In chemistry, there are three definitions of bases, namely, Arrhenius bases, Bronsted bases, and Lewis bases. All these definitions establish that bases are the substances that react with acids, which was proposed originally by G.-F. Rouelle in the mid-18th century.
History of Base
In 1884, Arrhenius proposed that a base is a substance that dissociates in an aqueous solution to produce hydroxide ions (OH−). These hydroxide ions can react with the hydrogen ions (H+, as per Arrhenius) from the dissociation of acids to produce water in an acid-base reaction. Therefore, a base is a metal hydroxide such as Ca(OH)2 or NaOH. Such types of aqueous hydroxide solutions were also defined by particular characteristic properties. They taste bitter, slippery to the touch, and change the colour of pH indicators (for example, it turns red litmus paper blue).
Properties of Bases
A few general properties of bases can be listed as follows.
The strong or concentrated bases are caustic on organic matter, and they react violently with the acidic substances.
Molten bases or the aqueous solutions dissociate in ions and conduct electricity.
The pH value of a basic solution at standard conditions is always greater than 7.
Reactions with the indicators: Bases turn a red litmus paper into blue, phenolphthalein pink, and keep bromothymol blue in its natural colour of blue, and then turn into methyl orange-yellow.
Bases are the bitter compounds.
Reactions Between Bases and Water
The reaction, below given, represents the general reaction between a base (B) and water to form a conjugate acid (BH+) and a conjugate base (OH-). The chemical reaction for this can be represented as follows:
B(aq) + H2O(l) ⇌ BH+(aq) + OH−(aq)
Kb, the equilibrium constant, for this reaction is estimated using the general equation, given below:
Kb = [BH+][OH−]/[B]
In the above equation, the base (B) and the extremely strong base (which is the conjugate base OH−) compete for protons. Resultantly, bases, which react with water, hold relatively small equilibrium constant values. Also, the base is weaker when it contains a lower equilibrium constant value.
Types of Bases
A strong base is defined as a basic chemical compound that can remove a proton (H+) from (or the deprotonate) a molecule of even a very weak acid (like water) in the acid-base reaction. The common examples of strong bases are given as hydroxides alkaline earth metals and alkali metals, such as Ca(OH)2 and NaOH, respectively. Because of their low solubility, a few of the bases, like alkaline earth hydroxides, can be used when the solubility constraint is not considered. An advantage of this low solubility is "several antacids were suspensions of the metal hydroxides like magnesium hydroxide and aluminium hydroxide."
Some strong bases are Sodium hydroxide - NaOH, Lithium hydroxide - LiOH, Potassium hydroxide - KOH.
Group 1 salts of hydrides, amides, and carbanions tend to be stronger bases because of their conjugate acid’s extreme weakness, which are amines, stable hydrocarbons, dihydrogen. In general, these bases are created by adding pure alkali metals like sodium into conjugate acid. They are known as superbases, and it is quite impossible to keep them in water because they are defined as the stronger bases compared to the hydroxide ion. So, they deprotonate conjugate acid water. For suppose, the ethoxide ion (which is the conjugate base of ethanol) in the presence of water undergoes the reaction given below:
CH3CH2O− + H2O → CH3CH2OH + OH−
Some examples of common superbases can be given as Sodium hydride (NaH), Sodium amide (NaNH2), and Butyl lithium (n-C4H9Li).
A weak base is defined as the one which does not fully ionize in an aqueous solution, or in the one, whose protonation is incomplete. For suppose, the ammonia compound transfers a proton to water as per the equation given below:
NH3(aq)+H2O ⇆ NH4++OH-(aq)
The equilibrium constant for the above-given reaction at 25 °C is given as 1.8 x 10−5 so that the extent of degree of ionization or reaction is quite small.
A Lewis base or the electron-pair donor is defined as a molecule having a high-energy pair of electrons that can be shared with the low-energy vacant orbital in an acceptor molecule to produce an adduct. In addition to the H+ molecule, the possible acceptors (which are called the Lewis acids) include neutral molecules like BF3 and metal ions like Fe3+ or Ag+. Whereas, adducts involving the metal ions are generally described as coordination complexes.
As per the original Lewis formulation, when a neutral base produces a bond with a neutral acid, an electric stress condition takes place. Then, acid and base share the electron pair that only belonged to the base formerly. As a result, there creates a high dipole moment, which can only be destroyed by the molecule rearrangement.