There is, in fact, a ubiquitous force in nature, called the van der Waals (vdw) force, which is one of the main forces acting between molecules and is responsible for holding together many condensed phases, such as solid and liquid hydrocarbons and polymers. It is responsible for about one third of the attractive force holding liquid water molecules together. This force was actually first observed as a correction to the ideal gas equation and is attractive even between neutral gas molecules, such as oxygen and nitrogen, in a vacuum, which is why they can be liquified. Although electromagnetic in origin (as we will see later), it is much weaker than the coulombic force acting between ions, such as in salt crystals.
THE FORCES INVOLVED IN COLLOIDAL STABILITY
Although van der Waals forces will always act to coagulate dispersed colloids, it is possible to generate an opposing repulsive force of comparable strength. This force arises because most materials, when dispersed in water, ionise to some degree or selectively adsorb ions from the solution and hence become charged. Two similarly charged colloids will repel each other via an electrostatic repulsion, which will oppose coagulation. The stability of a colloidal solution is therefore critically dependent on the charge generated at the surface of the particles. The combination of these two forces, attractive van der Waals and repulsive electrostatic forces, forms the fundamental basis for our understanding of the behaviour and stability of colloidal solutions. The corresponding theory is referred to as the DLVO (after Derjaguin, Landau, Verwey and Overbeek) theory of colloid stability, which we will consider in greater detail later. The stability of any colloidal dispersion is thus determined by the behaviour of the surface of the particle via its surface charge and its short‐range attractive van der Waals force.
Our understanding of these forces has led to our ability to selectively control the electrostatic repulsion and so create a powerful mechanism for controlling the properties of colloidal solutions. As an example, if we have a valuable mineral imbedded in a quartz rock, grinding the rock will separate out both pure individual quartz and the mineral particles, which can both be dispersed in water. The valuable mineral can then be selectively coagulated, whilst leaving the unwanted quartz in solution. This process is used widely in the mining industry as the first stage of mineral separation. The alternative of chemical processing, for example, by dissolving the quartz in hydrofluoric acid, would be both expensive and environmentally unfriendly.
Figure 1.1 Scanning electron microscope image of dried mono‐disperse silica colloids.
It should be realised, at the outset, that colloidal solutions (unlike true solutions) will almost always be in a metastable state. That is, an electrostatic repulsion prevents the particles from combining into their most thermodynamically stable state of aggregation into the macroscopic form, from which the colloidal dispersion was (artificially) created in the first place. On drying, colloidal particles will often remain separated by these repulsive forces, as illustrated by the scanning electron microscope picture of mono‐disperse silica colloids.
TYPES OF COLLOIDAL SYSTEMS
The term ‘colloid’ usually refers to particles within the approximate size range of 50 Å to 50 μm, but this, of course, is somewhat arbitrary. For example, blood could be considered as a colloidal solution in which large blood cells are dispersed in water. They are stabilized by their negative charge, and so oppositely charged ions, for example, those produced by an alum stick, will coagulate the cells and hence stop bleeding. Often we are interested in solid dispersions in aqueous solutions, but many other situations are also of interest and industrial importance. Some examples are given in the following table.
The properties of colloidal dispersions are intimately linked to the high surface area of the dispersed phase and the chemistry of these interfaces. This linkage is well illustrated by the titles of two of the main journals in this area: the Journal of Colloid and Interface Science and Colloids and Surfaces. The natural combination of colloid and surface chemistry represents a major area of both research activity and industrial development. It has been estimated that something like 20 per cent of all chemists in industry work in this area. The more recent term nano is also applied to these small scale materials, because of their typical nanometre size.
Table 1.1
| Dispersed phase | Dispersion medium | Name | Examples |
|---|---|---|---|
| Liquid | Gas | Liquid aerosol | Fogs, sprays |
| Solid | Gas | Solid aerosol | Smoke, dust |
| Gas | Liquid | Foam | Foams |
| Liquid | Liquid | Emulsion | Milk, mayonnaise |
| Solid | Liquid | ‘Sol’, | Au sol, AgI sol. |
| Paste at high | Toothpaste | ||
| concentration | |||
| Gas | Solid | Solid foam | Expanded polystyrene |
| Liquid | Solid | Solid emulsion | Opal, pearl |
| Solid | Solid | Solid suspension | Pigmented plastics |
THE LINK BETWEEN COLLOIDS AND SURFACES
The link between colloids and surfaces follows naturally from the fact that particulate matter has a high surface area to mass ratio. The surface area of a 1 cm diameter sphere (4πr2) is 3.14 cm2, whereas the surface area of the same amount of material but in the form of 0.1 micron diameter spheres (i.e., the size of the particles in latex paint) is 314,000 cm2. The enormous difference in surface area is one of the reasons why the properties of the surface become very important for colloidal solutions. One everyday example is that organic dye molecules or pollutants can be effectively removed from water by adsorption onto particulate or granular activated charcoal because of its high surface area. This process is widely used for water purification and in the oral treatment of poison victims.
Although it is easy to see that surface properties will determine the stability of colloidal dispersions, it is not so obvious why this can also be the case for some properties of macroscopic objects. As one important illustration, consider the interface between a liquid and its vapour: