Theories of acids and bases

 Example: Example HF + NH 3 ↔ NH F - The HF transfers a H + to the NH 3, so it acts as an acid; the NH 3 accepts the H + so it acts as a base.

 In Bronsted-Lowry theory, an acid can only behave as a proton donor if there is also a base present to act as a proton acceptor.  Consider the acid-base reaction between a generic acid HA and base B: HA + B ↔ A - + BH + We can see that HA acts as an acid, while B acts as a base. But, if we look at the reverse reaction, BH + acts as an acid (donates a proton) and A - acts as a base (accepts a proton). So, acids react to form bases and bases react to form acids. Acid- base pairs that are related to each other in this way are called conjugate acid-base pairs. Conjugate acid-base pairs differ by one proton.

 One example of a conjugate pair is H 2 O and H 3 O +, which is found in all acid-base reactions in aqueous solution.  H 2 O + H + ↔ H 3 O + -- this reaction takes place when an acid dissolves in water. H 3 O + is called the hydroxonium ion, the oxonium ion, or the hydronium ion.  For most reactions it is convenient to simply write H + (aq).  So, H 3 O + is the conjugate acid and H 2 O is the conjugate base

 Water can act as a base or as an acid, depending what it is reacting with.

Substances that can act as acids or bases are said to be amphoteric or amphiprotic.  To act as a Bronsted-Lowry acid, substances must be able to dissociate and release H+.  To act as a Bronsted-Lowry base, substances must be able to accept H+, which means, they must have a lone pair of electrons.  Example: HCO 3 - acts as acid and base  HCO 3 - (aq) + H 2 O(l) ↔ CO 3 2- (aq) + H 3 O + (aq)  HCO 3 - (aq) + H 2 O(l) ↔ H 2 CO 3 (aq) + OH - (aq)

 Lewis acids and Bases

 The Lewis definition of acids is broader than the Bronsted-Lowry theory.  By the Lewis definition, an acid is any species capable of accepting a lone pair of electrons (no longer restricted to just H + ).  Lewis acid-base reactions result in the formation of a covalent bond which will always be a dative bond, because both electrons come from the base.  For example: For example  BF 3 has an incomplete octet, so it is able to act as a Lewis acid by accepting a pair of electrons; NH 3 acts as a Lewis base by donating its lone pair of electrons. The arrow shows the bond is a coordinate covalent bond (dative bond)

 Other good examples of Lewis acid-base reactions are found in the chemistry of transition elements.  Transition metals often form ions with vacant orbitals. Thus they are able to act as Lewis acids and accept lone pairs of electrons.  The Lewis base, which donates the lone pair, is called a ligand, and usually surrounds the ion in a fixed number ratio.  Dative bonds form between the metal ion and the ligands, resulting in a complex ion that has a characteristic color.  Typical ligands found in complex ions include H 2 O, CN - and NH 3. Note that these all possess lone pairs of electrons.

TheoryDefinition of acidDefinition of base Bronsted-LowryProton donorProton acceptor LewisElectron pair acceptorElectron pair donor Although all Bronsted-Lowry acids are Lewis acids, not all Lewis acids are Bronsted-Lowry acids, so the term Lewis acids are usually reserved for those species which can only be described by Lewis theory (that is those that do not release H + ).