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CHAPTER 2 Water and Aqueous Solutions. 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

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chapter 2 water and aqueous solutions
CHAPTER 2Water and Aqueous Solutions
  • 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

Learning Objectives

physics of non covalent interactions
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
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
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

importance of 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
van der waals interactions
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

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
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
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
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 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
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 h 2 o s

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
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
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

Low Solubility of Hydrophobic Solutes

hydrophobic effect

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

hydrophobic effect favors ligand binding

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
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
effect of extracellular osmolarity
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)

ionization of water

H2O H+ + OH-

Ionization of Water

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
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

Proton Hopping

Hydrogen bonded networks form natural chains for rapid Proton transfer

ionization of water quantitative treatment
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
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+]

dissociation of weak electrolytes principle
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
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

Ax2 + bx + c = 0

dissociation of weak electrolytes simplification
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

p k a measures acidity

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

Buffers are mixtures of weak acids and their anions

biological buffer systems
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
chapter 2 summary
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