Compounds of lead

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Compounds of lead exist with lead in two main oxidation states: +2 and +4. The former is more common. Inorganic lead(IV) compounds are typically strong oxidants or exist only in highly acidic solutions. [1]

Contents

Chemistry

Various oxidized forms of lead are easily reduced to the metal. An example is heating PbO with mild organic reducing agents such as glucose. The mixture of the oxide and the sulfide heated together will also form the metal. [2]

2 PbO + PbS → 3 Pb + SO2

Metallic lead is attacked (oxidized) only superficially by air, forming a thin layer of lead oxide that protects it from further oxidation. The metal is not attacked by sulfuric or hydrochloric acids. It dissolves in nitric acid with the evolution of nitric oxide gas to form dissolved Pb(NO3)2.

3 Pb + 8 H+ + 8 NO−
3 → 3 Pb2+ + 6 NO−
3 + 2 NO + 4 H2O

When heated with nitrates of alkali metals, metallic lead oxidizes to form PbO (also known as litharge), leaving the corresponding alkali nitrite. PbO is representative of lead's +2 oxidation state. It is soluble in nitric and acetic acids, from which solutions it is possible to precipitate halide, sulfate, chromate, carbonate (PbCO3), and basic carbonate (Pb
3(OH)
2(CO
3)
2) salts of lead. The sulfide can also be precipitated from acetate solutions. These salts are all poorly soluble in water. Among the halides, the iodide is less soluble than the bromide, which, in turn, is less soluble than the chloride. [3]

Lead(II) oxide is also soluble in alkali hydroxide solutions to form the corresponding plumbite salt. [2]

PbO + 2 OH− + H2O → Pb(OH)2−
4

Chlorination of plumbite solutions causes the formation of lead's +4 oxidation state.

Pb(OH)2−
4 + Cl2 → PbO2 + 2 Cl− + 2 H2O

Lead dioxide is representative of the +4 oxidation state, and is a powerful oxidizing agent. The chloride of this oxidation state is formed only with difficulty and decomposes readily into lead(II) chloride and chlorine gas. The bromide and iodide of lead(IV) are not known to exist. [3] Lead dioxide dissolves in alkali hydroxide solutions to form the corresponding plumbates. [2]

PbO2 + 2 OH− + 2 H2O → Pb(OH)2−
6

Lead also has an oxide with mixed +2 and +4 oxidation states, red lead (Pb
3O
4), also known as minium.

Lead readily forms an equimolar alloy with sodium metal that reacts with alkyl halides to form organometallic compounds of lead such as tetraethyllead. [4]

Oxides and sulfide

There are three oxides known: PbO, Pb3O4 (sometimes called "minium"), and PbO2. The former has two allotropes: α-PbO and β-PbO, both with layer structure and tetracoordinated lead. The alpha allotrope is red-colored and has the Pb–O distance of 230 pm; the beta allotrope is yellow-colored and has the Pb–O distance of 221 and 249 pm (due to asymmetry). [5] Thanks to the similarity, both allotropes can exist under standard conditions (beta with small (10−5 relative) impurities, such as Si, Ge, Mo, etc.). PbO reacts with acids to form salts, and with alkalies to give plumbites, [Pb(OH)3]− or [Pb(OH)4]2−. [6]

The dioxide may be prepared by, for example, halogenization of lead(II) salts. The alpha allotrope is rhombohedral, and the beta allotrope is tetragonal. [6] Both allotropes are black-brown in color and always contain some water, which cannot be removed, as heating also causes decomposition (to PbO and Pb3O4). The dioxide is a powerful oxidizer: it can oxidize hydrochloric and sulfuric acids. It does not reacts with alkaline solution, but reacts with solid alkalies to give hydroxyplumbates, or with basic oxides to give plumbates. [6]

Reaction of lead with sulfur or hydrogen sulfide yields lead sulfide. The solid has the NaCl-like structure (simple cubic), which it keeps up to the melting point, 1114 °C (2037 °F). If the heating occurs in presence of air, the compounds decomposes to give the monoxide and the sulfate. [7] The compounds are almost insoluble in water, weak acids, and (NH4)2S/(NH4)2S2 solution is the key for separation of lead from analythical groups I to III elements, tin, arsenic, and antimony. The compounds dissolve in nitric and hydrochloric acids, to give elemental sulfur and hydrogen sulfide, respectively. [7] Heating mixtures of the monoxide and the sulfide forms the metal. [2]

2 PbO + PbS → 3 Pb + SO2↑

Halides and other salts

Heating lead carbonate with hydrogen fluoride yields the hydrofluoride, which decomposes to the difluoride when it melts. This white crystalline powder is more soluble than the diiodide, but less than the dibromide and the dichloride. No coordinated lead fluorides exist (except the unstable PbF+ cation). [8] The tetrafluoride, a yellow crystalline powder, is unstable.

Other dihalides are received upon heating lead(II) salts with the halides of other metals; lead dihalides precipitate to give white orthorhombic crystals (diiodide form yellow hexagonal crystals). They can also be obtained by direct elements reaction at temperature exceeding melting points of dihalides. Their solubility increases with temperature; adding more halides first decreases the solubility, but then increases due to complexation, with the maximum coordination number being 6. The complexation depends on halide ion numbers, atomic number of the alkali metal, the halide of which is added, temperature and solution ionic strength. [9] The tetrachloride is obtained upon dissolving the dioxide in hydrochloric acid; to prevent the exothermic decomposition, it is kept under concentrated sulfuric acid. The tetrabromide may not, and the tetraiodide definitely does not exist. [10] The diastatide has also been prepared. [11]

The metal is not attacked by sulfuric or hydrochloric acids. It dissolves in nitric acid with the evolution of nitric oxide gas to form dissolved Pb(NO3)2. [8] It is a well-soluble solid in water; it is thus a key to receive the precipitates of halide, sulfate, chromate, carbonate, and basic carbonate Pb3(OH)2(CO3)2 salts of lead. [3]

Chloride complexes

Diagram showing the forms of lead in chloride media. Lead complexes in chloride media.png
Diagram showing the forms of lead in chloride media.

Lead(II) forms a series of complexes with chloride, the formation of which alters the corrosion chemistry of the lead. This will tend to limit the solubility of lead in saline media.

Equilibrium constants for aqueous lead chloride complexes at 25 °C [13]
Pb2+ + Cl− → PbCl+K1 = 12.59
PbCl+ + Cl− → PbCl2K2 = 14.45
PbCl2 + Cl− → PbCl3−K3 = 3.98 ×10−1
PbCl3− + Cl− → PbCl42−K4 = 8.92 × 10−2

Organolead

The best-known compounds are the two simplest plumbane deratives: tetramethyllead (TML) and tetraethyllead (TEL); however, the homologs of these, as well as hexaethyldilead (HEDL), are of lesser stability. The tetralkyl deratives contain lead(IV); the Pb–C bonds are covalent. They thus resemble typical organic compounds. [14]

Lead readily forms an equimolar alloy with sodium metal that reacts with alkyl halides to form organometallic compounds of lead such as tetraethyllead. [15] The Pb–C bond energies in TML and TEL are only 167 and 145 kJ/mol; the compounds thus decompose upon heating, with first signs of TEL composition seen at 100 °C (210 °F). Pyrolysis yields elemental lead and alkyl radicals; their interreaction causes the synthesis of HEDL. [14] They also decompose upon sunlight or UV-light. [16] In presence of chlorine, the alkyls begin to be replaced with chlorides; the R2PbCl2 in the presence of HCl (a by-product of the previous reaction) leads to the complete mineralization to give PbCl2. Reaction with bromine follows the same principle. [16]

Phase diagrams of solubilities

Lead(II) sulfate is poorly soluble, as can be seen in the following diagram showing addition of SO42− to a solution containing 0.1 M of Pb2+. The pH of the solution is 4.5, as above that, Pb2+ concentration can never reach 0.1 M due to the formation of Pb(OH)2. Observe that Pb2+ solubility drops 10,000 fold as SO42− reaches 0.1 M.

PbSO4 solubility graph.png Lead sulphate pourdaix diagram.png
Plot showing aqueous concentration of dissolved Pb2+ as a function of SO42− [12] Diagram for lead in sulfate media [12]

The addition of chloride can lower the solubility of lead, though in chloride-rich media (such as aqua regia) the lead can become soluble again as anionic chloro-complexes.

PbCl2 solubility graph.png Lead chloride pourdiax diagram.png
Diagram showing the solubility of lead in chloride media. The lead concentrations are plotted as a function of the total chloride present. [12] Pourbaix diagram for lead in chloride (0.1 M) media [12]

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In chemistry, a salt is a chemical compound consisting of an ionic assembly of cations and anions. Salts are composed of related numbers of cations and anions so that the product is electrically neutral. These component ions can be inorganic, such as chloride (Cl−), or organic, such as acetate ; and can be monatomic, such as fluoride (F−) or polyatomic, such as sulfate.

Aqua regia

Aqua regia is a mixture of nitric acid and hydrochloric acid, optimally in a molar ratio of 1:3. Aqua regia is a yellow-orange fuming liquid, so named by alchemists because it can dissolve the noble metals gold and platinum, though not all metals.

Sodium carbonate

Sodium carbonate, Na2CO3, (also known as washing soda, soda ash and soda crystals) is the inorganic compound with the formula Na2CO3 and its various hydrates. All forms are white, water-soluble salts that yield moderately alkaline solutions in water. Historically it was extracted from the ashes of plants growing in sodium-rich soils. Because the ashes of these sodium-rich plants were noticeably different from ashes of wood (once used to produce potash), sodium carbonate became known as "soda ash." It is produced in large quantities from sodium chloride and limestone by the Solvay process.

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Lead(II) sulfate

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A selenide is a chemical compound containing a selenium anion with oxidation number of −2 (Se2−), much as sulfur does in a sulfide. The chemistry of the selenides and sulfides is similar. Similar to sulfide, in aqueous solution, the selenide ion, Se2−, is prevalent only in very basic conditions. In neutral conditions, hydrogen selenide ion, HSe−, is most common. In acid conditions, hydrogen selenide, H2Se, is formed.

Hydrazoic acid

Hydrazoic acid, also known as hydrogen azide or azoimide, is a compound with the chemical formula HN3. It is a colorless, volatile, and explosive liquid at room temperature and pressure. It is a compound of nitrogen and hydrogen, and is therefore a pnictogen hydride. It was first isolated in 1890 by Theodor Curtius. The acid has few applications, but its conjugate base, the azide ion, is useful in specialized processes.

Iodic acid

Iodic acid, HIO3. It is a white water-soluble solid. Its robustness contrasts with the instability of chloric acid and bromic acid. Iodic acid features iodine in the oxidation state +5 and is one of the most stable oxo-acids of the halogens. When heated, samples dehydrate to give iodine pentoxide. On further heating, the iodine pentoxide further decomposes, giving a mix of iodine, oxygen and lower oxides of iodine.

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Lead dioxide

Lead(IV) oxide is the inorganic compound with the formula PbO2. It is an oxide where lead is in an oxidation state of +4. It is a dark-brown solid which is insoluble in water. It exists in two crystalline forms. It has several important applications in electrochemistry, in particular as the positive plate of lead acid batteries.

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Arsenic pentasulfide

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Polonium dichloride

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Compounds of thorium

Many compounds of thorium are known: this is because thorium and uranium are the most stable and accessible actinides and are the only actinides that can be studied safely and legally in bulk in a normal laboratory. As such, they have the best-known chemistry of the actinides, along with that of plutonium, as the self-heating and radiation from them is not enough to cause radiolysis of chemical bonds as it is for the other actinides. While the later actinides from americium onwards are predominantly trivalent and behave more similarly to the corresponding lanthanides, as one would expect from periodic trends, the early actinides up to plutonium have relativistically destabilised and hence delocalised 5f and 6d electrons that participate in chemistry in a similar way to the early transition metals of group 3 through 8: thus, all their valence electrons can participate in chemical reactions, although this is not common for neptunium and plutonium.

Compounds of aluminium

Aluminium (or aluminum) combines characteristics of pre- and post-transition metals. Since it has few available electrons for metallic bonding, like its heavier group 13 congeners, it has the characteristic physical properties of a post-transition metal, with longer-than-expected interatomic distances. Furthermore, as Al3+ is a small and highly charged cation, it is strongly polarizing and aluminium compounds tend towards covalency; this behaviour is similar to that of beryllium (Be2+), an example of a diagonal relationship. However, unlike all other post-transition metals, the underlying core under aluminium's valence shell is that of the preceding noble gas, whereas for gallium and indium it is that of the preceding noble gas plus a filled d-subshell, and for thallium and nihonium it is that of the preceding noble gas plus filled d- and f-subshells. Hence, aluminium does not suffer the effects of incomplete shielding of valence electrons by inner electrons from the nucleus that its heavier congeners do. Aluminium's electropositive behavior, high affinity for oxygen, and highly negative standard electrode potential are all more similar to those of scandium, yttrium, lanthanum, and actinium, which have ds2 configurations of three valence electrons outside a noble gas core: aluminium is the most electropositive metal in its group. Aluminium also bears minor similarities to the metalloid boron in the same group; AlX3 compounds are valence isoelectronic to BX3 compounds (they have the same valence electronic structure), and both behave as Lewis acids and readily from adducts. Additionally, one of the main motifs of boron chemistry is regular icosahedral structures, and aluminium forms an important part of many icosahedral quasicrystal alloys, including the Al–Zn–Mg class.

References

  1. ↑ Polyanskiy 1986, pp. 14–15.
  2. 1 2 3 4 Pauling, Linus (1947). General Chemistry . W.H. Freeman. ISBN   978-0-486-65622-9.
  3. 1 2 3 Brady, James E.; Holum, John R. (1996). Descriptive Chemistry of the Elements. John Wiley and Sons. ISBN   978-0-471-13557-9.
  4. ↑ Windholz, Martha (1976). Merck Index of Chemicals and Drugs, 9th ed., monograph 8393. Merck. ISBN   978-0-911910-26-1.
  5. ↑ Polyanskiy 1986, p. 21.
  6. 1 2 3 Polyanskiy 1986, p. 22.
  7. 1 2 Polyanskiy 1986, p. 28.
  8. 1 2 Polyanskiy 1986, p. 32.
  9. ↑ Polyanskiy 1986, p. 33.
  10. ↑ Polyanskiy 1986, p. 34.
  11. ↑ Zuckerman, J. J.; Hagen, A. P. (1989). Inorganic Reactions and Methods, the Formation of Bonds to Halogens. John Wiley & Sons. p. 426. ISBN   978-0-471-18656-4.
  12. 1 2 3 4 5 Puigdomenech, Ignasi (2004). Hydra/Medusa Chemical Equilibrium Database and Plotting Software. KTH Royal Institute of Technology. Archived from the original on 2007-09-29.
  13. ↑ Ward, C. H.; Hlousek, Douglas A.; Phillips, Thomas A.; Lowe, Donald F. (2000). Remediation of Firing Range Impact Berms. CRC Press. ISBN   1566704626.
  14. 1 2 Polyanskiy 1986, p. 43.
  15. ↑ Windholz, Martha (1976). Merck Index of Chemicals and Drugs, 9th ed., monograph 8393. Merck. ISBN   0-911910-26-3.
  16. 1 2 Polyanskiy 1986, p. 44.

Bibliography

Polyanskiy, N. G. (1986). Fillipova, N. A (ed.). Аналитическая химия элементов: Свинец[Analytical Chemistry of the Elements: Lead] (in Russian). Nauka.CS1 maint: ref=harv (link)

See also

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