Equilibrium:Ionic equilibrium-ionization of acids and bases
Numerous equilibria involve only ionic exchange. The ionic exchange allows some chemical substance to behave as electrolytes that can conduct electricity through the movement of ions. Solution with weak electrolytes generates an ionic equilibrium of acid and base exchange.
Figure : Common uses of electrolytes – acids and bases
The Classical (Arrhenius) Definition of Acids and Bases state that an acid is a substance that has H in its formula and dissociates in waterto yield H3O+ and a base is a substance with a hydroxide (OH) in its formula that dissociates in waterto yield OH–.Arrhenius acids contain covalently bonded H atoms that ionize in water. Neutralization occurs when the H+ ion from the acid and the OH– ion from the base combine to form water.
The Brønsted-Lowry Definition of Acids and Bases states that an acid is a proton donor, that is, any species that donates an H+ ion. Accordingly, all Arrhenius acids are Brønsted-Lowry acids.A base is a proton acceptor, that is, any species that accepts an H+ ion.In the Brønsted-Lowry definition, an acid-base reaction is a proton transfer process.
acid1 + base2↔base1 + acid2
An acid reactant produces a base product and these two constitute an acid-base conjugate pair. Likewise, every acid has a conjugate base, and every base has a conjugate acid.
The Lewis Acid-Base definition states that an acid is an electron-pair acceptor and a base is an electron-pair donor. Lewis acids contain (or can generate) a vacant orbital.Metal ions act as Lewis acids when dissolved in water.
Consider the following example:
H2S and HS– are a conjugate acid-base pair where HS– is the conjugate base of the acid H2S.NH3 and NH4+ are another conjugate acid-base pair where NH4+ is the conjugate acid of the base NH3. The conjugate base of the pair has one fewer H and one more negative charge than the acid.The conjugate acid of the pair has one more H and one less negative charge than the base.
In general, an acid-base reaction proceeds to the greater extent in the direction in which a stronger acid and stronger base form aweaker acid and a weaker base.
Consider the following generalized acid-base reaction.
H+(aq) + OH– (aq) → H2O(l) ∆Horxn = -55.9 kJ
Strong acids dissociate completely into ions in water.
HA(g or l) + H2O(l)H3O+ (aq) + A– (aq) Kc>>1
Weak acids dissociate minimally into ions in water
HA(aq) + H2O(l) ↔ H3O+ (aq) + A– (aq) Kc<<1
In this equation, Kc stands for the equillibrium constant which can be used to define the acid dissociation constant, Ka.
Stronger acids have higher [H3O+] making larger Kawhile weaker acids have lower [H3O+] making smaller Ka.
In both cases, [H3O+] from the auto-ionization of water is negligible but a weak acid has a small Ka so we can assume that [HA]dissocis very small. Likewise,
[HA]eq = [HA]init – [HA]dissoc ≈ [HA]init
As the initial concentration of a weak acid decreases, the percent dissociation of the acid increases.
Figure : The order of ionic strength of acid and base