Ionic equilibrium: Ionization of Acids and Bases

The state that exists when acidic and basic ions in solution exactly neutralize each other is called ionization acid-base equilibrium. 

Consider an acid, HA, in water. According to the Brønsted-Lowry acid/base theory, the acid should protonate the water to form hydronium and the conjugate base, A–. There will be an equilibrium between the acid and water, hydronium and the conjugate base.    This equilibrium can be used to calculate the concentrations of species in the solution.

Acids are classified as either strong or weak acids based on their ionization in water.

A strong acid is an acid that is completely ionized in an aqueous solution. Hydrogen chloride (HCl) ionizes completely into hydrogen ions and chloride ions in water.

A weak acid is an acid that partially ionizes in an aqueous solution. Carbonic acid(H2CO3)  ionizes partially into hydrogen ions and carbonate ions in water.

Solubility Product

Ionization of acids and bases 

Bases are those compounds that furnish the hydroxyl ions in the aqueous medium. The degree of ionization of the acids and bases helps determine their strengths. On the basis of different acidic and basic compounds, the degree of ionization may differ.

Ionization of Acids

The degree of Ionization alludes to the strength of a corrosive or a base. A solid corrosive can totally ionize in water, while a feeble corrosive is said to just ionize to some degree. 

Since the ionization of a powerless corrosive is a balance, the compound condition and a steady articulation can be expressed as :

HA ( aq ) + H2O ( l ) sign of equilibrium reaction H3O+ ( aq ) + A–

Ka = [ H3O+ ] [A–] / [HA]

The balance Constant for ionization of a corrosive characterizes its Acid Ionization Constant (Ka). In any case, the more grounded the corrosive, the bigger will be the corrosive ionization steady (Ka). This implies that a solid corrosive is a superior proton benefactor. Because of the convergence of the item in the numerator of the Ka, the more grounded the corrosive, the bigger is the corrosive ionization steady (Ka).

Ionization of acetic acid

Strong acids ionize completely when they are dissolved in water, while weak acids ionize only slightly. As an example, glacial acetic acid has an acid dissociation constant of 1.75 x 10-5. A 10 M solution of acetic acid has a percent ionization of only 0.132 %.

Basic equilibrium equation for NH3

 Ammonia (NH3) is a gas that readily dissolves in water and behaves as a base. The ammonia equilibrium is described with the equation.

 NH3 + H2O = NH4(+) + OH(-). 

Formally, the acidity of the solution is expressed as pH. This is the logarithm of the concentration of hydrogen ions (protons, H+) in the solution.

Ionization of the weak acid

 In an equilibrium-controlled acid-base reaction, the equilibrium position always favors the formation of the weaker acid and the weaker base. This is because the weaker acid and the weaker base are the most stable species due to their lower potential energies.

 Direction of equilibrium

 Q can be used to determine which direction a reaction will shift to reach equilibrium. If K > Q, a reaction will proceed forward, converting reactants into products. If K < Q, the reaction will proceed in the reverse direction, converting products into reactants. If Q = K, then the system is already at equilibrium.

Calculation of KC and QC

Qc and Kc are calculated the same way, but Qc is used to determine which direction a reaction will proceed, while Kc is the equilibrium constant (the ratio of the concentrations of products and reactants when the reaction is at equilibrium). So, Qc could be = to Kc, but it may not be.

 Base equilibrium constant

 Historically, the equilibrium constant Kb for a base has been defined as the association constant for protonation of the base, B, to form the conjugate acid, HB+. In actuality, there is no need to define pKb separately from pKa, but it is done here because pKb values are found in some of the older chemistry literature.

Relationship between Ka and Kb

 The Ka is the acid dissociation constant. The larger the value of Kb, the stronger the base, and the larger the value of Ka, the stronger the acid. By multiplying Ka by Kb, you receive the Kw, or the dissociation constant for water, which is 1.0 x 10-14.

Relationship between Ka and Kb for a conjugate acid-base pair

The lower Ka for the acid indicates that it’s a weak acid that holds tightly onto the donatable proton. The weaker the acid, the stronger the base. The stronger the base, the higher the Kb. The weaker the acid, the lower the Ka.