Lead (symbol Pb, atomic number 82) is a dense, soft post-transition metal in group 14 of the periodic table. A heavy element, lead is solid at standard temperature and pressure (STP) and has a silvery-gray appearance that tarnishes to dull gray when exposed to air. Its atomic weight is about 207.2 u. Lead belongs to the carbon group but behaves differently from lighter congeners like carbon and silicon; it is classed as a poor metal or post-transition rather than a true transition metal. Common oxidation states of lead are +2 and +4 (in ionic compounds and complexes), with the +2 state by far the most stable due to the “inert pair” effect of the 6s electrons. Lead has no liquid phase at room conditions (it solidifies at 600.6 K) and is one of the few elements heavier than iron that has only stable isotopes at the heaviest end of the periodic table.

Key facts at a glance: Lead is exceptionally dense (11.34 g/cm³) and malleable. It conducts electricity fairly well (though not as well as copper or silver) and has a low melting point for a metal (600.6 K, about 327.5 °C), making it easy to cast and form. Lead remains largely unreactive in air or water at room temperature, thanks to a thin oxide layer that passivates its surface. Historically, lead has been mined and used extensively since ancient times, for example in plumbing (the word plumbing derives from the Latin plumbum for lead) and in construction, and later in printing and batteries. It is also notable for its extreme toxicity in biological systems, which has led to strict regulation of its use in paints, fuels, and other products.

Atomic Structure and Electron Configuration

Lead
Atomic number	82
Symbol	Pb
Group	14 (carbon group)
Electron configuration	[Xe] 4f14 5d10 6s2 6p2
Density	11.34 g/cm^3
Period	6
Main isotopes	206Pb, 207Pb, 208Pb
Melting point	327.5 °C
Block	p
Oxidation states	+2, +4
Wikidata	Q708

Lead atoms contain 82 protons and usually 82 electrons. The electron configuration of Pb in its ground state is [Xe] 4f^14 5d^10 6s^2 6p^2. This can be read as a closed-shell xenon core (54 electrons), filled 4f and 5d subshells, and two electrons in the 6s and two in the 6p orbitals. The electrons in the outer (6th) shell—that is, the 6s and 6p electrons—are considered the valence electrons, which participate in bonding and chemical reactivity. The filled 4f and 5d shells lie deeper inside and effectively shield the nucleus, contributing to lead’s relatively large atomic radius despite its high nuclear charge.

In the periodic table, lead resides in period 6 and group 14. As one moves down group 14 (carbon, silicon, germanium, tin, lead), atomic size increases and metallic character becomes more pronounced. Compared to tin (just above lead), lead’s valence electrons are held more loosely in a relativistic sense (the inner electrons move at speeds where relativity becomes significant), which contributes to the so-called inert pair effect. This effect helps explain why Pb(II) compounds are far more stable than Pb(IV) compounds under normal conditions.

Key atomic properties and periodic trends for lead: its atomic radius is on the order of 175 picometers (atomic radius) with a covalent radius around 146 pm. Lead’s Pauling electronegativity is about 2.33, placing it near the center of the scale (hydrogen is 2.20 for reference). The first ionization energy (energy to remove one electron) of lead is about 716 kJ/mol. These values are consistent with a relatively large, metallic atom with moderate attraction to its valence electrons.

Isotopes and Nuclear Properties

Naturally occurring lead consists of four stable isotopes: ^204Pb, ^206Pb, ^207Pb, and ^208Pb. Their approximate abundances in terrestrial lead are ~1.4%, 24.1%, 22.1%, and 52.4% respectively. Notably, three of these (^206Pb, ^207Pb, and ^208Pb) are the final decay products (endpoints) of long radioactive decay chains: uranium-238 decays to ^206Pb, uranium-235 to ^207Pb, and thorium-232 to ^208Pb. Thus these isotopes are both primordial (formed in supernovae) and radiogenic (formed by radioactive decay within the Earth). ^204Pb is unique as a primordial isotope not produced by radioactive decay, serving as a natural “baseline” in isotope geochemistry.

Lead has no long-lived radioactive isotopes of its own beyond trace amounts in decay series or short-lived transuranic products. The longest-lived radioactive lead nuclide is ^205Pb (not occurring naturally) with a half-life of ~17 million years. However, some relatively short-lived isotopes do appear in nature as decay intermediates. The most significant of these is ^210Pb (half-life ~22.3 years), which arises in the uranium-238 decay chain and commonly accumulates on Earth’s surface. Because ^210Pb decays moderately slowly, it is useful for dating environmental samples (for example, sediment layers or ice cores) on timescales of decades to centuries by measuring its ratio to stable lead isotopes.

In nuclear terms, lead-208 (^208Pb) is notable as a “doubly magic” nucleus (82 protons and 126 neutrons), which confers particularly high nuclear stability. The stable isotope ^207Pb (with spin 1/2) is sometimes used as a reference in nuclear magnetic resonance (NMR) studies of lead compounds. Overall, lead’s nuclear properties (stable heavy isotopes and well-understood decay series) make it useful in geochronology, particularly in uranium–lead and lead–lead dating methods for determining the age of rocks and minerals.

Allotropes and Typical Compounds

Metallic lead itself has only one normal form and no well-defined allotropes like carbon or phosphorus. In the solid state, pure lead crystallizes in a face-centered cubic (FCC) structure. It is very malleable, meaning it can be hammered or rolled into sheets. Freshly cut lead is bright and shiny, but it rapidly tarnishes in air as a thin oxide film forms, giving it a dull gray surface. There is no polymorphism or other allotrope of elemental lead known under ordinary conditions.

Lead’s characteristic compounds often reflect its two main oxidation states. In the +2 state (Pb(II) or plumbous), lead behaves much like a heavier analogue of tin (Sn(II)), forming relatively stable salts and complexes. In the +4 state (Pb(IV) or plumbic), lead forms compounds that are strong oxidizers and often unstable toward reduction back to +2. For example, lead(II) oxide (PbO) is a common yellow or red powder used in glassmaking and battery components, whereas lead(IV) oxide (PbO₂) is a dark brown solid known as a powerful oxidizing agent (it is, for instance, the active oxidant in lead–acid batteries). A well-known mixed-valence oxide is red lead (Pb₃O₄, sometimes written Pb^II₂Pb^IVO₄), historically used as pigment.

Common halides of lead include lead(II) chloride (PbCl₂), a white crystalline salt slightly soluble in hot water; lead(II) bromide and iodide (yellow solids); and lead(II) fluoride. Lead(IV) chloride (PbCl₄) is a volatile yellow liquid at room temperature, but it is so reactive that it decomposes on contact with water or even moist air, releasing chlorine gas and reducing to PbCl₂. Lead also forms a variety of oxyhalides and other oxycompounds (for example, lead(II) sulfate PbSO₄, lead carbonate PbCO₃ in mineral forms like cerussite, and lead nitrate Pb(NO₃)₂).

Lead hydride (PbH₄), also known as plumbane, is theoretically analogous to methane but is extremely unstable under normal conditions, decomposing spontaneously into lead and hydrogen. Organolead compounds were once important: for example, tetraethyl lead (Pb(C₂H₅)₄) was used as a gasoline anti-knock additive in the 20th century, and various lead-containing stabilizers were used in paints and plastics. These organolead compounds are quite toxic and most have been phased out due to safety concerns.

Chemically, lead(II) ions often exhibit a stereochemically active lone pair of electrons (the 6s^2 electrons), which can lead to distorted coordination geometries in complexes. Lead(II) salts typically form covalent bonds rather than purely ionic ones because of the polarizing power of Pb²⁺. Lead(IV) compounds, when stable, are strongly oxidizing: for instance, PbO₂ can oxidize chloride to chlorine. In many environments, however, Pb(IV) tends to reduce to Pb(II) and release electrons to any available acceptor.

Physical Properties

Lead is notable for its high density (11.34 g/cm³), which is about six times that of water and higher than most other common metals (denser than iron or copper, but slightly less dense than tungsten or gold). This high density makes lead good at blocking radiation (X-rays and gamma rays), since a dense material attenuates radiation effectively. It also gives lead a relatively large atomic volume and mass per atom.

At room temperature, lead is a shiny bluish-white metal (when freshly cut) that is soft and malleable – it can be dented with a fingernail and easily cut with a knife. Its bulk modulus (resistance to compression) and Young’s modulus (stiffness) are much lower than for harder metals like steel. Lead is nonmagnetic (it is diamagnetic) and conducts heat and electricity moderately well for a metal. Its electrical resistivity at 20 °C is about 192 nΩ·m (roughly 10 times higher resistivity than copper, making it a worse conductor than copper). The thermal conductivity of lead is also relatively low among metals.

Lead remains solid up to its melting point of 600.6 K. On heating, it expands and eventually melts into a silvery liquid, which then boils at a much higher temperature of 2022 K (about 1749 °C). Under normal pressures, lead has three solid phases: the ambient face-centered cubic form (called α-Pb), a hexagonal structure (β-Pb) that exists above about 4.2 GPa or under cryogenic cooling, and a tetragonal high-pressure form (γ-Pb). In everyday applications, only the common α (FCC) form is encountered.

Spectroscopically, lead atoms exhibit many lines in the ultraviolet region when excited, but it has relatively few visible emission lines. Analytical techniques cannot often rely on visible lead lines, so X-ray spectroscopy is sometimes used to detect lead (for instance, in X-ray fluorescence analysis, the characteristic L-series X-ray emission of lead is measured). The color or appearance of lead compounds can vary widely (for example, yellow in lead(II) iodide or white in lead carbonate) but these are due to electronic transitions in the compound, not the metal itself.

Thermally, lead has a low melting point and is easily fusible – in fact, a common lead alloy (lead–tin solder) melts at around 180–200 °C, well below the melting point of pure tin, allowing soldering. The thermal expansion coefficient of lead is relatively high, meaning it expands appreciably with temperature rise. Its heat capacity at room temperature is low enough that small amounts heat up quickly; the molar heat capacity is about 26 J/(mol·K) at 25 °C.

Chemical Reactivity and Trends

Lead metal is relatively unreactive compared to many other metals. At room temperature in dry air, it quickly forms a thin protective corrosion layer of lead(II) oxide (PbO) or basic lead carbonate (from carbon dioxide in the air), which prevents further reaction. This passivation explains why lead objects (like pipes or roofing) can survive for decades without significant corrosion. However, finely divided lead (powder) or heated lead reacts more readily with oxygen, forming thicker oxide layers or even melting to red lead (Pb₃O₄) at higher temperatures.

Put simply, lead does not dissolve in weak acids readily, because the acid attack is blocked by oxide or carbonate layers. In strong acids, especially oxidizing acids like nitric acid (HNO₃), lead will dissolve: for example, hot concentrated nitric acid oxidizes lead to lead nitrate, producing nitrogen dioxide gas. Non-oxidizing acids (such as hydrochloric acid, HCl) will react slowly with lead, producing soluble lead(II) chloride, but only if the surface is damaged or heated. Otherwise, dilute HCl has little effect on a lead surface because any HCl-produced lead(II) chloride tends to coat and protect the metal.

Lead is amphoteric in that its common oxidation state (+2) forms compounds that can react with both acids and bases. For instance, lead(II) oxide (PbO) reacts with acids to form salts (like lead nitrate) and also reacts with strong bases to form plumbite Pb(OH)₃]⁻ or [PbO₂ ]^2⁻ depending on conditions). A familiar example is lead(IV) oxide (PbO₂) dissolving in hot sodium hydroxide to give a sodium plumbate Na₂PbO₃] or [Na₂PbO₄.

In the metal reactivity series, lead lies below hydrogen (standard reduction potential of Pb²⁺/Pb is about –0.13 V), meaning lead does not spontaneously reduce water to hydrogen or dissolve in cold dilute acids without an oxidizing agent. This is why lead does not fizz or corrode like zinc or iron would in acid. However, lead is not immune: it can be dissolved by melting Zn(II) salts (forming Zn metal and dissolving Pb), or by reacting with molten metals like copper or tin to form alloys, which is how many lead-based alloys (such as solder) are made.

Lead(II) compounds are usually white or pale solids (lead sulfate, carbonate, or chloride), though they tend to darken on heating. Lead(IV) compounds, when they exist, are powerful oxidizers: PbO₂ reacts with HCl to produce chlorine gas, and potassium plumbate can, in principle, decompose water or oxidize organics. An example of redox trend in the lead family (carbon group) is that tin (Sn) also has +2 and +4 states, but Pb(IV) is less stable and more strongly oxidizing. In complexes, lead(II) ions often coordinate in asymmetric polyhedra (due to the stereochemically active lone pair), with coordination numbers from 4 up to 12 in some crystals.

Occurrence and Production

Lead is relatively rare in Earth’s crust (about 14 parts per million by weight, roughly the same abundance as silver). It is typically found as a component of mineral ores rather than freed. The most important lead ore is galena (lead sulfide, PbS), which often contains silver as an impurity. Other ore minerals include cerussite (lead carbonate, PbCO₃), anglesite (lead sulfate, PbSO₄), and litharge or massicot (forms of lead oxide). These minerals occur in hydrothermal veins, sedimentary deposits, and as weathering products of galena.

The largest lead deposits are found where ore-bearing veins concentrate minerals; historically, notable lead mining regions have included parts of Europe (e.g. the Mendips in England, the Harz in Germany, and parts of Croatia and Slovenia), as well as regions of the Americas, Australia, China, and Africa. Lead is often a byproduct of mining other metals: for example, it is produced in copper and zinc smelting processes or recovered from silver ores.

Extraction of lead generally involves first concentrating the ore and then smelting or roasting. In smelting, galena is roasted in air to convert lead sulfide into lead oxide and sulfur dioxide gas. The lead oxide is then reduced by carbon (coke) or a combination of carbon and carbon monoxide to produce metallic lead. Alternatively, some processes reduce sulfide ore directly. The molten lead is collected at the bottom of the furnace and can be cast into ingots or further refined to remove impurities (silver, gold, copper, etc.). Lead refining also occurs by electrolytic methods, though thermal (blast furnace) methods are more common for mass production.

Worldwide lead production is on the order of millions of tonnes per year. As of the early 21st century, annual global production of refined lead was around 10 million metric tons. About half of this output comes from recycling scrap lead, largely spent lead–acid batteries, which are collected and smelted to recover the lead. The remainder is mined ore. Major producers of lead (as of the 2020s) include China (the largest), followed by Australia, the United States, Peru, and Mexico. Because lead is relatively easy to recycle and separates from many alloys via gravity methods (due to its high density), recycling has become a key part of the lead supply chain.

In cosmic terms, lead is produced in ancient supernovae and in the late stages of stellar nucleosynthesis, and heavy lead is the stable end point of the thorium and uranium decay chains in stars. On Earth, roughly 70% of the lead we encounter has been recycled from older materials (construction, batteries, etc.) rather than newly mined, reflecting both its value and the environmental hazards of mining.

Applications and Technology

Lead’s unique combination of properties—high density, low melting point, and corrosion resistance—has led to many diverse applications, though health concerns have curtailed some traditional uses. Today, the most dominant commercial use of lead is in lead–acid rechargeable batteries, which account for about one-quarter to one-third of all lead consumption. These batteries (used in automobiles, backup power supplies, solar storage, etc.) rely on lead and lead dioxide plates submerged in sulfuric acid. Lead–acid batteries are heavy but cost-effective and reliable, and the lead is almost entirely recycled at battery end-of-life.

Another major application is radiation shielding: because lead is very effective at absorbing X-rays and gamma rays, it is used in protective aprons, shields around medical X-ray equipment, nuclear reactors, and even in space probes for electronics protection. Lead glass (glass with lead oxide added) is used for radiation shielding windows (e.g. in dentists’ offices or nuclear facilities) and in artistic crystal ware to increase refractive index. Similarly, lead–barium glass is used in cathode-ray tubes (older computer monitors and TVs) to block electron beams.

Historically, lead was used in uses now restricted or banned. For example, tetraethyl lead was added to gasoline in the 20th century to prevent engine knocking, but it was phased out from the 1970s–1980s onward in most countries because of its neurotoxicity. Lead was also used in paint (as white lead pigment) and in pipes and solder for plumbing; the awareness of lead poisoning led to bans on lead-based paints and replacement of lead pipes in drinking water systems in many places (the Flint, Michigan water crisis highlighted the dangers of lead pipes). Lead shot and bullets are widely used for hunting and ammunition, although this too is being restricted due to contamination of wildlife and ecosystems.

Other uses of lead include anti-corrosion coatings (abase lead is sometimes used on outdoor statues or rooftops), weights and counterbalances (for example, in radiation devices or scuba diving weights), and various alloys. Lead–tin alloys (solders) were once common for electrical soldering, but electronics today largely use lead-free solders. Lead’s low melting point and ease of casting also made it useful historically in casting type for printing and in making pewter (a tin–lead alloy) for utensils. Some pigments and stabilizers in plastics (such as lead stearate) are still used in certain applications, though replacing lead-based chemicals with safer alternatives is a major trend in industry.

In recent research applications, lead halide perovskites (materials containing lead and halogens) have drawn attention for high-efficiency solar cells and light emitters. These materials exploit lead’s electronic properties at the molecular scale, though their commercial use is tempered by concerns about lead leakage if devices break. In metallurgy, lead is sometimes used in bearing alloys or to impart machinability to special steels. It also has minor uses in chemical processes (for example, as a catalyst poison in octane cracking or in certain oxidation reactions) and in the nuclear industry as a coolant or spallation target, thanks to its stability under radiation.

Biology, Environment, and Safety

Lead is highly toxic to living organisms and has no known essential biological function. It is a potent neurotoxin: in humans and animals, lead interferes with nerve signaling and can cause developmental delays, cognitive deficits, and neurological disorders, especially in children. Lead inhibits many enzymes by binding to thiol (–SH) groups or by substituting for calcium or zinc in biochemical processes. It can cause anemia by disrupting hemoglobin synthesis, and it also damages the cardiovascular and renal systems. Chronic lead exposure accumulates in bones and tissues, remaining in the body for decades.

Because of these dangers, lead is tightly regulated in most countries. Occupational exposure limits (for inhaled dust and fumes) are enforced in industries that use lead, and products like paints, toys, and plumbing materials must meet strict lead limits. Drinking water regulations often set lead levels below 10–15 micrograms per liter, and blood lead levels above 5 micrograms per deciliter are considered a health concern. Lead paint in older buildings and lead-contaminated soil from past industrial sites require remediation to minimize exposure.

In the environment, lead persists because it is non-biodegradable. Poor waste handling can release lead into air (from smelting or burning) and water, where it settles into sediments. Acid rain can mobilize lead from soils. Plants and microbes do not use lead, but some can tolerate or even sequester small amounts. Lead in the food chain can harm animals; for instance, waterfowl that ingest lead shot or fishing weights are poisoned. Recent efforts to replace lead in ammunition aim to reduce such environmental loading.

Handling lead requires care: simple precautions include washing hands after contact, dust control work practices, and using protective equipment (gloves, respirators) where fine lead particulates or fumes are present. Physically, lead metal dust or oxide can be generated by machining; inhaling these particles is a severe hazard. Disposal of lead waste must prevent leaching into ground water (for example, by secure landfilling or recycling). Because lead smoke or fumes are particularly dangerous, industrial operations often use wet scrubbing or other technologies to capture and filter emissions.

On the positive side, modern processes recycle lead at very high efficiency. Over 95% of lead in spent batteries is typically recovered and reused. This circular use helps mitigate the need for new mining and reduces environmental release. Nonetheless, lead’s legacy as a historical material means that older paints, plumbing, and soils around highways and factories still pose risks. Public health efforts continue to lower exposure levels, as even small amounts of lead can impair neurological development in children.

History and Etymology

Lead is one of the few metals known since prehistoric times. Archeological evidence shows that humans first smelted lead from ore in the Near East over 7,000 years ago (circa 5,000 BC). The ancient Romans used lead for water pipes, wine cups, and cosmetics (though they did not fully realize the health danger, the Latin word for lead is plumbum, which gives the element its symbol Pb and is also the origin of the English word plumbing). Lead compounds like kohl (a black eye cosmetic) contained lead sulfide, and the Romans are thought to have suffered lead poisoning from lead-glazed pottery and lead-lined aqueducts, contributing to debates about lead and health in antiquity.

The etymology “Pb” reflects the long history: “plumbum” (Latin for lead) is the root of plumbing, and of words in Romance languages (French plomb, Spanish plomo). In English, lead’s Old High German ancestor was lead, possibly from the verb leathan (meaning to guide or go, reflecting lead workers’ craft). The numbering of its atomic symbol was assigned only after the periodic table was developed, but before then, alchemists associated lead with the planet Saturn.

Medieval and Renaissance alchemists considered lead one of the seven metals of alchemy. They attempted to transmute lead into gold (unsuccessfully) or to use it to cure the “Saturnine” temperament. In the 18th century, chemists began to distinguish lead from other similar elements (for example, Pierre Joséph Pelletier in 1753 determined that lead was distinct from bismuth, another heavy post-transition metal that was often confused with lead). By the 19th century, lead’s chemistry was well-understood: Humphry Davy in 1809 demonstrated that lead is an element by demonstrating its electrochemical reactions, and during the Industrial Revolution, mass smelting and refining of lead greatly expanded. The invention of the lead–acid battery by Gaston Planté in 1859 and its commercialization in the late 19th century vastly increased demand for refined lead.

During the 20th century, the story of lead is marked by both widespread use and growing awareness of its hazards. Tetraethyl lead was introduced in 1921 as a fuel additive for automobiles; it made engines run smoother but also released lead into air at scale. Concerns grew after World War II about lead in water systems and food. Beginning in the 1960s and 1970s, scientific reports led many governments to phase out leaded gasoline (the United States banned it for on-road vehicles by 1996, for example) and ban lead-based paints for household use. Public health milestones include the recognition of lead poisoning symptoms in the 19th century (such as “painters’ colic” in 19th-century London) and the modern lowering of acceptable blood lead thresholds. In 2011 the United Nations banned new lead-based paint globally.

In recent decades, lead has mostly shifted from old uses to niche applications that justify careful control (such as specialized batteries, electronics solders, and aviation fuel for small aircraft, where it remains one of the few high-octane options). New regulations on lead in products (electronics, batteries, and building materials) reflect both its toxic legacy and ongoing reliance in certain technologies.

Data Table

Property	Value
Symbol	Pb (from Latin plumbum)
Atomic number (Z)	82
Standard atomic weight	207.2
Group / period / block	14 (carbon group) / 6 / p-block
Category	Post-transition (poor) metal
Electron configuration	[Xe] 4f^14 5d^10 6s^2 6p^2
Valence electrons	4 (6s^2 6p^2)
Common oxidation states	+2, +4
Stable isotopes	^204Pb, ^206Pb, ^207Pb, ^208Pb (natural abundance: ~1.4%, 24.1%, 22.1%, 52.4%)
Atomic radius	~175 pm (atomic radius), 146 pm (covalent)
Electronegativity	2.33 (Pauling scale)
Ionization energy (1st)	715.6 kJ/mol (7.42 eV)
Density (ρ)	11.34 g/cm³ (at 20 °C)
Phase at STP	Solid
Melting point	600.6 K (327.5 °C)
Boiling point	2022 K (1749 °C)
Crystal structure	C (face-centered cubic)
Electrical resistivity	192 nΩ·m (20 °C)
Thermal conductivity	35.3 W/(m·K) (at 300 K)
Magnetic ordering	Diamagnetic (no permanent magnetism)