Nucleophiles and electrophiles

In this lesson, we will learn:

• To recall the meaning of the terms nucleophile and electrophile.
• To understand the complementary nature of nucleophiles and electrophiles in driving organic reactions.
• To explain the reactivity of nucleophiles with electrophiles using molecular orbital theory.

• Recalling How organic reactions occur?
  • Organic reactions involve movement of electrons which breaks bonds in the reactant(s) to make new bonds to form the product(s).
  • Interactions between opposite charges, and frontier orbitals (an occupied MO of one molecule and an empty MO of the other) both drive this process.
  • In most organic reactions, there is a mix of the two.
  
  
• In organic reactions, there are definitions based on if electrons are being donated$\,$ by or being accepted$\,$ by a molecule:
  • A nucleophile$\,$ (labelled Nu / Nu: / Nu- in reaction mechanisms) is an electron rich molecule that donates electrons, usually to an electron deficient molecule.
  • An electrophile$\,$ (labelled E / E⁺ in reaction mechanisms) is an electron poor molecule that accepts electrons.
  Using these, any organic reaction can be thought of as a nucleophile ‘attacking’ an electrophile and forming a new bond.
  
  
• Because organic reactions always involve breaking bonds and forming others, which is caused by electrons moving, which is caused by nucleophiles interacting with electrophiles, identifying nucleophiles and electrophiles is INCREDIBLY important in predicting organic reactions.
  
  For a simple example, in the reaction with H⁺ and OH^- there is a movement of electrons from the hydroxide ion to the H⁺ ion to form H₂O. The OH^- ion is the nucleophile as this is the electron rich molecule that donated electrons, while the H⁺ ion accepted them. The movement of a pair of electrons is shown by a curly arrow.
  
  See the diagram below.
  
  
• Identifying a good nucleophile can be done with the question:
  • How available are electrons in the nucleophile?
  We can use the idea of electronegativity to find this out.
  For example, compare ammonia (NH₃) and water (H₂O). Both molecules can act as a nucleophile by forming bonds using their lone pair, but one is better than the other:
  • NH₃ has a lone pair on nitrogen, which is less electronegative than oxygen.
  • H₂O has two lone pairs on oxygen, which is more electronegative than nitrogen and holds these electron pairs more tightly.
  Because nitrogen holds its lone pair less tightly than oxygen, it is better at forming new bonds with it – at donating the electron pair$\,$ to another molecule. Ammonia is therefore a better nucleophile than water.
  
  
• The strength of a nucleophile can be explained by molecular orbital theory.
  The more electronegative oxygen atom in water makes all the filled molecular orbitals, including the lone pair, lower energy than the filled molecular orbitals in ammonia, where nitrogen does not attract its electrons as strongly.
  This makes ammonia’s highest occupied molecular orbital (HOMO) higher in energy. A higher energy HOMO will have stronger interactions with an electrophile.
  
  
• The strength of an electrophile asks the opposite question:
  • How electron-deficient is the electrophile?
  The greater the positive charge, the stronger a potential bonding interaction with an electron-rich nucleophile would be. A vacant p orbital (e.g. boron) also makes a molecule electrophilic.
  Regardless, the nucleophile needs an electrophile’s empty orbital to interact with. In any molecule, the empty orbitals are empty because they are high in energy; the lowest unoccupied molecular orbital (LUMO) interacts most with the nucleophile. A lower energy LUMO will have stronger interactions with the HOMO of a nucleophile.
  
  
• The reactivity of two molecules, an electrophile and a nucleophile, can be reduced down to the HOMO-LUMO gap.
  • The smaller the gap, the greater the ‘splitting of levels’ – a bonding interaction between the molecules will be lower energy, and antibonding interactions higher.
  • The greater the gap, the less interaction between the orbitals. There will be less and less, and eventually no splitting of levels, and no bonding or antibonding interactions – no bond can or will form.
  
  See the diagram below.