interfacial phenomena

  1. Background

    When phases exist together, the boundary between two of them is termed an interface.
    Several types of interface can exist, depending on whether the two adjacent phases are in the solid, liquid, or gaseous state.
    Two most common type of interfaces are liquid interfaces and solid interfaces.

  2. Liquid Interfaces

    Surface and Interfacial Tensions: In the liquid state the cohesive forces between adjacent molecules are well developed. In a liquid drop suspended in air, molecules in the bulk of the liquid (middle of the liquid) are surrounded in all directions by other molecules for which they have an equal attraction.

    Molecules at the surface (i.e, attractive cohesive forces with other liquid molecules that are situated below and adjacent to them) They can develop adhesive forces of attraction with the molecules constituting the other phase involved in the interface, although, in the case of the liquid/gas interface, this adhesive force of attraction is small. The net effect is that the molecules at the surface of the liquid experience an inward force towards the bulk or middle of the liquid. Such a force pulls the molecules of the interface together and as a result contracts the surface. Liquid droplets therefore tend to assume a spherical shape since a sphere has the smallest surface area per unit volume. This tension in the surface is the force per unit length that must be applied parallel to the surface so as to counterbalance the net inward pull. This force is called surface tension, similar to a situation that exists when an object dangling over the edge of a cliff on a length of rope is pulled upwards by a man holding the rope and walking away from the edge of the top of the cliff.

    Interfacial tension is the force per unit length existing at the interface between two immiscible liquid phases. Interfacial tensions are less than surface tensions because the adhesive forces between two liquid phases forming an interface are greater than when a liquid and a gas phase exist together. It follows that if two liquids are completely miscible, no interfacial tension exists between them.

  3. Adsorption at Liquid Interfaces

    Surface free energy is defined as the work that must be done in order to increase the surface by unit area. As a result of such expansion, more molecules must be brought from the bulk to the interface. The more work that has to be expended to achieve this, the greater the surface free energy.

    Certain molecules and ions, when dispersed in the liquid, move of their own accord to the interface. Their concentration at the interface then exceeds their concentration in the bulk of the liquid the surface free energy and surface tension is reduced.

    Such a phenomenon, where the added molecules are partitioned in favor of the interface, is termed positive adsorption. Other materials (such as inorganic electrolytes) are partitioned in favor of the bulk (or centre of the liquid) leading to negative adsorption, and a corresponding increase in surface free energy and surface tension. Adsorption can also occur at solid interfaces.

    Adsorption should not be confused with absorption. Adsorption is solely a surface effect, whereas in absorption the liquid or gas being absorbed penetrates into the capillary spaces of the absorbent. The taking up of water by a sponge is absorption; the concentrating of alkaloid molecules on the surface of clay is adsorption.

    Molecules and ions that are adsorbed at interfaces are termed surface-active agents or surfactants. An alternative expression is amphiphile which suggests that the molecule or ion has a certain affinity for both polar and nonpolar solvents.

  4. Micelle Formation

    At low concentrations, the lipophilic portions of the SAA molecules "lie" on the surface of the aqueous phase. As the surfactant concentration increases, the degree of structure of the interface layer increases so that the maximum number of surfactant molecules is incorporated into that layer.

    When the interface is saturated, surfactant molecules form self-association referred to as a micelle, in which the lipophilic portion of the molecules form the interior of the structure and are surrounded by the hydrophilic heads of the molecules.

    See the following interactive module for an illustration of the concept.

    Micelles are referred to as association colloids since:

    1. They are aggregates of surfactant molecules.
    2. They are colloidal in size.

    Upon addition of surfactant to water, the overall concentration rises until two limits are reached:

    1. Maximum solubility of the monomer in water.
    2. Saturation of the interface (i.e. formation of a monomolecular layer of maximum possible desnity at the interface.)

    Formation of the monomolecular layer at the interface can be considered a entropy driven since removal of the hydrocarbon chains release the structured water which surrounded each chain in solution. Therefore, when the interface is saturated, addition of further surfactants would cause a drop in entropy if the additional monomers were to exist in solution with the inevitable layer of ordered water molecules surrounding their hydrophobic regions.

    They are roughly spherical at concentrations slightly above the CMC but shift to oval and eventually laminar shapes as they deform in order to accommodate more surfactant molecules.

    In a lipophilic phase additional surfactant molecules aggregate to form Reverse-Micelles following saturation of the interface. A hydrophilic core is formed from the hydrophilic head groups surrounded by the lipophilic tails.

    The following equilibria exist in a two phase system.

    In the aqueous phase

    Dissolved monomers <> Monomers in the interfacial layer

    Dissolved monomers <> Monomers in the micelles

    In the lipid phase

    Dissolved monomers <> Monomers in the interfacial layer

    Dissolved monomers <> Monomers in the reverse micelles

    Since the interfacial layer is a link between phases, it is possible for one molecule to move from one phase to the other by means of the existing equilibria.

  5. Considerations In Using Solubilized Drugs

    1. Toxicity of the surfactant.
      1. Due to the compound itself.
      2. Due to the physical effects on the patient, particularly membrane effects.
    2. Disagreeable odour or taste it used orally (after all, they are soaps).
    3. Drug toxicity it there is large increases in bioavailability,
    4. Effects on drug stability - micellar catalysis or stabilization.
    5. Physical stability of the system since small changes may affect the solubilizing efficiency of the vehicle.

  6. Temperature Effects on Surfactant Solubility

    1. The Krafft Point

      Definition: The temperature at which the solubility equals the Critical Micelle Concentration (CMC).

      Below the Krafft, ft is possible that even at the maximum solubility of the surfactant, the interface may not be saturated, therefore there is no reason for micelles to form. Above the Krafft point, micelles will form and due to their high solubility, there will be a dramatic increase in surfactant solubility.

      NB: this effect is not seen for all surfactants.

    2. The Cloud Point

      Definition: The temperature above which some surfactants begin to precipitate.

      For non-ionic surfactants whose hydrophilic portions consist of long hydrophilic chains, part of the molecule's solubility results from the hydration of those chains by water molecules in the solution.

      Increasing temperatures impart sufficient kinetic energy to the hydrating molecules so they are effectively lost into the bulk water. Their loss can produce a sufficient overall drop in the solubility of the surfactant that precipitation can occur.