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learning objectives
- Understand how molecular structure affects the strength of an acid or base.
We have seen that the strengths of acids and bases vary by several orders of magnitude. In this section, we review some of the structural and electronic factors that control the acidity or basicity of a molecule.
binding forces
In general, the stronger the \(\ce{A–H}\) or \(\ce{B–H^+}\) bond, the less likely the bond is broken and \(H^+\ ) ions and therefore the substance is less acidic. This phenomenon can be illustrated using the hydrogen halides:
| Relative acidity | HF | HCl | HBr | HELLO |
|---|---|---|---|---|
| H-X bond energy (kJ/mol) | 570 | 432 | 366 | 298 |
| pKa | 3.20 | −6,1 | −8,9 | −9,3 |
The trend in binding energies is due to a steady decrease in the overlap between the 1s orbital of hydrogen and the valence orbital of the halogen atom with increasing halogen size. The larger the atom to which the H is attached, the weaker the bond. Therefore, the bond between H and a large atom in a particular family, such as B. I or Te, weaker than the bond between H and a smaller atom of the same family, such as B. I or Te. B. F or O. As a result, acidic forces act The number of binary hydrides increases as we move down a column of the periodic table. For example, the acid series for its binary hydridesgroup 16The figures are as follows, with \(pK_a\) values in parentheses:
\[H_2O (14,00 = pK_w) < H_2S (7,05) < H_2Se (3,89) < H_2Te (2,6) \label{1}\]
Conjugate base stability
Whether we write an acid-base reaction as \(AH \rightleftharpoons A^−+H^+\) or \(BH^+ \rightleftharpoons B+H^+\), the conjugate base (\(A^− \ ) or \(B\)) contains one more electron pair than the original acid (\(AH\) or \(BH^+\)). Any agent that stabilizes the lone pair of electrons in the conjugate base favors the dissociation of \(H^+\) and makes the original acid a stronger acid. Let's see how this explains the relative acidity of the binary hydrides of the elements in the second row of the periodic table. The observed order of increasing acid is as follows, with pKa values given in parentheses:
\[CH_4 (~50) \ll NH_3 (~36) < H_2O (14,00) < HF (3,20) \label{2}\]
For example, consider the compounds at either end of this series: methane and hydrogen fluoride. The conjugate base of \(CH_4\) is \(CH_3^−\) and the conjugate base of \(HF\) is \(F^−\). Because fluorine is much more electronegative than carbon, fluorine can stabilize the negative charge on the \(F^−\) ion better than carbon can stabilize the negative charge on the CH3− ion. Consequently, \(\ce{HF}\) has a greater tendency to decay to form \(H^+\) and \(F^−\) than methane to form \(H^+\) and \(CH_3 ^ −\, making HF a much stronger acid than \(CH_4\).
The same trend is predicted by analyzing the properties of the conjugate acids. For a number of compounds of the general formula \(HE\), the E−H bond becomes more polar as E increases in electronegativity, favoring dissociation to form \(E^−\) and \(H ^+\ ). . Because of the increasing stability of the conjugate base and the increasing polarization of the E-H bond in the conjugate acid, the acid strengths of binary hydrides increase as we move from left to right up a row of the periodic table.
The acidity of binary hydrides increases as we move down a column or from left to right in a row of the periodic table.
the strongest known acid: the hydrosol cation
The stronger the acid, the weaker the covalent bond with a hydrogen atom. Therefore, the strongest possible acid is the molecule with the weakest bond. This is the hydrium(1+) cation, \(\ce{HeH^{+}}\), a positively charged ion formed by the reaction of a proton with a helium atom in the gas phase. It was first made in the laboratory in 1925 and is isoelectronic with molecular hydrogen (\ce{H2}}). With a proton affinity of 177.8 kJ/mol, it is the strongest known acid.
Ball and stick model of hydrosolIon (CC BY-SA 3.0;CCoil).
\(\ce{HeH^{+}}\) cannot be made in the condensed phase, as it will protonate any anion, molecule or atom it is associated with. However, it is possible to estimate ahypotheticaluse aqueous acidLei de Hess:
| NO+(G) | → | H+(G) | + is (G) | +178 kJ/mol |
| NO+(water) | → | NO+(G) | +973 kJ/mol | |
| H+(G) | → | H+(water) | −1530 kJ/mol | |
| Is(G) | → | Is(water) | +19 kJ/mol | |
| NO+(water) | → | H+(water) | + is (water) | -360 kJ/mol |
A change in dissociation free energy of -360 kJ/mol corresponds to a pKONEvon -63.
It has been suggested that \(\ce{HeH^{+}}\) should occur naturally in the interstellar medium, but it has not yet been discovered.
inductive effects
Atoms or groups of atoms in a molecule other than those to which the H is bonded can cause a change in the distribution of electrons within the molecule. This is called an induction effect and, like the coordination of water with a metal ion, can have a large impact on the acidity or basicity of the molecule. For example, hypohalo acids (general formula HOX, where X represents a halogen) have a hydrogen atom bonded to an oxygen atom. In aqueous solution everything creates the following equilibrium:
\[ HOX_{(aq)} \rightleftharpoons H^+_{(aq)} + OX^−{(aq)} \label{3}\]
However, the acidity of these acids varies by about three orders of magnitude due to the different electronegativity of the halogen atoms:
| HOX | Electronegativity of X | pKa |
|---|---|---|
| HOCl | 3,0 | 7.40 |
| HOBr | 2.8 | 8,55 |
| HOI | 2.5 | 10.5 |
As \(X\) increases in electronegativity, the electron density distribution within the molecule changes: Electrons are more strongly attracted to the halogen atom and in turn are drawn away from the H in the O-H bond, weakening the O- H and allowing the hydrogen to dissociate as \( H^+\).
The acidity of oxoacids with the general formula \(HOXO_n\) (with \(n\) = 0−3) strongly depends on the number of terminal oxygen atoms attached to the central atom \(X\). As shown in figure \(\PageIndex{1}\), the \(K_a\) values of chloroacids increase by a factor of \(10^4\) to \(10^6\). Oxygen because oxygen atoms are added one by one. The increase in acid strength with increasing number of terminal oxygen atoms is due to both an inductive effect and increased stabilization of the conjugate base.
Any inductive effect that removes electron density from an O-H bond increases the acidity of the compound.
Since oxygen is the second most electronegative element, adding terminal oxygen atoms will cause the O-H bond to lose electrons, making it weaker and thus increasing the strength of the acid. The colors in figure \(\PageIndex{1}\) show how the electrostatic potential, a measure of the strength of the interaction of a point charge anywhere on the surface of the molecule, changes as the number of terminal oxygen atoms increases. In figure \(\PageIndex{1}\) and figure \(\PageIndex{2}\), blue corresponds to low electron densities, while red corresponds to high electron densities. The oxygen atom in the O-H unit becomes less and less red from \(HClO\) to \(HClO_4\) (also written as \(HOClO_3\)), while the H atom becomes more and more blue, indicating that the electron density in the O-H unit decreases as the number of terminal oxygen atoms increases. A decrease in electron density in the O-H bond weakens it, facilitating the loss of hydrogen as \(H^) +\ ions, thereby increasing the strength of the acidic bond.
Less important, however, is the effect of shifting the negative charge to the conjugate base. As shown in Figure \(\PageIndex{2}\), the number of coordination structures that can be written for chlorooxoanions increases with increasing number of terminal oxygen atoms, so that the single negative charge can be successively displaced by more oxygen atoms.
The delocalization of electrons to the conjugate base increases the strength of the acid.
The electrostatic potential plots in Figure \(\PageIndex{2}\) show that the electron density on the terminal oxygen atoms decreases steadily as the number of atoms increases. The oxygen atom in ClO- is red, indicating that it is electron-rich, and in \(ClO_4^+\) the color of the oxygen changes progressively to green, indicating that the oxygen atoms become less and less electron-rich as the series progresses they become . For example, in the perchlorate ion (\(ClO_4^−\)) the single negative charge is distributed over all four oxygen atoms, while in the hypochlorite ion (\(OCl^−\)) the negative charge is widely spread over a single oxygen atom (Image \(\ Page Index{2}\)). As a result, the perchlorate ion has no localized negative charge for a proton to bind to. Consequently, the perchlorate anion has a much lower affinity for a proton than the hypochlorite ion, and perchloric acid is one of the strongest acids known.
As the number of terminal oxygen atoms increases, so does the number of coordination structures that can be written for the chlorooxoanions, and the single negative charge shifts to more oxygen atoms. As these electrostatic potential diagrams show, the electron density on the terminal oxygen atoms decreases steadily as the no. As the electron density on the oxygen atoms decreases, their affinity for a proton also decreases, making the anion less basic. This makes the original acetic acid more acidic.
Similar inductive effects also account for the acidity trend of oxoacids, which have the same number of oxygen atoms, as we move from left to right in a row of the periodic table. For example, \(H_3PO_4\) is a weak acid, \(H_2SO_4\) is a strong acid, and \(HClO_4\) is one of the strongest known acids. The number of terminal oxygen atoms increases steadily along the line, consistent with the observed increase in acidity. In addition, the electronegativity of the central atom from P to S in \(Cl\) increases steadily, causing electrons to be drawn from oxygen to the central atom, weakening the \(\ce{OH}\) bond increasing its strength oxyacidic.
A close look at the data in table \(\PageIndex{1}\) reveals two obvious anomalies: carbonic acid and phosphoric acid. If carbonic acid \((H_2CO_3\)) were a discrete molecule with the structure \(\ce{(HO)_2C=O}\), it would have a single terminal oxygen atom and should be comparable in acid strength to the hydrophosphoric acid (\ (H_3PO_4\)), for which pKa1 = 2.16. In contrast, the tabulated value of \(pK_{a1}\) for carbonic acid is 6.35, about 10,000 times weaker than expected. However, as we shall see, \(H_2CO_3\) is only a minor component of the aqueous solutions of \(CO_2\) known as carbonic acid. If phosphoric acid (\(H_3PO_3\)) actually had the structure \((HO)_3P\), there would be no terminal oxygen atoms attached to the phosphorus. It is therefore expected to be as strong an acid as \(HOCl\) (pKa = 7.40). In fact, \(pK_{a1}\) for phosphorous acid is 1.30 and the structure of phosphoric acid is \(\ce{(HO)_2P(=O)H}\) with one H atom attached directly to P is also bound a conjunction \(\ce{P=O}\). Therefore, the pKa1 of phosphoric acid is similar to that of other oxoacids with a terminal oxygen atom, e.g. B. \(H_3PO_4\). Fortunately, phosphoric acid is the only common oxo acid that has a hydrogen atom attached to the central atom instead of an oxygen.
| *\(H_2CO_3\) and \(H_2SO_3\) are at best minor components of aqueous solutions of \(CO_{2(g)}\) and \(SO_{2(g)}\) respectively, but such solutions are commonly referred to as carbonic acid or sulfuric acid. | ||||
|---|---|---|---|---|
| polyprotic acids | Type | \(pK_{a1}\) | \(pK_{a2}\) | \(pK_{a3}\) |
| Carbonic acid* | "\(H_2CO_3\)" | 6.35 | 10.33 | |
| citric acid | \(HO_2CCH-2C(OH)(CO_2H)CH_2CO_2H\) | 3.13 | 4,76 | 6.40 |
| malonic acid | \(HO-2CCH_2CO_2H\) | 2,85 | 5,70 | |
| oxalic acid | \(HO_2CCO_2H\) | 1,25 | 3,81 | |
| phosphoric acid | \(H_3PO_4\) | 2.16 | 7.21 | 12:32 |
| phosphoric acid | \(H_3PO_3\) | 1.3 | 6,70 | |
| succinic acid | \(HO_2CCH_2CH_2CO_2H\) | 4.21 | 5,64 | |
| sulfuric acid | \(H_2SO_4\) | −2,0 | 1,99 | |
| Sulfurous acid* | "\(H_2SO_3\)" | 1,85 | 7.21 | |
| polyprotic bases | Type | \(pK_{b1}\) | \(pK_{b2}\) | |
| Ethylenediamine | \(H_2N(CH_2)_2NH_2\) | 4.08 | 7.14 | |
| Piperazine | \(HN(CH_2CH_2)_2NH\) | 4.27 | 8,67 | |
| Propylenediamine | \(H_2N(CH_2)_3NH_2\) | 3,45 | 5.12 | |
Inductive effects are also observed in organic molecules containing electronegative substituents. The extent of the electron-withdrawing effect depends on the type and number of halogen substituents, as shown by the pKa values of various acetic acid derivatives:
\[pK_a CH_3CO_2H 4,76< CH_2ClCO_2H 2,87 As would be expected, fluorine, being more electronegative than chlorine, has a greater effect than chlorine, and the effect of three halogens is greater than the effect of two or one. From these data, note that the inductive effects can be quite large. For example, replacing the \(\ce{–CH_3}\) group in acetic acid with a \(\ce{–CF_3}\) group results in about a 10,000-fold increase in acidity! Example \(\PageIndex{1}\) Arrange the compounds in each row in order of increasing acid or base strength. The structures are shown here. given: connection string He asked: relative acid or base strengths Strategy: Use relative bond strengths, conjugate base stability, and induction effects to arrange the compounds in order of increasing ionization propensity in aqueous solution. Solution: Although sulfuric acid and sulfuric acid have two OH groups, the sulfur atom in sulfuric acid is attached to two terminal oxygen atoms, as opposed to one in sulfuric acid. Because oxygen is strongly electronegative, sulfuric acid is the stronger acid because the negative charge of the anion is stabilized by the extra oxygen atom. Comparing sulfuric acid and fluorosulfonic acid, we find that fluorine is more electronegative than oxygen. Thus, replacing an -OH with -F removes more electron density from the central S atom, which in turn removes electron density from both the S-OH bond and the O-H bond. Because the O-H bond is weaker, \(FSO_3H\) is a stronger acid than sulfuric acid. The predicted order of acid content given here is supported by the measured pKa values for these acids: \[pKa H_2SO_3 1,85 The structures of trifluoroamine and hydroxylamine are similar to those of ammonia. Trifluoroamine has all of the NH3 hydrogens replaced by fluorine, while hydroxylamine has one hydrogen atom replaced by OH. Replacing the three hydrogen atoms with fluorine strips N of its electron density, making the lone pair of electrons on N less available for bonding with a \(H^+\) ion. Therefore, \(NF_3\) is predicted to be a much weaker base than \(NH_3\). Because oxygen is more electronegative than hydrogen, replacing a hydrogen atom in \(NH_3\) with \(OH\) makes the amine less basic. Because oxygen is less electronegative than fluorine and only one hydrogen atom is replaced, the effect is smaller. The predicted order of increase in base power shown here is confirmed by the measured \(pK_b\) values: \[pK_bNF_3—< Trifluoroamine is such a weak base that it does not react with aqueous solutions of strong acids. Therefore, its fundamental ionization constant has never been measured. Exercise \(\PageIndex{1}\) Arrange the connections in each row in order \(HClO-3 > CH_3PO_3H_2 > H_3PO_4\) \(CF_3S^− < CH_3S^− < OH^−\)
Summary
Inductive effects and charge transfer significantly affect the acidity or basicity of a compound. The acid-base strength of a molecule depends largely on its structure. The weaker the A-H or B-H+ bond, the more likely it is to dissociate to form a \(H^+\) ion. Additionally, any agent that stabilizes the lone pair of electrons in the conjugate base favors the dissociation of \(H^+\), making the conjugate acid a stronger acid. Atoms or groups of atoms elsewhere in a molecule may also be important in determining acid or base strength through an induction effect, where a \(\ce{O-H}\) bond may be weakened and hydrogen is lost more easily than ( H^+\).
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