1. How H 2 0 interacts with: Itself –Hydrogen-bonding Ions and charged functional groups –Solvation, screening, dielectric value Non-polar groups –The hydrophobic effect

2. Hydrogen Bonding Atomic requirements –Donor: H attached to electronegative atom (typically O,N) [the book describes such groups as weak acids] –Acceptor: Lone pair on electronegative atom (e.g. O,N) [the book describes such groups as weak bases] Characteristics –Sort of intermediate between covalent and ionic: mainly electrostatic but with some directional (bonding) characteristics –Much weaker than covalent; even weaker than ionic, but ubiquitous

3. Hydrogen Bonding Geometry O N H C Ideal: roughly in-line Ideal: roughly according to lone pair geometry Distance between ‘heavy’ atoms: 2.8 to 3.5 Å The geometric parameters are much looser than for covalent bonding.

4. Hydrogen Bonding in Water In ice (see text, Fig. 2.3) In liquid water (see text, Fig. 2.4) Mechanism for proton transfer through chains of H-bonded water molecules (see text, Fig. 2.15)

5. Hydrogen Bonding in Water In ice –The diamond-like lattice in crystalline ice is unusually open, giving a lower density than liquid water, with important consequences for aquatic environments (see text, Fig. 2.3)

6. Hydrogen Bonding in Water In liquid water –rings are sometimes formed (this is sometimes a point of emphasis, though I don’t think it’s a particularly important point) –probably most important is the flexibility of the different H-bonding configurations possible; also the number of different partners (2 H- bonds from each O, one from each H, gives 4 partners in the best arrangements) (see text, Fig. 2.4)

7. Hydrogen Bonding in Water Mechanism for proton transfer through chains of H-bonded water molecules (see text, Fig. 2.15)

8. How H 2 0 interacts with ions: Solvation (see text, Figure 2.6) Screening and the dielectric value

9. How H 2 0 interacts with ions: Solvation –Water molecules orient themselves to make favorable electrostatic interactions with ions –This stabilizes the ions –Explains the solubility of ions in water (otherwise they would stay in an ionic solid where they make good ionic interactions) (see text, Fig. 2.6)

10. How H 2 0 interacts with ions: Screening –Owing to (1) the dipole moment of H 2 0 and (2) its ability to reorient, water molecules around and between ions screen or dampen the ionic forces between ions –The ability to screen ionic interactions is embodied in a term called the dielectric value or the dielectric constant, . –For water near 298K,  H20 = 78 (compared to 1 for a vacuum, or ~2 for organic solvents and hydrocarbons) i.e. ionic forces are diminished by a factor of 80 in water!

11. How H 2 0 interacts with non-polar solutes (the hydrophobic effect): Unfavorable transfer of non-polar groups from organic phase to water Non- polar solute organic aqueous  G 0 transfer = -RTlnK K=[solute] H20 /[ solute] org

12. How H 2 0 interacts with non-polar solutes (the hydrophobic effect): Mainly a surface effect –Presence of the non-polar group requires H 2 0 molecules to H- bond the each other in more ordered fashion than in bulk solution –The magnitude of the effect depends on the amount of surface area involved –Mainly an entropic effect (it’s not that the water has bad energetic interactions with the solute, but rather that its entropy is reduced) (see text, Fig. 2.8)

13. How H 2 0 interacts with non-polar solutes (the hydrophobic effect): Driving force for aggregation/association (minimizes surface area of interaction with water) Non-polar solutes or functional groups, or non- polar surfaces of larger molecules

14. How H 2 0 interacts with non-polar solutes (the hydrophobic effect): Biological relevance –Binding interactions between molecules –Protein folding –Formation of membrane lipid bilayer though water is described as the ‘universal solvent’, the insolubility of lipid molecules is one of the most vital features for cell structure: it makes the membrane possible

15. Acid-Base Equilibria in Water HA A - + H + or (HA + A + H + ) K a = [A - ][H + ]/[HA] [H + ] = K a [HA]/[A - ] (pH ≡ -log 10 [H + ], pK ≡ -log 10 K) pH = pK + log 10 [A - ]/[HA] (Henderson-Hasselbalch)

16. Acid-Base Equilibria in Water pH = pK + log 10 [A - ]/[HA] (Henderson-Hasselbalch) Explains: –buffering behavior of acid-base mixtures –distribution of acid vs. base forms of functional groups in molecules

17. Acid-Base Equilibria in Water pH = pK + log 10 [A - ]/[HA] (Henderson-Hasselbalch) –buffering behavior of acid-base mixtures when the numerator and denominator of the term in the log function are both similar, moderate changes to the numerator or denominator do not change the value of the log much, meaning the pH doesn’t change much when H+ or OH- are added, or produced by reactions in the cell the numerator and denominator are similar when the pH is near the pK, so buffers are only effective in a pH range near their pK value.

18. Acid-Base Equilibria in Water pH = pK + log 10 [A - ]/[HA] (Henderson-Hasselbalch) –distribution of acid vs. base forms of functional groups in molecules at pH 4, both forms are similarly present at pH 5, the base form is 10 times more abundant at pH 3, the acid form is 10 times more abundant The dependence of protonation state on pH (and environment) is important for how proteins function, since charged groups often play important roles in binding and catalysis. O C OH O C O-O- + H + pK  4