# pH > *For other uses of "pH", see the Wikipedia disambiguation page of that name.* **pH** is a logarithmic measure of how acidic or basic an aqueous [[Solution_(chemistry)|solution]] is. In the textbook form it is the negative base-10 logarithm of the [[Molar_concentration|molar concentration]] of [[Hydronium|hydronium]] ions, `pH = −log10[H3O+]`; in the definition adopted by the [[International_Union_of_Pure_and_Applied_Chemistry|IUPAC]] it is the negative logarithm of the hydrogen-ion *activity*, tied to standard buffers because the activity of a single [[Ion|ion]] cannot be measured on its own.[^iupac-goldbook] Pure [[Water|water]] at 25 °C has pH 7.00, lime juice sits near 2, coffee near 5 and household [[Ammonia|ammonia]] near 12.[^os-fig142] The number 7 is not what makes water neutral. Neutral means equal hydronium and hydroxide concentrations, and their product, the ion-product constant `Kw = [H3O+][OH−]`, rises with [[Temperature|temperature]] because the [[Self-ionization_of_water|self-ionization of water]] absorbs heat.[^os-kw] In the microsim below the reader slides the temperature of a beaker of pure water from 25 °C to 100 °C and watches the neutral mark on a pH ruler walk from 7.00 through 6.31 at 80 °C to about 6.13 at the boiling point, while the "acidic" and "basic" labels stay anchored to the comparison of [H3O⁺] with [OH⁻] rather than to the digit 7. The equation the sim answers is `pH + pOH = pKw(T)`, with the neutral point at `pH_neutral = pKw(T)/2`. On Wikitube's [[Chemistry]] flagship this page is the main article for Part VIII, *Acidity and basicity*; its siblings are [[Acid_dissociation_constant]] for weak acids, [[Buffer_solution]] for mixtures that hold their pH, and [[Titration]] for the curves traced when acid meets base. ## History The pH scale is a twentieth-century invention that arrived with electrochemical measurement. Before it, acidity was described by the colours of [[PH_indicator|indicator]] dyes and by the alkali a sample could neutralize, which ranked solutions without placing them on one scale. ### Origins and etymology The scale was introduced in 1909 by the Danish chemist S. P. L. Sørensen, working at the Carlsberg Laboratory in Copenhagen on how the hydrogen-ion concentration governs the action of [[Enzyme|enzymes]]. Concentrations in his experiments ranged over many powers of ten, and he handled them by writing the concentration as 10 raised to a negative power and quoting the exponent alone, printed as a "p" with a subscript "H".[^sorensen1909] He measured that exponent with a hydrogen electrode against a calomel reference, so from its first day the scale was an [[Electrochemistry|electrochemical]] quantity as much as a chemical one. The definition has since been rebuilt twice, first around the electromotive force of standard cells and then around the activity of the hydrogen ion, but the notation and the direction of the scale, in which larger numbers mean less acid, are Sørensen's.[^bates1960][^buck2002] Its convenience for biochemistry, where enzyme activity peaks in a narrow pH window, carried it into general use within a decade. ### Alternative origins What the "p" stands for has been argued about for a century; proposed expansions include the German *Potenz*, the French *puissance* and the Latin *pondus hydrogenii*. A 2000 examination of Sørensen's paper by Jens Nørby found no such word in it: Sørensen labelled the two solutions in his electrochemical cell *p* and *q*, and the letter attached to the hydrogen-ion exponent appears to be nothing more than that arbitrary label.[^norby2000] The reading of "p" as an operator, "take the negative logarithm of", is a later convenience that also gives pOH and p*K*a. ### First electronic measurement Electrical measurement of pH became possible when Fritz Haber and Zygmunt Klemensiewicz showed in 1909 that a thin [[Glass|glass]] membrane develops a potential that follows the hydrogen-ion concentration of the solution touching it, the principle of the glass electrode.[^haber1909] The enormous electrical resistance of the glass kept the method in specialist hands until electronic amplifiers arrived, because the cell could deliver only a minute current to a galvanometer. At the California Institute of Technology, Arnold Beckman, asked by a chemist in the citrus industry for a way to measure the acidity of lemon juice, combined a glass electrode with a vacuum-tube amplifier that boosted the signal twice; his first patent, filed in 1934, was for the amplifier, a marketable instrument existed by September 1935, and the "acidometer" was soon renamed the pH meter. The American Chemical Society designated the development of the Beckman pH meter a National Historic Chemical Landmark on March 24, 2004.[^acs-beckman] ## Definition Three quantities share the name: the IUPAC definition uses activity, p[H] is the concentration form that general chemistry teaches, and pOH is the mirror quantity for hydroxide. ### pH IUPAC defines pH as `pH = −lg a(H+)`, where a(H⁺) is the relative activity of the hydrogen ion, the product of its molality and an activity coefficient divided by the standard molality.[^iupac-goldbook] Because the activity coefficient of a single ion is not measurable, the definition is realized operationally. The primary method uses a Harned cell, `Pt | H2(g) | buffer, Cl−(aq) | AgCl | Ag`, a cell without liquid junction, and assigns pH values to a set of primary standard buffers from its electromotive force with an agreed convention for the chloride activity coefficient.[^buck2002] The 1960 report by Bates and Guggenheim fixed this approach; the 2002 IUPAC recommendations restated it with a full uncertainty budget, under which primary standards are known to a few thousandths of a unit and routine readings to a few hundredths.[^bates1960][^buck2002] The activity-based pH is a [[Dimensionless_quantity|dimensionless]] number, and one unit is a factor of ten in activity, not exactly in concentration. ### p[H] The concentration-based quantity `p[H] = −log10[H+]` is what Sørensen defined and what introductory texts call pH, writing the hydrated [[Proton|proton]] as hydronium: `pH = −log[H3O+]` and `pOH = −log[OH−]`.[^os-kw] In dilute solution the activity coefficient is close to one, so pH and p[H] differ by a few hundredths at most; in concentrated brines or strong acids they diverge. Every microsim on this portal works in p[H], as the Portal Books' examples do. ### pOH The counterpart quantity `pOH = −log10[OH−]` is linked to pH through the ion product of water. At every temperature `Kw = [H3O+][OH−]`, so taking logarithms gives `pH + pOH = pKw`, and neutral water, with [H3O⁺] = [OH⁻], sits at `pH = pKw/2`.[^os-kw] At 25 °C, Kw is 1.0×10⁻¹⁴, pKw is 14.00 and neutral pH is 7.00. The self-ionization is endothermic, so by [[Le_Chatelier's_principle|Le Chatelier's principle]] Kw grows with temperature and the neutral point drifts downward.[^os-kw] The Portal Book gives five values, collected here from its examples and exercises, with neutral pH computed as pKw/2:[^os-kw][^os-kw-T] | T (°C) | Kw | pKw | neutral pH | [H3O⁺] = [OH⁻] (mol/L) | |---|---|---|---|---| | 25 | 1.0×10⁻¹⁴ | 14.00 | 7.00 | 1.0×10⁻⁷ | | 40 | 2.9×10⁻¹⁴ | 13.54 | 6.77 | 1.7×10⁻⁷ | | 60 | 9.3×10⁻¹⁴ | 13.03 | 6.52 | 3.0×10⁻⁷ | | 80 | 2.4×10⁻¹³ | 12.62 | 6.31 | 4.9×10⁻⁷ | | 100 | 5.6×10⁻¹³ | 12.25 | 6.13 | 7.5×10⁻⁷ | This table is the whole of the article's main microsim. A pH ruler runs from 0 to 14 above a beaker of pure water; the single control is the water's temperature, 25–100 °C in half-degree steps, and the HUD carries `pH_neutral = pKw(T)/2`. As the slider moves, a marker labelled "neutral" slides down the ruler, 7.00 at 25 °C, 6.31 at 80 °C, about 6.13 at 100 °C, and the readouts of pH, pOH and pKw update together. The "acidic" and "basic" bands are drawn from the comparison of [H3O⁺] with [OH⁻], so at 80 °C a solution at pH 6.5 is shown as basic, which it is. Between the five book temperatures the sim interpolates log10 Kw linearly in 1/T, the straight line a [[Van_'t_Hoff_equation|van 't Hoff plot]] suggests; that interpolation is an authoring choice, labelled ILLUSTRATIVE on the canvas, and the sim does not extrapolate outside 25–100 °C. The lesson is one sentence: boiling-hot pure water has a pH near 6.1 and is exactly neutral. ## Measurement Routine pH is measured potentiometrically with a glass electrode paired with a reference electrode, the pair reading a [[Voltage|voltage]] that changes by about 59 mV per pH unit at 25 °C, the slope of the [[Nernst_equation|Nernst equation]].[^buck2002] The meter is calibrated against two or more standard buffers bracketing the expected value, because the electrode's slope and offset drift with age and temperature. Older methods use indicator dyes in solution or on paper. ### pH Indicators An indicator is itself a weak [[Acid|acid]], HIn, whose acid and conjugate-base forms have different colours. Its colour ratio follows its own dissociation constant, `pH = pKa(In) + log10([In−]/[HIn])`, so the eye sees a change over roughly one unit either side of pKa(In).[^os-indicators] Phenolphthalein is colourless below about pH 8.3 and pink above; methyl orange changes between red-orange near pH 4.2 and yellow near 6.3, which makes it unsuitable for a weak-acid titration whose steep step lies above 7.[^os-indicators] The [[Titration]] microsim colours its flask by exactly this ratio. ### Non-aqueous solutions The IUPAC definition and its standard buffers are for water.[^buck2002] In another solvent the proton's solvation energy, the solvent's own self-ionization constant and the electrode response all differ, so a number read with an aqueous-calibrated electrode has no fixed relation to the aqueous scale. Solvent-specific scales exist for methanol, acetonitrile and dimethyl sulfoxide, and an acid that is strong in water can be weak in a solvent that solvates the proton poorly. ### Unified absolute pH scale In 2010 Himmel, Goll, Leito and Krossing proposed a unified scale, pH_abs, defined through the absolute chemical potential of the proton in whatever phase it occupies, referenced to the proton in the ideal gas.[^himmel2010] On that axis aqueous pH values become one segment, and acidities in acetonitrile, ionic liquids, solids and gases can be placed beside them, with the aqueous zero recovered by a constant shift.[^himmel2010] It is a thermodynamic framework rather than a laboratory method, and it needs the proton's transfer energy between each solvent and the reference state. ### Extremes of pH measurements The scale is not bounded by 0 and 14. Those limits correspond to 1 mol/L of strong acid or base; more concentrated solutions push beyond them, and there activity and concentration part company. The most extreme natural values reported are from the Richmond Mine at Iron Mountain, California, where Nordstrom and colleagues measured acid mine waters with pH as low as −3.6, readings obtained only after the glass electrode was calibrated against sulfuric-acid standards of known activity.[^nordstrom2000] At the alkaline end, concentrated [[Sodium_hydroxide|sodium hydroxide]] exceeds pH 14 and attacks the glass membrane itself. ## Applications Because so many equilibria, from [[Solubility_equilibrium|solubility]] to enzyme activity and [[Standard_electrode_potential|electrode potentials]], involve the proton, pH is the most reported number in chemical, biological and environmental analysis. ### pH in soil Soil pH governs the charge on soil colloids and therefore what the soil can hold. The negative charge of humus and of iron and [[Aluminium|aluminium]] oxides depends on pH, and base saturation, the fraction of a soil's exchange capacity occupied by calcium, magnesium, potassium and sodium, rises roughly linearly between pH 4 and pH 7.[^soils118] Sodic soils, dispersed and slow to drain, typically sit above pH 8.5, while saline soils usually sit below it.[^soils118] In waterlogged soils the [[Redox|redox]] potential Eh and pH move together, Eh falling 59.16 mV per pH unit at 25 °C for couples that transfer equal numbers of protons and electrons.[^soils118] ### pH in plants Plants read soil pH mostly through what it makes available. A visible case is the bigleaf hydrangea: in acidic soil, pH 5.0 to 5.5, aluminium is soluble, the flowers take it up and bloom blue, while at pH 6.0 and above aluminium is locked away and the same plant blooms pink, so growers steer the colour with [[Sulfur|sulfur]] or aluminium sulfate on one side and lime on the other.[^clemson-hydrangea] ### pH in the ocean Surface seawater is mildly basic, near pH 8.2 before industrialization, because it is buffered by the carbonate system. Carbon dioxide dissolving from the [[Atmosphere_of_Earth|atmosphere]] forms carbonic acid and shifts that buffer; the surface ocean has already lost about 0.1 pH unit since the pre-industrial period, roughly a 30 percent rise in hydrogen-ion concentration.[^doney2009] Caldeira and Wickett's 2003 model calculation showed that continued fossil-fuel emissions could drive a further fall of several tenths of a unit within this century.[^caldeira2003] Carbonate-ion concentration falls with the pH, which is what threatens organisms that build shells of calcium carbonate.[^doney2009] ### pH in food Food-safety regulation uses pH as a bright line. Under United States federal regulation, "acid foods" have a natural pH of 4.6 or below, and "acidified foods" are low-acid foods to which acid has been added to reach a finished equilibrium pH of 4.6 or below with a water activity above 0.85.[^ecfr114] Many foods are naturally well inside the acid range: lime juice near pH 2, wine near 3.5, coffee near 5.[^os-fig142] ### pH of various body fluids Human arterial blood is held between pH 7.35 and 7.45; below 7.35 a person is in acidosis and above 7.45 in alkalosis.[^openstax-ap] At the normal 7.40 the bicarbonate-to-carbonic-acid ratio is 20:1, the [[Bicarbonate_buffer_system|bicarbonate buffer]] treated in the [[Buffer_solution]] article.[^openstax-ap] Other fluids are far from neutral: gastric juice sits near pH 1.5, stomach contents near 3, urine and saliva near 6.[^os-fig142][^openstax-ap] ## pH calculations The methods form a ladder: strong electrolytes need one logarithm, weak ones need an [[Equilibrium_constant|equilibrium constant]], and mixtures need every equilibrium solved at once. ### Strong acids and bases A strong acid such as HCl is fully ionized, so [H3O⁺] equals the acid's formal concentration and `pH = −log10 C`. A strong base gives [OH⁻] directly, pOH follows, and pH is 14.00 − pOH at 25 °C: 0.0125 mol/L of potassium hydroxide gives pOH 1.903 and pH 12.097.[^os-strong] The formula runs past the "ends" of the scale; hydrochloric acid at pH −1.07 is about 12 mol/L.[^os-strong] Below about 10⁻⁶ mol/L the water's own 10⁻⁷ mol/L of hydronium can no longer be ignored and the general method is needed. ### Weak acids and bases A weak acid ionizes only partly, and its pH follows from the [[Acid_dissociation_constant|acid dissociation constant]]. With x the hydronium concentration produced from an initial concentration C₀, `Ka = x²/(C0 − x)`; when x is small against C₀ the shortcut `x ≈ sqrt(Ka·C0)` serves, otherwise the quadratic is solved and its negative root discarded.[^os-weak] For 0.10 mol/L acetic acid, with Ka = 1.8×10⁻⁵, the shortcut gives pH 2.87 and the exact root 2.88; the book's measured pH of 2.89 corresponds to 1.3 percent ionization, so at this concentration the shortcut is safe.[^os-weak] Percent ionization is not a constant of the acid but rises on dilution, which is the point of the sibling article's microsim, where a slider on C₀ shows the shortcut's 5 percent test failing below about 7×10⁻³ mol/L. Weak bases are handled the same way with Kb, and `Ka·Kb = Kw` converts between an acid and its [[Conjugate_(acid-base_theory)|conjugate]] base, so the [[Acid_strength|strength]] of one fixes the weakness of the other.[^os-weak] ### General method For any mixture the exact pH comes from writing every relation the solution must obey: a mass balance for each acid–base pair, the equilibrium expression for each pair, the ion product of water, and the charge balance, which states that positive and negative charges in solution sum to zero. Eliminating the species concentrations leaves one equation in [H⁺], a polynomial whose degree grows with the number of equilibria. For sodium hydroxide added to a weak acid HA it reads `[H+] + [Na+] = [OH−] + CA·Ka/(Ka + [H+])`, with C_A the total acid concentration after dilution. Its left side rises and its right side falls as [H⁺] increases, so it has exactly one root, which the Titration and Buffer solution microsims find by bisection in log10[H⁺]. The reward is a [[Titration_curve|curve]] that stays finite at the [[Equivalence_point|equivalence point]], where the shortcuts divide by zero; the check is that the same charge balance reproduces the Portal Book's titration table to ±0.01 pH at every listed volume.[^os-titration] ## See also - [[Self-ionization_of_water]] - [[Acid–base_reaction]] - [[Brønsted–Lowry_acid–base_theory]] - [[Acid]] - [[Base_(chemistry)]] - [[Lewis_acids_and_bases]] - [[Hydronium]] - [[Acid_dissociation_constant]] - [[Buffer_solution]] - [[Titration]] - [[PH_indicator]] ## References [^iupac-goldbook]: IUPAC. "pH." *Compendium of Chemical Terminology* (the "Gold Book"), entry P04524. https://goldbook.iupac.org/terms/view/P04524 [^os-fig142]: Flowers, Paul; Neth, Edward; Robinson, William; et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 14 "Acid-Base Equilibria", §14.2 "pH and pOH", Figure 14.2, pH of common substances (page to pin). https://openstax.org/books/chemistry-atoms-first-2e/pages/14-2-ph-and-poh [^os-kw]: Flowers, Paul; Neth, Edward; Robinson, William; et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 14, §14.1–14.2, pp. 667–670: the ion product of water (1.0×10⁻¹⁴ at 25 °C, 2.4×10⁻¹³ at 80 °C), the definitions of pH and pOH, and pure water at 80 °C (4.9×10⁻⁷ mol/L, pH = pOH = 6.31). https://openstax.org/books/chemistry-atoms-first-2e/pages/14-2-ph-and-poh [^os-kw-T]: Flowers, Paul; Neth, Edward; Robinson, William; et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 14 end-of-chapter exercises, pp. 712–713: Kw = 2.9×10⁻¹⁴ at 40 °C and 9.3×10⁻¹⁴ at 60 °C; the 100 °C value of about 5.6×10⁻¹³ is given at p. 667. [^sorensen1909]: Sørensen, S. P. L. (1909). "Enzymstudien. II. Mitteilung. Über die Messung und die Bedeutung der Wasserstoffionenkonzentration bei enzymatischen Prozessen." *Biochemische Zeitschrift* 21: 131–304. [^norby2000]: Nørby, Jens G. (2000). "The origin and the meaning of the little p in pH." *Trends in Biochemical Sciences* 25 (1): 36–37. https://doi.org/10.1016/S0968-0004(99)01517-0 [^haber1909]: Haber, F.; Klemensiewicz, Z. (1909). "Über elektrische Phasengrenzkräfte." *Zeitschrift für Physikalische Chemie* 67U: 385–431. https://doi.org/10.1515/zpch-1909-6720 [^acs-beckman]: American Chemical Society. "Development of the Beckman pH Meter." National Historic Chemical Landmarks; designated at the Beckman Institute, California Institute of Technology, March 24, 2004. https://www.acs.org/education/whatischemistry/landmarks/beckman.html [^bates1960]: Bates, R. G.; Guggenheim, E. A. (1960). "Report on the standardization of pH and related terminology." *Pure and Applied Chemistry* 1 (1): 163–168. https://doi.org/10.1351/pac196001010163 [^buck2002]: Buck, R. P.; Rondinini, S.; Covington, A. K.; Baucke, F. G. K.; Brett, C. M. A.; Camões, M. F.; Milton, M. J. T.; Mussini, T.; Naumann, R.; Pratt, K. W.; Spitzer, P.; Wilson, G. S. (2002). "Measurement of pH. Definition, standards, and procedures (IUPAC Recommendations 2002)." *Pure and Applied Chemistry* 74 (11): 2169–2200. https://doi.org/10.1351/pac200274112169 [^himmel2010]: Himmel, Daniel; Goll, Sascha K.; Leito, Ivo; Krossing, Ingo (2010). "A Unified pH Scale for All Phases." *Angewandte Chemie International Edition* 49 (38): 6885–6888. https://doi.org/10.1002/anie.201000252 [^nordstrom2000]: Nordstrom, D. Kirk; Alpers, Charles N.; Ptacek, Carol J.; Blowes, David W. (2000). "Negative pH and Extremely Acidic Mine Waters from Iron Mountain, California." *Environmental Science & Technology* 34 (2): 254–258. https://doi.org/10.1021/es990646v [^soils118]: Canadian Society of Soil Science (2021). *Digging into Canadian Soils: An Introduction to Soil Science*. Soil chemistry: pH-dependent charge of humus and oxides, p. 185; base saturation against pH, p. 188; saline, sodic and saline-sodic classes, pp. 196–197; the Eh–pH slope of 59.16 mV per unit at 25 °C, p. 203. [^clemson-hydrangea]: Kluepfel, Marjan; Polomski, Robert F.; Williamson, Joey. "Hydrangea Care in South Carolina." Clemson Cooperative Extension, Home & Garden Information Center, factsheet HGIC 1067 (updated October 31, 2025). https://hgic.clemson.edu/factsheet/hydrangea/ [^doney2009]: Doney, Scott C.; Fabry, Victoria J.; Feely, Richard A.; Kleypas, Joan A. (2009). "Ocean Acidification: The Other CO₂ Problem." *Annual Review of Marine Science* 1: 169–192. https://doi.org/10.1146/annurev.marine.010908.163834 [^caldeira2003]: Caldeira, Ken; Wickett, Michael E. (2003). "Anthropogenic carbon and ocean pH." *Nature* 425: 365. https://doi.org/10.1038/425365a [^ecfr114]: Code of Federal Regulations, Title 21, Chapter I, Subchapter B, Part 114 "Acidified Foods", § 114.3 Definitions. https://www.ecfr.gov/current/title-21/chapter-I/subchapter-B/part-114 [^openstax-ap]: OpenStax (2022). *Anatomy and Physiology 2e*. §26.5 "Disorders of Acid-Base Balance": arterial blood pH 7.35–7.45; acidosis and alkalosis; the 20:1 bicarbonate to carbonic acid ratio at pH 7.40. https://openstax.org/books/anatomy-and-physiology-2e/pages/26-5-disorders-of-acid-base-balance [^os-indicators]: Flowers, Paul; Neth, Edward; Robinson, William; et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 14, §14.7 "Acid-Base Titrations", pp. 706–708: indicator equilibrium and colour-change interval, phenolphthalein and methyl orange ranges, choice of indicator for a weak-acid titration. https://openstax.org/books/chemistry-atoms-first-2e/pages/14-7-acid-base-titrations [^os-strong]: Flowers, Paul; Neth, Edward; Robinson, William; et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 14, §14.2, pp. 668–674: Examples 14.2–14.6 and their check problems (0.0125 mol/L KOH → pOH 1.903, pH 12.097; pH −1.07 ↔ 12 mol/L HCl). The printed results of these examples were lost in extraction and were recomputed from the book's inputs (sub-manual 05, App. A). https://openstax.org/books/chemistry-atoms-first-2e/pages/14-2-ph-and-poh [^os-weak]: Flowers, Paul; Neth, Edward; Robinson, William; et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 14, §14.3 "Relative Strengths of Acids and Bases", pp. 675–686: Ka, percent ionization (0.10 mol/L acetic acid at pH 2.89 → 1.3 %), the "x is small" shortcut and its 5 % test, Ka·Kb = Kw; Ka(CH₃CO₂H) = 1.80×10⁻⁵ at p. 716. https://openstax.org/books/chemistry-atoms-first-2e/pages/14-3-relative-strengths-of-acids-and-bases [^os-titration]: Flowers, Paul; Neth, Edward; Robinson, William; et al. (2019). *Chemistry: Atoms First 2e*. OpenStax. Chapter 14, §14.7, pp. 702–705: Table 14.2, pH during the titration of 25.00 mL of 0.100 mol/L HCl and of acetic acid with 0.100 mol/L NaOH. https://openstax.org/books/chemistry-atoms-first-2e/pages/14-7-acid-base-titrations ## External links - [pH — IUPAC Gold Book entry](https://goldbook.iupac.org/terms/view/P04524) - [Chemistry: Atoms First 2e, Chapter 14 — OpenStax](https://openstax.org/books/chemistry-atoms-first-2e/pages/14-2-ph-and-poh), the Portal Book behind this page's numbers - [Development of the Beckman pH Meter — ACS National Historic Chemical Landmark](https://www.acs.org/education/whatischemistry/landmarks/beckman.html) <!-- MATTERSIM:BEGIN g24 — Matter & Energy Cluster microsim (framework build, specs/sims/PH.json); do not hand-edit inside --> **Microsim — three.js (Wikitube framework), pending deploy:** *pH* will play here once `https://wikitube-3d-microsims.netlify.app/matter/PH.html` is live. <!-- pending: <div class="wt-sim" data-src="https://wikitube-3d-microsims.netlify.app/matter/PH.html" data-title="pH"></div> --> <!-- MATTERSIM:END --> ## Wikipedia : Wikitube **Strict pair:** [Wikipedia](https://en.wikipedia.org/wiki/PH) : [Wikitube](https://en.wikitube.io/wiki/PH) · pinned revision [1370268150](https://en.wikipedia.org/w/index.php?oldid=1370268150) · 2026-09-11 ## Previous hub tags Hubs: `Life_Physics`. Portals: [[PORTAL_Chemistry]]. --- *Matter & Energy Cluster child articles, wave 1 · 2026-09-11 · drafted · Chemistry row K39 · sim pending (matter/PH).*