the lattice energies of the corresponding salts of the other alkali metals

Source: The Elements(John Emsley) Page 116; (WSU Reference QD 466.e48 1999)

5. Aqueous chemistry of Lithium

There are numerous crystal structures of the Li+ ion with crown ethers. .. In fact most of these crystal structures are misleading in that the Li+ ion forms no complexes, or at best very weak complexes, with the familiar crown ethers such as 12-crown-4 in any known solvent

Source: Metal Complexes in Aqueous Solutions (A E Martell, R D Hancock) introduction page v (WSU library QD474.M375 1996)

6. Lithium-oxygen compounds

Lithium superoxide is a major product of the co-condensation reaction of an atomic beam of lithium with a jet of molecular oxygen at 4.2-15K.

Source: Inorganic Chemistry of the main Group Elements-volume 1(CC Addison et al) page 14 (WSU Library QD 146.I54 V.1)

The compound Li4XeO6·2H2O is Amorphous to X-rays, unlike the analogous sodium and potassium compounds, and is less stable but shows the absorption bands in the Ir as 640-720 and 2900-3600 cm-1 which are due to the XeO6- ion and H2O respectively.

Source: Inorganic Chemistry of the main Group Elements-volume 1(CC Addison et al) page 16 (WSU Library QD 146.I54 V.1)

7. Lithium-hydrogen compounds

Lithium hydride is prepared both industrially and in the laboratory by the reaction of molten lithium metal with hydrogen gas. Industrially lithium metal is heated to initiate its reaction with hydrogen gas.

Lithium hydride is also used commercially in the preparation of lithium aluminum hydride LiAlH₄ and to prepare an intermediate used in the synthesis of vitamin A.

LiH is a very strong base. It is insoluble in solvents except those where it reacts to form a soluble compound.

Source: Comprehensive Inorganic Chemistry (J.C. Bailar et al.) Page 331 (WSU Library QD 451.2 C64 V.1)

8. Lithium in the synthesis of organic compounds

The discovery by F.W. Stavely that metallic lithium and many organo-lithium compounds initiate the polymerization of isoprene to for an elastomer of nearly identical properties to those of natural rubber began a new era in rubber technology.

There is a broad range of polymerizations which can be initiated by lithium or it’s compounds. The most commonly used catalyst is butyl-lithium. The quantity of Li in the final product is of course minute.

Source: Specialty inorganic chemicals (R. Thompson) pages 112-113 (WSU library QD 151.2.S664)

Organolithium compounds are central to many aspects of synthetic organic chemistry and are primarily used as carbanions to construct carbon skeletons of a wide variety of organic compounds. Tremendous efforts have been devoted to the development of convenient methods of for generation of tailor made organolithium compounds.

Source: Main group metals in organic synthesis (H. Yamamoto, K Oshima) page 1 (WSU Library QD 411.7.S94 M33 2004 v.1)

Although [organo-lithium compounds] do not contain free carbanions they act as if they did. They are sources of R:^-. The very polar carbon-metal bond attacks all manner of lewis and bronstead acids. Water is more than sufficiently stron enough to protonate a Grignard or organo-lithium reagent.

The end product is a hydrocarbon. The sequence constitutes a new synthesis of hydrocarbons from alkyl, alkenyl, or aryl halides. … This reaction can be put to good use in the construction of isotypically labeled reagents because D₂O can be used in place of water to produce specifically deuteriolabled hydrocarbons.

CH₃CH₂Br + Li CH₃CH₂Li+D₂OCH₃CH₂D

Source: Organic Chemistry second edition (M. Jones) pages 734-735 (Stoker Library)

9. Lithium electrochemical cells

Lithium is the most electropositive of the metals with a standard electrode potential of 3.045v… Consequently Li metal cells have the potential to store the largest amount of power/unit weight or volume. … The greatest potential of course lies in the development of the electric vehicle and off-peak power storage.

Source: Specialty inorganic chemicals (R. Thompson) page 117 (WSU library QD 151.2.S664)

The lithium rechargeable battery used in portable computers and phones uses LiCoO₂ as the anode with a graphite cathode. Lithium ions are produced at the cathode and to maintain charge balance Co(III) is oxidized to Co(IV) at the anode.

LiCoO₂Li_1-xCoO₂+xLi⁺(solvent)+xe⁺

The lithium ions are intercalated into the graphite and return to the LiCoO₂ when the cell is discharged. The battery is rechargeable because for the cathode and the anode can act as hosts for the Li⁺ ions which can move back and forth between them.

Source: Inorganic Chemistry (Shriver & Atkins) page 261 (Stoker Library)

10. Toxicity of Lithium

Lithium is moderately toxic by ingestion but there are wide variations of tolerance. Even Lithium Carbonate which is used in psychiatry is prescribed at doses near to the toxic level. Some Lithium compounds are carcinogenic and teratogenic.

Source: The Elements (John Emsley) Page 116; (WSU Reference QD 466.e48 1999)

Lithium oxides, hydroxide, and carbonate are strong bases and their water solutions are very caustic. The toxicity of lithium compounds is a function of their solubility in water. Lithium ion is toxic to the central nervous system. Lithium is commonly ingested at dosages of 0.5g/d of lithium carbonate for treatment of bipolar disorders. Therapeutic vs. Toxic levels are small: 2mmol/L is toxic 4mmol/L is fatal. LiOH either directly or formed by the hydrolysis of other salts can cause caustic burns. Skin contact with lithium halides can result in skin dehydration. Organo-lithium compounds are often pyrophoric and require special handling.

Source: Encyclopedia of chemical technology fifth Edition volume 15 pages 134 to 135 (WSU Reference TP9.K54 2004 V.15)

Lithium sources

1. Source: Chemical Elements (David e.Newton) Pages 309-310 (WSU Reference QD 466.N464 1999V.2)

2. Source: Chemistry magic (Sweezey) page 72 (WSU Library QD 43 S8)

3. Source: Comprehensive Inorganic Chemistry (J.C. Bailar et al.) Page 331 (WSU Library QD 451.2 C64 V.1)

4. Source: Encyclopedia of chemical technology fifth Edition volume 15 pages 120 to 121 (WSU Reference TP9.K54 2004 V.15)

5. Source: Encyclopedia of inorganic chemistry (R. Bruce King) Pages 35-36 (WSU Library QD 148.E53 1994 V.1)

6. Source: Inorganic Chemistry of the main Group Elements-volume 1(CC Addison et al) page 14 (WSU Library QD 146.I54 V.1)

7. Source: Metal Complexes in Aqueous Solutions (A E Martell, R D Hancock) introduction page v (WSU library QD474.M375 1996)

8. Source: Specialty inorganic chemicals (R. Thompson) pages 112-113 (WSU library QD 151.2.S664)

9. Source: The chemical elements (H.M. Davis and G. Seaborg) page 22 (WSU Library QD 466.d35 1961)

10. Source: The Elements(John Emsley) Page 116; (WSU Reference QD 466.e48 1999)

11. Source: Inorganic Chemistry fourth edition (Shriver & Atkins) page 261 (Stoker Library)

12. Source: Organic Chemistry second edition (M. Jones) pages 734-735 (Stoker Library)

NJL

Magnesium and its Compounds

1. Discovery and Naming of Magnesium

In 1808 Sir Humphrey Davy reported the production of magnesium in the form of amalgam by electrolytic reduction of its oxide using a mercury cathode. In 1828 the French scientist A. Bussy fused magnesium chloride with metallic potassium and became the first to produce free magnesium. Michael Faraday in 1833 was the first to produce magnesium by electrolytic reduction from the chloride...

Source: Kirk-Othmer Encyclopedia Of Chemical Technology fifth Edition volume 15 page 320 (WSU Reference TP9.K54 2004 V.15)

The name magnesium goes back many centuries. It was selected in honor of a region in Greece known as Magnesia. The region contains large supplies of magnesium compounds.

Source: Chemical Elements (David E Newton) Pages 309-310 (WSU Reference QD 466.N464 1999V.2)

Three different substances were described during antiquity by terms related to our word "magnesia" (named from a district in Thessaly Greece). One, magnesius lapis, referred to magnetite, and another, magnesia niger, to pyrolusite. The third, magnesia, was probably applied to the soft white mineral steatite, also known as soapstone or talc, a hydrated magnesium silicate.

Source: Comprehensive Inorganic Chemistry (J.C. Bailar et al.) Page 593 (WSU Library QD 451.2 C64 V.1)

2. Occurrence and Extraction of Magnesium

Magnesium is the eighth most abundant element in the earth’s crust. It occurs naturally in a number of minerals such as dolomite CaCO3MgCO3 and magnesite MgCO3 and is the third most abundant element in seawater from which it is commonly extracted. The extraction from seawater relies on the fact that magnesium hydroxide is less soluble than calcium hydroxide, because the solubility of the salts of mononegative anions increases down the group. Either CaO (quick lime) or Ca(OH)2 (slaked lime) is added to the seawater and Mg(OH)2 precipitates. The hydroxide is converted to the chloride by treatment with hydrochloric acid.

The magnesium is then extracted by electrolysis of molten magnesium chloride.

Source: Inorganic Chemistry fourth edition (Shriver & Atkins) page 276 (Stoker Library)

Mg was formerly prepared by heating MgO and C to 2000 °C, at which temperature C reduces MgO. The gaseous mixture of Mg and CO was then cooled very rapidly to deposit the metal. This 'quenching' or 'shock-cooling' was necessary as the reaction is reversible, and if cooled slowly the reaction will come to equilibrium further to the left. MgO + C ↔ Mg + CO Magnesium is now produced by high temperature reduction, and by electrolysis.

Source: Concise Inorganic Chemistry fifth edition (J.D. Lee) page 328 (WSU Library QD 453.2.L44 1996)

3. Uses of Magnesium

Magnesium is the only Group 2 metal to be produced on a large scale. World production was 303000 tones in 1993. The largest producers were the USA 48%, the Soviet Union 12%, Canada 9% and Norway 8%. Mg is an extremely important lightweight structural metal because of its low density (1.74gcm-3 compared with steel 7.8gcm-3 or aluminum 2.7 g cm-3). Mg forms many binary alloys, often containing up to 9% AI, 3% Zn and 1 % Mn; traces of the lanthanides praseodymium Pr and neodymium Nd, and traces of thorium. The metal and its alloys can be cast, machined and welded quite easily. It is used to make the bodies of aircraft, aircraft parts and motor car engines. Up to 5% Mg is usually added to Al to improve its properties.

Source: Concise Inorganic Chemistry fifth edition page 328 (WSU Library QD 453.2.L44 1996)

Perhaps the best know magnesium compound is magnesium sulfate (MgSO4). It is popularly known as Epson salts.

One of the earliest stories about Epsom salts dates back to 1618. The town of Epsom, in Surrey, England, was suffering from a severe drought. A farmer named Henry Wicker brought his cattle to drink from a water hole on the town commons (central park). But the cattle would not drink the water. Wicker was surprised because he knew they were very thirsty. He tasted the water himself and found that it was very bitter. The bitterness was due to magnesium sulfate in the water. This compound became known as Epsom salts.

People soon learned that soaking in the natural waters that contained Epsom salts made them feel better. The salts seemed to have properties that soothed the body. Before long, soaking in these waters became very popular.

Source: Chemical Elements (David E Newton) Pages 309-310 (WSU Reference QD 466.N464 1999V.2)

4. General Properties of Magnesium

Magnesium is the easiest of all structural metals to machine (108). Because of this machine ability, it is sometimes used in applications where a large number of machining operations are required. The machine ability of magnesium alloys relative to that of other metals based on the lowest power required to remove 16 cm3 of metal and when magnesium is assigned a value of unity, is (108): Some of the advantages of excellent machine ability include reduced machining time, resulting in higher productivity for the machine tools and thus lower capital investment; greatly increased tool life; an excellent surface finish with a single large cut; well-broken chips which minimize handling costs; and less tool buildup.

If a coolant or cutting fluid must be used for a particular operation, mineral oil is usually preferred over water-based coolants. This is due to the fact that magnesium reacts with water to some degree, over time, producing hydrogen gas and magnesium hydroxide. As a result, wet chips should be treated with caution in order to prevent ignition of the evolved hydrogen, and to prevent partial drying of the wet mass, which may result in spontaneous ignition due to the heat evolved in the reaction in combination with poor heat-transfer characteristics of the partially dried mass. If wet chips are generated, the chips should be kept fully submerged in excess water.

Source: Kirk-Othmer Encyclopedia Of Chemical Technology fifth Edition volume 15 page 368 (WSU Reference TP9.K54 2004 V.15)

5. Magnesium Alloys

Aluminum alloys are the largest single consumer of magnesium. In 1992 over 50 percent of the magnesium shipped was consumed in this way. The aluminum beverage can contains about 4.5 percent magnesium in the lid and 1.1 percent in the can body. When combined with ferrosilicon, magnesium produces ductile iron. Automotive engines are a major consumer of ductile iron.

Source: MacMillian Encyclopedia of Chemistry (Lagowski) Page 879 (WSU Reference QD 4.M33 1997 V.3)

Magnesium has become important as a sttuctural material. Its great advantage is that it is very light, with a density of only 1.74 grams/cm3. For comparison's sake, water has a density of 1 gram/cm3. Magnesium has a density that is only about one-fifth that of iron and two-thirds that of aluminum. It is usually mixed with these metals to form an alloy. When alloyed with aluminum, magne- sium makes the latter metal stronger, lighter, and even more corrosion-resistant than it normally is. Many people, for example, have aluminum-magnesium alloy ladders in their homes. Its light weight also makes it ideal for fabricating automobile and aircraft parts, as well as power tools, lawn mower housings, and racing bikes.

Source: A Guide to the Elements second edition (Stwertka) Pages 57-58

6. Magnesium-Oxygen Compounds

In a close packing of oxide ions, the cations are inserted where, in accordance with their size, room is to be found. The inclusion of different cations thus changes the dimensions of the lattice of the oxide ions but little. The spaces in the latter are of two kinds, octahedral and tetrahedral. The octahedral spaces are the larger, so that the larger cations find room in them. The octahedral space provides room for the Mg2+ … ions. In order that such may enter a silicate lattice, larger spaces must be present between the closely packed oxide ions. These may be created either by an open arrangement in places, as in the feldspars or by layer structures.

Source: Structural Chemisry of inorganic compounds (Huckel) Page 772 (WSU Library QD 151 H833 V.2)

Magnesium Oxide, MgO, also known as magnesia, occurs in nature only infrequently as the mineral periclase, most commonly as a group of crystals in marble. Refractory magnesia represents the largest tonnage use of magnesium in compounds. The iron and steel industry is the largest consumer of these products in the United States and most other magnesia-consuming countries. Dead-burned magnesia from magnesite, seawater, or well and lake brines is used as a principal constituent in metallurgical furnace refractory products.

Source: Kirk-Othmer Encyclopedia of Chemical Technology fifth Edition volume 15 page 413 (WSU Reference TP9.K54 2004 V.15)

7. Magnesium-Halide Compounds

Chlorides Magnesium chloride is among the industrially most important of the magnesium salts. It is marketed in the anhydrous and hexahydrated forms. Both are deliquescent and form saturated solutions if left standing in a moist atmosphere.

Anhydrous magnesium chloride can be prepared by the direct chlorination of magnesium oxide, but it is usually prepared from a hydrate. The dehydration cannot be simply accomplished with heat because decomposition to the oxychloride occurs. After all the water has been removed, the compound is melted at 712°.

Approaching half a million tons of MgCl₂ is produced in the United States each year. Aside from its use as a source of the metal, it is used in oxychloride cements and wall plasters. It is also used in fire-proofing and dust-laying compositions.

Magnesium bromide is found in sea water, some mineral springs, the Dead Sea, the Utah Salt Lake and in deposits such as those at Strassfurt. At temperatures down to O°C, it crystallizes from aqueous solution as the hexahydrate which is very hygroscopic.

Magnesium iodide chemistry is very similar to that of the bromide.

Source: Comprehensive Inorganic Chemistry (J.C. Bailar et al.) Page 629-635 (WSU Library QD 451.2 C64 V.1)

Dicyclopentadienylmagnesium (magnesium cyclopentadienide), (C₅H₅)₂Mg, is a pyrophoric white fluffy solid that can be made directly from magnesium metal and cyclopentadiene at elevated temperatures: 500°C

Mg + 2C₅H₆ ↔ (C₅H₅)₂Mg + H₂

Dicyclopentadienylmagnesium is a useful reagent for the preparation of cycloIpentadienyl transition metal derivatives.

Source: Inorganic Chemistry of Main Group Elements (R. Bruce King) Page 271 (WSU Library QD 151.2.K5 1995 c2)

8. Magnesium in the Synthesis of Organic Compounds

[Grignard reagents] a very large and important group of alkyl and aryl magnesium halides together with the magnesium dialkyls and diaryls. There are two binary carbides of magnesium, which cannot be got pure enough for analysis, but from their behavior must have the compositions MgC2 and Mg2Ca. Neither can be made by heating magnesium or its oxide with carbon to a high temperature, and both are thermodynamically unstable above 800°.

In most cases the original magnesium compound is not isolated, but the solution immediately used for the synthesis, the second organic component, usually dissolved in ether or benzene, being added as soon as the magnesium has gone into solution. The reactivity essentially depends on the fact that magnesium has a much weaker affinity for carbon than for oxygen, nitro- gen, or the halogens, so that it very readily exchanges its hydrocarbon radical for one of the other elements. The number of reactions which the Grignard reagents can undergo is enormous.

Source: Chemical Elements and Their Compounds (Sidgwick) pages 224-225 (WSU Library QD 466 S53 v.1)

9. Magnesium in Life

An adult body contains about 25 g of magnesium, 60% of which is present in the skeleton while most of the remainder resides inside cells where it is the next most important cation after potassium. The daily requirement in the human diet is not known with certainty but has been estimated to be 200-300 mg.

All enzymes that utilize ATP in phosphate transfer, and many which aid either transfer of other groups or hydrolysis, require magnesium as the cofactor. For example, in phosphorylation reactions Mg²⁺ ion is thought to complex with oxygen on the various phosphate groups. As it is the smallest, doubly-charged cation which is biologically available, magnesium will normally form the strongest complexes of this type.

Source: Main Group Chemistry (A.G. Massey) page 142 (WSU Library QD 151.2.M37 2000)

Magnesium also plays a crucial role in the makeup of the green chlorophylls present in all green plant cells. The ability of the chlorophyll to capture solar energy and convert it by photosynthesis to energy is the ultimate source of all biological energy. The structure of the chlorophyll molecule, a ring of complex atomic structures around the magnesium atom, endows the molecule with its deep color and its ability to absorb light.

Source: A Guide to the Elements second edition (Stwertka) Pages 58

10. Magnesium and the Treatment of Waste Water

The largest use for magnesium hydroxide in the United States is for environmental applications. This portion of the market, which includes industrial water treatment, heavy-metals removal, and fluegas desulfurization, accounts for about more than half of total U.S. consumption.

For water treatment, magnesium hydroxide is supplied as a suspension containing about 58% solids, and it is used primarily to lower the pH of acidic solutions. In this market it competes with other acid-neutralizing compounds; the most common are lime and caustic soda. Magnesium hydroxide has advantages and disadvantages when compared to the other materials in this use. One of the advantages is that it is a pH buffer, and wastewater treated with it will not exceed a pH of 9.5 even if excess magnesium hydroxide is added. In contrast, excess addition of lime can raise the pH to 12,

Source: Kirk-Othmer Encyclopedia of Chemical Technology fifth Edition volume 15 page 406 (WSU Reference TP9.K54 2004 V.15)

Silica may be adsorbed on solid magnesium oxide or hydroxide, or on crushed dolomite. This is not a stoichiometric reaction involving the whole solid phase, but occurs to a variable extent on the surface only. The extent of adsorption is adequately expressed by the Freundlich adsorytion isotherm*: mass SiO2 adsorbed. This procedure can be used in conjunction with lime-soda softening since this results in a precipitate of Mg(OH)2. Effective adsorption of silica onto the Mg(OH)2 requires a temperature of at least 50 °C and preferably 100-110°C (the hot lime-soda process), and it may be necessary to add MgSO4 to get enough Mg(OH)2.

Source: Inorganic Chemistry and Industrial and Environmental Perspective (T.W. Swaddle) pages 274-275 (WSU Library QD 151.5.S93 1997)