Chem1 General Chemistry Virtual Textbook → States of matter → Colloids

States of matter:

Colloids: between the big and the small

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Computer-generated view of self-assembled micelles of surfactant molecules (lyotropic liquid crystal) obtained from neutron diffraction data. [source]

Individual molecules and ions are small — too small to see even through an optical microscope.

Sand, salt, and chalk dust are made up of tiny chunks of solid particles, each containing huge numbers of molecules. You can usually see the individual particles directly, although the smallest ones might require some magnification. There is, however, a vast but largely hidden world in between: particles so tiny that they cannot be resolved by an optical microscope, or molecules so large that they begin to constitute a phase of their own when they are suspended in a liquid.

This is the world of colloids which we will survey in this lesson.

Phases.... define, contrast large particle that can settle with molecules in solution

Somewhere between the two extremes of a solution and of a multiphase (heterogeneous) mixture, we encounter the realm of the colloidal dispersion.

Colloidal dispersions behave very much like solutions in that they appear to be homogeneous and the particles do not settle out under the influence of gravity. However, they exhibit a number of unique properties that distinguish them from true solutions.

Surface area increase with smaller size....

Although mass is conserved, surface area is not; as a solid is sliced up into smaller bits, more surfaces are created. These new surfaces are smaller, but their areas increase as the square of the particle size, so the total surface area increases faster than the particle size decreases.

... cube 6 sq cm sliced 10 in each direct give 1000 cm...Each new cube has a length of 0.10 cm, and thus a surface area
of 6 x (0.10)Û cmÛ = 0.06 cmÛ.
There are 1000 of these smaller cubes, so the total surface area is now 60 cmÛ -- quite a bit larger than it was originally!

Slicing the 1-cm cube into 1000 smaller 0.1-cm cubes, each having a surface area of 0.06 cmÛ, yields a total surface area of 60 cmÛ. The table shows how dividing the cube up even more finely can increase the surface to very large values. In practical situations, one can get areas of acres (hectares) per mole!

table: slices facelen,cm no_cubes surfarea

1 1 1 6

10 .1 1000 60 / 60,000cm /6000m etc

Now let's turn to the optical properties of colloids. Colloidal dispersions are distinguished from true solutions by their light-scattering properties. A beam of light passing through a solution tends retain its shape, as shown at the top.

John Tyndall discovered this effect in 1869. Tyndall scattering
(as it is more precisely known) scatters all wavelengths equally. This is in contrast to Rayleigh scattering, which scatters shorter wavelengths more, bringing us blue skies and red sunsets.

Tyndall scattering can be seen even in dispersions that are transparent. As the density of particles (or the particle size) increases, the light scattering may become great enough to produce a "cloudy" effect.

This is the reason that milk, fog, and clouds themselves appear to be white.
The individual water droplets in clouds (or the butterfat droplets in milk) are actually transparent, but the intense light scattering disperses the light in all directions, preventing us from seeing through them.

Colloidal particles are, like molecules, too small to be visible though an ordinary optical microscope. However, if one looks in a direction perpendicular to the light beam, a colloidal particle will "appear" over a dark background as a tiny speck due to the Tyndall scattering. A microscope specially designed for this application is known as an _______microcope. Bear in mind that the ultramicroscope (invented in Austria in 1902) does not really allow us to "see" the particle; the scattered light merely indicates where it is at any given moment.

If you observe a single colloidal particle through the ultramicroscope, you will notice that it is continually jumping around in an irregular manner. These movements are known as
_________ motion. Scottish botanist Robert Brown discovered this effect in 1827 when observing pollen particles floating in water through a microscope. (Pollen particles are larger than colloids, but they are still small enough to exhibit some Brownian motion.) Brownian motion arises from collisions of the liquid molecules with the solid particle. For large particles, the millions of collisions from different directions cancel out, so they remain stationary. The smaller the particle, the smaller the number of surrounding molecules able to collide with it, and the more likely will random fluctuations occur in the number of collisions from different sides. Simple statistics predicts that every once in a while, the imbalance in collisions from different directions will become great enough to give the particle a real kick!

It is worth noting that Albert Einstein's analysis of Brownian motion in 1901 constituted the first proof of the molecular theory of matter.

Very large polymeric molecules such as proteins, starches and other biological polymers, as well as many natural polymers, exhibit colloidal behavior. There is no absolute distinction between a "very large molecule" and a "particle".

Colloidal particles can be made of solids, liquid droplets, or even of tiny gas bubbles dispersed in a condensed phase.

Each type of dispersion has a special name, and almost all of them are represented by familiar examples.

### Stability and flocculation of colloids

Many colloidal dispersions consist of tiny particles of solids or liquids dispersed in a second phase such as water. Macroscopically, however, colloidal dispersions appear to be single-phase systems. For example colloidal solids (sols), unlike ordinary precipitates, do not settle out to the bottom of the container.

What keeps the colloidal particles suspended in the medium? How can we force the particles to settle out?
These questions, which have countless applications in everyday life, are discussed in this unit.

If we introduce some particles of colloidal dimension into a liquid, thermal motions will cause them to occasionally collide with each other. As they do so, we might expect the particles to remain "stuck together" owing to the operation of a universal attractive force which we call the dispersion force.

In the absence of any opposing forces, we would expect colloidal particles to "stick together" when they collide. This would quickly result in the growth of aggregates sufficiently large to exceed colloidal size and to fall to the bottom of the container. This process is called coagulation.

... so why can colloids remain dispersed for indefinitely long times without coagulating? The answer is that there must be some stronger force between the particles that opposes dispersion force attractions.

In one class of colloids, called lyophilic ("solvent loving") colloids, the particles contain chemical groups that are strongly attracted by the solvent. The resulting solvent sheath protects the particles coming together, so lyophilic colloids are stable and show no tendency to settle out.

Ordinary gelatine is a common example of a lyophilic colloid. It is in fact hydrophilic, since it forms strong hydrogen bonds with water. When you mix Jello or tapioca powder to make a gelatine dessert, the material takes up water and forms a stable colloidal gel (a dispersion of water in the gelatine.)

Lyophilic (hydrophilic) colloids are very common in biological systems and in foods, but most of the colloids we deal with exhibit very little attraction to water: think of oil emulsions or rock dust in river water! These colloids are said to be lyophobic.

Lyophobic collids are all inherently unstable; they will eventually coagulate. However, "eventually" can be a very long time (the settling time for some clay colloids in the ocean is 200-600 years!). To understand how lyophobic colloids can remain suspended for so long, it is necessary to know something about the electrical properties of phase boundaries

The basic principle here is that virtually all interfaces between two phases possess an electrical imbalance of some kind.

Particles composed of ionic or ionizable substances usually have on their surfaces charged atoms that do not have oppositvely-charged neighbors to produce electrical neutrality.

Non-ionic particles or droplets such as oils or latex will tend to selectively adsorb positive or negative ions present in solution, thus "coating themselves" with electrical charge.

In order to preserve electrical neutrality, the colloidal particle with its attached surface charges will be surrounded by shell of oppositely-charged ions. These outer ions are not really "attached" to the colloid in the same sense that the inner ions are, but are simply attracted from the bulk solution.

This entire assembly is called an electric double layer.
Electric double layers of one kind or another exist at all phase boundaries, but those associated with collids are especially important.

Each particle with its double layer is electrically neutral. However, when two particles approach each other, each "sees" mainly the outer part of the double layer of the other. These will always have the same charge sign (this depends on the type of colloid and the nature of the medium), so there will be an electrostatic repulsive force that opposes the dispersion force attractions. This mutual repulsion discourages collisions and allows dispersions of lyophobic collids to persist for very long times.

"Do not freeze"

Have you ever seen what happens to milk when it freezes?
Not a pretty sight! You will see this label on many foodstuffs and on colloidal consumer products such as latex house paint.

Freezing causes the ions in the double layer to come out of solution, thus causing the double layer to collapse. The particles can now approach close enough for dispersion forces to take
over, and once they do so, they never let go: the colloid is permanently coagulated!

Rivers carry millions of tonnes of colloidal clay into the oceans. If you fly over the mouth of a river, you can often see the difference in color as the clay colloids coagulate due to the action of the salt water.
The coagulated clay accumulates as sediments which eventually form a geographical feature called a river _______ .

All hydrophobic colloids (and in fact all hydrophobic surfaces) are surrounded by a polar region known as the electric double layer. Even if the entire assembly is electrically neutral, the charge on the outer part of the double layer will repel other colloidal particles of the same kind, preventing them from moving close enough together to allow dispsersion forces to kick in.

If the double layer is destroyed (typically by freezing or by raising the ionic content of the solution), the repulsion disappears and the colloidal dispersion coagulates.

### Surfactants

Surface-active agents are molecules that tend to concentrate at the interface between two phases and confer very special properties on them. Your most common encounter with surfactants occurs when you use a soap or a detergent. If you don't know how these work, now is your chance to find out!

First, we need to review a couple of terms relating to the manner in which a molecule interacts with water. Substances such as those shown at the left tend to interact strongly with water owing to their polarity and the ability to form hydrogen bonds with water. Such substances are said to be hydro_______.

Hydrophilic (Greek for "water loving") refers to the strong attractive forces between polar molecules and water, resulting in a fall in the potential energy when these molecules are brought into contact with water.

[... contrast with dodecane...]

When Hydrophobic ("water fearing") substances dissolve in water, they disrupt the water's hydrogen-bonding structure, costing energy and making them only slightly soluble. Hydrophobic solutes tend to get "squeezed out" of water, as evidenced by their higher-than- expected vapor pressures when in solution (positive deviations from Raoult's law.)

We can always make a non-polar molecule into a polar one by attaching a polar group to it. Thus replacing one of the terminal hydrogen atoms on octane with a hydroxyl group yields the polar octyl alcohol, or octanol

A detergent is just another name for a surfactant substance.
"Detergent" implies that the molecule has been specially selected or designed to exploit its surface-active properties.

One of the most commonly encountered detergents is ordinary laundry or dishwashing detergent. These have various formulas, but one of the oldest (and still very common ones) is based on sodium dodecyl sulfate:

Most detergents will dissolve in water to at least a small extent, but the molecules will tend to accumulate at the surface. In order to minimize the energy, the polar end of the detergent molecule immersed in the water, with the nonpolar ends sticking out of the surface.

As we add more surfactant molecules to the water, they tend to line up at the air-water interface with their nonpolar ends projecting above the surface as depicted here. The result is a monomolecular layer of detergent molecules. Those molecules that are unable to find space in the surface film will swim around in the bulk solution just like ordinary solutes...
almost!

Placing these "free" detergent molecules in the solvent is not very favorable energetically, owing to the way the nonpolar ends disrupt the water's hydrogen-bonding structure. After the concentration of deterget rises to a certain critical value, something rather dramatic happens.

What happens is the "free" surfactants join together to form tiny (colloidal-size) aggregates called micelles. This is far more favorable energetically than the alternative of having the same number of surfactant molecules immersed in the water individually, with the water trying to squeeze them back out.

The neat thing about micelles is that they allow surfactant molecules to masquerade as totally-polar species whose interaction with water is energetically favorable because the hydrophobic parts are hidden away inside the micelle. This is true even though the micelles are too large to behave as true solute molecules, and tend to have the properties of very stable colloidal emulsions.

But the real importance of micelles goes way beyond this!
In the presence of a non-aqueous "oil-like" phase that would ordinarily tend to separate out into two liquid layers, surfactants can incorporate the non-aqueous phase into their micelles, thereby allowing the mixture to behave very much like a single-phase aqueous system.

A surfactant used in this way is said to solubilize a non-aqueous phase and is sometimes called an emulsifying agent.

The ability of sufactants to solubilize hydrophobic substances has many practical applications. For example, have you ever wondered how the oils and fats you eat can be brought into contact with the dissolved enzymes that bring about their digestion?

Solubilization of fats is brought about by substances called bile salts, which act as the body's "digestive detergent". These substances form micelles containing lecithin and cholesterol and are stored in the gall bladder until needed.

After a meal, the micelles are secreted into the upper intestine, where they solubilize the oils and lipids, permitting them to undergo digestion.

The cleaning action of soaps and detergents works the same way.
Most of the "dirt" that cannot be washed away by plain water consists of an "oil" phase which gets solubilized by the detergent so that it can be carried away by the wash water.

Surfactant molecules essentially give the water a "new" surface. Attractions between the non-polar ends of adjacent surfactant molecules are much smaller than between hydrogen-bonded water molecules, reducing the surface tnesion and thus the tendency of the water to form a spherical drop.

The other important function of a detergent is to reduce the surface tension of the water so that it can "wet" the fibres and easily penetrate into the fabric.

Certain types of surfactant molecules can also arrange themselves into hollow micelle-like structures known as lipid bilayer vesicles.
These form the basic structure of biological cell walls. Lipid bilayers can be broken up by strong detergent solutions, which is the reason that soaps and detergents "kill germs".

Surfactants and micelle formation constitute one of the more interesting aspects of applied chemistry and have countless applications in food science, physiology, and industrial technology.

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