Basicity
A chemical species behaves as a base when it accepts protons (or donates electrons).
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In Organic Chemistry (the chemistry of carbon) the so-called 'HETEROATOMS', atoms other than carbon and hydrogen, are very abundant and usually bear lone electron pairs. Consequently, organic molecules usually contain many functional groups with electrons that can be shared with incoming protons, and thus show BASIC properties.
Functional groups bearing unshared electrons can accept protons and behave as more or less strong bases. Here are some examples:
The expression 'more or less strong base' is vague and of little scientific value. Is there any way to measure basicity and give it a number?
The pKb value is the numerical measure of basicity and is related to how far the acid-base equilibrium of a compound with water is shifted.
A:(⊖) + H2O Kb AH(⊕)conjugate acid + HO⊖
Kb = [AH(⊕)] [HO⊖][A:(⊖)]
pKb = − log Kb
The pKa and pKb values are related through the equilibrium constant of water with itself: the ionic product of water. Therefore, from the pKa value of an acid one can easily obtain the pKb value of its conjugate base, and vice versa.
A:(⊖) + H2O Kb AH(⊕)conjugate acid + HO⊖
AH(⊕) + H2O Ka A:(⊖)conjugate base + H3O⊕
H2O + H2O K = 10−14 H3O⊕ + HO⊖
pKa + pKb = 14
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What are the pKb values of the types of organic compounds mentioned above? Within each type the pKb values span a range that depends on each particular structure. We will address this in detail when we study each functional group but, for the time being, let us take a first look at pKb values.
Amines are the strongest bases in Organic Chemistry. Their pKb values are in the range 3-5. Their conjugate acids, ammonium salts, display an acidity (pKa = 9-11) similar to that of phenols. This is just a coincidence, because ammonium salts and phenols of course have nothing to do with one another.
Anilines are structurally very close to amines, but notice that the N atom is bonded to an aromatic ring. And that makes a big difference, because anilines are far less basic than amines. How many times less basic than amines are anilines, on average? Wrong! It is not 6 times but 10 to the power of 6 times. Remember that the definition of pKb involves a logarithm!!! Please note that the conjugate acids of anilines, anilinium salts, are as acidic as carboxylic acids.
Pyridines display a basicity similar to that of anilines.
Pyrroles, in contrast, although apparently closely related to pyridines, display an extremely low basicity. How can that be? We will fully solve the mystery when we study 'AROMATICITY'. Put simply, the N atom of pyridine bears a true unshared electron pair in an sp2 orbital, whereas the N atom of pyrrole has its 'apparent' unshared electron pair in a p orbital that overlaps strongly with the neighbouring double bonds. In the N atom of pyrrole, the lone electron pair is far, far less available than in pyridine.
Oxygen atoms bear two lone electron pairs, and the protons of a strong enough acid can bond to them. When OH or C=O groups are 'protonated', the resulting molecules are quite reactive. Many reactions of alcohols and ketones begin with their 'protonation' by a strong acid. We will see this in detail when we study these functional groups in depth.
The stronger the basicity of a compound, the more its equilibrium is shifted to the right. This suggests that the stronger the basicity, the more accessible the lone electron pair of the base. The more accessible the lone electron pair of the base, the more the equilibrium is shifted to the right.
In amines, the lone electron pair
is completely localized on the N atom, which has a moderate electronegativity. That explains their strong basicity, usually the highest in Organic Chemistry. They are 100% protonated by mineral acids. Anilines experimentally display a lower basicity than amines, which suggests that the lone electron pair on the N atom is less accessible in anilines than in amines. We can easily understand this with resonance forms in which the lone electron pair is delocalized towards the aromatic ring.
The N atom of pyridines CANNOT
delocalize its lone electron pair, because the possible resonance form would produce a kind of linear allenic structure that is impossible within a six-membered ring. Therefore the lone electron pair is available to an incoming proton and pyridines are basic compounds. In contrast, in pyrroles the lone electron pair can be delocalized towards the ring and its availability to an acid is extremely low. Both pyridines and pyrroles are 'aromatic' (synonymous with additional stability), but the big difference between them is that pyrroles DO need the lone electron pair to achieve it, whereas pyridines DO NOT. This will be fully explained in the AROMATICITY chapter. For the time being, please note that resonance forms give us a reasonable explanation for the very different basicity of pyridines and pyrroles.
The lone electron pairs on O atoms are weakly basic because they lie on a strongly electronegative atom. But with a sufficiently strong acid, the O atoms of a certain number of molecules can be 'protonated'. The few resulting protonated molecules are very reactive and can give rise to new products. The equilibrium will then regenerate more protonated molecules, which will react in turn, and so on. In many cases, the first step in the reaction of alcohols or carbonyl (C=O) compounds is their protonation.
