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

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

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

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

Hydrogen Bonds:

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

Biological Relevance of Hydrogen Bonds

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)

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

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

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.

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)

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

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-

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

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

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

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

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 

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.

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

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

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).

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)

Bound Water in Proteins

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.

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

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”

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

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

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+]

pH Scale: 1 unit = 10-fold

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!

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 = 0.001310, pH = 2.883 Ax2 + bx + c = 0

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 0.1 x x x = 0.00132, pH = 2.880

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

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

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.

Henderson–Hasselbalch Equation:Derivation

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

Water as a reactant in biochemistry

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