Uses
Barium azide is used in explosives. A saturated solution is generally used.

Physical Properties
Colorless monoclinic crystal; density 2.936 g/cm3; decomposes at 120°C; soluble in water, slightly soluble in ethanol.

Preparation
Barium azide may be prepared by reacting sodium azide with a soluble barium salt. The solution is concentrated to allow crystals grow. Crystals will explode if fully dried, or subject to friction. Product should be stored damp with ethanol.

Hazard
The dry solid is sensitive to shock, impact and friction. It decomposes explosively when heated to 275°C. Contact with acid can produce the explosive compound hydrazoic acid. Contact with lead, silver, and many other metals can form the explosive azides of those metals. Presence of sodium, potassium, barium and iron ions as impurities can enhance the shock sensitivity of barium azide. Barium azide also is a toxic compound. The toxic effects are similar to those of other soluble salts of barium.

Uses
Barium acetate is used as a mordant for printing textile fabrics; for drying paints and varnishes; in lubricating oil; in the preparation of other acetates; and as a catalyst in organic synthesis.

Physical Properties
White powdery solid; density 2.47g/cm3; decomposes on heating; highly soluble in water (55.8g /100g at 0°C), sparingly soluble in methanol (~1.43 g per liter).

Preparation
Barium acetate is made by the reaction of barium carbonate with acetic acid:

BaCO3 + 2CH3COOH → (CH3COO)2Ba + CO2 + H2O

The solution is concentrated and the anhydrous barium acetate crystallizes at a temperature above 41°C. At temperatures between 25 to 40°C, barium acetate monohydrate, Ba(C2H3O2)2•H2O [5908–64–5] (density 2.19 g/cm3) crystallizes out of solution.
Barium acetate also may be prepared by treating barium sulfide with acetic acid, followed by slow evaporation and subsequent crystallization of the salt from the solution:

BaS + 2CH3COOH → (CH3COO)2Ba + H2S

Reactions
Barium acetate converts to barium carbonate when heated in air at elevated temperatures. Reaction with sulfuric acid gives barium sulfate; with hydrochloric acid and nitric acid, the chloride and nitrate salts are obtained after evaporation of the solutions. It undergoes double decomposition reactions with salts of several metals. For example, it forms ferrous acetate when treated with ferrous sulfate solution and mercurous acetate when mixed with mercurous nitrate solution acidified with nitric acid. It reacts with oxalic acid forming barium oxalate.

Analysis
Elemental composition: Ba 53.77%, C 18.81%, H 2.37%, O 25.05%. The salt may be digested with nitric acid, diluted appropriately, and analyzed for barium. (See Barium.)

Toxicity
The salt or its aqueous solution is highly toxic. LD10 (oral) rabbit: 236 mg/kg; LD10 (subcutaneous) rabbit: 96 mg/kg. See Barium.

Occurrence
Barium was discovered in 1808 by Sir Humphrey Davy. Its abundance in the earth’s crust is about 0.0425% (425 mg/kg). The element also is found in sea water at trace concentration, 13 μg/L. It occurs in the minerals barite or heavy spar (as sulfate) and witherite (as carbonate).

Uses
The most important use of barium is as a scavenger in electronic tubes. The metal, often in powder form or as an alloy with aluminum, is employed to remove the last traces of gases from vacuum and television picture tubes. Alloys of barium have numerous applications. It is incorporated to lead alloy grids of acid batteries for better performance; and added to molten steel and metals in deoxidizing alloys to lower the oxygen content. Thin films of barium are used as lubricant suitable at high temperatures on the rotors of anodes in vacuum X-ray tubes and on alloys used for spark plugs. A few radioactive isotopes of this element find applications in nuclear reactions and spectrometry.

Physical Properties
Silvery-white metal; soft and ductile; density 3.51 g/cm3; melts at 727° C; vaporizes at 1897°C; vapor pressure 0.1 torr at 730°C; electrical resistivity 34.0 microohm-cm at 25°C; reacts with water.

Thermochemical Properties
ΔH°ƒ (cry) 0.0 kcal/mol
ΔH°ƒ (gas) 43.04 kcal/mol
G° ƒ (gas) 34.93 kcal/mol
S° (gas) 40.70 cal/degree mol
Cρ (gas) 4.97 cal/degree mol

Manufacture
The metal is obtained by the reduction of barium oxide with finely divided aluminum at temperatures between 1,100 to 1,200°C:

4 BaO + 2 Al → BaO•Al2O3 + 3Ba (gas)

Barium vapor is cooled by means of a water jacket and condensed into the solid metal. The solid block may be cast into rods or extruded into wires. Being a flammable solid, it is packaged under argon in steel containers or plastic bags.

Reactions
Barium metal reacts exothermically with oxygen at ambient temperatures forming barium oxide. The reaction is violent when the metal is present in powder form. It also reacts violently with water forming barium hydroxide and liberating hydrogen:

Ba + 2H2O → Ba(OH)2 + H2

Barium reacts violently with dilute acids, evolving hydrogen. Reactions with halogens give barium halides:

Ba + Cl2 → BaCl2

Barium is a strong reducing agent. The E° for the reaction:

Ba2+ (aq) + 2e– ←→ Ba(s) is – 2.90 V

It reduces oxidizing agents reacting violently. The metal combines with nitrogen and hydrogen at elevated temperatures producing barium nitride, Ba3N2, and barium hydride, BaH2, respectively.
Barium reduces oxides, chlorides and sulfides of less reactive metals producing the corresponding metals; e.g.,

Ba + CdO → BaO + Cd
Ba + ZnCl2 → BaCl2 + Zn
3Ba + Al2S3 → 3BaS + 2Al

When heated with nitrogen in the presence of carbon, it forms barium cyanide:

Ba + N2 + 2C → Ba(CN)2

Barium combines with several metals including aluminum, zinc, lead, and tin, forming a wide range of intermetallic compounds and alloys.

Hazard
The finely divided powder is pyrophoric. It can explode in contact with air or oxidizing gases. It is likely to explode when mixed and stirred with halogenated hydrocarbon solvents. It reacts violently with water.
All barium salts, especially the water and acid-soluble compounds, are highly toxic. Barium ion can cause death through ventricular fibrillation of the heart. It is a stimulant to the heart muscle. Intake of a few grams of barium salt can be lethal to humans. The insoluble salts such as barium sulfate, however, have little toxic action.

Analysis
The metal may be analyzed in the solid matrices by x-ray fluorescence or xray diffraction, and neutron activation techniques. Trace quantities in solution may be measured by flame or furnace atomic absorption spectrophotometry or by ICP emission technique. Measurements at further lower concentrations may be made by an ICP, coupled with a mass spectrometer (ICP/MS). Also, barium ion in solution may be measured by various wet methods, including gravimetry and volumetric analysis. In gravimetry, the metal is precipitated in slightly acidic solution as insoluble sulfate or chromate. Complexometric titration using the complexing agent, diethylenetriaminepentaacetic acid, and Eriochrome Black T as indicator, measures calcium and strontium along with barium and, therefore, is not suitable to analyze barium in a mixed solution.

Occurrence, Preparation and Uses
Actinium-227 occurs in uranium ore and is a decay product of uranium-235. It is found in equilibrium with its decay products. It is prepared by bombarding radium atoms with neutrons. Chemically, the metal is produced by reducing actinium fluoride with lithium vapor at 1,100°C to 1,300°C.

The element was discovered independently by A. Debierne and F. Giesel in 1899 and 1902, respectively. It is used in nuclear reactors as a source of neutrons.

Physical Properties
Silvery metal; cubic crystal; melts at 1,051°C; vaporizes at 3,198°C; density 10.0 g/cm3

Chemical Reactions
Actinium behaves like lanthanum forming mostly the trivalent salts of the metal. It is strongly electropositive, the first ionization potential being 5.17eV. Reacts with HCl forming AcCl3; also reacts with organic acids forming corresponding salts; combustion in air can produce oxide and nitride; susceptible to react with CO2 forming carbonate.

Analysis
The radioactivity can be measured by a beta counter. The metal at trace
concentrations can be determined by an atomic absorption or emission spectrophotometer.

Toxicity
Exposure to radiation can cause cancer.

Occurrence and Uses
Aluminum is the third most abundant element in the crust of the earth, accounting for 8.13% by weight. It does not occur in free elemental form in nature, but is found in combined forms such as oxides or silicates. It occurs in many minerals including bauxite, cryolite, feldspar and granite. Aluminum alloys have innumerable application; used extensively in electrical transmission lines, coated mirrors, utensils, packages, toys and in construction of aircraft and rockets.

Physical Properties
Silvery-white malleable metal, cubic crystal; melts at 660°C; b. p. 2520°C; density 2.70 g/cm3; insoluble in water, soluble in acids and alkalies.

Thermal, Electrochemical, and Thermochemical Properties
Specific heat 0.215 cal/g.°C (0.900 J/g.°C); heat capacity 5.81 cal/mol.°C (24.3 J/mol.°C); ΔHfus (2.54 kcal/mol (10.6 kJ/mol); ΔHvap 67.9 kcal/mol (284 kJ/mol); E° in aqueous soln. (acidic) at 25°C for the reaction
Al3+ + 3e– —› Al(s) , –1.66V; S°298 6.77 cal/degree mol. K (28.3 J/degree mol.K)

Production
Most aluminum is produced from its ore, bauxite, which contains between 40 to 60% alumina either as the trihydrate, gibbsite, or as the monohydrate, boehmite, and diaspore. Bauxite is refined first for the removal of silica and other impurities. It is done by the Bayer process. Ground bauxite is digested with NaOH solution under pressure, which dissolves alumina and silica, forming sodium aluminate and sodium aluminum silicate. Insoluble residues containing most impurities are filtered out. The clear liquor is then allowed to settle and starch is added to precipitate. The residue, so-called “red-mud”, is filtered out. After this “desilication,” the clear liquor is diluted and cooled. It is then seeded with alumina trihydrate (from a previous run) which promotes hydrolysis of the sodium aluminate to produce trihydrate crystals. The crystals are filtered out, washed, and calcined above 1,100°C to produce anhydrous alumina. The Bayer process, however, is not suitable for extracting bauxite that has high silica content (>10%). In the Alcoa process, which is suitable for highly silicious bauxite, the “red mud” is mixed with limestone and soda ash and calcined at 1,300°C. This produces “lime-soda sinter” which is cooled and treated with water. This leaches out water-soluble sodium alumnate, leaving behind calcium silicate and other impurites.
Alumina may be obtained from other minerals, such as nepheline, sodium potassium aluminum silicate, by similar soda lime sintering process.
Metal aluminum is obtained from the pure alumina at 950 to 1000°C electrolysis (Hall-Heroult process). Although the basic process has not changed since its discovery, there have been many modifications. Aluminum is also produced by electrolysis of anhydrous AlCl3.
Also, the metal can be obtained by nonelectrolytic reduction processes. In carbothermic process, alumina is heated with carbon in a furnace at 2000 to 2500°C. Similarly, in “Subhalide” process, an Al alloy, Al-Fe-Si-, (obtained by carbothermic reduction of bauxite) is heated at 1250°C with AlCl vapor. This forms the subchloride (AlCl), the vapor of which decomposes when cooled to 800°C.

Chemical Reactions
Reacts in moist air forming a coating of Al2O3; reacts with dilute mineral acids liberating H2,

2Al + 3H2SO4 ——›Al2(SO4)3 + 3H2↑

also reacts with steam to form H2; reduces a number of metals that are less active (in activity series), these include Fe, Mn, Cr, Zn, Co, Ni, Cu, Sn, Pb, etc.,

Al(s) + 3Ag+(aq) ——›Al3+(aq) + 3Ag(s)

Reactions, e.g., with alkyl halides in ether using Ziegler-Natta catalyst form alkyl aluminum halides, R3Al2X3, [R2AlX]2 and [RAlX]2; with bromine vapor forms anhydrous aluminum bromide,

2Al + 3Br2 ——› Al2Br6

Combines with iodine vapor forming aluminum iodide, AlI3; heating with HCl gas produces AlCl3,

Heating with Cl2 at 100°C also yields AlCl3,

When the metal is heated with AlCl3 at 1000°C it forms monovalent aluminum chloride, AlCl.
Produces aluminum carbide when the powder metal is heated with carbon at 2000°C or at 1000°C in presence of cryolite,

Heating the metal powder over 1000°C with sulfur, phosphorus, or selenium forms aluminum sulfide Al2S3, aluminum phosphide, AlP and aluminum selenide, Al2Se3, respectively,

Heating over 1100°C with N2 produces nitride, AlN; alkoxides are formed when the metal powder is treated with anhydrous alcohol, catalyzed by HgCl2

Reaction with CO at 1000°C produces the oxide Al2O3 and the carbide Al4C3.

Chemical Analysis
The metal may be analyzed by atomic absorption or emission spectrophotometry (at trace levels). Other techniques include X-ray diffraction, neutron activation analysis, and various colorimetric methods. Aluminum digested with nitric acid reacts with pyrocatechol violet or Eriochrome cyanide R dye to form a colored complex, the absorbance of which may be measured by a spectrophotometer at 535 nm.

Hazard
Finely divided aluminum dust is moderately flammable and explodes by heat or contact with strong oxidizing chemicals. Chronic inhalation of the powder can cause aluminosis, a type of pulmonary fibrosis. It is almost nontoxic by ingestion.

Uses
The anhydrous form is used as a catalyst for the Friedel-Crafts alkylation reaction. Its catalytic activity is similar to anhydrous AlCl3. Commercial applications, however, are few.

Physical Properties
Colorless crystalline solid in anhydrous form; melts at 97.5°C; boils at 256°C; density 3.01 g/cm3 at 25°C; moisture sensitive, fumes in air; soluble in water (reacts violently in cold water, and decomposes in hot water, alcohols, acetone, hexane, benzene, nitrobenzene, carbon disulfide and many other organic solvents).

Preparation
Prepared from bromine and metallic aluminum.

2Al + 3Br2 ——› Al2Br6 (anhydrous)

Thermochemical Properties

AlBr3 (cry)             ΔHƒ°            –126.0 kcal/mo
                             Cp               24.3 cal/degree
AlBr3 (gas)            ΔHƒ°            –101.6 kcal/mo
AlBr3 (aq)              ΔHƒ°            –214.0 kcal/mo
Al2Br6 (gas)          ΔHƒ°            –232.0 kcal/mo
AlBr3 (aq)              S°                –17.8 cal/degre
Al2Br6 (gas)         Hfusion       10.1 cal/g

Chemical Reactions
Decomposes upon heating in air to bromine and metallic aluminum.

Reacts with carbon tetrachloride at 100°C to form carbon tetrabromide;

4AlBr3 + 3CCl4 ——› 4AlCl3 + 3Br4

Reaction with phosgene yields carbonyl bromide and aluminum chlorobromide;

AlBr3 + COCl2 ——› COBr2 + AlCl2Br

Reacts violently with water; absorbs moisture forming hexahydrate,
AlBr3⋅6H2O [7784-27-2]

Chemical Analysis
Elemental composition, Al 10.11% and Br 89.89%; Al analyzed by AA spectrophotometry or colorimetric methods; Br– analyzed by iodometric titration or ion chromatography and then calculated stoichiometrically; solid may be dissolved in an organic solvent and determined by GC/MS, identified by mass ions
(AlBr3 )n where n is 2, 4 and 6.

Toxicity
Skin contact can cause tissue burn. It is moderately toxic by all routes of exposure. LD50 oral (rat and mouse): ~1600 mg/kg.

Uses
Aluminum chloride has extensive commercial applications. It is used primarily in the electrolytic production of aluminum. Another major use involves its catalytic applications in many organic reactions, including Friedel-Crafts alkylation, polymerization, isomerization, hydrocracking, oxidation, decarboxylation, and dehydrogenation. It is also used in the production of rare earth chlorides, electroplating of aluminum and in many metal finishing and metallurgical operations.

Physical Properties
White or light-yellow crystalline solid (or amorphous solid depending on the method of production); odor of HCl; hygroscopic; melts at 190°C at 2.5 atm; sublimes at 181.2°C; density 2.44 g/cm3 at 25°C; decomposes in water evolving heat; soluble in HCl; soluble in many organic solvents, including absolute ethanol, chloroform, carbon tetrachloride and ether; slightly soluble in benzene.

Thermochemical Data

ΔH°ƒ(s)      –168.3 kcal/mol
ΔG°ƒ(s)      –150.3 kcal/mol
S°              26.45 cal/deg mol
Hsoln.      –77.7 kcal/mol
Hfus         8.45 kcal/mol

Preparation
Aluminum chloride is made by chlorination of molten aluminum at temperatures between 650 to 750°C;

or by chlorination of alumina (bauxite or clay) at 800°C in the presence of a reducing agent, such as carbon or CO. It can be prepared by similar high temperature chlorination of bauxite in the presence of a chlorinated organic reductant such as CCl4.
A pelletized mixture of clay, lignite and a small amount of NaCl is chlorinated at 900°C, producing gaseous AlCl3 (Toth process). Alternatively, alumina is mixed with about 20% by weight carbon and a small amount of sodium salt. The mixture is chlorinated at 600°C (Bayer process). In the laboratory, anhydrous AlCl3 can be prepared by heating the metal with dry HCl gas at 150°C. The product sublimes and deposits in the cool air condenser. Unreacted HCl is vented out.

Reactions
Reacts with calcium and magnesium hydrides in tetrahydrofuran forming tetrahydro aluminates, Ca(AlH4)2; reacts with hydrides of alkali metals in ether forming aluminum hydride;

Hydrolyzes in chilled, dilute HCl forming aluminum chloride hexahydrate, AlCl3⋅6H2O; reacts violently with water, evolving HCl,

Hazard
Violent exothermic reactions can occur when mixed with water or alkene. Corrosive to skin.

Uses
The hexahydrate is used in the preparation of deodorant and antiperspirant. Also, it is applied in textile finishing to improve the antistatic characteristics and flammability ratings of various textile materials. Commercially, it is sold as crystalline powder or as a 28% solution in water.

Physical Properties
White or yellowish deliquescent powder; faint odor of HCl; density 2.40 g/cm3; soluble in water and polar organic solvents such as alcohol; aqueous solution acidic.

Preparation
Aluminum chloride hexahydrate is prepared by dissolving Al(OH)3 in conc. HCl and passing gaseous HCl through the solution at 0°C. The precipitate is washed with diethyl ether and dried. Alternatively, it is prepared by hydrolyzing anhydrous AlCl3 in cold dilute HCl.

Reactions
Decomposes to alumina when heated at 300°C;

Reacts with caustic soda solution forming gelatinous precipitate of aluminum hydroxide (hydrous aluminum oxide); yields aluminum monobasic stearate, Al(OH)2[OOC(CH2)16CH3] when its solution is mixed with a solution of sodium stearate.

Uses
It is used as a reducing agent, and also as a catalyst for polymerization reaction.

Physical and Thermochemical Properties
Colorless cubic crystal; very unstable; decomposes in water; ΔΗ°ƒ −11.0 kcal/mol (-46.0kJ/mol)

Preparation
Aluminum hydride is prepared by the reaction of lithium hydride with aluminum chloride in diethyl ether

Chemical Reactions
Aluminum hydride decomposes in air and water. Violent reactions occur with both. It forms a complex, aluminum diethyl etherate with diethyl ether. The product decomposes in water releasing heat.

AlH3 + (C2H5)2O ——›H3Al•O(C2H5)2

Similar complexes are likely to form with other lower aliphatic ethers. It also forms a 1:1 complex with trimethyl amine, H3Al•N(CH3)3 which reacts explosively with water (Ruff 1967).
Aluminum hydride reduces CO2 to methane under heating:

Reaction with lithium hydride in ether produces lithium aluminum hydride,

Safety
Many reactions of aluminum hydride or its complexes may proceed with explosive violence, especially with water or moist air.

Uses
The nonahydrate and other hydrated aluminum nitrates have many applications. These salts are used to produce alumina for preparation of insulating papers, in cathode tube heating elements, and on transformer core laminates. The hydrated salts are also used for extraction of actinide elements.

Physical Properties
White or colorless crystalline solid (nonahydrate – rhombic crystal); deliquescent; refractive index 1.54; melts at 73.5°C; decomposes at 150°C; highly soluble in cold water (63.7% at 25°C), decomposes in hot water, soluble in polar organic solvents.

Preparation
The nonahydrate is prepared by treating aluminum, aluminum hydroxide, aluminum oxide, or aluminous mineral with nitric acid. The nitrate is crystallized from the solution.

Reactions
Since Al(NO3)3 or its salt hydrates dissociates to Al3+ and NO3– ions in the aqueous solution, its reactions in solutions are those of Al3+ . It is partially hydrolyzed, producing H3O+ and thus accounting for the acidity of its solution in water. The products constitute a complex mixture of mono- and polynuclear hydroxo species.
Aluminum nitrate is soluble in bases, forming aluminates, [Al(OH)4(H2O)2]–. It decomposes to Al2O3 when heated at elevated temperatures.

Chemical Analysis
Elemental composition: Al 12.67%, N 19.73%, O 67.60%. Al may be analyzed by various instrumental techniques, including atomic absorption or emission spectroscopy, or colorimetry (see under Aluminum). The nitrate anion in aqueous phase may be measured by the NO3– ion selective electrode, ion chromatography, or reduction with cadmium or hydrazine, followed by colorimetric tests.