Biol/Chem 5310

Lecture: 2

August 27, 2002

Aqueous Solutions

Properties of Water

  • Water is the natural solvent for most biological molecules
  • Water is so common that its properties tend to be overlooked
  • Compare to NH3, CH4, and H2S
  • Highly polar---electronegativity difference between O and H is great

    • Hydrogen Bonds

  • Formed between H2O molecules, it is an electrostatic interaction
  • Donor is H, Acceptor is O
  • Vander waals contact distance (H to O) is 2.6 Å
  • H-bond distance is 1.8 Å, Intermediate between vanderwaals and covalent bond
  • covalent bond distance (H to O) is 0.958 Å
  • Energy is much less than covalent bond (20kJ/mol vs. 460 kJ/mol)
  • Other molecules with donors and acceptors can form H-bonds

    • Physical Properties of H2O

  • High melting point, boiling point
  • High heats of vaporization, sublimation
  • Expansion upon freezing
  • Properties are explained by networks of H-bonds in water.

    • Solvent properties

  • Water hydrates ions
  • The hydration of ionic species shields them from each other, and so diminishes the force of attraction between them that is described by Coulombs Law.
  • F is the Force, q's are the charges, k is the proportionality constant, r is the distance between charges, and D is the dielectric constant.
  • The dielectric constant of water is very high
  • Water forms H-bonds with other molecules in solution.
  • Water forces nonpolar molecules to self-associate: the hydrophobic effect
  • This is an entropy-driven process.
  • Proton mobility in water is very high relative to other cations.

    • Ionization of Water

  • Bronsted-Lowry formulation: Acid donates a proton, Base accepts a proton.

    • Strong acids/bases: complete dissociation

  • Strength determined by acid dissociation constant
  • Some acids dissociate completely, transfer all H+ to water

    • Weak acids/bases: partial dissociation

  • Weak acids, such as carboxyl and amino groups, are very important in biochemistry
  • Many biological molecules contain carboxyl or amino groups
  • Buffers are used to maintain constant pH, in vivo and in vitro
  • e.g., ionization of an amino acid

    • Henderson-Hasselbalch equation

    H-H equation is obtained from this:

  • Relates the pH of the solution to the pKa of the weak acid and the ratio of the dissociated to the undissociated forms.
  • The pH of a solution of an amino acid will depend upon the relative amounts of the various species (i.e. the different protonation states)
  • Assuming that we know the pKa, we can calculate the pH if we know the concentration of all the buffer species.   Alternatively, if the pH is known, we can determine the composition of the buffer.

    • Titrations

  • Show the change in pH as a function of the amount of acid or base added.
  • The midpoint of the titration represents the pKa, when [HA] = [A-]
  • A buffer is considered to be effective at a pH within 1 unit of its pKa value.

    • Animation of Fig. 2-15 Titration Curves for Acetic Acid, Phosphate and Ammonia

    Animation of Fig. 2-16 Titration of a Polyprotic Acid

    Link to an interesting site about Water

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    Last updated Friday, August 23, 2002

    Comments/questions: svik@mail.smu.edu

    Copyright 2002, Steven B. Vik, Southern Methodist University