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learning goal
- Learn how molecular structure affects the strength of an acid or base.
We have seen that the strength of acids and bases varies by many orders of magnitude. In this section, we will examine some of the structural and electronic factors that control the acidity or alkalinity of a molecule.
adhesive force
In general, the stronger the bond \(\ce{A–H}\) or \(\ce{B–H^+}\), the less likely the bond will break and a \(H^ +\) ion , so the substance is less acidic. This effect can be illustrated in terms of hydrogen halides:
| relative acid strength | High frequency | hydrochloric acid | waterstofbromide | Hoi |
|---|---|---|---|---|
| H-X bindingsenergie (kJ/mol) | 570 | 432 | 366 | 298 |
| pKa | 3.20 | −6.1 | −8,9 | −9,3 |
The trend of binding energy is due to the steadily decreasing overlap between the hydrogen 1s orbital and the valence orbital of the halogen atom as the size of the halogen increases. The larger the atom to which H is attached, the weaker the bond. Therefore, the bond between H and a large atom in a particular group (such as I or Te) is weaker than the bond between H and a smaller atom in the same group (such as F or O). Therefore, the acid strength of binary compounds increases as we move up the column of the periodic table of elements, hydrides. For example, the acidity of binary hydridesGroup 16The entries are as follows, and the value \(pK_a\) is 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 you write the acid-base reaction as \(AH \rightleftharpoons A^−+H^+\) or \(BH^+ \rightleftharpoons B+H^+\), the conjugate base (\(A^− \ ) or \(B\)) contains one pair more free electrons than the parent acid (\(AH\) or \(BH^+\)). Anything that stabilizes the lone pair on a conjugate base promotes the dissociation of \(H^+\) and makes the original acid a stronger acid. Let's see how this explains the relative acidity of the binary hydride elements in the second row of the periodic table. The observed order of acidity increase is as follows, with pKa values in parentheses:
\[CH_4 (~50) \ll NH_3 (~36) < H_2O (14,00) < HF (3,20) \label{2}\]
For example, consider compounds at both ends of the series: methane and hydrofluoric acid. The conjugate base \(CH_4\) is \(CH_3^−\) and the conjugate base \(HF\) is \(F^−\). Since fluorine is more electronegative than carbon, fluorine stabilizes negative charges on \(F^-\) ions better than carbon stabilizes negative charges on CH3 ions. So \(H^+\) and \(CH_3^−\) are formed with methane, making HF a much stronger acid than \(CH_4\).
The same trend was predicted by analyzing the properties of conjugated acids. For a number of compounds of the general formula \(HE\), as the electronegativity of E increases, the E-H bond becomes more polar, promoting dissociation to form \(E^−\) and \(H^+ \) . due to the increased stability of the conjugate base and the increased polarity of the E-H bond in the conjugate acid, the acid strength of the binary hydride increases as we move from left to right in the row of the periodic table.
The acid strength of binary hydrides increases as we move down a column or row of the periodic table from left to right.
Strongest known acid: hydrogen and helium cation
The stronger the acid, the weaker the covalent bond with the hydrogen atom. So the strongest acid is probably the molecule with the weakest bond. It is a hydrogen-helium cation (1+) \(\ce{HeH^{+}}\), a positively charged ion formed by the reaction of a proton with a helium atom in the gas phase. It was first produced in the laboratory in 1925 from molecular hydrogen (\ce{H2}}) and other electrons. It is the strongest known acid with a proton affinity of 177.8 kJ/mol.
Ball and stick model of hydrogen and heliumion. (CC BY-SA 3.0;Coil C).
\(\ce{HeH^{+}}\) cannot be prepared in the condensed phase because it protonates any anion, molecule or atom bound to it. However, you can get approximaginaryThe acidity of the aqueous solution usedHess's law:
| Hallo+(G) | → | H+(G) | + op (G) | +178 kJ/mol |
| Hallo+(water) | → | Hallo+(G) | +973 kJ/mol | |
| H+(G) | → | H+(water) | −1530 kJ/mol | |
| Op(G) | → | Op(water) | +19J/mol | |
| Hallo+(water) | → | H+(water) | + op (water) | -360 kJ/mol |
A dissociation-free energy change of −360 kJ/mol is equal to pPotassiumA−63。
It has been suggested that \(\ce{HeH^{+}}\) should occur naturally in the interstellar medium, but this has not been detected.
induction effect
Atoms or groups of atoms in a molecule other than H-bonded atoms or groups of atoms can cause changes in the distribution of electrons in the molecule. This is called the inductive effect, and like the coordination of water with metal ions, it can have a major effect on the acidity or alkalinity of a molecule. For example, hypohalogen acids (general formula HOX, where X is halogen) have one hydrogen atom bonded to one oxygen atom. In aqueous solution, both provide the following equilibrium:
\[ HOX_{(aq)} \rightleftharpoons H^+_{(aq)} + OX^−{(aq)} \label{3}\]
However, due to the difference in electronegativity of the halogen atoms, the acidities of these acids differ by about three orders of magnitude:
| HOX | X electronegativity | pKa |
|---|---|---|
| hypochlorous acid | 3.0 | 7:40 am |
| waterstofbromide | 2.8 | 8.55 |
| HOI | 2.5 | 10,5 |
As the electronegativity of \(X\) increases, the electron density distribution in the molecule changes: electrons are more strongly attracted to the halogen atoms and away from H in the OH bond, weakening the OH bond and allowing hydrogen to dissociate into \ (H^+\).
The acidity of an oxoacid with the general formula \(HOXO_n\) (where \(n\) = 0-3) largely depends on the number of terminal oxygen atoms attached to the central atom \(X\). As shown in the figure \(\PageIndex{1}\), the value of \(K_a\) of chlorooxyacids increases with each oxygen approximately \(10^4\) to \(10^6\) times with more oxygen atoms are added. The increase in acid strength with the number of terminal oxygen atoms is due to the inductive action of the conjugate base and increased stability.
Any inductive effect that removes the electron density from the OH bonds increases the acidity of the compound.
Since oxygen is the second most electronegative element, adding a terminal oxygen atom pulls an electron away from the OH bond, weakening it and thus increasing the strength of the acid. The colors in the graph \(\PageIndex{1}\) show how the electrostatic potential, which is a measure of the strength of interactions between point charges at any position on a molecule's surface, changes as the number of terminal oxygen atoms increases . In the graph \(\PageIndex{ 1}\) and the figure \(\PageIndex{2}\), blue corresponds to a low electron density and red to a high electron density. The oxygen atoms in the OH unit gradually turn red from \(HClO\) to \(HClO_4\) (also written \(HOClO_3\)), while the H atoms gradually turn blue, indicating that the density electron O– De H unit increases as terminal oxygen decreases as atomic number increases. The reduced electron density in the OH bond weakens it, facilitating the loss of hydrogen as \(H^+\) ions, increasing the strength of the acid.
At least as important, however, is the delocalization effect of the negative charge in the conjugate base. As shown in Figure \(\PageIndex{2}\), as the number of terminal oxygen atoms increases, the number of terminal oxygen atoms increases, the number of resonance structures that can be stored by the chlorooxy anion, so that a single negative charge is dissociated into more successive oxygen atoms in the domain atom.
Delocalization of electrons in the conjugate base increases the strength of the acid.
The electrostatic potential plot in figure \(\PageIndex{2}\) shows that the electron density on the terminal oxygen atoms steadily decreases as their number increases. The oxygen atom in ClO− is red, indicating that it is electron rich, while the oxygen color gradually turns green in \(ClO_4^+\), indicating that the oxygen atom is becoming less and less electron rich through the series. For example, in the perchlorate ion (\(ClO_4^-\)) the single negative charge is delocalized on all four oxygen atoms, while in the hypochlorite ion (\(OCl^-\)) the negative charge is mainly on the single oxygen atoms (Fig. .\(\PageIndex{2}\)). As a result, the perchlorate ion has no local negative charge to which protons can bind. Therefore, the affinity of the perchlorate anion for the proton is much less than that of the hypochlorite ion, and perchlorate is one of the strongest known acids.
As the number of terminal oxygen atoms increases, the number of resonance structures that can be written for the chlorooxyanion increases and a single negative charge is transferred to more oxygen atoms. As can be seen in these electrostatic potential maps, the electron density on the terminal oxygen atoms steadily decreases as their number increases. As the electron density on the oxygen atoms decreases, their affinity for the protons decreases, making the anion less basic. As a result, the original oxyacid is more acidic.
The acidity trends of oxyacids with the same number of oxygen atoms are also related to a similar effect caused when 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 steadily increases in line, consistent with the observed increase in acidity. In addition, the electronegativity of the central atom steadily increases from P to S to \(Cl\), which attracts electrons from the oxygen to the central atom, weakening the bond of \(\ce{O - H}\) and the strength of the oxygenated acid.
Close examination of the data in the \(\PageIndex{1}\) table reveals two clear anomalies: carbonic acid and phosphoric acid. If carbonic acid \((H_2CO_3\)) is a single molecule with the structure \(\ce{(HO)_2C=O}\), it has a single terminal oxygen atom and the strength of the acid should be comparable to phosphoric acid ( \(H_3PO_4 \) ), where pKa1 = 2.16. In contrast, the table value of carbonic acid \(pK_{a1}\) is 6.35, which is about 10,000 times less than expected. However, as we shall see, \(H_2CO_3\) is only a small fraction of the aqueous solution \(CO_2\) called carbonic acid. Similarly, if phosphoric acid (\(H_3PO_3\)) actually had the structure \((HO)_3P\), it would not have the terminal oxygen attached to the phosphorus. Therefore, it is expected to have an acid strength similar to that of \(HOCl\) (pKa = 7.40). In fact, \(pK_{a1}\) of phosphoric acid is 1.30, and the structure of phosphoric acid is \(\ce{(HO)_2P(=O)H}\), one H atom is bonded directly to P and one bond \(\ce{P=O}\). Thus, the pKa1 of phosphorous acid is comparable to that of other oxoacids with one terminal oxygen atom, such as \(H_3PO_4\). Fortunately, phosphoric acid is the only common oxyacid in which a hydrogen atom is bonded to the central atom instead of oxygen.
| *\(H_2CO_3\) and \(H_2SO_3\) are at most minor components of \(CO_{2(g)}\) and \(SO_{2(g)}\) aqueous solutions, respectively, but such solutions are well known containing carbonic acid and sulfurous acid. | ||||
|---|---|---|---|---|
| Polyacid | klerk | \(pK_{a1}\) | \(pK_{a2}\) | \(pK_{a3}\) |
| carbonated* | „\(H_2CO_3\)” | 6:35 am | 10.33 | |
| citric acid | \(HO_2CCH-2C(OH)(CO_2H)CH_2CO_2H\) | 3.13 | 4.76 | 6.40 |
| Please | \(HO-2CCH_2CO_2H\) | 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 | |
| Sulphuric acid | \(H_2SO_4\) | −2,0 | 1,99 | |
| sulphurous acid* | „\(H_2SO_3\)” | 1,85 | 7.21 | |
| policy | klerk | \(pK_{b1}\) | \(pK_{b2}\) | |
| the ethylenediamines | \(H_2N(CH_2)_2NH_2\) | 4.08 | 7.14 | |
| Piperazine | \(HN(CH_2CH_2)_2NH\) | 4.27 | 8.67 | |
| propyleendiamine | \(H_2N(CH_2)_3NH_2\) | 3.45 | 5.12 | |
Inductive effects have also been observed in organic molecules containing electronegative substituents. The magnitude of the electron-withdrawing effect depends on the nature 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 you might expect, fluorine, which is more electronegative than chlorine, has a greater effect than chlorine, and three halogens have a greater effect than two or one. As can be seen from this data, the induction effect can be very large. For example, replacing the \(\ce{–CH_3}\) acetic acid group with \(\ce{–CF_3}\) increases the acidity by about 10,000 times! Example \(\PageIndex{1}\) Arrange the compounds in each series in order of increasing acid or base strength. The structure is shown here. dany: serial relationship Need: relative acid-base strength strategy: Rank compounds by increasing tendency to ionize in aqueous solution based on relative binding strength, conjugate base stability, and inductive effect. solution: While both sulfuric acid and sulphurous acid have two -OH groups, in sulfuric acid the sulfur atom is bonded to two terminal oxygen atoms while in sulfuric acid there is only one. Since oxygen is very electronegative, sulfuric acid is a stronger acid because the anion's negative charge is stabilized by the extra oxygen atom. By comparing sulfuric acid and fluorosulfonic acid, we noticed that fluorine is more electronegative than oxygen. Therefore, replacing the -OH with -F removes more electron density from the central S atom, which in turn removes the electron density from the S-OH and OH bonds. Due to the weak O-H bond \(FSO_3H\) is a stronger acid than sulfuric acid. The sequence of acid strength predictions given here is confirmed by the measured pKa values for these acids: \[pKa H_2SO_3 1,85 Both trifluoramine and hydroxylamine are structurally similar to ammonia. In trifluoramine, all hydrogen atoms in NH3 are replaced by fluorine atoms, while in hydroxylamine, one hydrogen atom is replaced by OH. Replacing three hydrogen atoms with fluorine reduces the electron density of N, making it more difficult for the lone pairs of electrons in N to combine with \(H^+\) ions. Thus \(NF_3\) is predicted to be a much weaker base than \(NH_3\). Similarly, since oxygen is more electronegative than hydrogen, replacing one of the hydrogen atoms 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 sequence of base power increase shown here is confirmed by the measured values\(pK_b\): \[pK_bNF_3—< Trifluoramine is a weak base and does not react with aqueous solutions of strong acids. Therefore, the fundamental ionization constant has never been measured. Exercise \(\PageIndex{1}\) Order connections from any series \(HClO-3 > CH_3PO_3H_2 > H_3PO_4\) \(CF_3S^− < CH_3S^− < OH^−\)
to summarize
Inductive effects and charge delocalization significantly affect the acidity or alkalinity of a compound. The acid-base strength of a molecule largely depends on its structure. The weaker the A-H or B-H+ bond, the more likely it is to dissociate and form \(H^+\) ions. In addition, anything that stabilizes the lone pair on the conjugate base promotes the dissociation of \(H^+\), making the conjugate acid a stronger acid. Atoms or groups of atoms in other parts of the molecule are also important in determining the strength of an acid or base through inductive effects that can weaken \(\ce{O–H}\) bonds and facilitate the loss of hydrogen. because \( H^+\ ) ions.
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