Editing PH (section)

== Measurement ==

=== pH Indicators ===
{| {{table|class=floatright}}
|+ Average pH of common solutions
|-
!Substance
!pH range
!Type
|-
|[[Sulfuric acid|Battery acid]]
| style="background-color: #CC0000; text-align: center; color: #ffffff" |< 1
| rowspan="6" style="text-align: center" |[[Acid]]
|-
|[[Gastric acid]]
| style="background-color: #EE0000; text-align: center; color: #ffffff" |1.0–1.5
|-
|[[Vinegar]]
| style="background-color: #FF4000; text-align: center" |2–3
|-
|[[Orange juice]]
| style="background-color: #FF6600; text-align: center" |3.3–4.2
|-
|[[Coffee|Black coffee]]
| style="background-color: #ffff00 ; text-align: center" |5–5.03
|-
|[[Milk]]
| style="background-color: #99cc33; text-align: center" |6.5–6.8
|-
|[[Pure water]] at 25&nbsp;°C
| style="background-color: #339933; text-align: center; color: #ffffff" |7
| style="text-align: center" |Neutral
|-
|[[Sea water]]
| style="background-color: #19cdff; text-align: center; color: #000000" |7.5–8.4
| rowspan="4" style="text-align: center" |[[Base (chemistry)|Base]]
|-
|[[Ammonia]]
| style="background-color: #3333ff; text-align: center; color: #FFFFFF" |11.0–11.5
|-
|[[Bleach]]
| style="background-color: #330099; text-align: center; color: #FFFFFF" |12.5
|-
|[[Lye]]
| style="background-color: #330066; text-align: center; color: #FFFFFF" |14
|}
{{Main|pH indicator}}
pH can be measured using indicators, which change color depending on the pH of the solution they are in. By comparing the color of a test solution to a standard color chart, the pH can be estimated to the nearest whole number. For more precise measurements, the color can be measured using a [[Colorimeter (chemistry)|colorimeter]] or [[spectrophotometer]]. A [[Universal indicator]] is a mixture of several indicators that can provide a continuous color change over a range of pH values, typically from about pH 2 to pH 10. Universal indicator paper is made from absorbent paper that has been impregnated with a universal indicator.  An alternative method of measuring pH is using an electronic [[pH meter]], which directly measures the voltage difference between a pH-sensitive electrode and a reference electrode.

=== Non-aqueous solutions ===
pH values can be measured in non-aqueous solutions, but they are based on a different scale from aqueous pH values because the [[Standard state|standard states]] used for calculating hydrogen ion concentrations ([[Activity (chemistry)|activities]]) are different. The hydrogen ion activity, ''a''<sub>H<sup>+</sup></sub>, is defined<ref name="GoldBook2">{{GoldBookRef|title=activity (relative activity), ''a''|file=A00115}}</ref><ref name="GreenBook2">{{GreenBookRef2nd|pages=49–50}}</ref> as:
: <math chem="">a_\ce{H+} = \exp\left (\frac{\mu_\ce{H+} - \mu^{\ominus}_\ce{H+}}{RT}\right )</math>
where ''μ''<sub>H<sup>+</sup></sub> is the [[chemical potential]] of the hydrogen cation, <math chem="">\mu^{\ominus}_\ce{H+}</math> is its chemical potential in the chosen standard state, ''R'' is the [[molar gas constant]] and ''T'' is the [[thermodynamic temperature]]. Therefore, pH values on the different scales cannot be compared directly because of differences in the solvated proton ions, such as lyonium ions, which require an insolvent scale that involves the transfer activity coefficient of [[Lyonium ion|hydronium/lyonium ion]].

pH is an example of an [[acidity function]], but others can be defined. For example, the [[Hammett acidity function]], ''H''<sub>0</sub>, has been developed in connection with [[Superacid]]s.

=== Unified absolute pH scale ===
In 2010, a new approach to measuring pH was proposed, called the ''unified absolute pH scale''. This approach allows for a common reference standard to be used across different solutions, regardless of their pH range. The unified absolute pH scale is based on the absolute chemical potential of the hydrogen cation, as defined by the [[Lewis acids and bases|Lewis acid–base]] theory. This scale applies to liquids, gases, and even solids.<ref name="Krossing2">{{Cite journal |last1=Himmel |first1=Daniel |last2=Goll |first2=Sascha K. |last3=Leito |first3=Ivo |last4=Krossing |first4=Ingo |date=2010-08-16 |title=A Unified pH Scale for All Phases |journal=Angewandte Chemie International Edition |volume=49 |issue=38 |pages=6885–6888 |doi=10.1002/anie.201000252 |issn=1433-7851 |pmid=20715223}}</ref> The advantages of the unified absolute pH scale include consistency, accuracy, and applicability to a wide range of sample types. It is precise and versatile because it serves as a common reference standard for pH measurements. However, implementation efforts, compatibility with existing data, complexity, and potential costs are some challenges.

=== Extremes of pH measurements ===
{{Redirect|Negative pH|the band| Negative pH (band)}}
The measurement of pH can become difficult at extremely acidic or alkaline conditions, such as below pH 2.5 (ca. 0.003&nbsp;[[Mole (unit)|mol]]/dm<sup>3</sup> acid) or above pH 10.5 (above ca. 0.0003 mol/dm<sup>3</sup> alkaline). This is due to the breakdown of the [[Nernst equation]] in such conditions when using a glass electrode. Several factors contribute to this problem. First, [[liquid junction potential]]s may not be independent of pH.<ref name="Feldman2">{{cite journal |author=Feldman, Isaac |year=1956 |title=Use and Abuse of pH measurements |journal=Analytical Chemistry |volume=28 |issue=12 |pages=1859–1866 |doi=10.1021/ac60120a014}}</ref> Second, the high [[ionic strength]] of concentrated solutions can affect the electrode potentials. At high pH the glass electrode may be affected by "alkaline error", because the electrode becomes sensitive to the concentration of cations such as {{chem2|Na+}} and {{chem2|K+}} in the solution.<ref>{{VogelQuantitative}}, Section 13.19 The glass electrode</ref> To overcome these problems, specially constructed electrodes are available.

Runoff from mines or mine tailings can produce some extremely low pH values, down to −3.6.<ref>{{cite journal |last1=Nordstrom |first1=D. Kirk |last2=Alpers |first2=Charles N. |date=March 1999 |title=Negative pH, efflorescent mineralogy, and consequences for environmental restoration at the Iron Mountain Superfund site, California |url=http://digitalcommons.unl.edu/cgi/viewcontent.cgi?article=1495&context=usgsstaffpub |url-status=live |journal=Proceedings of the National Academy of Sciences of the United States of America |volume=96 |issue=7 |pages=3455–62 |bibcode=1999PNAS...96.3455N |doi=10.1073/pnas.96.7.3455 |pmc=34288 |pmid=10097057 |archive-url=https://web.archive.org/web/20170923012227/http://digitalcommons.unl.edu/cgi/viewcontent.cgi?article=1495&context=usgsstaffpub |archive-date=23 September 2017 |access-date=4 November 2018 |doi-access=free}}</ref>