Embed Size (px)
Citation preview
The nitrogen atom in an ammonia molecule has a lone pair of electrons.
The nitrogen atom in an ammonia molecule has a lone pair of electrons.
It can make a dative covalent bond with a hydrogen ion (a proton), forming an ammonium ion.
It can make a dative covalent bond with a hydrogen ion (a proton), forming an ammonium ion.
Ammonia is a Brønsted–Lowry base because it can accept protons.
The nitrogen atom in a methylamine molecule has a lone pair of electrons.
It can make a dative covalent bond with a proton, forming a methylammonium ion.
It can make a dative covalent bond with a proton, forming a methylammonium ion.
Methylamine is a Brønsted–Lowry base because it can accept protons.
Amines are Brønsted–Lowry bases because they can accept a proton.
When amines dissolve in water, they form alkaline solutions due to the presence of hydroxide ions.
Amines dissolve in dilute acids to form soluble ionic salts.
Phenylamine is almost insoluble in water.
But it dissolves in excess dilute hydrochloric acid.
The salt phenylammonium chloride is formed in the reaction.
Ammonia is a weak base.
It has a pKa value of 9.25 – the higher the value, the stronger the base.
These primary aliphatic amines have higher pKa values than ammonia.
These primary aliphatic amines have higher pKa values than ammonia.
These primary aliphatic amines have higher pKa values than ammonia.
They are stronger bases than ammonia. Why?
They are stronger bases than ammonia. Why?
The base strength of primary amines depends upon the availability of the lone pair of electrons on the nitrogen atom, which forms a bond with an H+ ion.
The base strength of primary amines depends upon the availability of the lone pair of electrons on the nitrogen atom, which forms a bond with an H+ ion.
The base strength of primary amines depends upon the availability of the lone pair of electrons on the nitrogen atom, which forms a bond with an H+ ion.
These amines contain alkyl groups.
These amines contain alkyl groups.
The alkyl groups are electron-releasing relative to nitrogen.
The alkyl groups are electron-releasing relative to nitrogen.
The alkyl groups are electron-releasing relative to nitrogen.
They increase the availability of the lone pair of electrons on the nitrogen atom, so primary aliphatic amines are stronger bases than ammonia.
They increase the availability of the lone pair of electrons on the nitrogen atom, so primary aliphatic amines are stronger bases than ammonia.
Primary aromatic amines are weaker bases than ammonia.
Ammonia has a pKa value of 9.25 but phenylamine has a pKa value of 4.62. Why?
The answer lies in the delocalised rings of electrons in the benzene ring and the lone pair of electrons on the nitrogen atom (seen here as the shape of a p orbital).
The answer lies in the delocalised rings of electrons in the benzene ring and the lone pair of electrons on the nitrogen atom (seen here as the shape of a p orbital).
The answer lies in the delocalised rings of electrons in the benzene ring and the lone pair of electrons on the nitrogen atom (seen here as the shape of a p orbital).
The answer lies in the delocalised rings of electrons in the benzene ring and the lone pair of electrons on the nitrogen atom (seen here as the shape of a p orbital).
The answer lies in the delocalised rings of electrons in the benzene ring and the lone pair of electrons on the nitrogen atom (seen here as the shape of a p orbital).
The answer lies in the delocalised rings of electrons in the benzene ring and the lone pair of electrons on the nitrogen atom (seen here as the shape of a p orbital).
The lone pair of electrons on the nitrogen atom interacts with these delocalised rings of electrons. The lone pair becomes less available for bonding to an H+ ion.
The lone pair of electrons on the nitrogen atom interacts with these delocalised rings of electrons. The lone pair becomes less available for bonding to an H+ ion.
The lone pair of electrons on the nitrogen atom interacts with these delocalised rings of electrons. The lone pair becomes less available for bonding to an H+ ion.
