The first category of acids are the proton donors, orBrønsted–Lowry acids. In the special case ofaqueous solutions, proton donors form thehydronium ion H3O+ and are known asArrhenius acids.Brønsted andLowry generalized the Arrhenius theory to include non-aqueoussolvents. A Brønsted–Lowry or Arrhenius acid usually contains a hydrogen atom bonded to a chemical structure that is still energetically favorable after loss of H+.
Aqueous Arrhenius acids have characteristic properties that provide a practical description of an acid.[2] Acids form aqueous solutions with a sour taste, can turn bluelitmus red, and react withbases and certain metals (likecalcium) to formsalts. The wordacid is derived from theLatinacidus, meaning 'sour'.[3] An aqueous solution of an acid has apH less than 7 and is colloquially also referred to as "acid" (as in "dissolved in acid"), while the strict definition refers only to thesolute.[1] A lower pH means a higheracidity, and thus a higher concentration of hydrogen cations in the solution. Chemicals or substances having the property of an acid are said to beacidic.
The second category of acids areLewis acids, which form a covalent bond with an electron pair. An example isboron trifluoride (BF3), whose boron atom has a vacantorbital that can form a covalent bond by sharing a lone pair of electrons on an atom in a base, for example the nitrogen atom inammonia (NH3).Lewis considered this as a generalization of the Brønsted definition, so that an acid is a chemical species that accepts electron pairs either directlyor by releasing protons (H+) into the solution, which then accept electron pairs. Hydrogen chloride, acetic acid, and most other Brønsted–Lowry acids cannot form a covalent bond with an electron pair, however, and are therefore not Lewis acids.[4] Conversely, many Lewis acids are not Arrhenius or Brønsted–Lowry acids. In modern terminology, anacid is implicitly a Brønsted acid and not a Lewis acid, since chemists almost always refer to a Lewis acid explicitly as such.[4]
Modern definitions are concerned with the fundamental chemical reactions common to all acids.
Most acids encountered in everyday life areaqueous solutions, or can be dissolved in water, so the Arrhenius and Brønsted–Lowry definitions are the most relevant.
The Brønsted–Lowry definition is the most widely used definition; unless otherwise specified, acid–base reactions are assumed to involve the transfer of a proton (H+) from an acid to a base.
Hydronium ions are acids according to all three definitions. Although alcohols and amines can be Brønsted–Lowry acids, they can also function asLewis bases due to the lone pairs of electrons on their oxygen and nitrogen atoms.
Arrhenius acids
Svante Arrhenius
In 1884,Svante Arrhenius attributed the properties of acidity to hydrogen cations (H+), later described asprotons orhydrons. AnArrhenius acid is a substance that, when added to water, increases the concentration of H+ ions in the water.[4][5] Chemists often write H+(aq) and refer to the hydrogen cation when describing acid–base reactions but the free hydrogen nucleus, aproton, does not exist alone in water, it exists as thehydronium ion (H3O+) or other forms (H5O2+, H9O4+). Thus, an Arrhenius acid can also be described as a substance that increases the concentration ofhydronium ions when added to water. Examples include molecular substances such as hydrogen chloride and acetic acid.
An Arrheniusbase, on the other hand, is a substance that increases the concentration ofhydroxide (OH−) ions when dissolved in water. This decreases the concentration of hydronium because the ions react to form H2O molecules:
H3O+ (aq) + OH− (aq) ⇌ H2O(liq) + H2O(liq)
Due to this equilibrium, any increase in the concentration of hydronium is accompanied by a decrease in the concentration of hydroxide. Thus, an Arrhenius acid could also be said to be one that decreases hydroxide concentration, while an Arrhenius base increases it.
In an acidic solution, the concentration of hydronium ions is greater than 10−7moles per liter. Since pH is defined as the negative logarithm of the concentration of hydronium ions, acidic solutions thus have a pH of less than 7.
Acetic acid, aweak acid, donates a proton (hydrogen ion, highlighted in green) to water in an equilibrium reaction to give theacetate ion and thehydronium ion. Red: oxygen, black: carbon, white: hydrogen.
While the Arrhenius concept is useful for describing many reactions, it is also quite limited in its scope. In 1923, chemistsJohannes Nicolaus Brønsted andThomas Martin Lowry independently recognized that acid–base reactions involve the transfer of a proton. ABrønsted–Lowry acid (or simply Brønsted acid) is a species that donates a proton to a Brønsted–Lowry base.[5] Brønsted–Lowry acid–base theory has several advantages over Arrhenius theory. Consider the following reactions ofacetic acid (CH3COOH), theorganic acid that gives vinegar its characteristic taste:
CH3COOH + H2O ⇌ CH3COO− + H3O+
CH3COOH + NH3 ⇌ CH3COO− + NH+4
Both theories easily describe the first reaction: CH3COOH acts as an Arrhenius acid because it acts as a source of H3O+ when dissolved in water, and it acts as a Brønsted acid by donating a proton to water. In the second example CH3COOH undergoes the same transformation, in this case donating a proton to ammonia (NH3), but does not relate to the Arrhenius definition of an acid because the reaction does not produce hydronium. Nevertheless, CH3COOH is both an Arrhenius and a Brønsted–Lowry acid.
Brønsted–Lowry theory can be used to describe reactions ofmolecular compounds in nonaqueous solution or the gas phase.Hydrogen chloride (HCl) and ammonia combine under several different conditions to formammonium chloride, NH4Cl. In aqueous solution HCl behaves ashydrochloric acid and exists as hydronium and chloride ions. The following reactions illustrate the limitations of Arrhenius's definition:
As with the acetic acid reactions, both definitions work for the first example, where water is the solvent and hydronium ion is formed by the HCl solute. The next two reactions do not involve the formation of ions but are still proton-transfer reactions. In the second reaction hydrogen chloride and ammonia (dissolved inbenzene) react to form solid ammonium chloride in a benzene solvent and in the third gaseous HCl and NH3 combine to form the solid.
A third, only marginally related concept was proposed in 1923 byGilbert N. Lewis, which includes reactions with acid–base characteristics that do not involve a proton transfer. ALewis acid is a species that accepts a pair of electrons from another species; in other words, it is an electron pair acceptor.[5] Brønsted acid–base reactions are proton transfer reactions while Lewis acid–base reactions are electron pair transfers. Many Lewis acids are not Brønsted–Lowry acids. Contrast how the following reactions are described in terms of acid–base chemistry:
In the first reaction afluoride ion, F−, gives up anelectron pair toboron trifluoride to form the producttetrafluoroborate. Fluoride "loses" a pair ofvalence electrons because the electrons shared in the B—F bond are located in the region of space between the two atomicnuclei and are therefore more distant from the fluoride nucleus than they are in the lone fluoride ion. BF3 is a Lewis acid because it accepts the electron pair from fluoride. This reaction cannot be described in terms of Brønsted theory because there is no proton transfer.
The second reaction can be described using either theory. A proton is transferred from an unspecified Brønsted acid to ammonia, a Brønsted base; alternatively, ammonia acts as a Lewis base and transfers a lone pair of electrons to form a bond with a hydrogen ion. The species that gains the electron pair is the Lewis acid; for example, the oxygen atom in H3O+ gains a pair of electrons when one of the H—O bonds is broken and the electrons shared in the bond become localized on oxygen.
Depending on the context, a Lewis acid may also be described as anoxidizer or anelectrophile. Organic Brønsted acids, such as acetic, citric, or oxalic acid, are not Lewis acids.[4] They dissociate in water to produce a Lewis acid, H+, but at the same time, they also yield an equal amount of a Lewis base (acetate, citrate, or oxalate, respectively, for the acids mentioned). This article deals mostly with Brønsted acids rather than Lewis acids.
Dissociation and equilibrium
Reactions of acids are often generalized in the formHA ⇌ H+ + A−, where HA represents the acid and A− is theconjugate base. This reaction is referred to asprotolysis. The protonated form (HA) of an acid is also sometimes referred to as thefree acid.[6]
Acid–base conjugate pairs differ by one proton, and can be interconverted by the addition or removal of a proton (protonation anddeprotonation, respectively). The acid can be the charged species and the conjugate base can be neutral in which case the generalized reaction scheme could be written asHA+ ⇌ H+ + A. In solution there exists anequilibrium between the acid and its conjugate base. Theequilibrium constantK is an expression of the equilibrium concentrations of the molecules or the ions in solution. Brackets indicate concentration, such that [H2O] meansthe concentration of H2O. Theacid dissociation constantKa is generally used in the context of acid–base reactions. The numerical value ofKa is equal to theproduct (multiplication) of the concentrations of the products divided by the concentration of the reactants, where the reactant is the acid (HA) and the products are the conjugate base and H+.
The stronger of two acids will have a higherKa than the weaker acid; the ratio of hydrogen cations to acid will be higher for the stronger acid as the stronger acid has a greater tendency to lose its proton. Because the range of possible values forKa spans many orders of magnitude, a more manageable constant, pKa is more frequently used, where pKa = −log10Ka. Stronger acids have a smaller pKa than weaker acids. Experimentally determined pKa at 25 °C in aqueous solution are often quoted in textbooks and reference material.
Nomenclature
Arrhenius acids are named according to theiranions. In the classical naming system, the ionic suffix is dropped and replaced with a new suffix, according to the table following. The prefix "hydro-" is used when the acid is made up of just hydrogen and one other element. For example, HCl haschloride as its anion, so the hydro- prefix is used, and the -ide suffix makes the name take the formhydrochloric acid.
In theIUPAC naming system, "aqueous" is simply added to the name of the ionic compound. Thus, for hydrogen chloride, as an acid solution, the IUPAC name is aqueous hydrogen chloride.
The strength of an acid refers to its ability or tendency to lose a proton. A strong acid is one that completely dissociates in water; in other words, onemole of a strong acid HA dissolves in water yielding one mole of H+ and one mole of the conjugate base, A−, and none of the protonated acid HA. In contrast, a weak acid only partially dissociates and at equilibrium both the acid and the conjugate base are in solution. Examples ofstrong acids arehydrochloric acid (HCl),hydroiodic acid (HI),hydrobromic acid (HBr),perchloric acid (HClO4),nitric acid (HNO3) andsulfuric acid (H2SO4). In water, each of these essentially ionizes 100%. The stronger an acid is, the more easily it loses a proton, H+. Two key factors that contribute to the ease of deprotonation are thepolarity of the H—A bond and the size of atom A, which determines the strength of the H—A bond. Acid strengths are also often discussed in terms of the stability of the conjugate base.
Sulfonic acids, which are organic oxyacids, are a class of strong acids. A common example istoluenesulfonic acid (tosylic acid). Unlike sulfuric acid itself, sulfonic acids can be solids. In fact,polystyrene functionalized into polystyrene sulfonate is a solid strongly acidic plastic that is filterable.
WhileKa measures the strength of an acid compound, the strength of an aqueous acid solution is measured by pH, which is an indication of the concentration of hydronium in the solution. The pH of a simple solution of an acid compound in water is determined by the dilution of the compound and the compound'sKa.
Lewis acid strength in non-aqueous solutions
Lewis acids have been classified in theECW model and it has been shown that there is no one order of acid strengths.[9] The relative acceptor strength of Lewis acids toward a series of bases, versus other Lewis acids, can be illustrated byC-B plots.[10][11] It has been shown that to define the order of Lewis acid strength at least two properties must be considered. For Pearson's qualitativeHSAB theory the two properties arehardness and strength while for Drago's quantitativeECW model the two properties are electrostatic and covalent.
Monoprotic acids, also known as monobasic acids, are those acids that are able to donate oneproton per molecule during the process ofdissociation (sometimes called ionization) as shown below (symbolized by HA):
Polyprotic acids, also known as polybasic acids, are able to donate more than one proton per acid molecule, in contrast to monoprotic acids that only donate one proton per molecule. Specific types of polyprotic acids have more specific names, such as diprotic (or dibasic) acid (two potential protons to donate), and triprotic (or tribasic) acid (three potential protons to donate). Some macromolecules such as proteins and nucleic acids can have a very large number of acidic protons.[12]
A diprotic acid (here symbolized by H2A) can undergo one or two dissociations depending on the pH. Each dissociation has its own dissociation constant, Ka1 and Ka2.
H2A (aq) + H2O (l) ⇌ H3O+ (aq) + HA− (aq)Ka1
HA− (aq) + H2O (l) ⇌ H3O+ (aq) + A2− (aq)Ka2
The first dissociation constant is typically greater than the second (i.e.,Ka1 >Ka2). For example,sulfuric acid (H2SO4) can donate one proton to form thebisulfate anion (HSO− 4), for whichKa1 is very large; then it can donate a second proton to form thesulfate anion (SO2− 4), wherein theKa2 is intermediate strength. The largeKa1 for the first dissociation makes sulfuric a strong acid. In a similar manner, the weak unstablecarbonic acid(H2CO3) can lose one proton to formbicarbonate anion(HCO− 3) and lose a second to formcarbonate anion (CO2− 3). BothKa values are small, butKa1 >Ka2 .
A triprotic acid (H3A) can undergo one, two, or three dissociations and has three dissociation constants, whereKa1 >Ka2 >Ka3.
H3A (aq) + H2O (l) ⇌ H3O+ (aq) + H2A− (aq)Ka1
H2A− (aq) + H2O (l) ⇌ H3O+ (aq) + HA2− (aq)Ka2
HA2− (aq) + H2O (l) ⇌ H3O+ (aq) + A3− (aq)Ka3
Aninorganic example of a triprotic acid is orthophosphoric acid (H3PO4), usually just calledphosphoric acid. All three protons can be successively lost to yield H2PO− 4, then HPO2− 4, and finally PO3− 4, the orthophosphate ion, usually just calledphosphate. Even though the positions of the three protons on the original phosphoric acid molecule are equivalent, the successiveKa values differ since it is energetically less favorable to lose a proton if the conjugate base is more negatively charged. Anorganic example of a triprotic acid iscitric acid, which can successively lose three protons to finally form thecitrate ion.
Although the subsequent loss of each hydrogen ion is less favorable, all of the conjugate bases are present in solution. The fractional concentration,α (alpha), for each species can be calculated. For example, a generic diprotic acid will generate 3 species in solution: H2A, HA−, and A2−. The fractional concentrations can be calculated as below when given either the pH (which can be converted to the [H+]) or the concentrations of the acid with all its conjugate bases:
A plot of these fractional concentrations against pH, for givenK1 andK2, is known as aBjerrum plot. A pattern is observed in the above equations and can be expanded to the generaln -protic acid that has been deprotonatedi -times:
whereK0 = 1 and the other K-terms are the dissociation constants for the acid.
Neutralization is the basis oftitration, where apH indicator shows equivalence point when the equivalent number of moles of a base have been added to an acid. It is often wrongly assumed that neutralization should result in a solution with pH 7.0, which is only the case with similar acid and base strengths during a reaction.
Neutralization with a base weaker than the acid results in a weakly acidic salt. An example is the weakly acidicammonium chloride, which is produced from the strong acidhydrogen chloride and the weak baseammonia. Conversely, neutralizing a weak acid with a strong base gives a weakly basic salt (e.g.,sodium fluoride fromhydrogen fluoride andsodium hydroxide).
In order for a protonated acid to lose a proton, the pH of the system must rise above the pKa of the acid. The decreased concentration of H+ in that basic solution shifts the equilibrium towards the conjugate base form (the deprotonated form of the acid). In lower-pH (more acidic) solutions, there is a high enough H+ concentration in the solution to cause the acid to remain in its protonated form.
Solutions of weak acids and salts of their conjugate bases formbuffer solutions.
To determine the concentration of an acid in an aqueous solution, an acid–base titration is commonly performed. A strong base solution with a known concentration, usually NaOH or KOH, is added to neutralize the acid solution according to the color change of the indicator with the amount of base added.[13] The titration curve of an acid titrated by a base has two axes, with the base volume on the x-axis and the solution's pH value on the y-axis. The pH of the solution always goes up as the base is added to the solution.
Example: Diprotic acid
This is an ideal titration curve foralanine, a diprotic amino acid.[14] Point 2 is the first equivalent point where the amount of NaOH added equals the amount of alanine in the original solution.
For each diprotic acid titration curve, from left to right, there are two midpoints, two equivalence points, and two buffer regions.[15]
Equivalence points
Due to the successive dissociation processes, there are two equivalence points in the titration curve of a diprotic acid.[16] The first equivalence point occurs when all first protons from the first ionization are titrated.[17] In other words, the amount of OH− added equals the original amount of H2A at the first equivalence point. The second equivalence point occurs when all protons are titrated. Therefore, the amount of OH− added equals twice the amount of H2A at this time. For a weak diprotic acid titrated by a strong base, the second equivalence point must occur at pH above 7 due to the hydrolysis of the resulted salts in the solution.[17] At either equivalence point, adding a drop of base will cause the steepest rise of the pH value in the system.
Buffer regions and midpoints
A titration curve for a diprotic acid contains two midpoints where pH=pKa. Since there are two different Ka values, the first midpoint occurs at pH=pKa1 and the second one occurs at pH=pKa2.[18] Each segment of the curve that contains a midpoint at its center is called the buffer region. Because the buffer regions consist of the acid and its conjugate base, it can resist pH changes when base is added until the next equivalent points.[5]
Applications of acids
In industry
Acids are fundamental reagents in treating almost all processes in modern industry. Sulfuric acid, a diprotic acid, is the most widely used acid in industry, and is also the most-produced industrial chemical in the world. It is mainly used in producing fertilizer, detergent, batteries and dyes, as well as used in processing many products such like removing impurities.[19] According to the statistics data in 2011, the annual production of sulfuric acid was around 200 million tonnes in the world.[20] For example, phosphate minerals react with sulfuric acid to producephosphoric acid for the production of phosphate fertilizers, andzinc is produced by dissolving zinc oxide into sulfuric acid, purifying the solution and electrowinning.
Acids are often used to remove rust and other corrosion from metals in a process known aspickling. They may be used as an electrolyte in awet cell battery, such assulfuric acid in acar battery.
In food
Carbonated water (H2CO3 aqueous solution) is commonly added to soft drinks to make them effervesce.
Tartaric acid is an important component of some commonly used foods like unripe mangoes and tamarind. Natural fruits and vegetables also contain acids.Citric acid is present in oranges, lemon and other citrus fruits.Oxalic acid is present in tomatoes, spinach, and especially incarambola andrhubarb; rhubarb leaves and unripe carambolas are toxic because of high concentrations of oxalic acid.Ascorbic acid (Vitamin C) is an essential vitamin for the human body and is present in such foods as amla (Indian gooseberry), lemon, citrus fruits, and guava.
Many acids can be found in various kinds of food as additives, as they alter their taste and serve as preservatives.Phosphoric acid, for example, is a component ofcola drinks.Acetic acid is used in day-to-day life as vinegar. Citric acid is used as a preservative in sauces and pickles.
Carbonic acid is one of the most common acid additives that are widely added insoft drinks. During the manufacturing process, CO2 is usually pressurized to dissolve in these drinks to generate carbonic acid. Carbonic acid is very unstable and tends to decompose into water and CO2 at room temperature and pressure. Therefore, when bottles or cans of these kinds of soft drinks are opened, the soft drinks fizz and effervesce as CO2 bubbles come out.[21]
Certain acids are used as drugs.Acetylsalicylic acid (Aspirin) is used as a pain killer and for bringing down fevers.
In human bodies
Acids play important roles in the human body. The hydrochloric acid present in the stomach aids digestion by breaking down large and complex food molecules. Amino acids are required for synthesis of proteins required for growth and repair of body tissues. Fatty acids are also required for growth and repair of body tissues. Nucleic acids are important for the manufacturing of DNA and RNA and transmitting of traits to offspring through genes. Carbonic acid is important for maintenance of pH equilibrium in the body.
Human bodies contain a variety of organic and inorganic compounds, among thosedicarboxylic acids play an essential role in many biological behaviors. Many of those acids areamino acids, which mainly serve as materials for the synthesis of proteins.[22] Other weak acids serve as buffers with their conjugate bases to keep the body's pH from undergoing large scale changes that would be harmful to cells.[23] The rest of the dicarboxylic acids also participate in the synthesis of various biologically important compounds in human bodies.
Acids are used ascatalysts in industrial and organic chemistry; for example,sulfuric acid is used in very large quantities in thealkylation process to produce gasoline. Some acids, such as sulfuric, phosphoric, and hydrochloric acids, also effectdehydration andcondensation reactions. In biochemistry, manyenzymes employ acid catalysis.[24]
Many biologically important molecules are acids.Nucleic acids, which contain acidicphosphate groups, includeDNA andRNA. Nucleic acids contain the genetic code that determines many of an organism's characteristics, and is passed from parents to offspring. DNA contains the chemical blueprint for the synthesis ofproteins, which are made up ofamino acid subunits.Cell membranes containfatty acidesters such asphospholipids.
An α-amino acid has a central carbon (the α oralpha carbon) that is covalently bonded to acarboxyl group (thus they arecarboxylic acids), anamino group, a hydrogen atom and a variable group. The variable group, also called the R group or side chain, determines the identity and many of the properties of a specific amino acid. Inglycine, the simplest amino acid, the R group is a hydrogen atom, but in all other amino acids it is contains one or more carbon atoms bonded to hydrogens, and may contain other elements such as sulfur, oxygen or nitrogen. With the exception of glycine, naturally occurring amino acids arechiral and almost invariably occur in theL-configuration.Peptidoglycan, found in some bacterialcell walls contains someD-amino acids. At physiological pH, typically around 7, free amino acids exist in a charged form, where the acidic carboxyl group (-COOH) loses a proton (-COO−) and the basic amine group (-NH2) gains a proton (-NH+ 3). The entire molecule has a net neutral charge and is azwitterion, with the exception of amino acids with basic or acidic side chains.Aspartic acid, for example, possesses one protonated amine and two deprotonated carboxyl groups, for a net charge of −1 at physiological pH.
Fatty acids and fatty acid derivatives are another group of carboxylic acids that play a significant role in biology. These contain long hydrocarbon chains and a carboxylic acid group on one end. The cell membrane of nearly all organisms is primarily made up of aphospholipid bilayer, amicelle of hydrophobic fatty acid esters with polar, hydrophilicphosphate "head" groups. Membranes contain additional components, some of which can participate in acid–base reactions.
Acid–base equilibrium plays a critical role in regulatingmammalian breathing.Oxygen gas (O2) drivescellular respiration, the process by which animals release the chemicalpotential energy stored in food, producingcarbon dioxide (CO2) as a byproduct. Oxygen and carbon dioxide are exchanged in thelungs, and the body responds to changing energy demands by adjusting the rate ofventilation. For example, during periods of exertion the body rapidly breaks down storedcarbohydrates and fat, releasing CO2 into the blood stream. In aqueous solutions such as blood CO2 exists in equilibrium withcarbonic acid andbicarbonate ion.
CO2 + H2O ⇌ H2CO3 ⇌ H+ + HCO−3
It is the decrease in pH that signals the brain to breathe faster and deeper, expelling the excess CO2 and resupplying the cells with O2.
Cell membranes are generally impermeable to charged or large, polar molecules because of thelipophilic fatty acyl chains comprising their interior. Many biologically important molecules, including a number of pharmaceutical agents, are organic weak acids that can cross the membrane in their protonated, uncharged form but not in their charged form (i.e., as the conjugate base). For this reason the activity of many drugs can be enhanced or inhibited by the use of antacids or acidic foods. The charged form, however, is often more soluble in blood andcytosol, both aqueous environments. When the extracellular environment is more acidic than the neutral pH within the cell, certain acids will exist in their neutral form and will be membrane soluble, allowing them to cross the phospholipid bilayer. Acids that lose a proton at theintracellular pH will exist in their soluble, charged form and are thus able to diffuse through the cytosol to their target.Ibuprofen,aspirin andpenicillin are examples of drugs that are weak acids.
Acarboxylic acid has the general formula R-C(O)OH, where R is an organic radical. The carboxyl group -C(O)OH contains acarbonyl group, C=O, and ahydroxyl group, O-H.
Normal carboxylic acids are the direct union of a carbonyl group and a hydroxyl group. Invinylogous carboxylic acids, a carbon-carbon double bond separates the carbonyl and hydroxyl groups.
^abcThe International Union of Pure and Applied Chemistry (IUPAC)."IUPAC - acid (A00071)".goldbook.iupac.org. Retrieved17 November 2025.
^Petrucci, R. H.; Harwood, R. S.; Herring, F. G. (2002).General Chemistry: Principles and Modern Applications (8th ed.). Prentice Hall. p. 146.ISBN0-13-014329-4.
^Lias, S. G.; Liebman, J. F.; Levin, R. D. (1984). "Evaluated Gas Phase Basicities and Proton Affinities of Molecules; Heats of Formation of Protonated Molecules".Journal of Physical and Chemical Reference Data.13 (3): 695.Bibcode:1984JPCRD..13..695L.doi:10.1063/1.555719.
^Laurence, C. and Gal, J-F. Lewis Basicity and Affinity Scales, Data and Measurement, (Wiley 2010) pp 50–51 ISBN 978-0-470-74957-9
^Cramer, R. E.; Bopp, T. T. (1977). "Graphical display of the enthalpies of adduct formation for Lewis acids and bases".Journal of Chemical Education.54:612–613.doi:10.1021/ed054p612. The plots shown in this paper used older parameters. Improved E&C parameters are listed inECW model.
^Wyman, Jeffries; Tileston Edsall, John. "Chapter 9: Polybasic Acids, Bases, and Ampholytes, Including Proteins".Biophysical Chemistry - Volume 1. p. 477.
^de Levie, Robert (1999).Aqueous Acid–Base Equilibria and Titrations. New York: Oxford University Press.
^Jameson, Reginald F. (1978). "Assignment of the proton-association constants for 3-(3,4-dihydroxyphenyl)alanine (L-dopa)".Journal of the Chemical Society, Dalton Transactions (1):43–45.doi:10.1039/DT9780000043.
^Graham, Timur (2006)."Acid Buffering".Acid Base Online Tutorial. University of Connecticut. Archived fromthe original on 13 February 2016. Retrieved6 February 2016.
