Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 431	Biochemistry	Fall 2001
Lecture Notes:: 7 September	© R. Paselk 2001
PREVIOUS		NEXT

Water, cont.

Because of the very strong dipole moments of these bonds and the very small size of the hydrogen substituents on water, a slight degree of orbital overlap occurs between adjacent water oxygens and hydrogens to give partial covalent bonds known a H-bonds (effectively, can only form with O, N, & F).

• Note that the partial covalent character means that they are directional! Figure 2.2 (p 24) shows a representation of H-bonds.
• Compare the bond length of water H bonds (0.18 nm) to the covalent bond-length between O and H of 0.096 nm - notice that the bond distance is nearly twice the true covalent bond distance, but significantly less than the van der Waals radius of 0.26 nm. [overhead 3.3, NP]

But in addition to covalent bonds there are a variety of non-covalent bonds/interactions as seen in the table below:

Interaction Type Example Average Strength, kcal/mol (kJ/mol) Range** Charge-charge (ionic)    5 (20) [in water solution] 1/r Charge-dipole     1/r2 Dipole-dipole      1/r3 Charge-induced dipole       Dipole-induced dipole*   0.1-0.2 (0.4-4)  1/r6 Dispersion*   0.1-0.2 (0.4-4) 1/r6 Hydrogen bond   3-8 (12-30)    van der Waals repulsion     1/r12

*van der Waals interactions, **from Zubay Biochemistry 3rd. Table 4.3, pg. 89.

Within solid bulk water (ice):

• every water molecule is bonded to 4 others, as in the ice structure seen in Figure 2.3 on pg. 24 [overhead 2.3, VV]
  • In liquid water the molecules are still bonded to a large degree (the heat of fusion for ice is only 13% of the heat of vaporization for ice, thus most of the H-bonds must survive melting).
  • Of course in liquid water the bonds are very unstable (average lifetime about 10 psec = 10^-11 sec), exchanging constantly to give a "flickering cluster" structure.
  • The various properties of water arise from this structure. (Note hi bp & mp, heat cap., viscosity, and, less obviously, that ice floats.
• Ice floats because the molecules are in an open lattice rather than close-packed. Garrett & Grisham (in their text, Biochemistry, 2nd ed) note that close-packed water molecules would only occupy about 57% of the volume of ice. This would lead one to expect that ice would float "high." It doesn't because most of the structure remains in the liquid phase at 0° C.)

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 on pg 26 shows the hexavalent 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. [Figure 2.8, p 28; overhead 7-51, VV]
  • 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 Figure 2.10 on pg 29. [overhead 2-7 VV]

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⁺) there will only be 10^-7 molar hydronium ions, in other words only about 2 of every billion water molecules will be protonated!

Ionization of Water, pH & Buffers

For aqueous solution [H⁺][OH^-]= 10^-14;

• pH = -log [H⁺]. Remember that a lo pH means a high concentration of protons.
• pK_a = -logK_atherefore pH + pOH = 14, where pOH = -log[OH^-].
• Brönsted acid: we will be using the Brönsted definition for acids and bases­an 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. For the general acid HA we can write:

   

  

   

  Where K_a is the acid dissociation constant. (Note that the definition of K_a is based on the Brønsted definition.) Values of K_a 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 K_a 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 K_a values of 10^-5 to 10^-15. Note that whether an acid is strong or weak is dependent of the solvent system! Strong acids have weak conjugate bases, and vice-versa.

• For reactions involving a strong acid or base we can assume, for practical purposes, that all of the strong acid or base added to a mixture will react until the base or acid originally present in solution is completely consumed. (Of course this is an approximation, all reactions actually approach an equilibrium condition, so that, in theory, there is always some reactant and some product present.) For example, if we start with a solution containing 0.100 mole of acetic acid and add 0.050 moles of sodium hydroxide the resulting mixture will contain 0.05 moles acetic acid, 0.050 moles sodium acetate and 0.000 moles sodium hydroxide (actually about 10^-10 moles, which is 0.000 for our two significant figure calculation).

The equilibrium equation for a mixture of a weak acid and its conjugate base can be rewritten by taking logs of both sides and rearranging to give the Henderson-Hasselbalch equation: pH = pK_a + log [A^-]/[HA]

We frequently represent the reaction of an acid with a base as a titration curve. (overheads) You should understand these curves and be able to label them for axis, percent dissociation at beginning, middle and end, buffer region, end-point, and how to find pK_a. An exercise to help you to review titrations curves is available.

Nitrogenous Bases and Nucleic Acids

Nucleotides: Nitrogenous base + sugar + phosphate

Nitrogenous bases (overhead; Table 3-1, p 43)

Purines: Fused 6 & 5 memebered hetero CN-ring system, usually unsaturated. Two common purines in biological systems, both used in DNA and RNA, as well as in energy carriers (ATP & GTP). A is also widely used as a recognition "handle" attached to vitamins etc. to aid enzymes and other proteins to find and bind these molecules.

• adenine
• guanine