91Ó°ÊÓ

Bases are defined in three different ways:

1. Arrhenius Bases: When we dissolve any substance in water and if it will dissociate and release hydroxide ions that is known as Arrhenius base.
2. Bronsted Bases: Bronsted- Lowery has defined bases as the species which accepts protons (${\mathrm{H}}^{+}$)
3. Lewis Bases: Lewis bases are defined as those species which can donate its lone pair of electrons to other species which are electron deficient.

Some of the most common examples of Arrhenius bases are: Sodium hydroxide, Potassium hydroxide, magnesium hydroxide, calcium hydroxide etc.

Examples of Bronsted bases are: Ammonia, dimethyl ether, Pyridine, THF etc.

Species which act as Lewis bases are: hydroxide ion, water, sulfate, ethane, benzene etc.

For a base to deprotonate (removal of proton) from any acidic compound it should be a stronger base and its conjugate acid should be weaker so that the conjugate acid of that base will not lose its protons and then base will remove proton from other acid. Acidic strength of any species depends on itsvalue. If any compound will have highervalue its acidic strength will be less i.e. it will be a weak acid and vice- versa.

Conjugate acid of KOH (potassium hydroxide) is water i.e. and its value is 15 and value of 1- butyne is 25 and as we know that lower the value, stronger will be the acid and conjugate acid of KOH is stronger acid therefore, base will not deprotonate 1- butyne.