Make sure that the proton in question is actually part of the functional group and is not simply attached to atoms that are attached to the functional group. Here is a short list of acidities of common functional groups in organic chemistry Note: that the pKa's are an approximate value for that functional group. The most stable conjugate base will be the strongest acid.
Method 1. These are two acids that are on the Table that we memorized. So, HCl is the stronger acid lower pKa. Method 2. So, HCl would be the stronger acid. Thus, the methoxide anion is the most stable lowest energy, least basic of the three conjugate bases, and the ethyl carbanion anion is the least stable highest energy, most basic. Conversely, ethanol is the strongest acid, and ethane the weakest acid.
When moving vertically within a given column of the periodic table, we again observe a clear periodic trend in acidity. This is best illustrated with the haloacids and halides: basicity, like electronegativity, increases as we move up the column. Vertical periodic trend in acidity and basicity. In order to make sense of this trend, we will once again consider the stability of the conjugate bases.
Because fluorine is the most electronegative halogen element, we might expect fluoride to also be the least basic halogen ion. But in fact, it is the least stable, and the most basic! It turns out that when moving vertically in the periodic table, the size of the atom trumps its electronegativity with regard to basicity. The atomic radius of iodine is approximately twice that of fluorine, so in an iodide ion, the negative charge is spread out over a significantly larger volume:.
We will see this idea expressed again and again throughout our study of organic reactivity, in many different contexts. For now, we are applying the concept only to the influence of atomic radius on base strength.
Because fluoride is the least stable most basic of the halide conjugate bases, HF is the least acidic of the haloacids, only slightly stronger than a carboxylic acid. HI, with a pK a of about -9, is almost as strong as sulfuric acid.
More importantly to the study of biological organic chemistry, this trend tells us that thiols are more acidic than alcohols. The pK a of the thiol group on the cysteine side chain, for example, is approximately 8. Recall that the driving force for a reaction is usually based on two factors: relative charge stability, and relative total bond energy. What explains this driving force? Consider first the charge factor: as we just learned, chloride ion on the product side is more stable than fluoride ion on the reactant side.
This partially accounts for the driving force going from reactant to product in this reaction: we are going from less stable ion to a more stable ion. What about total bond energy, the other factor in driving force? This also contributes to the driving force: we are moving from a weaker less stable bond to a stronger more stable bond. In the previous section we focused our attention on periodic trends — the differences in acidity and basicity between groups where the exchangeable proton was bound to different elements.
The first model pair we will consider is ethanol and acetic acid, but the conclusions we reach will be equally valid for all alcohol and carboxylic acid groups. Despite the fact that they are both oxygen acids, the pK a values of ethanol and acetic acid are strikingly different. What makes a carboxylic acid so much more acidic than an alcohol? As before, we begin by considering the stability of the conjugate bases, remembering that a more stable weaker conjugate base corresponds to a stronger acid.
In both species, the negative charge on the conjugate base is located on oxygen, so periodic trends cannot be invoked. For acetic acid, however, there is a key difference: two resonance contributors can be drawn for the conjugate base, and the negative charge can be delocalized shared over two oxygen atoms. This makes the ethoxide ion much less stable.
The delocalization of charge by resonance has a very powerful effect on the reactivity of organic molecules, enough to account for the difference of over 12 pK a units between ethanol and acetic acid and remember, pK a is a log expression, so we are talking about a factor of 10 12 between the K a values for the two molecules!
The resonance effect also nicely explains why a nitrogen atom is basic when it is in an amine, but not basic when it is part of an amide group. Recall that in an amide, there is significant double-bond character to the carbon-nitrogen bond, due to a minor but still important resonance contributor in which the nitrogen lone pair is part of a pi bond.
Notice that in this case, we are extending our central statement to say that electron density — in the form of a lone pair — is stabilized by resonance delocalization, even though there is not a negative charge involved. The lone pair on an amine nitrogen, by contrast, is not so comfortable — it is not part of a delocalized pi system, and is available to form a bond with any acidic proton that might be nearby.
Make a structural argument to account for its strength. Your answer should involve the structure of nitrate, the conjugate base of nitric acid. The negative charge on the oxygen that results from deprotonation of the acid is delocalized by resonance. A and B are ammonium groups, while C is an amine, so C is clearly the least acidic.
Looking at the conjugate base of B, we see that the lone pair electrons can be delocalized by resonance, making this conjugate base more stable than the conjugate base of A, where the electrons cannot be stabilized by resonance. Thus B is the most acidic. Often it requires some careful thought to predict the most acidic proton on a molecule. Ascorbic acid, also known as Vitamin C, has a pK a of 4. This overlap leads to a delocalization which extends from the ring out over the oxygen atom.
As a result, the negative charge is no longer entirely localized on the oxygen, but is spread out around the whole ion. Spreading the charge around makes the ion more stable than it would be if all the charge remained on the oxygen. However, oxygen is the most electronegative element in the ion and the delocalized electrons will be drawn towards it.
That means that there will still be a lot of charge around the oxygen which will tend to attract the hydrogen ion back again. That is why phenol is only a very weak acid. Why is phenol a much stronger acid than cyclohexanol? To answer this question we must evaluate the manner in which an oxygen substituent interacts with the benzene ring.
As noted in our earlier treatment of electrophilic aromatic substitution reactions, an oxygen substituent enhances the reactivity of the ring and favors electrophile attack at ortho and para sites.
It was proposed that resonance delocalization of an oxygen non-bonded electron pair into the pi-electron system of the aromatic ring was responsible for this substituent effect. A similar set of resonance structures for the phenolate anion conjugate base appears below the phenol structures. The resonance stabilization in these two cases is very different. An important principle of resonance is that charge separation diminishes the importance of canonical contributors to the resonance hybrid and reduces the overall stabilization.
The contributing structures to the phenol hybrid all suffer charge separation, resulting in very modest stabilization of this compound.
On the other hand, the phenolate anion is already charged, and the canonical contributors act to disperse the charge, resulting in a substantial stabilization of this species. The conjugate bases of simple alcohols are not stabilized by charge delocalization, so the acidity of these compounds is similar to that of water. An energy diagram showing the effect of resonance on cyclohexanol and phenol acidities is shown on the right.
Since the resonance stabilization of the phenolate conjugate base is much greater than the stabilization of phenol itself, the acidity of phenol relative to cyclohexanol is increased. Supporting evidence that the phenolate negative charge is delocalized on the ortho and para carbons of the benzene ring comes from the influence of electron-withdrawing substituents at those sites. In this reaction, the hydrogen ion has been removed by the strongly basic hydroxide ion in the sodium hydroxide solution.
Acids react with the more reactive metals to give hydrogen gas. Phenol is no exception - the only difference is the slow reaction because phenol is such a weak acid. Phenol is warmed in a dry tube until it is molten, and a small piece of sodium added.
There is some fizzing as hydrogen gas is given off. The mixture left in the tube will contain sodium phenoxide. Substitution of the hydroxyl hydrogen atom is even more facile with phenols, which are roughly a million times more acidic than equivalent alcohols.
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