Acid strength
Acid strength is the tendency of an acid, symbolised by the chemical formula , to dissociate into a proton, , and an
Examples of
A weak acid is only partially dissociated, with both the undissociated acid and its dissociation products being present, in solution, in equilibrium with each other.
Acetic acid () is an example of a weak acid. The strength of a weak acid is quantified by its acid dissociation constant, value.
The strength of a weak
Measures of acid strength
The usual measure of the strength of an acid is its acid dissociation constant (), which can be determined experimentally by titration methods. Stronger acids have a larger and a smaller logarithmic constant () than weaker acids. The stronger an acid is, the more easily it loses a proton, . Two key factors that contribute to the ease of deprotonation are the polarity of the bond and the size of atom A, which determine the strength of the bond. Acid strengths also depend on the stability of the conjugate base.
While the value measures the tendency of an acidic solute to transfer a proton to a standard solvent (most commonly water or
When the acidic medium in question is a dilute aqueous solution, the is approximately equal to the pH value, which is a negative logarithm of the concentration of aqueous in solution. The pH of a simple solution of an acid in water is determined by both and the acid concentration. For weak acid solutions, it depends on the degree of dissociation, which may be determined by an equilibrium calculation. For concentrated solutions of acids, especially strong acids for which pH < 0, the value is a better measure of acidity than the pH.
Strong acids
A strong acid is an acid that dissociates according to the reaction
where S represents a solvent molecule, such as a molecule of water or dimethyl sulfoxide (DMSO), to such an extent that the concentration of the undissociated species 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 is hydrochloric acid.
- (in aqueous solution)
Any acid with a value which is less than about -2 is classed as a strong acid. This results from the very high
The following are strong acids in aqueous and dimethyl sulfoxide solution. The values of , cannot be measured experimentally. The values in the following table are average values from as many as 8 different theoretical calculations.
Estimated pKa 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 acid = −1.6 [5]
- Sulfuric acid (first dissociation only, ≈ −3)[6]: (p. 171)
The following can be used as protonators in organic chemistry
- Fluoroantimonic acid
- Magic acid
- Carborane superacid
- Fluorosulfuric acid ( = −6.4)[7]
Weak acids
A weak acid is a substance that partially dissociates when it is dissolved in a solvent. In solution there is an equilibrium between the acid, , and the products of dissociation.
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 a dissociation constant, , defined as follows, where signifies the concentration of a chemical moiety, X.
When a numerical value of is known it can be used to determine the extent of dissociation in a solution with a given concentration of the acid, , by applying the law of conservation of mass.
where is the value of the
This equation shows that the pH of a solution of a weak acid depends on both its value and its concentration. Typical examples of weak acids include acetic acid and phosphorous acid. An acid such as oxalic acid () is said to be
For a more rigorous treatment of acid strength see acid dissociation constant. This includes acids such as the dibasic acid succinic acid, for which the simple method of calculating the pH of a solution, shown above, cannot be used.
Experimental determination
The experimental determination of a value is commonly performed by means of a titration.[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
until only the deprotonated species, , remains in solution. At each point in the titration pH is measured using a glass electrode and a pH meter. The equilibrium constant is found by fitting calculated pH values to the observed values, using the method of least squares.
Conjugate acid/base pair
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 = 1.75 x 10−5. Its conjugate base is the acetate ion with Kb = 10−14/Ka = 5.7 x 10−10 (from the relationship Ka × Kb = 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
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 to Brønsted–Lowry acid–base theory, the solvent S can accept a proton.
For example, hydrochloric acid is a weak acid in solution in pure acetic acid, , which is more acidic than water.
The extent of ionization of the hydrohalic acids decreases in the order . Acetic acid is said to be a differentiating solvent for the three acids, while water is not.[6]: (p. 217)
An important example of a solvent which is more basic than water is dimethyl sulfoxide, DMSO, . 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 values in solution in DMSO and other solvents can be found at Acidity–Basicity Data in Nonaqueous Solvents.
Superacids are strong acids even in solvents of low dielectric constant. Examples of superacids are fluoroantimonic acid and magic acid. Some superacids can be crystallised.[9] They can also quantitatively stabilize carbocations.[10]
Factors determining acid strength
The inductive effect
In organic carboxylic acids, an electronegative substituent can pull electron density out of an acidic bond through the inductive effect, resulting in a smaller value. The effect decreases, the further the electronegative element is from the carboxylate group, as illustrated by the following series of
Structure | Name | pKa |
---|---|---|
2-chlorobutanoic acid | 2.86 | |
3-chlorobutanoic acid | 4.0 | |
4-chlorobutanoic acid | 4.5 | |
butanoic acid | 4.5 |
Effect of oxidation state
In a set of
Structure | Name | Oxidation state |
pKa |
---|---|---|---|
perchloric acid | 7 | -8† | |
chloric acid | 5 | -1 | |
chlorous acid | 3 | 2.0 | |
hypochlorous acid | 1 | 7.53 |
† theoretical
References
- ^ 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.
- ISBN 0-13-841891-8
- ISBN 0-201-05660-7
- S2CID 29697201.
- ^ Bell, R. P. (1973), The Proton in Chemistry (2nd ed.), Ithaca, NY: Cornell University Press
- ^ ISBN 978-0-13-039913-7.
- ^ doi:10.1139/v78-385.
- ISBN 0-471-18817-4. Chapter 4: Experimental Procedure for Potentiometric pHMeasurement of Metal Complex Equilibria
- .
- .
- .
- ^ 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
- doi:10.1021/ed054p612. The plots shown in this paper used older parameters. Improved E&C parameters are listed in ECW model.
External links
- Titration of acids - freeware for data analysis and simulation of potentiometric titration curves