Definition and properties of acids and bases
Acids and bases posses in a sense some opposite properties.
- Any of the classes of substances whose aqueous solutions are characterized by the sour taste, the ability to turns blue litmus, and the ability to react with bases and certain metals to form salts are called acids.
- But a base is a substance that in aqueous solution is slippery to touch, that’s bitter, changed the color of the indicator, (litmus to blue) reacts with an acid to form salts.
The concepts of acid and base are developed one after another tends to make the acid-base definitions to more broad-based.
To describe the properties of acid and bases we need to familiar with five categories.
- Arrhenius acid-base theory
- Solvent system concept
- Protonic concept
- Lewis concept
- Hard soft acid-base principle
Arrhenius acid-base theory
Arrhenius was one of the early exponents of electrolytic dissociation theory to define acids and bases.
His classification of acids based on the theory that acids when dissolved in water, dissociate hydrogen ions and anions.
But bases, when dissolved in water, dissociate into hydroxyl ions and cations.
Thus sodium hydroxide neutralizes hydrochloric acid can be represented by a reaction involving the combination of hydrogen and hydroxyl ions to form water.
NaOH → Na+ + OH–
H+ + OH+ ⇆ H2O
Acid-base neutralization reaction
Arrhenius concept of acids and bases was highly useful in explaining the acid-base neutralization reaction in an aqueous medium.
|Strong acid||Strong base||Salt||Water|
where HA and BOH are strong electrolytes and completely dissociate in the solution.
H+ + A– + B+ + OH– ⇆ B+ + A– + H2O
Canceling likes ions
H+ + OH– ⇆ H2O
The heat of neutralization reaction
The neutralization reaction of all strong acids and bases used to the formation of 1-mole water from H+ ion and OH– ion. Thus the enthalpy change (ΔH) will also be the same for the reaction between strong acids and bases.
H+ + OH– ⇆ H2O + 13.6 kcal
So heat release during neutralization HCl, HClO4, HNO3, HBr, HI, and H2SO4 and strong base NaOH, KOH, RbOH, and Ca(OH)2) is the same, namely 13.4 kcal/mole or 56 KJ/mole.
Limitations of Arrhenius theory
- According to Arrhenius’s theory, HCl is an acid only when dissolved in water. But if we use solvent such as benzene or when it exists in the gaseous state, it does not consider as an acid.
- This theory cannot account for the acidic and basic properties of the materials in non-aqueous solvents.
For example, NH4NO3 in liquid ammonia is an acid, though it does not give H+ ions.
Similarly, many organic compounds in NH3, which do not give OH– ions at all, actually known to show basic character.
- The neutralization process limited to those reactions which can occur in aqueous solutions only, although the reactions involving salt formation do occur in many other solvents and even in the absence of solvents.
- This theory cannot explain the acidic character of certain salts such as AlCl3 in aqueous solution.
Acid-base chemistry of solvent
Protic solvent dissociates into two oppositely charged ions. If we consider water as the protic solvent, it dissociates H+ and OH– ions.
Thus hydrogen ion or a proton readily polarizes other anion or molecule, where solvated proton exists in the solution. So overall dissociation reaction of the solvent
Thus according to solvent chemistry, all those compounds which can give H3O+ ions in H2O will act as acids and all those compounds which can give OH– ions in H2O will act as bases.
Properties of liquid ammonia
Ammonia dissociates into two oppositely charged ions. Ammonia dissociates into NH4+ and NH2– ions.
Thus those compounds which give NH4+ ions in liquid ammonia will act as acids and all those compounds which can give NH₂⁻ ions in liquid ammonia act as bases.
Solvent system definition of acid and base
The dissociation of non-aqueous solvents directly responsible for the nature of the chemical reactions that can be initiated in such solvents.
According to the solvent system,
- An acid is a substance which by dissociation in the solvent forms the same cation as does the solvent itself due to auto-ionization.
- A base is one that, gives on dissociation in the solvent the same anion as does the solvent itself on its ionization.
Autoionization of some protonic and non-protonic solvents
|Solvent||Autoionization of solvent|
|Protic Solvent||H2O + H2O ⇆ H3O+ + OH–|
|NH3 + NH3 ⇆ NH4+ + NH2–|
|Aprotic Solvent||SO2 + SO2 ⇆ SO+ + SO3–|
|BrF3 + BrF3 ⇆ BrF2+ + BrF4–|
|N2O4 + N2O4 ⇆ 2NO+ + 2NO3–|
Neutralization reaction examples
Uses of solvent system theory
Evidently, this concept of the solvent system can be used to explain the acid-base neutralization reactions occurring in protic and aprotic solvents(protonic or non-protonic both).
Thus this theory can simply be said to be an extension of the Arrhenius theory.
Limitations of solvent system theory
- This theory does not consider a number of acid-base reactions included in the protonic definition.
- It limits acid-base phenomena to the solvent system only. Thus it does not explain the acid-base reactions which may occur in the absence of solvent.
- It can not explain acid-base neutralization reactions occurring without the presence of ions.