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Acid strength

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Measure of the tendency of an acid to dissociate

Acids and bases

Acceptor number Acid Acid–base reaction Acid–base homeostasis Acid strength Acidity function Amphoterism Base Buffer solutions Dissociation constant Donor number Equilibrium chemistry Extraction Hammett acidity function pH Proton affinity Self-ionization of water Titration Lewis acid catalysis Frustrated Lewis pair Chiral Lewis acid ECW model
Acid types
Brønsted–Lowry Lewis Mineral Organic Oxide Strong Superacids Weak Solid
Base types
Brønsted–Lowry Lewis Organic Oxide Strong Superbases Non-nucleophilic Weak
v t e

Measures of acid strength

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The usual measure of the strength of an acid is itsacid dissociation constant ( $K_{a}$ ), which can bedetermined experimentally bytitration methods. Stronger acids have a larger $K_{a}$ and a smaller logarithmic constant ( $\mathrm {p} K_{a}=-\log K_{a}$ ) than weaker acids. The stronger an acid is, the more easily it loses a proton,H⁺. Two key factors that contribute to the ease ofdeprotonation are thepolarity of theH−A bond and the size of atomA, which determine the strength of theH−A bond. Acid strengths also depend on the stability of the conjugate base.

While the $K_{a}$ value measures the tendency of an acidic solute to transfer a proton to a standard solvent (most commonly water orDMSO), the tendency of an acidic solvent to transfer a proton to a reference solute (most commonly a weakaniline base) is measured by itsHammett acidity function, the $H_{0}$ value. Although these two concepts of acid strength often amount to the same general tendency of a substance to donate a proton, the $K_{a}$ and $H_{0}$ values are measures of distinct properties and may occasionally diverge. For instance, hydrogen fluoride, whether dissolved in water ( $\mathrm {p} K_{a}=3.2$ ) or DMSO ( $\mathrm {p} K_{a}=15$ ), has $\mathrm {p} K_{a}$ values indicating that it undergoes incomplete dissociation in these solvents, making it a weak acid. However, as the rigorously dried, neat acidic medium, hydrogen fluoride has an $H_{0}$ value of –15,^[1] making it a more strongly protonating medium than 100% sulfuric acid and thus, by definition, asuperacid.^[2] (To prevent ambiguity, in the rest of this article, "strong acid" will, unless otherwise stated, refer to an acid that is strong as measured by its $\mathrm {p} K_{{\ce {a}}}$ value ( $\mathrm {p} K_{{\ce {a}}}<-1.74$ ). This usage is consistent with the common parlance of most practicingchemists.)

When the acidic medium in question is a dilute aqueous solution, the $H_{0}$ is approximately equal to thepH value, which is a negative logarithm of the concentration of aqueousH⁺ in solution. The pH of a simple solution of an acid in water is determined by both $K_{a}$ and the acid concentration. For weak acid solutions, it depends on thedegree of dissociation, which may be determined by an equilibrium calculation. For concentrated solutions of acids, especially strong acids for which pH < 0, the $H_{0}$ value is a better measure of acidity than the pH.

Strong acids

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Image of a strong acid mostly dissociating. The small red circles represent H⁺ ions.

Astrong acid is an acid that dissociates according to the reaction

HA + S ⇌ SH⁺ + A⁻

where S represents a solvent molecule, such as a molecule of water ordimethyl sulfoxide (DMSO), to such an extent that the concentration of the undissociated speciesHA is too low to be measured. For practical purposes a strong acid can be said to be completely dissociated. An example of a strong acid isperchloric acid.

HClO₄ → H⁺ + ClO−4 (in aqueous solution)

Any acid with a $\mathrm {p} K_{a}$ value which is less than about −2 behaves as a strong acid. This results from the very highbuffer capacity of solutions with apH value of 1 or less and is known as theleveling effect.^[3]

The following are strong acids in aqueous and dimethyl sulfoxide solution. As mentioned above, because the dissociation is so strongly favored, the concentrations ofHA and thus the values of $\mathrm {p} K_{a}$ cannot be measured experimentally. The values in the following table are average values from as many as 8 different theoretical calculations.

EstimatedpK_a values^[4]
Acid	Formula	in water	in DMSO
Hydrochloric acid	HCl	−5.9 ± 0.4	−2.0 ± 0.6
Hydrobromic acid	HBr	−8.8 ± 0.8	−6.8 ± 0.8
Hydroiodic acid	HI	−9.5 ± 1	−10.9 ± 1
Triflic acid	H[CF₃SO₃]	−14 ± 2	−14 ± 2
Perchloric acid	H[ClO₄]	−15 ± 2	−15 ± 2

Also, in water

Nitric acidHNO₃ $\mathrm {p} K_{a}=-1.6$ ^[5]
Sulfuric acidH₂SO₄ (first dissociation only, $\mathrm {p} K_{a1}\approx -3$ )^[6]^: 171

The following can be used as protonators inorganic chemistry

Fluoroantimonic acidH[SbF₆]
Magic acidH[FSO₃SbF₅]
Carborane superacidH[CHB₁₁Cl₁₁]
Fluorosulfuric acidH[FSO₃] ( $\mathrm {p} K_{a}=-6.4$ )^[7]

Sulfonic acids, such asp-toluenesulfonic acid (tosylic acid) are a class of strong organicoxyacids.^[7] Some sulfonic acids can be isolated as solids.Polystyrene functionalized intopolystyrene sulfonate is an example of a substance that is a solid strong acid.

Weak acids

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Main article:Acid dissociation constant

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Image of a weak acid partly dissociating

A weak acid is a substance that partially dissociates or partly ionizes when it is dissolved in a solvent. In solution, there is an equilibrium between the acid,HA, and the products of dissociation.

HA ⇌ H⁺ + A⁻

The solvent (e.g. water) is omitted from this expression when its concentration is effectively unchanged by the process of acid dissociation. The strength of a weak acid can be quantified in terms of adissociation constant, $K_{a}$ , defined as follows, where ${\ce {[X]}}$ signifies the concentration of a chemical moiety,X. $K_{a}={\frac {[\mathrm {H} ^{+}][\mathrm {A} ^{-}]}{[\mathrm {HA} ]}}$ When a numerical value of $K_{a}$ is known it can be used to determine the extent of dissociation in a solution with a given concentration of the acid, $T_{H}$ , by applying the law ofconservation of mass. ${\begin{aligned}T_{H}&=[\mathrm {H} ^{+}]+[\mathrm {HA} ]\\&=[\mathrm {H} ^{+}]+{\frac {[\mathrm {A} ^{-}][\mathrm {H} ^{+}]}{K_{a}}}\\&=[\mathrm {H} ^{+}]+{\frac {[\mathrm {H} ^{+}]^{2}}{K_{a}}}\end{aligned}}$ where $T_{H}$ is the value of theanalytical concentration of the acid. When all the quantities in this equation are treated as numbers, ionic charges are not shown and this becomes aquadratic equation in the value of the hydrogen ion concentration value,[H⁺]. ${\frac {[\mathrm {H} ^{+}]^{2}}{K_{a}}}+[\mathrm {H} ^{+}]-T_{H}=0$ This equation shows that the pH of a solution of a weak acid depends on both its $K_{a}$ value and its concentration. Typical examples of weak acids includeacetic acid andphosphorous acid. An acid such asoxalic acid (HOOC−COOH) is said to bedibasic because it can lose two protons and react with two molecules of a simple base.Phosphoric acid (H₃PO₄) is tribasic.

For a more rigorous treatment of acid strength seeacid dissociation constant. This includes acids such as the dibasic acidsuccinic acid, for which the simple method of calculating the pH of a solution, shown above, cannot be used.

Experimental determination

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Main article:Acid dissociation constant § Experimental determination

The experimental determination of a $\mathrm {p} K_{a}$ value is commonly performed by means of atitration.^[8] A typical procedure would be as follows. A quantity of strong acid is added to a solution containing the acid or a salt of the acid, to the point where the compound is fully protonated. The solution is then titrated with a strong base

HA + OH⁻ → A⁻ + H₂O

until only the deprotonated species,A⁻, remains in solution. At each point in the titration pH is measured using aglass electrode and apH meter. The equilibrium constant is found by fitting calculated pH values to the observed values, using the method ofleast squares.

Conjugate acid/base pair

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It is sometimes stated that "the conjugate of a weak acid is a strong base". Such a statement is incorrect. For example, acetic acid is a weak acid which has a $K_{a}=1.75\times 10^{-5}$ . Its conjugate base is theacetate ion with $K_{b}=10^{-14}$ and $K_{a}=5.7\times 10^{-10}$ (from the relationship $K_{a}\times K_{b}=10^{-14}$ ), which certainly does not correspond to a strong base. The conjugate of a weak acid is often a weak base and vice versa.

Acids in non-aqueous solvents

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The strength of an acid varies from solvent to solvent. An acid which is strong in water may be weak in a less basic solvent, and an acid which is weak in water may be strong in a more basic solvent. According toBrønsted–Lowry acid–base theory, the solvent S can accept a proton.

HA + S ⇌ A⁻ + HS⁺

For example, hydrochloric acid is a weak acid in solution in pureacetic acid,HO₂CCH₃, which is less basic than water.

HO₂CCH₃ + HCl ⇌ (HO)₂CCH+3 + Cl⁻

The extent of ionization of thehydrohalic acids decreases in the order HI > HBr > HCl. Acetic acid is said to be adifferentiating solvent for the three acids, while water is not.^[6]^: 217

An important example of a solvent which is more basic than water isdimethyl sulfoxide, DMSO,(CH₃)₂SO. A compound which is a weak acid in water may become a strong acid in DMSO.Acetic acid is an example of such a substance. An extensive bibliography of $\mathrm {p} K_{a}$ values in solution in DMSO and other solvents can be found atAcidity–Basicity Data in Nonaqueous Solvents.^{[inappropriate external link?]}

Superacids are strong acids even in solvents of low dielectric constant.^[9] Examples of superacids arefluoroantimonic acid andmagic acid. Some superacids can be crystallised.^[10] They can also quantitatively stabilizecarbocations.^[11]

Lewis acids reacting with Lewis bases in gas phase and non-aqueous solvents have been classified in theECW model, and it has been shown that there is no one order of acid strengths.^[12] The relative acceptor strength of Lewis acids toward a series of bases, versus other Lewis acids, can be illustrated byC-B plots.^[13]^[14] It has been shown that to define the order of Lewis acid strength at least two properties must be considered. For the qualitativeHSAB theory the two properties are hardness and strength while for the quantitativeECW model the two properties are electrostatic and covalent.

Factors determining acid strength

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The inductive effect

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In organic carboxylic acids, an electronegative substituent can pull electron density out of an acidic bond through theinductive effect, resulting in a smaller $\mathrm {p} K_{a}$ value. The effect decreases, the further the electronegative element is from the carboxylate group, as illustrated by the following series ofhalogenated butanoic acids.

Structure	Name	pK_a
	2-chlorobutanoic acid	2.86
	3-chlorobutanoic acid	4.0
	4-chlorobutanoic acid	4.5
	butanoic acid	4.5

Effect of oxidation state

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In a set ofoxoacids of an element, $\mathrm {p} K_{a}$ values decrease with the oxidation state of the element. The oxoacids of chlorine illustrate this trend.^[6]^{: (p. 171)}

Name	Oxidation state	pK_a
perchloric acid	7	−8^†
chloric acid	5	−1
chlorous acid	3	2.0
hypochlorous acid	1	7.53

† theoretical

References

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^Liang, Joan-Nan Jack (1976).The Hammett Acidity Function for Hydrofluoric Acid and some related Superacid Systems (Ph.D. Thesis)(PDF). Hamilton, Ontario: McMaster University. p. 94.
^Miessler, Gary L.; Tarr, Donald Arthur (1999).Inorganic chemistry (2nd ed.). Upper Saddle River, NJ: Prentice-Hall. p. 170.ISBN 978-0-13-841891-5.
^Porterfield, William W. (1984).Inorganic chemistry: an unified approach. Reading, Mass: Addison Wesley Pub. Co. p. 260.ISBN 978-0-201-05660-0.
^Trummal, Aleksander; Lipping, Lauri; Kaljurand, Ivari; Koppel, Ilmar A.; Leito, Ivo (2016). "Acidity of strong acids in water and dimethyl sulfoxide".J. Phys. Chem. A.120 (20):3663–3669.Bibcode:2016JPCA..120.3663T.doi:10.1021/acs.jpca.6b02253.PMID 27115918.S2CID 29697201.
^Bell, R. P. (1973),The Proton in Chemistry (2nd ed.), Ithaca, NY: Cornell University Press
^^a ^b ^cHousecroft, C. E.; Sharpe, A. G. (2004).Inorganic Chemistry (2nd ed.). Prentice Hall.ISBN 978-0-13-039913-7.
^^a ^bGuthrie, J.P. (1978)."Hydrolysis of esters of oxy acids: pK_a values for strong acids".Can. J. Chem.56 (17):2342–2354.doi:10.1139/v78-385.
^Martell, A. E.; Motekaitis, R. J. (1992). "Chapter 4: Experimental Procedure for Potentiometric pH Measurement of Metal Complex Equilibria".Determination and Use of Stability Constants. Wiley.ISBN 0-471-18817-4.
^Waters, Charlie (2023-12-21)."Superacid: The Strongest Acids in the World".Chemniverse. Retrieved2024-10-11.
^Zhang, Dingliang; Rettig, Stephen J.; Trotter, James; Aubke, Friedhelm (1996). "Superacid Anions: Crystal and Molecular Structures of Oxonium Undecafluorodiantimonate(V), [H₃O][Sb₂F₁₁], Cesium Fluorosulfate, CsSO₃F, Cesium Hydrogen Bis(fluorosulfate), Cs[H(SO₃F)₂], Cesium Tetrakis(fluorosulfato)aurate(III), Cs[Au(SO₃F)₄], Cesium Hexakis(fluorosulfato)platinate(IV), Cs₂[Pt(SO₃F)₆], and Cesium Hexakis(fluorosulfato)antimonate(V), Cs[Sb(SO₃F)₆]".Inorg. Chem.35 (21):6113–6130.doi:10.1021/ic960525l.
^George A. Olah, Schlosberg RH (1968). "Chemistry in Super Acids. I. Hydrogen Exchange and Polycondensation of Methane and Alkanes in FSO₃H–SbF₅ ("Magic Acid") Solution. Protonation of Alkanes and the Intermediacy of CH₅⁺ and Related Hydrocarbon Ions. The High Chemical Reactivity of "Paraffins" in Ionic Solution Reactions".Journal of the American Chemical Society.90 (10):2726–7.doi:10.1021/ja01012a066.
^Vogel G. C.; Drago, R. S. (1996). "The ECW Model".Journal of Chemical Education.73 (8):701–707.Bibcode:1996JChEd..73..701V.doi:10.1021/ed073p701.
^Laurence, C. and Gal, J-F. Lewis Basicity and Affinity Scales, Data and Measurement, (Wiley 2010) pp 50-51ISBN 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.

External links

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Titration of acids- freeware for data analysis and simulation of potentiometric titration curves

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