Nucleophiles and electrophiles: Neutral inorganic species
Polar bonds
Polar bonds are polarized such that the more electronegative atom has the greater share of the bonding electrons. As a result, it will be electron rich and will be a nucleophilic center. The less electronegative atom will be elec- tron deficient and will be an electrophilic center. Electronegative atoms are to the right of the periodic table and have lone pairs of electrons.
Nucleophilic strength
The relative strengths of neutral nucleophilic centers are determined by how well they can accommodate a positive charge. The more electronega- tive the atom, the less nucleophilic it will be. Therefore, nitrogen is more nucleophilic than oxygen, and oxygen is more nucleophilic than fluorine.
Electrophilic strength
The relative electrophilic strengths of hydrogen atoms in different mole- cules is determined by the stability of the ions formed. Hydrogen atoms attached to nitrogen are only weakly electrophilic, whereas hydrogens attached to halogen atoms are strongly electrophilic. Therefore, theelectrophilic strength of hydrogen atoms depends on the electronegativity of the neighboring atom.
Properties
It is possible to predict whether a molecule is likely to react as an elec- trophile or as a nucleophile, based on the strength of the nucleophilic or electrophilic centers present.
Polar bonds
If two atoms of quite different electronegativities are linked together, then the bond connecting them will be polar covalent such that the bonding electrons are biased towards the more electronegative atom. This will give the latter a slightly negative charge and make it a nucleophilic center. Conversely, the less electronegative atom will gain a slightly positive charge and be an electrophilic center (Fig. 1). The further to the right an element is in the periodic table, the more electronegative it is. Thus, fluorine is more electronegative than oxygen, which in turn is more electronegative than nitrogen. Note also that all the nucleophilic atoms identified above have lone pairs of electrons. This is another way of identifying nucleophilic atoms.
Nucleophilic strength
The molecules above have both nucleophilic and electrophilic centers and could react as nucleophiles or as electrophiles. However, it is usually found that there is a preference to react as one rather than the other. This is explained by considering the relative strengths of nucleophilic and electrophilic centers. First of all, let us consider the relative strengths of nucleophilic centers by comparing N, O, and F. If we compare the relative positions of these atoms in the periodic table, we find that fluorine is more electronegative than oxygen, which in turn is more electro-negative than nitrogen. However, when we compare the nucleophilic strengths of these atoms, we find that the nitrogen is more nucleophilic than oxygen, which in turn is more nucleophilic than fluorine.
The relative nucleophilic strengths of these atoms is explained by looking at the products which would be formed if these atoms were to act as nucleophiles. Let us compare the three molecules HF, H2O, and NH3 and see what happens if they were to form a bond to a proton (Fig. 2). Since the proton has no electrons, both electrons for the new bond must come from the nucleophilic centers (i.e. the F, O, and N). As a result, these atoms will gain a positive charge. If hydrogen fluoride acts as a nucleophile, then the fluorine atom gains a positive charge. Since the fluorine atom is strongly electronegative, it does not tolerate a positive charge. Therefore, this reaction does not take place. Oxygen is less electronegative and is able to tolerate the positive charge slightly better, such that an equilibrium is pos-sible between the charged and uncharged species. Nitrogen is the least elec-tronegative of the three atoms and tolerates the positive charge so well that the reaction is irreversible and a salt is formed.
Thus, nitrogen is strongly nucleophilic and will usually react as such, whereas halogens are weakly nucleophilic and will rarely react as such.
Lastly, it is worth noting that all these molecules are weaker nucleophiles than their corresponding anions, i.e. HF, H2O, and NH3 are weaker nucleophiles than F -, OH- and NH-2 respectively.
Electrophilic strength
The same argument can be used in reverse when looking at the relative electrophilic strengths of atoms in different molecules. Let us compare the elec-trophilic strengths of the hydrogens in HF, H2O, and NH3. In this case, reaction with a strong nucleophile or base would generate anions (Fig. 3). Fluorine being the most electronegative atom is best able to stabilize a negative charge and so the fluoride ion is the most stable ion of the three. Oxygen is also able to stabilize a negative charge, though not as well as fluorine.
Nitrogen is the least electronega-tive of the three atoms and has the least stabilizing influence on a negative charge and so the NH2 ion is unstable. The more stable the anion, the more easily it is formed and hence the hydrogen which is lost will be strongly electrophilic. This is the case for HF. In contrast, the hydrogen in ammonia is a very weak electrophilic center since the anion formed is unstable. As a result, nitrogen anions are only formed with very strong bases.
Properties
It is possible to predict whether molecules are more likely to react as nucleophiles or electrophiles depending on the strength of the nucleophilic and electrophilic centers present. For example, ammonia has both electrophilic and nucleophilic centers. However, it usually reacts as a nucleophile since the nitrogen atom is a strong nucleophilic center and the hydrogen atom is a weak electrophilic center. By contrast, molecules such as hydrogen fluoride or aluminum chloride prefer to react as electrophiles. This is because the nucleophilic centers in both these molecules (halogen atoms) are weak, whereas the electrophilic centers (H or Al) are strong. Water is a molecule which can react equally well as a nucleophile or as an electrophile. For example, water reacts as a nucleophile with a proton and as an electrophile with an anion (Fig. 4).
