Lead (metal)

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Lead,  ₈₂Pb

Lead
Pronunciation	/ˈlɛd/ ​(LED)
Appearance	metallic gray
Standard atomic weight Ar, std(Pb)	207.2(1)
Lead in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson Sn ↑ Pb ↓ Fl thallium ← lead → bismuth
Atomic number (Z)	82
Group	group 14 (carbon group)
Period	period 6
Block	p-block
Element category	post-transition metal
Electron configuration	[Xe] 4f14 5d10 6s2 6p2
Electrons per shell	2, 8, 18, 32, 18, 4
Physical properties
Phase at STP	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, −2, −1, +1, +2, +3, +4 (an amphoteric oxide)
Electronegativity	Pauling scale: 1.87 (+2)
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
Spectral lines of lead
Other properties
Natural occurrence	primordial
Crystal structure	​face-centered cubic (fcc)
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	−23.0×10−6 cm3/mol (at 298 K)
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	in the Middle East (7000 BCE)
Main isotopes of lead
Iso­tope Abun­dance Half-life (t1/2) Decay mode Pro­duct 204Pb 1.4% stable 206Pb 24.1% stable 207Pb 22.1% stable 208Pb 52.4% stable
Isotopic abundances vary greatly by sample

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

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

Lead is easily extracted from its ores; prehistoric people in Western Asia knew of it. Galena, a principal ore of lead, often bears silver, interest in which helped initiate widespread extraction and use of lead in ancient Rome. Lead production declined after the fall of Rome and did not reach comparable levels until the Industrial Revolution. In 2014, the annual global production of lead was about ten million tonnes, over half of which was from recycling. Lead's high density, low melting point, ductility and relative inertness to oxidation make it useful. These properties, combined with its relative abundance and low cost, resulted in its extensive use in construction, plumbing, batteries, bullets and shot, weights, solders, pewters, fusible alloys, white paints, leaded gasoline, and radiation shielding.

In the late 19th century, lead's toxicity was recognized, and its use has since been phased out of many applications. However, many countries still allow the sale of products that expose humans to lead, including some types of paints and bullets. Lead is a toxin that accumulates in soft tissues and bones, it acts as a neurotoxin damaging the nervous system and interfering with the function of biological enzymes, causing neurological disorders, such as brain damage and behavioral problems.

Physical properties

Atomic

A lead atom has 82 electrons, arranged in an electron configuration of [Xe]4f¹⁴5d¹⁰6s²6p². The sum of lead's 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 the carbon group. This is unusual; ionization energies generally fall going down a group, as an element's outer electrons become more distant from the nucleus, and more shielded by smaller orbitals. The similarity of ionization energies is caused by the lanthanide contraction—the decrease in element radii from lanthanum (atomic number 57) to lutetium (71), and the relatively small radii of the elements from hafnium (72) onwards. This is due to poor shielding of the nucleus by the lanthanide 4f electrons. The sum of the first four ionization energies of lead exceeds that of tin, contrary to what periodic trends would predict. Relativistic effects, which become significant in heavier atoms, contribute to this behavior. One such effect is the inert pair effect: the 6s electrons of lead become reluctant to participate in bonding, making the distance between nearest atoms in crystalline lead unusually long.

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

Bulk

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

A sample of lead solidified from the molten state

 

Lead's close-packed face-centered cubic structure and high atomic weight result in a density of 11.34 g/cm³, which is greater than that of common metals such as iron (7.87 g/cm³), copper (8.93 g/cm³), and zinc (7.14 g/cm³). This density is the origin of the idiom to go over like a lead balloon. Some rarer metals are denser: tungsten and gold are both at 19.3 g/cm³, and osmium—the densest metal known—has a density of 22.59 g/cm³, almost twice that of lead.

Lead is a very soft metal with a Mohs hardness of 1.5; it can be scratched with a fingernail. It is quite malleable and somewhat ductile. The bulk modulus of lead—a measure of its ease of compressibility—is 45.8 GPa. In comparison, that of aluminium is 75.2 GPa; copper 137.8 GPa; and mild steel 160–169 GPa. Lead's tensile strength, at 12–17 MPa, is low (that of aluminium is 6 times higher, copper 10 times, and mild steel 15 times higher); it can be strengthened by adding small amounts of copper or antimony.

The melting point of lead—at 327.5 °C (621.5 °F)—is very low compared to most metals. Its boiling point of 1749 °C (3180 °F) is the lowest among the carbon group elements. The electrical resistivity of lead at 20 °C is 192 nanoohm-meters, almost an order of magnitude higher than those of other industrial metals (copper at 15.43 nΩ·m; gold 20.51 nΩ·m; and aluminium at 24.15 nΩ·m). Lead is a superconductor at temperatures lower than 7.19 K; this is the highest critical temperature of all type-I superconductors and the third highest of the elemental superconductors.

Main isotopes of lead (₈₂Pb)

Iso­tope Decay abun­dance half-life (t1/2) mode pro­duct 202Pb syn 5.25(28)×104 y ε 202Tl 204Pb 1.4% stable 205Pb trace 1.53(7)×107 y ε 205Tl 206Pb 24.1% stable 207Pb 22.1% stable 208Pb 52.4% stable 209Pb trace 3.253(14) h β− 209Bi 210Pb trace 22.3(22) y β− 210Bi 211Pb trace 36.1(2) min β− 211Bi 212Pb trace 10.64(1) h β− 212Bi 214Pb trace 26.8(9) min β− 214Bi Isotopic abundances vary greatly by sample

Standard atomic weight Ar, standard(Pb)

207.2(1)

Isotopes

Natural lead consists of four stable isotopes with mass numbers of 204, 206, 207, and 208, and traces of five short-lived radioisotopes. The high number of isotopes is consistent with lead's atomic number being even. Lead has a magic number of protons (82), for which the nuclear shell model accurately predicts an especially stable nucleus. Lead-208 has 126 neutrons, another magic number, which may explain why lead-208 is extraordinarily stable.

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

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 chain, the actinium chain, and the thorium chain. Their isotopic concentrations in a natural rock sample depends greatly on the presence of these three parent uranium and thorium isotopes. For example, the relative abundance of lead-208 can range from 52% in normal samples to 90% in thorium ores; for this reason, the standard atomic weight of lead is given to only one decimal place. As time passes, the ratio of lead-206 and lead-207 to lead-204 increases, since the former two are supplemented by radioactive decay of heavier elements while the latter is not; this allows for lead–lead dating. As uranium decays into lead, their relative amounts change; this is the basis for uranium–lead dating. Lead-207 exhibits nuclear magnetic resonance, a property that has been used to study its compounds in solution and solid state, including in human body.

The Holsinger meteorite, 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.

 

Apart from the stable isotopes, which make up almost all lead that exists naturally, there are trace quantities of a few radioactive isotopes. One of them is lead-210; although it has a half-life of only 22.3 years, small quantities occur in nature because lead-210 is produced by a long decay series that starts with uranium-238 (which has been present for billions of years on Earth). Lead-211, -212, and -214 are present in the decay chains of uranium-235, thorium-232, and uranium-238, respectively, so traces of all three of these lead isotopes are found naturally. Minute traces of lead-209 arise from the very rare cluster decay of radium-223, one of the daughter products of natural uranium-235, and the decay chain of neptunium-237, traces of which are produced by neutron capture in uranium ores. Lead-210 is particularly useful for helping to identify the ages of samples by measuring its ratio to lead-206 (both isotopes are present in a single decay chain).

In total, 43 lead isotopes have been synthesized, with mass numbers 178–220. Lead-205 is the most stable radioisotope, with a half-life of around 1.5×10⁷ years. The second-most stable is lead-202, which has a half-life of about 53,000 years, longer than any of the natural trace radioisotopes.

Chemistry

Flame test: lead colors flame pale blue

 

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

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

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

Inorganic compounds

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

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

Lead(II) oxide

Lead(II)

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

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

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

 

Lead dihalides are well-characterized; this includes the diastatide, and mixed halides, such as PbFCl. The relative insolubility of the latter forms a useful basis for the gravimetric determination of fluorine. The difluoride was the first solid ionically conducting compound to be discovered (in 1834, by Michael Faraday). The other dihalides decompose on exposure to ultraviolet or visible light, especially the diiodide. Many lead(II) pseudohalides are known, such as the cyanide, cyanate, and thiocyanate. Lead(II) forms an extensive variety of halide coordination complexes, such as [PbCl₄]²⁻, [PbCl₆]⁴⁻, and the [Pb₂Cl₉]_n⁵ⁿ⁻ chain anion.

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

Lead(IV)

Few inorganic lead(IV) compounds are known. They are only formed in highly oxidizing solutions and do not normally exist under standard conditions. Lead(II) oxide gives a mixed oxide on further oxidation, Pb₃O₄. It is described as lead(II,IV) oxide, or structurally 2PbO·PbO₂, 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 and lead diselenide are only stable at high pressures. Lead tetrafluoride, a yellow crystalline powder, is stable, but less so than the difluoride. Lead tetrachloride (a yellow oil) decomposes at room temperature, lead tetrabromide is less stable still, and the existence of lead tetraiodide is questionable.

The capped square antiprismatic anion [Pb₉]⁴⁻ from [K(18-crown-6)]₂K₂Pb₉·(en)_1.5

Other oxidation states

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; this oxidation state is not stable, as both the lead(III) ion and the larger complexes containing it are radicals. The same applies for lead(I), which can be found in such radical species.

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 phases. Many of them show defective fluorite structures in which some oxygen atoms are replaced by vacancies: PbO can be considered as having such a structure, with every alternate layer of oxygen atoms absent.

Negative oxidation states can occur as Zintl phases, as either free lead anions, as in Ba₂Pb, with lead formally being lead(−IV), or in oxygen-sensitive ring-shaped or polyhedral cluster ions such as the trigonal bipyramidal Pb₅²⁻ ion, where two lead atoms are lead(−I) and three are lead(0). 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. They may be made in liquid ammonia via the reduction of lead by sodium.

Structure of a tetraethyllead molecule:
     Carbon
     Hydrogen
     Lead

Organolead

Lead can form multiply-bonded chains, a property it shares with its lighter homologs in the carbon group. Its capacity to do so is much less because the Pb–Pb bond energy is over three and a half times lower than that of the C–C bond. With itself, lead can build metal–metal bonds of an order up to three. With carbon, lead forms organolead compounds similar to, but generally less stable than, typical organic compounds (due to the Pb–C bond being rather weak). This makes the organometallic chemistry of lead far less wide-ranging than that of tin. Lead predominantly forms organolead(IV) compounds, even when starting with inorganic lead(II) reactants; very few organolead(II) compounds are known. The most well-characterized exceptions are Pb[CH(SiMe₃)₂]₂ and Pb(η⁵-C₅H₅)₂.

The lead analog of the simplest organic compound, methane, is plumbane. Plumbane may be obtained in a reaction between metallic lead and atomic hydrogen. Two simple derivatives, tetramethyllead and tetraethyllead, are the best-known organolead compounds. These compounds are relatively stable: tetraethyllead only starts to decompose if heated or if exposed to sunlight or ultraviolet light. (Tetraphenyllead is even more thermally stable, decomposing at 270 °C.) With sodium metal, lead readily forms an equimolar alloy that reacts with alkyl halides to form organometallic compounds such as tetraethyllead. The oxidizing nature of many organolead compounds is usefully exploited: lead tetraacetate is an important laboratory reagent for oxidation in organic synthesis, and tetraethyllead was once produced in larger quantities than any other organometallic compound. Other organolead compounds are less chemically stable. For many organic compounds, a lead analog does not exist.

Origin and occurrence

Solar System abundances

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

In space

Lead's per-particle abundance in the Solar System is 0.121 ppb (parts per billion). This figure is two and a half times higher than that of platinum, eight times more than mercury, and seventeen times more than gold. The amount of lead in the universe is slowly increasing as most heavier atoms (all of which are unstable) gradually decay to lead. The abundance of lead in the Solar System since its formation 4.5 billion years ago has increased by about 0.75%. The solar system abundances table shows that lead, despite its relatively high atomic number, is more prevalent than most other elements with atomic numbers greater than 40.

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- and r-processes.

In the s-process (s is for "slow"), captures are separated by years or decades, allowing less stable nuclei to undergo beta decay. A stable thallium-203 nucleus can capture a neutron and become thallium-204; this undergoes beta decay to give stable lead-204; on capturing another neutron, it becomes lead-205, which has a half-life of 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 beta decays to polonium-210, which alpha decays to lead-206. The cycle hence ends at lead-206, lead-207, lead-208, and bismuth-209.

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; the green arrow represents an alpha decay; 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, such as a supernova or the merger of two neutron stars. The neutron flux involved may be on the order of 10²² neutrons per square centimeter per second. The r-process does not form as much lead as the s-process. It tends to stop once neutron-rich nuclei reach 126 neutrons. At this point, the neutrons are arranged in complete shells in 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.

On Earth

Lead is classified as a chalcophile under the Goldschmidt classification, meaning it is generally found combined with sulfur. It rarely occurs in its native, metallic form. Many lead minerals are relatively light and, over the course of the Earth's history, have remained in the crust instead of sinking deeper into the Earth's interior. This accounts for lead's relatively high crustal abundance of 14 ppm; it is the 38th most abundant element in the crust.

The main lead-bearing mineral is galena (PbS), which is mostly found with zinc ores. Most other lead minerals are related to galena in some way; boulangerite, Pb₅Sb₄S₁₁, is a mixed sulfide derived from galena; anglesite, PbSO₄, is a product of galena oxidation; and cerussite or white lead ore, PbCO₃, is a decomposition product of galena. Arsenic, tin, antimony, silver, gold, copper, and bismuth are common impurities in lead minerals.

Lead is a fairly common element in the Earth's crust for its high atomic number (82). Most elements of atomic number greater than 40 are less abundant.

 

World lead resources exceed two billion tons. 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 88 million tons in 2016, of which Australia had 35 million, China 17 million, and Russia 6.4 million.

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.

Etymology

The modern 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). The Old English word is derived from the hypothetical reconstructed Proto-Germanic *lauda- ("lead"). According to linguistic theory, this word bore descendants in multiple Germanic languages of exactly the same meaning.

The origin of the Proto-Germanic *lauda- is not agreed in the linguistic community. One hypothesis suggests it is derived from Proto-Indo-European *lAudh- ("lead"; capitalization of the vowel is equivalent to the macron). Another hypothesis suggests it is borrowed from Proto-Celtic *ɸloud-io- ("lead"). This word is related to the Latin plumbum, which gave the element its chemical symbol Pb. The word *ɸloud-io- is thought to be the origin of Proto-Germanic *bliwa- (which also means "lead"), from which stemmed the German Blei.

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

History

World lead production peaking in the Roman period and the Industrial Revolution.

Prehistory and early history

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

Classical era

Because silver was extensively used as a decorative material and an exchange medium, lead deposits came to be worked in Asia Minor since 3000 BCE; later, lead deposits were developed in the Aegean and Laurion. These three regions collectively dominated production of mined lead until c. 1200 BCE. Since 2000 BCE, the Phoenicians worked deposits in the Iberian peninsula; by 1600 BCE, lead mining existed in Cyprus, Greece, and Sardinia.

Ancient Greek lead sling bullets with a winged thunderbolt molded on one side and the inscription "ΔΕΞΑΙ" ("take that" or "catch") on the other side

 

Rome's territorial expansion in Europe and across the Mediterranean, and its 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. Lead mining occurred in Central Europe, Britain, the Balkans, Greece, Anatolia, and Hispania, the latter accounting for 40% of world production.

Lead tablets were commonly used as a material for letters. Lead coffins, cast in flat sand forms, with interchangeable motifs to suit the faith of the deceased were used in ancient Judea.

Lead was used for making water pipes in the Roman Empire; the Latin word for the metal, plumbum, is the origin of the English word "plumbing". Its ease of working and resistance to corrosion ensured its widespread use in other applications including pharmaceuticals, roofing, currency, and warfare. Writers of the time, such as 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 flavor through verdigris formation.

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

The Roman author Vitruvius reported the health dangers of lead and modern writers have suggested that lead poisoning played a major role in the decline of the Roman Empire. Other researchers have criticized such claims, pointing out, for instance, that not all abdominal pain is caused by lead poisoning. According to archaeological research, Roman lead pipes increased lead levels in tap water but such an effect was "unlikely to have been truly harmful". When lead poisoning did occur, victims were called "saturnine", dark and cynical, after the ghoulish father of the gods, Saturn. By association, lead was considered the father of all metals. Its status in Roman society was low as it was readily available and cheap.

Roman lead pipes

Confusion with tin and antimony

During the classical era (and even up to the 17th century), tin was often not 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, its cognate олово (olovo) means "tin". To add to the confusion, lead bore a close relation to antimony: both elements commonly occur as sulfides (galena and stibnite), often together. Pliny incorrectly wrote that stibnite would give lead on heating, instead of antimony. In countries such as Turkey and India, the originally Persian name surma came to refer to either antimony sulfide or lead sulfide, and in some languages, such as Russian, gave its name to antimony (сурьма).

Middle Ages and the Renaissance

Lead mining in Western Europe declined after the fall of the Western Roman Empire, with Arabian Iberia being the only region having a significant output. The largest production of lead occurred in South and East Asia, especially China and India, where lead mining grew rapidly.

Elizabeth I of England was commonly depicted with a whitened face. Lead in face whiteners is thought to have contributed to her death.

 

In Europe, lead production began to increase in the 11th and 12th centuries, when it was again used for roofing and piping. Starting in the 13th century, lead was used to create stained glass. In the European and Arabian traditions of alchemy, lead (symbol in the European tradition) was considered an impure base metal which, by the separation, purification and balancing of its constituent essences, could be transformed to pure and incorruptible gold. During the period, lead was used increasingly for adulterating wine. The use of such wine was forbidden for use in Christian rites by a papal bull in 1498, but it continued to be imbibed and resulted in mass poisonings up to the late 18th century. Lead was a key material in parts of the printing press, which was invented around 1440; lead dust was commonly inhaled by print workers, causing lead poisoning. 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. Lead, in the form of Venetian ceruse, was extensively used in cosmetics by Western European aristocracy as whitened faces were regarded as a sign of modesty. This practice later 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 faces of women "came to represent their feminine virtue as Japanese women", with lead commonly used in the whitener.

Outside Europe and Asia

In the New World, lead production was recorded soon after the arrival of European settlers. The earliest record dates to 1621 in the English Colony of Virginia, fourteen years after its foundation. In Australia, the first mine opened by colonists on the continent was a lead mine, in 1841. In Africa, lead mining and smelting were known in the Benue Trough and the lower Congo Basin, where lead was used for trade with Europeans, and as a currency by the 17th century, well before the scramble for Africa. 

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

Industrial Revolution

In the second half of the 18th century, Britain, and later continental Europe and the United States, experienced the Industrial Revolution. This was the first time during which lead production rates exceeded those of Rome. 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. By 1900, the United States was the leader in global lead production, and other non-European nations—Canada, Mexico, and Australia—had begun significant production; production outside Europe exceeded that within. A great share of the demand for lead came from plumbing and painting—lead paints were in regular use. At this time, more (working class) people were exposed to the metal and 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; British physician Alfred Baring Garrod noted a third of his gout patients were plumbers and painters. The effects of chronic ingestion of lead, including mental disorders, were also studied in the 19th century. The first laws aimed at decreasing lead poisoning in factories were enacted during the 1870s and 1880s in the United Kingdom.

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

Modern era

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, lead blindness was documented, and the element was phased out of public use in the United States and Europe. The United Kingdom introduced mandatory factory inspections in 1878 and appointed the first Medical Inspector of Factories in 1898; as a result, a 25-fold decrease in lead poisoning incidents from 1900 to 1944 was reported. Most European countries banned lead paint—commonly used because of its opacity and water resistance—for interiors by 1930.

The last major human exposure to lead was the addition of tetraethyllead to gasoline as an antiknock agent, a practice that originated in the United States in 1921. It was phased out in the United States and the European Union by 2000.

In the 1970s, the United States and Western European countries introduced legislation to reduce lead air pollution. The impact was significant: while a study conducted by the Centers for Disease Control and Prevention in the United States in 1976–1980 showed that 77.8% of the population had elevated blood lead levels, in 1991–1994, a study by the same institute showed the share of people with such high levels dropped to 2.2%. The main product made of lead by the end of the 20th century was the lead–acid battery, which posed no direct threat to humans. 

From 1960 to 1990, lead output in the Western Bloc grew by about 31%. 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 the world's largest producer during the mid-1970s and the 1980s, and China starting major lead production in the late 20th century. 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. As was the case during European industrialization, lead has had a negative effect on health in China.

Production

Primary production of lead since 1840

 

As of 2014, production 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 2014, 4.58 million metric tons came from primary production and 5.64 million from secondary production. The top three producers of mined lead concentrate in that year were China, Australia, and the United States. The top three producers of refined lead were China, the United States, and India. According to the International Resource Panel's Metal Stocks in Society report of 2010, the total amount of lead in use, stockpiled, discarded, or dissipated into the environment, on a global basis, is 8 kg per capita. Much of this is in more developed countries (20–150 kg per capita) rather than less developed ones (1–4 kg per capita).

The primary and secondary lead production processes are similar. Some primary production plants now supplement their operations with scrap lead, and this trend is likely to increase in the future. Given adequate techniques, lead obtained via secondary processes is indistinguishable from lead obtained via primary processes. 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.

Primary

Most lead ores contain a low percentage of lead (rich ores have a typical content of 3–8%) which must be concentrated for extraction. During initial processing, ores typically undergo crushing, dense-medium separation, grinding, froth flotation, and drying. The resulting concentrate, which has a lead content of 30–80% by mass (regularly 50–60%), is then turned into (impure) lead metal. 

There are two main ways of doing this: a two-stage process involving roasting followed by blast furnace extraction, carried out in separate vessels; or a direct process in which the extraction of the concentrate occurs in a single vessel. The latter has become the most common route, though the former is still significant.

Two-stage process

First, the sulfide concentrate is roasted in air to oxidize the lead sulfide:

2 PbS(s) + 3 O₂(g) → 2 PbO(s) + 2 SO₂(g)↑

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

2 PbO(s) + C(s) → 2 Pb(s) + CO₂(g)↑

Impurities are mostly arsenic, antimony, bismuth, zinc, copper, silver, and gold. Typically they are removed in a series of pyrometallurgical processes. 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 float to the top of the melt and are skimmed off. Metallic silver and gold are removed and recovered economically by means of the Parkes process, in which zinc is added to lead. Zinc, which is immiscible in lead, dissolves the silver and gold. The zinc solution can be separated from the lead, and the silver and gold retrieved. De-silvered lead is freed of bismuth by the Betterton–Kroll process, treating it with metallic calcium and magnesium. The resulting bismuth dross can be skimmed off.

Alternatively to the pyrometallurgical processes, 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 lead fluorosilicate (PbSiF₆). Once electrical potential is applied, impure lead at the anode dissolves and plates onto the cathode, leaving the majority of the impurities in solution. This is a high-cost process and thus mostly reserved for refining bullion containing high percentages of impurities.

Direct process

In this process, lead bullion and slag is obtained directly from lead concentrates. The lead sulfide concentrate is melted in a furnace and oxidized, forming lead monoxide. Carbon (as coke or coal gas) is added to the molten charge along with fluxing agents. The lead monoxide is thereby reduced to metallic lead, in the midst of a slag rich in lead monoxide.

If the input is rich in lead, as much as 80% of the original lead can be obtained as bullion; the remaining 20% forms a slag rich in lead monoxide. For a low-grade feed, all of the lead can be oxidized to a high-lead slag. Metallic lead is further obtained from the high-lead (25–40%) slags via submerged fuel combustion or injection, reduction assisted by an electric furnace, or a combination of both.

Alternatives

Research on a cleaner, less energy-intensive lead extraction process continues; a major drawback is that either too much lead is lost as waste, or the alternatives result in a high sulfur content in the resulting lead metal. Hydrometallurgical extraction, in which anodes of impure lead are immersed into an electrolyte and pure lead is deposited onto a cathode, is a technique that may have potential, but is not currently economical except in cases where electricity is very cheap.

Secondary

Smelting, which is an essential part of the primary production, is often skipped during secondary production. It is only performed when metallic lead has undergone significant oxidation. The process is similar to that of primary production in either a blast furnace or a rotary furnace, with the essential difference being the greater variability of yields: blast furnaces produce hard lead (10% antimony) while reverberatory and rotary kiln furnaces produced semisoft lead (3–4% antimony). The Isasmelt process is a more recent smelting method that may act as an extension to primary production; battery paste from spent lead–acid batteries (containing lead sulfate and lead oxides) has its sulfate removed by treating it with alkali, and is then treated in a coal-fueled furnace in the presence of oxygen, which yields impure lead, with antimony the most common impurity. 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.

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

Applications

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

 

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).

Elemental form

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 are generally less common and more difficult to extract from parent ores. Lead's toxicity has led to its phasing out for some uses.

Lead has been used for bullets since their invention in the Middle Ages. It is inexpensive; its low melting point means small arms ammunition and shotgun pellets can be cast with minimal technical equipment; and it is denser than other common metals, which allows for better retention of velocity. It remains the main material for bullets, alloyed with other metals as hardeners. Concerns have been raised that lead bullets used for hunting can damage the environment.

Lead's high density and resistance to corrosion have been exploited in a number of related applications. It is used as ballast in sailboat keels; its density allows it to take up a small volume and minimize water resistance, thus counterbalancing the heeling effect of wind on the sails. It is used in scuba diving weight belts to counteract the diver's buoyancy. In 1993, the base of the Leaning Tower of Pisa was stabilized with 600 tonnes of lead. Because of its corrosion resistance, lead is used as a protective sheath for underwater cables.

A 17th-century gold-coated lead sculpture

 

Lead has many uses in the construction industry; lead sheets are used as architectural metals in roofing material, cladding, flashing, gutters and gutter joints, and on roof parapets. Detailed lead moldings are used as decorative motifs to fix lead sheet. Lead is still used in statues and sculptures, including for armatures. In the past it was often used to balance the wheels of cars; for environmental reasons this use is being phased out in favor of other materials.

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 in imperfections throughout the alloy, such as grain boundaries. In low concentrations, as well as acting as a lubricant, the globules hinder the formation of swarf as the alloy is worked, thereby improving machinability. Copper alloys with larger concentrations of lead are used in bearings. The lead provides lubrication, and the copper provides the load-bearing support.

Lead's high density, atomic number, and formability form the basis for use of lead as a barrier that absorbs sound, vibration, and radiation. Lead has no natural resonance frequencies; as a result, sheet-lead is used as a sound deadening layer in the walls, floors, and ceilings of sound studios. Organ pipes are often made from a lead alloy, mixed with various amounts of tin to control the tone of each pipe. Lead is an established shielding material from radiation in nuclear science and in X-ray rooms due to its denseness and high attenuation coefficient. Molten lead has been used as a coolant for lead-cooled fast reactors.

The largest use of lead in the early 21st century is in lead–acid batteries. The lead in batteries undergoes no direct contact with humans, so there are fewer toxicity concerns. People who work in battery production plants may be exposed to lead dust and inhale it. The reactions in the battery between lead, lead dioxide, and sulfuric acid provide a reliable source of voltage. Supercapacitors incorporating lead–acid batteries have been installed in kilowatt and megawatt scale applications in Australia, Japan, and the United States in frequency regulation, solar smoothing and shifting, wind smoothing, and other applications. These batteries have lower energy density and charge-discharge efficiency than lithium-ion batteries, but are significantly cheaper.

Lead glass

 

Lead is used in high voltage power cables as sheathing material to prevent water diffusion into insulation; this use is decreasing as lead is being phased out. Its use in solder for electronics is also being phased out by some countries to reduce the amount of environmentally hazardous waste. Lead is one of three metals used in the Oddy test for museum materials, helping detect organic acids, aldehydes, and acidic gases.

Compounds

In addition to being the main application for lead metal, lead-acid batteries are also the main consumer of lead compounds. The energy storage/release reaction used in these devices involves lead sulfate and lead dioxide:

Pb(s) + PbO
₂(s) + 2H
₂SO
₄(aq) → 2PbSO
₄(s) + 2H
₂O(l)

Other applications of lead compounds are very specialized and often fading. Lead-based coloring agents are used in ceramic glazes and glass, especially for red and yellow shades. While lead paints are phased out in Europe and North America, they remain in use in less developed countries such as China, India, or Indonesia. Lead tetraacetate and lead dioxide are used as oxidizing agents in organic chemistry. Lead is frequently used in the polyvinyl chloride coating of electrical cords. It can be used to treat candle wicks to ensure a longer, more even burn. Because of its toxicity, European and North American manufacturers use alternatives such as zinc. Lead glass is composed of 12–28% lead oxide, changing its optical characteristics and reducing the transmission of ionizing radiation. Lead-based semiconductors such as lead telluride and lead selenide are used in photovoltaic cells and infrared detectors.

Biological effects

Lead

Hazards
GHS pictograms
GHS signal word	Danger
GHS hazard statements	H302, H332, H351, H360Df, H373, H410
GHS precautionary statements	P201, P261, P273, P304, P340, P312, P308, P313, P391
NFPA 704	0 1 0

Lead has no confirmed biological role, and there is no confirmed safe level of lead exposure. A 2009 Canadian–American study concluded that even at levels that are considered to pose little to no risk, lead may cause "adverse mental health outcomes". Its prevalence in the human body—at an adult average of 120 mg—is nevertheless exceeded only by zinc (2500 mg) and iron (4000 mg) among the heavy metals. Lead salts are very efficiently absorbed by the body. A small amount of lead (1%) is stored in bones; the rest is excreted in urine and feces within a few weeks of exposure. Only about a third of lead is excreted by a child. Continual exposure may result in the bioaccumulation of lead.

Toxicity

Lead is a highly poisonous metal (whether inhaled or swallowed), affecting almost every organ and system in the human body. At airborne levels of 100 mg/m³, it is immediately dangerous to life and health. Most ingested lead is absorbed into the bloodstream. The primary cause of its 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, or mimicking and displacing other metals which act as cofactors in many enzymatic reactions. Among the essential metals that lead interacts with are calcium, iron, and zinc. High levels of calcium and iron tend to provide some protection from lead poisoning; low levels cause increased susceptibility.

Effects

Lead can cause severe damage to the brain and kidneys and, ultimately, death. By mimicking calcium, lead can cross the blood–brain barrier. It degrades the myelin sheaths of neurons, reduces their numbers, interferes with neurotransmission routes, and decreases neuronal growth. In the human body, lead inhibits porphobilinogen synthase and ferrochelatase, preventing both porphobilinogen formation and the incorporation of iron into protoporphyrin IX, the final step in heme synthesis. This causes ineffective heme synthesis and microcytic anemia.

Symptoms of lead poisoning

 

Symptoms of lead poisoning include nephropathy, colic-like abdominal pains, and possibly weakness in the fingers, wrists, or ankles. Small blood pressure increases, particularly in middle-aged and older people, may be apparent and can cause anemia. Several studies, mostly cross-sectional, found an association between increased lead exposure and decreased heart rate variability. In pregnant women, high levels of exposure to lead may cause miscarriage. Chronic, high-level exposure has been shown to reduce fertility in males.

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. Early childhood exposure has been linked with an increased risk of sleep disturbances and excessive daytime drowsiness in later childhood. High blood levels are associated with delayed puberty in girls. The rise and fall in exposure to airborne lead from the combustion of tetraethyl lead in gasoline during the 20th century has been linked with historical increases and decreases in crime levels, a hypothesis which is not universally accepted.

Exposure sources

Lead exposure is a global issue since lead mining and smelting, and battery manufacturing/disposal/recycling, are common in many countries. Lead enters the body via inhalation, ingestion, or skin absorption. Almost all inhaled lead is absorbed into the body; for ingestion, the rate is 20–70%, with children absorbing a higher percentage than adults.

Poisoning typically results from ingestion of food or water contaminated with lead, and less commonly after accidental ingestion of contaminated soil, dust, or lead-based paint. Seawater products can contain lead if affected by nearby industrial waters. Fruit and vegetables can be 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.

The use of lead for water pipes is problematic in areas with soft or acidic water. Hard water forms insoluble layers in the pipes whereas soft and acidic water dissolves the lead pipes. Dissolved carbon dioxide in the carried water may result in the formation of soluble lead bicarbonate; oxygenated water may similarly dissolve lead as lead(II) hydroxide. Drinking such water, over time, can cause health problems due to the toxicity of the dissolved lead. The harder the water the more calcium bicarbonate and sulfate it will contain, and the more the inside of the pipes will be coated with a protective layer of lead carbonate or lead sulfate.

Kymographic recording of the effect of lead acetate on frog heart experimental set up.

 

Ingestion of applied lead-based paint is the major source of exposure for children: a direct source is chewing on old painted window sills. Alternatively, as the applied dry paint deteriorates, it peels, is pulverized into dust and then enters the body through hand-to-mouth contact or contaminated food, water, or alcohol. Ingesting certain home remedies may result in exposure to lead or its compounds.

Inhalation is the second major exposure pathway, affecting smokers and especially workers in lead-related occupations. Cigarette smoke contains, among other toxic substances, radioactive lead-210.

Skin exposure may be significant for people working with organic lead compounds. The rate of skin absorption is lower for inorganic lead.

Treatment

Treatment for lead poisoning normally involves the administration of dimercaprol and succimer. Acute cases may require the use of disodium calcium edetate, the calcium chelate, and 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.

Environmental effects

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

 

The extraction, production, use, and disposal of lead and its products have caused significant contamination of the Earth's soils and waters. Atmospheric emissions of lead were at their peak during the Industrial Revolution, and the leaded gasoline period in the second half of the twentieth century. Lead releases originate from natural sources (i.e., concentration of the naturally occurring lead), industrial production, incineration and recycling, and mobilization of previously buried lead. Elevated concentrations of lead persist in soils and sediments in post-industrial and urban areas; industrial emissions, including those arising from coal burning, continue in many parts of the world, particularly in the developing countries.

Lead can accumulate in soils, especially those with a high organic content, where it remains for hundreds to thousands of years. Environmental lead can compete with other metals found in and on plants surfaces potentially inhibiting photosynthesis and at high enough concentrations, negatively affecting plant growth and survival. Contamination of soils and plants can allow lead to ascend the food chain affecting microorganisms and animals. In animals, lead exhibits toxicity in many organs, damaging the nervous, renal, reproductive, hematopoietic, and cardiovascular systems after ingestion, inhalation, or skin absorption. Fish uptake lead from both water and sediment; bioaccumulation in the food chain poses a hazard to fish, birds, and sea mammals.

Anthropogenic lead includes lead from shot and sinkers. These are among the most potent sources of lead contamination along with lead production sites. Lead was banned for shot and sinkers in the United States in 2017, although that ban was only effective for a month, and a similar ban is being considered in the European Union.

Analytical methods for the determination of lead in the environment include spectrophotometry, X-ray fluorescence, atomic spectroscopy and electrochemical methods. A specific ion-selective electrode has been developed based on the ionophore S,S'-methylenebis(N,N-diisobutyldithiocarbamate). An important biomarker assay for lead poisoning is δ-aminolevulinic acid levels in plasma, serum, and urine.

Restriction and remediation

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

 

By the mid-1980s, there was significant decline in the use of lead in industry. In the United States, environmental regulations reduced or eliminated the use of lead in non-battery products, including gasoline, paints, solders, and water systems. Particulate control devices were installed in coal-fired power plants to capture lead emissions. In 1992, U.S. Congress required the Environmental Protection Agency to reduce the blood lead levels of the country's children. Lead use was further curtailed by the European Union's 2003 Restriction of Hazardous Substances Directive. A large drop in lead deposition occurred in the Netherlands after the 1993 national ban on use of lead shot for hunting and sport shooting: from 230 tonnes in 1990 to 47.5 tonnes in 1995.

In the United States, the permissible exposure limit for lead in the workplace, comprising metallic lead, inorganic lead compounds, and lead soaps, was set at 50 μg/m³ over an 8-hour workday, and the blood lead level limit at 5 μg per 100 g of blood in 2012. Lead may still be found in harmful quantities in stoneware, vinyl (such as that used for tubing and the insulation of electrical cords), and Chinese brass. Old houses may still contain lead paint. White lead paint has been withdrawn from sale in industrialized countries, but specialized uses of other pigments such as yellow lead chromate remain. Stripping old paint by sanding produces dust which can be inhaled. Lead abatement programs have been mandated by some authorities in properties where young children live.

Lead waste, depending on the jurisdiction and the nature of the waste, may be treated as household waste (in order to facilitate lead abatement activities), or potentially hazardous waste requiring specialized treatment or storage. Lead is released to the wildlife in shooting places and a number of lead management practices, such as stewardship of the environment and reduced public scrutiny, have been developed to counter the lead contamination. Lead migration can be enhanced in acidic soils; to counter that, it is advised soils be treated with lime to neutralize the soils and prevent leaching of lead.

Research has been conducted on how to remove lead from biosystems by biological means: Fish bones are being researched for their ability to bioremediate lead in contaminated soil. The fungus Aspergillus versicolor is effective at absorbing lead ions from industrial waste before being released to water bodies. Several bacteria have been researched for their ability to remove lead from the environment, including the sulfate-reducing bacteria Desulfovibrio and Desulfotomaculum, both of which are highly effective in aqueous solutions.

Organic chemistry

From Wikipedia, the free encyclopedia

Methane, CH₄; the line-angle structural formula shows four carbon-hydrogen single bonds (σ, in black), and the typical 3D shape of tetrahedral molecules, with ~109° interior bond angles (in dashed-green).

 

Organic chemistry is a subdiscipline of chemistry that studies the structure, properties and reactions of organic compounds, which contain carbon in covalent bonding. Study of structure determines their chemical composition and formula. Study of properties includes physical and chemical properties, and evaluation of chemical reactivity to understand their behavior. The study of organic reactions includes the chemical synthesis of natural products, drugs, and polymers, and study of individual organic molecules in the laboratory and via theoretical (in silico) study. 

The range of chemicals studied in organic chemistry includes hydrocarbons (compounds containing only carbon and hydrogen) as well as compounds based on carbon, but also containing other elements, especially oxygen, nitrogen, sulfur, phosphorus (included in many biochemicals) and the halogens. Organometallic chemistry is the study of compounds containing carbon–metal bonds.

In addition, contemporary research focuses on organic chemistry involving other organometallics including the lanthanides, but especially the transition metals zinc, copper, palladium, nickel, cobalt, titanium and chromium. 

Line-angle representation

 

Ball-and-stick representation

 

Space-filling representation

 

Three representations of an organic compound, 5α-Dihydroprogesterone (5α-DHP), a steroid hormone. For molecules showing color, the carbon atoms are in black, hydrogens in gray, and oxygens in red. In the line angle representation, carbon atoms are implied at every terminus of a line and vertex of multiple lines, and hydrogen atoms are implied to fill the remaining needed valences (up to 4).

 

Organic compounds form the basis of all earthly life and constitute the majority of known chemicals. The bonding patterns of carbon, with its valence of four—formal single, double, and triple bonds, plus structures with delocalized electrons—make the array of organic compounds structurally diverse, and their range of applications enormous. They form the basis of, or are constituents of, many commercial products including pharmaceuticals; petrochemicals and agrichemicals, and products made from them including lubricants, solvents; plastics; fuels and explosives. The study of organic chemistry overlaps organometallic chemistry and biochemistry, but also with medicinal chemistry, polymer chemistry, and materials science.

History

Friedrich Wöhler

 

Before the nineteenth century, chemists generally believed that compounds obtained from living organisms were endowed with a vital force that distinguished them from inorganic compounds. According to the concept of vitalism (vital force theory), organic matter was endowed with a "vital force". During the first half of the nineteenth century, some of the first systematic studies of organic compounds were reported. Around 1816 Michel Chevreul started a study of soaps made from various fats and alkalis. He separated the different acids that, in combination with the alkali, produced the soap. Since these were all individual compounds, he demonstrated that it was possible to make a chemical change in various fats (which traditionally come from organic sources), producing new compounds, without "vital force". In 1828 Friedrich Wöhler produced the organic chemical urea (carbamide), a constituent of urine, from inorganic starting materials (the salts potassium cyanate and ammonium sulfate), in what is now called the Wöhler synthesis. Although Wöhler himself was cautious about claiming he had disproved vitalism, this was the first time a substance thought to be organic was synthesized in the laboratory without biological (organic) starting materials. The event is now generally accepted as indeed disproving the doctrine of vitalism.

In 1856 William Henry Perkin, while trying to manufacture quinine accidentally produced the organic dye now known as Perkin's mauve. His discovery, made widely known through its financial success, greatly increased interest in organic chemistry.

A crucial breakthrough for organic chemistry was the concept of chemical structure, developed independently in 1858 by both Friedrich August Kekulé and Archibald Scott Couper. Both researchers suggested that tetravalent carbon atoms could link to each other to form a carbon lattice, and that the detailed patterns of atomic bonding could be discerned by skillful interpretations of appropriate chemical reactions.

The era of the pharmaceutical industry began in the last decade of the 19th century when the manufacturing of acetylsalicylic acid—more commonly referred to as aspirin—in Germany was started by Bayer. By 1910 Paul Ehrlich and his laboratory group began developing arsenic-based arsphenamine, (Salvarsan), as the first effective medicinal treatment of syphilis, and thereby initiated the medical practice of chemotherapy. Ehrlich popularized the concepts of "magic bullet" drugs and of systematically improving drug therapies. His laboratory made decisive contributions to developing antiserum for diphtheria and standardizing therapeutic serums.

An example of an organometallic molecule, a catalyst called Grubbs' catalyst. Its formula is often given as RuCl₂(PCy₃)₂(=CHPh), where the ball-and-stick model is based on X-ray crystallography. The single metal atom ruthenium (Ru), (in turquoise), is at the very center of the structure; two chlorines (green), are bonded to the ruthenium atom—carbon atoms are black, hydrogens gray-white, and phosphorus orange. A phosphorus-ligand bond, tricyclohexyl phosphine, PCy, is below center; (another PCy ligand appears at the top of the image where its rings are obscuring one another). The ring group projecting to the right, an alkylidene, contains a metal-carbon double bond to ruthenium.

 

Early examples of organic reactions and applications were often found because of a combination of luck and preparation for unexpected observations. The latter half of the 19th century however witnessed systematic studies of organic compounds. The development of synthetic indigo is illustrative. The production of indigo from plant sources dropped from 19,000 tons in 1897 to 1,000 tons by 1914 thanks to the synthetic methods developed by Adolf von Baeyer. In 2002, 17,000 tons of synthetic indigo were produced from petrochemicals.

In the early part of the 20th century, polymers and enzymes were shown to be large organic molecules, and petroleum was shown to be of biological origin.

The multiple-step synthesis of complex organic compounds is called total synthesis. Total synthesis of complex natural compounds increased in complexity to glucose and terpineol. For example, cholesterol-related compounds have opened ways to synthesize complex human hormones and their modified derivatives. Since the start of the 20th century, complexity of total syntheses has been increased to include molecules of high complexity such as lysergic acid and vitamin B₁₂.

The total synthesis of vitamin B₁₂ marked a major achievement in organic chemistry.

 

The discovery of petroleum and the development of the petrochemical industry spurred the development of organic chemistry. Converting individual petroleum compounds into different types of compounds by various chemical processes led to organic reactions enabling a broad range of industrial and commercial products including, among (many) others: plastics, synthetic rubber, organic adhesives, and various property-modifying petroleum additives and catalysts. 

The majority of chemical compounds occurring in biological organisms are in fact carbon compounds, so the association between organic chemistry and biochemistry is so close that biochemistry might be regarded as in essence a branch of organic chemistry. Although the history of biochemistry might be taken to span some four centuries, fundamental understanding of the field only began to develop in the late 19th century and the actual term biochemistry was coined around the start of 20th century. Research in the field increased throughout the twentieth century, without any indication of slackening in the rate of increase, as may be verified by inspection of abstraction and indexing services such as BIOSIS Previews and Biological Abstracts, which began in the 1920s as a single annual volume, but has grown so drastically that by the end of the 20th century it was only available to the everyday user as an online electronic database.

Characterization

Since organic compounds often exist as mixtures, a variety of techniques have also been developed to assess purity, especially important being chromatography techniques such as HPLC and gas chromatography. Traditional methods of separation include distillation, crystallization, and solvent extraction. 

Organic compounds were traditionally characterized by a variety of chemical tests, called "wet methods", but such tests have been largely displaced by spectroscopic or other computer-intensive methods of analysis. Listed in approximate order of utility, the chief analytical methods are:

• Nuclear magnetic resonance (NMR) spectroscopy is the most commonly used technique, often permitting complete assignment of atom connectivity and even stereochemistry using correlation spectroscopy. The principal constituent atoms of organic chemistry – hydrogen and carbon – exist naturally with NMR-responsive isotopes, respectively ¹H and ¹³C.
• Elemental analysis: A destructive method used to determine the elemental composition of a molecule. See also mass spectrometry, below.
• Mass spectrometry indicates the molecular weight of a compound and, from the fragmentation patterns, its structure. High resolution mass spectrometry can usually identify the exact formula of a compound and is used in lieu of elemental analysis. In former times, mass spectrometry was restricted to neutral molecules exhibiting some volatility, but advanced ionization techniques allow one to obtain the "mass spec" of virtually any organic compound.
• Crystallography can be useful for determining molecular geometry when a single crystal of the material is available. Highly efficient hardware and software allows a structure to be determined within hours of obtaining a suitable crystal.

Traditional spectroscopic methods such as infrared spectroscopy, optical rotation, and UV/VIS spectroscopy provide relatively nonspecific structural information but remain in use for specific classes of compounds. Traditionally refractive index and density were also important for substance identification.

Properties

Physical properties of organic compounds typically of interest include both quantitative and qualitative features. Quantitative information includes melting point, boiling point, and index of refraction. Qualitative properties include odor, consistency, solubility, and color.

Melting and boiling properties

Organic compounds typically melt and many boil. In contrast, while inorganic materials generally can be melted, many do not boil, tending instead to degrade. In earlier times, the melting point (m.p.) and boiling point (b.p.) provided crucial information on the purity and identity of organic compounds. The melting and boiling points correlate with the polarity of the molecules and their molecular weight. Some organic compounds, especially symmetrical ones, sublime, that is they evaporate without melting. A well-known example of a sublimable organic compound is para-dichlorobenzene, the odiferous constituent of modern mothballs. Organic compounds are usually not very stable at temperatures above 300 °C, although some exceptions exist.

Solubility

Neutral organic compounds tend to be hydrophobic; that is, they are less soluble in water than in organic solvents. Exceptions include organic compounds that contain ionizable (which can be converted in ions) groups as well as low molecular weight alcohols, amines, and carboxylic acids where hydrogen bonding occurs. Organic compounds tend to dissolve in organic solvents. Solvents can be either pure substances like ether or ethyl alcohol, or mixtures, such as the paraffinic solvents such as the various petroleum ethers and white spirits, or the range of pure or mixed aromatic solvents obtained from petroleum or tar fractions by physical separation or by chemical conversion. Solubility in the different solvents depends upon the solvent type and on the functional groups if present in the solution.

Solid state properties

Various specialized properties of molecular crystals and organic polymers with conjugated systems are of interest depending on applications, e.g. thermo-mechanical and electro-mechanical such as piezoelectricity, electrical conductivity, and electro-optical (e.g. non-linear optics) properties. For historical reasons, such properties are mainly the subjects of the areas of polymer science and materials science.

Nomenclature

Various names and depictions for one organic compound.

 

The names of organic compounds are either systematic, following logically from a set of rules, or nonsystematic, following various traditions. Systematic nomenclature is stipulated by specifications from IUPAC. Systematic nomenclature starts with the name for a parent structure within the molecule of interest. This parent name is then modified by prefixes, suffixes, and numbers to unambiguously convey the structure. Given that millions of organic compounds are known, rigorous use of systematic names can be cumbersome. Thus, IUPAC recommendations are more closely followed for simple compounds, but not complex molecules. To use the systematic naming, one must know the structures and names of the parent structures. Parent structures include unsubstituted hydrocarbons, heterocycles, and monofunctionalized derivatives thereof. 

Nonsystematic nomenclature is simpler and unambiguous, at least to organic chemists. Nonsystematic names do not indicate the structure of the compound. They are common for complex molecules, which includes most natural products. Thus, the informally named lysergic acid diethylamide is systematically named (6aR,9R)-N,N-diethyl-7-methyl-4,6,6a,7,8,9-hexahydroindolo-[4,3-fg] quinoline-9-carboxamide.

With the increased use of computing, other naming methods have evolved that are intended to be interpreted by machines. Two popular formats are SMILES and InChI.

Structural drawings

Organic molecules are described more commonly by drawings or structural formulas, combinations of drawings and chemical symbols. The line-angle formula is simple and unambiguous. In this system, the endpoints and intersections of each line represent one carbon, and hydrogen atoms can either be notated explicitly or assumed to be present as implied by tetravalent carbon. 

History

By 1880 an explosion in the number of chemical compounds being discovered occurred assisted by new synthetic and analytical techniques. Grignard described the situation as "chaos le plus complet" as due to the lack of convention it was possible to have multiple names for the same compound. This led to the creation of the Geneva rules in 1892.

Classification of organic compounds

Functional groups

The family of carboxylic acids contains a carboxyl (-COOH) functional group. Acetic acid, shown here, is an example.

 

The concept of functional groups is central in organic chemistry, both as a means to classify structures and for predicting properties. A functional group is a molecular module, and the reactivity of that functional group is assumed, within limits, to be the same in a variety of molecules. Functional groups can have decisive influence on the chemical and physical properties of organic compounds. Molecules are classified on the basis of their functional groups. Alcohols, for example, all have the subunit C-O-H. All alcohols tend to be somewhat hydrophilic, usually form esters, and usually can be converted to the corresponding halides. Most functional groups feature heteroatoms (atoms other than C and H). Organic compounds are classified according to functional groups, alcohols, carboxylic acids, amines, etc.

Aliphatic compounds

The aliphatic hydrocarbons are subdivided into three groups of homologous series according to their state of saturation:

• Alkanes (paraffins): aliphatic hydrocarbons without any double or triple bonds, i.e. just C-C, C-H single bonds
• Alkenes (olefins): aliphatic hydrocarbons which contain one or more double bonds, i.e. di-olefins (dienes) or poly-olefins.
• Alkynes (acetylenes): aliphatic hydrocarbons which have one or more triple bonds.

The rest of the group is classed according to the functional groups present. Such compounds can be "straight-chain", branched-chain or cyclic. The degree of branching affects characteristics, such as the octane number or cetane number in petroleum chemistry. 

Both saturated (alicyclic) compounds and unsaturated compounds exist as cyclic derivatives. The most stable rings contain five or six carbon atoms, but large rings (macrocycles) and smaller rings are common. The smallest cycloalkane family is the three-membered cyclopropane ((CH₂)₃). Saturated cyclic compounds contain single bonds only, whereas aromatic rings have an alternating (or conjugated) double bond. Cycloalkanes do not contain multiple bonds, whereas the cycloalkenes and the cycloalkynes do.

Aromatic compounds

Benzene is one of the best-known aromatic compounds as it is one of the simplest and most stable aromatics.

 

Aromatic hydrocarbons contain conjugated double bonds. This means that every carbon atom in the ring is sp2 hybridized, allowing for added stability. The most important example is benzene, the structure of which was formulated by Kekulé who first proposed the delocalization or resonance principle for explaining its structure. For "conventional" cyclic compounds, aromaticity is conferred by the presence of 4n + 2 delocalized pi electrons, where n is an integer. Particular instability (antiaromaticity) is conferred by the presence of 4n conjugated pi electrons.

Heterocyclic compounds

The characteristics of the cyclic hydrocarbons are again altered if heteroatoms are present, which can exist as either substituents attached externally to the ring (exocyclic) or as a member of the ring itself (endocyclic). In the case of the latter, the ring is termed a heterocycle. Pyridine and furan are examples of aromatic heterocycles while piperidine and tetrahydrofuran are the corresponding alicyclic heterocycles. The heteroatom of heterocyclic molecules is generally oxygen, sulfur, or nitrogen, with the latter being particularly common in biochemical systems. 

Heterocycles are commonly found in a wide range of products including aniline dyes and medicines. Additionally, they are prevalent in a wide range of biochemical compounds such as alkaloids, vitamins, steroids, and nucleic acids (e.g. DNA, RNA). 

Rings can fuse with other rings on an edge to give polycyclic compounds. The purine nucleoside bases are notable polycyclic aromatic heterocycles. Rings can also fuse on a "corner" such that one atom (almost always carbon) has two bonds going to one ring and two to another. Such compounds are termed spiro and are important in a number of natural products.

Polymers

This swimming board is made of polystyrene, an example of a polymer.

 

One important property of carbon is that it readily forms chains, or networks, that are linked by carbon-carbon (carbon-to-carbon) bonds. The linking process is called polymerization, while the chains, or networks, are called polymers. The source compound is called a monomer.

Two main groups of polymers exist: synthetic polymers and biopolymers. Synthetic polymers are artificially manufactured, and are commonly referred to as industrial polymers. Biopolymers occur within a respectfully natural environment, or without human intervention. 

Since the invention of the first synthetic polymer product, bakelite, synthetic polymer products have frequently been invented.

Common synthetic organic polymers are polyethylene (polythene), polypropylene, nylon, teflon (PTFE), polystyrene, polyesters, polymethylmethacrylate (called perspex and plexiglas), and polyvinylchloride (PVC).

Both synthetic and natural rubber are polymers.

Varieties of each synthetic polymer product may exist, for purposes of a specific use. Changing the conditions of polymerization alters the chemical composition of the product and its properties. These alterations include the chain length, or branching, or the tacticity.

With a single monomer as a start, the product is a homopolymer.

Secondary component(s) may be added to create a heteropolymer (co-polymer) and the degree of clustering of the different components can also be controlled.

Physical characteristics, such as hardness, density, mechanical or tensile strength, abrasion resistance, heat resistance, transparency, colour, etc. will depend on the final composition.

Biomolecules

Maitotoxin, a complex organic biological toxin.

 

Biomolecular chemistry is a major category within organic chemistry which is frequently studied by biochemists. Many complex multi-functional group molecules are important in living organisms. Some are long-chain biopolymers, and these include peptides, DNA, RNA and the polysaccharides such as starches in animals and celluloses in plants. The other main classes are amino acids (monomer building blocks of peptides and proteins), carbohydrates (which includes the polysaccharides), the nucleic acids (which include DNA and RNA as polymers), and the lipids. In addition, animal biochemistry contains many small molecule intermediates which assist in energy production through the Krebs cycle, and produces isoprene, the most common hydrocarbon in animals. Isoprenes in animals form the important steroid structural (cholesterol) and steroid hormone compounds; and in plants form terpenes, terpenoids, some alkaloids, and a class of hydrocarbons called biopolymer polyisoprenoids present in the latex of various species of plants, which is the basis for making rubber.

Small molecules

Molecular models of caffeine.

 

In pharmacology, an important group of organic compounds is small molecules, also referred to as 'small organic compounds'. In this context, a small molecule is a small organic compound that is biologically active, but is not a polymer. In practice, small molecules have a molar mass less than approximately 1000 g/mol.

Fullerenes

Fullerenes and carbon nanotubes, carbon compounds with spheroidal and tubular structures, have stimulated much research into the related field of materials science. The first fullerene was discovered in 1985 by Sir Harold W. Kroto of the United Kingdom and by Richard E. Smalley and Robert F. Curl, Jr., of the United States. Using a laser to vaporize graphite rods in an atmosphere of helium gas, these chemists and their assistants obtained cagelike molecules composed of 60 carbon atoms (C60) joined together by single and double bonds to form a hollow sphere with 12 pentagonal and 20 hexagonal faces—a design that resembles a football, or soccer ball. In 1996 the trio was awarded the Nobel Prize for their pioneering efforts. The C60 molecule was named buckminsterfullerene (or, more simply, the buckyball) after the American architect R. Buckminster Fuller, whose geodesic dome is constructed on the same structural principles.

Others

Organic compounds containing bonds of carbon to nitrogen, oxygen and the halogens are not normally grouped separately. Others are sometimes put into major groups within organic chemistry and discussed under titles such as organosulfur chemistry, organometallic chemistry, organophosphorus chemistry and organosilicon chemistry.

Organic reactions

Organic reactions are chemical reactions involving organic compounds. Many of these reactions are associated with functional groups. The general theory of these reactions involves careful analysis of such properties as the electron affinity of key atoms, bond strengths and steric hindrance. These factors can determine the relative stability of short-lived reactive intermediates, which usually directly determine the path of the reaction. 

The basic reaction types are: addition reactions, elimination reactions, substitution reactions, pericyclic reactions, rearrangement reactions and redox reactions. An example of a common reaction is a substitution reaction written as:

Nu⁻ + C-X → C-Nu + X⁻

where X is some functional group and Nu is a nucleophile. 

The number of possible organic reactions is basically infinite. However, certain general patterns are observed that can be used to describe many common or useful reactions. Each reaction has a stepwise reaction mechanism that explains how it happens in sequence—although the detailed description of steps is not always clear from a list of reactants alone. 

The stepwise course of any given reaction mechanism can be represented using arrow pushing techniques in which curved arrows are used to track the movement of electrons as starting materials transition through intermediates to final products.

Organic synthesis

A synthesis designed by E.J. Corey for oseltamivir (Tamiflu). This synthesis has 11 distinct reactions.

 

Synthetic organic chemistry is an applied science as it borders engineering, the "design, analysis, and/or construction of works for practical purposes". Organic synthesis of a novel compound is a problem solving task, where a synthesis is designed for a target molecule by selecting optimal reactions from optimal starting materials. Complex compounds can have tens of reaction steps that sequentially build the desired molecule. The synthesis proceeds by utilizing the reactivity of the functional groups in the molecule. For example, a carbonyl compound can be used as a nucleophile by converting it into an enolate, or as an electrophile; the combination of the two is called the aldol reaction. Designing practically useful syntheses always requires conducting the actual synthesis in the laboratory. The scientific practice of creating novel synthetic routes for complex molecules is called total synthesis. 

Strategies to design a synthesis include retrosynthesis, popularized by E.J. Corey, which starts with the target molecule and splices it to pieces according to known reactions. The pieces, or the proposed precursors, receive the same treatment, until available and ideally inexpensive starting materials are reached. Then, the retrosynthesis is written in the opposite direction to give the synthesis. A "synthetic tree" can be constructed, because each compound and also each precursor has multiple syntheses.

Organosilicon

From Wikipedia, the free encyclopedia

A carbon–silicon bond present in all organosilicon compounds

 

Organosilicon compounds are organometallic compounds containing carbon–silicon bonds. Organosilicon chemistry is the corresponding science of their preparation and properties. Most organosilicon compounds are similar to the ordinary organic compounds, being colourless, flammable, hydrophobic, and stable to air. Silicon carbide is an inorganic compound.

Occurrence and applications

Organosilicon compounds are widely encountered in commercial products. Most common are sealants, caulks, adhesives, and coatings made from silicones. Others important uses include synthesis of polyhedral oligomeric silsesquioxanes, agricultural and plant control adjuvants commonly used in conjunction with herbicides and fungicides.

Silicone caulk, commercial sealants, are mainly composed of organosilicon compounds.

 

Polydimethylsiloxane (PDMS) is the principal component of silicones.

Biology and medicine

Carbon–silicon bonds are absent in biology. Silicates, on the other hand, have known existence in diatoms. Silafluofen is an organosilicon compound that functions as a pyrethroid insecticide. Several organosilicon compounds have been investigated as pharmaceuticals.

Properties of Si–C, Si–O, and Si–F bonds

In most organosilicon compounds, Si is tetravalent with tetrahedral molecular geometry. Carbon–silicon bonds compared to carbon–carbon bonds are longer (186 pm vs. 154 pm) and weaker with bond dissociation energy 451 kJ/mol vs. 607 kJ/mol. The C–Si bond is somewhat polarised towards carbon due to carbon's greater electronegativity (C 2.55 vs Si 1.90). The Si–C bond can be broken more readily than typical C–C bonds. One manifestation of bond polarization in organosilanes is found in the Sakurai reaction. Certain alkyl silanes can be oxidized to an alcohol in the Fleming–Tamao oxidation. 

Another manifestation is the β-silicon effect describes the stabilizing effect of a β-silicon atom on a carbocation with many implications for reactivity. 

Si–O bonds are much stronger (809 kJ/mol compared to 538 kJ/mol) than a typical C–O single bond. The favorable formation of Si–O bonds drives many organic reactions such as the Brook rearrangement and Peterson olefination. Compared to the strong Si–O bond, the Si–F bond is even stronger.

Preparation

The first organosilicon compound, tetraethylsilane, was prepared by Charles Friedel and James Crafts in 1863 by reaction of tetrachlorosilane with diethylzinc. 

The bulk of organosilicon compounds derive from organosilicon chlorides (CH₃)_4-xSiCl_x. These chlorides are produced by the "Direct process", which entails the reaction of methyl chloride with a silicon-copper alloy. The main and most sought-after product is dimethyldichlorosilane:

2 CH₃Cl + Si → (CH₃)₂SiCl₂

A variety of other products are obtained, including trimethylsilyl chloride and methyltrichlorosilane. About 1 million tons of organosilicon compounds are prepared annually by this route. The method can also be used for phenyl chlorosilanes.

Hydrosilylation

After the Direct Process, a second major method for the formation of Si-C bonds is hydrosilylation (also called hydrosilation). In this process, compounds with Si-H bonds (hydrosilanes) add to unsaturated substrates. Commercially, the main substrates are alkenes. Other unsaturated functional groups—alkynes, imines, ketones, and aldehydes. An example is the hydrosilation of phenylacetylene:

Hydrosilylation requires metal catalysts, especially those based on platinum group metals.

In the related silylmetalation, a metal replaces the hydrogen atom.

Functional groups

Silicon is a component of many functional groups. Most of these are analogous to organic compounds. The overarching exception is the rarity of multiple bonds to silicon, as reflected in the double bond rule.

Silanols, siloxides, and siloxanes

Silanols are analogues of alcohols. They are generally prepared by hydrolysis of silyl chlorides:

R₃SiCl + H₂O → R₃SiOH + HCl

Less frequently silanols are prepared by oxidation of silyl hydrides, a reaction that uses a metal catalyst:

2 R₃SiH + O₂ → 2 R₃SiOH

Many silanols have been isolated including (CH₃)₃SiOH and (C₆H₅)₃SiOH. They are about 500x more acidic than the corresponding alcohols. Siloxides are the deprotonated derivatives of silanols:

R₃SiOH + NaOH → R₃SiONa + H₂O

Silanols tend to dehydrate to give siloxanes:

2 R₃SiOH → R₃Si-O-SiR₃ + H₂O

Polymers with repeating siloxane linkages are called silicones. Compounds with an Si=O double bond called silanones are extremely unstable.

Silyl ethers

Silyl ethers have the connectivity Si-O-C. They are typically prepared by the reaction of alcohols with silyl chlorides:

(CH₃)₃SiCl + ROH → (CH₃)₃Si-O-R + HCl

Silyl ethers are extensively used as protective groups for alcohols. 

Exploiting the strength of the Si-F bond, fluoride sources such as tetra-n-butylammonium fluoride (TBAF) are used in deprotection of silyl ethers:

(CH₃)₃Si-O-R + F⁻ + H₂O → (CH₃)₃Si-F + H-O-R + OH⁻

Silyl chlorides

Organosilyl chlorides are important commodity chemicals. They are mainly used to produce silicone polymers as described above. Especially important silyl chlorides are dimethyldichlorosilane (Me₂SiCl₂), methyltrichlorosilane (MeSiCl₃), and trimethylsilyl chloride (Me₃SiCl). More specialized derivatives that find commercial applications include dichloromethylphenylsilane, trichloro(chloromethyl)silane, trichloro(dichlorophenyl)silane, trichloroethylsilane, and phenyltrichlorosilane. 

Although proportionately a minor outlet, organosilicon compounds are widely used in organic synthesis. Notably trimethylsilyl chloride Me₃SiCl is the main silylating agent. One classic method called the Flood reaction for the synthesis of this compound class is by heating hexaalkyldisiloxanes R₃SiOSiR₃ with concentrated sulfuric acid and a sodium halide.

Silyl hydrides

The silicon to hydrogen bond is longer than the C–H bond (148 compared to 105 pm) and weaker (299 compared to 338 kJ/mol). Hydrogen is more electronegative than silicon hence the naming convention of silyl hydrides. Commonly the presence of the hydride is not mentioned in the name of the compound. Triethylsilane has the formula Et₃SiH. Phenylsilane is PhSiH₃. The parent compound SiH₄ is called silane.

Silenes

Organosilicon compounds, unlike their carbon counterparts, do not have a rich double bond chemistry. Compounds with silene Si=C bonds (also known as alkylidenesilanes) are laboratory curiosities such as the silicon benzene analogue silabenzene. In 1967, Gusel'nikov and Flowers provided the first evidence for silenes from pyrolysis of dimethylsilacyclobutane. The first stable (kinetically shielded) silene was reported in 1981 by Brook.

Disilenes have Si=Si double bonds and disilynes are silicon analogues of an alkyne. The first Silyne (with a silicon to carbon triple bond) was reported in 2010.

Siloles

Chemical structure of silole

 

Siloles, also called silacyclopentadienes, are members of a larger class of compounds called metalloles. They are the silicon analogs of cyclopentadienes and are of current academic interest due to their electroluminescence and other electronic properties. Siloles are efficient in electron transport. They owe their low lying LUMO to a favorable interaction between the antibonding sigma silicon orbital with an antibonding pi orbital of the butadiene fragment.

Hypercoordinated silicon

Unlike carbon, silicon compounds can be coordinated to five atoms as well in a group of compounds ranging from so-called silatranes, such as phenylsilatrane, to a uniquely stable pentaorganosilicate:

The stability of hypervalent silicon is the basis of the Hiyama coupling, a coupling reaction used in certain specialized organic synthetic applications. The reaction begins with the activation of Si-C bond by fluoride:

R-SiR'₃ + R"-X + F⁻ → R-R" + R'₃SiF + X⁻

Various reactions

Certain allyl silanes can be prepared from allylic esters such as 1 and monosilylcopper compounds such as 2 in.

In this reaction type silicon polarity is reversed in a chemical bond with zinc and a formal allylic substitution on the benzoyloxy group takes place.

Environmental effects

Organosilicon compounds affect bee (and other insect) immune expression, making them more susceptible to viral infection.