AN INTRODUCTION TO SATURATED VAPOUR PRESSURE

Evaporation of the Liquid in an Open Container The average energy of the particles in a liquid is governed by the temperature The higher the temperature, the higher the average energy. But within that average, some particles have energies higher than the average, and others have energies lower than the average. Water molecules are simply breaking away from the surface layer. The energy which is lost as the particles evaporate is replaced from the surroundings. As the molecules in the water jostle with each other, new molecules will gain enough energy to escape from the surface.

Evaporation of the Liquid in an Closed Container first the molecules start to evaporate and hit the side of the container and cool down. After some time, the rate of evaporation and condensation becomes equal and the closed system reaches equilibrium. When these particles hit the walls of the container, they exert a pressure. This pressure is called the saturated vapour pressure (also known as saturation vapour pressure) of the liquid.

The variation of saturated vapour pressure with temperature The effect of temperature on the equilibrium between liquid and vapour. This can be seen in two ways: If you increase the temperature, you are increasing the average energy of the particles present. That means that more of them are likely to have enough energy to escape from the surface of the liquid. That will tend to increase the saturated vapour pressure.

OR:

The effect of temperature on the saturated vapour pressure of water Saturated vapour pressure and boiling point A liquid boils when its saturated vapour pressure becomes equal to the external pressure on the liquid. When that happens, it enables bubbles of vapour to form throughout the liquid - those are the bubbles you see when a liquid boils.

Saturated vapour pressure and solids

PHASE DIAGRAMS OF PURE SUBSTANCES

The phase diagram for water

The phase diagram for carbon dioxide

RAOULT'S LAW AND NON-VOLATILE SOLUTES

A simple explanation of why Raoult's Law works

Limitation of Raoults law Raoult's Law only works for ideal solutions. An ideal solution is defined as one which obeys Raoult's Law. Only very dilute solution obeys this laws. The forces of attraction between solvent and solute are exactly the same as between the original solvent molecules

Raoult's Law and melting and boiling points

Phase diagram water and a Non volatile substance

RAOULT'S LAW AND IDEAL MIXTURES OF LIQUIDS Raoult's Law and how it applies to mixtures of two volatile liquids Examples of ideal mixtures There is actually no such thing as an ideal mixture! However, some liquid mixtures get fairly close to being ideal. These are mixtures of two very closely similar substances. Commonly quoted examples include: hexane and heptane benzene and methylbenzene propan-1-ol and propan-2-ol

Ideal mixtures and intermolecular forces

Vapour pressure / composition diagrams Suppose you have an ideal mixture of two liquids A and B. Each of A and B is making its own contribution to the overall vapour pressure of the mixture - as we've seen above. Let's focus on one of these liquids - A, for example. Suppose you double the mole fraction of A in the mixture (keeping the temperature constant). According to Raoult's Law, you will double its partial vapour pressure. If you triple the mole fraction, its partial vapour pressure will triple - and so on. In other words, the partial vapour pressure of A at a particular temperature is proportional to its mole fraction. If you plot a graph of the partial vapour pressure of A against its mole fraction, you will get a straight line.

Now we'll do the same thing for B - except that we will plot it on the same set of axes. The mole fraction of B falls as A increases so the line will slope down rather than up. As the mole fraction of B falls, its vapour pressure will fall at the same rate.

Notice that the vapour pressure of pure B is higher than that of pure A. That means that molecules must break away more easily from the surface of B than of A. B is the more volatile liquid. To get the total vapour pressure of the mixture, you need to add the values for A and B together at each composition. The net effect of that is to give you a straight line as shown in the next diagram.

Boliling point / composition diagrams If a liquid has a high vapour pressure at a particular temperature, it means that its molecules are escaping easily from the surface. If, at the same temperature, a second liquid has a low vapour pressure, it means that its molecules aren't escaping so easily. What does that imply about the boiling points of the two liquids? There are two ways of looking at this.

Either: If the molecules are escaping easily from the surface, it must mean that the intermolecular forces are relatively weak. That means that you won't have to supply so much heat to break them completely and boil the liquid. The liquid with the higher vapour pressure at a particular temperature is the one with the lower boiling point. Or: Liquids boil when their vapour pressure becomes equal to the external pressure. If a liquid has a high vapour pressure at some temperature, you won't have to increase the temperature very much until the vapour pressure reaches the external pressure. On the other hand if the vapour pressure is low, you will have to heat it up a lot more to reach the external pressure.

Constructing a boiling point / composition diagram

To make this diagram really useful (and finally get to the phase diagram we've been heading towards), we are going to add another line. This second line will show the composition of the vapour over the top of any particular boiling liquid. If you boil a liquid mixture, you would expect to find that the more volatile substance escapes to form a vapour more easily than the less volatile one. That means that in the case we've been talking about, you would expect to find a higher proportion of B (the more volatile component) in the vapour than in the liquid. You can discover this composition by condensing the vapour and analysing it. That would give you a point on the diagram.

Using the phase diagram if you boil a liquid mixture C1, it will boil at a temperature T1 and the vapour over the top of the boiling liquid will have the composition C2. All you have to do is to use the liquid composition curve to find the boiling point of the liquid, and then look at what the vapour composition would be at that temperature. Notice again that the vapour is much richer in the more volatile component B than the original liquid mixture was.

Suppose that you collected and condensed the vapour over the top of the boiling liquid and reboiled it. You would now be boiling a new liquid which had a composition C2. That would boil at a new temperature T2, and the vapour over the top of it would have a composition C3. You can see that we now have a vapour which is getting quite close to being pure B. If you keep on doing this (condensing the vapour, and then reboiling the liquid produced) you will eventually get pure B.

FRACTIONAL DISTILLATION OF IDEAL MIXTURES OF LIQUIDS

Fractional distillation in the lab -The apparatus

Some notes on the apparatus The fractionating column is packed with glass beads (or something similar) to give the maximum possible surface area for vapour to condense on. You will see why this is important in a minute. Some fractionating columns have spikes of glass sticking out from the sides which serve the same purpose. If you sketch this, make sure that you don't completely seal the apparatus. There has to be a vent in the system otherwise the pressure build-up when you heat it will blow the apparatus apart. In some cases, where you are collecting a liquid with a very low boiling point, you may need to surround the collecting flask with a beaker of cold water or ice. The mixture is heated at such a rate that the thermometer is at the temperature of the boiling point of the more volatile component. Notice that the thermometer bulb is placed exactly at the outlet from the fractionating column.

Relating what happens in the fractionating column to the phase diagram

Fractional distillation industrially

The column contains a number of trays that the liquid collects on as the vapour condenses. The up-coming hot vapour is forced through the liquid on the trays by passing through a number of bubble caps. This produces the maximum possible contact between the vapour and liquid. This all makes the boiling-condensing-reboiling process as efficient as possible. The overflow pipes are simply a controlled way of letting liquid trickle down the column. If you have a mixture of lots of liquids to separate (such as in petroleum fractionation), it is possible to tap off the liquids from some of the trays rather than just collecting what comes out of the top of the column. That leads to simpler mixtures such as gasoline, kerosene and so on.

NON-IDEAL MIXTURES OF LIQUIDS

Positive deviations from Raoult's Law In mixtures showing a positive deviation from Raoult's Law, the vapour pressure of the mixture is always higher than you would expect from an ideal mixture. The deviation can be small behave just like ideal mixtures as far as distillation is concerned But some liquid mixtures have very large positive deviations from Raoult's Law, and in these cases, the curve becomes very distorted.

Notice that mixtures over a range of compositions have higher vapour pressures than either pure liquid. The maximum vapour pressure is no longer that of one of the pure liquids.

Explaining the deviations The fact that the vapour pressure is higher than ideal in these mixtures means that molecules are breaking away more easily than they do in the pure liquids. That is because the intermolecular forces between molecules of A and B are less than they are in the pure liquids. You can see this when you mix the liquids. Less heat is evolved when the new attractions are set up than was absorbed to break the original ones. Heat will therefore be absorbed when the liquids mix. The enthalpy change of mixing is endothermic. The classic example of a mixture of this kind is ethanol and water. This produces a highly distorted curve with a maximum vapour pressure for a mixture containing 95.6% of ethanol by mass.

Negative deviations from Raoult's Law the deviation are much greater giving a minimum value for vapour pressure lower than that of either pure component

Explaining the deviations These are cases where the molecules break away from the mixture less easily than they do from the pure liquids. New stronger forces must exist in the mixture than in the original liquids. You can recognise this happening because heat is evolved when you mix the liquids - more heat is given out when the new stronger bonds are made than was used in breaking the original weaker ones. Many (although not all) examples of this involve actual reaction between the two liquids. The example of a major negative deviation that we are going to look at is a mixture of nitric acid and water. These two covalent molecules react to give hydroxonium ions and nitrate ions. You now have strong ionic attractions involved.

A large positive deviation from Raoult's Law: ethanol and water mixtures If a mixture has a high vapour pressure it means that it will have a low boiling point. The molecules are escaping easily and you won't have to heat the mixture much to overcome the intermolecular attractions completely. The implication of this is that the boiling point / composition curve will have a minimum value lower than the boiling points of either A or B. In the case of mixtures of ethanol and water, this minimum occurs with 95.6% by mass of ethanol in the mixture. The boiling point of this mixture is 78.2°C, compared with the boiling point of pure ethanol at 78.5°C, and water at 100°C. You might think that this 0.3°C doesn't matter much, but it has huge implications for the separation of ethanol / water mixtures.

the boiling point / composition curve for ethanol / water mixtures

Distillation of an ethanol and water Mixture

Notes What happens if you reboil that liquid? The liquid curve and the vapour curve meet at that point. The vapour produced will have that same composition of 95.6% ethanol. If you condense it again, it will still have that same composition. You have hit a barrier. It is impossible to get pure ethanol by distiling any mixture of ethanol and water containing less than 95.6% of ethanoll. This particular mixture of ethanol and water boils as if it were a pure liquid. It has a constant boiling point, and the vapour composition is exactly the same as the liquid. It is known as a constant boiling mixture or an azeotropic mixture or an azeotrope.

To summarise Distilling a mixture of ethanol containing less than 95.6% of ethanol by mass lets you collect: a distillate containing 95.6% of ethanol in the collecting flask (provided you are careful with the temperature control, and the fractionating column is long enough); pure water in the boiling flask. What if you distil a mixture containing more than 95.6% ethanol? Work it out for yourself using the phase diagram, and starting with a composition to the right of the azeotropic mixture. You should find that you get: pure ethanol in the boiling flask.

A large negative deviation from Raoult's Law: nitric acid and water mixtures Nitric acid and water form mixtures in which particles break away to form the vapour with much more difficulty than in either of the pure liquids. You can see this from the vapour pressure / composition curve. Mixtures of nitric acid and water can have boiling points higher than either of the pure liquids because it needs extra heat to break the stronger attractions in the mixture. In the case of mixtures of nitric acid and water, there is a maximum boiling point of 120.5°C when the mixture contains 68% by mass of nitric acid. That compares with the boiling point of pure nitric acid at 86°C, and water at 100°C. Notice the much bigger difference this time due to the presence of the new ionic interactions.

vapour pressure / composition curve of Nitric Acid and water

Distilling dilute nitric acid The vapour produced is richer in water than the original acid. If you condense the vapour and reboil it, the new vapour is even richer in water. Fractional distillation of dilute nitric acid will enable you to collect pure water from the top of the fractionating column. As the acid loses water, it becomes more concentrated. Its concentration gradually increases until it gets to 68% by mass of nitric acid. At that point, the vapour produced has exactly the same concentration as the liquid, because the two curves meet. You produce a constant boiling mixture (or azeotropic mixture or azeotrope). If you distil dilute nitric acid, that's what you will eventually be left with in the distillation flask. You can't produce pure nitric acid from the dilute acid by distilling it.

Distilling nitric acid more concentrated than 68% by mass The vapour formed is richer in nitric acid. If you condense and reboil this, you will get a still richer vapour. If you continue to do this all the way up the fractionating column, you can get pure nitric acid out of the top. As far as the liquid in the distillation flask is concerned, it is gradually losing nitric acid. Its concentration drifts down towards the azeotropic composition. Once it reaches that, there can't be any further change, because it then boils to give a vapour with the same composition as the liquid. Distilling a nitric acid / water mixture containing more than 68% by mass of nitric acid gives you pure nitric acid from the top of the fractionating column and the azeotropic mixture left in the distillation flask.