1. CHAPTER 2 Water and Aqueous Solutions
Learning goals:
What kind of interactions occur between molecules
Why water is a good medium for life
Why nonpolar moieties aggregate in water
How dissolved molecules alter properties of water
How weak acids and bases behave in water
How buffers work and why we need them
How water participates in biochemical reactions

2. Structure of the Water Molecule
Octet rule dictates that there are four electron pairs around an oxygen atom in water
These electrons are in four sp3 orbitals
Two of these pairs covalently link two hydrogen atoms to a central oxygen atom
The two remaining pairs remain nonbonding (lone pairs)
Water geometry is a distorted tetrahedron
The electronegativity of the oxygen atom induces a net dipole moment
Because of the dipole moment, water can serve as both a hydrogen bond donor and acceptor

3. Hydrogen Bonds
Strong dipole-dipole interaction that arises between an acid (proton donor) and a base (proton acceptor)
Typically 4–6 kJ/mol for bonds with neutral atoms,
and 6–10 kJ/mol for bonds with one charged atom
Can only form between H and N, O, F
Hydrogen bonds are strongest when the bonded molecules are oriented to maximize electrostatic interaction
Ideally the three atoms involved are in a line

4. Hydrogen Bonding in Water
Water can serve as both
an H donor
an H acceptor
Up to four H-bonds per water molecule gives water its
anomalously high boiling point
anomalously high melting point
unusually large surface tension
Hydrogen bonding in water is cooperative
Hydrogen bonds between neighboring molecules are weak (20 kJ/mol) relative to the H–O covalent bonds (420 kJ/mol)

5. FIGURE 2-1b Structure of the water molecule
FIGURE 2-1b Structure of the water molecule. (b) Two H2O molecules joined by a hydrogen bond (designated here, and throughout this book, by three blue lines) between the oxygen atom of the upper molecule and a hydrogen atom of the lower one. Hydrogen bonds are longer and weaker than covalent O—H bonds.

6. Why is this? The interaction energies and entropies are too low
Water as a Solvent
Water is a good solvent for charged and polar substances
amino acids and peptides
small alcohols
carbohydrates
Water is a poor solvent for nonpolar substances
nonpolar gases
aromatic moieties
aliphatic chains
Why is this? The interaction energies and entropies are too low

8. Water dissolves many salts
High dielectric constant reduces attraction between oppositely charged ions in salt crystal; almost no attraction at large (> 40 nm) distances
Essentially, the water “shorts out” the interaction b/w two ions
Strong electrostatic interactions between the solvated ions and water molecules lower the energy of the system
Entropy increases as ordered crystal lattice is dissolved

9. FIGURE 2-6 Water as solvent
FIGURE 2-6 Water as solvent. Water dissolves many crystalline salts by hydrating their component ions. The NaCl crystal lattice is disrupted as water molecules cluster about the C– and Na+ ions. The ionic charges are partially neutralized, and the electrostatic attractions necessary for lattice formation are weakened.

10. Backup: Intermolecular Forces
Four Basic Types of Non Covalent Interactions
Ion-Dipole (z1m2/r3)
Dipole-Dipole (m1m2/r3)
Dipole-Induced dipole (m1a2/r6)
London Forces (a1a2/r6)
The strength of the interactions in each force is inversely proportional to the distance between the interacting species (That’s the 1/r term next to each name)

11. Coulomb’s Law and Interacting Species
The force between two interacting ions is equal to the magnitude of the charges divided by the square of the distance between them.
k = proportionality constant = 8.99x109 JmC-2
D is the dielectric constant
The dielectric constant is a measure of the ability of the solvent or medium containing the charges to keep them apart (D for a vacuum is 1, D for water is ~80, D for a brick wall is ∞)

12. Interion and Intermolecular Forces

13. Physics of Noncovalent Interactions
Noncovalent interactions do not involve sharing a pair of electrons. Based on their physical origin, one can distinguish between:
Ionic (Coulombic) Interactions (Ionic and Ion/Dipole)
Electrostatic interactions between permanently charged species, or between the ion and a permanent dipole
Dipole Interactions (Dipole/Dipole and Dipole/Ind. Dipole)
Electrostatic interactions between uncharged, but polar molecules
van der Waals Interactions (London Forces)
Weak interactions between all atoms, regardless of polarity
Attractive (dispersion) and repulsive (steric) component
Hydrophobic Effect (Rationale behind a lot of Biochemistry)
Complex phenomenon associated with the ordering of water molecules around nonpolar substances

14. Examples of Noncovalent Interactions
H-bonds are a special type of dipole/dipole interaction
Between amino acids = Salt Bridges
Hydrophobic, Dipole/Induced Dipole and London Force Interactions Combined make van der Waals
“Hydrophobic Exclusions”

15. Importance of Hydrogen Bonds
Source of unique properties of water
Structure and function of proteins
Structure and function of DNA
Structure and function of polysaccharides
Binding of substrates to enzymes
Binding of hormones to receptors
Matching of mRNA and tRNA
“I believe that as the methods of structural chemistry are further applied to physiological problems, it will be found that the significance of the hydrogen bond for physiology is greater than that of any other single structural feature.”
–Linus Pauling, The Nature of the Chemical Bond, 1939

16. Hydrogen Bonds: Examples
FIGURE 2-3 Common hydrogen bonds in biological systems. The hydrogen acceptor is usually oxygen or nitrogen; the hydrogen donor is another electronegative atom.

17. Biological Relevance of Hydrogen Bonds
FIGURE 2-4 Some biologically important hydrogen bonds.

18. van der Waals Interactions
van der Waals interactions have two components:
Attractive force (London dispersion) depends on the polarizability
Repulsive force (Steric repulsion) depends on the size of atoms
Attraction dominates at longer distances (typically 0.4–0.7 nm)
Repulsion dominates at very short distances
There is a minimum energy distance (van der Waals contact distance)

20. Origin of the London Dispersion Force
Quantum mechanical origin
Instantaneous polarization by fluctuating charge distributions
Universal and always attractive
Stronger in polarizable molecules
Important only at a short range

21. Biochemical Significance of van der Waals Interactions
Weak individually
easily broken, reversible
Universal
occur between any two atoms that are near each other
Importance
determines steric complementarity
stabilizes biological macromolecules (stacking in DNA)
facilitates binding of polarizable ligands

22. The Hydrophobic Effect
Refers to the association or folding of nonpolar molecules in the aqueous solution
Is one of the main factors behind:
protein folding
protein-protein association
formation of lipid micelles
binding of steroid hormones to their receptors
Does not arise because of some attractive direct force between two nonpolar molecules

23. Solubility of Polar and Nonpolar Solutes
Why are nonpolar molecules poorly soluble in water?

24. Low solubility of hydrophobic solutes can be explained by entropy
Bulk water has little order:
– high entropy
Water near a hydrophobic solute is highly ordered:
– low entropy
Low entropy is thermodynamically unfavorable, thus hydrophobic solutes have low solubility.
Remember to consider the Entropy of the Solute AND the Entropy of the Solvent since both are part of the System
Seemingly unfavorable or unlikely Solute behaviour can often be explained when considering the changes in the Solvent behaviour.

25. Water surrounding nonpolar solutes has lower entropy
FIGURE 2-7a Amphipathic compounds in aqueous solution. (a) Long-chain fatty acids have very hydrophobic alkyl chains, each of which is surrounded by a layer of highly ordered water molecules.
Clathrate Cages

26. Origin of the Hydrophobic Effect (1)
Consider amphipathic lipids in water
Lipid molecules disperse in the solution; nonpolar tail of each lipid molecule is surrounded by ordered water molecules
Entropy of the system decreases
System is now in an unfavorable state

27. FIGURE 2-7b (part 1) Amphipathic compounds in aqueous solution
FIGURE 2-7b (part 1) Amphipathic compounds in aqueous solution. (b) By clustering together in micelles, the fatty acid molecules expose the smallest possible hydrophobic surface area to the water, and fewer water molecules are required in the shell of ordered water. The energy gained by freeing immobilized water molecules stabilizes the micelle.

28. Origin of the Hydrophobic Effect (2)
Nonpolar portions of the amphipathic molecule aggregate so that fewer water molecules are ordered
The released water molecules will be more random and the entropy increases
All nonpolar groups are sequestered from water, and the released water molecules increase the entropy further
Only polar “head groups” are exposed and make energetically favorable H-bonds

29. FIGURE 2-7b (part 2) Amphipathic compounds in aqueous solution
FIGURE 2-7b (part 2) Amphipathic compounds in aqueous solution. (b) By clustering together in micelles, the fatty acid molecules expose the smallest possible hydrophobic surface area to the water, and fewer water molecules are required in the shell of ordered water. The energy gained by freeing immobilized water molecules stabilizes the micelle.

30. Hydrophobic effect favors ligand binding
Binding sites in enzymes and receptors are often hydrophobic
Such sites can bind hydrophobic substrates and ligands such as steroid hormones
Many drugs are designed to take advantage of the hydrophobic effect

31. Ionization of Water  H2O H+ + OH- 
O-H bonds are polar and can dissociate heterolytically
Products are a proton (H+) and a hydroxide ion (OH–)
Dissociation of water is a rapid reversible process
Most water molecules remain un-ionized, thus pure water has very low electrical conductivity (resistance: 18 M•cm)
The equilibrium is strongly to the left
Extent of dissociation depends on the temperature

32. Proton Hydration Protons do not exist free in solution.
They are immediately hydrated to form hydronium (oxonium) ions.
A hydronium ion is a water molecule with a proton associated with one of the non-bonding electron pairs.
Hydronium ions are solvated by nearby water molecules.
The covalent and hydrogen bonds are interchangeable. This allows for an extremely fast mobility of protons in water via “proton hopping.”

33. Proton Hopping
FIGURE 2-14 Proton hopping. Short “hops” of protons between a series of hydrogen-bonded water molecules result in an extremely rapid net movement of a proton over a long distance. As a hydronium ion (upper left) gives up a proton, a water molecule some distance away (lower right) acquires one, becoming a hydronium ion. Proton hopping is much faster than true diffusion and explains the remarkably high ionic mobility of H+ ions compared with other monovalent cations such as Na+ and K+.

34. Water bound to proteins is essential for their function
FIGURE 2-9 Water binding in hemoglobin. (PDB ID 1A3N) The crystal structure of hemoglobin, shown (a) with bound water molecules (red spheres) and (b) without the water molecules. The water molecules are so firmly bound to the protein that they affect the x-ray diffraction pattern as though they were fixed parts of the crystal. The two α subunits of hemoglobin are shown in gray, the two β subunits in blue. Each subunit has a bound heme group (red stick structure), visible only in the β subunits in this view. The structure and function of hemoglobin are discussed in detail in Chapter 5.

35. Water bound to proteins is essential for their function
FIGURE 2–10 Water chain in cytochrome f. Water is bound in a proton channel of the membrane protein cytochrome f, which is part of the
energy-trapping machinery of photosynthesis in chloroplasts (see Fig. 19–61). Five water molecules are hydrogen-bonded to each other and to
functional groups of the protein: the peptide backbone atoms of valine, proline, arginine, and alanine residues, and the side chains of three
asparagine and two glutamine residues. The protein has a bound heme (see Fig. 5–1), its iron ion facilitating electron flow during photosynthesis.
Electron flow is coupled to the movement of protons across the membrane, which probably involves “proton hopping” (see Fig. 2–14) through
this chain of bound water molecules.

36. Water bound to proteins is essential for their function
FIGURE 2–11 Hydrogen-bonded water as part of a protein's sugarbinding site. In the L-arabinose-binding protein of the bacterium E. coli, five water molecules are essential components of the hydrogen-bonded network of interactions between the sugar arabinose (center) and at
least 13 amino acid residues in the sugar-binding site. Viewed in three dimensions, these interacting groups constitute two layers of binding
moieties; amino acid residues in the first layer are screened in red, those in the second layer in green. Some of the hydrogen bonds are drawn longer than others for clarity; they are not actually longer than the others.

37. Ionization of Water: Quantitative Treatment
Concentrations of participating species in an equilibrium process
are not independent but are related via the equilibrium constant:
[H+]•[OH-] 
H2O H+ + OH- 
Keq = ————
[H2O]
Keq can be determined experimentally, it is 1.8•10–16 M at 25C.
[H2O] can be determined from water density, it is 55.5 M.
Ionic product of water:
In pure water [H+] = [OH–] = 10–7 M

38. What is pH?
pH is defined as the negative logarithm of the hydrogen ion concentration
Simplifies equations
The pH and pOH must always add to 14
In neutral solution, [H+] = [OH–] and the pH is 7
pH can be negative ([H+] = 6 M)
pH = -log[H+]

39. pH scale is logarithmic:
1 unit = 10-fold

40. Dissociation of Weak Electrolytes: Principle
Weak electrolytes dissociate only partially in water.
Extent of dissociation is determined by the acid dissociation constant Ka.
We can calculate the pH if the Ka is known. But some algebra is needed!

41. Dissociation of Weak Electrolytes: Example
What is the final pH of a solution when 0.1 moles of acetic acid is added to water to a final volume of 1L? The Ka of acetic acid is 1.74x10-5
We assume that the only source of H+ is the weak acid
To find the [H+], a quadratic equation must be solved
0.1 – x x x
x = , pH = 2.883

42. Dissociation of Weak Electrolytes: Simplification
The equation can be simplified if the amount of dissociated species is much less than the amount of undissociated acid
Approximation works for sufficiently weak acids and bases
Check that x < [total acid]
0.1 – x x x
x x
x = , pH = 2.880
Nearly identical to the answer determined by the quadratic formula!

43. pKa measures acidity
pKa = –log Ka (strong acid  large Ka  small pKa)
FIGURE 2-16 Conjugate acid-base pairs consist of a proton donor and a proton acceptor. Some compounds, such as acetic acid and ammonium ion, are monoprotic; they can give up only one proton. Others are diprotic (carbonic acid and glycine) or triprotic (phosphoric acid). The dissociation reactions for each pair are shown where they occur along a pH gradient. The equilibrium or dissociation constant (Ka) and its negative logarithm, the pKa, are shown for each reaction. *For an explanation of apparent discrepancies in pKa values for carbonic acid (H2CO3), see p. 63.

45. Buffers are mixtures of weak acids and their anions (conjugate base)
Buffers resist change in pH
At pH = pKa, there is a 50:50 mixture of acid and anion forms of the compound
Buffering capacity of acid/anion system is greatest at pH = pKa
Buffering capacity is lost when the pH differs from pKa by more than 1 pH unit

46. Acetic Acid-Acetate as a Buffer System
FIGURE 2–19 The acetic acid–acetate pair as a buffer system. The system is capable of absorbing either H+ or OH- through the reversibility of the dissociation of acetic acid. The proton donor, acetic acid (HAc), contains a reserve of bound H+, which can be released to neutralize an addition of OH- to the system, forming H2O. This happens because the product [H+][OH-] transiently exceeds Kw (1 x M2). The equilibrium quickly adjusts to restore the product to 1 x M2 (at 25 °C), thus transiently reducing the concentration of H+. But now the quotient [H+][Ac-]/[HAc] is less than Ka, so HAc dissociates further to restore equilibrium. Similarly, the conjugate base, Ac-, can react with H+ ions added to the system; again, the two ionization reactions simultaneously come to equilibrium. Thus a conjugate acid-base pair, such as acetic
acid and acetate ion, tends to resist a change in pH when small amounts of acid or base are added. Buffering action is simply the consequence of two reversible reactions taking place simultaneously and reaching their points of equilibrium as governed by their equilibrium constants, Kw and Ka.

47. FIGURE 2-17 The titration curve of acetic acid
FIGURE 2-17 The titration curve of acetic acid. After addition of each increment of NaOH to the acetic acid solution, the pH of the mixture is measured. This value is plotted against the amount of NaOH added, expressed as a fraction of the total NaOH required to convert all the acetic acid (CH3COOH) to its deprotonated form, acetate (CH3COO–). The points so obtained yield the titration curve. Shown in the boxes are the predominant ionic forms at the points designated. At the midpoint of the titration, the concentrations of the proton donor and proton acceptor are equal, and the pH is numerically equal to the pKa. The shaded zone is the useful region of buffering power, generally between 10% and 90% titration of the weak acid.

48. Weak acids have different pKas
FIGURE 2–18 Comparison of the titration curves of three weak acids. Shown here are the titration curves for CH3COOH, H2PO-4, and NH+4 . The predominant ionic forms at designated points in the titration are given in boxes. The regions of buffering capacity are indicated at the right. Conjugate acid-base pairs are effective buffers between approximately 10% and 90% neutralization of the proton-donor species.

49. Henderson–Hasselbalch Equation: How to Find the pH of a weak acid during its titration
For a weak acid:
HA H+ + A-  
and, therefore
We want to know the pH, so solve for [H+]

50. How to solve any weak acid/base problem
Identify the Acid and the Base and the Ionization Reaction
2) If it is a titration:
(Eg: Find the pH at each of the following points in the titration of 25 mL of 0.3 M Acetic acid with 0.3 M NaOH. The Ka value is 1.74x10-5: The initial pH, After adding 10 mL of 0.3 M NaOH, After adding mL of 0.3 M NaOH, After adding 25 mL of 0.3 M NaOH, After adding 26 mL of 0.3 M NaOH)
Calculate how much of your starting material was reacted and subtract that from the starting number of moles of acid or base
Determine the new total volume and calculate the molarity of the acid or base
Use the Ka equation or the Henderson-Hasselbach equation to determine the molarity of the remaining species in the beaker or the pH
If it is not a titration (Eg: What is the concentration of acetate in 500 mL of a 0.05M Acetate buffer at pH 5.5?)
Determine how many moles of acid or base you have

51. Biological Buffer Systems
Maintenance of intracellular pH is vital to all cells
Enzyme-catalyzed reactions have optimal pH
Solubility of polar molecules depends on H-bond donors and acceptors
Equilibrium between CO2 gas and dissolved HCO3– depends on pH
Buffer systems in vivo are mainly based on
phosphate, concentration in millimolar range
bicarbonate, important for blood plasma
histidine, efficient buffer at neutral pH
Buffer systems in vitro are often based on sulfonic acids of cyclic amines
HEPES
PIPES
CHES

52. Chapter 2: Summary In this chapter, we learned about:
The nature of intermolecular forces
The properties and structure of liquid water
The behavior of weak acids and bases in water
The way water can participate in biochemical reactions