The acidity of protons depends on the electronegativity of the atoms to which they are attached. The more electronegative the atom, the more acidic the proton will be. Therefore, a hydrogen atom attached to a halogen atom will be more acidic than a hydrogen atom attached to oxygen. A hydrogen atom attached to oxygen will be more acidic than a hydrogen atom attached to nitrogen. Hydrogen atoms attached to carbon are not usually acidic at all.
pKa is a measure of the strength of an acid. The lower the value of pKa the stronger the acid.pKa is the negative logarithm of Ka which is a measure of the dissociation or ionization of the acid. The larger the value of Ka, the stronger the acid.
Inductive effects can affect the stability of the conjugate base by stabilizing or destabilizing the negative charge. Electron-withdrawing groups such as halogens diminish the charge and stabilize the conjugate base, resulting in a stronger acid. Electron-donating groups (e.g. alkyl groups) will increase the charge and destabilize the conjugate base, resulting in a weaker acid.
A negative charge can be stabilized by resonance, resulting in delocalization of the charge over two or more atoms. Carboxylic acids are acidic because the resulting carboxylate ion can be stabilized by delocalization of the charge between two oxygen atoms. Phenols are acidic because the resulting phenolate ion can be stabilized by delocalization of the charge between the oxygen and three carbon atoms. Alcohols are only weakly acidic because the charge on the resulting alkoxide ion is localized on the oxygen and destabilized by the inductive effect of the alkyl group.
Amines and amides are very weak acids. However, amides are more acidic than amines due to resonance and inductive effects.
The acidic protons of various molecules are not equally acidic and their relative acidity depends on a number of factors, one of which is the electronegativity of the atom to which they are attached. For example, consider hydrofluoric acid, ethanoic acid, and methylamine (Fig. 1). Hydrofluoric acid has the most acidic proton since the hydrogen is attached to a stronglyelectronegative fluorine. The fluorine strongly polarizes the H–F bond such that the hydrogen becomes highly electron deficient and is easily lost. Once the proton is lost, the fluoride ion can stabilize the resulting negative charge.
The acidic protons on methylamine are attached to nitrogen which is less elec-tronegative than fluorine. Therefore, the N–H bonds are less polarized, and the protons are less electron deficient. If one of the protons is lost, the nitrogen is left with a negative charge which it cannot stabilize as efficiently as a halide ion. All of this means that methylamine is a much weaker acid than hydrogen fluoride.
Ethanoic acid is more acidic than methylamine but less acidic than hydrofluoric acid. This is because the electronegativity of oxygen lies between that of a halogen and that of a nitrogen atom.
These differences in acid strength can be demonstrated if the three molecules above are placed in water. Mineral acids such as HF, HCl, HBr, and HI are strong acids and dissociate or ionize completely (Fig. 2).
Ethanoic acid (acetic acid) partially dissociates in water and an equilibrium is set up between the carboxylic acid (termed the free acid) and the carboxylate ion (Fig. 3). An acid which only partially ionizes in this manner is termed a weak acid.
If methylamine is dissolved in water, none of the acidic protons are lost at all and the amine behaves as a weak base instead of an acid, and is in equilibrium with its protonated form (Fig. 4).
Methylamine can act as an acid but it has to be treated with a strong base such as butyl lithium (Fig. 5).
Lastly, hydrogen atoms attached to carbon are not usually acidic since carbon atoms are not electronegative. There are exceptions to this rule.
Acids can be described as being weak or strong and the pKa is a measure of this. Dissolving acetic acid in water, results in an equilibrium between the carboxylic acid and the carboxylate ion (Fig. 6).
Ethanoic acid on the left hand of the equation is termed the free acid, while the carboxylate ion formed on the right hand side is termed its conjugate base. The extent of ionization or dissociation is defined by the equilibrium constant (Keq );
Keq is normally measured in a dilute aqueous solution of the acid and so the concentration of water is high and assumed to be constant. Therefore, we can rewrite the equilibrium equation in a simpler form where Ka is the acidity con- stant and includes the concentration of pure water (55.5 M).
The acidity constant is also a measure of dissociation and of how acidic a particu-lar acid is. The stronger the acid, the more it is ionized and the greater the con-centration of products in the above equation. This means that a strong acid has a high Ka value. The Ka values for the following ethanoic acids are in brackets and demonstrate that the strongest acid in the series is trichloroacetic acid.
Kavalues are awkward to work with and so it is more usual to measure theacidic strength as a pKa value rather than Ka. The pKa is the negative logarithm of Ka (pKa = - log10Ka) and results in more manageable numbers. The pKavalues for each of the above ethanoic acids is shown in brackets below. The strongest acid (trichloroacetic acid) has the lowest pKa value.
Therefore the stronger the acid, the higher the value of Ka, and the lower the value of pKa. An amine such as ethylamine (CH3CH2NH2) is an extremely weak acid (pKa = 40) compared to ethanol (pKa = 16). This is due to the relative electronega-tivities of oxygen and nitrogen as described above. However, the electronegativ-ity of neighboring atoms is not the only influence on acidic strength. For example, the pKa values of ethanoic acid (4.76), ethanol (16), and phenol (10) show that ethanoic acid is more acidic than phenol, and that phenol is more acidic than ethanol. The difference in acidity is quite marked, yet hydrogen is attached to oxygen in all three structures.
Similarly, the ethanoic acids Cl3CCO2H (0.63), Cl2CHCO2H (1.26), ClCH2CO2H (2.87), and CH3CO2H (4.76) have significantly different pKa values and yet the acidic hydrogen is attached to an oxygen in each of these structures. Therefore, factors other than electronegativity have a role to play in determining acidic strength.
Stabilizing the negative charge of the conjugate base is important in determining the strength of the acid and so any effect which stabilizes the charge will result in a stronger acid. Substituents can help to stabilize a negative charge and do so by an inductive effect. This is illustrated by comparing the pKa values of the alcohols CF3CH2OH and CH3CH2OH (12.4 and 16, respectively) where CF3CH2OH is more acidic than CH3CH2OH. This implies that the anion CF3CH2O is more stable than CH3CH2O (Fig. 7).
Fluorine atoms are strongly electronegative and this means that each C–F bond is strongly polarized such that the carbon bearing the fluorine atoms becomes strongly electropositive. Since this carbon atom is now electron deficient, it will ‘demand’ a greater share of the electrons in the neighboring C–C bond. This results in electrons being withdrawn from the neighboring carbon, making it elec-tron deficient too. This inductive effect will continue to be felt through the various bonds of the structure. It will decrease through the bonds but it is still significant enough to be felt at the negatively charged oxygen. Since the inductive effect is electron withdrawing it will decrease the negative charge on the oxygen and help to stabilize it. This means that the original fluorinated alcohol will lose its proton more readily and will be a stronger acid.
This inductive effect explains the relative acidities of the chlorinated ethanoic acids Cl3CCO2H (0.63), Cl2CHCO2H (1.26), ClCH2CO2H (2.87), and CH3CO2H (4.76). Trichloroethanoic acid is the strongest acid since its conjugate base (the carboxylate ion) is stabilized by the inductive effect created by three electronegative chlorine atoms. As the number of chlorine atoms decrease, so does the inductive effect .
Inductive effects also explain the difference between the acid strengths of ethylamine (pKa ~ 40) and ammonia (pKa ~ 33). The pKa values demonstrate that ammonia is a stronger acid than ethylamine. In this case, the inductive effect is electron donating. The alkyl group of ethylamine enhances the negative charge of the conjugate base and so destabilizes it, making ethylamine a weaker acid than ammonia (Fig. 8).
The negative charge on some conjugate bases can be stabilized by resonance. Resonance involves the movement of valence electrons around a structure, resulting in the sharing of charge between different atoms – a process called delocalization. The effects of resonance can be illustrated by comparing theacidities of ethanoic acid (pKa 4.76), phenol (pKa 10.0) and ethanol (pKa 12.4). The pKa values illustrate that ethanoic acid is a stronger acid than phenol, and that phenol is a stronger acid than ethanol.
The differing acidic strengths of ethanoic acid, phenol and ethanol can be explained by considering the relative stabilities of their conjugate bases (Fig. 9).
The charge of the carboxylate ion is on an oxygen atom, and since oxygen is electronegative, the charge is stabilized. However, the charge can be shared with the other oxygen leading to delocalization of the charge. This arises by a resonance interaction between a lone pair of electrons on the negatively charged oxygen and the π electrons of the carbonyl group (Fig. 10). A lone pair of electrons on the ‘bottom’ oxygen forms a new π bond to the neighboring carbon. At the same time as this takes place, the weak π bond of the carbonyl group breaks. This is essential or else the carbonyl carbon would end up with five bonds and that is not permit-ted. Both electrons in the original π bond now end up on the ‘top’ oxygen which means that this oxygen ends up with three lone pairs and gains a negative charge. Note that the π bond and the charge have effectively ‘swapped places’. Both the structures involved are called resonance structures and are easily interconvertible. The negative charge is now shared or delocalized equally between both oxygens and is stabilized. Therefore, ethanoic acid is a stronger acid than one would expect based on the electronegativity of oxygen alone.
Phenol is less acidic than ethanoic acid but is more acidic than ethanol. Once again, resonance can explain these differences. The conjugate base of phenol is called the phenolate ion. In this case, the resonance process can be carried out sev-eral times to place the negative charge on four separate atoms – the oxygen atom and three of the aromatic carbon atoms (Fig. 11). The fact that the negative charge can be spread over four atoms might suggest that the phenolate anion should be more stable than the carboxylate anion, since the charge is spread over more atoms. However, with the phenolate ion, three of the resonance structures place the charge on a carbon atom which is much less electronegative than an oxygen atom. These resonance structures will be far less important than the resonance structure having the charge on oxygen. As a result, delocalization is weaker for the phenolate ion than it is for the ethanoate ion. Nevertheless, a certain amount of delocalization still takes place which is why a phenolate ion is more stable than an ethoxide ion.
Lastly, we turn to ethanol. The conjugate base is the ethoxide ion which cannot be stabilized by delocalizing the charge, since resonance is not possible. There is no π bond available to participate in resonance. Therefore, the negative charge is localized on the oxygen. Furthermore, the inductive donating effect of the neigh-boring alkyl group (ethyl) enhances the charge and destabilizes it (Fig. 12). This makes the ethoxide ion the least stable (or most reactive) of the three anions we have studied. As a result, ethanol is the weakest acid.
Amines and amides are very weak acids and only react with very strong bases.
The pKa values for ethanamide and ethylamine are 15 and 40, respectively, which means that ethanamide has the more acidic proton (Fig. 13). This can be explained by resonance and inductive effects (Fig. 14).
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