Editing PH (section)

== Applications ==
Pure water has a pH of 7 at 25&nbsp;°C, meaning it is neutral. When an [[acid]] is dissolved in water, the pH will be less than 7, while a [[Base (chemistry)|base]], or [[alkali]], will have a pH greater than 7. A strong acid, such as [[hydrochloric acid]], at concentration 1&nbsp;mol/L has a pH of 0, while a strong alkali like [[sodium hydroxide]], at the same concentration, has a pH of 14. Since pH is a logarithmic scale, a difference of one in pH is equivalent to a tenfold difference in hydrogen ion concentration.

Neutrality is not exactly 7 at 25&nbsp;°C, but 7 serves as a good approximation in most cases. Neutrality occurs when the concentration of hydrogen cations ([{{chem2|H+}}]) equals the concentration of hydroxide ions ([{{chem2|OH−}}]), or when their activities are equal. Since [[self-ionization of water]] holds the product of these concentration [{{chem2|H+}}] × [{{chem2|OH−}}] = ''K''<sub>w</sub>, it can be seen that at neutrality [{{chem2|H+}}] = [{{chem2|OH−}}] = {{radic|''K''<sub>w</sub>}}, or pH = p''K''<sub>w</sub>/2. p''K''<sub>w</sub> is approximately 14 but depends on ionic strength and temperature, and so the pH of neutrality does also. Pure water and a solution of [[Sodium chloride|NaCl]] in pure water are both neutral, since [[Self-ionization of water|dissociation of water]] produces equal numbers of both ions. However the pH of the neutral NaCl solution will be slightly different from that of neutral pure water because the hydrogen and hydroxide ions' activity is dependent on [[ionic strength]], so ''K''<sub>w</sub> varies with ionic strength.

When pure water is exposed to air, it becomes mildly acidic. This is because water absorbs [[carbon dioxide]] from the air, which is then slowly converted into [[bicarbonate]] and hydrogen cations (essentially creating [[carbonic acid]]).
:<chem>CO2 + H2O <=> HCO3^- + H^+</chem>

=== pH in soil ===
{{See also|Soil pH}}

[[File:Soil_pH_effect_on_nutrient_availability.svg|thumb|Nutritional elements availability within soil varies with pH. Light blue color represents the ideal range for most plants.]]
The United States Department of Agriculture [[Natural Resources Conservation Service]], formerly Soil Conservation Service classifies [[soil pH]] ranges as follows:<ref>{{cite web |author=Soil Survey Division Staff |title=Soil survey manual.1993. Chapter 3, selected chemical properties. |url=http://soils.usda.gov/technical/manual/contents/chapter3.html |url-status=dead |archive-url=https://web.archive.org/web/20110514151830/http://soils.usda.gov/technical/manual/contents/chapter3.html |archive-date=14 May 2011 |access-date=2011-03-12 |publisher=Soil Conservation Service. U.S. Department of Agriculture Handbook 18}}</ref>
{{Table alignment}}

{| {{table|class=defaultright}}
! scope="col" |Denomination
! scope="col" |pH range
|-
|Ultra acidic
|< 3.5
|-
|Extremely acidic
|3.5–4.4
|-
|Very strongly acidic
|4.5–5.0
|-
|Strongly acidic
|5.1–5.5
|-
|Moderately acidic
|5.6–6.0
|-
|Slightly acidic
|6.1–6.5
|-
|Neutral
|6.6–7.3
|-
|Slightly alkaline
|7.4–7.8
|-
|Moderately alkaline
|7.9–8.4
|-
|Strongly alkaline
|8.5–9.0
|-
|Very strongly alkaline
|9.0–10.5
|-
|Hyper alkaline
|> 10.5
|}
Topsoil pH is influenced by soil parent material, erosional effects, climate and vegetation. A recent map<ref>{{Cite journal |last1=Ballabio |first1=Cristiano |last2=Lugato |first2=Emanuele |last3=Fernández-Ugalde |first3=Oihane |last4=Orgiazzi |first4=Alberto |last5=Jones |first5=Arwyn |last6=Borrelli |first6=Pasquale |last7=Montanarella |first7=Luca |last8=Panagos |first8=Panos |date=2019 |title=Mapping LUCAS topsoil chemical properties at European scale using Gaussian process regression |journal=Geoderma |language=en |volume=355 |pages=113912 |bibcode=2019Geode.35513912B |doi=10.1016/j.geoderma.2019.113912 |pmc=6743211 |pmid=31798185 |doi-access=free}}</ref> of topsoil pH in Europe shows the alkaline soils in Mediterranean, Hungary, East Romania, North France. Scandinavian countries, Portugal, Poland and North Germany have more acid soils.

=== pH in plants ===
[[File:Lemon_-_whole_and_split.jpg|thumb|[[Lemon juice]] tastes sour because it contains 5% to 6% [[citric acid]] and has a pH of 2.2 (high acidity).]]
Plants contain pH-dependent [[Plant pigment|pigments]] that can be used as [[PH indicator|pH indicators]], such as those found in [[hibiscus]], [[red cabbage]] ([[anthocyanin]]), and grapes ([[red wine]]). [[Citrus]] fruits have acidic juice primarily due to the presence of [[citric acid]], while other [[Carboxylic acid|carboxylic acids]] can be found in various living systems. The [[protonation]] state of [[phosphate]] derivatives, including [[Adenosine triphosphate|ATP]], is pH-dependent. [[Hemoglobin]], an oxygen-transport enzyme, is also affected by pH in a phenomenon known as the [[Root effect]].

=== pH in the ocean ===
{{Anchor|Seawater}}{{See also|Seawater#pH value|Ocean#pH and alkalinity|Ocean acidification}}

The pH of [[seawater]] plays an important role in the ocean's [[Carbon cycle#Ocean|carbon cycle]]. There is evidence of ongoing [[ocean acidification]] (meaning a drop in pH value): Between 1950 and 2020, the average pH of the ocean surface fell from approximately 8.15 to 8.05.<ref>{{Cite journal |last1=Terhaar |first1=Jens |last2=Frölicher |first2=Thomas L. |last3=Joos |first3=Fortunat |date=2023 |title=Ocean acidification in emission-driven temperature stabilization scenarios: the role of TCRE and non-{{CO2}} greenhouse gases |journal=Environmental Research Letters |language=en |volume=18 |issue=2 |pages=024033 |bibcode=2023ERL....18b4033T |doi=10.1088/1748-9326/acaf91 |issn=1748-9326 |s2cid=255431338|doi-access=free }}</ref> [[Carbon dioxide emissions]] from human activities are the primary cause of ocean acidification, with [[Carbon dioxide in Earth's atmosphere|atmospheric carbon dioxide levels]] at 430 ppm {{CO2}} at [[Mauna Loa]] observatory in 2025.<ref name="NOAA_CO2">{{cite web | title=Trends in CO<sub>2</sub> | website=NOAA Global Monitoring Laboratory | date=2025-04-22 | url=https://gml.noaa.gov/ccgg/trends/monthly.html | access-date=2025-04-22}}</ref> In 2024, the annual atmospheric {{CO2}} increase measured by the [[National Oceanic and Atmospheric Administration|NOAA]]’s Global Monitoring Laboratory was 3.75 ppm {{CO2}}/year.<ref name="Berwyn2025">{{cite web | last1=Berwyn | first1=Bob | title=A grim signal: Atmospheric CO<sub>2</sub> soared in 2024 | website=Ars Technica | date=2025-04-25 | url=https://arstechnica.com/science/2025/04/a-grim-signal-atmospheric-co2-soared-in-2024/ | access-date=2025-04-27}}</ref> CO<sub>2</sub> from the [[atmosphere]] is absorbed by the oceans. This produces [[carbonic acid]] (H<sub>2</sub>CO<sub>3</sub>) which dissociates into a [[bicarbonate ion]] ({{Chem|HCO|3|-|}}) and a [[Hydron|hydrogen cation]] (H<sup>+</sup>). The presence of free hydrogen cations (H<sup>+</sup>) lowers the pH of the ocean. 

==== Three pH scales in oceanography ====
The measurement of pH in seawater is complicated by the [[Chemical property|chemical properties]] of seawater, and three distinct pH scales exist in [[chemical oceanography]].<ref name="zeebe2">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> In practical terms, the three seawater pH scales differ in their pH values up to 0.10, differences that are much larger than the accuracy of pH measurements typically required, in particular, in relation to the ocean's [[Total inorganic carbon|carbonate system]].<ref name="zeebe2" /> 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.

As part of its [[operational definition]] of the pH scale, the [[IUPAC]] defines 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]] (≈&nbsp;0.1) compared to that of seawater (≈&nbsp;0.7), and, as a consequence, are not recommended for use in characterizing 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>{{cite journal |author=Hansson, I. |year=1973 |title=A new set of pH-scales and standard buffers for seawater |journal=Deep-Sea Research |volume=20 |issue=5 |pages=479–491 |bibcode=1973DSRA...20..479H |doi=10.1016/0011-7471(73)90101-0}}</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]], {{chem2|H+}} + {{chem|SO|4|2-|↔ HSO|4|-}}, such that the total scale includes the effect of both [[Proton|protons]] (free hydrogen cations) and hydrogen sulfate ions:
: [{{chem2|H+}}]<sub>T</sub> = [{{chem2|H+}}]<sub>F</sub> + [{{chem|HSO|4|-}}]

An alternative scale, the ''free scale'', often denoted pH<sub>F</sub>, omits this consideration and focuses solely on [{{chem2|H+}}]<sub>F</sub>, in principle making it a simpler representation of hydrogen ion concentration. Only [{{chem2|H+}}]<sub>T</sub> can be determined,<ref>{{cite journal |author=Dickson, A. G. |year=1984 |title=pH scales and proton-transfer reactions in saline media such as sea water |journal=Geochim. Cosmochim. Acta |volume=48 |issue=11 |pages=2299–2308 |bibcode=1984GeCoA..48.2299D |doi=10.1016/0016-7037(84)90225-4}}</ref> therefore [{{chem2|H+}}]<sub>F</sub> must be estimated using the [{{chem|SO|4|2-}}] and the stability constant of {{chem|HSO|4|-}}, {{nowrap|''K''{{su|b=S|p=*}}}}:
: [{{chem2|H+}}]<sub>F</sub> = [{{chem2|H+}}]<sub>T</sub> − [{{chem|HSO|4|-}}] = [{{chem2|H+}}]<sub>T</sub> ( 1 + [{{chem|SO|4|2-}}] / ''K''{{su|b=S|p=*}} )<sup>−1</sup>

However, it is difficult to estimate ''K''{{su|b=S|p=*}} 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 cations and [[fluoride]] ions, {{chem2|H+}} + {{chem2|F-}} ⇌ HF. Resulting in the following expression for [{{chem2|H+}}]<sub>SWS</sub>:
: [{{chem2|H+}}]<sub>SWS</sub> = [{{chem2|H+}}]<sub>F</sub> + [{{chem|HSO|4|-}}] + [HF]

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 (>&nbsp;400 times) than those of fluoride. As a consequence, for most practical purposes, the difference between the total and seawater scales is very small.

The following three equations summarize the three scales of pH:
: pH<sub>F</sub> = −log<sub>10</sub>[{{chem2|H+}}]<sub>F</sub>
: pH<sub>T</sub> = −log<sub>10</sub>([{{chem2|H+}}]<sub>F</sub> + [{{chem|HSO|4|-}}]) = −log<sub>10</sub>[{{chem2|H+}}]<sub>T</sub>
: pH<sub>SWS</sub> = −log<sub>10</sub>({{chem2|H+}}]<sub>F</sub> + [{{chem|HSO|4|-}}] + [HF]) = −log<sub>10</sub>[v]<sub>SWS</sub>

=== pH in food ===
The pH level of food influences its flavor, texture, and [[shelf life]].<ref>{{cite journal |last1=Bello |first1=Andrez |last2=Palacios |first2=Baretto |date=2013 |title=Effect of pH on Color and Texture of Food Products |url=https://link.springer.com/article/10.1007/s12393-013-9067-2 |journal=Food Eng Rev |volume=5 |issue=3 |pages=158–170 |doi=10.1007/s12393-013-9067-2 |hdl=10251/57961 |access-date=June 20, 2024|url-access=subscription }}</ref> Acidic foods, such as [[citrus fruits]], tomatoes, and [[vinegar]], typically have a pH below 4.6<ref name="okla">{{cite web |url=https://extension.okstate.edu/fact-sheets/the-importance-of-food-ph-in-commercial-canning-operations.html#ph-and-microbial-growth |title=The Importance of Food pH in Commercial Canning Operations |last=McGlynn |first=William |website=Oklahoma State University |date=July 1, 2016 |access-date=June 20, 2024}}</ref> with sharp and tangy taste, while basic foods taste bitter or soapy.<ref>{{cite web |url=https://www.samaterials.com/ph-acids-bases-and-common-materials.html |title=PH Scale: Acids, Bases, and Common Materials |last=Trento |first=Chin |website= Stanford Advanced Materials|date=Dec 27, 2023 |access-date=June 20, 2024}}</ref>  Maintaining the appropriate pH in foods is essential for preventing the growth of harmful [[microorganisms]].<ref name="okla" /> The alkalinity of vegetables such as [[spinach]] and [[kale]] can also influence their texture and color during cooking.<ref>{{cite journal |last1=Akdas |first1=Zelal |last2=Bakkalbasi |first2=Emre |date=2017 |title=Influence of different cooking methods on color, bioactive compounds, and antioxidant activity of kale |journal=International Journal of Food Properties |volume=20 |issue=4 |pages=877–887 |doi=10.1080/10942912.2016.1188308}}</ref> The pH also influences the [[Maillard reaction]], which is responsible for the browning of food during cooking, impacting both flavor and appearance.<ref>{{cite journal |last1=Rauh |first1=Valentin |last2=Xiao |first2=Yinghua |date=2022 |title=The shelf life of heat-treated dairy products |url=https://www.sciencedirect.com/science/article/abs/pii/S0958694621002636 |journal=International Dairy Journal |volume=125 |doi=10.1016/j.idairyj.2021.105235 |access-date=June 20, 2024|url-access=subscription }}</ref>

=== pH of various body fluids ===
: {| class="wikitable"
|+pH of various body fluids<ref name=Boron2012>{{cite book |last1=Boron |first1=Walter F. |url= https://books.google.com/books?id=54mxMgO5H_YC&dq=pH%20in%20living%20systems&pg=PA652 |title=Medical Physiology: A Cellular And Molecular Approach |last2=Boulpaep |first2=Emile L. |date=13 January 2012 |publisher=[[Elsevier Health Sciences]], Saunders |isbn=9781455711819 |pages=652–671 |oclc=1017876653 |access-date=8 May 2022 |archive-date=8 May 2022 |archive-url=https://web.archive.org/web/20220508051939/https://www.google.co.in/books/edition/_/54mxMgO5H_YC?hl=en&gbpv=1&dq=pH+in+living+systems&pg=PA652 |url-status=live |edition=2nd }}</ref>
|-
! Compartment
! pH
|-
| [[Gastric acid]] || 1.5–3.5<ref>{{cite news|url=https://www.ucsfhealth.org/medical-tests/stomach-acid-test|title=Stomach acid test|publisher=University of California San Francisco|access-date=21 February 2024}}</ref><ref>{{cite book |last1=Marieb |first1=Elaine N. |url=https://books.google.com/books?id=BxDfnQEACAAJ |title=Human anatomy & physiology |last2=Mitchell |first2=Susan J. |date=30 June 2011 |publisher=[[Benjamin Cummings]] |isbn=9780321735287 |location=San Francisco |access-date=8 May 2022 |archive-date=8 May 2022 |archive-url=https://web.archive.org/web/20220508051937/https://www.google.co.in/books/edition/_/BxDfnQEACAAJ?hl=en |url-status=live }}</ref>
|-
| [[Lysosome]]s || 4.5<ref name=Boron2012/>
|-
| [[Human skin]] || 4.7<ref>{{Cite journal|last1=Lambers|first1=H.|last2=Piessens|first2=S.|last3=Bloem|first3=A.|last4=Pronk|first4=H.|last5=Finkel|first5=P.|date=2006-10-01|url=https://onlinelibrary.wiley.com/doi/10.1111/j.1467-2494.2006.00344.x|url-status=live|title=Natural skin surface pH is on average below 5, which is beneficial for its resident flora|journal=International Journal of Cosmetic Science| volume= 28 |issue= 5|pages=359–370|doi=10.1111/j.1467-2494.2006.00344.x|issn=1468-2494|pmid=18489300|s2cid=25191984|access-date=8 May 2022|archive-date=21 March 2022|archive-url= https://web.archive.org/web/20220321033318/https://onlinelibrary.wiley.com/doi/10.1111/j.1467-2494.2006.00344.x|url-access=subscription}}</ref>
|-
| Granules of [[chromaffin cell]]s || 5.5
|-
| [[Urine]] || 6.0
|-
| [[Breast milk]] || 7.0–7.45<ref>{{cite journal | url=https://pubmed.ncbi.nlm.nih.gov/3748680/#:~:text=Thereafter%2C%20the%20pH%20of%20milk,with%20the%20concentration%20of%20lactose | pmid=3748680 | date=1986 | last1=Morriss Jr | first1=F. H. | last2=Brewer | first2=E. D. | last3=Spedale | first3=S. B. | last4=Riddle | first4=L. | last5=Temple | first5=D. M. | last6=Caprioli | first6=R. M. | last7=West | first7=M. S. | title=Relationship of human milk pH during course of lactation to concentrations of citrate and fatty acids | journal=Pediatrics | volume=78 | issue=3 | pages=458–464 | doi=10.1542/peds.78.3.458 }}</ref>
|-
| [[Cytosol]] || 7.2
|-
| [[Blood]] (natural pH) || 7.34–7.45<ref name=Boron2012/>
|-
| [[Cerebrospinal fluid]] (CSF) || 7.5
|-
| [[Mitochondrial matrix]] || 7.5
|-
| [[Pancreas]] secretions || 8.1
|}

In living organisms, the pH of various [[body fluid]]s, cellular compartments, and organs is tightly regulated to maintain a state of acid–base balance known as [[acid–base homeostasis]]. [[Acidosis]], defined by blood pH below 7.35, is the most common disorder of acid–base homeostasis and occurs when there is an excess of acid in the body. In contrast, [[alkalosis]]  is characterized by excessively high blood pH.

Blood pH is usually slightly alkaline, with a pH of 7.365, referred to as physiological pH in biology and medicine. [[Dental plaque|Plaque]] formation in teeth can create a local acidic environment that results in [[tooth decay]] through demineralization. [[Enzyme]]s and other [[Protein]]s have an optimal pH range for function and can become inactivated or [[Denaturation (biochemistry)|denatured]] outside this range.