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pH

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</style><div role="note" class="hatnote navigation-not-searchable">For other uses, see <a href="PH_(disambiguation)" class="mw-disambig" title="PH (disambiguation)">PH (disambiguation)</a>.</div>

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</style><table class="sidebar nomobile nowraplinks"><tbody><tr><th class="sidebar-title" style="background:#d3d3d3;">Acids and bases</th></tr><tr><td class="sidebar-image" style="background:light-dark(transparent,#999);"></td></tr><tr><td class="sidebar-content hlist" style="padding:0.2em 0 0.75em;">

<ul><li><a href="Acceptor_number" class="mw-redirect" title="Acceptor number">Acceptor number</a></li>

<li><a href="Acid%E2%80%93base_reaction" title="Acid–base reaction">Acid–base reaction</a></li>

<li><a href="Acid%E2%80%93base_homeostasis" title="Acid–base homeostasis">Acid–base homeostasis</a></li>

<li><a href="Acid_strength" title="Acid strength">Acid strength</a></li>

<li><a href="Acidity_function" title="Acidity function">Acidity function</a></li>

<li><a href="Amphoterism" title="Amphoterism">Amphoterism</a></li>

<li><a href="Buffer_solution" title="Buffer solution">Buffer solutions</a></li>

<li><a href="Dissociation_constant" title="Dissociation constant">Dissociation constant</a></li>

<li><a href="Donor_number" title="Donor number">Donor number</a></li>

<li><a href="Equilibrium_chemistry" title="Equilibrium chemistry">Equilibrium chemistry</a></li>

<li><a href="Acid%E2%80%93base_extraction" title="Acid–base extraction">Extraction</a></li>

<li><a href="Hammett_acidity_function" title="Hammett acidity function">Hammett acidity function</a></li>

<li><a href="Proton_affinity" title="Proton affinity">Proton affinity</a></li>

<li><a href="Self-ionization_of_water" title="Self-ionization of water">Self-ionization of water</a></li>

<li><a href="Acid%E2%80%93base_titration" title="Acid–base titration">Titration</a></li>

<li><a href="Lewis_acid_catalysis" title="Lewis acid catalysis">Lewis acid catalysis</a></li>

<li><a href="Frustrated_Lewis_pair" title="Frustrated Lewis pair">Frustrated Lewis pair</a></li>

<li><a href="Chiral_Lewis_acid" title="Chiral Lewis acid">Chiral Lewis acid</a></li>

<li><a href="ECW_model" title="ECW model">ECW model</a></li></ul></td>

</tr><tr><th class="sidebar-heading" style="background:#e5e5e5;">

<a href="Acid" title="Acid">Acid</a> types</th></tr><tr><td class="sidebar-content hlist" style="padding:0.2em 0 0.75em;">

<ul><li><a href="Br%C3%B8nsted%E2%80%93Lowry_acid%E2%80%93base_theory" title="Brønsted–Lowry acid–base theory">Brønsted–Lowry</a></li>

<li><a href="Lewis_acids_and_bases" title="Lewis acids and bases">Lewis</a></li>

<li><a href="Mineral_acid" title="Mineral acid">Mineral</a></li>

<li><a href="Organic_acid" title="Organic acid">Organic</a></li>

<li><a href="Acidic_oxide" title="Acidic oxide">Oxide</a></li>

<li><a href="Strong_acid" class="mw-redirect" title="Strong acid">Strong</a></li>

<li><a href="Superacid" title="Superacid">Superacids</a></li>

<li><a href="Solid_acid" title="Solid acid">Solid</a></li></ul></td>

</tr><tr><th class="sidebar-heading" style="background:#e5e5e5;">

<a href="Base_(chemistry)" title="Base (chemistry)">Base</a> types</th></tr><tr><td class="sidebar-content hlist" style="padding:0.2em 0 0.75em;">

<ul><li><a href="Br%C3%B8nsted%E2%80%93Lowry_acid%E2%80%93base_theory" title="Brønsted–Lowry acid–base theory">Brønsted–Lowry</a></li>

<li><a href="Lewis_acids_and_bases" title="Lewis acids and bases">Lewis</a></li>

<li><a href="Organic_base" title="Organic base">Organic</a></li>

<li><a href="Basic_oxide" title="Basic oxide">Oxide</a></li>

<li><a href="Base_(chemistry)#Strong_bases" title="Base (chemistry)">Strong</a></li>

<li><a href="Superbase" title="Superbase">Superbases</a></li>

<li><a href="Non-nucleophilic_base" title="Non-nucleophilic base">Non-nucleophilic</a></li>

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</style></td></tr></tbody></table>In <a href="Chemistry" title="Chemistry">chemistry</a>, pH (/piːˈeɪtʃ/ pee-AYCH) is a <a href="Logarithmic_scale" title="Logarithmic scale">logarithmic scale</a> used to specify the <a href="Acid" title="Acid">acidity</a> or <a href="Base_(chemistry)" title="Base (chemistry)">basicity</a> of <a href="Aqueous_solution" title="Aqueous solution">aqueous solutions</a>. Acidic solutions (solutions with higher concentrations of hydrogen (<a href="Hydrogen_ion#Cation_(positively_charged)" title="Hydrogen ion"><style data-mw-deduplicate="TemplateStyles:r1123817410">

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</style>H+</a>) <a href="Cation" class="mw-redirect" title="Cation">cations</a>) are measured to have lower pH values than basic or <a href="Alkali" title="Alkali">alkaline</a> solutions. Historically, pH denotes "<a href="Chemical_potential" title="Chemical potential">potential</a> of <a href="Hydrogen" title="Hydrogen">hydrogen</a>" (or "power of hydrogen").<a href="#cite_note-1">[1]</a>

The pH scale is logarithmic and inversely indicates the <a href="Thermodynamic_activity" title="Thermodynamic activity">activity</a> of <a href="Hydronium" title="Hydronium">hydrogen cations</a> in the solution

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<annotation encoding="application/x-tex">{\displaystyle {\ce {pH}}=-\log _{10}(a_{{\ce {H+}}})\thickapprox -\log _{10}([{\ce {H+}}]/{\text{M}})}</annotation>

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</math><img src="./fe31ce01f515ddddc195f7a5b6a13dd32c80a7a8.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -0.838ex; width:38.675ex; height:3.009ex;" alt="{\displaystyle {\ce {pH}}=-\log _{10}(a_{{\ce {H+}}})\thickapprox -\log _{10}([{\ce {H+}}]/{\text{M}})}" loading="lazy"></dd></dl>

where [H+] is the <a href="Equilibrium_chemistry" title="Equilibrium chemistry">equilibrium</a> <a href="Molar_concentration" title="Molar concentration">molar concentration</a> of H+ (in M = <a href="Mole_(unit)" title="Mole (unit)">mol</a>/<a href="Litre" title="Litre">L</a>) in the solution. At 25 <a href="Celsius" title="Celsius">°C</a> (77 <a href="Fahrenheit" title="Fahrenheit">°F</a>), solutions of which the pH is less than 7 are acidic, and solutions of which the pH is greater than 7 are basic. Solutions with a pH of 7 at 25 °C are neutral (i.e. have the same concentration of H+ ions as OH− ions, i.e. the same as <a href="Pure_water" class="mw-redirect" title="Pure water">pure water</a>). The neutral value of the pH depends on the temperature and is lower than 7 if the temperature increases above 25 °C. The pH range is commonly given as zero to 14, but a pH value can be less than 0 for very concentrated <a href="Acid_strength" title="Acid strength">strong acids</a> or greater than 14 for very concentrated <a href="Base_(chemistry)#Strong_bases" title="Base (chemistry)">strong bases</a>.<a href="#cite_note-2">[2]</a>

The pH scale is <a href="Measurement_traceability" class="mw-redirect" title="Measurement traceability">traceable</a> to a set of standard solutions whose pH is established by international agreement.<a href="#cite_note-covington3-3">[3]</a> Primary pH standard values are determined using a <a href="Galvanic_cell" title="Galvanic cell">concentration cell with transference</a> by measuring the potential difference between a <a href="Hydrogen_electrode" class="mw-redirect" title="Hydrogen electrode">hydrogen electrode</a> and a <a href="Standard_electrode_potential_(data_page)" title="Standard electrode potential (data page)">standard electrode</a> such as the <a href="Silver_chloride_electrode" title="Silver chloride electrode">silver chloride electrode</a>. The pH of aqueous solutions can be measured with a <a href="Glass_electrode" title="Glass electrode">glass electrode</a> and a <a href="PH_meter" title="PH meter">pH meter</a> or a color-changing <a href="PH_indicator" title="PH indicator">indicator</a>. Measurements of pH are important in <a href="Chemistry" title="Chemistry">chemistry</a>, <a href="Agronomy" title="Agronomy">agronomy</a>, medicine, water treatment, and many other applications.

<div class="mw-heading mw-heading2"><h2 id="History">History</h2></div>

In 1909, the <a href="Danish_people" class="mw-redirect" title="Danish people">Danish</a> chemist <a href="S._P._L._S%C3%B8rensen" title="S. P. L. Sørensen">Søren Peter Lauritz Sørensen</a> introduced the concept of pH at the <a href="Carlsberg_Laboratory" title="Carlsberg Laboratory">Carlsberg Laboratory</a>,<a href="#cite_note-Sørensen2-4">[4]</a> originally using the notation "pH•", with H• as a subscript to the lowercase p. The concept was later revised in 1924 to the modern pH to accommodate definitions and measurements in terms of <a href="Electrochemical_cells" class="mw-redirect" title="Electrochemical cells">electrochemical cells</a>.<blockquote>For the sign p, I propose the name 'hydrogen ion exponent' and the symbol pH•. Then, for the hydrogen ion exponent (pH•) of a solution, the negative value of the <a href="Common_logarithm" title="Common logarithm">Briggsian logarithm</a> of the related hydrogen ion <a href="Equivalent_concentration" title="Equivalent concentration">normality factor</a> is to be understood.<a href="#cite_note-Sørensen2-4">[4]</a></blockquote>Sørensen did not explain why he used the letter p, and the exact meaning of the letter is still disputed.<a href="#cite_note-5">[5]</a><a href="#cite_note-6">[6]</a> Sørensen described a way of measuring pH using potential differences, and it represents the negative power of 10 in the concentration of hydrogen ions. The letter p could stand for the French puissance, German Potenz, or Danish potens, all meaning "power", or it could mean "potential". All of these words start with the letter p in <a href="French_language" title="French language">French</a>, <a href="German_language" title="German language">German</a>, and <a href="Danish_language" title="Danish language">Danish</a>, which were the languages in which Sørensen published: Carlsberg Laboratory was French-speaking; German was the dominant language of scientific publishing; Sørensen was Danish. He also used the letter q in much the same way elsewhere in the paper, and he might have arbitrarily labelled the test solution "p" and the reference solution "q"; these letters are often paired with e4 then e5.<a href="#cite_note-7">[7]</a> Some literature sources suggest that "pH" stands for the <a href="Latin_language" class="mw-redirect" title="Latin language">Latin term</a> pondus hydrogenii (quantity of hydrogen) or potentia hydrogenii (power of hydrogen), although this is not supported by Sørensen's writings.<a href="#cite_note-Otterson2-8">[8]</a><a href="#cite_note-Lian2-9">[9]</a><a href="#cite_note-Bradley2-10">[10]</a>

In modern <a href="Chemistry" title="Chemistry">chemistry</a>, the p stands for "the negative <a href="Common_logarithm" title="Common logarithm">decimal logarithm</a> of", and is used in the term pKa for <a href="Acid_dissociation_constant" title="Acid dissociation constant">acid dissociation constants</a>,<a href="#cite_note-Jens2-11">[11]</a> so pH is "the negative <a href="Common_logarithm" title="Common logarithm">decimal logarithm of</a> H+ ion concentration", while pOH is "the negative decimal logarithm of OH− ion concentration".

American bacteriologist <a href="Alice_Catherine_Evans" title="Alice Catherine Evans">Alice Catherine Evans</a>, who influenced <a href="Dairy" title="Dairy">dairying</a> and <a href="Food_safety" title="Food safety">food safety</a>, credited <a href="William_Mansfield_Clark" title="William Mansfield Clark">William Mansfield Clark</a> and colleagues, including herself, with developing pH measuring methods in the 1910s, which had a wide influence on laboratory and industrial use thereafter. In her memoir, she does not mention how much, or how little, Clark and colleagues knew about Sørensen's work a few years prior.<a href="#cite_note-Evans-Memoirs2-12">[12]</a> She said:<blockquote>In these studies [of bacterial metabolism] Dr. Clark's attention was directed to the effect of acid on the growth of bacteria. He found that it is the intensity of the acid in terms of hydrogen-ion concentration that affects their growth. But existing methods of measuring acidity determined the quantity, not the intensity, of the acid. Next, with his collaborators, Dr. Clark developed accurate methods for measuring hydrogen-ion concentration. These methods replaced the inaccurate titration method of determining the acid content in use in biologic laboratories throughout the world. Also they were found to be applicable in many industrial and other processes in which they came into wide usage.<a href="#cite_note-Evans-Memoirs2-12">[12]</a></blockquote>The first <a href="Electronics" title="Electronics">electronic</a> method for measuring pH was invented by <a href="Arnold_Orville_Beckman" class="mw-redirect" title="Arnold Orville Beckman">Arnold Orville Beckman</a>, a professor at the <a href="California_Institute_of_Technology" title="California Institute of Technology">California Institute of Technology</a> in 1934.<a href="#cite_note-13">[13]</a> It was in response to a request from the local citrus grower <a href="Sunkist_Growers%2C_Incorporated" title="Sunkist Growers, Incorporated">Sunkist</a>, which wanted a better method for quickly testing the pH of lemons they were picking from their nearby orchards.<a href="#cite_note-14">[14]</a>

<div class="mw-heading mw-heading2"><h2 id="Definition">Definition</h2></div>

The pH of a solution is defined as the decimal <a href="Logarithm" title="Logarithm">logarithm</a> of the reciprocal of the <a href="Hydron" title="Hydron">hydrogen ion</a> <a href="Activity_(chemistry)" class="mw-redirect" title="Activity (chemistry)">activity</a>, aH+.<a href="#cite_note-covington3-3">[3]</a> Mathematically, pH is expressed as:

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<annotation encoding="application/x-tex">{\displaystyle {\ce {pH}}=-\log _{10}(a_{{\ce {H+}}})=\log _{10}\left({\frac {1}{a_{{\ce {H+}}}}}\right)}</annotation>

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</math><img src="./932bace628e5b8db56e6220cc82210336491b3ab.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -2.505ex; width:34.984ex; height:6.176ex;" alt="{\displaystyle {\ce {pH}}=-\log _{10}(a_{{\ce {H+}}})=\log _{10}\left({\frac {1}{a_{{\ce {H+}}}}}\right)}" loading="lazy"></dd></dl>

For example, for a solution with a hydrogen ion activity of 5×10−6 <a href="Mole_(unit)" title="Mole (unit)">mol</a>/<a href="Litre" title="Litre">L</a> (i.e., the concentration of hydrogen cations), the pH of the solution can be calculated as follows:

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<annotation encoding="application/x-tex">{\displaystyle {\ce {pH}}=-\log _{10}(5\times 10^{-6})=5.3}</annotation>

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</math><img src="./1cf76a8ab825684516143408b22f4c3f8435fa8d.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -0.838ex; width:29.717ex; height:3.176ex;" alt="{\displaystyle {\ce {pH}}=-\log _{10}(5\times 10^{-6})=5.3}" loading="lazy"></dd></dl>

The concept of pH was developed because <a href="Ion-selective_electrodes" class="mw-redirect" title="Ion-selective electrodes">ion-selective electrodes</a>, which are used to measure pH, respond to activity. The electrode potential, E, follows the <a href="Nernst_equation" title="Nernst equation">Nernst equation</a> for the hydrogen cation, which can be expressed as:

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<annotation encoding="application/x-tex">{\displaystyle E=E^{0}+{\frac {RT}{F}}\ln(a_{{\ce {H+}}})=E^{0}-{\frac {RT\ \ln {10}}{F}}{\ce {pH}}\approx E^{0}-{\frac {2.303\ RT}{F}}{\ce {pH}}}</annotation>

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</math><img src="./52baacfee9e6ff5c415c2194743c1c96d5050834.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -1.838ex; width:66.446ex; height:5.343ex;" alt="{\displaystyle E=E^{0}+{\frac {RT}{F}}\ln(a_{{\ce {H+}}})=E^{0}-{\frac {RT\ \ln {10}}{F}}{\ce {pH}}\approx E^{0}-{\frac {2.303\ RT}{F}}{\ce {pH}}}" loading="lazy"></dd></dl>

where E is a measured potential, E0 is the standard electrode potential, R is the <a href="Molar_gas_constant" class="mw-redirect" title="Molar gas constant">molar gas constant</a>, T is the thermodynamic temperature, F is the <a href="Faraday_constant" title="Faraday constant">Faraday constant</a>. For H+, the number of electrons transferred is one. The electrode potential is proportional to pH when pH is defined in terms of activity.

The precise measurement of pH is presented in International Standard <a href="ISO_31-8" title="ISO 31-8">ISO 31-8</a> as follows:<a href="#cite_note-15">[15]</a> A <a href="Galvanic_cell" title="Galvanic cell">galvanic cell</a> is set up to measure the <a href="Electromotive_force" title="Electromotive force">electromotive force</a> (e.m.f.) between a reference electrode and an electrode sensitive to the hydrogen ion activity when they are both immersed in the same aqueous solution. The reference electrode may be a <a href="Silver_chloride_electrode" title="Silver chloride electrode">silver chloride electrode</a> or a <a href="Saturated_calomel_electrode" title="Saturated calomel electrode">calomel electrode</a>, and the hydrogen-ion selective electrode is a <a href="Standard_hydrogen_electrode" title="Standard hydrogen electrode">standard hydrogen electrode</a>.

<dl><dd>Reference electrode | concentrated solution of KCl || test solution | H2 | Pt</dd></dl>

Firstly, the cell is filled with a solution of known hydrogen ion activity and the electromotive force, ES, is measured. Then the electromotive force, EX, of the same cell containing the solution of unknown pH is measured.

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<annotation encoding="application/x-tex">{\displaystyle {\ce {pH(X)}}={\ce {pH(S)}}+{\frac {E_{{\ce {S}}}-E_{{\ce {X}}}}{z}}}</annotation>

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</math><img src="./61d7f14649cf22c1e0a8f70a165821e5367fa721.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -1.838ex; width:28.382ex; height:5.343ex;" alt="{\displaystyle {\ce {pH(X)}}={\ce {pH(S)}}+{\frac {E_{{\ce {S}}}-E_{{\ce {X}}}}{z}}}" loading="lazy"></dd></dl>

The difference between the two measured electromotive force values is proportional to pH. This method of calibration avoids the need to know the <a href="Standard_electrode_potential" title="Standard electrode potential">standard electrode potential</a>. The proportionality constant, 1/z, is ideally equal to <math xmlns="http://www.w3.org/1998/Math/MathML" alttext="{\displaystyle {\frac {F}{RT\ln {10}}}\ }">

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<annotation encoding="application/x-tex">{\displaystyle {\frac {F}{RT\ln {10}}}\ }</annotation>

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</math><img src="./da90052e3fc4ff151c53de27ac3d0802bb2f50bb.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -2.005ex; width:9.856ex; height:5.343ex;" alt="{\displaystyle {\frac {F}{RT\ln {10}}}\ }" loading="lazy">, the "Nernstian slope".

In practice, a <a href="Glass_electrode" title="Glass electrode">glass electrode</a> is used instead of the cumbersome hydrogen electrode. A combined glass electrode has an in-built reference electrode. It is calibrated against <a href="Buffer_solution" title="Buffer solution">Buffer solutions</a> of known hydrogen ion (H+) activity proposed by the International Union of Pure and Applied Chemistry (<a href="IUPAC" class="mw-redirect" title="IUPAC">IUPAC</a>).<a href="#cite_note-covington3-3">[3]</a> Two or more buffer solutions are used in order to accommodate the fact that the "slope" may differ slightly from ideal. To calibrate the electrode, it is first immersed in a standard solution, and the reading on a <a href="PH_meter" title="PH meter">pH meter</a> is adjusted to be equal to the standard buffer's value. The reading from a second standard buffer solution is then adjusted using the "slope" control to be equal to the pH for that solution. Further details, are given in the <a href="IUPAC" class="mw-redirect" title="IUPAC">IUPAC</a> recommendations.<a href="#cite_note-covington22-16">[16]</a> When more than two buffer solutions are used the electrode is calibrated by fitting observed pH values to a straight line with respect to standard buffer values. Commercial standard buffer solutions usually come with information on the value at 25 °C and a correction factor to be applied for other temperatures.

The pH scale is logarithmic and therefore pH is a <a href="Dimensionless_quantity" title="Dimensionless quantity">dimensionless quantity</a>.<a href="#cite_note-17">[17]</a>

This was the original definition of Sørensen in 1909,<a href="#cite_note-Sor2-18">[18]</a> which was superseded in favor of pH in 1924. [H] is the concentration of hydrogen ions, denoted [H+] in modern chemistry. More correctly, the <a href="Thermodynamic_activity" title="Thermodynamic activity">thermodynamic activity</a> of H+ (aH+) in dilute solution should be replaced by [H+]/c0, where the standard state concentration c0 = 1 mol/L. This ratio is a pure number whose logarithm can be defined.

It is possible to measure the concentration of hydrogen cations directly using an electrode calibrated in terms of hydrogen ion concentrations. One common method is to <a href="Titration" title="Titration">titrate</a> a solution of known concentration of a strong acid with a solution of known concentration of strong base in the presence of a relatively high concentration of background electrolyte. By knowing the concentrations of the acid and base, the concentration of hydrogen cations can be calculated and the measured potential can be correlated with concentrations. The calibration is usually carried out using a <a href="Gran_plot#Electrode_calibration" title="Gran plot">Gran plot</a>.<a href="#cite_note-19">[19]</a> This procedure makes the activity of hydrogen cations equal to the numerical value of concentration of these ions.

The glass electrode (and other <a href="Ion_selective_electrode" class="mw-redirect" title="Ion selective electrode">Ion selective electrodes</a>) should be calibrated in a medium similar to the one being investigated. For instance, if one wishes to measure the pH of a seawater sample, the electrode should be calibrated in a solution resembling seawater in its chemical composition.

The difference between p[H] and pH is quite small, and it has been stated that pH = p[H] + 0.04.<a href="#cite_note-20">[20]</a> However, it is common practice to use the term "pH" for both types of measurement.

pOH is sometimes used as a measure of the concentration of hydroxide ions, OH−. By definition, pOH is the negative logarithm (to the base 10) of the hydroxide ion concentration (mol/L). pOH values can be derived from pH measurements and vice-versa. The concentration of hydroxide ions in water is related to the concentration of hydrogen cations by

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<annotation encoding="application/x-tex">{\displaystyle [{\ce {OH^-}}]={\frac {K_{{\ce {W}}}}{[{\ce {H^+}}]}}}</annotation>

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</math><img src="./0e914d2d8665e01b649b6ecc09acdb2bb823b33c.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -2.838ex; width:14.838ex; height:6.343ex;" alt="{\displaystyle [{\ce {OH^-}}]={\frac {K_{{\ce {W}}}}{[{\ce {H^+}}]}}}" loading="lazy"></dd></dl>

where KW is the <a href="Self-ionization_of_water" title="Self-ionization of water">self-ionization</a> constant of water. Taking <a href="Logarithm" title="Logarithm">logarithms</a>,

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<annotation encoding="application/x-tex">{\displaystyle {\ce {pOH}}={\ce {p}}K_{{\ce {W}}}-{\ce {pH}}.}</annotation>

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</math><img src="./759d32a09b7693aeef989fecebaa74adce9da236.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -0.671ex; width:19.652ex; height:2.509ex;" alt="{\displaystyle {\ce {pOH}}={\ce {p}}K_{{\ce {W}}}-{\ce {pH}}.}" loading="lazy"></dd></dl>

So, at room temperature, pOH ≈ 14 − pH. (Under otherwise standard conditions, the equation is exact at approximately 24.87 °C.) However this relationship is not strictly valid in other circumstances, such as in measurements of <a href="Alkaline_soils" class="mw-redirect" title="Alkaline soils">soil alkalinity</a>.

<div class="mw-heading mw-heading2"><h2 id="Measurement">Measurement</h2></div>

<div class="mw-heading mw-heading3"><h3 id="pH_Indicators">pH Indicators</h3></div>

<caption>Average pH of common solutions

</caption>

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<td><a href="Sulfuric_acid" title="Sulfuric acid">Battery acid</a>

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<td style="background-color: #CC0000; text-align: center; color: #ffffff">< 1

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<td><a href="Gastric_acid" title="Gastric acid">Gastric acid</a>

</td>

<td style="background-color: #EE0000; text-align: center; color: #ffffff">1.0–1.5

</td></tr>

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<td><a href="Vinegar" title="Vinegar">Vinegar</a>

</td>

<td style="background-color: #FF4000; text-align: center">2–3

</td></tr>

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<td><a href="Orange_juice" title="Orange juice">Orange juice</a>

</td>

<td style="background-color: #FF6600; text-align: center">3.3–4.2

</td></tr>

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<td><a href="Coffee" title="Coffee">Black coffee</a>

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<td style="background-color: #ffff00 ; text-align: center">5–5.03

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<td style="background-color: #99cc33; text-align: center">6.5–6.8

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<td><a href="Pure_water" class="mw-redirect" title="Pure water">Pure water</a> at 25 °C

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<td style="background-color: #339933; text-align: center; color: #ffffff">7

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<td style="text-align: center">Neutral

</td></tr>

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<td><a href="Sea_water" class="mw-redirect" title="Sea water">Sea water</a>

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<td style="background-color: #19cdff; text-align: center; color: #000000">7.5–8.4

</td>

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<td style="background-color: #3333ff; text-align: center; color: #FFFFFF">11.0–11.5

</td></tr>

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<td><a href="Bleach" title="Bleach">Bleach</a>

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<td style="background-color: #330099; text-align: center; color: #FFFFFF">12.5

</td></tr>

<tr>

</td>

<td style="background-color: #330066; text-align: center; color: #FFFFFF">14

</td></tr></tbody></table>

<div role="note" class="hatnote navigation-not-searchable">Main article: <a href="PH_indicator" title="PH indicator">pH indicator</a></div>

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 <a href="Colorimeter_(chemistry)" title="Colorimeter (chemistry)">colorimeter</a> or <a href="Spectrophotometer" class="mw-redirect" title="Spectrophotometer">spectrophotometer</a>. A <a href="Universal_indicator" title="Universal indicator">Universal indicator</a> 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 <a href="PH_meter" title="PH meter">pH meter</a>, which directly measures the voltage difference between a pH-sensitive electrode and a reference electrode.

<div class="mw-heading mw-heading3"><h3 id="Non-aqueous_solutions">Non-aqueous solutions</h3></div>

pH values can be measured in non-aqueous solutions, but they are based on a different scale from aqueous pH values because the <a href="Standard_state" title="Standard state">standard states</a> used for calculating hydrogen ion concentrations (<a href="Activity_(chemistry)" class="mw-redirect" title="Activity (chemistry)">activities</a>) are different. The hydrogen ion activity, aH+, is defined<a href="#cite_note-GoldBook2-21">[21]</a><a href="#cite_note-GreenBook2-22">[22]</a> as:

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<annotation encoding="application/x-tex">{\displaystyle a_{{\ce {H+}}}=\exp \left({\frac {\mu _{{\ce {H+}}}-\mu _{{\ce {H+}}}^{\ominus }}{RT}}\right)}</annotation>

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</math><img src="./c2554db9edcefbf9fbe64b965caf52d88fbef11c.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -3.171ex; width:26.043ex; height:7.509ex;" alt="{\displaystyle a_{{\ce {H+}}}=\exp \left({\frac {\mu _{{\ce {H+}}}-\mu _{{\ce {H+}}}^{\ominus }}{RT}}\right)}" loading="lazy"></dd></dl>

where μH+ is the <a href="Chemical_potential" title="Chemical potential">chemical potential</a> of the hydrogen cation, <math xmlns="http://www.w3.org/1998/Math/MathML" alttext="{\displaystyle \mu _{{\ce {H+}}}^{\ominus }}">

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<annotation encoding="application/x-tex">{\displaystyle \mu _{{\ce {H+}}}^{\ominus }}</annotation>

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</math><img src="./4b94d142b00deab67cbdaa710bcd956bb12cc916.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -1.338ex; width:4.069ex; height:3.509ex;" alt="{\displaystyle \mu _{{\ce {H+}}}^{\ominus }}" loading="lazy"> is its chemical potential in the chosen standard state, R is the <a href="Molar_gas_constant" class="mw-redirect" title="Molar gas constant">molar gas constant</a> and T is the <a href="Thermodynamic_temperature" title="Thermodynamic temperature">thermodynamic temperature</a>. 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 <a href="Lyonium_ion" title="Lyonium ion">hydronium/lyonium ion</a>.

pH is an example of an <a href="Acidity_function" title="Acidity function">acidity function</a>, but others can be defined. For example, the <a href="Hammett_acidity_function" title="Hammett acidity function">Hammett acidity function</a>, H0, has been developed in connection with <a href="Superacid" title="Superacid">Superacids</a>.

<div class="mw-heading mw-heading3"><h3 id="Unified_absolute_pH_scale">Unified absolute pH scale</h3></div>

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 <a href="Lewis_acids_and_bases" title="Lewis acids and bases">Lewis acid–base</a> theory. This scale applies to liquids, gases, and even solids.<a href="#cite_note-Krossing2-23">[23]</a> 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.

<div class="mw-heading mw-heading3"><h3 id="Extremes_of_pH_measurements">Extremes of pH measurements</h3></div>

<div role="note" class="hatnote navigation-not-searchable">"Negative pH" redirects here. For the band, see <a href="Negative_pH_(band)" title="Negative pH (band)">Negative pH (band)</a>.</div>

The measurement of pH can become difficult at extremely acidic or alkaline conditions, such as below pH 2.5 (ca. 0.003 <a href="Mole_(unit)" title="Mole (unit)">mol</a>/dm3 acid) or above pH 10.5 (above ca. 0.0003 mol/dm3 alkaline). This is due to the breakdown of the <a href="Nernst_equation" title="Nernst equation">Nernst equation</a> in such conditions when using a glass electrode. Several factors contribute to this problem. First, <a href="Liquid_junction_potential" title="Liquid junction potential">liquid junction potentials</a> may not be independent of pH.<a href="#cite_note-Feldman2-24">[24]</a> Second, the high <a href="Ionic_strength" title="Ionic strength">ionic strength</a> 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 Na+ and K+ in the solution.<a href="#cite_note-25">[25]</a> 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.<a href="#cite_note-26">[26]</a>

<div class="mw-heading mw-heading2"><h2 id="Applications">Applications</h2></div>

Pure water has a pH of 7 at 25 °C, meaning it is neutral. When an <a href="Acid" title="Acid">acid</a> is dissolved in water, the pH will be less than 7, while a <a href="Base_(chemistry)" title="Base (chemistry)">base</a>, or <a href="Alkali" title="Alkali">alkali</a>, will have a pH greater than 7. A strong acid, such as <a href="Hydrochloric_acid" title="Hydrochloric acid">hydrochloric acid</a>, at concentration 1 mol/L has a pH of 0, while a strong alkali like <a href="Sodium_hydroxide" title="Sodium hydroxide">sodium hydroxide</a>, 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 °C, but 7 serves as a good approximation in most cases. Neutrality occurs when the concentration of hydrogen cations ([H+]) equals the concentration of hydroxide ions ([OH−]), or when their activities are equal. Since <a href="Self-ionization_of_water" title="Self-ionization of water">self-ionization of water</a> holds the product of these concentration [H+] × [OH−] = Kw, it can be seen that at neutrality [H+] = [OH−] = √Kw, or pH = pKw/2. pKw is approximately 14 but depends on ionic strength and temperature, and so the pH of neutrality does also. Pure water and a solution of <a href="Sodium_chloride" title="Sodium chloride">NaCl</a> in pure water are both neutral, since <a href="Self-ionization_of_water" title="Self-ionization of water">dissociation of water</a> 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 <a href="Ionic_strength" title="Ionic strength">ionic strength</a>, so Kw varies with ionic strength.

When pure water is exposed to air, it becomes mildly acidic. This is because water absorbs <a href="Carbon_dioxide" title="Carbon dioxide">carbon dioxide</a> from the air, which is then slowly converted into <a href="Bicarbonate" title="Bicarbonate">bicarbonate</a> and hydrogen cations (essentially creating <a href="Carbonic_acid" title="Carbonic acid">carbonic acid</a>).

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<annotation encoding="application/x-tex">{\displaystyle {\ce {CO2 + H2O <=> HCO3^- + H^+}}}</annotation>

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</math><img src="./c9a2e896d9c84824826a17d8ed1d066dbdb23fd0.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -1.005ex; width:29.469ex; height:3.343ex;" alt="{\displaystyle {\ce {CO2 + H2O <=> HCO3^- + H^+}}}" loading="lazy"></dd></dl>

The United States Department of Agriculture <a href="Natural_Resources_Conservation_Service" title="Natural Resources Conservation Service">Natural Resources Conservation Service</a>, formerly Soil Conservation Service classifies <a href="Soil_pH" title="Soil pH">soil pH</a> ranges as follows:<a href="#cite_note-27">[27]</a>

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<th scope="col">Denomination

</th>

<th scope="col">pH range

</th></tr>

<tr>

<td>Ultra acidic

</td>

<td>< 3.5

</td></tr>

<tr>

<td>Extremely acidic

</td>

<td>3.5–4.4

</td></tr>

<tr>

<td>Very strongly acidic

</td>

<td>4.5–5.0

</td></tr>

<tr>

<td>Strongly acidic

</td>

<td>5.1–5.5

</td></tr>

<tr>

<td>Moderately acidic

</td>

<td>5.6–6.0

</td></tr>

<tr>

<td>Slightly acidic

</td>

<td>6.1–6.5

</td></tr>

<tr>

<td>Neutral

</td>

<td>6.6–7.3

</td></tr>

<tr>

<td>Slightly alkaline

</td>

<td>7.4–7.8

</td></tr>

<tr>

<td>Moderately alkaline

</td>

<td>7.9–8.4

</td></tr>

<tr>

<td>Strongly alkaline

</td>

<td>8.5–9.0

</td></tr>

<tr>

<td>Very strongly alkaline

</td>

<td>9.0–10.5

</td></tr>

<tr>

<td>Hyper alkaline

</td>

<td>> 10.5

</td></tr></tbody></table>

Topsoil pH is influenced by soil parent material, erosional effects, climate and vegetation. A recent map<a href="#cite_note-28">[28]</a> 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.

<div class="mw-heading mw-heading3"><h3 id="pH_in_plants">pH in plants</h3></div>

Plants contain pH-dependent <a href="Plant_pigment" class="mw-redirect" title="Plant pigment">pigments</a> that can be used as <a href="PH_indicator" title="PH indicator">pH indicators</a>, such as those found in <a href="Hibiscus" title="Hibiscus">hibiscus</a>, <a href="Red_cabbage" title="Red cabbage">red cabbage</a> (<a href="Anthocyanin" title="Anthocyanin">anthocyanin</a>), and grapes (<a href="Red_wine" title="Red wine">red wine</a>). <a href="Citrus" title="Citrus">Citrus</a> fruits have acidic juice primarily due to the presence of <a href="Citric_acid" title="Citric acid">citric acid</a>, while other <a href="Carboxylic_acid" title="Carboxylic acid">carboxylic acids</a> can be found in various living systems. The <a href="Protonation" title="Protonation">protonation</a> state of <a href="Phosphate" title="Phosphate">phosphate</a> derivatives, including <a href="Adenosine_triphosphate" title="Adenosine triphosphate">ATP</a>, is pH-dependent. <a href="Hemoglobin" title="Hemoglobin">Hemoglobin</a>, an oxygen-transport enzyme, is also affected by pH in a phenomenon known as the <a href="Root_effect" title="Root effect">Root effect</a>.

<div class="mw-heading mw-heading3"><h3 id="pH_in_the_ocean">pH in the ocean</h3></div>

<div role="note" class="hatnote navigation-not-searchable">See also: <a href="Seawater#pH_value" title="Seawater">Seawater § pH value</a>, <a href="Ocean#pH_and_alkalinity" title="Ocean">Ocean § pH and alkalinity</a>, and <a href="Ocean_acidification" title="Ocean acidification">Ocean acidification</a></div>

The pH of <a href="Seawater" title="Seawater">seawater</a> plays an important role in the ocean's <a href="Carbon_cycle#Ocean" title="Carbon cycle">carbon cycle</a>. There is evidence of ongoing <a href="Ocean_acidification" title="Ocean acidification">ocean acidification</a> (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.<a href="#cite_note-29">[29]</a> <a href="Carbon_dioxide_emissions" class="mw-redirect" title="Carbon dioxide emissions">Carbon dioxide emissions</a> from human activities are the primary cause of ocean acidification, with <a href="Carbon_dioxide_in_Earth's_atmosphere" class="mw-redirect" title="Carbon dioxide in Earth's atmosphere">atmospheric carbon dioxide levels</a> at 430 ppm CO2 at <a href="Mauna_Loa" title="Mauna Loa">Mauna Loa</a> observatory in 2025.<a href="#cite_note-NOAA_CO2-30">[30]</a> In 2024, the annual atmospheric CO2 increase measured by the <a href="National_Oceanic_and_Atmospheric_Administration" title="National Oceanic and Atmospheric Administration">NOAA</a>’s Global Monitoring Laboratory was 3.75 ppm CO2/year.<a href="#cite_note-Berwyn2025-31">[31]</a> CO2 from the <a href="Atmosphere" title="Atmosphere">atmosphere</a> is absorbed by the oceans. This produces <a href="Carbonic_acid" title="Carbonic acid">carbonic acid</a> (H2CO3) which dissociates into a <a href="Bicarbonate_ion" class="mw-redirect" title="Bicarbonate ion">bicarbonate ion</a> (HCO− 3) and a <a href="Hydron" title="Hydron">hydrogen cation</a> (H+). The presence of free hydrogen cations (H+) lowers the pH of the ocean.

<div class="mw-heading mw-heading4"><h4 id="Three_pH_scales_in_oceanography">Three pH scales in oceanography</h4></div>

The measurement of pH in seawater is complicated by the <a href="Chemical_property" title="Chemical property">chemical properties</a> of seawater, and three distinct pH scales exist in <a href="Chemical_oceanography" class="mw-redirect" title="Chemical oceanography">chemical oceanography</a>.<a href="#cite_note-zeebe2-32">[32]</a> 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 <a href="Total_inorganic_carbon" title="Total inorganic carbon">carbonate system</a>.<a href="#cite_note-zeebe2-32">[32]</a> 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 <a href="Operational_definition" title="Operational definition">operational definition</a> of the pH scale, the <a href="IUPAC" class="mw-redirect" title="IUPAC">IUPAC</a> defines a series of <a href="Buffer_solution" title="Buffer solution">Buffer solutions</a> across a range of pH values (often denoted with <a href="National_Bureau_of_Standards" class="mw-redirect" title="National Bureau of Standards">National Bureau of Standards</a> (NBS) or <a href="National_Institute_of_Standards_and_Technology" title="National Institute of Standards and Technology">National Institute of Standards and Technology</a> (NIST) designation). These solutions have a relatively low <a href="Ionic_strength" title="Ionic strength">ionic strength</a> (≈ 0.1) compared to that of seawater (≈ 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 <a href="Standard_electrode_potential" title="Standard electrode potential">electrode potential</a>. To resolve this problem, an alternative series of buffers based on <a href="Artificial_seawater" title="Artificial seawater">artificial seawater</a> was developed.<a href="#cite_note-33">[33]</a> 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 pHT. The total scale was defined using a medium containing <a href="Sulfate" title="Sulfate">sulfate</a> ions. These ions experience <a href="Protonation" title="Protonation">protonation</a>, H+ + SO2− 4↔ HSO− 4, such that the total scale includes the effect of both <a href="Proton" title="Proton">protons</a> (free hydrogen cations) and hydrogen sulfate ions:

An alternative scale, the free scale, often denoted pHF, omits this consideration and focuses solely on [H+]F, in principle making it a simpler representation of hydrogen ion concentration. Only [H+]T can be determined,<a href="#cite_note-34">[34]</a> therefore [H+]F must be estimated using the [SO2− 4] and the stability constant of HSO− 4, K* S:

However, it is difficult to estimate K* S in seawater, limiting the utility of the otherwise more straightforward free scale.

Another scale, known as the seawater scale, often denoted pHSWS, takes account of a further protonation relationship between hydrogen cations and <a href="Fluoride" title="Fluoride">fluoride</a> ions, H+ + F− ⇌ HF. Resulting in the following expression for [H+]SWS:

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. 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:

<dd>pHSWS = −log10(H+]F + [HSO− 4] + [HF]) = −log10[v]SWS</dd></dl>

The pH level of food influences its flavor, texture, and <a href="Shelf_life" title="Shelf life">shelf life</a>.<a href="#cite_note-35">[35]</a> Acidic foods, such as <a href="Citrus_fruits" class="mw-redirect" title="Citrus fruits">citrus fruits</a>, tomatoes, and <a href="Vinegar" title="Vinegar">vinegar</a>, typically have a pH below 4.6<a href="#cite_note-okla-36">[36]</a> with sharp and tangy taste, while basic foods taste bitter or soapy.<a href="#cite_note-37">[37]</a> Maintaining the appropriate pH in foods is essential for preventing the growth of harmful <a href="Microorganisms" class="mw-redirect" title="Microorganisms">microorganisms</a>.<a href="#cite_note-okla-36">[36]</a> The alkalinity of vegetables such as <a href="Spinach" title="Spinach">spinach</a> and <a href="Kale" title="Kale">kale</a> can also influence their texture and color during cooking.<a href="#cite_note-38">[38]</a> The pH also influences the <a href="Maillard_reaction" title="Maillard reaction">Maillard reaction</a>, which is responsible for the browning of food during cooking, impacting both flavor and appearance.<a href="#cite_note-39">[39]</a>

<div class="mw-heading mw-heading3"><h3 id="pH_of_various_body_fluids">pH of various body fluids</h3></div>

<caption>pH of various body fluids<a href="#cite_note-Boron2012-40">[40]</a>

</caption>

<th>Compartment

</th>

<th>pH

</th></tr>

<tr>

<td><a href="Gastric_acid" title="Gastric acid">Gastric acid</a></td>

</td></tr>

<tr>

<td><a href="Lysosome" title="Lysosome">Lysosomes</a></td>

</td></tr>

<tr>

<td><a href="Human_skin" title="Human skin">Human skin</a></td>

</td></tr>

<tr>

<td>Granules of <a href="Chromaffin_cell" title="Chromaffin cell">chromaffin cells</a></td>

<td>5.5

</td></tr>

<tr>

<td><a href="Urine" title="Urine">Urine</a></td>

<td>6.0

</td></tr>

<tr>

<td><a href="Breast_milk" title="Breast milk">Breast milk</a></td>

</td></tr>

<tr>

<td><a href="Cytosol" title="Cytosol">Cytosol</a></td>

<td>7.2

</td></tr>

<tr>

<td><a href="Blood" title="Blood">Blood</a> (natural pH)</td>

</td></tr>

<tr>

<td><a href="Cerebrospinal_fluid" title="Cerebrospinal fluid">Cerebrospinal fluid</a> (CSF)</td>

<td>7.5

</td></tr>

<tr>

<td><a href="Mitochondrial_matrix" title="Mitochondrial matrix">Mitochondrial matrix</a></td>

<td>7.5

</td></tr>

<tr>

<td><a href="Pancreas" title="Pancreas">Pancreas</a> secretions</td>

<td>8.1

</td></tr></tbody></table></dd></dl>

In living organisms, the pH of various <a href="Body_fluid" title="Body fluid">body fluids</a>, cellular compartments, and organs is tightly regulated to maintain a state of acid–base balance known as <a href="Acid%E2%80%93base_homeostasis" title="Acid–base homeostasis">acid–base homeostasis</a>. <a href="Acidosis" title="Acidosis">Acidosis</a>, 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, <a href="Alkalosis" title="Alkalosis">alkalosis</a> 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. <a href="Dental_plaque" title="Dental plaque">Plaque</a> formation in teeth can create a local acidic environment that results in <a href="Tooth_decay" title="Tooth decay">tooth decay</a> through demineralization. <a href="Enzyme" title="Enzyme">Enzymes</a> and other <a href="Protein" title="Protein">Proteins</a> have an optimal pH range for function and can become inactivated or <a href="Denaturation_(biochemistry)" title="Denaturation (biochemistry)">denatured</a> outside this range.

<div class="mw-heading mw-heading2"><h2 id="pH_calculations">pH calculations</h2></div>

When calculating the pH of a solution containing acids or bases, a <a href="Determination_of_equilibrium_constants#Speciation_calculations" title="Determination of equilibrium constants">chemical speciation calculation</a> is used to determine the concentration of all chemical species present in the solution. The complexity of the procedure depends on the nature of the solution. Strong acids and bases are compounds that are almost completely dissociated in water, which simplifies the calculation. However, for weak acids, a <a href="Quadratic_equation" title="Quadratic equation">quadratic equation</a> must be solved, and for weak bases, a cubic equation is required. In general, a set of <a href="Non-linear" class="mw-redirect" title="Non-linear">non-linear</a> <a href="Simultaneous_equation" class="mw-redirect" title="Simultaneous equation">simultaneous equations</a> must be solved.

Water itself is a weak acid and a weak base, so its dissociation must be taken into account at high pH and low solute concentration (see <a href="Amphoterism" title="Amphoterism">Amphoterism</a>). It <a href="Self-ionization_of_water" title="Self-ionization of water">dissociates</a> according to the equilibrium

with a <a href="Acid_dissociation_constant" title="Acid dissociation constant">dissociation constant</a>, Kw defined as

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<annotation encoding="application/x-tex">{\displaystyle K_{\text{w}}={\ce {[H+][OH^{-}]}}}</annotation>

</semantics>

</math><img src="./b3cd9093573ebd180683045ee786fc8c3fd2260c.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -0.838ex; width:17.394ex; height:3.009ex;" alt="{\displaystyle K_{\text{w}}={\ce {[H+][OH^{-}]}}}" loading="lazy"></dd></dl>

where [H+] stands for the concentration of the aqueous <a href="Hydronium_ion" class="mw-redirect" title="Hydronium ion">hydronium ion</a> and [OH−] represents the concentration of the <a href="Hydroxide_ion" class="mw-redirect" title="Hydroxide ion">hydroxide ion</a>. This equilibrium needs to be taken into account at high pH and when the solute concentration is extremely low.

<div class="mw-heading mw-heading3"><h3 id="Strong_acids_and_bases">Strong acids and bases</h3></div>

<a href="Strong_acid" class="mw-redirect" title="Strong acid">Strong acids</a> and <a href="Strong_base" class="mw-redirect" title="Strong base">bases</a> are compounds that are essentially fully dissociated in water. This means that in an acidic solution, the concentration of hydrogen cations (H+) can be considered equal to the concentration of the acid. Similarly, in a basic solution, the concentration of hydroxide ions (OH−) can be considered equal to the concentration of the base. The pH of a solution is defined as the negative logarithm of the concentration of H+, and the pOH is defined as the negative logarithm of the concentration of OH−. For example, the pH of a 0.01 M solution of hydrochloric acid (HCl) is equal to 2 (pH = −log10(0.01)), while the pOH of a 0.01 M solution of sodium hydroxide (NaOH) is equal to 2 (pOH = −log10(0.01)), which corresponds to a pH of about 12.

However, self-ionization of water must also be considered when concentrations of a strong acid or base is very low or high. For instance, a 5×10−8 M solution of HCl would be expected to have a pH of 7.3 based on the above procedure, which is incorrect as it is acidic and should have a pH of less than 7. In such cases, the system can be treated as a mixture of the acid or base and water, which is an <a href="Amphoteric" class="mw-redirect" title="Amphoteric">amphoteric</a> substance. By accounting for the self-ionization of water, the true pH of the solution can be calculated. For example, a 5×10−8 M solution of HCl would have a pH of 6.89 when treated as a mixture of HCl and water. The self-ionization equilibrium of solutions of sodium hydroxide at higher concentrations must also be considered.<a href="#cite_note-45">[45]</a>

<div class="mw-heading mw-heading3"><h3 id="Weak_acids_and_bases">Weak acids and bases</h3></div>

A <a href="Weak_acid" class="mw-redirect" title="Weak acid">weak acid</a> or the conjugate acid of a weak base can be treated using the same formalism.

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<annotation encoding="application/x-tex">{\displaystyle {\ce {HA <=> H+ + A-}}}</annotation>

</semantics>

</math><img src="./f7495efc7639cfaaeed95f8c49fd6d6c7a34a55e.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -0.505ex; width:17.482ex; height:2.843ex;" alt="{\displaystyle {\ce {HA <=> H+ + A-}}}" loading="lazy"></li>

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<annotation encoding="application/x-tex">{\displaystyle {\ce {HA+ <=> H+ + A}}}</annotation>

</semantics>

</math><img src="./aeac633a5b16a7e1b8cf7e818e7ca2596354fc0a.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -0.505ex; width:17.482ex; height:2.843ex;" alt="{\displaystyle {\ce {HA+ <=> H+ + A}}}" loading="lazy"></li></ul>

First, an acid dissociation constant is defined as follows. Electrical charges are omitted from subsequent equations for the sake of generality

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<annotation encoding="application/x-tex">{\displaystyle K_{a}={\frac {{\ce {[H] [A]}}}{{\ce {[HA]}}}}}</annotation>

</semantics>

</math><img src="./f29a403e204523dcc5d4345e9da4125d6bccc8e0.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -2.671ex; width:13.083ex; height:6.509ex;" alt="{\displaystyle K_{a}={\frac {{\ce {[H] [A]}}}{{\ce {[HA]}}}}}" loading="lazy"></dd></dl>

and its value is assumed to have been determined by experiment. This being so, there are three unknown concentrations, [HA], [H+] and [A−] to determine by calculation. Two additional equations are needed. One way to provide them is to apply the law of <a href="Mass_conservation" class="mw-redirect" title="Mass conservation">mass conservation</a> in terms of the two "reagents" H and A.

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<annotation encoding="application/x-tex">{\displaystyle C_{{\ce {A}}}={\ce {[A]}}+{\ce {[HA]}}}</annotation>

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</math><img src="./9ebde93fda69a21943e8842a9a32da905b2f645e.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -0.838ex; width:16.882ex; height:2.843ex;" alt="{\displaystyle C_{{\ce {A}}}={\ce {[A]}}+{\ce {[HA]}}}" loading="lazy"></dd>

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<annotation encoding="application/x-tex">{\displaystyle C_{{\ce {H}}}={\ce {[H]}}+{\ce {[HA]}}}</annotation>

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</math><img src="./920b90697a63b1c14c8302e733fc525ba0d9585b.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -0.838ex; width:16.882ex; height:2.843ex;" alt="{\displaystyle C_{{\ce {H}}}={\ce {[H]}}+{\ce {[HA]}}}" loading="lazy"></dd></dl>

C stands for <a href="Analytical_concentration" class="mw-redirect" title="Analytical concentration">analytical concentration</a>. In some texts, one mass balance equation is replaced by an equation of charge balance. This is satisfactory for simple cases like this one, but is more difficult to apply to more complicated cases as those below. Together with the equation defining Ka, there are now three equations in three unknowns. When an acid is dissolved in water CA = CH = Ca, the concentration of the acid, so [A] = [H]. After some further algebraic manipulation an equation in the hydrogen ion concentration may be obtained.

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<annotation encoding="application/x-tex">{\displaystyle [{\ce {H}}]^{2}+K_{a}[{\ce {H}}]-K_{a}C_{a}=0}</annotation>

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</math><img src="./46348613cafd9be4dcf2b772575b5af47f080d54.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -0.838ex; width:25.983ex; height:3.176ex;" alt="{\displaystyle [{\ce {H}}]^{2}+K_{a}[{\ce {H}}]-K_{a}C_{a}=0}" loading="lazy"></dd></dl>

Solution of this <a href="Quadratic_equation" title="Quadratic equation">quadratic equation</a> gives the hydrogen ion concentration and hence p[H] or, more loosely, pH. This procedure is illustrated in an <a href="ICE_table" class="mw-redirect" title="ICE table">ICE table</a> which can also be used to calculate the pH when some additional (strong) acid or alkaline has been added to the system, that is, when CA ≠ CH.

For example, what is the pH of a 0.01 M solution of <a href="Benzoic_acid" title="Benzoic acid">benzoic acid</a>, pKa = 4.19?

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<annotation encoding="application/x-tex">{\displaystyle K_{a}=10^{-4.19}=6.46\times 10^{-5}}</annotation>

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</math><img src="./8ba5721f63f313ffed3e2d4f1c4c0470276b2445.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -0.671ex; width:27.663ex; height:3.009ex;" alt="{\displaystyle K_{a}=10^{-4.19}=6.46\times 10^{-5}}" loading="lazy"></li>

<li>Step 2: Set up the quadratic equation. <math xmlns="http://www.w3.org/1998/Math/MathML" alttext="{\displaystyle [{\ce {H}}]^{2}+6.46\times 10^{-5}[{\ce {H}}]-6.46\times 10^{-7}=0}">

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<annotation encoding="application/x-tex">{\displaystyle [{\ce {H}}]^{2}+6.46\times 10^{-5}[{\ce {H}}]-6.46\times 10^{-7}=0}</annotation>

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</math><img src="./3b00dbc68e7860d66750e413f778916b6b915cd9.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -0.838ex; width:40.334ex; height:3.176ex;" alt="{\displaystyle [{\ce {H}}]^{2}+6.46\times 10^{-5}[{\ce {H}}]-6.46\times 10^{-7}=0}" loading="lazy"></li>

<li>Step 3: Solve the quadratic equation. <math xmlns="http://www.w3.org/1998/Math/MathML" alttext="{\displaystyle [{\ce {H+}}]=7.74\times 10^{-4};\quad \mathrm {pH} =3.11}">

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<annotation encoding="application/x-tex">{\displaystyle [{\ce {H+}}]=7.74\times 10^{-4};\quad \mathrm {pH} =3.11}</annotation>

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</math><img src="./d23e6dd79b1f843749965771698d8a5403a236a4.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -0.838ex; width:32.903ex; height:3.176ex;" alt="{\displaystyle [{\ce {H+}}]=7.74\times 10^{-4};\quad \mathrm {pH} =3.11}" loading="lazy"></li></ul>

For alkaline solutions, an additional term is added to the mass-balance equation for hydrogen. Since the addition of hydroxide reduces the hydrogen ion concentration, and the hydroxide ion concentration is constrained by the self-ionization equilibrium to be equal to <math xmlns="http://www.w3.org/1998/Math/MathML" alttext="{\displaystyle {\frac {K_{w}}{{\ce {[H+]}}}}}">

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<annotation encoding="application/x-tex">{\displaystyle {\frac {K_{w}}{{\ce {[H+]}}}}}</annotation>

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</math><img src="./671c1d10d6bce20ed4de57be093835a6aad5019f.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -2.838ex; width:5.384ex; height:6.176ex;" alt="{\displaystyle {\frac {K_{w}}{{\ce {[H+]}}}}}" loading="lazy">, the resulting equation is:

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<annotation encoding="application/x-tex">{\displaystyle C_{\ce {H}}={\frac {[{\ce {H}}]+[{\ce {HA}}]-K_{w}}{\ce {[H]}}}}</annotation>

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</math><img src="./e4eaeb404231cb0a9a118cd2972dc991b67adaab.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -2.671ex; width:23.941ex; height:6.509ex;" alt="{\displaystyle C_{\ce {H}}={\frac {[{\ce {H}}]+[{\ce {HA}}]-K_{w}}{\ce {[H]}}}}" loading="lazy"></dd></dl>

<div class="mw-heading mw-heading3"><h3 id="General_method">General method</h3></div>

Some systems, such as with <a href="Polyprotic" class="mw-redirect" title="Polyprotic">polyprotic</a> acids, are amenable to spreadsheet calculations.<a href="#cite_note-46">[46]</a> With three or more reagents or when many complexes are formed with general formulae such as ApBqHr, the following general method can be used to calculate the pH of a solution. For example, with three reagents, each equilibrium is characterized by an equilibrium constant, β.

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<annotation encoding="application/x-tex">{\displaystyle [{\ce {A}}_{p}{\ce {B}}_{q}{\ce {H}}_{r}]=\beta _{pqr}[{\ce {A}}]^{p}[{\ce {B}}]^{q}[{\ce {H}}]^{r}}</annotation>

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</math><img src="./0116cc100cbe49e89c845e94bfb7cfa9d41efdf4.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -1.005ex; width:28.452ex; height:3.009ex;" alt="{\displaystyle [{\ce {A}}_{p}{\ce {B}}_{q}{\ce {H}}_{r}]=\beta _{pqr}[{\ce {A}}]^{p}[{\ce {B}}]^{q}[{\ce {H}}]^{r}}" loading="lazy"></dd></dl>

Next, write down the mass-balance equations for each reagent:

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<annotation encoding="application/x-tex">{\displaystyle {\begin{aligned}C_{\ce {A}}&=[{\ce {A}}]+\Sigma p\beta _{pqr}[{\ce {A}}]^{p}[{\ce {B}}]^{q}[{\ce {H}}]^{r}\\C_{\ce {B}}&=[{\ce {B}}]+\Sigma q\beta _{pqr}[{\ce {A}}]^{p}[{\ce {B}}]^{q}[{\ce {H}}]^{r}\\C_{\ce {H}}&=[{\ce {H}}]+\Sigma r\beta _{pqr}[{\ce {A}}]^{p}[{\ce {B}}]^{q}[{\ce {H}}]^{r}-K_{w}[{\ce {H}}]^{-1}\end{aligned}}}</annotation>

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</math><img src="./3d204f0d1baca5254f67602e86ba0091a004b673.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -4.338ex; width:43.079ex; height:9.676ex;" alt="{\displaystyle {\begin{aligned}C_{\ce {A}}&=[{\ce {A}}]+\Sigma p\beta _{pqr}[{\ce {A}}]^{p}[{\ce {B}}]^{q}[{\ce {H}}]^{r}\\C_{\ce {B}}&=[{\ce {B}}]+\Sigma q\beta _{pqr}[{\ce {A}}]^{p}[{\ce {B}}]^{q}[{\ce {H}}]^{r}\\C_{\ce {H}}&=[{\ce {H}}]+\Sigma r\beta _{pqr}[{\ce {A}}]^{p}[{\ce {B}}]^{q}[{\ce {H}}]^{r}-K_{w}[{\ce {H}}]^{-1}\end{aligned}}}" loading="lazy"></dd></dl>

There are no approximations involved in these equations, except that each stability constant is defined as a quotient of concentrations, not activities. Much more complicated expressions are required if activities are to be used.

There are three <a href="Simultaneous_equation" class="mw-redirect" title="Simultaneous equation">simultaneous equations</a> in the three unknowns, [A], [B] and [H]. Because the equations are non-linear and their concentrations may range over many powers of 10, the solution of these equations is not straightforward. However, many computer programs are available which can be used to perform these calculations. There may be more than three reagents. The calculation of hydrogen ion concentrations, using this approach, is a key element in the <a href="Determination_of_equilibrium_constants" title="Determination of equilibrium constants">determination of equilibrium constants</a> by <a href="Potentiometric_titration" title="Potentiometric titration">potentiometric titration</a>.

<ul><li><a href="PH_indicator" title="PH indicator">pH indicator</a></li>

<li><a href="Arterial_blood_gas" class="mw-redirect" title="Arterial blood gas">Arterial blood gas</a></li>

<li><a href="Chemical_equilibrium" title="Chemical equilibrium">Chemical equilibrium</a></li>

<div class="mw-heading mw-heading2"><h2 id="References">References</h2></div>

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<li id="cite_note-Boron2012-40">^ <a href="#cite_ref-Boron2012_40-0">a</a> <a href="#cite_ref-Boron2012_40-1">b</a> <a href="#cite_ref-Boron2012_40-2">c</a> <cite id="CITEREFBoronBoulpaep2012" class="citation book cs1">Boron, Walter F.; Boulpaep, Emile L. (13 January 2012). <a rel="nofollow" class="external text" href="https://books.google.com/books?id=54mxMgO5H_YC&dq=pH%20in%20living%20systems&pg=PA652">Medical Physiology: A Cellular And Molecular Approach</a> (2nd ed.). <a href="Elsevier_Health_Sciences" class="mw-redirect" title="Elsevier Health Sciences">Elsevier Health Sciences</a>, Saunders. pp. 652–671. <a href="ISBN_(identifier)" class="mw-redirect" title="ISBN (identifier)">ISBN</a> <bdi>9781455711819</bdi>. <a href="OCLC_(identifier)" class="mw-redirect" title="OCLC (identifier)">OCLC</a> <a rel="nofollow" class="external text" href="https://search.worldcat.org/oclc/1017876653">1017876653</a>. <a rel="nofollow" class="external text" href="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">Archived</a> from the original on 8 May 2022. Retrieved 8 May 2022.</cite>

</li>

<li id="cite_note-41"><a href="#cite_ref-41">^</a> <cite class="citation news cs1"><a rel="nofollow" class="external text" href="https://www.ucsfhealth.org/medical-tests/stomach-acid-test">"Stomach acid test"</a>. University of California San Francisco. Retrieved 21 February 2024.</cite>

</li>

<li id="cite_note-42"><a href="#cite_ref-42">^</a> <cite id="CITEREFMariebMitchell2011" class="citation book cs1">Marieb, Elaine N.; Mitchell, Susan J. (30 June 2011). <a rel="nofollow" class="external text" href="https://books.google.com/books?id=BxDfnQEACAAJ">Human anatomy & physiology</a>. San Francisco: <a href="Benjamin_Cummings" title="Benjamin Cummings">Benjamin Cummings</a>. <a href="ISBN_(identifier)" class="mw-redirect" title="ISBN (identifier)">ISBN</a> <bdi>9780321735287</bdi>. <a rel="nofollow" class="external text" href="https://web.archive.org/web/20220508051937/https://www.google.co.in/books/edition/_/BxDfnQEACAAJ?hl=en">Archived</a> from the original on 8 May 2022. Retrieved 8 May 2022.</cite>

</li>

<li id="cite_note-43"><a href="#cite_ref-43">^</a> <cite id="CITEREFLambersPiessensBloemPronk2006" class="citation journal cs1">Lambers, H.; Piessens, S.; Bloem, A.; Pronk, H.; Finkel, P. (1 October 2006). <a rel="nofollow" class="external text" href="https://onlinelibrary.wiley.com/doi/10.1111/j.1467-2494.2006.00344.x">"Natural skin surface pH is on average below 5, which is beneficial for its resident flora"</a>. International Journal of Cosmetic Science. 28 (5): 359–370. <a href="Doi_(identifier)" class="mw-redirect" title="Doi (identifier)">doi</a>:<a rel="nofollow" class="external text" href="https://doi.org/10.1111%2Fj.1467-2494.2006.00344.x">10.1111/j.1467-2494.2006.00344.x</a>. <a href="ISSN_(identifier)" class="mw-redirect" title="ISSN (identifier)">ISSN</a> <a rel="nofollow" class="external text" href="https://search.worldcat.org/issn/1468-2494">1468-2494</a>. <a href="PMID_(identifier)" class="mw-redirect" title="PMID (identifier)">PMID</a> <a rel="nofollow" class="external text" href="https://pubmed.ncbi.nlm.nih.gov/18489300">18489300</a>. <a href="S2CID_(identifier)" class="mw-redirect" title="S2CID (identifier)">S2CID</a> <a rel="nofollow" class="external text" href="https://api.semanticscholar.org/CorpusID:25191984">25191984</a>. <a rel="nofollow" class="external text" href="https://web.archive.org/web/20220321033318/https://onlinelibrary.wiley.com/doi/10.1111/j.1467-2494.2006.00344.x">Archived</a> from the original on 21 March 2022. Retrieved 8 May 2022.</cite>

</li>

<li id="cite_note-44"><a href="#cite_ref-44">^</a> <cite id="CITEREFMorriss_JrBrewerSpedaleRiddle1986" class="citation journal cs1">Morriss Jr, F. H.; Brewer, E. D.; Spedale, S. B.; Riddle, L.; Temple, D. M.; Caprioli, R. M.; West, M. S. (1986). "Relationship of human milk pH during course of lactation to concentrations of citrate and fatty acids". Pediatrics. 78 (3): 458–464. <a href="Doi_(identifier)" class="mw-redirect" title="Doi (identifier)">doi</a>:<a rel="nofollow" class="external text" href="https://doi.org/10.1542%2Fpeds.78.3.458">10.1542/peds.78.3.458</a>. <a href="PMID_(identifier)" class="mw-redirect" title="PMID (identifier)">PMID</a> <a rel="nofollow" class="external text" href="https://pubmed.ncbi.nlm.nih.gov/3748680">3748680</a>.</cite>

</li>

<li id="cite_note-45"><a href="#cite_ref-45">^</a> <cite id="CITEREFMaloney" class="citation web cs1">Maloney, Chris. <a rel="nofollow" class="external text" href="http://sinophibe.blogspot.com/2011/03/ph-calculation-of-very-small.html">"pH calculation of a very small concentration of a strong acid"</a>. <a rel="nofollow" class="external text" href="https://web.archive.org/web/20110708062942/http://sinophibe.blogspot.com/2011/03/ph-calculation-of-very-small.html">Archived</a> from the original on 8 July 2011. Retrieved 13 March 2011.</cite>

</li>

<li id="cite_note-46"><a href="#cite_ref-46">^</a> <cite id="CITEREFBillo2011" class="citation book cs1">Billo, E.J. (2011). EXCEL for Chemists (3rd ed.). Wiley-VCH. <a href="ISBN_(identifier)" class="mw-redirect" title="ISBN (identifier)">ISBN</a> <bdi>978-0-470-38123-6</bdi>.</cite>

</li>

</ol></div></div>

<div class="mw-heading mw-heading2"><h2 id="External_links">External links</h2></div>

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<div class="side-box-text plainlist"><a href="Wikidata" title="Wikidata">Wikidata</a> has the property:

<ul><li> <a href="https://www.wikidata.org/wiki/Property_talk:P9440" class="extiw external" title="d:Property talk:P9440">pH value (P9440)</a> (see <a class="external text external" href="https://query.wikidata.org/embed.html#SELECT%20%3FWikidata_item_%20%3FWikidata_item_Label%20%3Fvalue%20%3FvalueLabel%20%3FEnglish_Wikipedia_article%20%23Show%20data%20in%20this%20order%0A%7B%0A%09%3FWikidata_item_%20wdt%3AP9440%20%3Fvalue%20.%20%23Collecting%20all%20items%20which%20have%20P9440%20data%2C%20from%20whole%20Wikidata%20item%20pages%0A%09OPTIONAL%20%7B%3FEnglish_Wikipedia_article%20schema%3Aabout%20%3FWikidata_item_%3B%20schema%3AisPartOf%20%3Chttps%3A%2F%2Fen.wikipedia.org%2F%3E%20.%7D%20%23If%20collected%20item%20has%20link%20to%20English%20Wikipedia%2C%20show%20that%0A%09SERVICE%20wikibase%3Alabel%20%7B%20bd%3AserviceParam%20wikibase%3Alanguage%20%22en%22%20%20%7D%20%23Show%20label%20in%20this%20language.%20%22en%22%20is%20English.%20%20%20%0A%7D%0ALIMIT%201000">uses</a>)</li></ul></div></div>

</div>

<ul><li> The dictionary definition of <a href="https://en.wiktionary.org/wiki/%E3%8F%97" class="extiw external" title="wiktionary:㏗">pH</a> at Wiktionary</li>

<li> Media related to <a href="https://commons.wikimedia.org/wiki/Category:%E3%8F%97" class="extiw external" title="commons:Category:㏗">pH</a> at Wikimedia Commons</li>

<li> Learning materials related to <a href="https://en.wikiversity.org/wiki/Special:Search/PH" class="extiw external" title="v:Special:Search/PH">PH</a> at Wikiversity</li></ul>

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</style></div><div role="navigation" class="navbox" aria-labelledby="Wastewater267" style="padding:3px"><table class="nowraplinks hlist mw-collapsible autocollapse navbox-inner" style="border-spacing:0;background:transparent;color:inherit"><tbody><tr><th scope="col" class="navbox-title" colspan="2"><div id="Wastewater267" style="font-size:114%;margin:0 4em"><a href="Wastewater" title="Wastewater">Wastewater</a></div></th></tr><tr><th scope="row" class="navbox-group" style="width:1%">Sources and types</th><td class="navbox-list-with-group navbox-list navbox-odd" style="width:100%;padding:0"><div style="padding:0 0.25em">

<ul><li><a href="Acid_mine_drainage" title="Acid mine drainage">Acid mine drainage</a></li>

<li><a href="Ballast_water_discharge_and_the_environment" title="Ballast water discharge and the environment">Ballast water</a></li>

<li><a href="Bathroom" title="Bathroom">Bathroom</a></li>

<li><a href="Blackwater_(coal)" title="Blackwater (coal)">Blackwater (coal)</a></li>

<li><a href="Blackwater_(waste)" title="Blackwater (waste)">Blackwater (waste)</a></li>

<li><a href="Boiler_blowdown" title="Boiler blowdown">Boiler blowdown</a></li>

<li><a href="Brine" title="Brine">Brine</a></li>

<li><a href="Combined_sewer" title="Combined sewer">Combined sewer</a></li>

<li><a href="Cooling_tower#Wet_cooling_tower_material_balance" title="Cooling tower">Cooling tower</a></li>

<li><a href="Water_cooling" title="Water cooling">Cooling water</a></li>

<li><a href="Fecal_sludge_management" title="Fecal sludge management">Fecal sludge</a></li>

<li><a href="Greywater" title="Greywater">Greywater</a></li>

<li><a href="Infiltration/Inflow" class="mw-redirect" title="Infiltration/Inflow">Infiltration/Inflow</a></li>

<li><a href="Industrial_wastewater_treatment#Sources_of_industrial_wastewater" title="Industrial wastewater treatment">Industrial wastewater</a></li>

<li><a href="Ion_exchange#Regeneration_wastewater" title="Ion exchange">Ion exchange</a></li>

<li><a href="Leachate" title="Leachate">Leachate</a></li>

<li><a href="Concentrated_animal_feeding_operation" title="Concentrated animal feeding operation">Manure</a></li>

<li><a href="Environmental_effects_of_paper" class="mw-redirect" title="Environmental effects of paper">Papermaking</a></li>

<li><a href="Produced_water" title="Produced water">Produced water</a></li>

<li><a href="Return_flow" title="Return flow">Return flow</a></li>

<li><a href="Reverse_osmosis#Waste_stream_considerations" title="Reverse osmosis">Reverse osmosis</a></li>

<li><a href="Sanitary_sewer" title="Sanitary sewer">Sanitary sewer</a></li>

<li><a href="Septage" class="mw-redirect" title="Septage">Septage</a></li>

<li><a href="Sewage" title="Sewage">Sewage</a></li>

<li><a href="Sewage_sludge" title="Sewage sludge">Sewage sludge</a></li>

<li><a href="Toilet" title="Toilet">Toilet</a></li>

<li><a href="Urban_runoff" title="Urban runoff">Urban runoff</a></li></ul>

</div></td></tr><tr><th scope="row" class="navbox-group" style="width:1%"><a href="Wastewater_quality_indicators" title="Wastewater quality indicators">Quality indicators</a></th><td class="navbox-list-with-group navbox-list navbox-even" style="width:100%;padding:0"><div style="padding:0 0.25em">

<ul><li><a href="Adsorbable_organic_halides" title="Adsorbable organic halides">Adsorbable organic halides</a></li>

<li><a href="Biochemical_oxygen_demand" title="Biochemical oxygen demand">Biochemical oxygen demand</a></li>

<li><a href="Chemical_oxygen_demand" title="Chemical oxygen demand">Chemical oxygen demand</a></li>

<li><a href="Coliform_index" title="Coliform index">Coliform index</a></li>

<li><a href="Oxygen_saturation" title="Oxygen saturation">Oxygen saturation</a></li>

<li><a href="Heavy_metals" title="Heavy metals">Heavy metals</a></li>

<li><a href="Salinity" title="Salinity">Salinity</a></li>

<li><a href="Temperature" title="Temperature">Temperature</a></li>

<li><a href="Total_dissolved_solids" title="Total dissolved solids">Total dissolved solids</a></li>

<li><a href="Total_suspended_solids" title="Total suspended solids">Total suspended solids</a></li>

<li><a href="Turbidity" title="Turbidity">Turbidity</a></li>

<li><a href="Wastewater_surveillance" title="Wastewater surveillance">Wastewater surveillance</a></li></ul>

</div></td></tr><tr><th scope="row" class="navbox-group" style="width:1%"><a href="Wastewater_treatment" title="Wastewater treatment">Treatment options</a></th><td class="navbox-list-with-group navbox-list navbox-odd" style="width:100%;padding:0"><div style="padding:0 0.25em">

<ul><li><a href="Activated_sludge" title="Activated sludge">Activated sludge</a></li>

<li><a href="Aerated_lagoon" title="Aerated lagoon">Aerated lagoon</a></li>

<li><a href="Agricultural_wastewater_treatment" title="Agricultural wastewater treatment">Agricultural wastewater treatment</a></li>

<li><a href="API_oil%E2%80%93water_separator" title="API oil–water separator">API oil–water separator</a></li>

<li><a href="Carbon_filtering" title="Carbon filtering">Carbon filtering</a></li>

<li><a href="Water_chlorination" title="Water chlorination">Chlorination</a></li>

<li><a href="Clarifier" title="Clarifier">Clarifier</a></li>

<li><a href="Constructed_wetland" title="Constructed wetland">Constructed wetland</a></li>

<li><a href="Decentralized_wastewater_system" title="Decentralized wastewater system">Decentralized wastewater system</a></li>

<li><a href="Extended_aeration" title="Extended aeration">Extended aeration</a></li>

<li><a href="Facultative_lagoon" title="Facultative lagoon">Facultative lagoon</a></li>

<li><a href="Fecal_sludge_management" title="Fecal sludge management">Fecal sludge management</a></li>

<li><a href="Filtration" title="Filtration">Filtration</a></li>

<li><a href="Imhoff_tank" title="Imhoff tank">Imhoff tank</a></li>

<li><a href="Industrial_wastewater_treatment" title="Industrial wastewater treatment">Industrial wastewater treatment</a></li>

<li><a href="Ion_exchange" title="Ion exchange">Ion exchange</a></li>

<li><a href="Membrane_bioreactor" title="Membrane bioreactor">Membrane bioreactor</a></li>

<li><a href="Reverse_osmosis" title="Reverse osmosis">Reverse osmosis</a></li>

<li><a href="Rotating_biological_contactor" title="Rotating biological contactor">Rotating biological contactor</a></li>

<li><a href="Secondary_treatment" title="Secondary treatment">Secondary treatment</a></li>

<li><a href="Sedimentation_(water_treatment)" title="Sedimentation (water treatment)">Sedimentation</a></li>

<li><a href="Septic_tank" title="Septic tank">Septic tank</a></li>

<li><a href="Settling_basin" title="Settling basin">Settling basin</a></li>

<li><a href="Sewage_sludge_treatment" title="Sewage sludge treatment">Sewage sludge treatment</a></li>

<li><a href="Sewage_treatment" title="Sewage treatment">Sewage treatment</a></li>

<li><a href="Sewer_mining" title="Sewer mining">Sewer mining</a></li>

<li><a href="Stabilization_pond" class="mw-redirect" title="Stabilization pond">Stabilization pond</a></li>

<li><a href="Trickling_filter" title="Trickling filter">Trickling filter</a></li>

<li><a href="Ultraviolet_germicidal_irradiation" title="Ultraviolet germicidal irradiation">Ultraviolet germicidal irradiation</a></li>

<li><a href="Vermifilter" title="Vermifilter">Vermifilter</a></li>

<li><a href="Wastewater_treatment_plant" class="mw-redirect" title="Wastewater treatment plant">Wastewater treatment plant</a></li></ul>

</div></td></tr><tr><th scope="row" class="navbox-group" style="width:1%">Disposal options</th><td class="navbox-list-with-group navbox-list navbox-even" style="width:100%;padding:0"><div style="padding:0 0.25em">

<ul><li><a href="Combined_sewer" title="Combined sewer">Combined sewer</a></li>

<li><a href="Evaporation_pond" title="Evaporation pond">Evaporation pond</a></li>

<li><a href="Groundwater_recharge" title="Groundwater recharge">Groundwater recharge</a></li>

<li><a href="Infiltration_basin" title="Infiltration basin">Infiltration basin</a></li>

<li><a href="Injection_well" title="Injection well">Injection well</a></li>

<li><a href="Irrigation" title="Irrigation">Irrigation</a></li>

<li><a href="Marine_dumping" class="mw-redirect" title="Marine dumping">Marine dumping</a></li>

<li><a href="Marine_outfall" title="Marine outfall">Marine outfall</a></li>

<li><a href="Reclaimed_water" title="Reclaimed water">Reclaimed water</a></li>

<li><a href="Sanitary_sewer" title="Sanitary sewer">Sanitary sewer</a></li>

<li><a href="Septic_drain_field" title="Septic drain field">Septic drain field</a></li>

<li><a href="Sewage_farm" title="Sewage farm">Sewage farm</a></li>

<li><a href="Storm_drain" title="Storm drain">Storm drain</a></li>

<li><a href="Surface_runoff" title="Surface runoff">Surface runoff</a></li>

<li><a href="Vacuum_sewer" title="Vacuum sewer">Vacuum sewer</a></li></ul>

</div></td></tr><tr><td class="navbox-abovebelow" colspan="2"><div>

<ul><li> Category: Sewerage</li></ul>

</div></td></tr></tbody></table></div>

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</style></div><div role="navigation" class="navbox authority-control" aria-labelledby="Authority_control_databases_frameless&#124;text-top&#124;10px&#124;alt=Edit_this_at_Wikidata&#124;link=https&#58;//www.wikidata.org/wiki/Q40936#identifiers&#124;class=noprint&#124;Edit_this_at_Wikidata1626" style="padding:3px"><table class="nowraplinks hlist mw-collapsible autocollapse navbox-inner" style="border-spacing:0;background:transparent;color:inherit"><tbody><tr><th scope="col" class="navbox-title" colspan="2"><div id="Authority_control_databases_frameless&#124;text-top&#124;10px&#124;alt=Edit_this_at_Wikidata&#124;link=https&#58;//www.wikidata.org/wiki/Q40936#identifiers&#124;class=noprint&#124;Edit_this_at_Wikidata1626" style="font-size:114%;margin:0 4em">Authority control databases </div></th></tr><tr><th scope="row" class="navbox-group" style="width:1%">National</th><td class="navbox-list-with-group navbox-list navbox-odd" style="width:100%;padding:0"><div style="padding:0 0.25em"><ul><li><a rel="nofollow" class="external text" href="https://d-nb.info/gnd/4189273-2">Germany</a></li><li><a rel="nofollow" class="external text" href="https://id.loc.gov/authorities/sh85063427">United States</a></li><li><a rel="nofollow" class="external text" href="https://catalogue.bnf.fr/ark:/12148/cb11944522b">France</a></li><li><a rel="nofollow" class="external text" href="https://data.bnf.fr/ark:/12148/cb11944522b">BnF data</a></li><li><a rel="nofollow" class="external text" href="https://aleph.nkp.cz/F/?func=find-c&local_base=aut&ccl_term=ica=ph381261&CON_LNG=ENG">Czech Republic</a></li><li><a rel="nofollow" class="external text" href="https://www.nli.org.il/en/authorities/987007533739905171">Israel</a></li></ul></div></td></tr><tr><th scope="row" class="navbox-group" style="width:1%">Other</th><td class="navbox-list-with-group navbox-list navbox-even" style="width:100%;padding:0"><div style="padding:0 0.25em"><ul><li><a rel="nofollow" class="external text" href="https://lux.collections.yale.edu/view/concept/707e5e3f-0705-40a7-811a-2bf1d4ac40aa">Yale LUX</a></li></ul></div></td></tr></tbody></table></div></div><div><div class="zim-footer">

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