1. CHAPTER 2 Water and Aqueous Solutions
Learning Objectives
Types of non-covalent interactions between molecules
Properties of water – THE medium for life
Hydrophobic -- nonpolar -- moieties aggregate in water
Solute effects on bulk properties of water
Weak acids and bases
Buffers theory and practice
Water as participant in biochemical reactions

2. Physics of Non-covalent Interactions
Non-covalent interactions do not involve sharing a pair of electrons. Based on their physical origin, one can distinguish between
Ionic (Coulombic) Interactions
Electrostatic interactions between permanently charged species,
or between the ion and a permanent dipole
Dipole Interactions
Electrostatic interactions between uncharged, but polar molecules
Van der Waals Interactions
Weak interactions between all atoms, regardless of polarity
Attractive (dispersion) and repulsive (steric) component
Hydrophobic Effect
Complex phenomenon associated with the ordering of water molecules around non-polar substances

3. Noncovalent Forces and Interactions
Hydrogen bonds
Ion-Ion 1/r
Ion-dipole 1/r2
Dipole-dipole 1/r3
Dipole - Induced dipole - 1/r5
ID – ID (Van der Waals) - 1/r6
Hydrophobic

4. Hydrogen Bonds
Strong dipole-dipole or charge-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
Typically involves two electronegative atoms (frequently nitrogen and oxygen)
Hydrogen bonds are strongest when
the bonded molecules are oriented to
maximize electrostatic interaction.
Ideally the three atoms involved are in a line

5. Hydrogen Bonds:

6. 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 a substrates to enzymes
Binding of hormones to receptors
Matching of mRNA and tRNA

7. Biological Relevance of Hydrogen Bonds

8. 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 nm)
Repulsion dominates at very short distances
There is a minimum energy distance (van der Waals contact distance)

9. 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

10. Water is the Medium for Life
Life evolved in water (UV protection)
Organisms typically contain 70-90% water
Chemical reactions occur in aqueous milieu
Water is a critical determinant of the structure and function of proteins, nucleic acids, and membranes

11. Structure of the Water Molecule
Four electron pairs on four
sp3 orbitals (distorted tetrahedron)
Two pairs covalently link hydrogen atoms to a central oxygen atom.
Two remaining pairs remain nonbonding (lone pairs)
The electronegativity of the oxygen atom induces a net dipole moment
Water can serve as both a hydrogen bond donor and acceptor.

12. Hydrogen Bonding in Water
Up to four H-bonds H2O
high boiling point
high melting point
large surface tension
Hydrogen bonding in water is cooperative.
Hydrogen bonds between neighboring molecules are weak (20 kJ/mole) relative to the H–O covalent bonds (420 kJ/mol)

13. 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

15. Water Dissolves Many Salts
High dielectric constant of water (ε = 80 ) shields oppositely charged ions;
Almost no attraction > 40 nm
Electrostatics of solvation lowers the energy of the system
Entropy increases as ordered crystal lattice is dissolved
NaCl(s) <=>Na+ + Cl-

16. Ice: H2O(s)
Water has many different crystal forms; the hexagonal ice is the most common
Hexagonal ice forms a regular lattice, and thus has a low entropy
Hexagonal ice has lower density than liquid water; ice floats

17. The Hydrophobic Effect
Refers to the association or folding of non-polar 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 non-polar molecules

18. Solubility of Polar and Non-polar Solutes
Why are non-polar molecules poorly soluble in water?

19. Low Solubility of Hydrophobic Solutes
Disruption of H-bonded H2O networks
“Ordered” Water near a hydrophobic solute
Cavity formation in a medium with high surface tension

20. Hydrophobic Effect
Lipid molecules disperse in the solution; nonpolar tail of each lipid molecule is surrounded by ordered water molecules
Lipid aggregates – Water released, surface area reduced 

21. FIGURE 2-7b (part 3) Amphipathic compounds in aqueous solution
FIGURE 2-7b (part 3) 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.

22. 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

23. Colligative Properties
Some properties of solution — boiling point, melting point, and osmolarity — do not depend strongly on the nature of the dissolved substance. These are called colligative properties
Other properties — viscosity, surface tension, taste, and color, among other — depend strongly on the chemical nature of the solute. These are non-colligative properties.
Cytoplasm of cells are highly concentrated solutions and have high osmotic pressure

24. FIGURE 2-11 Osmosis and the measurement of osmotic pressure
FIGURE 2-11 Osmosis and the measurement of osmotic pressure. (a) The initial state. The tube contains an aqueous solution, the beaker contains pure water, and the semipermeable membrane allows the passage of water but not solute. Water flows from the beaker into the tube to equalize its concentration across the membrane. (b) The final state. Water has moved into the solution of the nonpermeant compound, diluting it and raising the column of water within the tube. At equilibrium, the force of gravity operating on the solution in the tube exactly balances the tendency of water to move into the tube, where its concentration is lower. (c) Osmotic pressure (Π) is measured as the force that must be applied to return the solution in the tube to the level of that in the beaker. This force is proportional to the height, h, of the column in (b).

25. Effect of Extracellular Osmolarity
Osmotic Pressure
For a single solute
Π = RT (ic)
Where i is extent of dissociation and c is concentration.
For mixtures
Π = RT Σ (ic)

26. Bound Water in Proteins

27. FIGURE 2-10 Water chain in cytochrome f
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 Figure 19-64). 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 Figure 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 Figure 2-13) through this chain of bound water molecules.

28. 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 H2O H+ + OH- is strongly to the left
Extent of dissociation depends on the temperature

29. 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”

30. Proton Hopping
Hydrogen bonded networks form natural chains for rapid Proton transfer

31. 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

32. 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
pH can be negative ([H+] = 6 M)
In neutral solution, [H+] = [OH-] and the pH is 7
pH = -log[H+]

33. pH Scale: 1 unit = 10-fold

34. 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!

35. Dissociation of Weak Electrolytes: Example
What is the final pH of a solution when 0.1 moles of acetic acid is adjusted to 1 L of water?
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
Ax2 + bx + c = 0

36. 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

37. pKa = -log Ka (strong acid  large Ka  small pKa)
pKa measures acidity
pKa = -log Ka (strong acid  large Ka  small pKa)

38. Buffers are mixtures of weak acids and their anions
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

39. FIGURE 2-16 The titration curve of acetic acid
FIGURE 2-16 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.

40. Henderson–Hasselbalch Equation:Derivation

41. 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

42. Water as a reactant in biochemistry

43. Chapter 2: Summary
The goal of this chapter was to help you to better understand:
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