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how a gas can interact with a liquid.

Several household examples illustrate the interaction between a gas and a liquid. For example, put some cold tap water in a saucepan and heat it gently on the stove. We have all done this and seen the result. As the water warms, thousands of tiny bubbles appear and cling to the side of the pan, as illustrated in Figure 2.1. This happens long before the water boils. The bubbles are obviously gas − but what gas is it, and where did it come from?

      Gases dissolve less at higher temperature. Heating the water expels the dissolved gas, visible as a cluster of small bubbles.

Figure 2.1 Gases Dissolve in Liquids.

      A moment's thought should reveal the answer. The gas can't be steam because steam bubbles can't exist under water unless the water is boiling. The gas can't be hydrogen because water doesn't decompose at this low temperature. No, the gas must be coming from something dissolved in the water, most likely oxygen and nitrogen from the air.

This first example demonstrates two important principles:

       Gases dissolve in liquids.

       Gases dissolve less in hot liquids than they do in cold liquids.

Another common example is the bottle of champagne illustrated in Figure 2.2. Closely inspecting the unopened bottle, we see very little gas; the bottle is nearly full of liquid. Yet, upon popping the cork, an enormous quantity of gas suddenly appears. This gas is carbon dioxide, a different gas than in our first example. The liquid is not quite the same either, but it's mainly water. And we can be sure that pure soda water would behave in exactly the same way.

      Gases dissolve less at lower pressure. Popping the cork lets the pressure drop, releasing lots of bubbles.

Figure 2.2 A Different Gas.

      Again, where did the gas come from? There is only one place that it could have been hiding; it was dissolved in the liquid. This confirms our theory that gases dissolve in liquids, but far more was dissolved this time; the solubility of carbon dioxide in water is greater than the solubility of air in water.

      This second example demonstrates two additional principles:

       Gases dissolve less at low pressure than they do at high pressure.

       Some gases dissolve in a given liquid more than other gases do.

      A final example (not illustrated) is what happens when you heat cooking oil in a pan. No bubbles appear. Apparently, the oil doesn't dissolve any air − and it won't dissolve much carbon dioxide either!

      This last example gives a fifth principle:

       Some liquids dissolve more of a gas than other liquids do.

      These five principles are all the science you will need to truly understand what happens in a column. They remind you that a gas can dissolve in a liquid and that the amount of gas that can dissolve depends on only four simple variables; the temperature and pressure, the type of gas, and the type of liquid.

      That's it. Nothing else affects the solubility of a gas in a liquid.

      The four variables are very easy to understand, yet they are the hidden foundation of all gas chromatography. Let's see how that can be …

      In most process gas chromatographs, three of the four variables are closely controlled and do not vary:

       The column temperature is held constant.

       The carrier gas pressure is held constant.

       The liquid phase is predetermined and doesn't change.

      The fourth and most important variable is due to the different gases in the sample. And this is the real cause of chromatographic separation:

       Different gases have different solubility in the liquid phase.

Gas chromatography works because each component to be separated has a different solubility in the liquid phase. We shall see that the less soluble peaks move quickly through the column while the more soluble peaks take longer to get through.

      This is the process of separation. It's all about solubility.

      Chemists call the liquid phase a solvent and each dissolved component a solute. But no real chemistry is involved. If a chemical reaction occurred, it might destroy some of the molecules that we are trying to measure and likely would cause a gradual and irreversible deterioration of the column itself.

      Before moving on, a quick reminder. The discussion in this chapter focuses on the most common kind of column; one that has a liquid stationary phase. As noted earlier, another kind of column uses a solid stationary phase. The solute molecules can't dissolve in a solid, but they can and do adhere to its solid surface, and the final outcome is much the same.

      Troubleshooting tips

      The household examples used above may provide some valuable help with troubleshooting and are worth remembering:

       When a column works at higher temperature, gas solubility is reduced, and all the peaks come out earlier on the chromatogram, thereby reducing their separation. For an easy way to remember this, recall the heated water!

       When a column works at higher pressure, gas solubility is increased and all the peaks come out later on the chromatogram, thereby increasing their separation. For an easy way to remember this, recall the bubbly champagne!

      These troubleshooting tips assume that the carrier gas flow rate is held constant. Later chapters discuss the effect of other variables.

      The rest of this chapter explains how solubility causes the classic peak shape. The following chapter examines how a difference in solubility will cause the peaks to become separated from each other.

How peaks form

      Forming an equilibrium

To examine the interaction between a gas and a liquid, consider a small enclosed space that's internally divided into a gas space and a liquid space, as in Figure 2.3a. For explanatory purposes, the diagrams in Figure 2.3 show the gas space and liquid space as deep layers that would not work in practice. In a real column, the gas and liquid layers are very shallow, so the sample molecules can move quickly between them.

Figure 2.3 Forming an Equilibrium.

Let's

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