Lead(II) nitrate is a colourless inorganic compound with the chemical formula Pb(NO3)2. Unlike most other lead(II) salts, it is soluble in water. The use of lead nitrate as a white pigment in paint has been discontinued due to concerns regarding its toxicity.

Lead(II) nitrate
Lead(II) nitrate 1.jpg
Names
IUPAC name
Lead(II) nitrate
Other names
Lead nitrate
Plumbous nitrate
Lead dinitrate
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.030.210
RTECS number
  • OG2100000
UNII
UN number 1469
Properties
Pb(NO3)2
Molar mass 331.2 g/mol[1]
Appearance Colourless crystals
Density 4.53 g/cm3 (20 °C)[1]
Melting point 470 °C (878 °F; 743 K)[1] decomposes
See data page
−74.0·10−6 cm3/mol[2]
1.782[3]
Structure
Face-centred cubic, cP36
Pa3, No. 205[4]
a = 0.78586 nm[4]
0.4853 nm3
4
Hazards
Safety data sheet See: data page
ICSC 1000
Lethal dose or concentration (LD, LC):
500 mg/kg (guinea pig, oral)[5]
Supplementary data page
Refractive index (n),
Dielectric constantr), etc.
Thermodynamic
data
Phase behaviour
solid–liquid–gas
UV, IR, NMR, MS
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

HistoryEdit

In 1597, the German alchemist Andreas Libavius first described the compound, coining the medieval names of plumb dulcis and calx plumb dulcis, meaning "sweet lead", because of its taste.[6] The production of lead(II) nitrate from either metallic lead or lead oxide in nitric acid was small-scale, for direct use in making other lead compounds. It was produced as a raw material for the production of coloured pigments in lead paints, such as chrome yellow (lead(II) chromate), chrome orange (lead(II) hydroxide chromate) and similar lead compounds. These pigments were used for dyeing and printing calico and other textiles.[7] Although originally not understood during the following centuries, the decrepitation property of lead(II) nitrate led to its use in matches and special explosives.[8]

In the 19th century lead(II) nitrate began to be produced commercially. The production process is chemically straightforward: metallic lead is dissolved in nitric acid. The compound is obtained by crystallization from concentrated solution. The main use was as a white pigment in paint, but the use of lead paint has been superseded by the use of less toxic paints that use titanium dioxide as the white pigment.

Preparation and propertiesEdit

 
Unit cell in Pb(NO3)2 crystal. Black Pb2+, blue and red NO3-

Lead(II) nitrate can be obtained by reaction of lead(II) oxide with nitric acid[9]

PbO + 2 HNO3 → Pb(NO3)2 + H2O

When concentrated nitric acid is used, crystals of lead nitrate are obtained as the common ion effect reduces the solubility of lead nitrate in this medium.[10]

The compound crystallizes in the face-centred cubic system, space group Pa3Z=4 (Bravais lattice notation), with unit cell length 784 pm.[11] There is no evidence for free internal rotation of the nitrate groups within the crystal lattice at elevated temperatures.

In nitric acid treatment of lead-containing wastes, e.g., in the processing of lead–bismuth wastes from lead refineries, impure solutions of lead(II) nitrate are formed as by-product. These solutions are reported to be used in the gold cyanidation process.[12]

When concentrated sodium hydroxide solution is added to a nitrate solution, basic nitrates may be formed. Up through the half equivalence point, Pb(NO3)2·Pb(OH)2 predominates, then after this point Pb(NO3)2·5Pb(OH)2 is formed. Simple Pb(OH)2 is not formed up to at least pH 12.[13][14]

Lead(II) has a standard reduction potential (E0) of −0.125 V and the nitrate ion has an E0 of +0.956 V.[15] These properties show that lead(II) nitrate can behave as an oxidizing agent only towards easily oxidized substrates.

When heated, lead(II) nitrate crystals decompose to lead(II) oxide, oxygen and nitrogen dioxide.

2 Pb(NO3)2 (s) → 2 PbO (s) + 4 NO2 (g) + O2 (g)

Because of this property, lead nitrate is sometimes used in fireworks.[8]

ComplexationEdit

Lead nitrate has been used to make complexes involving lead(II) because of its relatively high solubility, compared to other lead salts, in various solvents. For example, combining lead nitrate and pentaethylene glycol (EO5) in a solution of acetonitrile and methanol followed by slow evaporation produces the complex [Pb(NO3)2(EO5)].[16] In the crystal structure for this compound, the EO5 chain is wrapped around the lead ion in an equatorial plane. The two nitrate ligands are both bidentate; one is situated above the plane and the other below. The total coordination number is 10. Reaction of the tripodal ligand 2,4,6-tris[4-(imidazol-1-ylmethyl)phenyl]-1,3,5-triazine (timpt) with lead(II) nitrate produced a polycatenated structure in which the lead atom has a stereochemically active lone pair of electrons.[17] The nitrate ion acts as a bridging ligand in this complex.

ApplicationsEdit

Because of the toxicity of lead(II) salts, the production lead paints has all but ceased. Titanium dioxide is now the preferred substance to use as a white pigment in paint.[18] Other historical applications of lead(II) nitrate, such as in matches and fireworks, have declined or ceased as well.

Current applications of lead(II) nitrate include use as a heat stabilization in nylon and polyesters, in thermographic printing paper, and in rodenticides.[9] To improve the leaching process in the gold cyanidation process, lead(II) nitrate solution is added. Although a bulk process, only limited amounts (10 to 100 milligrams lead(II) nitrate per kilogram gold) are required.[19][20] Both the cyanidation itself, as well as the use of lead compounds in the process, are deemed controversial due to the compounds' toxic nature.

On a laboratory scale, lead(II) nitrate may be used to make nitrogen dioxide. The dry compound is heated in a steel vessel, producing nitrogen dioxide gas, which dimerizes to dinitrogen tetroxide when condensed to a liquid or when it is dissolved in an organic solvent.

Pb(NO3)2 → PbO2 + 2 NO2

In organic chemistry, lead(II) nitrate has been used as an oxidant, for example as an alternative to the Sommelet reaction for oxidation of benzylic halides to aldehydes.[21] It has also found use in the preparation of isothiocyanates from dithiocarbamates.[22] Because of its toxicity it has largely fallen out of favour, but it still finds occasional use, for example as a bromide scavenger during SN1 substitution.[23]

SafetyEdit

Lead(II) nitrate is toxic. It must be handled and stored with the appropriate safety precautions to prevent inhalation, ingestion and skin contact. Due to its hazardous nature, the limited applications of lead(II) nitrate are under constant scrutiny.

Ingestion of lead(II) nitrate will lead to acute lead poisoning.[24] All inorganic lead compounds are classified by the International Agency for Research on Cancer (IARC) as probably carcinogenic to humans (Category 2A).[25] They have been linked to renal cancer and glioma in experimental animals and to renal cancer, brain cancer and lung cancer in humans, although studies of workers exposed to lead are often complicated by concurrent exposure to arsenic.[26] Lead is known to substitute for zinc in a number of enzymes, including δ-aminolevulinic acid dehydratase (porphobilinogen synthase) in the haem biosynthetic pathway and pyrimidine-5′-nucleotidase, important for the correct metabolism of DNA and can therefore cause fetal damage.[27]

Material Safety Data SheetsEdit

ReferencesEdit

  1. ^ a b c Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.70. ISBN 1439855110.
  2. ^ Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.133. ISBN 1439855110.
  3. ^ Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds. McGraw-Hill. p. 475. ISBN 0-07-049439-8.
  4. ^ a b Nowotny, H.; Heger, G. (1986). "Structure refinement of lead nitrate". Acta Crystallographica Section C. 42 (2): 133. doi:10.1107/S0108270186097032.
  5. ^ "Lead compounds (as Pb)". Immediately Dangerous to Life and Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  6. ^ Libavius, Andreas (1595). Alchemia Andreæ Libavii. Francofurti: Iohannes Saurius.
  7. ^ Partington, James Riddick (1950). A Text-book of Inorganic Chemistry. MacMillan. p. 838.
  8. ^ a b Barkley, J. B. (October 1978). "Lead nitrate as an oxidizer in blackpowder". Pyrotechnica. Post Falls, Idaho: Pyrotechnica Publications. 4: 16–18.
  9. ^ a b Greenwood, Norman N.; Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. pp. 388, 456. ISBN 0-7506-3365-4.
  10. ^ Ferris, L. M. (1959). "Lead nitrate—Nitric acid—Water system". Journal of Chemical & Engineering Data. 5 (3): 242. doi:10.1021/je60007a002.
  11. ^ Nowotny, H.; G. Heger (1986). "Structure refinement of lead nitrate". Acta Crystallogr. C. 42 (2): 133–35. doi:10.1107/S0108270186097032.
  12. ^ "Product catalog; other products". Tilly, Belgium: Sidech. Archived from the original on 2007-07-01. Retrieved 2008-01-05.
  13. ^ Othmer, D. F. (1967). Kirk-Othmer Encyclopedia of Chemical Technology. 12 (Iron to Manganese) (second completely revised ed.). New York: John Wiley & Sons. p. 272. ISBN 0-471-02040-0.
  14. ^ Pauley, J. L.; M. K. Testerman (1954). "Basic Salts of Lead Nitrate Formed in Aqueous Media". Journal of the American Chemical Society. 76 (16): 4220–4222. doi:10.1021/ja01645a062.
  15. ^ Hill, John W.; Petrucci, Ralph H. (1999). General Chemistry (2nd ed.). Upper Saddle River, New Jersey: Prentice Hall. p. 781. ISBN 0-13-010318-7.
  16. ^ Rogers, Robin D.; Andrew H. Bond; Debra M. Roden (1996). "Structural Chemistry of Poly (ethylene glycol). Complexes of Lead(II) Nitrate and Lead(II) Bromide". Inorg. Chem. 35 (24): 6964–6973. doi:10.1021/ic960587b. PMID 11666874.
  17. ^ Shuang-Yi Wan; Jian Fan; Taka-aki Okamura; Hui-Fang Zhu; Xing-Mei Ouyang; Wei-Yin Sun & Norikazu Ueyama (2002). "2D 4.82 Network with threefold parallel interpenetration from tripodal ligand and lead(II) nitrate". Chem. Commun. (21): 2520–2521. doi:10.1039/b207568g.
  18. ^ "Historical development of titanium dioxide". Millennium Inorganic Chemicals. Archived from the original on October 21, 2007. Retrieved 2008-01-04.
  19. ^ Habashi, Fathi (1998). "Recent advances in gold metallurgy". Revisa de la Facultad de Ingeniera, Universidad Central de Venezuela. 13 (2): 43–54.
  20. ^ "Auxiliary agents in gold cyanidation". Gold Prospecting and Gold Mining. Retrieved 2008-01-05.
  21. ^ Schulze, K. E. (1884). "Über α- und β-Methylnaphtalin". Chemische Berichte. 17: 1530. doi:10.1002/cber.188401701384.
  22. ^ Dains, F. B.; Brewster, R. Q.; Olander, C. P. "Phenyl isothiocyanate". Organic Syntheses.; Collective Volume, 1, p. 447
  23. ^ Rapoport, H.; Jamison, T. (1998). "(S)-N-(9-Phenylfluoren-9-yl)alanine and (S)-Dimethyl-N-(9-phenylfluoren-9-yl)aspartate". Organic Syntheses.; Collective Volume, 9, p. 344
  24. ^ "Lead nitrate, Chemical Safety Card 1000". International Labour Organization, International Occupational Safety and Health Information Centre. March 1999. Retrieved 2008-01-19.
  25. ^ "Inorganic and Organic Lead Compounds" (PDF). IARC Monographs on the Evaluation of Carcinogenic Risks to Humans. International Agency for Research on Cancer. Suppl. 7: 239. 1987. Archived from the original (PDF) on 2008-03-06. Retrieved 2008-01-19.
  26. ^ World Health Organization, International Agency for Research on Cancer (2006). "Inorganic and Organic Lead Compounds" (PDF). IARC Monographs on the Evaluation of Carcinogenic Risks to Humans. International Agency for Research on Cancer. 87. ISBN 92-832-1287-8. Archived from the original (PDF) on 2007-10-21. Retrieved 2008-01-01.
  27. ^ Mohammed-Brahim, B.; Buchet, J.P.; Lauwerys, R. (1985). "Erythrocyte pyrimidine 5'-nucleotidase activity in workers exposed to lead, mercury or cadmium". Int Arch Occup Environ Health. 55 (3): 247–52. doi:10.1007/BF00383757. PMID 2987134.