The reactions of amines with acids

The reactions of amines with acids These are most easily considered using the Bronsted-Lowry theory of acids and bases  - the base is a hydrogen ion acceptor. We'll do a straight comparison between amines  and the familiar ammonia reactions. A reminder about the ammonia reactions Ammonia reacts with acids to produce ammonium ions. The ammonia molecule picks up  a hydrogen ion from the acid and attaches it to the lone pair on the nitrogen. If the reaction is in solution in water (using a dilute acid), the ammonia takes a hydrogen  ion (a proton) from a hydroxonium ion. (Remember that hydrogen ions present in solutions  of acids in water are carried on water molecules as hydroxonium ions, H3O+.) If the acid was hydrochloric acid, for example, you would end up with a solution containing  ammonium chloride - the chloride ions, of course, coming from the hydrochloric acid. You could also write this last equation as: . . . but if you do it this way, you must include the state symbols. If you write H+ on its own,  it implies an unattached hydrogen ion - a proton. Such things don't exist on their own in  solution in water. If the reaction is happening in the gas state, the ammonia accepts a proton directly from  the hydrogen chloride: This time you produce clouds of white solid ammonium chloride. The corresponding reactions with amines The nitrogen lone pair behaves exactly the same. The fact that one (or more) of the  hydrogens in the ammonia has been replaced by a hydrocarbon group makes no  difference. For example, with ethylamine: If the reaction is done in solution, the amine takes a hydrogen ion from a hydroxonium  ion and forms an ethylammonium ion. Or: The solution would contain ethylammonium chloride or sulphate or whatever. Alternatively, the amine will react with hydrogen chloride in the gas state to produce the  same sort of white smoke as ammonia did - but this time of ethylammonium chloride. These examples have involved a primary amine. It would make no real difference if you  used a secondary or tertiary one. The equations would just look more complicated. The product ions from diethylamine and triethylamine would be diethylammonium ions and triethylammonium ions respectively. The reactions of amines with water Again, it is easiest to use the Bronsted-Lowry theory and, again, it is useful to do a straight  comparison with ammonia. A reminder about the ammonia reaction with water Ammonia is a weak base and takes a hydrogen ion from a water molecule to produce  ammonium ions and hydroxide ions. However, the ammonia is only a weak base, and doesn't hang on to the hydrogen ion very  successfully. The reaction is reversible, with the great majority of the ammonia at any one  time present as free ammonia rather than ammonium ions. The presence of the hydroxide ions from this reaction makes the solution alkaline. The corresponding reaction with amines The amine still contains the nitrogen lone pair, and does exactly the same thing. For example, with ethylamine, you get ethylammonium ions and hydroxide ions produced. There is, however, a difference in the position of equilibrium. Amines are usually stronger  bases than ammonia. (There are exceptions to this, though - particularly if the amine group  is attached directly to a benzene ring.)
The reactions of amines with copper(II) ions Just like ammonia, amines react with copper(II) ions in two separate stages. In the first step, we can go on using the Bronsted-Lowry theory (that a base is a hydrogen ion acceptor). The second  stage of the reaction can only be explained in terms of the Lewis theory (that a base is an  electron pair donor).