Chemistry · Chapter 2
Study notes aligned to the official NEB syllabus.
Ionic equilibrium is the equilibrium established in solutions of weak electrolytes between the un-ionised molecules and the ions they produce. Strong electrolytes (strong acids, strong bases, most salts) ionise almost completely; weak electrolytes (weak acids like $\ce{CH3COOH}$, weak bases like $\ce{NH3}$) ionise only partly and set up a dynamic equilibrium.
$$\ce{CH3COOH(aq) <=> CH3COO-(aq) + H+(aq)}$$
Arrhenius defined an acid as a substance giving $\ce{H+}$ and a base as one giving $\ce{OH-}$ in water. Limitations: it applies only to aqueous solutions, cannot explain the basicity of $\ce{NH3}$ (no $\ce{OH-}$) or acidic/basic behaviour of salts, and treats the bare proton as free (it is really hydrated as $\ce{H3O+}$).
An acid is a proton ($\ce{H+}$) donor; a base is a proton acceptor. This works in any solvent:
$$\ce{HCl + H2O -> H3O+ + Cl-}$$
Here $\ce{HCl}$ is the acid, $\ce{H2O}$ the base.
When an acid donates a proton it becomes its conjugate base; when a base accepts a proton it becomes its conjugate acid. They differ by one $\ce{H+}$:
$$\ce{CH3COOH <=> CH3COO- + H+}$$
Here $\ce{CH3COOH}$ (acid) and $\ce{CH3COO-}$ (its conjugate base) form one conjugate pair, while $\ce{H2O}$/$\ce{H3O+}$ form another.
A strong acid has a weak conjugate base and vice versa. Water is amphoteric (can act as acid or base).
A Lewis acid is an electron-pair acceptor (e.g. $\ce{BF3}$, $\ce{AlCl3}$, $\ce{H+}$); a Lewis base is an electron-pair donor (e.g. $\ce{NH3}$, $\ce{OH-}$). This is the broadest definition:
$$\ce{BF3 + NH3 -> BF3.NH3}$$
(the nitrogen lone pair forms a coordinate bond to the electron-deficient boron).
Strength is the extent of ionisation. Strong acids ($\ce{HCl, H2SO4, HNO3}$) ionise fully; weak acids ($\ce{CH3COOH}$) ionise partly. Strength is measured by the dissociation constant ($K_a$, $K_b$).
Ionic equilibrium is the equilibrium established in solutions of weak electrolytes between the un-ionised molecules and the ions they produce. Strong electrolytes (strong acids, strong bases, most salts) ionise almost completely; weak electrolytes (weak acids like , weak bases like ) ionise only partly and set up a dynamic equilibrium.
Arrhenius defined an acid as a substance giving and a base as one giving in water. Limitations: it applies only to aqueous solutions, cannot explain the basicity of (no ) or acidic/basic behaviour of salts, and treats the bare proton as free (it is really hydrated as ).
An acid is a proton () donor; a base is a proton acceptor. This works in any solvent:
Here is the acid, the base.
When an acid donates a proton it becomes its conjugate base; when a base accepts a proton it becomes its conjugate acid. They differ by one :
Here (acid) and (its conjugate base) form one conjugate pair, while / form another.
A strong acid has a weak conjugate base and vice versa. Water is amphoteric (can act as acid or base).
A Lewis acid is an electron-pair acceptor (e.g. , , ); a Lewis base is an electron-pair donor (e.g. , ). This is the broadest definition:
(the nitrogen lone pair forms a coordinate bond to the electron-deficient boron).
Strength is the extent of ionisation. Strong acids () ionise fully; weak acids () ionise partly. Strength is measured by the dissociation constant (, ).