Barium hydroxide

Last updated
Barium hydroxide
Ba(OH)2monohydrate.tif
Ba(OH)2 octahydrate.JPG
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.037.470 OOjs UI icon edit-ltr-progressive.svg
EC Number
  • 241-234-5
846955
PubChem CID
RTECS number
  • CQ9200000
UNII
  • InChI=1S/Ba.2H2O/h;2*1H2/q+2;;/p-2 Yes check.svgY
    Key: RQPZNWPYLFFXCP-UHFFFAOYSA-L Yes check.svgY
  • InChI=1/Ba.2H2O/h;2*1H2/q+2;;/p-2
    Key: RQPZNWPYLFFXCP-NUQVWONBAR
  • [Ba+2].[OH-].[OH-]
Properties
Ba(OH)2
Molar mass 171.34 g/mol (anhydrous)
189.355 g/mol (monohydrate)
315.46 g/mol (octahydrate)
Appearancewhite solid
Density 3.743 g/cm3 (monohydrate)
2.18 g/cm3 (octahydrate, 16 °C)
Melting point 78 °C (172 °F; 351 K) (octahydrate)
300 °C (monohydrate)
407 °C (anhydrous)
Boiling point 780 °C (1,440 °F; 1,050 K)
mass of BaO (not Ba(OH)2):
1.67 g/100 mL (0 °C)
3.89 g/100 mL (20 °C)
4.68 g/100 mL (25 °C)
5.59 g/100 mL (30 °C)
8.22 g/100 mL (40 °C)
11.7 g/100 mL (50 °C)
20.94 g/100 mL (60 °C)
101.4 g/100 mL (100 °C)[ citation needed ]
Solubility in other solventslow
Basicity (pKb)0.15 (first OH), 0.64 (second OH) [1]
−53.2·10−6 cm3/mol
1.50 (octahydrate)
Structure
octahedral
Thermochemistry [2]
944.7 kJ·mol−1
Enthalpy of fusion fHfus)
16 kJ·mol−1
Hazards
GHS labelling:
GHS-pictogram-acid.svg GHS-pictogram-exclam.svg
Danger
H302, H314, H332, H412
NFPA 704 (fire diamond)
3
0
0
Flash point Non-flammable
Related compounds
Other anions
Barium oxide
Barium peroxide
Other cations
Calcium hydroxide
Strontium hydroxide
Supplementary data page
Barium hydroxide (data page)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
X mark.svgN  verify  (what is  Yes check.svgYX mark.svgN ?)

Barium hydroxide is a chemical compound with the chemical formula Ba(OH)2. The monohydrate (x = 1), known as baryta or baryta-water, is one of the principal compounds of barium. This white granular monohydrate is the usual commercial form.

Contents

Preparation and structure

Barium hydroxide can be prepared by dissolving barium oxide (BaO) in water:

BaO + H2O → Ba(OH)2

It crystallises as the octahydrate, which converts to the monohydrate upon heating in air. At 100 °C in a vacuum, the monohydrate will yield BaO and water. [3] The monohydrate adopts a layered structure (see picture above). The Ba2+ centers adopt a square anti-prismatic geometry. Each Ba2+ center is bound by two water ligands and six hydroxide ligands, which are respectively doubly and triply bridging to neighboring Ba2+ centre sites. [4] In the octahydrate, the individual Ba2+ centers are again eight coordinate but do not share ligands. [5]

Coordination sphere about an individual barium ion in Ba(OH)2.H2O. Ba(OH)2O2HCoordSph.tif
Coordination sphere about an individual barium ion in Ba(OH)2.H2O.

Uses

Industrially, barium hydroxide is used as the precursor to other barium compounds. The monohydrate is used to dehydrate and remove sulfate from various products. [6] This application exploits the very low solubility of barium sulfate. This industrial application is also applied to laboratory uses.

Laboratory uses

Barium hydroxide is used in analytical chemistry for the titration of weak acids, particularly organic acids. Its clear aqueous solution is guaranteed to be free of carbonate, unlike those of sodium hydroxide and potassium hydroxide, as barium carbonate is insoluble in water. This allows the use of indicators such as phenolphthalein or thymolphthalein (with alkaline colour changes) without the risk of titration errors due to the presence of carbonate ions, which are much less basic. [7]

Barium hydroxide is occasionally used in organic synthesis as a strong base, for example for the hydrolysis of esters [8] and nitriles, [9] [10] [11] and as a base in aldol condensations.

Barium hydroxide-catalyzed 2-carboxy-1,3-dihydroxynaphthalene preparation.svg
Barium hydroxide-catalyzed methylsuccinic acid preparation.svg

There are several uses for barium hydroxide such as to hydrolyse one of the two equivalent ester groups in dimethyl hendecanedioate. [12]

Barium hydroxide has also been used, as well, in the decarboxylation of amino acids liberating barium carbonate in the process. [13]

It is also used in the preparation of cyclopentanone, [14] diacetone alcohol [15] and D-gulonic γ-lactone. [16]

Cyclopentanone prepn.png
Barium hydroxide-catalyzed diacetone alcohol preparation.svg

Reactions

Barium hydroxide decomposes to barium oxide when heated to 800 °C. Reaction with carbon dioxide gives barium carbonate. Its aqueous solution, being highly alkaline, undergoes neutralization reactions with acids due to it being a strong base. It is especially useful on reactions that require the titrations of weak organic acids. Thus, it forms barium sulfate and barium phosphate with sulfuric and phosphoric acids, respectively. Reaction with hydrogen sulfide produces barium sulfide. Precipitation of many insoluble, or less soluble barium salts, may result from double replacement reaction when a barium hydroxide aqueous solution is mixed with many solutions of other metal salts. [17]

Reactions of barium hydroxide with ammonium salts are strongly endothermic. The reaction of barium hydroxide octahydrate with ammonium chloride [18] [19] or [20] ammonium thiocyanate [20] [21] is often used as a classroom chemistry demonstration, producing temperatures cold enough to freeze water and enough water to dissolve the resulting mixture.

Safety

Barium hydroxide presents the same hazards such as skin irritation and burns as well as eye damage, just as the other strong bases and as other water-soluble barium compounds: it is corrosive and toxic. [ citation needed ]

See also

Related Research Articles

<span class="mw-page-title-main">Barium</span> Chemical element, symbol Ba and atomic number 56

Barium is a chemical element with the symbol Ba and atomic number 56. It is the fifth element in group 2 and is a soft, silvery alkaline earth metal. Because of its high chemical reactivity, barium is never found in nature as a free element.

<span class="mw-page-title-main">Hydroxide</span> Chemical compound

Hydroxide is a diatomic anion with chemical formula OH. It consists of an oxygen and hydrogen atom held together by a single covalent bond, and carries a negative electric charge. It is an important but usually minor constituent of water. It functions as a base, a ligand, a nucleophile, and a catalyst. The hydroxide ion forms salts, some of which dissociate in aqueous solution, liberating solvated hydroxide ions. Sodium hydroxide is a multi-million-ton per annum commodity chemical. The corresponding electrically neutral compound HO is the hydroxyl radical. The corresponding covalently bound group –OH of atoms is the hydroxy group. Both the hydroxide ion and hydroxy group are nucleophiles and can act as catalysts in organic chemistry.

<span class="mw-page-title-main">Alkaline earth metal</span> Group of chemical elements

The alkaline earth metals are six chemical elements in group 2 of the periodic table. They are beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). The elements have very similar properties: they are all shiny, silvery-white, somewhat reactive metals at standard temperature and pressure.

<span class="mw-page-title-main">Sodium hydroxide</span> Chemical compound with formula NaOH

Sodium hydroxide, also known as lye and caustic soda, is an inorganic compound with the formula NaOH. It is a white solid ionic compound consisting of sodium cations Na+ and hydroxide anions OH.

<span class="mw-page-title-main">Base (chemistry)</span> Type of chemical substance

In chemistry, there are three definitions in common use of the word base, known as Arrhenius bases, Brønsted bases, and Lewis bases. All definitions agree that bases are substances which react with acids as originally proposed by G.-F. Rouelle in the mid-18th century.

<span class="mw-page-title-main">Copper(II) sulfate</span> Chemical compound

Copper(II) sulfate, also known as copper sulphate, is an inorganic compound with the chemical formula CuSO4. It forms hydrates CuSO4·nH2O, where n can range from 1 to 7. The pentahydrate, a bright blue crystal, is the most commonly encountered hydrate of copper(II) sulfate. Older names for the pentahydrate include blue vitriol, bluestone, vitriol of copper, and Roman vitriol. It exothermically dissolves in water to give the aquo complex [Cu(H2O)6]2+, which has octahedral molecular geometry. The structure of the solid pentahydrate reveals a polymeric structure wherein copper is again octahedral but bound to four water ligands. The Cu(II)(H2O)4 centers are interconnected by sulfate anions to form chains. Anhydrous copper sulfate is a light grey powder.

<span class="mw-page-title-main">Neutralization (chemistry)</span> Chemical reaction in which an acid and a base react quantitatively

In chemistry, neutralization or neutralisation is a chemical reaction in which acid and a base react quantitatively with each other. In a reaction in water, neutralization results in there being no excess of hydrogen or hydroxide ions present in the solution. The pH of the neutralized solution depends on the acid strength of the reactants.

The Brønsted–Lowry theory (also called proton theory of acids and bases) is an acid–base reaction theory which was proposed independently by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923. The fundamental concept of this theory is that when an acid and a base react with each other, the acid forms its conjugate base, and the base forms its conjugate acid by exchange of a proton (the hydrogen cation, or H+). This theory is a generalization of the Arrhenius theory.

Classical qualitative inorganic analysis is a method of analytical chemistry which seeks to find the elemental composition of inorganic compounds. It is mainly focused on detecting ions in an aqueous solution, therefore materials in other forms may need to be brought to this state before using standard methods. The solution is then treated with various reagents to test for reactions characteristic of certain ions, which may cause color change, precipitation and other visible changes.

<span class="mw-page-title-main">Hexafluorosilicic acid</span> Octahedric silicon compound

Hexafluorosilicic acid is an inorganic compound with the chemical formula H
2
SiF
6
. Aqueous solutions of hexafluorosilicic acid consist of salts of the cation and hexafluorosilicate anion. These salts and their aqueous solutions are colorless.

The Kjeldahl method or Kjeldahl digestion (Danish pronunciation: [ˈkʰelˌtɛˀl]) in analytical chemistry is a method for the quantitative determination of nitrogen contained in organic substances plus the nitrogen contained in the inorganic compounds ammonia and ammonium (NH3/NH4+). Without modification, other forms of inorganic nitrogen, for instance nitrate, are not included in this measurement. Using an empirical relation between Kjeldahl nitrogen content and protein content it is an important method for analyzing proteins. This method was developed by Johan Kjeldahl in 1883.

<span class="mw-page-title-main">Thermometric titration</span>

A thermometric titration is one of a number of instrumental titration techniques where endpoints can be located accurately and precisely without a subjective interpretation on the part of the analyst as to their location. Enthalpy change is arguably the most fundamental and universal property of chemical reactions, so the observation of temperature change is a natural choice in monitoring their progress. It is not a new technique, with possibly the first recognizable thermometric titration method reported early in the 20th century. In spite of its attractive features, and in spite of the considerable research that has been conducted in the field and a large body of applications that have been developed; it has been until now an under-utilized technique in the critical area of industrial process and quality control. Automated potentiometric titration systems have pre-dominated in this area since the 1970s. With the advent of cheap computers able to handle the powerful thermometric titration software, development has now reached the stage where easy to use automated thermometric titration systems can in many cases offer a superior alternative to potentiometric titrimetry.

<span class="mw-page-title-main">Barium chlorate</span> Chemical compound

Barium chlorate, Ba(ClO3)2, is the barium salt of chloric acid. It is a white crystalline solid, and like all soluble barium compounds, irritant and toxic. It is sometimes used in pyrotechnics to produce a green color. It also finds use in the production of chloric acid.

<span class="mw-page-title-main">Barium bromide</span> Chemical compound

Barium bromide is the chemical compound with the formula BaBr2. Like barium chloride, it dissolves well in water and is toxic.

<span class="mw-page-title-main">Barium ferrate</span> Chemical compound

Barium ferrate is the chemical compound of formula BaFeO4. This is a rare compound containing iron in the +6 oxidation state. The ferrate(VI) ion has two unpaired electrons, making it paramagnetic. It is isostructural with BaSO4, and contains the tetrahedral [FeO4]2− anion.

Barium perchlorate is a powerful oxidizing agent, with the formula Ba(ClO4)2. It is used in the pyrotechnic industry.

<span class="mw-page-title-main">Tetraethylammonium chloride</span> Chemical compound

Tetraethylammonium chloride (TEAC) is a quaternary ammonium compound with the chemical formula (C2H5)4N+Cl, sometimes written as Et4N+Cl. In appearance, it is a hygroscopic, colorless, crystalline solid. It has been used as the source of tetraethylammonium ions in pharmacological and physiological studies, but is also used in organic chemical synthesis.

Barium permanganate is a chemical compound, with the formula Ba(MnO4)2. It forms violet to brown crystals that are sparingly soluble in water.

References

  1. "Sortierte Liste: pKb-Werte, nach Ordnungszahl sortiert. - Das Periodensystem online" (in German).
  2. Lide, David R., ed. (2009). CRC Handbook of Chemistry and Physics (90th ed.). Boca Raton, Florida: CRC Press. ISBN   978-1-4200-9084-0.
  3. (1960). Gmelins Handbuch der anorganischen Chemie (8. Aufl.), Weinheim: Verlag Chemie, p. 289.
  4. Kuske, P.; Engelen, B.; Henning, J.; Lutz, H.D.; Fuess, H.; Gregson, D. "Neutron diffraction study of Sr(OH)2(H2O) and beta-Ba(OH)2*(H2O)" Zeitschrift für Kristallographie (1979-2010) 1988, vol. 183, p319-p325.
  5. Manohar, H.; Ramaseshan, S. "The crystal structure of barium hydroxide octahydrate Ba (OH)2(H2O)8" Zeitschrift für Kristallographie, Kristallgeometrie, Kristallphysik, Kristallchemie 1964. vol. 119, p357-p374
  6. Robert Kresse, Ulrich Baudis, Paul Jäger, H. Hermann Riechers, Heinz Wagner, Jochen Winkler, Hans Uwe Wolf, "Barium and Barium Compounds" in Ullmann's Encyclopedia of Industrial Chemistry, 2007 Wiley-VCH, Weinheim. doi : 10.1002/14356007.a03_325.pub2
  7. Mendham, J.; Denney, R. C.; Barnes, J. D.; Thomas, M. J. K. (2000), Vogel's Quantitative Chemical Analysis (6th ed.), New York: Prentice Hall, ISBN   0-582-22628-7
  8. Meyer, K.; Bloch, H. S. (1945). "Naphthoresorcinol". Org. Synth. 25: 73; Coll. Vol.3: 637.
  9. Brown, G. B. (1946). "Methylsuccinic acid". Org. Synth. 26: 54; Coll. Vol.3: 615.
  10. Ford, Jared H. (1947). "β-Alanine". Org. Synth. 27: 1; Coll. Vol.3: 34.
  11. Anslow, W. K.; King, H.; Orten, J. M.; Hill, R. M. (1925). "Glycine". Org. Synth. 4: 31; Coll. Vol.1: 298.
  12. Durham, L. J.; McLeod, D. J.; Cason, J. (1958). "Methyl hydrogen hendecanedioate". Org. Synth. 38:55; Coll. Vol.4:635.
  13. Chaudhari, M. R.; Kulkarni, Y. A.; Gokhale, S. B. (6 October 2008). Biochemistry and Clinical Pathology. ISBN   9788185790169.
  14. Thorpe, J. F.; Kon, G. A. R. (1925). "Cyclopentanone". Org. Synth. 5: 37; Coll. Vol.1: 192.
  15. Conant, J. B.; Tuttle, Niel. (1921). "Diacetone alcohol". Org. Synth. 1: 45; Coll. Vol.1: 199.
  16. Karabinos, J. V. (1956). "γ-lactone". Org. Synth. 36: 38; Coll. Vol.4: 506.
  17. Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN   0-07-049439-8
  18. "Endothermic Reactions of Hydrated Barium Hydroxide and Ammonium Chloride". UC San Diego. Retrieved 2 April 2014.
  19. Endothermic Solid-Solid Reactions
  20. 1 2 Camp, Eric. "Endothermic Reaction". Univertist of Washington. Retrieved 2 April 2014.
  21. "Endothermic solid-solid reactions" (PDF). Classic Chemistry Demonstrations. The Royal Society of Chemistry. Archived from the original (PDF) on 7 April 2014. Retrieved 2 April 2014.