This is a good article. Click here for more information.

Lead

From Wikipedia, the free encyclopedia
Jump to: navigation, search
This article is about the metal. For other uses, see Lead (disambiguation).
Lead,  82Pb
Lead electrolytic and 1cm3 cube.jpg
General properties
Name, symbol lead, Pb
Pronunciation /ˈlɛd/
LED
Appearance metallic gray
Lead in the periodic table
Hydrogen (diatomic nonmetal)
Helium (noble gas)
Lithium (alkali metal)
Beryllium (alkaline earth metal)
Boron (metalloid)
Carbon (polyatomic nonmetal)
Nitrogen (diatomic nonmetal)
Oxygen (diatomic nonmetal)
Fluorine (diatomic nonmetal)
Neon (noble gas)
Sodium (alkali metal)
Magnesium (alkaline earth metal)
Aluminium (post-transition metal)
Silicon (metalloid)
Phosphorus (polyatomic nonmetal)
Sulfur (polyatomic nonmetal)
Chlorine (diatomic nonmetal)
Argon (noble gas)
Potassium (alkali metal)
Calcium (alkaline earth metal)
Scandium (transition metal)
Titanium (transition metal)
Vanadium (transition metal)
Chromium (transition metal)
Manganese (transition metal)
Iron (transition metal)
Cobalt (transition metal)
Nickel (transition metal)
Copper (transition metal)
Zinc (transition metal)
Gallium (post-transition metal)
Germanium (metalloid)
Arsenic (metalloid)
Selenium (polyatomic nonmetal)
Bromine (diatomic nonmetal)
Krypton (noble gas)
Rubidium (alkali metal)
Strontium (alkaline earth metal)
Yttrium (transition metal)
Zirconium (transition metal)
Niobium (transition metal)
Molybdenum (transition metal)
Technetium (transition metal)
Ruthenium (transition metal)
Rhodium (transition metal)
Palladium (transition metal)
Silver (transition metal)
Cadmium (transition metal)
Indium (post-transition metal)
Tin (post-transition metal)
Antimony (metalloid)
Tellurium (metalloid)
Iodine (diatomic nonmetal)
Xenon (noble gas)
Caesium (alkali metal)
Barium (alkaline earth metal)
Lanthanum (lanthanide)
Cerium (lanthanide)
Praseodymium (lanthanide)
Neodymium (lanthanide)
Promethium (lanthanide)
Samarium (lanthanide)
Europium (lanthanide)
Gadolinium (lanthanide)
Terbium (lanthanide)
Dysprosium (lanthanide)
Holmium (lanthanide)
Erbium (lanthanide)
Thulium (lanthanide)
Ytterbium (lanthanide)
Lutetium (lanthanide)
Hafnium (transition metal)
Tantalum (transition metal)
Tungsten (transition metal)
Rhenium (transition metal)
Osmium (transition metal)
Iridium (transition metal)
Platinum (transition metal)
Gold (transition metal)
Mercury (transition metal)
Thallium (post-transition metal)
Lead (post-transition metal)
Bismuth (post-transition metal)
Polonium (post-transition metal)
Astatine (metalloid)
Radon (noble gas)
Francium (alkali metal)
Radium (alkaline earth metal)
Actinium (actinide)
Thorium (actinide)
Protactinium (actinide)
Uranium (actinide)
Neptunium (actinide)
Plutonium (actinide)
Americium (actinide)
Curium (actinide)
Berkelium (actinide)
Californium (actinide)
Einsteinium (actinide)
Fermium (actinide)
Mendelevium (actinide)
Nobelium (actinide)
Lawrencium (actinide)
Rutherfordium (transition metal)
Dubnium (transition metal)
Seaborgium (transition metal)
Bohrium (transition metal)
Hassium (transition metal)
Meitnerium (unknown chemical properties)
Darmstadtium (unknown chemical properties)
Roentgenium (unknown chemical properties)
Copernicium (transition metal)
Nihonium (unknown chemical properties)
Flerovium (post-transition metal)
Moscovium (unknown chemical properties)
Livermorium (unknown chemical properties)
Tennessine (unknown chemical properties)
Oganesson (unknown chemical properties)
Sn

Pb

Fl
thalliumleadbismuth
Atomic number (Z) 82
Group, block group 14 (carbon group), p-block
Period period 6
Element category   post-transition metal
Standard atomic weight (±) (Ar) 207.2(1)[1]
Electron configuration [Xe] 4f14 5d10 6s2 6p2
Electrons per shell
2, 8, 18, 32, 18, 4
Physical properties
Phase solid
Melting point 600.61 K ​(327.46 °C, ​621.43 °F)
Boiling point 2022 K ​(1749 °C, ​3180 °F)
Density near r.t. 11.34 g/cm3
when liquid, at m.p. 10.66 g/cm3
Heat of fusion 4.77 kJ/mol
Heat of vaporization 179.5 kJ/mol
Molar heat capacity 26.650 J/(mol·K)
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 978 1088 1229 1412 1660 2027
Atomic properties
Oxidation states 4, 3, 2, 1, −1, −2, −4 ​(an amphoteric oxide)
Electronegativity Pauling scale: 1.87
Ionization energies 1st: 715.6 kJ/mol
2nd: 1450.5 kJ/mol
3rd: 3081.5 kJ/mol
Atomic radius empirical: 175 pm
Covalent radius 146±5 pm
Van der Waals radius 202 pm
Miscellanea
Crystal structure face-centered cubic (fcc)
Face-centered cubic crystal structure for lead
Speed of sound thin rod 1190 m/s (at r.t.) (annealed)
Thermal expansion 28.9 µm/(m·K) (at 25 °C)
Thermal conductivity 35.3 W/(m·K)
Electrical resistivity 208 nΩ·m (at 20 °C)
Magnetic ordering diamagnetic
Magnetic susceptibility (χmol) −23.0·10−6 cm3/mol (at 298 K)[2]
Young's modulus 16 GPa
Shear modulus 5.6 GPa
Bulk modulus 46 GPa
Poisson ratio 0.44
Mohs hardness 1.5
Brinell hardness 38–50 MPa
CAS Number 7439-92-1
History
Discovery by Middle Easterns (7000 BCE)
Most stable isotopes of lead
iso NA half-life DM DE (MeV) DP
202Pb syn 53000 y ε 0.0497 202Tl
204Pb 1.4% is stable with 122 neutrons
205Pb trace 1.53×107 y ε 0.051 205Tl
206Pb 24.1% is stable with 124 neutrons
207Pb 22.1% is stable with 125 neutrons
208Pb 52.4% is stable with 126 neutrons
210Pb trace 22.3 y β 0.064 210Bi
| references | in Wikidata

Lead (/lɛd/) (from the Old English léad) is a chemical element with atomic number 82 and symbol Pb (after the Latin plumbum). When freshly cut, it has a bluish-white color that soon tarnishes to a dull gray upon exposure to air. Lead is a soft, malleable, and heavy metal; its density of 11.34 g/cm3 exceeds that of most common materials. Lead has the second highest atomic number of all practically stable elements. As such, it is located at the end of three major decay chains of heavier elements, which, in part, accounts for lead's relative abundance: its stability exceeds those of other similarly numbered elements.

Lead is a post-transition metal and is relatively inert unless powdered. Its weak metallic character is illustrated by its general amphoteric nature: lead and lead oxides react with both acids and bases. Lead also displays a tendency toward covalent bonding. Compounds of lead are most commonly found in the +2 oxidation state, rather than +4, unlike the lighter group 14 elements of the periodic table; exceptions are mostly limited to organolead compounds. Like the lighter group 14 elements, lead exhibits a tendency to bond to itself; it can form chains, rings, and polyhedral structures.

Lead is easily extracted from ore, and it was known to prehistoric people in Western Asia. A principal ore of lead, galena, often bears silver, and interest in silver helped initiate widespread lead production and use in ancient Rome. Lead production declined after the fall of Rome and did not reach comparable levels until the Industrial Revolution. Today, lead is produced in quantities of around ten thousand tonnes annually; secondary production from recycling is gaining ground, accounting for around half of that figure.

Lead has several properties that make it useful: high density, low melting point, ductility, and relative inertness to oxidation. Combined with its relative abundance and low cost, these factors have led to its widespread use, including in building construction, batteries, bullets and shot, weights, solders, pewters, fusible alloys, and for radiation shielding. In the late nineteenth century, lead came to be recognized as poisonous, and since that time, lead has been and is being phased out for many applications. Lead is a neurotoxin that accumulates in soft tissues and bones, damaging the nervous system and causing brain disorders. Lead can also cause blood disorders in mammals.

Physical properties[edit]

Atomic[edit]

A neutral lead atom has 82 electrons, arranged in an electronic configuration of [Xe]4f145d106s26p2. The combined first and second ionization energies—the total energy required to remove the two 6p electrons—is close to that of tin, lead's upper neighbor in group 14. This is unusual since ionization energies generally fall going down a group as an element's electrons become more distant from its nucleus. The similarity is attributable to the lanthanide contraction— the decrease in element radii from atomic number 57, lanthanum, to 71, lutetium, and the relatively small ionic radii of the subsequent elements starting with 72, hafnium. This contraction results from poor shielding of the nucleus by the lanthanide 4f electrons; the outer electrons are drawn towards the nucleus, resulting in smaller atomic radii. The combined first four ionization energies of lead exceed those of tin,[3] contrary to what the periodic trends would predict. For this reason lead, unlike tin,[4] rarely has a +4 oxidation state in inorganic compounds,[4] and exhibits the +2 state instead. Such behavior is attributable to relativistic effects, which become particularly prominent at the bottom of the periodic table;[4] the result is that the 6s electrons of lead become reluctant to participate in bonding[a] (a phenomenon referred to as the inert pair effect), and the distance between nearest atoms in crystalline lead is unusually long.[6]

Aside from lead, the lighter elements in group 14 have a stable or metastable allotrope in which they crystallize, via covalent bonding, in the diamond cubic structure. Each atom is tetrahedrally coordinated, indicating that all four bonds are equivalent and have attained the lowest possible energy. Despite the fact that two of the electrons are in s-orbitals and the other two in higher-energy p-orbitals, the orbitals are close enough in energy to allow mixing into four hybrid sp3 orbitals. In lead, on the other hand, the inert pair effect means that the separation is larger, and is not enough to be compensated by the energy that would be released from the additional bonds formed.[7] Thus, rather than having the diamond-cubic covalent 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 (like that of the similarly sized divalent metals calcium and strontium).[8] In 2013, quasicrystalline allotropes of lead with pentagonal symmetry have also been characterized.[b]

Bulk[edit]

Freshly prepared or fractured lead has a bright silvery appearance with a very slight hint of blue.[11] Lead otherwise tarnishes on contact with moist air, giving it a dull appearance the hue of which will vary depending on the prevailing conditions. The characteristic properties of lead include high density, softness, malleability, ductility, poor electrical conductivity compared to other metals, high resistance to corrosion (conferred by its surface patina), and a propensity to react with organic reagents.[11]

A sample of lead solidified from the molten state

Lead's face-centered cubic structure and high atomic weight give it a density[12] of 11.34 g/cm3. This figure exceeds that of common metals such as iron (7.87 g/cm3), copper (8.93 g/cm3), and zinc (7.14 g/cm3), [13] and is the origin of the idiom to go over like a lead balloon.[14] Some rarer metals are denser: tungsten and gold are both 19.3 g/cm3, and osmium— the densest metal known—has a density of 22.59 g/cm3, almost twice that of lead.[15]

Lead is a very soft metal with a Mohs hardness of 1.5; it can be scratched with a fingernail.[16] It is malleable and ductile,[c] with its malleability exceeding its ductility.[17] The bulk modulus—a measure of the ease of compressibility of a material—of lead is 45.8 GPa. (For comparison, that of aluminium is 75.2 GPa; copper 137.8 GPa; and mild steel 160–169 GPa.)[18] Lead's tensile strength is comparatively low: 12–17 MPa (that of aluminium is 6 times higher; copper 10 times higher; mild steel 15 times higher); this value is easily improved by adding small concentrations of other metals or metalloids, such as copper or antimony.[19]

The melting point of lead—at 327.5 °C (621.5 °F)[20]—is low compared to most metals.[21][d] Its boiling point is 1749 °C (3180 °F).[23] The electrical resistivity of lead at 20 °C is 208 nanoohm-meters, almost an order of magnitude higher than those of other industrial metals (that of copper is 17.12 nΩ·m; gold 22.55 nΩ·m; aluminium 27.09 nΩ·m). Lead is a superconductor at temperatures lower than 7.19 K;[24] this is the highest critical temperature of all type-I superconductors and the third highest of all elemental superconductors.

Isotopes[edit]

Main article: Isotopes of lead

Lead has four stable isotopes, lead-204, lead-206, lead-207, and lead-208.[25] The high number of such isotopes is due to lead's atomic number of 82 being even,[e] as well as a magic number (meaning lead's protons are arranged into complete shells within its atomic nucleus).[f] With its high atomic number, lead is the second-heaviest element that occurs naturally in the form of isotopes regarded as stable: bismuth has a higher atomic number of 83, but its only primordial isotope was found in 2003 to decay at an extremely gradual rate.[g] 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;[26] accordingly, their predicted half-lives are extremely long, ranging up to over 10100 years.[29][h] Lead is therefore often quoted as the heaviest stable element.[26]

The Holsinger meteorite, representing the largest piece of the Canyon Diablo meteorite. Uranium–lead dating and lead–lead dating on this meteorite allowed refinement of the age of the Earth to 4.55 billion ± 70 million years.

Three of the stable isotopes are found in three of the four major decay chains: lead-206, lead-207, and lead-208 are the final decay products of uranium-238, uranium-235, and thorium-232, respectively; these decay chains are called the uranium series, actinium series, and thorium series. Their isotopic concentration in a natural rock sample depends on the presence of other elements: for example, the relative amount of lead-208 can range from 52.4% in normal samples to 90% in thorium ores.[30] (For this reason, the atomic weight of lead is given to only one decimal place.[31]) As time passes, the ratio of lead-206 and lead-207 to lead-204 increase, since the former two are supplemented by radioactive decay of heavier elements and the latter is not; this allows for lead–lead dating. Analogously, as uranium decays (eventually) into lead, their relative amounts change; this is the basis for uranium–lead dating.[32]

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 22.3 years[26]—a period too short to allow any primordial lead-210 to exist—some small non-primordial quantities of it occur in nature. This is because lead-210 is found in the uranium series. Even though it constantly decays away, it is regenerated by decay of its polonium-214 parent which, in turn, is supplemented by the decay of its parent, and so on, all the way up to original uranium-238 (which has been present for billions of years on Earth). Lead-214, -212, and -211 are present in the decay chains of natural uranium-238; thorium-232, and uranium-235; therefore, traces of all three of these lead isotopes are also found naturally. Lastly, very minute traces of lead-209 are present from the cluster decay of radium-223, one of the daughter products of natural uranium-235. Hence, natural lead consists of not only the four stable isotopes but also minute traces of another five short-lived radioisotopes.[33] Lead-210 is particularly useful for helping to identify ages of samples containing it, which is performed by measuring lead-210 to lead-206 ratios (both isotopes are present in a single decay chain).[34][35]

In total, thirty-eight lead sotopes have been synthesized, with mass numbers 178–215.[26] Lead-205 is the most stable radioisotope of lead, with a half-life of around 1.5×107 years.[i] The second-most stable is the synthetic lead-202, which has a half-life of about 53,000 years, longer than any of the natural trace radioisotopes.

Inorganic compounds[edit]

Lead shows two main oxidation states: +4 and +2. The tetravalent state is common for group 14. The divalent state is rare for carbon and silicon, minor for germanium, important (but not prevailing) for tin, and is the more important for lead.[37] 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, for example, 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. This is not quite as applicable to compounds in which lead forms covalent bonds to elements of similar electronegativity such as carbon in organolead compounds. In them, the 6s and 6p orbitals remain similarly sized and sp3 hybridization in compounds is still energetically favorable. Lead, like carbon, is predominantly tetravalent in such compounds.[38]

Crystalline form of lead(II) sulfate

There is a sizeable difference in the electronegativity values for 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 down group 14; tin, by comparison, has electronegativity values of 1.80 and 1.96 in the +2 and +4 oxidation states.[39]

Lead(II)[edit]

Lead(II) compounds are characteristic of the inorganic chemistry of lead. Even strong oxidizing agents like fluorine and chlorine react with lead at room temperature to give only PbF2 and PbCl2.[40] It forms binary compounds with many nonmetals, but not all of them; for example, there is no known lead carbide.[41]

Most lead(II) compounds are ionic, but they are not as ionic as those of many other metals and many are thus water-insoluble. In solution, lead(II) ions are colorless, but under specific conditions, lead is capable of changing its color.[42] Unlike tin(II) ions, they do not react as reducing agents in solution. Lead(II) ions partially hydrolyze in aqueous solution to form Pb(OH)+ and finally Pb4(OH)4 (in which the hydroxyls ions act as bridging ligands).[43]

     Lead and      oxygen in a tetragonal unit cell of lead(II,IV) oxide

Lead monoxide exists in two polymorphs, red α-PbO and yellow β-PbO, the latter being stable only above around 488 °C. It is the most commonly used compound of lead.[44] Its hydroxide counterpart, lead(II) hydroxide, is not capable of existence outside of solution; in solution, it is known to form plumbite anions.

Lead commonly reacts with the heavier chalcogens. Lead sulfide can only be dissolved in strong acids.[45] It is a semiconductor, a photoconductor, and an extremely sensitive infrared radiation detector. A mixture of the monoxide and the monosulfide, when heated, forms the metal.[46] The other two chalcogenides are likewise photo-conducting. They are unusual in that their color becomes lighter down the group.[40]

Lead dihalides are well-characterized; this includes the diastatide,[47] and mixed examples, such as PbFCl. The relative insolubility of the latter forms a useful basis for the gravimetric determination of fluorine. The difluoride was the first ionically conducting compound to be discovered (in 1838, by Michael Faraday). The other dihalides decompose on exposure to ultraviolet or visible light, especially the diiodide.[48] Many pseudohalides are also known.[40] Lead(II) forms a tremendous variety of coordination complexes, such as [PbCl4]2−, [PbCl6]4−, and the chain anion [Pb2Cl9]n5n, noting most are not yet adequately characterized structurally.[48]

Lead(II) sulfate is well known for its insolubility in water, like the sulfates of the other heavy divalent cations. Lead(II) nitrate and lead(II) acetate, in contrast, are very soluble, and this property is exploited in the synthesis of other lead compounds.[49]

The 5s electron pair tends to be stereochemically active in tin(II) compounds but is much less so in lead(II) compounds. Consequently, there are often structural similarities between lead(II) compounds and analogous compounds of the divalent cations of calcium, strontium, barium, europium, and ytterbium.[48]

Lead(IV)[edit]

Few inorganic lead(IV) compounds are known, and they are typically strong oxidants or exist only in highly acidic solutions.[4] Lead(II) oxide gives a mixed oxide on further oxidation, Pb
3
O
4
. It is described as lead(II,IV) oxide, or structurally 2PbOPbO
2
, and is the best-known mixed valence lead compound. Lead dioxide is a strong oxidizing agent, capable of oxidizing hydrochloric acid to chlorine gas. This is because the expected PbCl4 that would be produced is unstable and spontaneously decomposes to PbCl2 and Cl2. Analogously to lead monoxide, lead dioxide is capable of forming plumbate anions. Lead disulfide, like the monosulfide, is a semiconductor.[50] Lead(IV) selenide is also known.[51] Lead tetrafluoride, a yellow crystalline powder, is stable, but less stable 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.[52][53]

Other oxidation states[edit]

See also: Plumbide
The capped square antiprismatic anion [Pb9]4− from [K(18-crown-6)]2K2Pb9·(en)1.5[54]

Some lead compounds exist in formal oxidation states other than +4 or +2. Lead(III) may be obtained as an intermediate between lead(II) and lead(IV), in larger organolead complexes (rather than on its own).[55][56] This oxidation state is not specifically stable, as the lead(III) ion (and, consequently, the larger complexes containing it) is a radical; the same applies for lead(I), which can also be found in such species.[57]

Numerous mixed lead(II,IV) oxides are known. When PbO2 is heated in air, it becomes Pb12O19 at 293 °C, Pb12O17 at 351 °C, Pb3O4 at 374 °C, and finally PbO at 605 °C. A further sesquioxide Pb2O3 can be obtained at high pressure, along with several non-stoichiometric phrases. Many of them show defect fluorite structures in which some oxygen atoms are replaced by vacancies: for instance, PbO can be considered as having such a structure, with every alternate layer of oxygen atoms absent.[58]

Negative oxidation states can occur as Zintl phases, as either free lead anions, for example, in Ba
2
Pb
, with lead formally being lead(−IV),[59] or in oxygen-sensitive cluster ions, for example, in a trigonal bipyramidal Pb2−
5
ion, where two lead atoms are lead(−I) and three are lead(0).[60] In such anions, each atom is at a polyhedral vertex and contributes two electrons to each covalent bond along an edge from their sp3 hybrid orbitals, the other two being an external lone pair.[43] They may be made in liquid ammonia via the reduction of lead by sodium.[61]

Organolead[edit]

Main article: Organolead compound
Structure of a tetraethyllead molecule:
     Carbon
     Hydrogen
     Lead

Lead can form long singly or multiply bonded chains—catenas—and so shares some covalent chemistry with its lighter homolog carbon. This tendency is much lower for lead because the Pb–Pb bond energy (98 kJ/mol) is much lower than that of the C–C bond (356 kJ/mol).[62] Lead atoms can build metal–metal bonds of an order up to three.[63] Lead also forms covalent bonds with carbon to produce organolead compounds similar to, but generally less stable than, typical organic compounds,[64] as the Pb–C bond is rather weak.[43] Consequently the organometallic chemistry of lead is comparably narrow: it is far less wide-ranging than that of tin.[65] It predominantely forms organolead(IV) compounds. Very few organolead(II) compounds are known: even starting with inorganic lead(II) reactants always results in organolead(IV) products. The most well-characterized exceptions are the purple bis(disyl)plumbylene, Pb[CH(SiMe)3)2]2 and lead cyclopentadienide, Pb(η5-C5H5)2.[65]

The simplest organic compound of lead is plumbane, the analog of methane. Plumbane may be obtained in a reaction between metallic lead and atomic (not molecular) hydrogen.[66] It is unstable but two simple derivatives, tetramethyllead and tetraethyllead, are the best-known organolead compounds. They may be made by the addition of trimethyllead or triethyllead to alkenes or alkynes; these precursors may themselves be made from the corresponding lead halides and lithium aluminium hydride at −78 °C. These compounds are relatively stable—tetraethyllead only starts to decompose at 100 °C (210 °F)[64]—or if exposed to sunlight or ultraviolet light.[67] (Tetraphenyllead is even more thermally stable, decomposing only at 270 °C.)[65] With sodium metal, lead readily forms an equimolar alloy that reacts with alkyl halides to form organometallic compounds such as tetraethyllead.[68] The oxidizing nature of many organolead compounds is usefully exploited: lead tetraacetate is an important laboratory reagent for oxidation in organic chemistry;[69] tetraethyllead was once produced in larger quantities than any other organometallic compound.[70] Other organolead compounds, including homologs of said compounds, are less chemically stable;[64] a lead analog of the next alkaneethane—is not even known.[66] Polyplumbanes are not well-characterized and are generally highly thermally unstable and reactive.[65]

Origin and occurrence[edit]

In space[edit]

Solar System abundances[71]
Atomic
number
Element Relative
amount
42 Molybdenum 0.798
46 Palladium 0.440
50 Tin 1.146
78 Platinum 0.417
80 Mercury 0.127
82 Lead 1
90 Thorium 0.011
92 Uranium 0.003

Lead is uncommon in space—its Solar System per-particle abundance is 0.121 ppb (parts per billion).[71][j] Even so, this figure is two and a half times higher than that of platinum, eight times that of mercury, and seventeen times that of gold.[71] The amount of lead in the universe is slowly increasing[72] since most heavier atoms (all of which are unstable) gradually decay to lead. As an example, the abundance of lead in the Solar System since its formation some 4.5 billion years ago has increased by about 0.75%.[71]

Primordial lead—which comprises the isotopes lead-204, lead-206, lead-207, and lead-208—was mostly created as a result of repetitive neutron capture processes occurring in stars. The two main modes of capture are the s-process and the r-process.

In the s-process (s is for "slow"), captures are separated by years or decades, allowing less stable nuclei to beta decay. For example, a stable thallium-203 nucleus captures a neutron and becomes thallium-204; this is unstable, and undergoes beta decay to give stable lead-204; on capturing another neutron, it becomes lead-205, which is stable enough to generally last longer than a capture takes (its half-life is around 15 million years). Further captures result in lead-206, lead-207, and lead-208. On capturing another neutron, lead-208 becomes lead-209, which quickly decays into bismuth-209. On capturing another neutron, bismuth-209 becomes bismuth-210 and this undergoes alpha decay into thallium-206 (which would beta decay into lead-206), or beta decay to yield polonium-210 (which would inevitably alpha decay to lead-206). The cycle ends at lead-206, lead-207, lead-208, and bismuth-209.[73]

Chart of the final part of the s-process, from mercury to polonium. Red lines and circles represent neutron captures; blue arrows represent beta decays; green arrows represent alpha decays; cyan arrows represent electron captures.

In the r-process (r is for "rapid"), captures happen faster than nuclei can decay. This occurs in environments with a high neutron density, possibly in a supernova or during the merger of two neutron stars. The neutron flux involved may be on the order of 1022 neutrons/(cm2·second).[74] The r-process does not form as much lead as the s-process. This is because the r-process tends to stop once very neutron-rich nuclei reach 126 neutrons. At this point, the neutrons are arranged in complete shells within the atomic nucleus, and it becomes harder to energetically accommodate more of them. When the neutron flux subsides, these nuclei beta decay into stable isotopes of osmium, iridium, and platinum.[73]

On Earth[edit]

Lead is classified as a chalcophile under the Goldschmidt classification, meaning it is generally found combined with sulfur. Many lead minerals are relatively light and, over the course of the Earth's history, have remained in the crust instead of sinking into the Earth's interior. This accounts for lead's relatively high crustal abundance of 14 ppm.[75]

Lead is a fairly common element in the Earth's crust for its high atomic number.

The main lead-bearing mineral is galena (PbS), which is mostly found with zinc ores.[76] Most other lead minerals are related to galena in some way; for example, boulangerite, Pb
5
Sb
4
S
11
, is a mixed sulfide derived from galena; anglesite, PbSO
4
, is a product of galena oxidation; and cerussite or white lead ore, PbCO
3
, is a decomposition product of galena. Arsenic, tin, antimony, silver, gold, and bismuth are other common impurities in lead minerals.[76]

Lead is easily extracted from its ores usually together with zinc and copper, which are also chalcophiles.[76] Metallic lead occurs in nature but is rare. World lead resources exceed 2 billion tons.[77] Significant deposits are located in Australia, China, Ireland, Mexico, Peru, Portugal, Russia, and the United States. Global reserves—resources that are economically feasible to extract—totaled 89 million tons in 2015, of which Australia had 35 million, China 15.8 million, and Russia 9.2 million.[77]

Typical background concentrations of lead do not exceed 0.1 μg/m3 in the atmosphere; 100 mg/kg in soil; and 5 μg/L in freshwater and seawater.[78]

Etymology[edit]

The English word "lead" is of Germanic origin; it comes from the Middle English leed and Old English lēad (with the macron above the "e" signifying that the vowel sound of that letter is long).[79] The Old English word is derived from the hypothetical reconstructed Proto-Germanic *lauda- ("lead").[80] In turn, this is thought to have originated in either the c. 3500 BCE Proto-Indo-European *lAudh- ("lead"; capitalization of the vowel is equivalent to the macron),[81] or the later Proto-Celtic *ɸloud-io- ("lead").[k]

The name of the chemical element is not related to the verb of the same spelling, which is instead derived from (eventually) the Proto-Germanic *laidijan- ("to lead").[83]

History[edit]

World lead production peaking in the Roman period and the rising Industrial Revolution
(Years B.P. = years before 1950)[84]

Prehistory and early history[edit]

Metallic lead beads dating back to at least 7000–6500 BCE have been found in Asia Minor and may represent the first example of metal smelting.[85] At this time lead had few (if any) applications due to its softness and dull appearance.[85] The major reason for the spread of lead production, rather than its utility, was its association with silver, which may be obtained by burning galena (a widespread lead mineral).[86][87] The Ancient Egyptians were the first to use lead in cosmetics, an application that would spread to Ancient Greece and beyond;[88] the Egyptians may have used lead for sinkers in fishing nets, in glazes, glasses and enamels, and for ornaments.[86] Various civilizations of the Fertile Crescent used lead as a writing material, as currency, and for construction.[86] Lead was used in the Ancient Chinese royal court as a stimulant,[86] as currency,[89] and as a contraceptive;[90] the Indus Valley civilization and the Mesoamericans[86] used it for making amulets; and the eastern and southern Africa peoples used lead in wire drawing.[91]

Classical era[edit]

This metal was by far the most used material in classical antiquity, and it is appropriate to refer to the (Roman) Lead Age. Lead was to the Romans what plastic is to us.

Heinz Eschnauer and Markus Stoeppler
"Wine—An enological specimen bank", 1992[92]

Because silver was extensively used as a decorative material and an exchange medium, lead deposits came to be worked in Asia Minor from 3000 BCE[87] and, subsequently, from 2000 BCE in the Iberian peninsula by the Phoenicians;[93] and in Athens, Carthage, and Sicily.[86]

Rome's territorial expansion in Europe and across the Mediterranean, and its concurrent development of mining led to it becoming the greatest producer of lead during the classical era, with an estimated annual output peaking at 80,000 tonnes. Like their predecessors, the Romans obtained lead mostly as a by-product of silver smelting.[84][94][95] Lead mining occurred in Central Europe, Britain, the Balkans, Greece, Anatolia, and Hispania, and accounted for 40% of world production.[84]

Lead was used for making water pipes in the Roman Empire; so much so that the Latin word for the metal, plumbum, was the origin of the English word "plumbing" and its derivatives. Its formability and resistance to corrosion[96] ensured its widespread use in other applications ranging from pharmaceuticals to roofing, and currency to warfare.[97][98][99] Writers of the time, such as the Cato the Elder, Columella, and Pliny the Elder, recommended lead (or lead-coated) vessels for the preparation of sweeteners and preservatives added to wine and food. The lead conferred an agreeable taste due to the formation of "sugar of lead" (lead(II) acetate) whereas copper or bronze vessels could impart a bitter flavour on account of verdigris formation.[100] Vitruvius, in contrast, reported the health dangers of lead.[101]

Ancient pipes in a museum case
Lead Roman pipes[l]

It has been suggested that lead poisoning played a major role in the decline of the Roman Empire.[102][103][m] Other researchers have strongly criticized such claims, citing lack of the errors in linking fall of Rome to lead poisoning and even "false evidence";[105][106] according to archaeological research, Roman lead pipes increased lead levels in tap water but such an effect was "unlikely to have been truly harmful".[107][108] Lead poisoning—a condition in which one becomes dark and cynical—was called "saturnine" after the ghoulish father of the gods, Saturn. By association, lead was considered the father of all metals.[109] In contrast, its social status was low as it was readily available in Roman society.[110]

Confusion with tin and antimony[edit]

During the classical era (and even up to the 17th century), tin was often not clearly distinguished from lead: Romans called lead plumbum nigrum ("black lead"), and tin plumbum candidum ("bright lead"). The association of lead and tin can be seen in other languages: the word olovo in Czech translates to "lead", but in Russian the cognate олово (olovo) means "tin".[111] Lead also bore a close relation to antimony: both elements commonly occur as sulfides (galena and stibnite), often together. Pliny declared that stibnite would give lead on heating, whereas the mineral produced on heating was actually antimony.[112] The originally South Asian surma—"galena" in English—spread across Asia with that meaning, and gave its name to antimony in a number of Central Asian languages, and in Russian.

Middle Ages and the Renaissance[edit]

Elizabeth I of England was one European monarch commonly depicted with a whitened face. Lead in face whiteners is thought to be one reason for her death.[113]

After the fall of the Western Roman Empire and into the medieval era, lead continued to be used in plumbing in Western Europe,[114] but lead mining in Europe declined, with the only region having a significant production being Arabian Iberia.[115][116] The largest production of lead occurred in South and East Asia, especially China and India, where lead output underwent strong growth.[116] In Europe, lead production only began to revive in the 11th and 12th centuries, where it was again used for roofing and piping; from the 13th century, it was used to create stained glass.[116] During the period, lead was used increasingly for adulterating wine. This practice was declared forbidden in 1498 by a papal bull, but it continued long past that time and resulted in numerous mass poisonings up to the late 18th century.[115] Lead was a key material in parts of the printing press, which was invented around 1440, and lead dust was commonly inhaled by press operators, causing lead poisoning.[117] Firearms were invented at around the same time, and lead, despite being more expensive than iron, became the chief material for making bullets. It was less damaging to iron gun barrels, had a higher density (which allowed for better retention of velocity) and its lower melting point made the production of bullets easier as they could be made using a wood fire.[118] Lead was extensively used in cosmetics by Western European aristocracy, as whitened faces were regarded as a sign of modesty.[88] This practice eventually expanded to white wigs and eyeliners, and only faded out with the French Revolution in the late 18th century. A similar fashion appeared in Japan in the 18th century with the emergence of the geishas, a practice that continued long into the 20th century. The white face became a "symbol of a Japanese woman"; lead was commonly used as the whitener.[119][120][121]

Outside Europe and Asia[edit]

In the New World, lead was produced soon after the arrival of European settlers. The earliest recorded lead production dates to 1621, in the English Colony of Virginia, fourteen years after its foundation.[122] In Australia, the first mine opened by colonists on the continent was a lead mine, in 1841.[123] Centuries before the Europeans were able to start colonizing Africa in the late 19th century, lead mining was known in the Benue Trough[124] and the lower Congo basin, where lead was used for trade with Europeans, and as a currency.[125][n]

The Industrial Revolution[edit]

Lead mining in the upper Mississippi River region in the United States in 1865

In the second half of the 18th century Britain, and later continental Europe and the United States, experienced the Industrial Revolution. During the period, lead mining proved important; the Industrial Revolution was the first time during which lead production rates exceeded those of Rome.[84] Britain was the leading producer, losing this status by the mid-19th century with the depletion of its mines and the development of lead mining in Germany, Spain, and the United States.[126] Lead production in the United States dominated by 1900;[127] other non-European nations—Canada, Mexico, and Australia—started massive lead production activities, and by 1900, Europe's output of lead fell below that elsewhere.[128] A great share of the demand for lead came from plumbing and painting—lead paints had been invented and were regularly used.[129] At this time, more people—the working class—contacted the metal; lead poisoning cases escalated. This led to research into the effects of lead intake: lead was proven to be more dangerous in its fume form than as a solid metal; lead poisoning and gout were linked (Alfred Baring Garrod noted a third of his gout patients were plumbers and painters); and the effects of chronic ingestion of lead, including mental disorders, were all studied in the 19th century. The first laws to decrease the degree of lead poisoning in factories followed during the 1870s and 1880s in the United Kingdom.[129]

Into the modern era[edit]

Promotional poster for Dutch Boy lead paint, United States, 1912

Further evidence of the threat that lead posed to humans was discovered in the late 19th and early 20th centuries—mechanisms of harm were better understood, and lead blindness was documented[130]—and countries in Europe and the United States started efforts to reduce the amount of lead that people came into contact with. The last major innovation to expose humans to contact with lead was the adding of tetraethyllead to gasoline as an antiknock agent, a practice that originated in the United States in 1921; it was phased out there, and in the European Union, by 2000.[129] Most European countries banned usage of lead paint—commonly used to this point because of its opacity and water resistance[131]—for interiors by 1930.[132] The result of many regulations and bans put on lead products was significant: in the last quarter of the 20th century, the percentage of people with excessive lead blood levels dropped from over three-quarters of the population to slightly over two percent in the United States.[129] By the end of the 20th century, the main product made of lead was the lead–acid battery,[133] which possesses no direct threat to humans. This facilitated consistent lead production in industrialized countries. From 1960 to 1990, lead output in the Western Bloc grew by 31%.[134] The share of the world's lead production by the Eastern Bloc increased from 10% to 30% from 1950 to 1990, with the Soviet Union being world's largest producer during the mid-1970s and the 1980s, and China starting massive lead production in the late 20th century.[135] Unlike the European communist countries, China was largely unindustrialized by the mid-20th century; in 2004, China surpassed Australia as the largest producer of lead.[136] As with European industrialization, lead has had a negative effect on health in China.[137]

Production[edit]

Historical evolution of the production of lead, as extracted in different countries

Production and consumption of lead is increasing worldwide (due to its use in lead–acid batteries). There are two major categories of production: primary, from mined ores; and secondary from scrap. In 2013, 4.74 million metric tons came from primary production and 5.74 million tons from secondary production. The top mining countries for lead in that year were China, Australia, Russia, India, Bolivia, Sweden, North Korea, South Africa, Poland, and Ireland. The top lead producing countries were China, United States, India, South Korea, Germany, Mexico, United Kingdom, Canada, Japan, and Australia.[138] According to the International Resource Panel's Metal Stocks in Society report of 2010, the global per capita stock of lead in use in society is 8 kg. Much of this is in more developed countries (20–150 kg per capita) rather than less developed countries (1–4 kg per capita).[139]

Production processes for primary and secondary lead are similar. Some primary production plants now also use scrap lead, and this trend is likely to increase in the future. Given adequate techniques, secondary lead is indistinguishable from primary lead. Scrap lead from the building trade is usually fairly clean and is re-melted without the need for smelting, though refining is sometimes needed. Secondary lead production is therefore cheaper, in terms of energy requirements, than is primary production, often by 50% or more.[140]

Primary[edit]

Most lead ores contain only a very low percentage of lead, which must be concentrated during processing.[141] During initial processing, ores typically undergo crushing, dense-medium separation, grinding, froth flotation, and drying. The resulting concentrate, which has a lead content fraction of 30–80%, [142] is then turned into (impure) lead metal. The main route for doing so involves a two-stage process. First, the sulfide concentrate is roasted in the air to oxidize the lead sulfide:[143]

2 PbS + 3 O2 → 2 PbO + 2 SO2
World's largest mining countries of lead, 2015[77]
Country Output
(thousand
tons)
 China 2,300
 Australia 633
 United States 385
 Peru 300
 Mexico 240
 India 130
 Russia 90
 Bolivia 82
 Sweden 76
 Turkey 54
 North Korea 45
 Poland 40
 South Africa 40
 Kazakhstan 38
 Ireland 33
Other countries 226

As the original concentrate was not pure lead sulfide, roasting yields lead oxide and a mixture of sulfates and silicates of lead and other metals contained in the ore.[142] This impure lead oxide is reduced in a coke-fired blast furnace to the (again, impure) metal: [144][145]

2 PbO + C → Pb + CO2

Research on a cleaner less energy intensive process continues, with some success; a major drawback is that the alternatives result in either an exceedingly high sulfur content in the resulting lead metal, or too much lead is lost as waste. A promising alternative involves direct smelting without an intermediate compound involved; hydrometallurgical extraction, in which anodes of impure lead and cathodes of pure lead are dissolved in an electrolyte, is another technique that is being explored.[146]

Impurities in the resulting metal are still significant; these are mostly arsenic, antimony, bismuth, zinc, copper, silver, and gold. The melt is treated in a reverberatory furnace with air, steam, and sulfur, which oxidizes the impurities except for silver, gold, and bismuth. Oxidized contaminants are removed by drossing, where they float to the top and are skimmed off.[147][148] Metallic silver and gold are removed and recovered economically by means of the Parkes process, in which zinc is added to lead. The zinc adsorbs silver and gold both of which, being immiscible in lead, can be separated and retrieved.[46][148] De-silvered lead is freed of bismuth according to the Betterton–Kroll process by treating it with metallic calcium and magnesium. The resulting bismuth dross can be skimmed off.[148]

Very pure lead can be obtained by processing smelted lead electrolytically using the Betts process. Anodes of impure lead and cathodes of pure lead are placed in an electrolyte of silica fluoride. Once electrical potential is applied, impure lead at the anode dissolves and plates out on the cathode leaving the impurities remain in solution.[148][149] This technique could potentially be applied to the original concentrate but doing so would be too costly despite attempts to make it cheaper; it is only currently used for refining lead.[146]

Secondary[edit]

Smelting, which is an essential part of the primary production, is often skipped during secondary production. The reason for this is that scrap lead itself is commonly reduced to its metallic form. Smelting is only performed when metallic lead had undergone significant chemical transformation, such as oxidation or rusting.[140] When smelting is performed, the process is similar to that of the primary production in either a blast furnace or a rotary furnace (with the essential difference being the greater variability of possible yields from the primary process). The Isasmelt process is a more recent method that may act as an extension to primary production; the essence of this process is that battery paste from spent lead-acid batteries is deprived of its sulfur content (by, for example, treating it with alkalis) and then treated in a coal-fueled furnace in the presence of oxygen, which eventually yields impure lead, with antimony being the most common impurity.[150] Refining of secondary lead is similar to that of primary lead; some refining processes may be skipped depending on the material recycled and its potential contamination, with bismuth and silver most commonly being accepted as impurities.[150]

Of the sources of lead for recycling, lead–acid batteries are most important; lead pipe, sheet, and cable sheathing are other significant sources.[140]

Applications[edit]

Bricks of lead (alloyed with 4% antimony) are used as radiation shielding.[151]

Contrary to popular belief, pencil leads in wooden pencils have never been made from lead. When the pencil originated as a wrapped graphite writing tool, the particular type of graphite used was named plumbago (literally, act for lead or lead mockup).[152]

Elemental form[edit]

Lead metal has several useful mechanical properties, including high density, low melting point, ductility, and relative inertness. Many metals are superior to lead in some of these aspects but lead is more common than most of these metals, and lead-bearing minerals are easier to mine and process than those of many other metals. One disadvantage of using lead is its toxicity, which explains why it has been or is being phased out for some uses.

Lead has been used for bullets since their invention (see above); with the development of firearms, round bullets became pointed and later, lead was jacketed with, for example, copper.[153] Concerns have been raised over whether lead bullets used for hunting can damage the environment.[o]

Because of its high density and resistance to corrosion, lead is used as ballast in sailboat keels.[155] Its high density allows it to counterbalance the heeling effect of wind on the sails at the same time occupying a small volume and thus minimizing underwater resistance. Similarly, lead is used in scuba diving weight belts to counteract the diver's buoyancy.[156] In 1993, 600 tonnes of lead were used to stabilize the base of the Leaning Tower of Pisa.[157] Given its corrosion resistance, lead is used as a protective sheath for (seabed) submarine cables.[158]

Lead is added to copper alloys such as brass and bronze, to improve machinability and for its lubricating qualities. Being practically insoluble in copper the lead forms solid globules permeated throughout imperfections within the alloy, such as grain boundaries. In low concentrations, as well as acting as lubricants, these globules hinder the formation of large chips as the alloy is worked, thereby improving machinability. Copper alloys with larger concentrations of lead are used in bearings. The lead provides lubrication; the copper provides the load bearing support.[159]

Lead is used to form glazing bars for stained glass or other multi-lit windows. The practice has become less common, not due to concerns about lead toxicity but for stylistic reasons. Sheet-lead is used as a sound deadening layer in some areas in wall, floor and ceiling design in sound studios.[160][161] It is the traditional base metal of organ pipes, mixed with various amounts of tin to control the tone of the pipe.[162][163]

Lead has many uses in the construction industry (for example, lead sheets are used as architectural metals in roofing material, cladding, flashing, gutters and gutter joints, and on roof parapets).[164][165] Detailed lead moldings are used as decorative motifs to fix lead sheet. Lead is still used in statues and sculptures.[166] It is often used to balance the wheels of a car; for environmental reasons this use is being phased out in favour of other materials.[167]

Multicolor lead-glazing in a Tang dynasty Chinese sancai ceramic cup dating from the 8th century CE

Apart from its mechanical properties, lead is useful in lead–acid batteries. The reactions in the battery between lead, lead dioxide, and sulfuric acid provide a reliable source of voltage.[p] This has been the largest use of lead in early 21st century since the lead in batteries undergoes no direct contact with humans (and thus there are no immediate toxicity concerns).

Lead is used in electrodes for the process of electrolysis. Its use in solder for electronics is being phased out by some countries to reduce the amount of environmentally hazardous waste, and in high voltage power cables as sheathing material to prevent water diffusion into insulation. Lead is one of three metals used in the Oddy test for museum materials, helping detect organic acids, aldehydes, and acidic gases. It is also used as shielding from radiation (in X-ray rooms, for example).[169] Molten lead is used as a coolant for lead cooled fast reactors).[170]

Compounds[edit]

Lead compounds are used as, or in, coloring agents, oxidants, plastic, candles, glass, and semiconductors. Lead-based coloring agents are used in ceramic glazes, notably for red and yellow shades.[171] Lead tetraacetate (LTA) and lead dioxide have been used as oxidizing agents in organic chemistry. Lead is frequently used in polyvinyl chloride (PVC) plastic, which coats electrical cords.[172][173] Lead is used to treat some candle wicks to ensure a longer, more even burn. Because of its toxicity, European and North American manufacturers use alternatives such as zinc.[174][175] Lead glass is composed of 12–28% lead oxide. It changes the optical characteristics of the glass and reduces the transmission of ionizing radiation.[176] Lead-based semiconductors, such as lead telluride, lead selenide and lead antimonide are finding applications in photovoltaic (solar energy) cells and infrared detectors.[177]

Biological and environmental effects[edit]

Main article: Lead poisoning

Biological[edit]

Symptoms of lead poisoning

Along with such elements as cadmium and mercury,[178] lead has no biological role.[179] Despite this, it is relatively abundant in the human body, being one of the most abundant heavy mentals that have no such role. Lead in human bodies comes from food, drinking water, and the environment. The rates vary greatly by country.[180]

Lead is considered a highly poisonous metal (whether inhaled or swallowed), affecting almost every organ and system in the body.[181] The main target for lead toxicity in humans is the central nervous system. By mimicking calcium, lead can cross the blood-brain barrier. It subsequently degrades the myelin sheaths of neurons, reduces their numbers, interferes with neurotransmission routes, and decreases neuronal growth.[182] In a child's developing brain, lead interferes with synapse formation in the cerebral cortex, neurochemical development (including that of neurotransmitters), and the organization of ion channels.[183]

The primary cause of lead's toxicity is its predilection for interfering with the proper functioning of enzymes. It does so by binding to the sulfhydryl groups found on many enzymes,[182] or mimicking and displacing other metals which act as cofactors in many enzymatic reactions.[184] Lead salts are very quickly and efficiently absorbed by the body, accumulating in it and leading to both chronic and acute poisoning.[185]

Among the essential metals with which lead interacts are calcium, iron, and zinc.[186] High levels of calcium and iron tend to provide some protections from lead poisoning; low levels confer increased susceptibility.[185]

A small amount of ingested lead (1%) will be stored in bones, and the rest will be excreted by an adult through urine and feces within a few weeks of exposure. Only about a third of lead will be excreted by a child.[187]

Chronic exposure to lead or its salts (especially soluble salts or the strong oxidant PbO2) in adults can result in decreased performance in some tests that measure nervous system functions.[188] Symptoms include nephropathy, and colic-like abdominal pains and possibly weakness in the fingers, wrists, or ankles. Lead exposure also causes small increases in blood pressure, particularly in middle-aged and older people and can cause anemia. Exposure to high lead levels can cause severe damage to the brain and kidneys in adults or children and ultimately cause death. In pregnant women, high levels of exposure to lead may cause miscarriage. Chronic, high-level exposure has been shown to reduce fertility in males.[189]

Lead also damages nervous connections (especially in young children) and causes blood and brain disorders. Lead poisoning nowadays typically results from ingestion of food or water contaminated with lead, but may also occur after accidental ingestion of contaminated soil, dust, or lead-based paint.[190] It is rapidly absorbed into the bloodstream and is believed to have adverse effects on the central nervous system, the cardiovascular system, kidneys, and the immune system.[191] In the 20th century the air was commonly contaminated with lead (in the form of TEL from gasoline) and various observations have led to the hypothesis of a link between lead and crime levels (which is not universally accepted).[192]

NFPA 704
"fire diamond"
Flammability code 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g., canola oil Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
Fire diamond for lead granules

Treatment for lead poisoning normally involves the administration of dimercaprol and succimer.[193] Acute cases may require the use of disodium calcium edetate, this being the calcium chelate of the disodium salt of ethylenediaminetetraacetic acid (EDTA). It has a greater affinity for lead than calcium with the result that lead chelate is formed by exchange and excreted in the urine, leaving behind harmless calcium.[194]

The role of extremely low levels of lead in causing permanent cognitive deficits in children has brought about a widespread reduction in its use.[195][196] Early childhood exposure has further been linked with an increased risk of sleep disturbances and excessive daytime sleepiness in later childhood.[197] High blood levels are associated with delayed puberty in girls.[198]

Despite the toxicity of lead in significant amounts, there is some evidence that trace amounts are beneficial in pigs and rats, and that its absence causes deficiencies such as depressed growth, anemia, and disturbed iron metabolism. If this finding holds for humans it would make lead an essential element, one with a threshold of toxicity so low that lead toxicity would remain a much higher priority than lead deficiency.[199][200][201]

Exposure sources[edit]

Battery collection site in Dakar, Senegal, where at least 18 children died of lead poisoning in 2008

Lead can be ingested through fruits and vegetables contaminated by high levels of lead in the soils they were grown in. Soil can be contaminated through particulate accumulation from lead in pipes, lead paint, and residual emissions from leaded gasoline (before use of the latter was generally phased out).[202] The use of lead for water pipes is problematic in areas with soft or (and) acidic water. Hard water forms insoluble layers in the pipes whereas soft and acidic water dissolves the lead pipes.[203] Ingesting certain home remedy medicines may result in exposure to lead or lead compounds.[204] Ingestion of lead-based paint is the major source of exposure for children. As the paint deteriorates, it peels, is pulverized into dust and then enters the body through hand-to-mouth contact or contaminated food, water, or alcohol.

Inhalation is the second major exposure pathway, especially for workers in lead-related occupations. Most cases of adult elevated blood lead levels are workplace-related.[205] Almost all inhaled lead is absorbed into the body; for ingestion, the rate is 20–70%. with children absorbing lead at a higher rate than adults.[204]

Dermal exposure may be significant for a narrow category of people working with organic lead compounds. The rate of skin absorption is low for inorganic lead.[204]

Environmental[edit]

The extraction, production, use, and disposal of lead and its products have caused significant contamination of the Earth's soils and waters, posing a hazard to living organisms because of its toxicity. Atmospheric emissions of lead were at their peak during the Industrial Revolution and the period of leaded petrol in the second half of the twentieth century. These periods are over but elevated concentrations of lead persist in soils and sediments in post-industrial and urban areas, and industrial emissions continue in many parts of the world.[206]

Radiography of a swan found dead in Condé-sur-l'Escaut (northern France), highlighting lead shot. The amount of lead is exceptionally high (some hundreds of pellets; a dozen is enough to kill an adult swan within a few days). Such bodies are sources of environmental contamination by lead.

Lead accumulates in soils, especially those with a high organic content, where it remains for a long time (hundreds and thousands of years). It can take the place of other metals within plants and can accumulate on their surfaces, thereby retarding photosynthesis, and preventing the growth of the plant or killing it. Contamination of soils and plants, in turn, affects microorganisms and animals. Affected animals have a reduced ability to synthesize red blood cells. Sources of lead contamination are therefore being curtailed.[207][q]

Research has been conducted on how to remove lead from biosystems via biological organisms. Fish bones are being researched for their ability to bioremediate lead in contaminated soil.[209][210] The fungus Aspergillus versicolor is particularly effective at removing lead ions.[211] Several bacteria have been researched for their ability to reduce lead including the sulfate reducing bacteria Desulfovibrio and Desulfotomaculum, both of which are highly effective in aqueous solutions.[212]

Restriction of lead usage[edit]

During the 20th century, the use of lead in paint pigments was sharply curtailed because of the danger of lead poisoning, especially to children.[213] By the mid-1980s, a significant shift in lead end-use patterns had taken place. Much of this shift was a result of compliance, in the United States, with environmental regulations that significantly reduced or eliminated the use of lead in non-battery products, including gasoline, paints, solders, and water systems. Lead use is being further curtailed by the European Union's Restriction of Hazardous Substances Directive.[214] Lead may be found in harmful quantities in stoneware,[215] vinyl[216] (such as that used for tubing and the insulation of electrical cords), and Chinese brass. Old houses may contain substantial amounts of lead paint.[216] White lead paint has been withdrawn from sale in industrialized countries, but yellow lead chromate is still in use. Old paint should not be stripped by sanding, as this produces inhalable dust.[217]

In the United States, the Occupational Safety and Health Administration has set the permissible exposure limit for lead exposure in the workplace as 0.05 mg/m3 over an 8-hour workday, which applies to metallic lead, inorganic lead compounds, and lead soaps. The National Institute for Occupational Safety and Health has set a recommended exposure limit of 0.05 mg/m3 over an 8-hour workday and recommends that workers' blood concentrations of lead stay below 0.06 mg per 100 g blood. At levels of 100 mg/m3, lead is immediately dangerous to life and health.[218]

See also[edit]

Notes[edit]

  1. ^ About 10% of the lanthanide contraction has also been attributed to relativistic effects.[5]
  2. ^ In the same year, it was reported that lead formed a nanoscale thin-film, quasicrystalline allotrope, with pentagonal symmetry, when deposited on the surface of an icosahedral silver-indium-ytterbium quasicrystal. The electronic nature of the allotrope—whether it was metallic or an insulator (or something in-between)—was not recorded.[9][10]
  3. ^ The difference between the two terms is that malleability refers to deformability under compression (i.e., pressing a tablet of a material into a sheet) whereas ductility refers to its ability to stretch (i.e., elongating a rod of a material into a wire).
  4. ^ A (wet) finger can be dipped into molten lead without risk of a burning injury.[22]
  5. ^ An even number of either protons or neutrons generally increases the nuclear stability of isotopes, compared to isotopes with odd numbers. For example, elements with odd atomic numbers have no more than two stable isotopes, while even-numbered elements have multiple stable isotopes, with tin (element 50) having the highest number of isotopes of all elements, ten.[26] See Even and odd atomic nuclei for more details.
  6. ^ Lead 208 is doubly magic, and especially stable against decay, as its 126 neutrons also form a complete shell.
  7. ^ The half-life found in the experiment was 1.9×1019 years.[27] A kilogram of natural bismuth would thus be radioactive with an activity value of approximately 0.003 becquerels (decays per second). For comparison, the activity value of natural radiation within the human body is around 65 becquerels per kilogram of body weight (4500 becquerels on average).[28]
  8. ^ These predicted half-lives are as follows:[29]
    204Pb: 2.3×1035–1.2×1037 y
    206Pb: 1.8×1065–6.7×1068 y
    207Pb: 3.6×10152–3.4×10189 y
    208Pb: 1.2×10124–7.4×10132 y
  9. ^ It decays solely via electron capture, which means when there are no electrons available and lead is accordingly fully ionized—has all 82 electrons removed—it cannot decay and becomes stable. Fully ionized thallium-205, the isotope lead-205 would decay to, becomes unstable with respect to decaying into a bound state of lead-205.[36]
  10. ^ Abundances in the source are listed relative to silicon rather than in per-particle notation. The sum of all elements per each 106 parts of silicon is 2.6682×1010 parts; lead alone comprises 3.258 parts.
  11. ^ The latter word is related to the Latin plumbum, which gave the element its chemical symbol Pb. It may also be the origin of the Proto-Germanic *bliwa- (which also means "lead"), from which stemmed the German Blei.[82]
  12. ^ The inscription reads: "Made when the Emperor Vespasian was consul for the ninth term and the Emperor Titus was consul for the seventh term, when Gnaeus Iulius Agricola was imperial governor (of Britain)."
  13. ^ The fact that Julius Caesar fathered only one child, as well as the alleged sterility of his successor, Caesar Augustus, have been attributed to lead poisoning.[104]
  14. ^ It is not known when mining was first performed in the region because no tradition of keeping written records was in place, but there are European 17th century records of trade with the Congolese, which indicates lead was first smelted no later than then.[125]
  15. ^ For instance, the U.S. state of California banned lead bullets for hunting on that basis in April 2015.[154]
  16. ^ See [168] for details on how a lead–acid battery works.
  17. ^ For example, in the Netherlands, the use of lead shot for hunting and sport shooting was banned in 1993, which caused a large drop in lead emission, from 230 tonnes in 1990 to 47.5 tonnes in 1995, two years after the ban.[208]

References[edit]

  1. ^ CIAAW (2015). "Standard Atomic Weights". Commission on Isotopic Abundances and Atomic Weights (CIAAW). Retrieved 18 February 2017. 
  2. ^ Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. p. E110. ISBN 0-8493-0464-4. 
  3. ^ Lide 2004, p. 10-179.
  4. ^ a b c d Polyanskiy 1986, pp. 14–15.
  5. ^ Pyykko, P. (1988). "Relativistic effects in structural chemistry". Chemical Reviews. 88 (3): 563–594. doi:10.1021/cr00085a006. 
  6. ^ Norman, N. C. (1996). Periodicity and the s- and p-Block Elements. Oxford: Oxford University Press. p. 36. ISBN 978-0-19-855961-0. 
  7. ^ Greenwood & Earnshaw 1998, p. 227.
  8. ^ Christensen, N. E. (2002). "Relativistic Solid State Theory". In Schwerdtfeger, P. Relativistic Electronic Structure Theory - Fundamentals. Elsevier. pp. 867–868. ISBN 978-0-08-054046-7. 
  9. ^ H. R, Sharma; K., Nozawa; Smerdon, J. A.; et al. (2013). "Templated three-dimensional growth of quasicrystalline lead". Nature Communications. 4. doi:10.1038/ncomms3715. 
  10. ^ Sharma, H. R.; Smerdon, J. A.; Nugent, P. J.; et al. (2014). "Crystalline and quasicrystalline allotropes of Pb formed on the fivefold surface of icosahedral Ag-In-Yb". The Journal of Chemical Physics. 140: 174710. doi:10.1063/1.4873596. 
  11. ^ a b Polyanskiy 1986, p. 18.
  12. ^ Thornton, Radu & Brush 2001, p. 6.
  13. ^ Lide 2004, pp. 12-35–12-37.
  14. ^ Jones, P. A. (2014). "Lead balloon". Jedburgh Justice and Kentish Fire: The Origins of English in Ten Phrases and Expressions. London: Constable. p. 42. ISBN 978-1-4721-1622-2. 
  15. ^ Lide 2004, pp. 4-39–4-96.
  16. ^ Vogel, N. A.; Achilles, R. (2013). The Preservation and Repair of Historic Stained and Leaded Glass (PDF) (Report). United States Department of the Interior. p. 8. Retrieved 30 October 2016. 
  17. ^ Anderson, J. (1869). "Malleability and ductility of metals". Scientific American. 21 (22): 341–343. doi:10.1038/scientificamerican11271869-341. 
  18. ^ Gale, W. F.; Totemeier, T. C. (2003). Smithells Metals Reference Book. Butterworth-Heinemann. pp. 15–2–15–3. ISBN 978-0-08-048096-1. 
  19. ^ Thornton, Radu & Brush 2001, p. 8.
  20. ^ Lide 2004, p. 12-220.
  21. ^ Koshal, D. (2014). Manufacturing Engineer's Reference Book. Butterworth-Heinemann. p. 1/92. ISBN 978-0-08-052395-8. 
  22. ^ "The physics behind four amazing demonstrations - CSI". Skeptical Inquirer. 23 (6). 1999. Retrieved 6 September 2016. 
  23. ^ Lide 2004, p. 12-219.
  24. ^ Blakemore, J. S. (1985). Solid State Physics. Cambridge University Press. p. 272. ISBN 978-0-521-31391-9. 
  25. ^ Polyanskiy 1986, p. 16.
  26. ^ a b c d e Audi, G.; Wapstra, A. H.; Thibault, C.; et al. (2003). "The NUBASE evaluation of nuclear and decay properties" (PDF). Nuclear Physics A. 729 (1): 3–128. Bibcode:2003NuPhA.729....3A. doi:10.1016/j.nuclphysa.2003.11.001. Archived from the original (PDF) on 24 July 2013. 
  27. ^ Marcillac, P. de; Coron, N.; Dambier, G.; et al. (2003). "Experimental detection of α-particles from the radioactive decay of natural bismuth". Nature. 422 (6934): 876–878. Bibcode:2003Natur.422..876D. doi:10.1038/nature01541. PMID 12712201. 
  28. ^ "Nuclear Radiation and Health Effects". World Nuclear Association. 2015. Retrieved 12 November 2015. 
  29. ^ a b Beeman, J. W.; Bellini, F.; Cardani, L.; et al. (2013). "New experimental limits on the α decays of lead isotopes". The European Physical Journal A. 49 (50). Retrieved 21 August 2016. 
  30. ^ Smirnov, A. Yu.; Borisevich, V. D.; Sulaberidze, A. (2012). "Evaluation of specific cost of obtainment of lead-208 isotope by gas centrifuges using various raw materials". Theoretical Foundations of Chemical Engineering. 46 (4): 373–378. doi:10.1134/s0040579512040161. 
  31. ^ Greenwood & Earnshaw 1998, p. 368.
  32. ^ Boltwood, B. B. (1907). "On the ultimate disintegration products of the radio-active elements. Part II. The disintegration products of uranium". American Journal of Science. 23: 77–88. doi:10.2475/ajs.s4-23.134.78. 
  33. ^ University of California Berkeley Nuclear Forensic Search Project. "Decay Chains". Nuclear Forensics: A Scientific Search Problem. Retrieved 23 November 2015. 
  34. ^ Fiorini, E. (June 2010). "2.000 years-old Roman Lead for physics" (PDF). ASPERA: 7–8. Retrieved 29 October 2016. 
  35. ^ Nosengo, N. (2010). "Roman ingots to shield particle detector". Nature News. doi:10.1038/news.2010.186. 
  36. ^ Takahashi, K.; Boyd, R. N.; Mathews, G. J.; et al. (1987). "Bound-state beta decay of highly ionized atoms" (PDF). Physical Review C. American Institute of Physics for the American Physical Society. 36 (4). ISSN 0556-2813. OCLC 1639677. Archived from the original (PDF) on 21 October 2014. Retrieved 27 August 2013. 
  37. ^ Greenwood & Earnshaw 1998, p. 373.
  38. ^ Kaupp, M. (2014). "Chemical bonding of main-group elements". In Frenking, G.; Shaik, S. The Chemical Bond: Chemical Bonding Across the Periodic Table (PDF). John Wiley & Sons. pp. 9–10. doi:10.1002/9783527664658.ch1. 
  39. ^ Rappoport, Zvi; Marek, Ilan (2010). The Chemistry of Organocopper Compounds. John Wiley & Sons. p. 509. ISBN 978-0-470-77296-6. 
  40. ^ a b c Greenwood & Earnshaw 1998, p. 389.
  41. ^ Percy, J. (1870). The Metallurgy of Lead, including Desiverization and Cupellation. Murray, J. p. 67. 
  42. ^ Ensafi, A. A.; Far, A. K.; Meghdadi, S. (2009). "Highly selective optical-sensing film for lead(II) determination in water samples". Journal of Hazardous Materials. 172 (2–3): 1069–75. doi:10.1016/j.jhazmat.2009.07.112. PMID 19709813. 
  43. ^ a b c King, R. B. (1995). Inorganic Chemistry of Main Group Elements. Wiley-VCH. pp. 43–63. ISBN 978-0-471-18602-1. 
  44. ^ Greenwood & Earnshaw 1998, p. 384.
  45. ^ Lewis, A. E. (2010). "Review of metal sulphide precipitation" (PDF). Hydrometallurgy. 104 (2): 222–234. doi:10.1016/j.hydromet.2010.06.010. Retrieved 14 October 2014. 
  46. ^ a b Pauling, L. (1947). General Chemistry. W. H. Freeman and Company. ISBN 978-0-486-65622-9. 
  47. ^ 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. 
  48. ^ a b c Greenwood & Earnshaw 1998, p. 382.
  49. ^ Greenwood & Earnshaw 1998, p. 388.
  50. ^ Cava, R. J.; Hor, Y. S.; Cava, R. J. (2011). "Pressure stabilized Se–Se dimer formation in PbSe2". Solid State Sciences. 13: 38–41. Bibcode:2011SSSci..13...38B. doi:10.1016/j.solidstatesciences.2010.10.003. 
  51. ^ Silverman, M. S. (1966). "High-pressure (70-kilobar) synthesis of new crystalline lead dichalcogenides". Inorganic Chemistry. 5 (11): 2067–2069. doi:10.1021/ic50045a056. 
  52. ^ Greenwood & Earnshaw 1998, p. 398.
  53. ^ Macomber, R. S. (1996). Organic Chemistry. University Science Books. p. 230. ISBN 978-0-935702-90-3. 
  54. ^ Yong, L.; Hoffmann, S. D.; Fässler, T. F. (2006). "A low-dimensional arrangement of [Pb9]4− clusters in [K(18-crown-6)]2K2Pb9·(en)1.5". Inorganica Chimica Acta. Elsevier. 359 (15): 4774–4778. doi:10.1016/j.ica.2006.04.017. 
  55. ^ Becker, M.; Förster, C.; Franzen, C.; et al. (2008). "Persistent radicals of trivalent tin and lead". Inorganic Chemistry. 47 (21): 9965–9978. doi:10.1021/ic801198p. PMID 18823115. 
  56. ^ Mosseri, S.; Henglein, A.; Janata, E. (1990). "Trivalent lead as an intermediate in the oxidation of lead(II) and the reduction of lead(IV) species". The Journal of Physical Chemistry. 94 (6): 2722–2726. doi:10.1021/j100369a089. 
  57. ^ Chia, S.-P.; X., H.-W.; Li, Y.; et al. (2013). "A base-stabilized lead(I) dimer and an aromatic plumbylidenide anion". Angewandte Chemie International Edition. 52 (24): 6298–6301. doi:10.1002/anie.201301954. 
  58. ^ Greenwood & Earnshaw 1998, p. 386.
  59. ^ Universität Freiburg. "Binäre Zintl-Phasen" [Binary Zintl Phases] (in German). Retrieved 18 February 2017. 
  60. ^ Alsfasser, R. (2007). Moderne anorganische Chemie [Modern inorganic chemistry] (in German). Walter de Gruyter. pp. 261–263. ISBN 978-3-11-019060-1. 
  61. ^ Greenwood & Earnshaw 1998, p. 393.
  62. ^ Greenwood & Earnshaw 1998, p. 374.
  63. ^ Stabenow, F.; Saak, W.; Weidenbruch, M. (2003). "Tris(triphenylplumbyl)plumbate: An anion with three stretched lead–lead bonds". Chemical Communications (18): 2342. doi:10.1039/B305217F. 
  64. ^ a b c Polyanskiy 1986, p. 43.
  65. ^ a b c d Greenwood & Earnshaw 1998, p. 404.
  66. ^ a b Wiberg, E.; Wiberg, N.; Holleman, A. F. (2001). Inorganic Chemistry. Academic Press. p. 918. ISBN 978-0-12-352651-9. 
  67. ^ Polyanskiy 1986, p. 44.
  68. ^ Windholz, M. (1976). Merck Index of Chemicals and Drugs (9th ed.). Merck & Co. ISBN 978-0-911910-26-1. Monograph 8393. 
  69. ^ Zýka, J. (1966). "Analytical study of the basic properties of lead tetraacetate as oxidizing agent" (PDF). Pure and Applied Chemistry. 13 (4): 569–581. doi:10.1351/pac196613040569. Retrieved 19 December 2013. 
  70. ^ Greenwood & Earnshaw 1998, p. 405.
  71. ^ a b c d Lodders, K. (2003). "Solar System abundances and condensation temperatures of the elements". The Astrophysical Journal. 591 (2): 1220–1247. doi:10.1086/375492. ISSN 0004-637X. 
  72. ^ Roederer, I. U.; Kratz, K.-L.; Frebel, A.; et al. (2009). "The end of nucleosynthesis: Production of lead and thorium in the early galaxy". The Astrophysical Journal. The American Astronomical Society. 698 (2): 1963–1980. Bibcode:2009ApJ...698.1963R. doi:10.1088/0004-637X/698/2/1963. Retrieved 18 July 2016. 
  73. ^ a b Burbidge, E. M.; Burbidge, G. R.; Fowler, W. A.; et al. (1957). "Synthesis of the Elements in Stars" (PDF). Reviews of Modern Physics. 29 (4): 547. Bibcode:1957RvMP...29..547B. doi:10.1103/RevModPhys.29.547. 
  74. ^ Frebel, A. (2015). Searching for the Oldest Stars: Ancient Relics from the Early Universe. Princeton University. pp. 114–115. ISBN 978-0-691-16506-6. 
  75. ^ "Lead - Pb". environmentalchemistry.com. Retrieved 22 February 2017. 
  76. ^ a b c Sutherland et al. 2005, p. 5.
  77. ^ a b c United States Geological Survey (2016). Lead (PDF) (Report). Retrieved 20 February 2016. 
  78. ^ Rieuwerts, J. (2015). The Elements of Environmental Pollution. Routledge. p. 225. ISBN 978-0-415-85919-6. 
  79. ^ Merriam-Webster. "Definition of LEAD". www.merriam-webster.com. Retrieved 12 August 2016. 
  80. ^ Kroonen 2013, *lauda-.
  81. ^ Nikolayev, Sergei, ed. (2012). "*lAudh-". Indo-European Etymology. starling.rinet.ru. Retrieved 21 August 2016. 
  82. ^ Kroonen 2013, *bliwa- 2.
  83. ^ Kroonen 2013, *laidijan-.
  84. ^ a b c d Hong, S.; Candelone, J.-P.; Patterson, C. C.; et al. (1994). "Greenland ice evidence of hemispheric lead pollution two millennia ago by Greek and Roman civilizations" (PDF). Science. 265 (5180): 1841–1843. Bibcode:1994Sci...265.1841H. doi:10.1126/science.265.5180.1841. PMID 17797222. 
  85. ^ a b Rich 1994, p. 4.
  86. ^ a b c d e f Winder, C. (1993). "The history of lead - Part 3". LEAD Action News. 2 (3). Archived from the original on 31 August 2007. Retrieved 12 February 2016. 
  87. ^ a b Rich 1994, p. 5.
  88. ^ a b "A History of Cosmetics from Ancient Times | Cosmetics Info". www.cosmeticsinfo.org. Retrieved 18 July 2016. 
  89. ^ Yu, L.; Yu, H. (2004). Chinese Coins: Money in History and Society. Long River Press. p. 26. ISBN 978-1-59265-017-0. 
  90. ^ "Toronto museum explores history of contraceptives". ABC News. 2003. Retrieved 13 February 2016. 
  91. ^ Bisson & Vogel 2000, p. 105.
  92. ^ Eschnauer, H. R.; Stoeppler, M. (1992). "Wine—An enological specimen bank". In Stoeppler, M. Hazardous Materials in the Environment. Elsevier Science. pp. 49–72 (58). ISBN 0-444-89078-5. 
  93. ^ Sutherland et al. 2005, p. 2.
  94. ^ Callataÿ, F. de (2005). "The Graeco-Roman economy in the super long-run: Lead, copper, and shipwrecks". Journal of Roman Archaeology. 18: 361–372. doi:10.1017/S104775940000742X. 
  95. ^ Settle, D. M.; Patterson, C. C. (1980). "Lead in Albacore: Guide to Lead Pollution in Americans". Science. 207 (4436): 1167–1176. Bibcode:1980Sci...207.1167S. doi:10.1126/science.6986654. PMID 6986654.  see 1170f.
  96. ^ Rich 1994, p. 6.
  97. ^ Thornton, Radu & Brush 2001, pp. 179–184.
  98. ^ Bisel, S. C.; Bisel, J. F. (2002). "Health and nutrition at Herculaneum". In Jashemski, W. F.; Meyer, F. G. The Natural History of Pompeii. Cambridge University. pp. 451–475 (460). ISBN 978-0-521-80054-9. 
  99. ^ Retief, F.; Cilliers, L. P. (2006). "Lead poisoning in ancient Rome". Acta Theologica. 26 (2): 147–164 (149–151). doi:10.4314/actat.v26i2.52570. 
  100. ^ Grout, J. (2017). "Lead poisoning and Rome". Encyclopaedia Romana. Retrieved 15 February 2017. 
  101. ^ Hodge, T. A. (1981). "Vitruvius, lead pipes and lead poisoning". American Journal of Archaeology. Archaeological Institute of America. 85 (4): 486–491. doi:10.2307/504874. JSTOR 504874. 
  102. ^ Gilfillan, S. C. (1965). "Lead poisoning and the fall of Rome". Journal of Occupational Medicine. 7 (2): 53–60. ISSN 0096-1736. PMID 14261844. 
  103. ^ Nriagu, J. O. (1983). "Saturnine gout among Roman aristocrats". New England Journal of Medicine. 308 (11): 660–663. doi:10.1056/NEJM198303173081123. ISSN 0028-4793. PMID 6338384. 
  104. ^ Frankenburg, F. R. (2014). Brain-Robbers: How Alcohol, Cocaine, Nicotine, and Opiates Have Changed Human History. ABC-CLIO. p. 16. ISBN 978-1-4408-2932-1. 
  105. ^ Scarborough, J. (1984). "The myth of lead poisoning among the Romans: An essay review". Journal of the History of Medicine and Allied Sciences. 39 (4): 469–475. doi:10.1093/jhmas/39.4.469. ISSN 0022-5045. 
  106. ^ Waldron, H. A. (1985). "Lead and lead poisoning in antiquity". Medical History. 29 (1): 107–108. ISSN 0025-7273. PMC 1139494Freely accessible. 
  107. ^ Reddy, A.; Braun, C. L. (2010). "Lead and the Romans". Journal of Chemical Education. 87 (10): 1052–1055. doi:10.1021/ed100631y. ISSN 0021-9584. 
  108. ^ Delile, H.; Blichert-Toft, J.; Goiran, J.-P.; et al. (2014). "Lead in ancient Rome's city waters". Proceedings of the National Academy of Sciences. 111 (18): 6594–6599. doi:10.1073/pnas.1400097111. ISSN 0027-8424. PMC 4020092Freely accessible. PMID 24753588. 
  109. ^ Finger, S. (2006). Doctor Franklin's Medicine. Philadelphia: University of Pennsylvania Press. p. 184. ISBN 978-0-8122-3913-3. 
  110. ^ Lewis, J. "Lead Poisoning: A Historical Perspective". USEPA. Retrieved 31 January 2017. 
  111. ^ Polyanskiy 1986, p. 8.
  112. ^ Thomson, T. (1830). The History of Chemistry. Henry Colburn and Richard Bentley (publishers). p. 74. 
  113. ^ Kellett, C. (2012). Poison and Poisoning: A Compendium of Cases, Catastrophes and Crimes. Accent Press. pp. 106–107. ISBN 978-1-909335-05-9. 
  114. ^ Squatriti, P., ed. (2000). Working with Water in Medieval Europe: Technology and Resource Use. Brill. pp. 134 ff. ISBN 978-90-04-10680-2. 
  115. ^ a b Winder, C. (1993). "The history Of lead part 1". LEAD Action News. The LEAD Group. Archived from the original on 31 August 2007. Retrieved 5 February 2016. 
  116. ^ a b c Rich 1994, p. 7.
  117. ^ Sinha, S. P.; Shelly; Sharma, V.; et al. (1993). "Neurotoxic effects of lead exposure among printing press workers". Bulletin of Environmental Contamination and Toxicology. 51 (4). doi:10.1007/BF00192162. 
  118. ^ Ramage, C. K. (1980). Lyman Cast Bullet Handbook (3rd ed.). Lyman Publications. p. 8. 
  119. ^ Nakashima, T.; Matsuno, K.; Matsushita, T. (2007). "Lifestyle-determined gender and hierarchical differences in the lead contamination of bones from a feudal town of the Edo period". Journal of Occupational Health. 49 (2): 134–139. doi:10.1539/joh.49.134. PMID 17429171. 
  120. ^ Nakashima, T.; Hayashi, H.; Tashiro, H.; et al. (1998). "Gender and hierarchical differences in lead-contaminated Japanese bone from the Edo period". Journal of Occupational Health. 40: 55–60. doi:10.1539/joh.40.55. 
  121. ^ Ashikari, M. (2003). "The memory of the women's white faces: Japaneseness and the ideal image of women". Japan Forum. 15: 55–79. doi:10.1080/0955580032000077739. 
  122. ^ Beard, M. E. (1995). Lead in Paint, Soil, and Dust: Health Risks, Exposure Studies, Control Measures, Measurement Methods, and Quality Assurance. ASTM International. p. 66. ISBN 978-0-8031-1884-3. 
  123. ^ "Mining History". www.mininghistory.asn.au. Retrieved 15 February 2016. 
  124. ^ Bisson & Vogel 2000, p. 85.
  125. ^ a b Bisson & Vogel 2000, pp. 131–132.
  126. ^ "Lead mining". The Northern Echo. Retrieved 16 February 2016. 
  127. ^ Sohn, E. "Lead: Versatile Metal, Long Legacy". Dartmouth Toxic Metals Superfund Research Program. Retrieved 16 February 2016. 
  128. ^ Rich 1994, p. 11.
  129. ^ a b c d Riva, M. A.; Lafranconi, A.; d'Orso, M. I.; et al. (2012). "Lead poisoning: Historical aspects of a paradigmatic "occupational and environmental disease"". Safety and Health at Work. 3 (1): 11–16. doi:10.5491/SHAW.2012.3.1.11. PMC 3430923Freely accessible. PMID 22953225. 
  130. ^ Hernberg, S. (2000). "Lead poisoning in a historical perspective" (PDF). American Journal of Industrial Medicine. 38: 244–254. doi:10.1002/1097-0274(200009)38:3<244::AID-AJIM3>3.0.CO;2-F. PMID 10940962. 
  131. ^ "Why use lead in paint?". Chemistry World. Royal Society of Chemistry. Retrieved 22 February 2017. 
  132. ^ Markowitz, G.; Rosner, D. (2000). ""Cater to the children": the role of the lead industry in a public health tragedy, 1900–1955.". American Journal of Public Health. 90: 36–46. doi:10.2105/ajph.90.1.36. PMC 1446124Freely accessible. PMID 10630135. 
  133. ^ Rich 1994, p. 117.
  134. ^ Rich 1994, p. 17.
  135. ^ Rich 1994, pp. 91–92.
  136. ^ United States Geological Survey (2005). Lead (PDF) (Report). Retrieved 20 February 2016. 
  137. ^ Zhang, X.; Yang, L.; Li, Y.; et al. (2012). "Impacts of lead/zinc mining and smelting on the environment and human health in China". Environmental Monitoring and Assessment. 184 (4): 2261–2273. doi:10.1007/s10661-011-2115-6. ISSN 1573-2959. PMID 21573711. 
  138. ^ Guberman, D. E. (2015). "Lead". 2013 Minerals Yearbook (PDF) (Report). United States Geological Survey. Retrieved 2 November 2016. 
  139. ^ Graedel, T. E.; et al. (2010). "Metal stocks in Society – Scientific Synthesis" (PDF). International Resource Panel. Retrieved 2 July 2012. 
  140. ^ a b c Thornton, Radu & Brush 2001, p. 56.
  141. ^ Greenwood & Earnshaw 1998, p. 369.
  142. ^ a b Sutherland et al. 2005, pp. 7–19.
  143. ^ Thornton, Radu & Brush 2001, p. 51.
  144. ^ Thornton, Radu & Brush 2001, pp. 51–52.
  145. ^ "Primary Extraction of Lead Technical Notes". LDA International. Archived from the original on 22 March 2007. Retrieved 7 April 2007. 
  146. ^ a b Thornton, Radu & Brush 2001, pp. 52–53.
  147. ^ Sutherland et al. 2005, pp. 19–27.
  148. ^ a b c d "Primary Lead Refining Technical Notes". LDA International. Archived from the original on 22 March 2007. Retrieved 7 April 2007. 
  149. ^ Sutherland et al. 2005, p. 26.
  150. ^ a b Thornton, Radu & Brush 2001, p. 57.
  151. ^ Street & Alexander 1998, p. 181.
  152. ^ Evans, J. W. (1908). "V.— The meanings and synonyms of plumbago". Transactions of the Philological Society. 26 (2): 133–179. doi:10.1111/j.1467-968X.1908.tb00513.x. 
  153. ^ Klatt, E. C. "Firearms Tutorial". library.med.utah.edu. Retrieved 5 February 2017. 
  154. ^ Bastasch, M. (2015). "California officially bans hunters from using lead bullets". The Daily Caller. Retrieved 4 July 2016. 
  155. ^ Parker, R. B. (2005). The New Cold-Molded Boatbuilding: From Lofting to Launching. WoodenBoat Books. pp. 194–195. ISBN 978-0-937822-89-0. 
  156. ^ Krestovnikoff, M.; Halls, M. (2006). Scuba Diving. Penguin. p. 70. ISBN 978-0-7566-4063-7. 
  157. ^ Street & Alexander 1998, p. 182.
  158. ^ Jensen, C. F. (2013). Online Location of Faults on AC Cables in Underground Transmission. Springer. p. 136. ISBN 978-3-319-05397-4. 
  159. ^ Copper Development Association. "Leaded Coppers". copper.org. Retrieved 10 July 2016. 
  160. ^ Guruswamy, S. (2000). Engineering properties and applications of lead alloys. Marcel Dekker. p. 31. ISBN 978-0-8247-8247-4. 
  161. ^ Lansdown, R.; Yule, William, eds. (1986). The Lead debate : the environment, toxicology, and child health. Croom Helm. p. 240. ISBN 978-0-7099-1653-6. 
  162. ^ Audsley, G. A. (1965). The Art of Organ Building. 2. Courier. pp. 250–251. ISBN 978-0-486-21315-6. 
  163. ^ Palmieri, R., ed. (2006). The Organ. Garland. pp. 412–413. ISBN 978-0-415-94174-7. 
  164. ^ "Think Lead research summary" (PDF). The Lead Sheet Association. Retrieved 20 February 2017. 
  165. ^ "Weatherings to Parapets and Cornices". The Lead Sheet Association. Retrieved 20 February 2017. 
  166. ^ See, for example: "Lead garden ornaments". H. Crowther Ltd. 2016. Retrieved 20 February 2017. …producing quality garden ornament from our studio in West London for over a century. 
  167. ^ "Lead" (PDF). Mineral Commodities Summaries. U.S. Geological Survey. 2016. p. 97. Retrieved 20 February 2017. 
  168. ^ Progressive Dynamics, Inc. "How Lead Acid Batteries Work: Battery Basics". progressivedyn.com. Retrieved 3 July 2016. 
  169. ^ National Council on Radiation Protection and Measurements (2004). Structural Shielding Design for Medical X-ray Imaging Facilities. National Council on Radiation Protection and Measurement. pp. 16–17. ISBN 978-0-929600-83-3. 
  170. ^ Tuček, K.; Carlsson, J.; Wider, H. (2006). "Comparison of sodium and lead-cooled fast reactors regarding reactor physics aspects, severe safety and economical issues" (PDF). Nuclear Engineering and Design. 236 (14–16): 1589–1598. doi:10.1016/j.nucengdes.2006.04.019. 
  171. ^ Leonard, A. R.; Lynch, G. (1958). "Dishware as a possible source for lead poisoning". California Medicine. 89 (6): 414–416. PMC 1512529Freely accessible. PMID 13608300. 
  172. ^ Zweifel, H. (2009). Plastics Additives Handbook. Hanser. p. 438. ISBN 978-3-446-40801-2. 
  173. ^ Wilkes, C. E.; Summers, J. W.; Daniels, C. A.; et al. (2005). PVC Handbook. Hanser. p. 106. ISBN 978-1-56990-379-7. 
  174. ^ Randerson, J. (2002). "Candle pollution". newscientist.com (2348). Retrieved 7 April 2007. 
  175. ^ Nriagu, J.; Kim, M. J. (2000). "Emissions of lead and zinc from candles with metal-core wicks". The Science of the Total Environment. 250 (1–3): 37–41. doi:10.1016/S0048-9697(00)00359-4. PMID 10811249. 
  176. ^ Amstock, J. S. (1997). Handbook of Glass in Construction. McGraw-Hill Professional. pp. 116–119. ISBN 978-0-07-001619-4. 
  177. ^ Rogalski, Antonio (2010). Infrared Detectors, Second Edition. CRC Press. pp. 485–541. ISBN 978-1-4200-7672-1. Retrieved 19 November 2016. 
  178. ^ Roberts, Stephen M.; James, Robert C.; Williams, Phillip L. (2014). Principles of Toxicology: Environmental and Industrial Applications. Wiley. p. 289. ISBN 978-1-118-98251-8. 
  179. ^ John Emsley (2001). Nature's Building Blocks: An A-Z Guide to the Elements. Oxford University. p. 226. ISBN 978-0-19-850341-5. 
  180. ^ "Lead". Air quality guidelines for Europe (PDF). World Health Organization, Regional Office for Europe. 2000. p. 3. ISBN 978-92-890-1358-1. OCLC 475274390. 
  181. ^ U.S. Food and Drug Administration (2015). Q3D Elemental Impurities Guidance for Industry (PDF) (Report). U. S. Department of Health and Human Services. p. 41. Retrieved 15 February 2017. 
  182. ^ a b Rudolph, Abraham M.; Rudolph, Colin D.; Hostetter, Margaret K.; Lister, George E.; Siegel, Norman J. (2003). "Lead". Rudolph's Pediatrics (21st ed.). McGraw-Hill Professional. p. 369. ISBN 978-0-8385-8285-5. 
  183. ^ Mycyk, M.; Hryhorczuk, D.; Amitai, Y.; et al. (2005). "Lead". In Erickson, T. B.; Ahrens, W. R.; Aks, S. Pediatric Toxicology: Diagnosis and Management of the Poisoned Child. McGraw-Hill Professional. p. 462. ISBN 978-0-07-141736-5. 
  184. ^ Dart, R. C.; Hurlbut, K. M.; Boyer-Hassen, L. V. (2004). "Lead". In Dart, R. C. Medical Toxicology (3rd ed.). Lippincott Williams & Wilkins. p. 1426. ISBN 0-7817-2845-2. 
  185. ^ a b Venugopal, B. (2013). Physiologic and Chemical Basis for Metal Toxicity. Springer. pp. 177–178. ISBN 978-1-4684-2952-7. 
  186. ^ Kosnett, M. J. (2006). "Lead". In Olson, K. R. Poisoning and Drug Overdose (5th ed.). McGraw-Hill Professional. p. 238. ISBN 978-0-07-144333-3. 
  187. ^ "Toxic Substances Portal – Lead". Agency for Toxic Substance and Disease Registry. Archived from the original on 6 June 2011. 
  188. ^ Unites States Environmental Protection Agency (2016). "Basic Information about Lead Air Pollution". epa.gov. Retrieved 9 November 2016. 
  189. ^ Sokol, R. C. (2005). "Summary". In Golub, M. S. Metals, Fertility, and Reproductive Toxicity. Taylor and Francis. p. 153. ISBN 978-0-415-70040-5. 
  190. ^ "ToxFAQs: CABS/Chemical Agent Briefing Sheet: Lead" (PDF). Agency for Toxic Substances and Disease Registry/Division of Toxicology and Environmental Medicine. 2006. Archived from the original (PDF) on 4 March 2010. 
  191. ^ Bergeson, Lynn L. (2008). "The proposed lead NAAQS: Is consideration of cost in the clean air act's future?". Environmental Quality Management. 18: 79–84. doi:10.1002/tqem.20197. 
  192. ^ Casciani, D. (2014). "Did removing lead from petrol spark a decline in crime?". BBC News. Retrieved 2017-01-30. 
  193. ^ Jagadish Prasad, P. (2010). Conceptual Pharmacology. Universities Press. p. 652. ISBN 978-81-7371-679-9. Retrieved 21 June 2012. 
  194. ^ Masters, S. B.; Trevor, A. J.; Katzung, B. G. (2008). Katzung & Trevor's Pharmacology: Examination & Board Review (8th ed.). McGraw Hill Medical. pp. 481–483. ISBN 978-0-07-148869-3. 
  195. ^ Needleman, H. L.; Schell, A.; Bellinger, D.; et al. (1990). "The long-term effects of exposure to low doses of lead in childhood. An 11-year follow-up report". New England Journal of Medicine. 322 (2): 83–88. doi:10.1056/NEJM199001113220203. PMID 2294437. 
  196. ^ Hu, H. (1991). "Knowledge of diagnosis and reproductive history among survivors of childhood plumbism". American Journal of Public Health. 81 (8): 1070–1072. doi:10.2105/AJPH.81.8.1070. PMC 1405695Freely accessible. PMID 1854006. 
  197. ^ Liu, J.; et al. (2015). "Early blood lead levels and sleep disturbance in preadolescence". PubMed. 38 (12): 1869–1874. doi:10.5665/sleep.5230. PMC 4667382Freely accessible. PMID 26194570. 
  198. ^ Schoeters, G.; Den Hond, E.; Dhooge, W.; et al. (2008). "Endocrine disruptors and abnormalities of pubertal development". Basic & Clinical Pharmacology & Toxicology. 102 (2): 168–175. doi:10.1111/j.1742-7843.2007.00180.x. PMID 18226071. 
  199. ^ Berdanier, C. D.; Dwyer, J. T.; Heber, D. (2016). Handbook of Nutrition and Food (3rd ed.). CRC Press. pp. 211–226. ISBN 978-1-4665-0572-8. Retrieved 3 July 2016. 
  200. ^ Gottschlich, M. M. (2001). The Science and Practice of Nutrition Support: A Case-based Core Curriculum. Kendall Hunt. p. 98. ISBN 978-0-7872-7680-5. Retrieved 9 July 2016. 
  201. ^ Insel, P. M.; Turner, R. Elaine; Ross, Don (2004). Nutrition. Jones & Bartlett Learning. p. 499. ISBN 978-0-7637-0765-1. Retrieved 10 July 2016. 
  202. ^ "Information for the Community Lead Toxicity". Agency for Toxic Substances and Disease Registry (MP4 webcast, 82 MB). Retrieved 11 February 2017. 
  203. ^ Moore, M. R. (1977). "Lead in drinking water in soft water areas—health hazards". Science of the Total Environment. 7 (2): 109–115. doi:10.1016/0048-9697(77)90002-X. PMID 841299. 
  204. ^ a b c Cite error: The named reference health was invoked but never defined (see the help page).
  205. ^ "NIOSH Adult Blood Lead Epidemiology and Surveillance". United States National Institute for Occupational Safety and Health. Retrieved 4 October 2007. 
  206. ^ United Nations Environment Programme, Chemicals Branch, Division of Technology, Industry and Economics (2010). Final review of scientific information on lead (PDF). United Nations Environment Programme. pp. 11–33. Retrieved 31 January 2017. 
  207. ^ Greene, D. (10 May 2014). "Effects of lead on the environment". lead.org.au. Retrieved 30 October 2016. 
  208. ^ Deltares; Netherlands Organisation for Applied Scientific Research (2016). Lood en zinkemissies door jacht [Lead and zinc emissions from hunting] (PDF) (Report) (in Dutch). Retrieved 18 February 2017. 
  209. ^ Freeman, K. S. (January 2012). "Remediating soil lead with fishbones". Environmental Health Perspectives. 120 (1): a20–a21. doi:10.1289/ehp.120-a20a. PMC 3261960Freely accessible. PMID 22214821. 
  210. ^ Young, S. (9 July 2012). "Battling lead contamination, one fish bone at a time". Compass. United States Coast Guard. Retrieved 11 February 2017. 
  211. ^ Bairagi, H.; Khan, M.; Ray, L.; et al. (2011). "Adsorption profile of lead on Aspergillus versicolor: A mechanistic probing". Journal of Hazardous Materials. 186 (1): 756–764. doi:10.1016/j.jhazmat.2010.11.064. PMID 21159429. 
  212. ^ Park, J. H.; Bolan, N.; Meghara, M.; et al. (2011). "Bacterial-assisted immobilization of lead in soils: Implications for remediation" (PDF). Pedologist: 162–174. Archived from the original (PDF) on 26 November 2015. 
  213. ^ "Lead Paint Information". Master Painters Australia. Archived from the original on 12 February 2008. Retrieved 7 April 2007. 
  214. ^ Smith, D. R.; Flegal, A. R. (1995). "Lead in the biosphere: Recent trends". AMBIO. 24: 21–23. JSTOR 4314280. 
  215. ^ Grandjean, P. (1978). "Widening perspectives of lead toxicity". Environmental Research. 17 (2): 303–321. doi:10.1016/0013-9351(78)90033-6. PMID 400972. 
  216. ^ a b Levin, R.; Brown, M. J.; Kashtock, M. E.; et al. (2008). "Lead exposures in U.S. children, 2008: Implications for prevention". Environmental Health Perspectives. 116 (10): 1285–1293. doi:10.1289/ehp.11241. PMC 2569084Freely accessible. PMID 18941567. 
  217. ^ Marino, P. E.; Landrigan, P. J.; Graef, J.; et al. (1990). "A case report of lead paint poisoning during renovation of a Victorian farmhouse". American Journal of Public Health. 80 (10): 1183–1185. doi:10.2105/AJPH.80.10.1183. PMC 1404824Freely accessible. PMID 2119148. 
  218. ^ The National Institute for Occupational Safety and Health. "CDC – NIOSH Pocket Guide to Chemical Hazards - Lead". www.cdc.gov. Retrieved 18 November 2016. 

Bibliography[edit]

Further reading[edit]

External links[edit]