Strong and Weak Acids and Bases

Intro to Acids and Bases

Arrhenius Theory of Acids and Bases  An acid is a substance that dissociates in water to produce one or more hydrogen ions (H+)  A base is a substance that dissociates in water to form one or more hydroxide ions. (OH-)

Examples: + - Acid: HCl(aq) H (aq) + Cl (aq) + - Base: LiOH  Li (aq) + OH (aq)

Limitations:  Classified based on chemical formula  Some substances do not have OH- in their chemical formulas but still - yield OH when they react with water. E.g. NH3 (ammonia)

Bronsted-Lowry Theory of Acids and Bases  An acid is a proton (H+) donor and must have H in its formula.  A base is a proton acceptor and must have a lone pair of electrons to form a bond with H+

 Two molecules or ions that are related by the transfer of a proton are called a conjugate acid-base pair.  Conjugate acid of a base is the particle that results when the base receives the proton from the acid.  Conjugate base of the acid is the particle that results when the acid donates a proton. Identify the conjugate acid/base pairs in the following:

+ - NH3(aq) + H2O(l) ⇄ NH4 (ag) + OH (aq)

Amphiprotic: Can act as either an acid or a base i.e has both a lone pair and an H-atom

- - HCO3 (aq) ) + H2O(l) ⇄ H2CO3(aq) + OH (aq)

- 2- + HCO3 (aq) + H2O(l) ⇄ CO3 (aq) + H3O (aq)

Strong and Weak Acids and Bases Strong Acids and Bases  Completely dissociate in water into their ions

100% + - HCl(aq) + H2O(aq)  H3O (aq) + Cl (aq)

100% - LiOH + H2O(aq)  LiOH(aq) + OH (aq)

+  As a result the [H3O ] in a solution of a strong acid is equal to the concentration of the acid.

 Strong acids include HClO4 (perchloric), HI, HBr, HCl, H2SO4

(sulfuric), and HNO3 (nitric)  Strong bases include all oxides and hydroxides of alkali metals as well as alkaline earth metal oxides and hydroxides below beryllium. Weak Acids and Bases  Do NOT completely dissociate in water into their ions

1% + - CH3COOH(aq) + H2O(aq) ↔ H3O (aq) + CH3COO (aq)

1% + - NH3(aq) + H2O(aq) ↔ NH4 (aq) + OH (aq)

+  As a result, the concentration [H3O ] in a solution of a weak acid is always less than the concentration of the dissolved acid.

Percent Ionization

+ [H (aq) ] % ionization = 100 % ionization [HA ] ҙ [H + ] = [HA ] (aq) (aq) 100 ҙ (aq)

Polyprotic Acids:

Monoprotic acids contain only a single hydrogen ion that can dissociate. For example HCl

Diprotic acids contain two hydrogen ions that can dissociate. For example H2SO4

Triprotic acids contain three hydrogen ions that can dissociate. For example H3PO4 Acid/Base Equilibrium Constants (dissociation constants)

(Table on p. 803)

Weak acids and bases are dealt with as equilibrium systems and therefore, will have an equilibrium constant that will be determined by the concentrations of each component in the system when equilibrium is established:

For acids the dissociation constant is called Ka, while for bases it is called Kb.

- + CH3COOH(aq) + H2O(l) CH3COO (aq) + H3O (aq)

- + Ka = [CH3COO ] [H3O ]

[CH3COOH]

+ - NH3(aq) + H2O(l) NH4 (aq) + OH (aq)

+ - Kb = [NH4 ] [OH ]

The smaller the value of K = the weaker the acid or base

Strong acids/bases have very large values of K and are not considered as equilibrium systems (assumed 100% dissociation)