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Hydrogen ion

2025-07-28T20:41:20Z

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{{Short description|Hydrogen atom that has gained or lost an electron}}
{{more citations needed|date=October 2014}}
A '''hydrogen ion''' is created when a [[hydrogen]] atom loses or gains an [[electron]]. A positively charged hydrogen ion (or [[proton]]) can readily combine with other particles and therefore is only seen isolated when it is in a gaseous state or a nearly particle-free space.<ref>{{cite web|url=https://www.britannica.com/science/hydrogen-ion|title=Hydrogen ion - chemistry|website=britannica.com|access-date=18 March 2018}}</ref> Due to its extremely high [[charge density]] of approximately {{val|2e10}} times that of a sodium ion, the bare hydrogen ion cannot exist freely in solution as it readily hydrates, i.e., bonds quickly.<ref>due to its extremely high charge density of approximately {{val|2e10}} times that of a [[sodium]] ion</ref> The hydrogen ion is recommended by [[International Union of Pure and Applied Chemistry|IUPAC]] as a general term for all [[ion]]s of [[hydrogen]] and its [[isotope]]s.<ref>''Compendium of Chemical Terminology'', 2nd edition McNaught, A.D. and Wilkinson, A. Blackwell Science, 1997 {{ISBN|0-86542-684-8}}, also [http://www.iupac.org/publications/compendium/index.html online] {{webarchive|url=https://web.archive.org/web/20051212164525/http://www.iupac.org/publications/compendium/index.html |date=2005-12-12 }}</ref>
Depending on the [[electric charge|charge]] of the ion, two different classes can be distinguished: positively charged ions (hydrons) and negatively charged (hydride) ions.

==Cation (positively charged)==
{{Main|Hydron (chemistry)}}

[[Image:Zundel-cation.JPG|thumb|Zundel cation]]
A hydrogen atom is made up of a nucleus with charge +1, and a single electron. Therefore, the only positively charged ion possible has charge +1. It is notated H+.

Depending on the isotope in question, the hydrogen cation has different names:

* [[Hydron (chemistry)|Hydron]]: general name referring to the positive ion of any hydrogen isotope (H+)
* [[Proton]]: 1H+ (i.e. the cation of [[Hydrogen atom|protium]])
* [[Deuteron]]: 2H+, D+
* [[Triton (physics)|Triton]]: 3H+, T+
In addition, the ions produced by the reaction of these cations with water as well as their [[hydrates]] are called hydrogen ions:

* [[Hydronium]] ion: {{chem2|H3O+}}
* [[Zundel cation]]: {{chem2|H5O2+}} (named for [[Georg Zundel]])
* [[Eigen cation]]: {{chem2|H9O4+}} (or {{chem2|H3O+*3H2O}}) (named for [[Manfred Eigen]])

Zundel cations and Eigen cations play an important role in proton diffusion according to the [[Grotthuss mechanism]].

In connection with acids, "hydrogen ions" typically refers to hydrons.

[[Image:Ions.svg|thumb|left|upright=2.5]]In the image at left the hydrogen atom (center) contains a single proton and a single electron. Removal of the electron gives a cation (left), whereas addition of an electron gives an anion (right). The hydrogen anion, with its loosely held two-electron cloud, has a larger radius than the neutral atom, which in turn is much larger than the bare proton of the cation. Hydrogen forms the only cation that has no electrons, but even cations that (unlike hydrogen) still retain one or more electrons are still smaller than the neutral atoms or molecules from which they are derived.
{{clear}}

==Anion (negatively charged)==
{{main|Hydrogen anion}}

Hydrogen [[anion]]s are formed when additional electrons are acquired:

* [[Hydride]]: general name referring to the negative ion of any hydrogen isotope (H−)
* Protide: 1H−
* Deuteride: 2H−, D−
* Tritide: 3H−, T−

==Uses==
Hydrogen ions drive [[ATP synthase]] in [[photosynthesis]]. This happens when hydrogen ions get pushed across the membrane creating a high concentration inside the [[thylakoid membrane]] and a low concentration in the cytoplasm. However, because of osmosis, the H+ will force itself out of the membrane through ATP synthase. Using their [[kinetic energy]] to escape, the protons will spin the ATP synthase which in turn will create [[Adenosine triphosphate|ATP]]. This happens in [[cellular respiration]] as well though the concentrated membrane will instead be the inner membrane of the [[mitochondria]].

Hydrogen ions concentration, measured as pH, is also responsible for the [[acidic]] or [[basic (chemistry)|basic]] nature of a compound. Water molecules split to form H+ and hydroxide anions. This process is referred to as the [[self-ionization of water]].

==Ocean acidification==
The concentration of hydrogen ions and pH are inversely proportional; in an aqueous solution, an increased concentration of hydrogen ions yields a low pH, and subsequently, an acidic product. By definition, an acid is an ion or molecule that can donate a proton, and when introduced to a solution it will react with water molecules (H2O) to form a hydronium ion (H3O+), a conjugate acid of water.<ref>OpenStax, Chemistry. OpenStax CNX. Jun 20, 2016 http://cnx.org/contents/85abf193-2bd2-4908-8563-90b8a7ac8df6@9.311.</ref> For simplistic reasoning, the hydrogen ion (H+) is often used to abbreviate the hydronium ion.

[[Ocean acidification]] is the direct consequence of elevated concentrations of hydrogen ions and carbonate saturation from significant absorption of carbon dioxide (CO2) by the world's oceans.<ref>W.S. Broecker, T. Takahashi (1997) Neutralization of fossil fuel CO2 by marine calcium carbonate</ref> The pre-industrial state of the ocean's carbonate chemistry has been notably stable, including the balance of its pH.<ref>P.N. Pearson, M.R. Palmer (2000) Atmospheric carbon dioxide concentrations over the past 60 million years Nature, 406, pp. 695-699</ref> Following the industrial revolution, anthropogenic emissions of burning [[fossil fuels]], cement production, and changes in land use, have increased the oceans uptake of carbon dioxide from the atmosphere by 30%.<ref>C.L. Sabine, et al. (2004). The oceanic sink for anthropogenic CO2
Science, 305 (5682), pp. 367-371</ref> In the ocean, the [[blue carbon|absorption capacity of this greenhouse gas]] is 59 times higher than in the atmosphere;<ref>Lal R. (2008). Carbon sequestration. Philosophical transactions of the Royal Society of London. Series B, Biological sciences, 363(1492), 815–830. https://doi.org/10.1098/rstb.2007.2185</ref> the ocean acts as the largest [[carbon sink]] on the planet, playing a significant role in climate regulation.<ref>Ben I. Mcneil & Richard J. Matear (2007). Climate change feedbacks on future oceanic acidification, Tellus B: Chemical and Physical Meteorology, 59:2, 191-198</ref> In addition to carbon fluxes, the natural process of carbon sequestration from the atmosphere into the deep ocean is facilitated by two systems, the biological pump and the solubility pump. The solubility pump is a physico-chemical process that transfers CO2 at the air-sea interface.<ref>Hessen, D., Ågren, G., Anderson, T., Elser, J., & De Ruiter, P. (2004). Carbon Sequestration in Ecosystems: The Role of Stoichiometry. Ecology, 85(5), 1179-1192. Retrieved November 22, 2020, from http://www.jstor.org/stable/3450161</ref> Based on Henry's Law, the amount of dissolved CO2 in an aqueous solution is directly proportional to the partial pressure of CO2 in the atmosphere.<ref>Avishay DM, Tenny KM. Henry's Law. [Updated 2020 Sep 7]. In: StatPearls [Internet]. Treasure Island (FL): StatPearls Publishing; 2020 Jan-. Available from: https://www.ncbi.nlm.nih.gov/books/NBK544301/</ref> To maintain equilibrium, a state of high atmospheric partial pressure of CO2 leads to an increased oceanic exchange of this gas by molecular diffusion.

In the surface waters, dissolved atmospheric carbon dioxide (CO2(aq)) reacts with water molecules to form carbonic acid (H2CO3), a weak diprotic acid. Diprotic acids consist of two ionizable hydrogen atoms in each molecule.<ref>OpenStax, Chemistry. OpenStax CNX. Jun 20, 2016 http://cnx.org/contents/85abf193-2bd2-4908-8563-90b8a7ac8df6@9.311.</ref> In an aqueous solution, partial dissociation of carbonic acid releases a hydrogen proton (H+) and a bicarbonate ion ({{chem2|HCO3-}}), and subsequently, the bicarbonate ion dissociates into an additional hydrogen proton and a carbonate ion ({{chem2|CO3(2-)}}).<ref>OpenStax, Chemistry. OpenStax CNX. Jun 20, 2016 http://cnx.org/contents/85abf193-2bd2-4908-8563-90b8a7ac8df6@9.311.</ref> The dissolving and dissociating of these inorganic carbon species generate an increase in the concentration of hydrogen ions and inversely lowers ambient surface ocean pH. The carbonate buffering system governs the acidity of seawater by maintaining dissolved inorganic carbon species in chemical equilibrium.

The chemical equation consists of reactants and products that may react in either direction. More reactants added to a system yield more product production (the chemical reaction shifts to the right) and if more product is added, additional reactants will form, shifting the chemical reaction to the left. Therefore, in this model, a high concentration of the beginning reactant, carbon dioxide, produces an increased amount of end-product (H+ and {{chem2|CO3(2-)}}), thus lowering pH and creating a more acidic solution. The natural buffering system of the ocean resist the change in pH by producing more bicarbonate ions generated by free acid protons reacting with carbonate ions to form an alkaline character.<ref>Middelburg, J. J., Soetaert, K., & Hagens, M. (2020). Ocean Alkalinity, Buffering and Biogeochemical Processes. Reviews of geophysics (Washington, D.C. : 1985), 58(3), e2019RG000681. https://doi.org/10.1029/2019RG000681</ref> However, increasing atmospheric CO2 concentrations may exceed the buffering capacity threshold, consequently resulting in higher rates of ocean acidification. Shifts in the ocean's carbonate chemistry has the potential to manipulate [[ocean biogeochemical cycle]]s for many elements and compounds causing profound impacts on marine ecosystems. Furthermore, the solubility of CO2 is temperature dependent; elevated surface water temperatures reduce CO2 solubility. A continual rise in atmospheric partial pressure of CO2 could potentially convert the ocean from acting as sink (the vertical transport of carbon to the depths of the ocean) to becoming a source (CO2 degassing from the ocean), further increasing global temperatures.<ref>Matsumoto, K. (2007). Biology-mediated temperature control on atmosphericpCO2and ocean biogeochemistry. Geophysical Research Letters, 34(20). doi:10.1029/2007gl031301</ref>

==See also==
* [[Acid]]
* [[Protonation]]
* [[Dihydrogen cation]]
* [[Trihydrogen cation]]

==References==
{{reflist}}

{{Set index article}}

[[Category:Hydrogen physics|Ocean Acidification]]

Lysocline

2025-07-28T20:34:04Z

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{{short description|Depth in the ocean below which the rate of dissolution of calcite increases dramatically}}
[[File:Surface ocean present-day omega calcite, GLODAPv2.png|thumb|The graphic presents the present-day annual mean surface omega calcite: the normalised saturation state of calcite. Areas with a value less an 1 indicate a likeliness for dissolution (undersaturated) while a value over 1 indicates areas less likely for dissolution (oversaturation).]]
The '''lysocline''' is the depth in the [[ocean]] dependent upon the [[carbonate compensation depth]] (CCD), usually around 5 km, below which the rate of [[Solvation|dissolution]] of [[calcite]] increases dramatically because of a pressure effect. While the lysocline is the upper bound of this transition zone of calcite saturation, the CCD is the lower bound of this zone.<ref name=":02">{{Citation|last=Broecker|first=W. S.|title=6.19 – The Oceanic CaCO3 Cycle|date=2003|url=http://www.sciencedirect.com/science/article/pii/B0080437516061193|work=Treatise on Geochemistry|pages=529–549|editor-last=Holland|editor-first=Heinrich D.|publisher=Pergamon|doi=10.1016/b0-08-043751-6/06119-3|isbn=9780080437514|access-date=2019-10-17|editor2-last=Turekian|editor2-first=Karl K.|url-access=subscription}}</ref>

CaCO3 content in sediment varies with different depths of the ocean, spanned by levels of separation known as the transition zone. In the mid-depth area of the ocean, sediments are rich in CaCO3, content values reaching 85–95%.<ref name=":02" /> This area is then spanned hundreds of meters by the transition zone, ending in the abyssal depths with 0% concentration. The lysocline is the upper bound of the transition zone, where amounts of CaCO3 content begins to noticeably drop from the mid-depth 85–95% sediment. The CaCO3 content drops to 0% [[concentration]] at the lower bound, known as the calcite compensation depth.<ref name=":02" />

Shallow marine waters are generally [[Supersaturation|supersaturated]] in calcite, CaCO3, because as [[Marine life|marine organisms]] (which often have shells made of calcite or its polymorph, [[aragonite]]) die, they tend to fall downwards without dissolving.<ref>{{Cite journal|last=Shiraiwa|first=Y.|date=2003|title=Physiological regulation of carbon fixation in the photosynthesis and calcification of coccolithophorids|journal=Comparative Biochemistry and Physiology Part B: Biochemistry and Molecular Biology|volume=136|issue=4|pages=775–783|doi=10.1016/S1096-4959(03)00221-5|pmid=14662302|issn=1096-4959}}</ref> As depth and pressure increases within the [[water column]], calcite solubility increases, causing supersaturated water above the saturation depth, allowing for preservation and burial of CaCO3 on the seafloor.<ref name=":1">{{Cite journal|last1=Sigman|first1=D. M.|last2=Boyle|first2=E. A.|date=2000|title=Glacial/interglacial variations in atmospheric carbon dioxide|journal=Nature|language=en|volume=407|issue=6806|pages=859–869|doi=10.1038/35038000|pmid=11057657|bibcode=2000Natur.407..859S|s2cid=7136822|issn=1476-4687}}</ref> However, this creates undersaturated seawater below the saturation depth, preventing CaCO3 burial on the [[Seabed|sea floor]] as the shells start to dissolve.

The equation
:<math>\Omega = \frac{ \left[\text{Ca}^{2+}\right] \left[\text{CO}_3^{2-}\right] } {K'_\mathrm{sp}}</math>
expresses the CaCO3 saturation state of seawater.<ref name=":2">{{Cite journal|last=Zeebe|first=R. E.|date=2012|title=History of Seawater Carbonate Chemistry, Atmospheric CO2, and Ocean Acidification|journal=Annual Review of Earth and Planetary Sciences|volume=40|issue=1|pages=141–165|doi=10.1146/annurev-earth-042711-105521|bibcode=2012AREPS..40..141Z|s2cid=18682623|issn=0084-6597}}</ref> The calcite saturation horizon is where Ω = 1; dissolution proceeds slowly below this depth. The lysocline is the depth that this dissolution impacts is again notable, also known as the inflection point with sedimentary CaCO3 versus various water depths.<ref name=":2" />

== Calcite compensation depth ==
The [[Carbonate compensation depth|calcite compensation depth]] (CCD) occurs at the depth that the rate of calcite to the sediments is balanced with the dissolution flux, the depth at which the CaCO3 content are values 2–10%.<ref name=":2" /> Hence, the lysocline and CCD are not equivalent. The lysocline and compensation depth occur at greater depths in the [[Atlantic Ocean|Atlantic]] (5000–6000 m) than in the [[Pacific Ocean|Pacific]] (4000–5000 m), and at greater depths in [[Tropics|equatorial regions]] than in [[Polar regions of Earth|polar regions]].<ref>{{Cite journal|last1=Volat|first1=J. L.|last2=Pastouret|first2=L.|last3=V. G.|first3=Colette|date=1980|title=Dissolution and carbonate fluctuations in Pleistocene deep-sea cores: A review|journal=Marine Geology|volume=34|issue=1|pages=1–28|doi=10.1016/0025-3227(80)90138-3|bibcode=1980MGeol..34....1V|issn=0025-3227}}</ref>

The depth of the CCD varies as a function of the chemical composition of the seawater and its temperature.<ref>{{Cite journal|last=Broecker|first=W. S.|date=2009|title=Wally's Quest to Understand the Ocean's CaCO3 Cycle|journal=Annual Review of Marine Science|volume=1|issue=1|pages=1–18|doi=10.1146/annurev.marine.010908.163936|pmid=21141027|bibcode=2009ARMS....1....1B|s2cid=45348785|issn=1941-1405}}</ref> Specifically, it is the deep waters that are undersaturated with [[calcium carbonate]] primarily because its solubility increases strongly with increasing pressure and [[salinity]] and decreasing temperature. As the atmospheric concentration of [[carbon dioxide]] continues to increase, the CCD can be expected to decrease in depth, as the ocean's acidity rises.<ref name=":1" />

==See also==
* [[Biological pump]]
* [[Carbonate compensation depth]]
* [[Ocean acidification]]

==References==
{{Reflist}}

[[Category:Geochemistry]]
[[Category:Oceanography]]

Cerussite

2025-07-28T19:45:37Z

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{{Short description|Lead carbonate mineral}}
{{Distinguish|Sericite}}
{{Infobox mineral
| name = Cerussite
| category = [[Carbonate mineral]]
| image = Cerussite - Nakhlak mine, Anarak, Esfahan, Iran.jpg
| imagesize = 280px
| caption =
| formula = Lead carbonate: PbCO3
| IMAsymbol=Cer<ref>{{Cite journal|last=Warr|first=L.N.|date=2021|title=IMA–CNMNC approved mineral symbols|journal=Mineralogical Magazine|volume=85|issue=3|pages=291–320|doi=10.1180/mgm.2021.43|bibcode=2021MinM...85..291W|s2cid=235729616|doi-access=free}}</ref>
| molweight =
| strunz = 5.AB.15
| system = [[Orthorhombic]]
| class = Dipyramidal (mmm) [[H-M symbol]]: (2/m 2/m 2/m)
| symmetry = ''Pnma [62]''
| color = Colorless, white, gray, blue, or green
| habit = Massive granular, reticulate, tabular to equant crystals
| twinning = Simple or cyclic contact twins
| cleavage = Good [110] and [021]
| fracture = Brittle conchoidal
| mohs = 3 to 3.5
| luster = Adamantine, vitreous, resinous
| refractive = nα = 1.803, nβ = 2.074, nγ = 2.076
| opticalprop = Biaxial (−)
| birefringence = δ = 0.273
| pleochroism =
| streak = White
| gravity = 6.53–6.57
| density =
| melt =
| fusibility =
| diagnostic =
| solubility =
| diaphaneity = Transparent to translucent
| other = May fluoresce yellow under LW UV
| references = <ref>[https://www.mineralienatlas.de/lexikon/index.php/MineralData?mineral=Cerussite Mineralienatlas]</ref><ref>[https://rruff.info/doclib/hom/cerussite.pdf Cerussite]. Handbook of Mineralogy. (PDF) Retrieved on 2011-10-10.</ref><ref>[https://www.mindat.org/min-934.html Cerussite]. Mindat. Retrieved on 2011-10-10.</ref>
}}

'''Cerussite''' (also known as '''lead carbonate''' or '''white lead ore''') is a [[mineral]] consisting of [[lead carbonate]] with the chemical formula PbCO3, and is an important [[ore]] of lead. The name is from the [[Latin]] ''cerussa'', [[white lead]]. ''Cerussa nativa'' was mentioned by [[Conrad Gessner]] in 1565, and in 1832 [[François Sulpice Beudant|F. S. Beudant]] applied the name ''céruse'' to the mineral, whilst the present form, cerussite, is due to [[W. Haidinger]] (1845). Miners' names in early use were lead-spar and white-lead-ore.{{sfn|Spencer|1911}}

Cerussite [[crystal]]lizes in the [[orthorhombic crystal system]] and is isomorphous with [[aragonite]]. Like aragonite it is very frequently [[crystal twinning|twinned]], the compound crystals being pseudo-hexagonal in form. Three crystals are usually twinned together on two faces of the prism, producing six-rayed stellate groups with the individual crystals intercrossing at angles of nearly 60°. Crystals are of frequent occurrence and they usually have very bright and smooth faces. The mineral also occurs in compact granular masses, and sometimes in fibrous forms. The mineral is usually colorless or white, sometimes grey or greenish in tint and varies from transparent to translucent with an adamantine lustre. It is very brittle, and has a [[conchoidal fracture]]. It has a [[Mohs hardness]] of 3 to 3.75 and a [[specific gravity]] of 6.5. A variety containing 7% of zinc carbonate, replacing lead carbonate, is known as iglesiasite, from Iglesias in [[Sardinia]], where it is found.{{sfn|Spencer|1911}}

The mineral may be readily recognized by its characteristic twinning, in conjunction with the adamantine lustre and high specific gravity. It dissolves with effervescence in dilute [[nitric acid]]. A [[blowpipe (tool)|blowpipe test]] will cause it to fuse very readily, and gives indications for lead.{{sfn|Spencer|1911}}

Finely crystallized specimens have been obtained from the [[Lahnstein#Friedrichssegen|Friedrichssegen]] mine in [[Lahnstein]] in [[Rhineland-Palatinate]], [[Johanngeorgenstadt]] in [[Saxony]], [[Stříbro]] in the [[Czech Republic]], [[Phoenixville]] in [[Pennsylvania]], [[Broken Hill, New South Wales|Broken Hill in New South Wales]], and several other localities. Delicate [[Acicular (crystal habit)|acicular]] crystals of considerable length were found long ago in the Pentire Glaze mine near [[St Minver]] in [[Cornwall]].{{sfn|Spencer|1911}} Cerussite is often found in considerable quantities, and has a lead content of up to 77.5%.<ref name="LeadAmount">{{Cite FTP |url=ftp://ftp.conservation.ca.gov/pub/dmg/pubs/cg/1949/02_11.pdf |title=Mineral Information Service |date=November 1949 |accessdate=4 June 2016 |server=California Division of Mines |url-status=dead |pages=2 }}</ref>

Lead(II) carbonate is practically insoluble in neutral water (solubility product [Pb2+][{{chem2|CO3(2−)}}] ≈ {{val|1.5e-13}} at 25 °C), but will dissolve in dilute acids.

==Commercial uses==
"[[White lead]]" is the key ingredient in (now discontinued) [[lead paint]]s. Ingestion of lead-based paint chips is the most common cause of [[lead poisoning]] in children.<ref>{{cite web |title=Lead Poisoning in Children |url=http://familydoctor.org/617.xml |url-status=dead |archive-url=https://web.archive.org/web/20060925103949/http://familydoctor.org/617.xml |archive-date=September 25, 2006}}</ref><ref>{{cite web |url=http://www.calpoison.org/public/lead.html |title=California Poison Control System: Lead Poisoning |url-status=dead |archive-url=https://web.archive.org/web/20070114141758/http://www.calpoison.org/public/lead.html |archive-date=2007-01-14}}</ref><ref>{{cite journal |title=Lightening the lead load in children |journal=Am Fam Physician |date=August 2000 |pmid=10950212|last1=Ellis |first1=M. R. |last2=Kane |first2=K. Y. |volume=62 |issue=3 |pages=545-54, 559-60 }}</ref>

Both "white lead" and [[Lead(II) acetate|lead acetate]] have been used in cosmetics throughout history, though this practice has ceased in Western countries.<ref>Gunn, Fenja. (1973). The Artificial Face: A History of Cosmetics. — as cited in [http://websites.umich.edu/~ece/student_projects/leisure/leadfacepowder.html Leisure Activities of an 18th Century Lady] and [http://leda.law.harvard.edu/leda/data/788/Schaffer06.pdf Reading Our Lips: The History of Lipstick Regulation in Western Seats of Power] {{webarchive|url=https://web.archive.org/web/20060905135116/http://leda.law.harvard.edu/leda/data/788/Schaffer06.pdf |date=2006-09-05}}</ref>

==Gallery==
<gallery widths="145px" heights="130px">
File:cerussite09.jpg|[[Crystal]]s of cerussite, a secondary lead ore
File:Light of the Desert.jpg|At 890 carats, the [[Light of the Desert]] (located at [[Toronto]]'s [[Royal Ontario Museum]]) is the world's largest faceted cerussite.<ref>{{cite web |title=Iconic: Light of the Desert |date=2009-11-26 |url=https://www.rom.on.ca/en/collections-research/rom-channel/iconic-light-desert |accessdate=2012-07-25}}</ref>
File:Cerussite-Malachite-Mimetite-158529.jpg|Colorless cerussite crystal that has been included by wisps of light green [[malachite]]
File:Cerussite-18566.jpg|Fine example of reticulated growth
</gallery>

==See also==
*[[Venetian ceruse]] – Cerussite-based cosmetic popularly thought to be worn by [[Elizabeth I of England]]

==References==
*{{EB1911|wstitle=Cerussite |volume=5 |page=762 |first=Leonard James|last=Spencer|author-link=Leonard James Spencer}}
{{Reflist}}

==External links==
* [https://web.archive.org/web/20040404192328/http://mineral.galleries.com/minerals/carbonat/cerussit/cerussit.htm Mineral galleries]

[[Category:Carbonate minerals]]
[[Category:Gemstones]]
[[Category:Lead minerals]]
[[Category:Luminescent minerals]]
[[Category:Minerals in space group 62]]
[[Category:Orthorhombic minerals]]
[[Category:Aragonite group]]
[[Category:Minerals described in 1845]]