Aqueous Solutions

  1. Aqueous solutions
    1. Most biochemical reactions involve solutes dissolved in water.
    2. Two important properties of aqueous solutions:
      1. Solute concentration.
      2. pH.
    3. Solute Concentration
      1. Molecular weight = Sum of the weight of all atoms in a molecule (expressed in Daltons).
      2. Mole = Amount of a substance that has a mass in grams numerically equivalent to its molecular weight in Daltons.
      3. Molarity = Number of moles of solute per liter of solution.
      4. Advantage to measuring in moles:
        1. Rescales weighing of single molecules in Daltons in grams, which is more practical for laboratory use.
        2. A mole of one substance has the same number of molecules as a mole of any other substance (6.02 x 1023, Avogadro's number).
        3. Allows one to combine substances in fixed ratios of molecules.
    4. Acids, Bases and pH
      1. Dissociation of Water Molecules: Occasionally, the hydrogen atom that is shared in a hydrogen bond between two water molecules, shifts from the oxygen atom to which it is covalently bonded to the unshared orbitals of the oxygen atom to which it is hydrogen bonded.
        1. Only a hydrogen ion (proton with a +1 charge) is actually transferred.
        2. Transferred proton binds to an unshared orbital of the second water molecule creating a hydronium ion (H3O+).
        3. Water molecule that lost a proton has a net negative charge and is called a hydroxide ion (OH-).
        4. By convention, ionization of water is expressed as the dissociation into H+ and OH-.
        5. Reaction is reversible.
        6. At equilibrium, most of the water is not ionized.
      2. Acids and Bases--At equilibrium in pure water:
        1. Number of H+ ions = number of OH- ions.
        2. NOTE: This is a good place to point out how few water molecules are actually dissociated (only 1 out of 554,000,000 molecules).
          1. ACID
            • Substance that increases the relative [H+] of a solution.
            • Also removes OH- because it tends to combine with H+ to form water.
          2. BASE
            • Substance that reduces that relative [H+] of a solution.
            • May alternately increase [OH-].
            • For example:
            • A base may reduce [H+] directly; NH3 + H+ --> NH4
            • A base may reduce [H+] indirectly; NaOH --> Na+ + OH- then OH- + H+ --> H2O
          3. A solution in which: (Brackets indicate molar concentration.)
            • [H+] = [OH-] is a neutral solution.
            • [H+] greater than [OH-] is an acidic solution.
            • [H+] less than [OH-] is a basic solution.
        3. The pH Scale: In any aqueous solution: [H+][OH-] = 10e-14 Molar
        4. For example:
          1. In a neutral solution, [H+] = 10-7 M and [OH-] = 10-7 M.
          2. In an acid solution where the [H+] = 10-5 M and [OH-] = 10-9 M.
          3. In a basic solution where the [H+] = 10-9 M and [OH-] = 10-5 M.
        5. pH scale = Scale used to measure degree of acidity. It ranges from 0 to 14.
          1. pH = Negative log10 of the [H+] expressed in moles per liter.
          2. pH of 7 is a neutral solution.
          3. pH < 7 is an acidic solution.
          4. pH > 7 is a basic solution.
          5. Most biological fluids were within the pH range of 6 to 8. There are some exceptions such as stomach acid with pH = 1.5.
          6. Each pH unit represents a tenfold difference (scale is logarithmic), so a slight change in pH represents a large change in actual [H+].
    5. Buffers
      1. Organisms must maintain pH of body fluids within a narrow range (usually pH 6-8).
      2. Buffer = Substance that prevents large sudden changes in pH.
        1. Are combinations of H+-donor and H+-acceptor forms of weak acids or bases.
      3. Buffers work by:
        1. Accepting H+ ions from solution when they are in excess.
        2. Donating H+ ions to the solution when they have been depleted.
        3. For example: Explain bicarbonate buffer.