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| {{About|orthophosphoric acid|other acids commonly called "phosphoric acid"|Phosphoric acids and phosphates}}
| | A fast stomach is the think of all menopausal women. One of the most distressing effects of menopause is what seems to be the abrupt appearance of unsightly plus possibly flabby abdomen weight. If you look in the mirror, you long for the days whenever the belly was flat and firm. Not just does that belly fat change how we feel inside a clothing, it also puts you at risk for certain severe health challenges.<br><br>Your 3rd measurement is what's called the kilt length. This 1 is tricky, incredibly alone, plus so I provide you permission to have somebody aid we with it. What you need to do is measure the distance from a navel (perfectly, the spot at which we took the waist measuring, wherever that could have been) to the middle of the waist hip ratio kneecap. Centuries of kilt creating have revealed a preferred way of taking this measurement. Kneel found on the ground, and keep the front of your body well straight. Hold the end of the tape at the navel, plus have your helper find the point at that it strikes the ground, while falling perfectly straight. This length is the length of the kilt.<br><br>Once you've completed which, and before we really buy any fabric, keep in your mind it is going to have to accomodate a waist hip ratio calculator kilt length. If you have a kilt length of 30", you can't buy fabric that's just 20" wide. It won't function. So, be cautious.<br><br>If you realize that loosing weight is piece of your healthy-body program, you have lots of resources to help you. Weight Watchers offers support like menus, group meetings plus weigh-ins. The ChooseMyPlate website can inform we regarding health goals, and also tricks and inspirations for planning.<br><br>But new analysis place out by the Mayo Clinic indicates which the BMI would not become the best indicator for weight following all. An athletic, muscular guy whom is 5'10" plus 200 pounds comes up "overweight" on the BMI scale as a result of his muscle mass. Muscle weighs over fat, so merely utilizing height plus fat is not an accurate measure of the person's fitness. BMIs which indicate a individual is overweight (25-29) or obese (29 and up) will affect a person's wellness and existence insurance status erroneously.<br><br>Needless to say you'll easily learn in the event you need to lose abdomen fat by looking at the mirror. However, it really is still a superior idea to check when you are having too much fat around your waist. You can check your [http://safedietplansforwomen.com/waist-to-hip-ratio waist hip ratio]. It is the ratio of your waist circumference to the hip circumference. In many cases, you need to get rid of stomach fat in the event you have a ratio larger than 0.85 (for women) or 1.0 (for men).<br><br>Gather the materials. We will need the following: measuring tape (the type chosen by a tailor, not a carpenter), a sewing needle, thread of any color best matches the desired kilt, a pencil, a excellent preponderance of protection pins, an iron, plus a can of spray-on starch. Oh, and, of course, the kilt material. How much we require usually have to be determined by how big you may be, plus how broad you create your pleats. I am an incredibly large fellow, nevertheless I furthermore used particularly broad pleats. I bought four yards of fabric, and had plenty left over. More on this later.<br><br>What utilized to be called "middle-aged spread" has become a ubiquitous phenomenon over all age groups inside America, including our youngsters. It is time for us to take a closer consider our lifestyle, for it happens to be shortening the lifetime spans. |
| {{chembox
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| | Watchedfields = changed
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| | verifiedrevid = 470622096
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| | ImageFile = Phosphoric-acid-2D-dimensions.png
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| | ImageSize = 180px
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| | ImageName = Structural formula of phosphoric acid, showing dimensions
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| | ImageFile1 = Phosphoric-acid-3D-balls.png
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| | ImageSize1 = 150px
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| | ImageName1 = Ball-and-stick model
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| | ImageFile2 = Phosphoric-acid-3D-vdW.png
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| | ImageSize2 = 150px
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| | ImageName2 = Space-filling model
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| | IUPACName = trihydroxidooxidophosphorus<br />phosphoric acid
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| | OtherNames = Orthophosphoric acid<br />trihydroxylphosphine oxide
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| | Section1 = {{Chembox Identifiers
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| | PubChem = 1004
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| | UNII_Ref = {{fdacite|correct|FDA}}
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| | UNII = E4GA8884NN
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| | KEGG_Ref = {{keggcite|correct|kegg}}
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| | KEGG = D05467
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| | InChI = 1/H3O4P/c1-5(2,3)4/h(H3,1,2,3,4)
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| | InChIKey = NBIIXXVUZAFLBC-UHFFFAOYAI
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| | ChEBI_Ref = {{ebicite|correct|EBI}}
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| | ChEBI = 26078
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| | SMILES = OP(=O)(O)O
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| | ChEMBL_Ref = {{ebicite|correct|EBI}}
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| | ChEMBL = 1187
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| | StdInChI_Ref = {{stdinchicite|correct|chemspider}}
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| | StdInChI = 1S/H3O4P/c1-5(2,3)4/h(H3,1,2,3,4)
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| | StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
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| | StdInChIKey = NBIIXXVUZAFLBC-UHFFFAOYSA-N
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| | CASNo = 7664-38-2
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| | CASNo_Ref = {{cascite|correct|CAS}}
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| | CASOther = <br />16271-20-8 (hemihydrate)
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| | ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
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| | ChemSpiderID = 979
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| | EINECS = 231-633-2
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| | UNNumber = 1805
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| | RTECS = TB6300000
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| }}
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| | Section2 = {{Chembox Properties
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| | Formula = H<sub>3</sub>PO<sub>4</sub>
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| | MolarMass = 97.995 g/mol
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| | Appearance = white solid or colourless, viscous liquid (>42 °C) <br> [[deliquescent]]
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| | Odor - odorless
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| | Density = 1.885 g/mL (liquid)<br />1.685 g/mL (85% solution)<br />2.030 g/mL (crystal at 25 °C)
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| | MeltingPt = 42.35 °C (anhydrous)<br />29.32 °C (hemihydrate)
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| | BoilingPt = 158 °C (decomposition)
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| | Solubility = 5.48 g/mL
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| | SolubleOther = soluble in [[ethanol]]
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| | pKa = 2.148, 7.198, 12.319
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| | Viscosity = 2.4–9.4 [[Poise|cP]] (85% aq. soln.)<br />147 [[Poise|cP]] (100%)
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| | RefractIndex = 1.34203
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| }}
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| | Section3 = {{Chembox Structure
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| | CrystalStruct = monoclinic
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| }}
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| | Section4 = {{Chembox Thermochemistry
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| | DeltaHf = -1288 kJ·mol<sup>−1</sup><ref name=b1>{{cite book| author = Zumdahl, Steven S.|title =Chemical Principles 6th Ed.| publisher = Houghton Mifflin Company| year = 2009| isbn = 0-618-94690-X|page=A22}}</ref>
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| | Entropy = 158 J·mol<sup>−1</sup>·K<sup>−1</sup><ref name=b1/>
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| }}
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| | Section7 = {{Chembox Hazards
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| | ExternalMSDS = [http://www.inchem.org/documents/icsc/icsc/eics1008.htm ICSC 1008]
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| | EUIndex = 015-011-00-6
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| | EUClass = Corrosive ('''C''')
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| | NFPA-H = 3
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| | NFPA-R = 0
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| | NFPA-F = 0
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| | NFPA-O =
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| | RPhrases = {{R34}}
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| | SPhrases = {{S1/2}} {{S26}} {{S45}}
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| | FlashPt = Non-flammable
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| | LD50 = 1530 mg/kg (rat, oral)
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| }}
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| | Section8 = {{Chembox Related
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| | OtherFunctn = [[Hypophosphorous acid]]<br />[[Phosphorous acid]]<br />[[Pyrophosphoric acid]]<br />[[Triphosphoric acid]]<br />[[Perphosphoric acid]]<br />[[Permonophosphoric acid]]
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| | Function = [[phosphorus]] [[oxoacid]]s
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| }}
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| }}
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| | |
| '''Phosphoric acid''' (also known as '''orthophosphoric acid''' or '''phosphoric(V) acid''') is a [[mineral acid|mineral (inorganic) acid]] having the [[chemical formula]] [[Hydrogen|H]]<sub>3</sub>[[Phosphorus|P]][[Oxygen|O]]<sub>4</sub>. Orthophosphoric [[acid]] molecules can combine with themselves to form a variety of compounds which are also referred to as '''phosphoric acids''', but in a more general way. The term ''phosphoric acid'' can also refer to a [[chemical]] or [[reagent]] consisting of phosphoric acids, such as ''[[pyrophosphoric acid]]'' or ''[[triphosphoric acid]]'', but usually orthophosphoric acid.
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| The [[conjugate acid|conjugate base]] of phosphoric acid is the [[phosphate#Chemical properties|dihydrogen phosphate]] ion, {{chem|H|2|PO|4|−}}, which in turn has a conjugate base of [[phosphate#Chemical properties|hydrogen phosphate]], {{chem|HPO|4|2−}}, which has a conjugate base of [[phosphate]], {{chem|PO|4|3−}}.
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| In addition to being a chemical reagent, phosphoric acid has a wide variety of uses, including as a rust inhibitor, food additive, dental and orthop(a)edic etchant, electrolyte, flux, dispersing agent, industrial etchant, fertilizer feedstock, and component of home cleaning products.
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| The most common source of phosphoric acid is an 85% [[aqueous]] [[solution]]; such solutions are colourless, odourless, and non-[[Volatility (chemistry)|volatile]]. Rather viscous, syrupy [[liquid]]s, but still pourable. Because it is a concentrated acid, an 85% solution can be [[corrosive]], although nontoxic when diluted. Because of the high percentage of phosphoric acid in this reagent, at least some of the orthophosphoric acid is condensed into [[polyphosphoric acid]]s. For the sake of labeling and simplicity, the 85% represents H<sub>3</sub>PO<sub>4</sub> as if it were all orthophosphoric acid. Dilute aqueous solutions of phosphoric acid exist in the ortho- form.
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| ==Reactions==
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| Pure [[anhydrous]] phosphoric acid is a white solid that melts at 42.35 °C to form a colorless, viscous liquid.
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| Orthophosphoric acid as ''phosphoric acid'', which is the [[IUPAC nomenclature|IUPAC name]] for this compound. The prefix ''[[ortho-|ortho]]'' is used to distinguish the acid from related phosphoric acids, called polyphosphoric acids. Orthophosphoric acid is a non-[[toxicity|toxic]], [[inorganic]], rather weak triprotic [[acidity|acid]], which, when pure, is a [[solid]] at room [[temperature]] and [[pressure]]. The [[chemical structure]] of orthophosphoric acid is shown above in the data table. Orthophosphoric acid is a very [[Chemical polarity|polar]] [[molecule]]; therefore it is highly soluble in water. The [[oxidation state]] of [[phosphorus]] (P) in ortho- and other phosphoric acids is +5; the oxidation state of all the [[oxygen]] atoms (O) is −2 and all the [[hydrogen]] atoms (H) is +1. Triprotic means that an orthophosphoric acid molecule can dissociate up to three times, giving up an H<sup>+</sup> each time, which typically combines with a [[Water (molecule)|water molecule]], H<sub>2</sub>O, as shown in these [[chemical reaction|reactions]]:
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| :H<sub>3</sub>PO<sub>4</sub><sub>(s)</sub> + H<sub>2</sub>O<sub>(l)</sub> {{eqm}} H<sub>3</sub>O<sup>+</sup><sub>(aq)</sub> + H<sub>2</sub>PO<sub>4</sub><sup>−</sup><sub>(aq)</sub> ''K''<sub>a1</sub>= 7.25×10<sup>−3</sup>
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| :H<sub>2</sub>PO<sub>4</sub><sup>−</sup><sub>(aq)</sub>+ H<sub>2</sub>O<sub>(l)</sub> {{eqm}} H<sub>3</sub>O<sup>+</sup><sub>(aq)</sub> + HPO<sub>4</sub><sup>2−</sup><sub>(aq)</sub> ''K''<sub>a2</sub>= 6.31×10<sup>−8</sup>
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| :HPO<sub>4</sub><sup>2−</sup><sub>(aq)</sub>+ H<sub>2</sub>O<sub>(l)</sub> {{eqm}} H<sub>3</sub>O<sup>+</sup><sub>(aq)</sub> + PO<sub>4</sub><sup>3−</sup><sub>(aq)</sub> ''K''<sub>a3</sub>= 4.80×10<sup>−13</sup>
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| The [[ion#anions and cations|anion]] after the first dissociation, H<sub>2</sub>PO<sub>4</sub><sup>−</sup>, is the ''dihydrogen phosphate'' anion. The anion after the second dissociation, HPO<sub>4</sub><sup>2−</sup>, is the ''hydrogen phosphate'' anion. The anion after the third dissociation, PO<sub>4</sub><sup>3−</sup>, is the '''[[phosphate]]''' or '''orthophosphate''' anion. For each of the dissociation reactions shown above, there is a separate [[acid dissociation constant]], called ''K''<sub>a1</sub>, ''K''<sub>a2</sub>, and ''K''<sub>a3</sub> given at 25 °C. Associated with these three dissociation constants are corresponding p''K''<sub>a1</sub>=2.12, p''K''<sub>a2</sub>=7.21, and p''K''<sub>a3</sub>=12.67 values at 25 °C. Even though all three [[hydrogen]] (H) atoms are equivalent on an orthophosphoric acid molecule, the successive ''K''<sub>a</sub> values differ since it is energetically less favorable to lose another H<sup>+</sup> if one (or more) has already been lost and the molecule/ion is more negatively charged.
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| Because the triprotic dissociation of orthophosphoric acid, the fact that its [[conjugate base]]s (the phosphates mentioned above) cover a wide [[pH]] range, and, because phosphoric acid/phosphate [[solution]]s are, in general, non-toxic, mixtures of these types of phosphates are often used as [[buffering agents]] or to make [[buffer solution]]s, where the desired pH depends on the proportions of the phosphates in the mixtures. Similarly, the non-toxic, [[ion#anions and cations|anion]] [[salt]]s of triprotic [[organic compound|organic]] [[citric acid]] are also often used to make buffers. Phosphates are found pervasively in biology, especially in the compounds derived from phosphorylated [[sugar]]s, such as [[DNA]], [[RNA]], and [[adenosine triphosphate]] (ATP). There is a separate article on [[phosphate]] as an anion or its salts.
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| Upon heating orthophosphoric acid, condensation of the phosphoric units can be induced by driving off the water formed from condensation. When one molecule of water has been removed for each two molecules of phosphoric acid, the result is [[pyrophosphoric acid]] (H<sub>4</sub>P<sub>2</sub>O<sub>7</sub>). When an average of one molecule of water per phosphoric unit has been driven off, the resulting substance is a glassy solid having an empirical formula of '''HPO<sub>3</sub>''' and is called '''metaphosphoric acid'''.<ref>[http://web.archive.org/web/20070403074509/http://www.bartleby.com/65/ph/phsphracid.html phosphoric acid]. The Columbia Encyclopedia, Sixth Edition.</ref> Metaphosphoric acid is a singly anhydrous version of orthophosphoic acid and is sometimes used as a water- or moisture-absorbing reagent. Further [[dehydration|dehydrating]] is very difficult, and can be accomplished only by means of an extremely strong [[desiccant]] (and not by heating alone). It produces ''[[phosphorus pentoxide|phosphoric anhydride]]'' (phosphorus pentoxide), which has an empirical formula P<sub>2</sub>O<sub>5</sub>, although an actual molecule has a chemical formula of P<sub>4</sub>O<sub>10</sub>. Phosphoric anhydride is a solid, which is very strongly moisture-absorbing and is used as a [[desiccant]].
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| In the presence of superacids (acids stronger than [[Sulfuric Acid|{{chem|H|2|S|O|4}}]]), {{chem|H|3|P|O|4}} reacts to form mystery products, perhaps corrosive, acidic salts of the hypothetical<ref>{{cite doi|10.1063/1.475628}}</ref> '''tetrahydroxylphosphonium''' ion, which is [[isoelectronic]] with [[silicic acid|orthosilicic acid]]. The suspected reaction with [[fluoroantimonic acid|{{chem|H|Sb|F|6}}]], for example, is supposed to go:
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| :{{chem|H|3|P|O|4}} + {{chem|H|Sb|F|6}} <nowiki>→</nowiki> [{{chem|P}}({{chem|O|H}})<sub>4</sub><sup>+</sup>] [{{chem|Sb|F|6}}]<sup>−</sup>
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| ===Aqueous solution===
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| For a given total acid concentration [A] = [H<sub>3</sub>PO<sub>4</sub>] + [H<sub>2</sub>PO<sub>4</sub><sup>−</sup>] + [HPO<sub>4</sub><sup>2−</sup>] + [PO<sub>4</sub><sup>3−</sup>] ([A] is the total number of moles of pure H<sub>3</sub>PO<sub>4</sub> which have been used to prepare 1 liter of solution), the composition of an aqueous solution of phosphoric acid can be calculated using the equilibrium equations associated with the three reactions described above together with the [H<sup>+</sup>] [OH<sup>−</sup>] = 10<sup>−14</sup> relation and the electrical neutrality equation. Possible concentrations of polyphosphoric molecules and ions is neglected. The system may be reduced to a fifth degree equation for [H<sup>+</sup>] which can be solved numerically, yielding:
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| {| class="wikitable"
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| ! style="width:60px; text-align:center;"|'''[A] (mol/L)'''
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| ! style="width:50px; text-align:center;"|'''pH'''
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| ! style="width:110px; text-align:center;"|'''[H<sub>3</sub>PO<sub>4</sub>]/[A] (%)'''
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| ! style="width:110px; text-align:center;"|'''[H<sub>2</sub>PO<sub>4</sub><sup>−</sup>]/[A] (%)'''
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| ! style="width:110px; text-align:center;"|'''[HPO<sub>4</sub><sup>2−</sup>]/[A] (%)'''
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| ! style="width:110px; text-align:center;"|'''[PO<sub>4</sub><sup>3−</sup>]/[A] (%)'''
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| |-
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| |1||1.08||'''91.7'''||8.29||6.20×10<sup>−6</sup>||1.60×10<sup>−17</sup>
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| |10<sup>−1</sup>||1.62||'''76.1'''||23.9||6.20×10<sup>−5</sup>||5.55×10<sup>−16</sup>
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| |-
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| |10<sup>−2</sup>||2.25||'''43.1'''||'''56.9'''||6.20×10<sup>−4</sup>||2.33×10<sup>−14</sup>
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| |-
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| |10<sup>−3</sup>||3.05||10.6||'''89.3'''||6.20×10<sup>−3</sup>||1.48×10<sup>−12</sup>
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| |-
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| |10<sup>−4</sup>||4.01||1.30||'''98.6'''||6.19×10<sup>−2</sup>||1.34×10<sup>−10</sup>
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| |-
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| |10<sup>−5</sup>||5.00||0.133||'''99.3'''||0.612||1.30×10<sup>−8</sup>
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| |-
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| |10<sup>−6</sup>||5.97||1.34×10<sup>−2</sup>||'''94.5'''||5.50||1.11×10<sup>−6</sup>
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| |-
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| |10<sup>−7</sup>||6.74||1.80×10<sup>−3</sup>||'''74.5'''||25.5||3.02×10<sup>−5</sup>
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| |-
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| |10<sup>−10</sup>||7.00||8.24×10<sup>−4</sup>||'''61.7'''||'''38.3'''||8.18×10<sup>−5</sup>
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| |}
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| For strong acid concentrations, the solution is mainly composed of H<sub>3</sub>PO<sub>4</sub>. For [A] = 10<sup>−2</sup>, the pH is close to p''K''<sub>a1</sub>, giving an equimolar mixture of H<sub>3</sub>PO<sub>4</sub> and H<sub>2</sub>PO<sub>4</sub><sup>−</sup>. For [A] below 10<sup>−3</sup>, the solution is mainly composed of H<sub>2</sub>PO<sub>4</sub><sup>−</sup> with [HPO<sub>4</sub><sup>2−</sup>] becoming non negligible for very dilute solutions. [PO<sub>4</sub><sup>3−</sup>] is always negligible. Since this analysis does not take into account ion activity coefficients, the pH and molarity of a real phosphoric acid solution may deviate substantially from the above values.
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| ==Preparation==
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| Phosphoric acid can be prepared by three routes – the thermal process and the wet process. The wet process dominates in the commercial sector. The more expensive thermal process produces a purer product that is used for applications in the food industry.
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| ===Wet===
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| Wet process phosphoric acid is prepared by adding [[sulfuric acid]] to [[tricalcium phosphate]] rock, typically found in nature as [[apatite]]. The reaction is:
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| :Ca<sub>5</sub>(PO<sub>4</sub>)<sub>3</sub>X + 5 H<sub>2</sub>SO<sub>4</sub> + 10 H<sub>2</sub>O → 3 H<sub>3</sub>PO<sub>4</sub> + 5 CaSO<sub>4</sub>·2H<sub>2</sub>O + HX
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| :where X may include OH, F, Cl, and Br
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| The initial phosphoric acid solution may contain 23–33% P<sub>2</sub>O<sub>5</sub>, but can be concentrated by the evaporation of water to produce ''commercial-'' or ''merchant-grade'' phosphoric acid, which contains about 54% [[P2O5|P<sub>2</sub>O<sub>5</sub>]]. Further evaporation of water yields ''superphosphoric acid'' with a P<sub>2</sub>O<sub>5</sub> concentration above 70%.<ref>Thomas, W P and Lawton, W S "Stable ammonium polyphosphate liquid fertilizer from merchant grade phosphoric acid" {{US Patent|4721519}}, Issue date: January 26, 1988</ref><ref>{{cite web |url=http://msds.simplot.com/datasheets/12002.pdf |title=Super Phosphoric Acid 0-68-0 Material Safety Data Sheet |publisher= J.R. Simplot Company |accessdate=4 May 2010 |date=May 2009}}</ref>
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| Digestion of the phosphate ore using sulfuric acid yields the insoluble [[calcium sulfate]] (gypsum), which is filtered and removed as [[phosphogypsum]]. Wet-process acid can be further purified by removing fluorine to produce animal-grade phosphoric acid, or by solvent extraction and arsenic removal to produce food-grade phosphoric acid.
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| The [[nitrophosphate process]] is similar to the wet process except that it uses nitric acid in place of sulfuric acid. The advantage to this route is that the coproduct, calcium nitrate is also a plant fertilizer. This method is rarely employed.
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| ===Thermal===
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| This very pure phosphoric acid is obtained by burning elemental [[phosphorus]] to produce [[phosphorus pentoxide]], which is subsequently dissolved in dilute phosphoric acid. This route produces a very pure phosphoric acid, since most impurities present in the rock have been removed when extracting phosphorus from the rock in a furnace. The end result is food-grade, thermal phosphoric acid; however, for critical applications, additional processing to remove arsenic compounds may be needed.
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| Elemental [[phosphorus]] is produced by an electric furnace. At a high temperature, a mixture of phosphate ore, silica and carbonaceous material (coke, coal etc...) produces calcium silicate, phosphorus gas and [[carbon monoxide]]. The P and CO off-gases from this reaction are cooled under water to isolate solid phosphorus. Alternatively, the P and CO off-gases can be burned with air to produce [[phosphorus pentoxide]] and carbon dioxide.
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| ==Uses==
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| The dominant use of phosphoric acid is for [[fertilizer]]s, consuming approximately 90% of production.<ref name=Ullmann>Klaus Schrödter, Gerhard Bettermann, Thomas Staffel, Friedrich Wahl, Thomas Klein, Thomas Hofmann "Phosphoric Acid and Phosphates" in ''Ullmann’s Encyclopedia of Industrial Chemistry'' 2008, Wiley-VCH, Weinheim. {{DOI|10.1002/14356007.a19_465.pub3}}</ref>
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| {| class="wikitable"
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| |-
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| ! application
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| ! demand (2006) in thousands of tons
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| ! main phosphate derivatives
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| | soaps and detergents ||1836 || [[Sodium triphosphate|STPP]]
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| | food industry|| 309 ||[[Sodium triphosphate|STPP]] (Na<sub>5</sub>P<sub>3</sub>O<sub>10</sub>), [[Sodium hexametaphosphate|SHMP]], [[Trisodium phosphate|TSP]], [[Disodium pyrophosphate|SAPP]], [[Sodium aluminium phosphate|SAlP]] (NaA, [[Monocalcium phosphate|MCP]], [[Disodium phosphate|DSP]] (Na<sub>2</sub>HPO<sub>4</sub>), H<sub>3</sub>PO<sub>4</sub>
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| | [[water treatment]] || 164 ||SHMP, [[Sodium triphosphate|STPP]], [[Tetrasodium pyrophosphate|TSPP]], [[Monosodium phosphate|MSP]] (NaH<sub>2</sub>PO<sub>4</sub>), DSP
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| | [[toothpaste]]s || 68 || [[Dicalcium phosphate|DCP]] (CaHPO<sub>4</sub>), IMP, SMFP
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| | other applications || 287 || [[Sodium triphosphate|STPP]] (Na<sub>3</sub>P<sub>3</sub>O<sub>9</sub>), TCP, APP, DAP, [[zinc phosphate]] (Zn<sub>3</sub>(PO<sub>4</sub>)<sub>2</sub>), [[aluminium phosphate]] (AlPO<sub>4</sub>, H<sub>3</sub>PO<sub>4</sub>)
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| |}
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| ===Food additive===
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| Food-grade phosphoric acid (additive [[E number|E338]]<ref>{{cite web|url=http://www.food.gov.uk/policy-advice/additivesbranch/enumberlist#h_7|title=Current EU approved additives and their E Numbers|date=14 March 2012|publisher=Foods Standards Agency|accessdate=22 July 2012}}</ref>) is used to acidify foods and beverages such as various [[cola]]s, but not without controversy regarding its health effects. It provides a tangy or sour taste, and being a mass-produced chemical is available cheaply and in large quantities. The low cost and bulk availability is unlike more expensive seasonings that give comparable flavors, such as [[citric acid]] which is obtainable from [[citrus]], but usually fermented by ''[[Aspergillus niger]]'' mold from scrap [[molasses]], waste [[starch]] [[hydrolysis|hydrolysates]] and phosphoric acid.<ref>{{cite pmid|17317159 }}</ref> Various phosphates, e.g., [[monocalcium phosphate]], are used as [[leavening agent]]s.
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| ===Niche uses===
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| Phosphoric acid and its derivatives are pervasive and find many niche applications.
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| ====Rust removal====
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| Phosphoric acid may be used as a "rust converter", by direct application to rusted iron, steel tools, or surfaces. The phosphoric acid converts reddish-brown [[iron(III) oxide]], Fe<sub>2</sub>O<sub>3</sub> ([[rust]]) to black [[Iron(III) phosphate|ferric phosphate]], FePO<sub>4</sub>.
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| An empirical formula for this reaction is:
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| :<math> 2 H_3PO_4 + Fe_2 O_3 \longrightarrow 2 FePO_4 + 3 H_2 O </math>
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|
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| "Rust converter" is sometimes a greenish liquid suitable for dipping (in the same sort of acid bath as is used for [[Pickling (metal)|pickling metal]]), but it is more often formulated as a [[gel]], commonly called "naval jelly". It is sometimes sold under other names, such as "rust remover" or "rust killer". As a thick gel, it may be applied to sloping, vertical, or even overhead surfaces.
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| After treatment, the black ferric phosphate coating can be scrubbed off, leaving a fresh metal surface. Multiple applications of phosphoric acid may be required to remove all rust. The black phosphate coating can also be left in place, where it will provide moderate further corrosion resistance (such protection is also provided by the superficially similar [[Parkerize|Parkerizing]] and [[bluing (steel)|blued]] electrochemical [[conversion coating]] processes).
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| ====In medicine====
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| Phosphoric acid is used in [[dentistry]] and [[orthodontics]] as an [[etching]] solution, to clean and roughen the surfaces of teeth where dental appliances or fillings will be placed.
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| Phosphoric acid is also an ingredient in over-the-counter anti-nausea medications that also contain high levels of [[sugar]] ([[glucose]] and [[fructose]]). This acid is also used in many teeth whiteners to eliminate plaque that may be on the teeth before application.
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| ====Other applications====
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| Among other applications, phosphoric acid is used:
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| * As an external standard for phosphorus-31 [[Nuclear magnetic resonance]] (NMR).
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| * For [[high-performance liquid chromatography]].
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| * As a chemical oxidizing agent for [[activated carbon]] production, as used in the Wentworth Process.<ref>{{cite doi|10.1016/S0008-6223(96)00093-0}}</ref>
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| * As the [[electrolyte]] in [[phosphoric acid fuel cell]]s.
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| * With distilled water (2–3 drops per gallon) as an electrolyte in oxyhydrogen generators.
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| * As a catalyst in the hydration of alkenes to produce alcohols, predominantly ethanol.
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| * As an [[electrolyte]] in [[copper]] [[electropolishing]] for burr removal and circuit board planarization.
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| * As a [[flux (metallurgy)|flux]] by hobbyists (such as model railroaders) as an aid to [[soldering]].
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| * In [[compound semiconductor]] processing, phosphoric acid is a common wet etching agent: for example, in combination with hydrogen peroxide and water it is used to etch [[Indium gallium arsenide|InGaAs]] selective to [[indium phosphide|InP]].<ref>[http://terpconnect.umd.edu/~browns/wetetch.html Wet chemical etching.] umd.edu</ref>
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| * Heated in [[microfabrication]] to etch [[silicon nitride]] (Si<sub>3</sub>N<sub>4</sub>). It is highly selective in etching Si<sub>3</sub>N<sub>4</sub> instead of SiO<sub>2</sub>, [[silicon dioxide]].<ref name="Wolf">{{cite book |title =Silicon processing for the VLSI era: Volume 1 – Process technology |last = Wolf |first = S. |coauthors = R.N. Tauber |year=1986 |page=534 |isbn=0-9616721-6-1}}</ref>
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| * As a cleaner by [[construction]] trades to remove mineral deposits, cementitious smears, and hard water stains.
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| * As a [[chelation|chelant]] in some household cleaners aimed at similar cleaning tasks.
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| * In [[hydroponics]] pH solutions to lower the pH of nutrient solutions. While other types of acids can be used, phosphorus is a nutrient used by plants, especially during flowering, making phosphoric acid particularly desirable.
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| * As a pH adjuster in cosmetics and skin-care products.<ref>{{cite web|publisher = Paula's Choice|title = Ingredient dictionary: P|work = Cosmetic ingredient dictionary|accessdate = 16 November 2007|url = http://www.cosmeticscop.com/learn/cosmetic_dictionary.asp?id=21&letter=P}}</ref>
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| * As a dispersing agent in detergents and leather treatment.
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| * As an additive to stabilize acidic aqueous solutions within a wanted and specified pH range.
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| ==Biological effects==
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| {{Expand section|1=health effects beyond soft drinks|date=January 2013}}
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| ===In soft drinks===
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| Phosphoric acid, used in many soft drinks (primarily [[cola]]), has been linked in epidemiological studies to (1) chronic kidney disease and (2) lower bone density.
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| (1)
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| A study performed by the Epidemiology Branch of the US National Institute of Environmental Health Sciences, concludes that drinking 2 or more colas per day was associated with doubling the risk of chronic kidney disease.<ref>http://www.ncbi.nlm.nih.gov/pubmed/17525693?ordinalpos=1&itool=EntrezSystem2.PEntrez.Pubmed.Pubmed_ResultsPanel.Pubmed_DefaultReportPanel.Pubmed_RVDocSum</ref>
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| (2)
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| A study<ref>{{cite pmid|17023723}}</ref> using dual-energy X-ray absorptiometry rather than a questionnaire about breakage, provides reasonable evidence to support the theory that drinking cola results in lower bone density. This study was published in the [[American Journal of Clinical Nutrition]]. A total of 1672 women and 1148 men were studied between 1996 and 2001. Dietary information was collected using a food frequency questionnaire that had specific questions about the number of servings of cola and other carbonated beverages and that also made a differentiation between regular, caffeine-free, and diet drinks. The paper cites significant statistical evidence to show that women who consume cola daily have lower bone density. Total phosphorus intake was not significantly higher in daily cola consumers than in nonconsumers; however, the calcium-to-phosphorus ratios were lower.
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| On the other hand, another study suggests that insufficient intake of [[phosphorus]] leads to lower [[bone density]]. The study does not examine the effect of phosphoric acid, which binds with [[magnesium]] and [[calcium]] in the digestive tract to form salts that are not absorbed, but rather studies general phosphorus intake.<ref>{{cite pmid|10024903}}</ref>
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| A clinical study by Heaney and Rafferty using calcium-balance methods found no impact of carbonated soft drinks containing phosphoric acid on calcium excretion.<ref name="heaney">{{cite pmid|11522558}}</ref> The study compared the impact of water, milk, and various soft drinks (two with caffeine and two without; two with phosphoric acid and two with citric acid) on the calcium balance of 20- to 40-year-old women who customarily consumed ~3 or more cups (680 mL) of a carbonated soft drink per day. They found that, relative to water, only milk and the two caffeine-containing soft drinks increased urinary calcium, and that the calcium loss associated with the caffeinated soft drink consumption was about equal to that previously found for caffeine alone. Phosphoric acid without [[caffeine]] had no impact on urine calcium, nor did it augment the urinary calcium loss related to caffeine. Because studies have shown that the effect of caffeine is compensated for by reduced calcium losses later in the day,<ref>{{cite pmid|2403065}}</ref> Heaney and Rafferty concluded that the net effect of carbonated beverages—including those with caffeine and phosphoric acid—is negligible, and that the skeletal effects of carbonated soft drink consumption are likely due primarily to milk displacement.
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| Other chemicals such as [[caffeine]] (also a significant component of popular common cola drinks) were also suspected as possible contributors to low bone density, due to the known effect of caffeine on [[calciuria]]. One other study, involving 30 women over the course of a week, suggests that phosphoric acid in colas has no such effect, and postulates that caffeine has only a temporary effect, which is later reversed. The authors of this study conclude that the skeletal effects of carbonated beverage consumption are likely due primarily to milk displacement<ref name="heaney"/> (another possible [[confounding factor]] may be an association between high soft drink consumption and sedentary lifestyle){{Citation needed|date=January 2013}}.
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| ==See also==
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| * Phosphate [[fertilizer]]s, such as [[ammonium phosphate]] fertilizers
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| ==References==
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| {{reflist|35em}}
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| ==External links==
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| {{Commons category|phosphoric acid}}
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| * [http://www.inchem.org/documents/icsc/icsc/eics1008.htm International chemical safety card 1008]
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| * [http://www.npi.gov.au/substances/phosphoric-acid/index.html National pollutant inventory – Phosphoric acid fact sheet]
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| * [http://www.cdc.gov/niosh/npg/npgd0506.html NIOSH Pocket guide to chemical hazards]
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| * [http://www2.iq.usp.br/docente/gutz/Curtipot_.html Excel spreadsheet containing phosphoric acid titration curve, distribution diagram and buffer pH calculation]
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| {{Hydrogen compounds}}
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| {{Use dmy dates|date=October 2010}}
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| {{DEFAULTSORT:Phosphoric Acid}}
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| [[Category:Mineral acids]]
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| [[Category:Flavors]]
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| [[Category:Food acidity regulators]]
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| [[Category:Food antioxidants]]
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| [[Category:Glassforming liquids and melts]]
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| [[Category:Phosphates]]
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| [[Category:Phosphorus oxoacids]]
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| [[Category:Hydrogen compounds]]
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