PH
24530
226169751
2008-07-17T03:38:44Z
Vsmith
84417
Reverted edits by [[Special:Contributions/76.123.221.182|76.123.221.182]] ([[User talk:76.123.221.182|talk]]) to last version by Closedmouth
{{otheruses|PH (disambiguation)}}
{{lowercase|title=pH}}
{{Acids and Bases}}
'''''pH''''' is the measure of the [[Acid|acidity]] or [[Base (chemistry)|alkalinity]] of a [[solution]]. It is formally a measure of the [[Activity (chemistry)|activity]] of dissolved [[hydrogen ion]]s (H<sup>+</sup>), but for very dilute solutions, the [[Concentration#Molarity|molarity]] ([[molar concentration]]) of H<sup>+</sup> may be used as a substitute with little loss of accuracy.<ref>[http://chem.lapeer.org/Chem2Docs/pHFacts.html pH Facts<!-- Bot generated title -->]</ref> In solution, hydrogen ions occur as a number of [[cation]]s including [[hydronium]] ions (H<sub>3</sub>O<sup>+</sup>).<ref>[http://www.jp.horiba.com/story_e/ph/ph01_03.htm HORIBA : The story of pH<!-- Bot generated title -->]</ref>
In pure water at 25 [[celsius|°C]], the concentration of H<sup>+</sup> equals the concentration of [[hydroxide]] ions (OH<sup>-</sup>). This is defined as "neutral" and corresponds to a pH level of 7.0. Solutions in which the concentration of H<sup>+</sup> exceeds that of OH<sup>-</sup> have a pH value lower than 7.0 and are known as acids. Solutions in which OH<sup>-</sup> exceeds H<sup>+</sup> have a pH value greater than 7.0 known as bases. Because pH is dependent on ionic activity, a property which cannot be measured easily or fully predicted theoretically, it is difficult to determine an accurate value for the pH of a solution. The pH reading of a solution is usually obtained by comparing unknown solutions to those of known pH, and there are [[#Measurement|several ways]] to do so.
The concept of pH was first introduced by [[Danish people|Danish]] [[chemist]] [[S. P. L. Sørensen]] at the [[Carlsberg Laboratory]] in [[1909]].<ref>Carlsberg Research Centre history page, http://www.crc.dk/history.shtml</ref> Sørensen suggested the notation "PH" for convenience, standing for "power of hydrogen" <ref>Carlsberg Group Company History Page, http://www.carlsberggroup.com/Company/Research/Pages/pHValue.aspx</ref>, using the negative logarithm of the concentration of hydrogen ions in solution <ref>Sørensen, http://www.geocities.com/bioelectrochemistry/sorensen.htm</ref>.
== Definition ==
The [[operational definition]] of pH is officially defined by International Standard [[ISO 31-8]] as follows: <ref>Quantities and units – Part 8: Physical chemistry and molecular physics, Annex C (normative): pH. [[International Organization for Standardization]], 1992.</ref> For a solution X, first measure the [[electromotive force]] ''E''<sub>X</sub> of the [[galvanic cell]]
:reference electrode | concentrated solution of KCl || solution X | H<sub>2</sub> | Pt
and then also measure the electromotive force ''E''<sub>S</sub> of a galvanic cell that differs from the above one only by the replacement of the solution X of unknown pH, pH(X), by a solution S of a known standard pH, pH(S). The pH of X is then
:<math>\text{pH(X)} = \text{pH(S)} + \frac{(E_{\text{S}} - E_{\text{X}})F}{RT \ln 10}</math>
where
:''F'' is the [[Faraday constant]];
:''R'' is the [[molar gas constant]];
:''T'' is the [[thermodynamic temperature]].
Defined this way, pH is a [[dimensionless quantity]]. Values pH(S) for a range of standard solutions S, along with further details, are given in the relevant [[IUPAC]] recommendation.<ref>[http://www.iupac.org/publications/pac/1985/pdf/5703x0531.pdf Definitions of pH scales, standard reference values, measurement of pH, and related terminology]. Pure Appl. Chem. (1985), 57, pp 531–542.</ref>
pH has no fundamental meaning as a unit; its official definition is a practical one. However in the restricted range of dilute [[aqueous solution]]s having an amount-of-dissolved-substance concentrations less than 0.1 mol/L, and being neither strongly [[alkaline]] nor strongly [[acidic]] (2 < pH < 12), the definition is such that
:<math>\text{pH} = -\log_{10}\left[\frac{\gamma_1 [\text{H}^+]) }{ \text{1 mol L}^{-1} } \right] \pm 0.02</math>
where [H<sup>+</sup>] denotes the amount-of-substance concentration of hydrogen ion H<sup>+</sup> and ''γ''<sub>1</sub> denotes the [[activity coefficient]] of a typical [[univalent]] [[electrolyte]] in the solution.
;Simplified definition
pH is a measurement of the concentration of hydrogen ions in a solution. Because of its mathematical formulation, low pH values are associated with solutions with high concentrations of hydrogen ions, while high pH values occur for solutions with low concentrations of hydrogen ions. Pure water has a pH of 7.0, and other solutions are usually described with reference to this value. Acids are defined as those solutions that have a pH less than 7 (i.e. more hydrogen ions than water); while bases are defined as those solutions that have a pH greater than 7 (i.e. less hydrogen ions than water).
The definitions of [[weak acid|weak]] and [[strong acid]]s, and [[weak base|weak]] and [[strong base]]s do not refer to pH, but instead describe whether an acid or base [[ionization|ionizes]] in solution.
==Explanation==
[[Image:PH scale.png|thumb|right|Visual representation of the pH scale.]]
[[Image:PH scale 2.png|thumb|right|Another visual representation of the pH scale.]]
In simpler terms, the number arises from a measure of the [[Activity (chemistry)|activity]] of [[hydrogen]] ions or their equivalent in the solution. The pH scale is an inverse [[logarithm]]ic representation of hydrogen proton ([[Hydrogen ion|H<sup>+</sup>]]) concentration. Unlike linear scales, which have a constant relationship between the item being measured (H<sup>+</sup> concentration in this case) and the value reported, each individual pH unit is a factor of 10 different than the next higher or lower unit. For example, a change in pH from 2 to 3 represents a 10-fold decrease in H<sup>+</sup> concentration, and a shift from 2 to 4 represents a one-hundred (10 × 10)-fold decrease in H<sup>+</sup> concentration. The formula for calculating pH is:
:<math>\mbox{pH} = -\log_{10} \alpha_{\mathrm{H}^+}</math>
Where ''α''<sub>H<sup>+</sup></sub> denotes the activity of H<sup>+</sup> ions, and is [[Dimensionless number|dimensionless]]. In solutions containing other ions, activity and concentration will not generally be the same. Activity is a measure of the effective concentration of hydrogen ions, rather than the actual concentration; it includes the fact that other ions surrounding hydrogen ions will shield them and affect their ability to participate in chemical reactions. These other ions change the effective amount of hydrogen ion concentration in any process that involves H<sup>+</sup>.
In dilute solutions such as [[tap water]], activity is approximately equal to the numeric value of the concentration of the H<sup>+</sup> ion, denoted as [H<sup>+</sup>] (<nowiki>[</nowiki>[[hydronium|H<sub>3</sub>O<sup>+</sup>]]]), measured in [[mole (unit)|moles]] per [[litre]] (also known as [[molarity]]). Therefore, it is often convenient to define pH as:
:<math>\mbox{pH} \approx -\log_{10}{\frac{[\mathrm{H^+}]}{1~\mathrm{mol/L}}} </math>
For both definitions, log<sub>10</sub> denotes the base-10 [[logarithm]], therefore pH defines a [[logarithmic scale]] of acidity. For example, if one makes a [[lemonade]] with a H<sup>+</sup> concentration of 0.0050 moles per litre, its pH would be:
:<math>\mbox{pH}_{\mathrm{lemonade}} \approx -\log_{10}{(0.0050)} \approx 2.3</math>
A solution of pH = 8.2 will have an [H<sup>+</sup>] concentration of 10<sup>−8.2</sup> mol/L, or about 6.31 × 10<sup>−9</sup> mol/L. Thus, its hydrogen activity α<sub>H<sup>+</sup></sub> is around 6.31 × 10<sup>−9</sup>. A solution with an [H<sup>+</sup>] concentration of 4.5 × 10<sup>−4</sup> mol/L will have a pH value of 3.35.
In solution at 25 °C, a pH of 7 indicates neutrality (i.e. the pH of pure water) because [[water]] naturally dissociates into H<sup>+</sup> and OH<sup>−</sup> ions with equal concentrations of 1×10<sup>−7</sup> mol/L. A lower pH value (for example pH 3) indicates increasing strength of acidity, and a higher pH value (for example pH 11) indicates increasing strength of basicity. Note, however, that pure water, when exposed to the atmosphere, will take in [[carbon dioxide]], some of which reacts with water to form [[carbonic acid]] and H<sup>+</sup>, thereby lowering the pH to about 5.7.
Neutral pH at 25 °C is not ''exactly'' 7. pH is an experimental value, so it has an associated error. Since the [[dissociation constant]] of water is (1.011 ± 0.005) × 10<sup>−14</sup>, pH of water at 25 °C would be 6.998 ± 0.001. The value is consistent, however, with neutral pH being 7.00 to two [[significant figure]]s, which is near enough for most people to assume that it is exactly 7. The pH of water gets smaller with higher temperatures. For example, at 50 °C, pH of water is 6.55 ± 0.01. This means that a diluted solution is neutral when its pH at 50 °C is around 6.55, and also that a pH of 7.00 is very slightly basic.
Most substances have a pH in the range 0 to 14, although [[Superacid|extremely acidic]] or [[Superbase|extremely basic]] substances may have pH less than 0 or greater than 14. An example is acid mine runoff, with a pH = –3.6. Note that this does not translate to a molar concentration of 3981 M; such high activity values are the result of the extremely high value of the activity coefficient while concentrations are within a "reasonable" range.<ref name="Nordstrom">Nordstrom, DK ''et al'' (2000) Negative pH and extremely acidic mine waters from Iron Mountain California. ''Environ Sci Technol'','''34''', 254-258.</ref> E.g. a 7.622 molal H<sub>2</sub>SO<sub>4</sub> solution has a pH = -3.13, hydrogen activity α<sub>H<sup>+</sup></sub> around 1350 and [[activity coefficient]] γ<sub>H<sup>+</sup></sub> = 165.4 when using the MacInnes convention for scaling Pitzer single ion activity coefficient.<ref name="Nordstrom"/>
Arbitrarily, the pH is <math>-\log_{10}{([\mbox{H}^+])}</math>. Therefore,
:<math>\mbox{pH} = -\log_{10}{[{\mbox{H}^+}]}</math>
or, by substitution,
:<math>\mbox{pH} = \frac{\epsilon}{0.059}</math>.
The "pH" of any other substance may also be found (e.g. the potential of silver ions, or pAg<sup>+</sup>) by deriving a similar equation using the same process. These other equations for potentials will not be the same, however, as the number of moles of electrons transferred (n) will differ for the different reactions.
== Calculation of pH for weak and strong acids ==
Values of pH weak and strong acids can be approximated using certain theories and assumptions.
Under the [[acid-base reaction theories|Brønsted-Lowry theory]], stronger or weaker acids are a relative concept. But here we define a strong acid as a species which is a much stronger acid than the hydronium (H<sub>3</sub>O<sup>+</sup>) ion. In that case the dissociation reaction (strictly HX+H<sub>2</sub>O↔H<sub>3</sub>O<sup>+</sup>+X<sup>−</sup> but simplified as HX↔H<sup>+</sup>+X<sup>−</sup>) goes to completion, i.e. no unreacted acid remains in solution. Dissolving the strong acid [[hydrochloric acid|HCl]] in water can therefore be expressed:
:HCl(aq) → H<sup>+</sup> + Cl<sup>−</sup>
This means that in a 0.01 mol/L solution of HCl it is approximated that there is a concentration of 0.01 mol/L dissolved hydrogen ions. From above, the pH is: pH = −log<sub>10</sub> [H<sup>+</sup>]:
:pH = −log (0.01)
which equals 2.
For weak acids, the dissociation reaction does not go to completion. An [[chemical equilibrium|equilibrium]] is reached between the hydrogen ions and the [[conjugate acid|conjugate base]]. The following shows the equilibrium reaction between [[formic acid|methanoic acid]] and its ions:
:HCOOH(aq) ⇌ H<sup>+</sup> + HCOO<sup>−</sup>
It is necessary to know the value of the [[equilibrium constant]] of the reaction for each acid in order to calculate its pH. In the context of pH, this is termed the ''[[acidity constant]]'' of the acid but is worked out in the same way (see [[chemical equilibrium]]):
:''K''<sub>a</sub> = [hydrogen ions][acid ions] / [acid]
For HCOOH, ''K''<sub>a</sub> = 1.6 × 10<sup>−4</sup>
When calculating the pH of a weak acid, it is usually assumed that the water does not provide any hydrogen ions. This simplifies the calculation, and the concentration provided by water, 1×10<sup>−7</sup> mol/L, is usually insignificant.
With a 0.1 mol/L solution of methanoic acid (HCOOH), the acidity constant is equal to:
:''K''<sub>a</sub> = [H<sup>+</sup>][HCOO<sup>−</sup>] / [HCOOH]
Given that an unknown amount of the acid has dissociated, [HCOOH] will be reduced by this amount, while [H<sup>+</sup>] and [HCOO<sup>−</sup>] will each be increased by this amount. Therefore, [HCOOH] may be replaced by 0.1 − ''x'', and [H<sup>+</sup>] and [HCOO<sup>−</sup>] may each be replaced by ''x'', giving us the following equation:
:<math>1.6\times 10^{-4} = \frac{x^2}{0.1-x}.</math>
Solving this for ''x'' yields 3.9×10<sup>−3</sup>, which is the concentration of hydrogen ions after dissociation. Therefore the pH is −log(3.9×10<sup>−3</sup>), or about 2.4.
== Measurement ==
{| align=right style="border: 1px solid gray; margin: 0 0 0.5em 1em;"
<caption>'''Representative pH values'''{{Fact|date=March 2008}}</caption>
! Substance || pH
|-
| [[Hydrochloric acid]], 10M ||bgcolor=#BB0000|<center><font color=#FFFFFF>-1.0
|-
| [[Lead-acid battery]] ||bgcolor=#CC0000|<center><font color=#FFFFFF>0.5
|-
| [[Gastric acid]] ||bgcolor=#EE0000|<center><font color=#FFFFFF>1.5 – 2.0
|-
| [[Lemon| Lemon juice]] ||bgcolor=#FF3300|<center><font color=#FFFFFF>2.4
|-
| [[Cola]] ||bgcolor=#FF6600|<center>2.5
|-
| [[Vinegar]] ||bgcolor=#FF9900|<center>2.9
|-
| [[Orange (fruit)|Orange]] or [[apple]] juice ||bgcolor=#FFCC00|<center>3.5
|-
| [[Tomato Juice]] ||bgcolor=yellow|<center>4.0
|-
| [[Beer]] ||bgcolor=yellow|<center>4.5
|-
| [[Acid Rain]] ||bgcolor=yellow|<center><5.0
|-
| [[Coffee]] ||bgcolor=yellow|<center>5.0
|-
| [[Tea]] or healthy [[skin]] ||bgcolor=#669900|<center>5.0
|-
| [[Urine]] ||bgcolor=#339955|<center><font color=#FFFFFF>6.0
|-
| [[Milk]] ||bgcolor=#339933|<center><font color=#FFFFFF>6.5
|-
| [[Water|Pure Water]] ||bgcolor=green|<center><font color=#FFFFFF>7.0
|-
| Healthy [[human]] [[saliva]] ||bgcolor=green|<center><font color=#FFFFFF>6.5 – 7.4
|-
| [[Blood]] ||bgcolor=#009966|<center><font color=#FFFFFF>7.34 – 7.45
|-
| [[Seawater]] ||bgcolor=#006699|<center><font color=#FFFFFF>7.7 – 8.3
|-
| [[Soap| Hand soap]] ||bgcolor=blue|<center><font color=#FFFFFF>9.0 – 10.0
|-
| [[Ammonia| Household ammonia]] ||bgcolor=blue|<center><font color=#FFFFFF>11.5
|-
| [[Bleach]] ||bgcolor=#0000CC|<center><font color=#FFFFFF>12.5
|-
| [[Sodium hydroxide|Household lye]] ||bgcolor=#000099|<center><font color=#FFFFFF>13.5
|}
pH can be measured:<br />
* by addition of a [[pH indicator]] into the solution under study. The indicator color varies depending on the pH of the solution. Using indicators, qualitative determinations can be made with universal indicators that have broad color variability over a wide pH range and quantitative determinations can be made using indicators that have strong color variability over a small pH range. Precise measurements can be made over a wide pH range using indicators that have multiple equilibriums in conjunction with [[spectrophotometry|spectrophotometric]] methods to determine the relative abundance of each pH-dependent component that make up the color of solution ,{{Fact|date=August 2007}} or
* by using a [[pH meter]] together with pH-selective electrodes ([[pH glass electrode]], [[hydrogen electrode]], [[quinhydrone electrode]], [[ion sensitive field effect transistor]] and others).
* by using pH paper, indicator paper that turns color corresponding to a pH on a color key. pH paper is usually strips of paper that has been soaked in an indicator solution, and is used for approximations.
As the pH scale is logarithmic, it does not start at zero. Thus the most acidic of liquids encountered can have a pH as low as −5.{{Fact|date=April 2008}} The most alkaline typically has pH of 14. Measurement of extremely low pH values has various complications. Calibration of the electrode in such cases can be done with standard solutions of concentrated sulfuric acid whose pH values can be calculated with the Pitzer model.<ref name="Nordstrom"/>
As an example of home application, the measurement of pH value can be used to quantify the amount of acid in a swimming pool.
== pOH ==
There is also pOH, in a sense the opposite of pH, which measures the concentration of OH<sup>−</sup> ions, or the [[Base (chemistry)|alkalinity]]. Since water self ionizes, and notating [OH<sup>−</sup>] as the concentration of [[hydroxide]] ions, we have
:<math> K_w = a_{{\rm{H}}^ * } a_{{\rm{OH}}^ - }= 10^{ - 14}</math> (*)
where K<sub>w</sub> is the ionization constant of water.
Now, since
:<math>\log _{10} K_w = \log _{10} a_{{\rm{H}}^ + } + \log _{10} a_{{\rm{OH}}^ - }</math>
by [[list of logarithmic identities|logarithmic identities]], we then have the relationship:
:<math>- 14 = {\rm{log}}_{{\rm{10}}} \,a_{{\rm{H}}^{\rm{ + }} } + \log _{10} \,a_{{\rm{OH}}^ - } </math>
and thus
:<math>{\rm{pOH}} = - \log _{10} \,a_{{\rm{OH}}^ - } = 14 + \log _{10} \,a_{{\rm{H}}^ + } = 14 - {\rm{pH}} </math>
Also, this formula is valid exactly for temperature = 298.15 K (25 °C) only, but is acceptable for most lab calculations.
== Indicators ==
[[Image:Hydrangea macrophylla - Hortensia hydrangea.jpg|right|thumbnail|250px|The ''[[Hydrangea macrophylla]]'' blossoms in [[pink]] or [[blue]], depending on soil pH. In acidic soils, the flowers are blue; in alkaline soils, the flowers are pink.]]
An [[pH indicator|indicator]] is used to measure the pH of a substance. Common indicators are [[litmus paper]], [[phenolphthalein]], [[methyl orange]], [[phenol red]], [[bromothymol blue]], [[bromocresol green]] and [[bromocresol purple]]. To demonstrate the principle with common household materials, [[red cabbage]], which contains the dye [[anthocyanin]], is used.<ref>[http://chemistry.about.com/library/weekly/aa012803a.htm chemistry.about.com]</ref>
In addition to [[red cabbage]], some flower petals (such as [[hibiscus]] and [[marigold]]) impart a bluish stain when crushed onto white paper, and may be used as a 'homemade litmus'. Addition of acidic substances will turn the paper red, after which alkaline substances will return it to blue.
==Seawater==
In [[chemical oceanography]] pH measurement is complicated by the [[chemical property|chemical properties]] of [[seawater]], and several distinct pH scales exist.<ref name=zeebe>Zeebe, R. E. and Wolf-Gladrow, D. (2001) ''CO<sub>2</sub> in seawater: equilibrium, kinetics, isotopes'', Elsevier Science B.V., Amsterdam, Netherlands (ISBN 0 444 50946 1).</ref>
As part of its [[operational definition]] of the pH scale, the [[IUPAC]] define a series of [[buffer solution]]s across a range of pH values (often denoted with [[National Bureau of Standards|NBS]] or [[National Institute of Standards and Technology|NIST]] designation). These solutions have a relatively low [[ionic strength]] (~0.1) compared to that of seawater (~0.7), and consequently are not recommended for use in characterising the pH of seawater since the ionic strength differences cause changes in [[standard electrode potential|electrode potential]]. To resolve this problem, an alternative series of buffers based on [[artificial seawater]] was developed.<ref>Hansson, I. (1973) A new set of pH-scales and standard buffers for seawater. ''Deep Sea Research'', '''20''': 479-491.</ref> This new series resolves the problem of ionic strength differences between samples and the buffers, and the new pH scale is referred to as the '''total scale''', often denoted as '''pH<sub>T</sub>'''.
The total scale was defined using a medium containing [[sulfate]] ions. These ions experience [[protonation]], H<sup>+</sup> + SO<sub>4</sub><sup>2−</sup> {{unicode|⇌}} HSO<sub>4</sub><sup>−</sup>, such that the total scale includes the effect of both [[proton]]s (free hydrogen ions) and hydrogen sulfate ions:
:[H<sup>+</sup>]<sub>T</sub> = [H<sup>+</sup>]<sub>F</sub> + [HSO<sub>4</sub><sup>−</sup>] <br>
An alternative scale, the '''free scale''', often denoted '''pH<sub>F</sub>''', omits this consideration and focuses solely on [H<sup>+</sup>]<sub>F</sub>, in principle making it a simpler representation of hydrogen ion concentration. Analytically, only [H<sup>+</sup>]<sub>T</sub> can be determined,<ref>Dickson, A. G. (1984) pH scales and proton-transfer reactions in saline media such as sea water. ''Geochim. Cosmochim. Acta'', '''48''': 2299–2308.</ref> therefore, [H<sup>+</sup>]<sub>F</sub> must be estimated using the [SO<sub>4</sub><sup>2−</sup>] and the stability constant of HSO<sub>4</sub><sup>−</sup>, K<sub>S</sub><sup>*</sup>:
: [H<sup>+</sup>]<sub>F</sub> = [H<sup>+</sup>]<sub>T</sub> − [HSO<sub>4</sub><sup>−</sup>] = [H<sup>+</sup>]<sub>T</sub> ( 1 + [SO<sub>4</sub><sup>2−</sup>] / K<sub>S</sub><sup>*</sup> )<sup>−1</sup> <br>
However, it is difficult to estimate K<sub>S</sub><sup>*</sup> in seawater, limiting the utility of the otherwise more straightforward free scale.
Another scale, known as the '''seawater scale''', often denoted '''pH<sub>SWS</sub>''', takes account of a further protonation relationship between hydrogen ions and [[fluoride]] ions, H<sup>+</sup> + F<sup>−</sup> {{unicode|⇌}} HF. Resulting in the following expression for [H<sup>+</sup>]<sub>SWS</sub>:
:[H<sup>+</sup>]<sub>SWS</sub> = [H<sup>+</sup>]<sub>F</sub> + [HSO<sub>4</sub><sup>−</sup>] + [HF] <br>
However, the advantage of considering this additional complexity is dependent upon the abundance of fluoride in the medium. In seawater, for instance, sulfate ions occur at much greater concentrations (> 400 times) than those of fluoride. Consequently, for most practical purposes, the difference between the total and seawater scales is very small.
The following three equations summarise the three scales of pH:
: pH<sub>F</sub> = − log [H<sup>+</sup>]<sub>F</sub> <br>
: pH<sub>T</sub> = − log ( [H<sup>+</sup>]<sub>F</sub> + [HSO<sub>4</sub><sup>−</sup>] ) = − log [H<sup>+</sup>]<sub>T</sub> <br>
: pH<sub>SWS</sub> = − log ( [H<sup>+</sup>]<sub>F</sub> + [HSO<sub>4</sub><sup>−</sup>] + [HF] ) = − log [H<sup>+</sup>]<sub>SWS</sub> <br>
In practical terms, the three seawater pH scales differ in their values by up to 0.12 pH units, differences that are much larger than the accuracy of pH measurements typically required, particularly in relation to the ocean's [[Total inorganic carbon|carbonate system]].<ref name=zeebe /> Since it omits consideration of sulfate and fluoride ions, the free scale is significantly different from both the total and seawater scales. Because of the relative unimportance of the fluoride ion, the total and seawater scales differ only very slightly.
==Body fluids==
{{Expand-section|date=March 2008}}
{| class="wikitable" align="right"
|+pH in body fluids <ref name=boron> Unless else specified in table: {{cite book |author=Walter F., PhD. Boron |title=Medical Physiology: A Cellular And Molecular Approaoch |publisher=Elsevier/Saunders |location= |year= |pages= |isbn=1-4160-2328-3 |oclc= |doi=}} Page 634 </ref>
|-
! Fluid
! pH
|-
| [[gastric acid]] || 0.7
|-
| [[lysosome]] || 5.5
|-
| granule of [[chromaffin cell]] || 5.5
|-
| Neutral H<sub>2</sub>O at 37°C || 6.81
|-
| [[cytosol]] || 7.2
|-
| [[cerebrospinal fluid|CSF]] || 7.3
|-
| [[arterial]] [[blood plasma]] || 7.4
|-
| [[mitochondrial matrix]] || 7.5
|-
| exocrine secretions of [[pancreas]] || 8.1
|}
The pH of different [[body fluids]], including urine, saliva, and blood, varies with function and other factors. They are mostly tightly regulated systems to keep the [[acid-base homeostasis]]. A notable acidic substance in the body is plaque. [[Dental_plaque|Plaque]]'s pH is low and will dissolve teeth if not removed. The pH of blood is known to be slightly basic, at a value of 7.4. pH is vital in maintaining the functioning of cells. For example, enzymes are heavily affected by changes in pH, and have an optimum pH at which they operate. Outside a small range they can denature and cease to catalyse vital reactions.
==See also==
*[[Acidosis]]
*[[Alkalosis]]
*[[S. P. L. Sørensen]]
== References ==
<!--See http://en.wikipedia.org/wiki/Wikipedia:Footnotes for an explanation of how to generate footnotes using the <ref(erences/)> tags-->
{{reflist}}
== External links ==
* [http://www.chembuddy.com/?left=BATE&right=dissociation_constants Some Ka values]
* [http://www.iupac.org/goldbook/P04524.pdf Definition of pH] IUPAC [[Gold Book]]
*[http://www.iupac.org/reports/provisional/abstract01/rondinini_prs.pdf The Measurement of pH - Definition, Standards and Procedures] – Report of the Working Party on pH, IUPAC Provisional Recommendation (a proposal to revise the current IUPAC 1985 and ISO 31-8 definition of pH described above)
[[Category:Acid-base chemistry]]
[[Category:Units of measure]]
[[Category:Water quality indicators]]
[[af:PH]]
[[ar:أس هيدروجيني]]
[[zh-min-nan:PH]]
[[bs:PH]]
[[bg:Водороден показател]]
[[ca:PH]]
[[cs:Kyselost]]
[[cy:PH]]
[[da:PH og pOH]]
[[de:PH-Wert]]
[[et:Vesinikeksponent]]
[[el:PH]]
[[es:PH]]
[[eo:PH (kemia parametro)]]
[[eu:PH]]
[[fa:پهاش]]
[[fr:Potentiel hydrogène]]
[[gl:PH]]
[[ko:수소 이온 농도]]
[[hr:PH]]
[[io:PH]]
[[id:PH]]
[[ia:Potential de hydrogeno]]
[[is:Sýrustig]]
[[it:PH]]
[[he:PH]]
[[pam:PH]]
[[ku:Nirxa pH]]
[[la:Pondus Hydrogenii]]
[[lv:PH]]
[[lt:PH]]
[[hu:PH]]
[[mk:Водороден показател]]
[[ml:പി.എച്ച്. മൂല്യം]]
[[ms:PH]]
[[nl:PH]]
[[ja:水素イオン指数]]
[[no:PH]]
[[nn:PH]]
[[nov:PH]]
[[pl:Skala pH]]
[[pt:PH]]
[[ro:PH]]
[[ru:Водородный показатель]]
[[simple:PH]]
[[sk:Kyslosť]]
[[sl:PH]]
[[so:PH]]
[[sr:PH вредност]]
[[sh:PH]]
[[su:PH]]
[[fi:Happamuus]]
[[sv:PH]]
[[th:ค่า pH]]
[[vi:PH]]
[[tr:PH]]
[[uk:PH]]
[[zh:PH值]]