Editing Citric acid (section)

==Chemical characteristics==
[[File:Citric acid speciation.svg|thumb|left|[[Ion speciation|Speciation]] diagram for a 10-millimolar solution of citric acid]]

Citric acid can be obtained as an [[anhydrous]] (water-free) form or as a [[hydrate|monohydrate]]. The anhydrous form crystallizes from hot water, while the monohydrate forms when citric acid is crystallized from cold water. The monohydrate can be converted to the anhydrous form at about 78&nbsp;°C. Citric acid also dissolves in absolute (anhydrous) [[ethanol]] (76 parts of citric acid per 100 parts of ethanol) at 15&nbsp;°C. It [[decarboxylation|decomposes]] with loss of carbon dioxide above about 175&nbsp;°C.

Citric acid is a triprotic [[acid]], with [[Acid dissociation constant|pK<sub>a</sub>]] values, extrapolated to zero ionic strength, of 3.128, 4.761, and 6.396 at 25&nbsp;°C.<ref>{{cite journal
| first1= Robert N.| last1= Goldberg| first2= Nand| last2= Kishore| first3= Rebecca M.| last3= Lennen| s2cid= 94614267 |url=https://www.nist.gov/system/files/documents/srd/jpcrd615.pdf#page=35&zoom=auto,-76,487 |title= Thermodynamic Quantities for the Ionization Reactions of Buffers| journal= J. Phys. Chem. Ref. Data| year = 2002| volume= 31| issue= 1| pages= 231–370| doi= 10.1063/1.1416902| bibcode= 2002JPCRD..31..231G}} (Link added 4 August 2024)</ref> The pK<sub>a</sub> of the hydroxyl group has been found, by means of [[Carbon-13 NMR spectroscopy|<sup>13</sup>C NMR spectroscopy]], to be 14.4.<ref>{{cite journal| first1= Andre M. N.| last1= Silva
| first2= Xiaole| last2= Kong| first3 =Robert C.| last3=Hider| title= Determination of the pKa value of the hydroxyl group in the α-hydroxycarboxylates citrate, malate and lactate by <sup>13</sup>C NMR: implications for metal coordination in biological systems| journal= Biometals| year = 2009| volume= 22| issue= 5| pages= 771–778| doi= 10.1007/s10534-009-9224-5| pmid=19288211| s2cid= 11615864}}</ref> 
The speciation diagram shows that solutions of citric acid are [[buffer solution]]s between about pH&nbsp;2 and pH&nbsp;8. In biological systems around pH 7, the two species present are the citrate ion and mono-hydrogen citrate ion. The SSC 20X hybridization buffer is an example in common use.<ref>{{cite web | url=http://openwetware.org/wiki/SSC | title=SSC - OpenWetWare }}</ref><ref>Maniatis, T.; Fritsch, E. F.; Sambrook, J. 1982. Molecular Cloning: A Laboratory Manual. Cold Spring Harbor Laboratory, Cold Spring Harbor, NY.</ref> Tables compiled for biochemical studies are available.<ref>{{cite book | doi = 10.1016/0076-6879(55)01020-3 | chapter = [16] Preparation of buffers for use in enzyme studies
| title = Methods in Enzymology Volume 1 | year = 1955| last1 = Gomori| first1 = G.| isbn = 9780121818012| volume = 1| pages = 138–146 | chapter-url = https://static.igem.org/mediawiki/2008/8/84/Protein_Buffers.pdf | url = https://archive.org/details/methodsinenzymol01acad/page/138}}</ref>

Conversely, the pH of a 1&nbsp;mM solution of citric acid will be about 3.2. The pH of fruit juices from [[citrus fruit]]s like oranges and lemons depends on the citric acid concentration, with a higher concentration of citric acid resulting in a lower pH.

[[Acid salt]]s of citric acid can be prepared by careful adjustment of the pH before crystallizing the compound. See, for example, [[sodium citrate]].

The citrate ion forms complexes with metallic cations. The [[stability constants of complexes|stability constant]]s for the formation of these complexes are quite large because of the [[chelate effect]]. Consequently, it forms complexes even with alkali metal cations. However, when a chelate complex is formed using all three carboxylate groups, the chelate rings have 7 and 8 members, which are generally less stable thermodynamically than smaller chelate rings. In consequence, the hydroxyl group can be deprotonated, forming part of a more stable 5-membered ring, as in [[ammonium ferric citrate]], {{chem2|[NH4+]5Fe(3+)(C6H4O7(4−))2*2H2O}}.<ref>{{cite journal| first1= M.| last1= Matzapetakis| first2= C. P. | last2= Raptopoulou| first3 =A. | last3=Tsohos| first4 =V. | last4=Papaefthymiou| first5 =S. N. | last5=Moon| first6 =A. | last6=Salifoglou| title= Synthesis, Spectroscopic and Structural Characterization of the First Mononuclear, Water Soluble Iron−Citrate Complex, (NH<sub>4</sub>)<sub>5</sub>Fe(C<sub>6</sub>H<sub>4</sub>O<sub>7</sub>)<sub>2</sub>·2H<sub>2</sub>O| journal= J. Am. Chem. Soc.| year = 1998| volume= 120| issue= 50| pages= 13266–13267| doi= 10.1021/ja9807035}}</ref>

Citric acid can be [[ester]]ified at one or more of its three [[carboxylic acid]] groups to form any of a variety of mono-, di-, tri-, and mixed esters.<ref>{{cite journal |doi=10.1016/S0040-4020(96)01061-7|title=Total synthesis of rhizoferrin, an iron chelator|year=1997|last1=Bergeron|first1=Raymond J.|last2=Xin|first2=Meiguo|last3=Smith|first3=Richard E.|last4=Wollenweber|first4=Markus|last5=McManis|first5=James S.|last6=Ludin|first6=Christian|last7=Abboud|first7=Khalil A.|journal=Tetrahedron|volume=53|issue=2|pages=427–434}}</ref>