Acids and bases - review
Definitions of acids and bases
Acids and bases – review
Definition by Svante Arrhenius:
• An acid contains H that ionises in water to give H+ ion; • A base contains OH that ionises in water to give OH– ion.
Note that this is fine for regular acids and bases, but excludes other common ones, such as ammonia and pyridine. Brønsted-Lowry definition:
• An acid is a proton donor; • A base is a proton acceptor.
This implies that an acid-base neutralisation reaction is a proton transfer reaction.
Some Brønsted acids and bases
The following are diagrams of strong acids. They are, from left to right, nitric acid, sulfuric acid, perchloric acid and hydrochloric acid. These all undergo complete H+ dissociation in water.
The following are diagrams of weak acids. They are, from left to right, carbonic, acetic acid, sulfurous acid and hydrofluoric acid. These only partially dissociate in water.
The following are diagrams of strong bases. They are, from left to right, sodium hydroxide, carbonate, phosphate and sulfide. They all are proton acceptors.
The following are diagrams of weak bases. They are, from left to right, ammonia, pyridine, H2PO4– and HSO4–.
Note that acid-base reactions do not just occur in water; they can also occur in gas or other liquid or solid states. For instance, opening a bottle of hydrochloric acid next to a bottle of ammonia will cause vapours from both substances to escape and react, forming NH4Cl(s). Conjugate acid-base pairs
For Brønsted-Lowry acid-base reactions, both the forward and reverse reactions are acid-base reactions. There are always two sets of species on either side of the equation that differ only by one proton. These are called conjugate acid-base pairs. In the equation below the pairs are of the same colour. base
acid
HCl + H2O ⇌ H3O + Cl–
Examples
acid
base
HF (acid) + H2O (base) ⇌ F– (conjugate base) + H3O+ (conjugate acid) CH3CO2H (acid) + CN– (base) ⇌ CH3CO2– (conjugate base) + HCN (conjugate acid) HPO42– (acid) + SO32– (base) ⇌ PO43– (conjugate base) + HSO3– (conjugate acid) Lewis acids and bases
There is another definition of acids and bases, as proposed by Gilbert N Lewis. This states that a Lewis acids is an electron-pair acceptor and a Lewis base is an electron-pair donor. A reaction of a Lewis acid with a Lewis base gives a compound called an adduct.
The strength of Lewis acids and bases is not as readily quantifiable as their Brønsted-Lowry counterparts. The nature of H+(aq)
Firstly, consider that H+ is simply a bare proton. These bare protons will not exist in water. This is because in any compound, some of the acids will have lone pairs. A bare proton will attach to one of the lone pairs. As a consequence, H+(aq) cannot exist in solution; it is always solvated. The simplest species involving this is H(H2O)5+, as shown. It is a H3O+ cation surrounded by four water molecules, forming hydrogen bonds.
Acids and bases – review
Definition by Svante Arrhenius:
• An acid contains H that ionises in water to give H+ ion; • A base contains OH that ionises in water to give OH– ion.
Note that this is fine for regular acids and bases, but excludes other common ones, such as ammonia and pyridine. Brønsted-Lowry definition:
• An acid is a proton donor; • A base is a proton acceptor.
This implies that an acid-base neutralisation reaction is a proton transfer reaction.
Some Brønsted acids and bases
The following are diagrams of strong acids. They are, from left to right, nitric acid, sulfuric acid, perchloric acid and hydrochloric acid. These all undergo complete H+ dissociation in water.
The following are diagrams of weak acids. They are, from left to right, carbonic, acetic acid, sulfurous acid and hydrofluoric acid. These only partially dissociate in water.
The following are diagrams of strong bases. They are, from left to right, sodium hydroxide, carbonate, phosphate and sulfide. They all are proton acceptors.
The following are diagrams of weak bases. They are, from left to right, ammonia, pyridine, H2PO4– and HSO4–.
Note that acid-base reactions do not just occur in water; they can also occur in gas or other liquid or solid states. For instance, opening a bottle of hydrochloric acid next to a bottle of ammonia will cause vapours from both substances to escape and react, forming NH4Cl(s). Conjugate acid-base pairs
For Brønsted-Lowry acid-base reactions, both the forward and reverse reactions are acid-base reactions. There are always two sets of species on either side of the equation that differ only by one proton. These are called conjugate acid-base pairs. In the equation below the pairs are of the same colour. base
acid
HCl + H2O ⇌ H3O + Cl–
Examples
acid
base
HF (acid) + H2O (base) ⇌ F– (conjugate base) + H3O+ (conjugate acid) CH3CO2H (acid) + CN– (base) ⇌ CH3CO2– (conjugate base) + HCN (conjugate acid) HPO42– (acid) + SO32– (base) ⇌ PO43– (conjugate base) + HSO3– (conjugate acid) Lewis acids and bases
There is another definition of acids and bases, as proposed by Gilbert N Lewis. This states that a Lewis acids is an electron-pair acceptor and a Lewis base is an electron-pair donor. A reaction of a Lewis acid with a Lewis base gives a compound called an adduct.
The strength of Lewis acids and bases is not as readily quantifiable as their Brønsted-Lowry counterparts. The nature of H+(aq)
Firstly, consider that H+ is simply a bare proton. These bare protons will not exist in water. This is because in any compound, some of the acids will have lone pairs. A bare proton will attach to one of the lone pairs. As a consequence, H+(aq) cannot exist in solution; it is always solvated. The simplest species involving this is H(H2O)5+, as shown. It is a H3O+ cation surrounded by four water molecules, forming hydrogen bonds.
