|Lecture Notes: 10 September
© R. Paselk 2008
Water is an excellent solvent for polar substances since its dipolar structure enables it to insulate them from each other and it can make good dipole-dipole and dipole-charge bonds. Figure 2-6 shows the liganding of water to sodium and chloride ions to form hydration shells (For sodium ions, the waters in the inner hydration-shell exchange every 2-4 nsec.). Anything which can H-bond will also of course be quite soluble.
How does water interact with non-polar molecules?
- The problem here is that in order to dissolve in water a non-polar molecule must disrupt a series of H-bonds and no new bonds of equal strength are substituted.
- Thus water tends to exclude non-polar substances. When we forcefully disperse a non-polar substance into water then the water must form a cage around the molecule to maximize H-bonds for each water molecule. [text Figure 2-7a]
- Now an additional problem arises-the waters hydrating the non-polar group are "locked" in place - they can't easily flip about because there is no interior bond for substitution!
- Thus the entropy of these water molecules is greatly reduced. So the insolubility of non-polar groups in water has both enthalpic and entropic factors!
- Note that many non-polar molecules may dissolve in water by the formation of micelles as in text Figure 2-7b.
Finally, recall that water is a good nucleophile and so will participate in many chemical reactions-readily hydrolyzes esters, amides, anhydrides etc.
Ionization of Water, pH & Buffers
Dissociation of water molecules: One aspect of water we have yet to talk about is its dissociation or ionization. In normal aqueous solution there is a certain probability that a hydrogen nucleus (a proton) can exchange between two hydrogen bonded molecules:
(Of course the hydronium ion, H3O+, will be associated with additional water molecules as well through H-bonding. For simplicity we will just write H+, with the understanding that it refers in fact to hydrated hydronium ions in aqueous solution. ) Note the reaction is not highly favored, in neutral solution (no excess H+or OH-) there will only be 10-7 molar hydronium ions, in other words only about 2 of every billion water molecules will be protonated!
For aqueous solution [H+][OH-]= 10-14
- pH = -log [H+]. Remember that a lo pH means a high concentration of protons.
[text Figure 2-14]
- pKa = -logKatherefore pH + pOH = 14, where pOH = -log[OH-].
- Brönsted acid: we will be using the Brönsted definition for acids and bases: an acid is a proton donor, while a base is a proton acceptor. Recall the corollary that acids and bases therefore exist as conjugate acid base pairs. Note that when an acid by this definition gives up a proton it becomes a base, since the reverse reaction would be accepting a proton:
- Thus the acetic acid in the first reaction becomes its conjugate base acetate ion, while the base, hydroxide ion, becomes its conjugate acid, water. In the reverse reaction the nomenclature also reverses. Note that a molecule such as water can be both an acid, donating a proton to become its conjugate base hydroxide ion, or it can be a base, accepting a proton to become its conjugate acid, a hydronium ion.
The strengths (ability to donate protons) of acids vary considerably. [text Figure 2-15]
- For the general acid HA we can write:
- Where Ka is the acid dissociation constant. (Note that the definition of Ka is based on the Brønsted definition.) Values of Ka can vary tremendously (1015 to 10-60) - after all anything with at least one proton can be considered an acid under some circumstances with this definition. The common definition of a strong acid is an acid which dissociates completely in a 1 M solution. The common strong acids in aqueous solution, such as sulfuric, nitric and hydrochloric acids have Ka values (for the first dissociation in the case of sulfuric) of 102 to 109. Thus they all dissociate completely (first dissociation only for sulfuric) in aqueous solution, though they will have different strengths in some other solvents. Most common organic acids are weak in aqueous solution, having Ka values of 10-5 to 10-15. Note that whether an acid is strong or weak is dependent on the solvent system! Strong acids have weak conjugate bases, and vice-versa.
Last modified 10 September 2008