1 Mm Lead

[Kemal Allan]

Lead is a chemical element; it has symbol Pb (from Latin plumbum) and atomic number 82. It is a heavy metal that is denser than most common materials. Lead is soft and malleable, and also has a relatively low melting point. When freshly cut, lead is a shiny gray with a hint of blue. It tarnishes to a dull gray color when exposed to air. Lead has the highest atomic number of any stable element and three of its isotopes are endpoints of major nuclear decay chains of heavier elements.

Lead is a relatively unreactive post-transition metal. Its weak metallic character is illustrated by its amphoteric nature; lead and lead oxides react with acids and bases, and it tends to form covalent bonds. Compounds of lead are usually found in the +2 oxidation state rather than the +4 state common with lighter members of the carbon group. Exceptions are mostly limited to organolead compounds. Like the lighter members of the group, lead tends to bond with itself; it can form chains and polyhedral structures.

Since lead is easily extracted from its ores, prehistoric people in the Near East were aware of it. Galena is a principal ore of lead which often bears silver. Interest in silver helped initiate widespread extraction and use of lead in ancient Rome. Lead production declined after the fall of Rome and did not reach comparable levels until the Industrial Revolution. Lead played a crucial role in the development of the printing press, as movable type could be relatively easily cast from lead alloys.[8] In 2014, the annual global production of lead was about ten million tonnes, over half of which was from recycling. Lead's high density, low melting point, ductility and relative inertness to oxidation make it useful. These properties, combined with its relative abundance and low cost, resulted in its extensive use in construction, plumbing, batteries, bullets and shot, weights, solders, pewters, fusible alloys, white paints, leaded gasoline, and radiation shielding.

Lead is a neurotoxin that accumulates in soft tissues and bones. It damages the nervous system and interferes with the function of biological enzymes, causing neurological disorders ranging from behavioral problems to brain damage, and also affects general health, cardiovascular, and renal systems. Lead's toxicity was first documented by ancient Greek and Roman writers, who noted some of the symptoms of lead poisoning, but became widely recognized in Europe in the late 19th century.

The sum of the first four ionization energies of lead exceeds that of tin,[9] contrary to what periodic trends would predict. This is explained by relativistic effects, which become significant in heavier atoms,[10] which contract s and p orbitals such that lead's 6s electrons have larger binding energies than its 5s electrons.[11] A consequence is the so-called inert pair effect: the 6s electrons of lead become reluctant to participate in bonding, stabilising the +2 oxidation state and making the distance between nearest atoms in crystalline lead unusually long.[12]

Lead's lighter carbon group congeners form stable or metastable allotropes with the tetrahedrally coordinated and covalently bonded diamond cubic structure. The energy levels of their outer s- and p-orbitals are close enough to allow mixing into four hybrid sp3 orbitals. In lead, the inert pair effect increases the separation between its s- and p-orbitals, and the gap cannot be overcome by the energy that would be released by extra bonds following hybridization.[13] Rather than having a diamond cubic structure, lead forms metallic bonds in which only the p-electrons are delocalized and shared between the Pb2+ ions. Lead consequently has a face-centered cubic structure[14] like the similarly sized[15] divalent metals calcium and strontium.[16][a][b][c]

Pure lead has a bright, shiny gray appearance with a hint of blue.[21] It tarnishes on contact with moist air and takes on a dull appearance, the hue of which depends on the prevailing conditions. Characteristic properties of lead include high density, malleability, ductility, and high resistance to corrosion due to passivation.[22]

With its high atomic number, lead is the heaviest element whose natural isotopes are regarded as stable; lead-208 is the heaviest stable nucleus. (This distinction formerly fell to bismuth, with an atomic number of 83, until its only primordial isotope, bismuth-209, was found in 2003 to decay very slowly.)[h] The four stable isotopes of lead could theoretically undergo alpha decay to isotopes of mercury with a release of energy, but this has not been observed for any of them; their predicted half-lives range from 1035 to 10189 years[41] (at least 1025 times the current age of the universe).

Apart from the stable isotopes, which make up almost all lead that exists naturally, there are trace quantities of a few radioactive isotopes. One of them is lead-210; although it has a half-life of only 22.2 years,[37] small quantities occur in nature because lead-210 is produced by a long decay series that starts with uranium-238 (that has been present for billions of years on Earth). Lead-211, -212, and -214 are present in the decay chains of uranium-235, thorium-232, and uranium-238, respectively, so traces of all three of these lead isotopes are found naturally. Minute traces of lead-209 arise from the very rare cluster decay of radium-223, one of the daughter products of natural uranium-235, and the decay chain of neptunium-237, traces of which are produced by neutron capture in uranium ores. Lead-213 also occurs in the decay chain of neptunium-237. Lead-210 is particularly useful for helping to identify the ages of samples by measuring its ratio to lead-206 (both isotopes are present in a single decay chain).[50]

Bulk lead exposed to moist air forms a protective layer of varying composition. Lead(II) carbonate is a common constituent;[52][53][54] the sulfate or chloride may also be present in urban or maritime settings.[55] This layer makes bulk lead effectively chemically inert in the air.[55] Finely powdered lead, as with many metals, is pyrophoric,[56] and burns with a bluish-white flame.[57]

Fluorine reacts with lead at room temperature, forming lead(II) fluoride. The reaction with chlorine is similar but requires heating, as the resulting chloride layer diminishes the reactivity of the elements.[55] Molten lead reacts with the chalcogens to give lead(II) chalcogenides.[58]

Lead metal resists sulfuric and phosphoric acid but not hydrochloric or nitric acid; the outcome depends on insolubility and subsequent passivation of the product salt.[59] Organic acids, such as acetic acid, dissolve lead in the presence of oxygen.[55] Concentrated alkalis will dissolve lead and form plumbites.[60]

Lead shows two main oxidation states: +4 and +2. The tetravalent state is common for the carbon group. The divalent state is rare for carbon and silicon, minor for germanium, important (but not prevailing) for tin, and is the more important of the two oxidation states for lead.[55] This is attributable to relativistic effects, specifically the inert pair effect, which manifests itself when there is a large difference in electronegativity between lead and oxide, halide, or nitride anions, leading to a significant partial positive charge on lead. The result is a stronger contraction of the lead 6s orbital than is the case for the 6p orbital, making it rather inert in ionic compounds. The inert pair effect is less applicable to compounds in which lead forms covalent bonds with elements of similar electronegativity, such as carbon in organolead compounds. In these, the 6s and 6p orbitals remain similarly sized and sp3 hybridization is still energetically favorable. Lead, like carbon, is predominantly tetravalent in such compounds.[61]

There is a relatively large difference in the electronegativity of lead(II) at 1.87 and lead(IV) at 2.33. This difference marks the reversal in the trend of increasing stability of the +4 oxidation state going down the carbon group; tin, by comparison, has values of 1.80 in the +2 oxidation state and 1.96 in the +4 state.[62]

Lead(II) compounds are characteristic of the inorganic chemistry of lead. Even strong oxidizing agents like fluorine and chlorine react with lead to give only PbF2 and PbCl2.[55] Lead(II) ions are usually colorless in solution,[63] and partially hydrolyze to form Pb(OH)+ and finally [Pb4(OH)4]4+ (in which the hydroxyl ions act as bridging ligands),[64][65] but are not reducing agents as tin(II) ions are. Techniques for identifying the presence of the Pb2+ ion in water generally rely on the precipitation of lead(II) chloride using dilute hydrochloric acid. As the chloride salt is sparingly soluble in water, in very dilute solutions the precipitation of lead(II) sulfide is instead achieved by bubbling hydrogen sulfide through the solution.[66]

Lead monoxide exists in two polymorphs, litharge α-PbO (red) and massicot β-PbO (yellow), the latter being stable only above around 488 C. Litharge is the most commonly used inorganic compound of lead.[67] There is no lead(II) hydroxide; increasing the pH of solutions of lead(II) salts leads to hydrolysis and condensation.[68]Lead commonly reacts with heavier chalcogens. Lead sulfide is a semiconductor, a photoconductor, and an extremely sensitive infrared radiation detector. The other two chalcogenides, lead selenide and lead telluride, are likewise photoconducting. They are unusual in that their color becomes lighter going down the group.[69]

Lead(II) sulfate is insoluble in water, like the sulfates of other heavy divalent cations. Lead(II) nitrate and lead(II) acetate are very soluble, and this is exploited in the synthesis of other lead compounds.[74]

Few inorganic lead(IV) compounds are known. They are only formed in highly oxidizing solutions and do not normally exist under standard conditions.[75] Lead(II) oxide gives a mixed oxide on further oxidation, Pb3O4. It is described as lead(II,IV) oxide, or structurally 2PbOPbO2, and is the best-known mixed valence lead compound. Lead dioxide is a strong oxidizing agent, capable of oxidizing hydrochloric acid to chlorine gas.[76] This is because the expected PbCl4 that would be produced is unstable and spontaneously decomposes to PbCl2 and Cl2.[77] Analogously to lead monoxide, lead dioxide is capable of forming plumbate anions. Lead disulfide[78] and lead diselenide[79] are only stable at high pressures. Lead tetrafluoride, a yellow crystalline powder, is stable, but less so than the difluoride. Lead tetrachloride (a yellow oil) decomposes at room temperature, lead tetrabromide is less stable still, and the existence of lead tetraiodide is questionable.[80]

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