Colloids

Several important acids can be classified as polyprotic acids, which can lose more than one H⁺ ion when they act as Bronsted acids. Diprotic acids, such as sulfuric acid (H₂SO₄), carbonic acid (H₂CO₃), hydrogen sulfide (H₂S), chromic acid (H₂CrO₄), and oxalic acid (H₂C₂O₄) have two acidic hydrogen atoms. Triprotic acids, such as phosphoric acid (H₃PO₄) and citric acid (C₆H₈O₇), have three.

There is usually a large difference in the ease with which these acids lose the first and second (or second and third) protons. When sulfuric acid is classified as a strong acid, students often assume that it loses both of its protons when it reacts with water. That isn't a legitimate assumption. Sulfuric acid is a strong acid because K_a for the loss of the first proton is much larger than 1. We therefore assume that essentially all the H₂SO₄ molecules in an aqueous solution lose the first proton to form the HSO₄^-, or hydrogen sulfate, ion.

H₂SO₄(aq) + H₂O(l) H₃O⁺(aq) + HSO₄^-(aq)

K_a1 = 1 x 10³

But K_a for the loss of the second proton is only 10^-2 and only 10% of the H₂SO₄ molecules in a 1 M solution lose a second proton.

HSO₄^-(aq) + H₂O(l) H₃O⁺(aq) + SO₄^2-(aq)

K_a2 = 1.2 x 10^-2

H₂SO₄ only loses both H⁺ ions when it reacts with a base, such as ammonia.

The large difference between the values of K_a for the sequential loss of protons by a polyprotic acid is important because it means we can assume that these acids dissociate one step at a time

an assumption known as stepwise dissociation.

As every next dissociation constant is smaller, usually it is enough to do the pH calculation treating acid as if it was only monoprotic (ie omitting effects of the next dissociation step). Further if the acid is weak even in respect to the first dissociation step and the second constants is much smaller, we can neglect second dissociation step completely.

Lewis acids and bases

The most general definition of acids and bases, which encompasses the Arrhenius and Bronsted-Lowry definitions is due to our old friend, Lewis and his dot structures. A Lewis acid is defined to be any species that accepts lone pair electrons. A Lewis base is any species that donates lone pair electrons. Thus, H⁺ is a Lewis acid, since it can accept a lone pair, while OH^- and NH₃are Lewis bases, both of which donate a lone pair:

Interestingly, however, is that species which have no hydrogen to donate (a la the Bronsted-Lowry scheme) can still be acids according to the Lewis scheme. As an example, consider the molecule BF₃. If we determine Lewis structure of BF₃, we find that B is octet deficient and can accept a lone pair. Thus it can act as a Lewis acid.

Oxoacids have the hydrogen atoms in the acid bonded to oxygen. The more oxygen atoms there are in the acid the stronger the acid will be.

If we compared acid strength with acids where the atoms bonded to hydrogen were in the same period, the difference in the size of the atoms would not be significant. In that case the electronegativity becomes the deciding factor. The greater the electronegativity of the atom the stronger will be the acid.

Acid Strength and Atom Size: When the bond breaks between a Hydrogen atom and another atom, both bonding electrons will go to the other atom giving it a negative charge. The larger this atom the greater will be its surface area. The negative charge can be dispersed over this larger surface area and thereby be stabilized. The more stabilized this anion the less likely it will react with the Hydrogen ion to reverse the process. In general if we compare acids with the other atoms are in the same group of the periodic table, the further down in the group the atom is the stronger will be the acid.

Selenious acid (H₂SeO₃), a diprotic acid, has K₁ = 3.0 x 10^-3 and K₂ = 5.0 x 10^-8. What is the [HSeO₄^-] of a 0.50 M solution of selenious acid?

3.0 x 10^-3 = x²/0.50-x x² + 3.0x10^-3x – 0.0015 = 0 x = 0.037 = [HSeO₄^-]