Electrophiles are neutral or positively charged species with an empty atomic orbital (such as the empty p orbital in borane) or a low-energy antibonding orbital that can easily accept electrons. The simplest electrophile is the hydrogen cation, H+, usually named for what it is, a proton. H+ is a species without any electrons at all and a vacant, very low energy, 1s orbital. It is so reactive that it is hardly ever found and almost any nucleophile will react with it. Acid solu tions containing H+ are neutralized by the nucleophile hydroxide, for example, and strong acid goes on to protonate water as well, the water acting as a nucleophile and the proton as the electrophile. The product is the hydronium ion, H3O+, the true acidic species in all aqueous strong acids. Here’s the reaction between hydroxide and H+ with the electron movement from the nucleophile to the electrophile represented by curly arrows. The arrows start on the hydroxide’s negative charge, which represents one of the oxygen’s pairs of electrons:

Other electrophiles with empty atomic orbitals include borane, which you met on p. 103, and related compounds such as boron trifluoride and aluminium trichloride. BF₃ reacts with ethers, as shown below, to form stable complexes. This time the arrow starts on the lone pair.

Few organic compounds have vacant atomic orbitals and in most organic electrophiles the LUMOs are instead low-energy antibonding orbitals associated with electronegative atoms. These antibonding orbitals can be either π* orbitals or σ* orbitals—in other words, mol ecules which make good electrophiles might have a double or a single bond to an electronegative atom such as O, N, Cl, or Br. It’s important that an electronegative atom is involved in order to lower the energy of the orbital (see p. 96) and make it ready to accept electrons.

Carbon's place in the electronegativity scale

 Here is a summary of electronegativities for atoms commonly involved in organic reactions.

This bar chart makes it clear why carbon is just so special: it can form strong bonds to almost anything, especially itself. Elements at either end of the scale form weak bonds to similar elements (metal–metal bonds are weak, as are halogen or O–O bonds), but elements in the middle can form strong bonds to other elements at either end of the scale or elements in the middle. Being in the middle also gives C versatile reactivity: it is electrophilic when bonded to a more electronegative element and nucleophilic when bonded to a more nucleophilic element.

The most important molecules with a double bond to an electronegative atom are carbonyl compounds. In fact, carbonyl groups are the most important functional groups in organic chemistry. We looked at their orbitals on p. 103 and we devote the next chapter, Chapter 6, to a detailed study of their reactivity. The low-energy π* orbital is available to accept electrons, and its electrophilicity is further enhanced by the partial positive charge at carbon which arises from the C=O dipole. Here’s an example of a carbonyl compound, acetone, reacting with an anionic nucleophile—we’ll choose borohydride in this case. Notice   how the arrow does not start on the negative charge, as the charge does not represent a pair of electrons here.

The arrows showing electron movement are a little more involved this time, but the explanation is straightforward. The fi rst arrow shows the electrons moving from the nucleophile’s HOMO (the B–H σ orbital) to the electrophile’s LUMO (the C=O π* orbital). The new feature in this mechanism is a second arrow showing the electrons moving from the double bond onto the oxygen atom. This is easy to explain. Since the reaction is putting electrons into an antibonding orbital (the π*), a bond has to break. That breaking bond is the C=O π bond (the σ bond remains intact). The electrons in the bond have to go somewhere and they end up as an extra lone pair (represented by the negative charge) on oxygen. The product has a new C–H σ bond in place of the C=O π bond. Molecules with a single bond to electronegative atoms can also make good electrophiles. In compounds such as HCl or CH3Br, the σ* orbital is low in energy because of the electro negative Cl or Br (see p. 95) and the dipole attracts the electrons of the nucleophile to the H or C atom. Here’s an example of hydrogen chloride acting as an electrophile with ammonia as the nucleophile. As with the carbonyl example above, we are putting electrons into an antibonding orbital, so a bond must break. This time the antibonding orbital is the H–Cl σ*, so the bond which breaks is the H–Cl σ bond.

Some σ bonds are electrophilic even though they have no dipole at all. The bonds in the halogens I2, Br2, and Cl2 are a case in point. Bromine, for example, is strongly electrophilic because it has a weak Br–Br bond with a low energy σ* orbital. Why is the σ* low in energy? Well, bromine is slightly electronegative, but it is also large: it has to use 4s and 4p atomic orbitals for bonding, but these orbitals are large and diffuse, and overlap poorly, meaning the σ* molecular orbital is not raised far in energy and can easily accept electrons. How different the situation is with a C–C bond: C–C single bonds are almost never electrophilic.

Bromine reacts with many nucleophiles, for example in the reaction shown below between a sulfide and bromine. Lone pair electrons are donated from sulfur into the Br–Br σ* orbital, which makes a new bond between S and Br, and breaks the old Br–Br bond.

● Electrophiles accept electrons into empty low-energy orbitals represented by one of the following: