Citric acid

Citric acid
Names
IUPAC name
2-hydroxypropane-1,2,3-tricarboxylic acid
Other names
Citric Acid
3-carboxy-3-hydroxypentanedioic acid
2-hydroxy-1,2,3-propanetricarboxylic acid[1]
Identifiers
ATC code A09AB04
77-92-9 Yes
ChEBI CHEBI:30769 Yes
ChEMBL ChEMBL1261 Yes
ChemSpider 305 Yes
DrugBank DB04272 Yes
EC number 201-069-1
Jmol-3D images Image
KEGG D00037 Yes
PubChem 311
RTECS number GE7350000
UNII XF417D3PSL Yes
Properties
Molecular formula
 C6H8O7
Molar mass 192.12 g·mol−1
Appearance crystalline white solid
Odor odorless
Density 1.665 g/cm3 (anhydrous)
1.542 g/cm3 (18 °C, monohydrate)
Melting point 156 °C (313 °F; 429 K)
Boiling point 310 °C (590 °F; 583 K) decomposes from 175 °C[2]
117.43 g/100 mL (10 °C)
147.76 g/100 mL (20 °C)
180.89 g/100 mL (30 °C)
220.19 g/100 mL (40 °C)
382.48 g/100 mL (80 °C)
547.79 g/100 mL (100 °C)[3]
Solubility soluble in alcohol, ether, ethyl acetate, DMSO
insoluble in C6H6, CHCl3, CS2, toluene[2]
Solubility in ethanol 62 g/100 g (25 °C)[2]
Solubility in amyl acetate 4.41 g/100 g (25 °C)[2]
Solubility in diethyl ether 1.05 g/100 g (25 °C)[2]
Solubility in 1,4-Dioxane 35.9 g/100 g (25 °C)[2]
log P -1.64
Acidity (pKa) pKa1 = 3.13[4]
pKa2 = 4.76[4]
pKa3 = 6.39,[5] 6.40[6]
1.493 - 1.509 (20 °C)[3]
1.46 (150 °C)[2]
Viscosity 6.5 cP (50% aq. sol.)[3]
Structure
Crystal structure Monoclinic
Thermochemistry
Specific
heat capacity (C)
 226.51 J/mol·K (26.85 °C)[7]
252.1 J/mol·K[7]
Std enthalpy of
formation (ΔfHo298)
 -1548.8 kJ/mol[3]
Std enthalpy of
combustion (ΔcHo298)
 -1960.6 kJ/mol[7]
-1972.34 kJ/mol (monohydrate)[3]
Hazards
MSDS HMDB
Main hazards skin and eye irritant
GHS pictograms [4]
GHS signal word Warning
H319[4]
P305+351+338[4]
EU classification Xi C
R-phrases R34, R36/37/38, R41
S-phrases S24/25, S26, S36/37/39, S45
NFPA 704
Flash point 155 °C (311 °F; 428 K)
345 °C (653 °F; 618 K)
Explosive limits 8%[4]
3000 mg/kg (rats, oral)
Related compounds
Related compounds
 Sodium citrate
Calcium citrate
Except where noted otherwise, data is given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
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Infobox references

Citric acid is a weak organic acid with the formula C6H8O7. It is a natural preservative which occurs naturally in citrus fruits and is also used to add an acidic or sour taste to foods and drinks. In biochemistry, the conjugate base of citric acid, citrate, is important as an intermediate in the citric acid cycle, which occurs in the metabolism of all aerobic organisms. It consists of 3 carboxyl (R-COOH) groups.

Citric acid is a commodity chemical, and more than a million tons are produced every year by fermentation. It is used mainly as an acidifier, as a flavoring, and as a chelating agent.

Properties

Citric acid crystal under polarized light, enlarged 200x
Concentration pH
1M ≈1.57
0.5M ≈1.72
0.1M ≈2.57

At room temperature, citric acid is a white hygroscopic crystalline powder. It can exist either in an anhydrous (water-free) form or as a 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 by heating above 78 °C. Citric acid also dissolves in absolute (anhydrous) ethanol (76 parts of citric acid per 100 parts of ethanol) at 15 °C.

In chemical structure, citric acid shares the properties of other carboxylic acids. When heated above 175 °C, it decomposes through the loss of carbon dioxide and water (see decarboxylation).

Citric acid is a slightly stronger acid than typical carboxylic acids because the anion can be stabilized by intramolecular hydrogen-bonding from other protic groups on citric acid.

Discovery and production

Lemons, oranges, limes, and other citrus fruits possess high concentrations of citric acid

Medieval scholars in Europe were aware of the acidity of lemon and lime juices; such knowledge is recorded in the 13th century encyclopedia Speculum Maius (The Great Mirror), compiled by Vincent of Beauvais. Citric acid was first isolated in 1784 by the chemist Carl Wilhelm Scheele, who crystallized it from lemon juice.[8][9]
Industrial-scale citric acid production first began in 1890 based on the Italian citrus fruit industry, where the juice was treated with hydrated lime (calcium hydroxide) to precipitate calcium citrate, which was isolated and converted back to the acid using diluted sulfuric acid.[8]

3Ca(OH)2(s) + 2C6H8O7(l) → Ca3(C6H5O7)2(s) + 3H2O(l)
3H2SO4(l) + Ca3(C6H5O7)2(s) → 2C6H8O7(l) + 3CaSO4(s)

In 1893, C. Wehmer discovered Penicillium mold could produce citric acid from sugar. However, microbial production of citric acid did not become industrially important until World War I disrupted Italian citrus exports.
In 1917, American food chemist James Currie discovered certain strains of the mold Aspergillus niger could be efficient citric acid producers, and the pharmaceutical company Pfizer began industrial-level production using this technique two years later, followed by Citrique Belge in 1929.

In this production technique, which is still the major industrial route to citric acid used today, cultures of A. niger are fed on a sucrose or glucose-containing medium to produce citric acid. The source of sugar is corn steep liquor, molasses, hydrolyzed corn starch or other inexpensive sugary solutions.[10] After the mold is filtered out of the resulting solution, citric acid is isolated by precipitating it with calcium hydroxide to yield calcium citrate salt, from which citric acid is regenerated by treatment with sulfuric acid, as in the direct extraction from citrus fruit juice.

In 1977, a patent was granted to Lever Brothers for the chemical synthesis of citric acid starting either from aconitic or isocitrate/alloisocitrate calcium salts under high pressure conditions. This produced citric acid in near quantitative conversion under what appeared to be a reverse non-enzymatic Krebs cycle reaction.[11]

In 2007, world wide annual production stood at approximately 1,600,000 tons.[12] More than 50% of this volume was produced in China. More than 50% was used as acidity regulator in beverages, some 20% in other food applications, 20% for detergent applications and 10% for related applications other than food, such as cosmetics, pharmaceutics and in the chemical industry.

Occurrence

Citric acid exists in greater than trace amounts in a variety of fruits and vegetables, most notably citrus fruits. Lemons and limes have particularly high concentrations of the acid; it can constitute as much as 8% of the dry weight of these fruits (about 47 g/L in the juices[13]). The concentrations of citric acid in citrus fruits range from 0.005 mol/L for oranges and grapefruits to 0.30 mol/L in lemons and limes. Within species, these values vary depending on the cultivar and the circumstances in which the fruit was grown.

Biochemistry

Citric acid cycle

Main article: Citric acid cycle

Citrate, the conjugate base of citric acid is one of a series of compounds involved in the physiological oxidation of acetate from fats, proteins, and carbohydrates. The acetate from these macronutrients are converted into the intracellular energy of ATP, as well as the common byproducts carbon dioxide, and water.

This series of chemical reactions is central to nearly all metabolic reactions, and is the source of two-thirds of the food-derived energy in higher organisms. Hans Adolf Krebs received the 1953 Nobel Prize in Physiology or Medicine for the discovery. The series of reactions is known by various names, including the "citric acid cycle", the "Krebs cycle" or "Szent-Györgyi — Krebs cycle", and the "tricarboxylic acid (TCA) cycle".

Other biological roles

Citrate is a vital component of bone, helping to regulate the size of calcium crystals.[14]

Applications

The dominant use of citric acid is as a flavoring and preservative in food and beverages, especially soft drinks.[8] Within the European Union it is denoted by E number E330. Citrate salts of various metals are used to deliver those minerals in a biologically available form in many dietary supplements. The buffering properties of citrates are used to control pH in household cleaners and pharmaceuticals. In the United States the purity requirements for citric acid as a food additive are defined by the Food Chemicals Codex, which is published by the United States Pharmacopoeia (USP).

Foods, other

Citric acid can be added to ice cream as an emulsifying agent to keep fats from separating, to caramel to prevent sucrose crystallization, or in recipes in place of fresh lemon juice. Citric acid is used with sodium bicarbonate in a wide range of effervescent formulae, both for ingestion (e.g., powders and tablets) and for personal care (e.g., bath salts, bath bombs, and cleaning of grease). Citric acid is also often used in cleaning products and sodas or fizzy drinks.

Citric acid sold in a dry powdered form is commonly sold in markets and groceries as "sour salt", due to its physical resemblance to table salt. It has use in culinary applications where an acid is needed for either its chemical properties or for its sour flavor, but a dry ingredient is needed and additional flavors are unwanted (e.g., instead of vinegar or lemon juice).

Cleaning and chelating agent

Citric acid is an excellent chelating agent, binding metals. It is used to remove limescale from boilers and evaporators.[8] It can be used to soften water, which makes it useful in soaps and laundry detergents. By chelating the metals in hard water, it lets these cleaners produce foam and work better without need for water softening. Citric acid is the active ingredient in some bathroom and kitchen cleaning solutions. A solution with a 6% concentration of citric acid will remove hard water stains from glass without scrubbing. In the industry, it is used to dissolve rust from steel. Citric acid can be used in shampoo to wash out wax and coloring from the hair.

Illustrative of its chelating abilities, citric acid was the first successful eluant used for total ion-exchange separation of the lanthanides, during the Manhattan Project in the 1940s. In the 1950s, it was replaced by the far more efficient EDTA. It can be used to substantially slow setting of Portland cement.

Cosmetics and pharmaceuticals

Citric acid is widely used as a pH adjusting agent in creams and gels of all kinds. In this role, it is classified in most jurisdictions as a processing aid and so does not need to be listed on ingredient lists.

Citric acid is an alpha hydroxy acid and used as an active ingredient in chemical peels.

Citric acid is commonly used as a buffer to increase the solubility of brown heroin. Single-use citric acid sachets have been used as an inducement to get heroin users to exchange their dirty needles for clean needles in an attempt to decrease the spread of HIV and hepatitis.[15] Other acidifiers used for brown heroin are ascorbic acid, acetic acid, and lactic acid; in their absence, a drug user will often substitute lemon juice or vinegar.

Citric acid is used as one of the active ingredients in the production of antiviral tissues.[16]

Dyeing

Citric acid can be used in food coloring to balance the pH level of a normally basic dye. It is used as an odorless alternative to white vinegar for home dyeing with acid dyes.

Qualitative analysis

Sodium citrate, the sodium salt of citric acid, is used as a chelating agent and is present in the Benedict's reagent, used for identification both qualitatively and quantitatively, of reducing sugars.

Industrial and construction

Citric acid can be used as a successful alternative to nitric acid in passivation of stainless steel.[17]

Photography

Citric acid can be used as a lower-odor stop bath as part of the process for developing photographic film. Photographic developers are alkaline, so a mild acid is used to neutralize and stop their action quickly, but commonly used acetic acid leaves a strong vinegar odor in the darkroom.[18]

Synthesize solid materials from small molecules

In materials science, the Citrate-gel method is similar process to sol-gel method which is a method for producing solid materials from small molecules. During the synthetic process, metal salts or alkoxides are introduced into a citric acid solution. The formation of citric complexes is believed to balance the difference in individual behaviour of ions in solution, which results in a better distribution of ions and prevents the separation of components at later process stages. The polycondensation of ethylene glycol and citric acid starts above 100ºС, resulting in polymer citrate gel formation.

Safety

As a weak acid, exposure to pure citric acid can cause adverse effects: inhalation may cause cough, shortness of breath, or sore throat; ingestion may cause abdominal pain and sore throat; exposure to skin or eyes may cause redness or pain.[19] Long term or repeated consumption may cause erosion of tooth enamel.[19][20][21]

Compendial status

See also

References

  1. ↑ David R. Lide, ed. (2005). "Physical Constants of Organic Compounds". CRC Handbook of Chemistry and Physics (Internet Version ed.). Boca Raton, FL: CRC Press.
  2. ↑ 2.0 2.1 2.2 2.3 2.4 2.5 2.6 http://chemister.ru/Database/properties-en.php?dbid=1&id=1624
  3. ↑ 3.0 3.1 3.2 3.3 3.4 CID 311 from PubChem
  4. ↑ 4.0 4.1 4.2 4.3 4.4 4.5 Sigma-Aldrich Co., Citric acid. Retrieved on 2014-06-02.
  5. ↑ "Data for Biochemical Research". ZirChrom Separations, Inc. Retrieved January 11, 2012.
  6. ↑ "Ionization Constants of Organic Acids". Michigan State University. Retrieved January 11, 2012.
  7. ↑ 7.0 7.1 7.2 Citric acid in Linstrom, P.J.; Mallard, W.G. (eds.) NIST Chemistry WebBook, NIST Standard Reference Database Number 69. National Institute of Standards and Technology, Gaithersburg MD. http://webbook.nist.gov (retrieved 2014-06-02)
  8. ↑ 8.0 8.1 8.2 8.3 Frank H. Verhoff (2005), "Citric Acid", Ullmann's Encyclopedia of Industrial Chemistry, Weinheim: Wiley-VCH
  9. ↑ Graham, Thomas (1842). Elements of chemistry, including the applications of the science in the arts. Hippolyte Baillière, foreign bookseller to the Royal College of Surgeons, and to the Royal Society, 219, Regent Street. p. 944. Retrieved June 4, 2010.
  10. ↑ Lotfy, Walid A.; Ghanem, Khaled M.; El-Helow, Ehab R. (2007). "Citric acid production by a novel Aspergillus niger isolate: II. Optimization of process parameters through statistical experimental designs". Bioresource Technology 98 (18): 3470–3477. doi:10.1016/j.biortech.2006.11.032. PMID 17317159.
  11. ↑ US 4056567-V.Lamberti and E.Gutierrez
  12. ↑ Berovic, M.; Legisa, M. (2007). "Citric acid production". Biotechnology Annual Review Volume 13. Biotechnology Annual Review 13. pp. 303–343. doi:10.1016/S1387-2656(07)13011-8. ISBN 9780444530325. PMID 17875481.
  13. ↑ Penniston KL, Nakada SY, Holmes RP, Assimos DG (2008). "Quantitative Assessment of Citric Acid in Lemon Juice, Lime Juice, and Commercially-Available Fruit Juice Products" (PDF). Journal of Endourology 22 (3): 567–570. doi:10.1089/end.2007.0304. PMC 2637791. PMID 18290732.
  14. ↑ Hu, Y.-Y.; Rawal, A.; Schmidt-Rohr, K. (December 2010). "Strongly bound citrate stabilizes the apatite nanocrystals in bone". Proceedings of the National Academy of Sciences 107 (52): 22425–22429. doi:10.1073/pnas.1009219107. PMC 3012505. PMID 21127269. Retrieved July 28, 2012.
  15. ↑ Garden, J., Roberts, K., Taylor, A., and Robinson, D. (2003). "Evaluation of the Provision of Single Use Citric Acid Sachets to Injecting Drug Users" (pdf). Scottish Center for Infection and Environmental Health.
  16. ↑ "Tissues that fight germs". CNN. July 14, 2004. Retrieved May 8, 2008.
  17. ↑ "Pickling and Passivating Stainless Steel" (PDF). Euro-inox.org. Retrieved 2013-01-01.
  18. ↑ Anchell, Steve. "The Darkroom Cookbook: 3rd Edition (Paperback)". Focal Press. Retrieved 2013-01-01.
  19. ↑ 19.0 19.1 "Citric acid". International Chemical Safety Cards. NIOSH.
  20. ↑ J. Zheng, F. Xiao, L.M. Qian, Z.R. Zhou (December 2009). "Erosion behavior of human tooth enamel in citric acid solution". Tribology International 42 (11–12): 1558–1564. doi:10.1016/j.triboint.2008.12.008.
  21. ↑ "Effect of Citric Acid on Tooth Enamel".
  22. ↑ British Pharmacopoeia Commission Secretariat (2009). "Index, BP 2009" (PDF). Retrieved February 4, 2010.
  23. ↑ "Japanese Pharmacopoeia, Fifteenth Edition" (PDF). 2006. Retrieved 4 Februally 2010. Check date values in: |accessdate= (help)

External links

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