Arrhenius concept of Acids and bases

Hydrochloric acid, found in gastric juice and vinegar, is essential for digestive processes and can be found in lemon and orange juices. Acids, such as acetic and tartaric acids, taste sour and can turn blue litmus paper red or liberate dihydrogen when reacting with certain metals. Bases, on the other hand, turn red litmus paper blue, taste bitter, and feel soapy. When mixed in the right proportion, acids and bases react to form salts, such as sodium chloride, barium sulphate, and sodium nitrate.

Sodium chloride, a common salt, is formed by the reaction between hydrochloric acid and sodium hydroxide. Hydrochloric acid and acetic acid are polar covalent molecules, with hydrochloric acid being completely ionized into its constituent ions and acetic acid being only partially ionized (< 5%). Ionization occurs when neutral molecules split into charged ions in the solution.

According to Arrhenius theory, acids dissociate in water to give hydrogen ions H+ (aq) and bases produce hydroxyl ions OH–(aq). The ionization of an acid HX (aq) can be represented as HX → H+ (aq) + X– (aq) or HX(aq) + H2O(l) → H3O+ (aq) + X–(aq).

However, the Arrhenius concept of acid and base is limited to aqueous solutions and does not account for the bassist of substances like ammonia without a hydroxyl group.

 

Bronsted Lowry Concept

 

Acids and bases are substances that donate and accept hydrogen ions, respectively. Acids are proton donors, while bases are proton acceptors. The basic solution is formed due to the presence of hydroxyl ions. Water and ammonia are called Lowry-Brönsted acids and bases, respectively. In the reverse reaction of NH3 in H2O, H+ is transferred from NH4+ to OH–, forming a Bronsted acid and a Brönsted base.

Acid-base pairs differ only by one proton, known as conjugate acid-base pairs. The conjugate base of an acid H2O and the conjugate acid of the base NH3 are different. If a Brönsted acid is a strong acid, its conjugate base is a weak base, and vice versa. Water acts as both an acid and a base in the ionization of hydrochloric acid in water. In the case of HCl, water acts as a base, while in the case of ammonia, it acts as an acid by donating a proton.

G.N. Lewis defined acids as species that accept electron pairs and bases that donate electron pairs in 1923. There is little difference between the Brönsted-Lowry bases and Lewis bases concepts, as the base provides a lone pair in both cases. Electron deficient species like AlCl3, Co3+, and Mg2+ can act as Lewis acids, while species like H2O, NH3, and OH– can act as Lewis bases.

 

Ionization Constant of Acids

 

Strong acids like perchloric acid (HClO4), hydrochloric acid (HCl), hydrobromic acid (HBr), hyrdoiodic acid (HI), nitric acid (HNO3) and sulphuric acid (H2SO4) are termed strong because they are almost completely dissociated into their constituent ions in an aqueous medium, thereby acting as proton (H+) donors. Strong bases like lithium hydroxide (LiOH), sodium hydroxide (NaOH), potassium hydroxide (KOH), caesium hydroxide (CsOH) and barium hydroxide Ba(OH)2  are almost completely dissociated into ions in an aqueous medium giving hydroxyl ions, OH–.

According to Arrhenius concept they are strong acids and bases as they are able to completely dissociate and produce H3O+ and OH– ions respectively in the medium. strong acid means a good proton donor and a strong base implies a good proton acceptor.

Consider, the acid-base dissociation equilibrium of a weak acid HA, acid (or base) dissociation equilibrium is dynamic involving a transfer of proton in forward and reverse directions. Tendency of donating a proton over the other shall be termed as the stronger acid and the equilibrium will shift in thedirection of weaker acid.

If HA is a stronger acid than H3O+, then HA will donate protons and not H3O+ , and the solution will mainly contain A– and H3O+ ions.

The equilibrium moves in the direction of formation of weaker acid and weaker base. The stronger acid donates a proton to the stronger base.

A strong acid dissociates completely in water, the resulting base formed would be very weak i.e., strong acids havevery weak conjugate bases.

Strong acids like perchloric acid (HClO4), hydrochloric acid (HCl), hydrobromic acid (HBr), hydroiodic acid (HI), nitric acid (HNO3) and sulphuric acid (H2SO4) will give conjugate base ions ClO4– , Cl, Br–, I– , NO3– and HSO4 – , which are much weaker bases than H2O.A very strong base would give a very weak conjugate acid.

Weak acids are nitrous acid (HNO2), hydrofluoric acid (HF) and acetic acid (CH3COOH). It should be noted that the weak

Acids have very strong conjugate bases.

Water is unique in its ability to act both as an acid and a base.

In pure water, one H2O molecule donates a protons and acts as an acid and another water molecule accepts a proton and acts as a base at the same time.

The dissociation constant is represented by, K = [H3O+] [OH–] / [H2O]

The concentration of water is omitted from the denominator as water is a pure liquid and its concentration remains constant. [H2O] is incorporated within the equilibrium constant to give a new constant, Kw, which is called the ionic product of water.

Kw = [H+][OH–]

The concentration of H+ has been found out experimentally as 1.0 × 10–7 M at 298 K. Dissociation of water produces equal number of H+  and OH– ions,

the concentration of hydroxyl ions, [OH–] = [H+] = 1.0 × 10–7 M.

the value of Kw at 298K,Kw = [H3O+][OH–] = (1 × 10–7)2 = 1 × 10–14 M2

The value of Kw is temperature dependent as it is an equilibrium constant.

From this the molarity of pure water can be given as,[H2O] = (1000 g /L)(1 mol/18.0 g) = 55.55 M.

We can distinguish acidic, neutral and basic aqueous solutions by the relative values of the H3O+ and OH– concentrations:

Acidic: [H3O+] > [OH– ]

Neutral: [H3O+] = [OH– ]

Basic : [H3O+] < [OH–]