Editing Lead (section)

== Chemistry ==
[[File:FlammenfärbungPb.png|thumb|right|upright=0.5|[[Flame test]]: lead colors flame pale blue|alt=A flame with a small metal rod penetrating it; the flame near the rod is pale blue.]]
Bulk lead exposed to moist air forms a protective layer of varying composition. [[Lead carbonate|Lead(II) carbonate]] is a common constituent;{{sfn|Thürmer|Williams|Reutt-Robey|2002|pp=2033–2035}}{{sfn|Tétreault|Sirois|Stamatopoulou|1998|pp=17–32}}{{sfn|Thornton|Rautiu|Brush|2001|pp=10–11}} the [[lead(II) sulfate|sulfate]] or [[lead(II) chloride|chloride]] may also be present in urban or maritime settings.{{sfn|Greenwood|Earnshaw|1998|p=373}} This layer makes bulk lead effectively chemically inert in the air.{{sfn|Greenwood|Earnshaw|1998|p=373}} Finely powdered lead, as with many metals, is [[pyrophoricity|pyrophoric]],{{sfn|Bretherick|2016|p=1442}} and burns with a bluish-white flame.{{sfn|Harbison|Bourgeois|Johnson|2015|p=132}}

[[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.{{sfn|Greenwood|Earnshaw|1998|p=373}} Molten lead reacts with the [[chalcogen]]s to give lead(II) chalcogenides.{{sfn|Greenwood|Earnshaw|1998|p=374}}

Lead metal resists [[sulfuric acid|sulfuric]] and [[phosphoric acid]] but not [[hydrochloric acid|hydrochloric]] or [[nitric acid]]; the outcome depends on insolubility and subsequent passivation of the product salt.{{sfn|Thornton|Rautiu|Brush|2001|pp=11–12}} Organic acids, such as [[acetic acid]], dissolve lead in the presence of oxygen.{{sfn|Greenwood|Earnshaw|1998|p=373}} Concentrated [[alkali]]s dissolve lead and form [[plumbite]]s.{{sfn|Polyanskiy|1986|p=20}}

=== Inorganic compounds ===
{{see also|Lead compounds}}

Lead shows two main oxidation states: +4 and +2. The [[Valence (chemistry)|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.{{sfn|Greenwood|Earnshaw|1998|p=373}} This is attributable to [[relativistic quantum chemistry|relativistic effects]], specifically the [[Inert-pair effect|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 sp<sup>3</sup> hybridization is still energetically favorable. Lead, like carbon, is predominantly tetravalent in such compounds.{{sfn|Kaupp|2014|pp=9–10}}

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.{{sfn|Dieter|Watson|2009|p=509}}

==== Lead(II) ====
[[File:Oxid_olovnatý.JPG|left|thumb|[[Lead(II) oxide]]|alt=Cream powder]]Lead(II) compounds are characteristic of the inorganic chemistry of lead. Even strong [[oxidizing agent]]s like fluorine and chlorine react with lead to give only [[Lead(II) fluoride|PbF<sub>2</sub>]] and [[Lead(II) chloride|PbCl<sub>2</sub>]].{{sfn|Greenwood|Earnshaw|1998|p=373}} Lead(II) ions are usually colorless in solution,{{sfn|Hunt|2014|p=215}} and partially hydrolyze to form Pb(OH)<sup>+</sup> and finally [Pb<sub>4</sub>(OH)<sub>4</sub>]<sup>4+</sup> (in which the [[Hydroxy group|hydroxyl]] ions act as [[bridging ligand]]s),{{sfn|King|1995|pp=43–63}}{{sfn|Bunker|Casey|2016|p=89}} but are not [[reducing agent]]s as tin(II) ions are. [[Qualitative inorganic analysis|Techniques]] for identifying the presence of the Pb<sup>2+</sup> ion in water generally rely on the precipitation of lead(II) chloride using dilute hydrochloric acid. As the chloride salt is sparingly soluble in water, in very dilute solutions the precipitation of lead(II) sulfide is instead achieved by bubbling [[hydrogen sulfide]] through the solution.{{sfn|Whitten|Gailey|David|1996|pp=904–905}}

[[Lead(II) oxide|Lead monoxide]] exists in two [[Crystal polymorphism|polymorphs]], [[litharge]] α-PbO (red) and [[massicot]] β-PbO (yellow), the latter being stable only above around 488&nbsp;°C. Litharge is the most commonly used inorganic compound of lead.{{sfn|Greenwood|Earnshaw|1998|p=384}} There is no lead(II) hydroxide; increasing the pH of solutions of lead(II) salts leads to hydrolysis and condensation.{{sfn|Greenwood|Earnshaw|1998|p=387}} Lead commonly reacts with heavier chalcogens. [[Lead sulfide]] is a [[semiconductor]], a [[Photoconductivity|photoconductor]], and an extremely sensitive [[Particle detector|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.{{sfn|Greenwood|Earnshaw|1998|p=389}}

[[File:Red-lead-unit-cell-3D-balls.png|right|thumb|upright| Lead and [[oxygen]] in a tetragonal [[Crystal structure#Unit cell|unit cell]] of [[lead(II,IV) oxide]]|alt=Alternating dark gray and red balls connected by dark gray-red cylinders]]
Lead dihalides are well-characterized; this includes the diastatide{{sfn|Zuckerman|Hagen|1989|p=426}} and mixed halides, such as PbFCl. The relative insolubility of the latter forms a useful basis for the [[Gravimetric analysis|gravimetric]] determination of fluorine. The difluoride was the first solid [[Ionic conductivity (solid state)|ionically conducting]] compound to be discovered (in 1834, by [[Michael Faraday]]).{{sfn|Funke|2013}} The other dihalides decompose on exposure to ultraviolet or visible light, especially [[Lead(II) iodide|the diiodide]].{{sfn|Greenwood|Earnshaw|1998|p=382}} Many lead(II) [[Pseudohalogen|pseudohalides]] are known, such as the cyanide, cyanate, and [[Lead(II) thiocyanate|thiocyanate]].{{sfn|Greenwood|Earnshaw|1998|p=389}}{{sfn|Bharara|Atwood|2006|p=4}} Lead(II) forms an extensive variety of halide [[coordination complex]]es, such as [PbCl<sub>4</sub>]<sup>2−</sup>, [PbCl<sub>6</sub>]<sup>4−</sup>, and the [Pb<sub>2</sub>Cl<sub>9</sub>]<sub>''n''</sub><sup>5''n''−</sup> chain anion.{{sfn|Greenwood|Earnshaw|1998|p=382}}

[[Lead(II) sulfate]] is insoluble in water, like the sulfates of other heavy divalent [[Ion|cations]]. [[Lead(II) nitrate]] and [[lead(II) acetate]] are very soluble, and this is exploited in the synthesis of other lead compounds.{{sfn|Greenwood|Earnshaw|1998|p=388}}

==== 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.{{sfn|Toxicological Profile for Lead|2007|p=277}} Lead(II) oxide gives a mixed oxide on further oxidation, Pb<sub>3</sub>O<sub>4</sub>. It is described as [[lead(II,IV) oxide]], or structurally 2PbO·PbO<sub>2</sub>, and is the best-known mixed valence lead compound. [[Lead dioxide]] is a strong oxidizing agent, capable of oxidizing hydrochloric acid to chlorine gas.{{sfn|Downs|Adams|2017|p=1128}} This is because the expected PbCl<sub>4</sub> that would be produced is unstable and spontaneously decomposes to PbCl<sub>2</sub> and Cl<sub>2</sub>.{{sfn|Brescia|2012|p=234}} Analogously to [[Lead(II) oxide|lead monoxide]], lead dioxide is capable of forming [[plumbate]] anions. [[Lead(IV) sulfide|Lead disulfide]]{{sfn|Macintyre|1992|p=3775}} and lead diselenide{{sfn|Silverman|1966|pp=2067–2069}} are only stable at high pressures. [[Lead tetrafluoride]], a yellow crystalline powder, is stable, but less so than the [[Lead(II) fluoride|difluoride]]. [[Lead(IV) chloride|Lead tetrachloride]] (a yellow oil) decomposes at room temperature, lead tetrabromide is less stable still, and the existence of lead tetraiodide is questionable.{{sfn|Greenwood|Earnshaw|1998|p=381}}

==== Other oxidation states ====
{{see also|Plumbide}}

[[File:Nonaplumbide-anion-from-xtal-3D-balls.png|thumb|left|upright|The [[gyroelongated square pyramid|capped square antiprismatic]] anion [Pb<sub>9</sub>]<sup>4−</sup> from [K(18-crown-6)]<sub>2</sub>K<sub>2</sub>Pb<sub>9</sub>·(en)<sub>1.5</sub>{{sfn|Yong|Hoffmann|Fässler|2006|pp=4774–4778}}|alt=Nine dark gray spheres connected by cylinders of the same color forming a convex shape]]

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 [[radical (chemistry)|radicals]].{{sfn|Becker|Förster|Franzen|Hartrath|2008|pp=9965–9978}}{{sfn|Mosseri|Henglein|Janata|1990|pp=2722–2726}}{{sfn|Konu|Chivers|2011|pp=391–392}} The same applies for lead(I), which can be found in such radical species.{{sfn|Hadlington|2017|p=59}}

Numerous mixed lead(II,IV) oxides are known. When PbO<sub>2</sub> is heated in air, it becomes Pb<sub>12</sub>O<sub>19</sub> at 293&nbsp;°C, Pb<sub>12</sub>O<sub>17</sub> at 351&nbsp;°C, Pb<sub>3</sub>O<sub>4</sub> at 374&nbsp;°C, and finally PbO at 605&nbsp;°C. A further [[sesquioxide]], Pb<sub>2</sub>O<sub>3</sub>, 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.{{sfn|Greenwood|Earnshaw|1998|pp=384–386}}

Negative oxidation states can occur as [[Zintl phase|Zintl phases]], as either free lead anions, as in Ba<sub>2</sub>Pb, with lead formally being <!--to avoid false positives in search for 'being led/lead' typo -->lead(−IV),{{sfn|Röhr|2017}} or in oxygen-sensitive ring-shaped or polyhedral cluster ions such as the [[trigonal bipyramidal molecular geometry|trigonal bipyramidal]] Pb<sub>5</sub><sup>2−</sup> ion, where two lead atoms are lead(−I) and three are lead(0).{{sfn|Alsfasser|2007|pp=261–263}} In such anions, each atom is at a polyhedral vertex and contributes two electrons to each covalent bond along an edge from their sp<sup>3</sup> hybrid orbitals, the other two being an external [[lone pair]].{{sfn|King|1995|pp=43–63}} They may be made in [[Ammonia|liquid ammonia]] via the reduction of lead by [[sodium]].{{sfn|Greenwood|Earnshaw|1998|p=393}}

{{Clear}}

=== Organolead ===
{{Main|Organolead compound}}

[[File:Tetraethyllead-3D-balls.png|right|thumb|upright|Structure of a [[tetraethyllead]] molecule:<br />
{{Color box|#505050}} [[Carbon]]<br />
{{Color box|#eeeeee}} [[Hydrogen]]<br />
{{Color box|#577179}} Lead|alt=A gray-green sphere linked to four black spheres, each, in turn, linked also to three white ones]]

Lead can form [[catenation|multiply-bonded chains]], a property it shares with its lighter [[Homologous series|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]].{{sfn|Greenwood|Earnshaw|1998|p=374}} With itself, lead can build metal–metal bonds of an order up to three.{{sfn|Stabenow|Saak|Weidenbruch|2003}} With carbon, lead forms organolead compounds similar to, but generally less stable<!--as can be seen below--> than, typical organic compounds{{sfn|Polyanskiy|1986|p=43}} (due to the Pb–C bond being rather weak).{{sfn|King|1995|pp=43–63}} This makes the [[organometallic chemistry]] of lead far less wide-ranging than that of tin.{{sfn|Greenwood|Earnshaw|1998|p=404}} 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(SiMe<sub>3</sub>)<sub>2</sub>]<sub>2</sub> and [[plumbocene]].{{sfn|Greenwood|Earnshaw|1998|p=404}}

The lead analog of the simplest [[organic compound]], [[methane]], is [[plumbane]]. Plumbane may be obtained in a reaction between metallic lead and atomic hydrogen.{{sfn|Wiberg|Wiberg|Holleman|2001|p=918}} Two simple derivatives, [[tetramethyllead]] and [[tetraethyllead]], are the best-known [[Organolead chemistry|organolead]] compounds. These compounds are relatively stable: tetraethyllead only starts to decompose if heated{{sfn|Toxicological Profile for Lead|2007|p=287}} or if exposed to sunlight or ultraviolet light.{{sfn|Polyanskiy|1986|p=44}}{{efn|[[Tetraphenyllead]] is even more thermally stable, decomposing at 270&nbsp;°C.{{sfn|Greenwood|Earnshaw|1998|p=404}}}} With sodium metal, lead readily forms an equimolar alloy that reacts with [[Haloalkane|alkyl halides]] to form [[Organometallic chemistry|organometallic]] compounds such as tetraethyllead.{{sfn|Windholz|1976}} The oxidizing nature of many organolead compounds is usefully exploited: [[Lead(IV) acetate|lead tetraacetate]] is an important laboratory reagent for oxidation in organic synthesis.{{sfn|Zýka|1966|p=569}} Tetraethyllead, once added to automotive gasoline, was produced in larger quantities than any other organometallic compound,{{sfn|Greenwood|Earnshaw|1998|p=404}} and is still widely used in [[avgas|fuel for small aircraft]].<ref>{{cite web|url=https://www.flyingmag.com/when-will-we-see-unleaded-av-gas/|title=When will we see unleaded AvGas?|date=5 August 2019 |access-date=2024-05-26}}</ref>
Other organolead compounds are less chemically stable.{{sfn|Polyanskiy|1986|p=43}} For many organic compounds, a lead analog does not exist.{{sfn|Wiberg|Wiberg|Holleman|2001|p=918}}