Relative Strengths of Ionization of Acids and Bases

Overview of Acids and Bases- A Lewis acid is a molecule or ion that is capable of either donating a proton (i.e., hydrogen ion, H+) or creating a covalent bond with an electron pair.

Proton donors, or Bronsted–Lowry acids, are the first class of acids. Proton donors, in the peculiar situation of aqueous solutions, generate the hydronium ion H3O+ and are referred to as Arrhenius acids. Bronsted and Lowry extended Arrhenius’s theory beyond aqueous solvents. Typically, a Bronsted or Arrhenius acid comprises a hydrogen atom bound to a chemical structure that retains a favourable energy structure after the loss of H+.

Aqueous Arrhenius acids exhibit distinctive features that serve as a useful description of an acid. Acids produce sour aqueous solutions, can discolour blue litmus paper, and react with bases and some metals (such as calcium) to generate salts. The term ‘acid’ comes from the Latin acids/acre, which means ‘sour’. Aqueous solutions of acids with a pH less than 7 are generally referred to as “acid” (as in “dissolved in acid”), but the formal definition relates exclusively to the solute. A lower pH indicates increased acidity, which results in a greater concentration of positive hydrogen ions in the solution. Acidic chemicals or compounds are those that exhibit the properties of an acid.

Svante Arrhenius postulated in 1884 that a base is a material that dissociates into Hydroxide ions OH in aqueous solution. These ions can combine with hydrogen ions (H+) formed during acid dissociation to create water in an acid–base reaction. As a result, a base was defined as a metal hydroxide such as sodium hydroxide or calcium hydroxide. Additionally, some features of these aqueous hydroxide solutions were characterised. They are slick to the touch, have a bitter taste, and cause pH metres to change colour (e.g., turn red litmus paper blue).

By modifying the autoionization equilibrium in water, bases produce solutions with a lower hydrogen ion activity than pure water, i.e., water with a pH greater than 7.0 under typical circumstances. If a soluble base includes and releases OH ions quantitatively, it is referred to as an alkali. Metal oxides, hydroxides, and particularly alkoxides are basic, whereas weak acids’ conjugate bases are weak bases.

Acids and bases are considered chemical polar opposites because acids increase the concentration of hydronium (H3O+) in water, whereas bases decrease it. The term “neutralisation” refers to the interaction between aqueous solutions of an acid and a base that results in a solution of water and a salt in which the salt separates into its component ions. When an aqueous solution becomes saturated with a particular salt solute, any remaining salt precipitates out of the solution.

The Relative Strengths of Ionization of Acids and Bases

When an acid or base is dissolved in water, its relative strength is determined by its ability to ionise. When the ionisation process is nearly complete, the acid or base is said to be strong; when only a little amount of ionisation occurs, the acid or base is said to be weak. As will become clear throughout this chapter, there are many more weak acids and bases than strong acids and bases. Table 1 is a list of the most frequently used strong acids and bases.

Table 1. Some Common Strong acids and Strong Bases

Acids’ relative strengths can be assessed by determining their equilibrium constants in water. Stronger acids ionise to a larger degree than weaker acids in solutions of the same concentration, resulting in higher quantities of hydronium ions. The acid-ionization constant, Ka, is the equilibrium constant for an acid. For an acidic response, HA:

HA (aq)+H2O(l)⇌H3O+(aq) +A−(aq) HA(aq)+H2O(l)⇌H3O+(aq)+A−(aq)

The acid ionization constant is written as- 

Ka=[H3O+][A−][HA]Ka=[H3O+][A−][HA]

where the equilibrium concentrations are used. Although water is a reactant in the process, it also acts as a solvent, therefore we omit [H2O] from the equation. The greater the Ka of an acid, the greater the concentration of hydrogen sulphide. O+ and A in relation to the nonionized acid concentration, HA, in an equilibrium mixture, and the stronger the acid. When an acid undergoes full ionisation, the concentration of HA becomes zero and the acid ionisation constant becomes unimaginably huge (Ka ). Acids that are only partly ionised are referred to as “weak” acids, and their acid ionisation constants may be determined experimentally. Ionization Constants of Weak Acids is a table of weak acid ionisation constants.

Three acid ionisation equations and their associated Ka values are presented below to demonstrate this concept. The ionisation constants decrease as the order of the above equations rises, suggesting that the relative acid strength decreases. CH₂CO₂H<HNO₂ <HSO₁² :

CH,CO,H{ag) +H,O(0)=HO (aq) + CH,CO, (aq) K =18 × 10%

HNO₂ (aq) + H₂O (1) H₂O (aq) + NO₂ (aq)  K₂=4.6 x 10-4

HSO, (aq) + H₂O (aq) H₂O (aq) + SO (aq)  K₁=1.2× 10 ²

Another indicator of an acid’s potency is its percent ionisation. The percent ionisation of a weak acid is described in terms of an equilibrium mixture’s composition:

% ionization=[H3O+]eq[HA]0×100% ionization=[H3O+]eq[HA]0×100

where the numerator equals the concentration of the acid’s conjugate base ([A] = [H3O+] via stoichiometry). In contrast to the Ka value, the percent ionisation of a weak acid changes with the initial acid concentration, often decreasing as the acid concentration increases. Calculations of equilibrium, such as those presented later in this chapter, can be used to corroborate this behaviour.

We may evaluate the relative strengths of bases by measuring their base-ionization constant, (Kb), in aqueous solutions, much as we did with acids. Stronger bases ionise to a larger degree in solutions of the same concentration, resulting in higher hydroxide ion concentrations than weaker bases. The ionisation constant of a stronger base is higher than that of a weaker base. B:

B (aq) 1₂0 (1) HB (aq) OH (aq)

The ionisation constant equation is written as: 

Kb = [HB] [OH-]/[B]

The concentrations at equilibrium are those. Because water is the solvent, we don’t include [H2O] in the equation. The three bases’ chemical reactions and ionisation constants are as follows:

CH3 CO2 H(aq)+H2 O(l) H3 O + (aq)+CH3 CO2– (aq) K_{a} = 1.8 * 10– 5

HNO₂ (aq) + H₂O (1) H₂O+ (aq) + NO₂– (aq) K_{a} = 4.6 * 10– 4

HSO4 (aq) + H₂O (aq) H3O+ (aq) + SO42- (aq) K_{a} = 1.2 * 10– 2

Definition of ionization

Ionization is the process of an atom or molecule gaining or losing electrons to gain or lose a charge. Ions are electrically charged atoms or molecules. Ionization can occur when an electron is lost through a collision with a subatomic particle, another atom, molecule, or ion, or by electromagnetic radiation. In heterolytic substitution reactions and bond breaking, ion pairs can be created. An excited nucleus can eject an inner-shell electron during the internal conversion process of radioactive decay. It ionises materials.

Conclusion

Any material containing hydrogen that is capable of transferring a proton (hydrogen ion) to another chemical is considered an acid. A base is a molecule or an ion that is capable of accepting a hydrogen ion emitted by an acid. Generally, acidic compounds have a sour flavour.