Vapour Pressure

INTRODUCTION

THEORY

Atkins and de Paula (2001 p. 192) indicates that a phase is a state of matter that is uniformed in chemical composition and physical state. They further go on to say that phases are difficult to identify for substances and thus have to be carefully analyzed before a conclusion can be made.

Atkins and de Paula (2001 p.192) indicates that a constituent is a chemical species that is present. It can be an “ion or a molecule”. They further went on to say that a component is a “chemical independent constituent”.

(A) THE ORIGIN OF SATURATED VAPOUR PRESSURE

(i) THE EVAPORATION OF A LIQUID

The “average energy” of particles found in a liquid are attributed to the temperature of system. He explains that higher temperatures are responsible for increases in the “average energy”. However, he also states that there are particles that will have both higher and lower energies with respect to the average. These particles that have enough energy will evaporate and therefore leave the surface of the liquid. (Clark 2004)

“Evaporation” occurs on a liquid's surface due to the particle's energy levels as opposed to “boiling” which occurs because of a surmount of energy that “disrupt[s] the attractive forces” within a liquid. This can be attributed to the “bubbles of gas being formed through the liquid”. A contrastive analysis to water evaporating was made and it stated that the energy lost as evaporation is “replaced from the surrounding”. It was stated that the molecules interact with each other by “jostling” and in lieu of this; other molecules will be able to gain enough energy to escape. (Clark 2004)

(ii) THE EVAPORATION OF A LIQUID IN A CLOSED CONTAINER

Clark (2004) considered an analysis of a liquid in a closed container, but more specifically considers water. He states that “common sense” allow you to notice that it evaporation does not occur owing to the container being closed. However, was explicit in saying that there is a “constant evaporation” at the surface of the liquid. He indicates also that “particles continue to break away” but are “trapped in the space above the liquid” which also is owed to the container being closed. This is illustrated below:

Clark (2004) states that as the “gaseous particles bounce around” one would expect that a few comes into contact with the liquid's surface and is then “trapped there”. Consequently, a rapid equilibrium ensues which causes the “particles leaving the surface” to be “exactly balanced by the number rejoining it”.

Clark (2004) states that a “fixed number” of particles will be present in the “space above the liquid” and further explains that the saturated vapour pressure occurs when the particles “hits the walls of the container”.

(B) THE VARIATION OF SATURATED VAPOUR PRESSURE WITH TEMPERATURE

(i) THE EFFECT OF TEMPERATURE ON THE EQUILIBRIUM BETWEEN LIQUID AND VAPOUR

Clark (2004) indicates that this can be explained in two ways. One is based on a common sense approach and the other based on Le Chatelier's Principle.

The common sense approach is based on the concept that the average energy of the particles present will increase as the temperature increases (this was mentioned before). It consequently leads to the conjecture that more of these particles will be able to escape from the liquid's surface and therefore directly cause the saturated vapour pressure to increase. Clark (2004)

With regards to Le Chatelier's Principle, one needs to consider the equilibrium between a liquid and a vapour as follows (Clark 2004):

Clark (2004) attributes to the “forward change” as endothermic due to the fact that heat is needed in the process transforming “liquid into vapour”. He relates Le Chatelier's principle to say that “increasing the temperature” when a system is in “dynamic equilibrium” will “favour [sic] the endothermic change”. It leads to the inference that “increasing the temperature increases the amount of vapour present” and consequently “increases the saturated vapour pressure”.

(ii) SATURATED VAPOUR PRESSURE AND BOILING POINT

Clark (2004) explains that a liquid will boil when its “vapour pressure is equal to the external pressure on the liquid” and when this happens “it enables bubbles of vapour” to be seen within and “throughout the liquid”.

Clark also explains what occurs when the external pressure is higher than the vapour pressure and states that the bubbles mentioned earlier are not allowed to form and evaporation takes place at the liquid's surface.

(C) IDEAL MIXTURES OF LIQUIDS

(i) RAOULT'S LAW AND IDEAL MIXTURES OF LIQUIDS

Clark (2004) elucidates the concept of Raoult's law and provides an analysis of how it can be related to the mixture of two volatile miscible liquids “in all proportions”, to consequently give a “single liquid”.

Monk (2004 pg. 228) defines an ideal mixture as one that obeys Raoult's law.

Raoult's law is defined as “the partial vapour pressure of a component in a mixture is equal to the vapour pressure of the pure component at that temperature multiplied by its mole fraction in the mixture”. It should be noted that the law only works for ideal mixtures. Clark (2004)

Atkins and de Paula (2001 p.195) relates the partial vapour pressures of the components of an ideal solution. They indicate the following:

Where:

ü and are the vapour pressures of pure A and B.

ü and are the mole fractions of A and B.

ü and are the partial pressures of A and B.

In a mixture of gases, each of them will exert an individual and independent pressure and is called its partial pressure. The total vapour pressure of a mixture is the sum of the partial pressures of all the gases present. If two is being considered then: (Clark 2004)

Where:

ü is the total vapour pressure.

ü is the partial pressure of a gas A.

ü is the partial pressure of gas B.

The compositions of the liquid and vapour are not necessarily in equilibrium. It is suggested that the more volatile component has the richer vapour (Atkins, 2001 p.196). It follows from Dalton's law that:

Atkins (2001 p. 196) explains that when the mixture is ideal, the “partial pressures and the total pressure” can be expressed in terms of the mole fractions as follows:

a) IDEAL MIXTURES AND INTERMOLECULAR FORCES

Clark (2004) explains that in a pure liquid: the molecules with more energy are able to “overcome intermolecular attractions” and leave the surface forming a “vapour” in the process. He indicates that when the intermolecular forces are small a lot more molecules are able to escape at a fixed temperature.

Clark also elucidates that if a second liquid is present then above can be applied as well and goes on to say that at “any particular temperature” molecules in a “certain proportion” will have “enough energy” to be able to exit the liquid's surface.

When a mixture is ideal (the liquids obey Raoult's law), the “tendency” of the molecules of these two liquids (considered as one) to “escape” is unchanged.

Clark's analysis of the diagram suggests one might think “only half as many of each molecule [is] escaping” but the “proportion” of the molecules that are escaping will be the same. Clark further states that the diagram shows a 50/50 mixture of two liquids and means that only a partial (half) amount of the molecules of each liquid are on the surface as “pure liquids”. Now when the proportion of the molecules is the same, only “half as many” molecules are able to escape in a “given time”. If one set of molecules have the “same tendency to escape” as opposed to the other, then “the intermolecular forces” between two molecules will be the same for both liquids. On the other hand, if the “forces were different”, the “tendency to escape” would vary somewhat significantly. If both had different intermolecular interactions or forces, then it would be expected that the escape rates would be very different indeed. This intermolecular force is often referred to van der Waals interactions.

b) IDEAL MIXTURES AND ENTHALPY CHANGE OF MIXING

Mixing two liquids involve breaking “existing intermolecular attractions” and this process requires energy (endothermic). It then requires making new intermolecular interactions and this process releases energy (exothermic). However, if the interactions are the same then no energy will be required or released and means that there will be 0 enthalpy of mixing. It then leads to an inference that: if there is a temperature change then the mixture is not ideal. (Clark 2004)

(ii) VAPOUR PRESSURE / COMPOSITION DIAGRAMS

Atkins and de Paula (2001 pg. 195) to relate the partial pressures of A and B, giving the total vapour pressure, p as:

Since:

Therefore:

Therefore from, it can be said that “the total vapour pressure (at some fixed temperature) changes linearly with composition from to as changes from 0 to 1” (Atkins 2001 p. 195).

On another note, if you have an ideal mixture of two liquids A and B, then both A and B will contribute to the total vapour pressure of the mixture. According to Raoult's law, the partial vapour pressures of A and B at a particular temperature are proportional to the mole fraction. Therefore plotting partial vapour pressure of A or B against its mole fraction, separately, will allow straight lines to be ascertained as indicated below: (Clark 2004)

The vapour pressure of pure B is higher than that of pure A. It means that the molecules must break away more easily from the surface of B than of A and therefore leads to the inference that B is the more volatile liquid as opposed to A. (Clark 2004)

This all leads to being able to determine the total vapour pressure of the mixture. This is done by adding the vapour pressures of A and B at each composition (mole fraction) and the resultant effect is a straight line as shown below: (Clark 2004)

(iii) TEMPERATURE COMPOSITION DIAGRAMS

Negi and Anand (1985 pg. 367) indicates that based on Raoult's law, different concentrations of solutions have different vapour pressures; hence the solutions of different concentrations will boil at different temperatures. Solutions whose components have low vapour pressures will boil at higher temperatures than solutions in which the components have high vapour pressures. This is due to the fact that the external pressure is attainted at a lower temperature in case of volatile components. Thus the solution with a high vapour pressure starts boiling at a lower temperature.

(iv) BOLILING POINT / COMPOSITION DIAGRAMS

Negi and Anand (1985 pg. 367) gives the following analysis of the figure that shows a boiling point composition curve for an ideal solution. They stated that the “liquid curve” is below the “vapour curve” and this can be garnered from the diagram as well. They further went on to say that it is due to the fact that at “constant pressure” above the “boiling point” only “vapours will exist”. In the figure A'LB' is the boiling point liquid composition curve and A'VB' represents the boiling point - vapour composition curve. Suppose a solution X is heated, on heating the vapour pressure of the system will increase until it reaches a constant value equal to the external pressure. At this stage, the liquid begins to boil. The temperature therefore corresponds to the boiling point of the solution of composition X. At, the vapours in equilibrium with the liquid will have the composition Y which is richer in the more volatile component B than X. If the vapour corresponding to Y is condensed, a liquid will result which will boil at a temperature. The composition of the vapours in equilibrium with this condensate is given by . It is again richer in B than the original vapour. If this process of condensing and redistilling is repeated a number of times, ultimately vapour containing pure B can be obtained.

The initial residue left after the removal of the vapours will be richer in the less volatile component A, say X'. The residue corresponding to X' will boil at a temperature Tx' (Tx' > Tx). The corresponding composition of the vapours in equilibrium with solution of composition X' is given by Y'. These vapours are again richer in B and the residual liquid richer in A. In this manner, if the vapour at one boiling point is condensed and redistilled at a higher temperature and the process is repeated again and again, the residue becomes richer and richer in A than the original solution until pure A is obtained.

a) THE RELATIONSHIP BETWEEN BOILING POINT AND VAPOUR PRESSURE

Clark's (2004) reasoning has lead to the development of the following constructs:

* A liquid with a high vapour pressure at a temperature T implies that the molecules are escaping quite easily from the surface of the said liquid.

* Another liquid with a low vapour pressure at the same temperature T implies that the molecules of this liquid are NOT escaping at the same rate as opposed to the other.

Two explanations, however simple, give an indication of how vapour pressures and boiling points are related:

* The molecules of the liquid that will escape easily from the surface imply that the intermolecular forces are “relatively weak” and therefore do not require “much heat” to break the bonds (or interactions) to boil the liquid. Therefore, the liquid with the high vapour pressure will be the one with the low boiling point as compared to the other liquid with the lower vapour pressure. (Clark 2004)

* Another analogy is based on that liquids will boil when the vapour pressures becomes equal to the external pressure, as was mentioned earlier. A liquid with a high vapour pressure means that a high temperature will not be required for this liquid to become equal with atmospheric pressure and thus boil quickly. On the other hand, if the vapour pressure of the other liquid is low, then a higher temperature will be required to raise the vapour pressure to the external pressure so that the liquid boils. (Clark 2004)

b) CONSTRUCTING A BOILING POINT / COMPOSITION DIAGRAM

Starting with the boiling points of pure A and B. B has the higher vapour pressure. For mixtures of A and B, the boiling points form a curve.

Adding another line will show the composition of the vapour over the top of any particular boiling liquid. If you a boil a liquid mixture (binary), the more volatile liquid escapes quite easily as compared to the other. Therefore, one can expect to find a “higher proportion of B”, which is more volatile in the vapour as opposed to it being in the liquid. This composition can be identified by “condensing the vapour and analyzing it” giving the following diagram: (Clark 2004)

When you boil a particular mixture of A and B notice that the vapour over the top of the boiling liquid has a composition which is much richer in B, the more volatile component. A vapour composition curve can be ascertained when different compositions are used of the liquid mixtures. The final diagram is as follows: (Clark 2004)

c) USING THE PHASE DIAGRAM

If you boil a liquid mixture, you can find out the temperature it boils at, and the composition of the vapour over the boiling liquid. In the proceeding diagram, boiling a liquid mixture C1, one can be able to find the temperature it boils at, and that corresponds to the temperature T1. In addition, the vapour pressure of the boiling liquid will have a composition denoted by C2. (Clark 2004)

(D) NON - IDEAL MIXTURES OF LIQUIDS

(I) VAPOUR PRESSURE / COMPOSITION DIAGRAMS FOR NON-IDEAL MIXTURES

If you plot the vapour pressure of an ideal mixture of two liquids against their composition the following graph is ascertained (as mentioned before): (Clark 2004)

In this case, pure A has the higher vapour pressure and so is the more volatile liquid. Raoult's Law only works for ideal mixtures. In ideal mixtures, the interactions between the particles are same and the tendency for the molecules to escape is the same in pure liquids. However, for ideal mixtures, this is far from true. (Clark 2004)

(iii) POSITIVE DEVIATIONS FROM RAOULT'S LAW

Mixtures that show a positive deviation from Raoult's law has vapour pressures that are higher than expected with respect to an ideal mixture. These are usually small deviations and is illustrated below: (Clark 2004)

There are cases when the deviations are greater as in the following diagram:

The mixtures within a range of compositions have very high vapour pressures as compared to the pure liquids and the maximum vapour pressure is no longer one of the pure liquids. (Clark 2004)

(a) EXPLAINING THE DEVIATIONS

The high vapour pressure can be explained in terms of molecules breaking away more easily as in pure solutions and thus are able to go into vapour more quickly. It also leads to the inference that the intermolecular forces between the two molecules (A and B) are less than in pure liquids. Consequently, less heat is absorbed when the “new attractions” are set up as compared to the energy absorbed to break the prior attractions and heat will be absorbed more contributing to the enthalpy of mixing being “endothermic”. (Clark 2004)

(iv) NEGATIVE DEVIATIONS FROM RAOULT'S LAW

There are mixtures with vapour pressures less than expected as compared to ideal vapour pressures and this deviates negatively from Raoult's law. Like, positive deviations, some cases have small deviations whilst others have more significant and noticeable deviations. (Clark 2004)

(a) EXPLAINING THE DEVIATIONS

When molecules go into the vapour state less easily compared to pure liquids then there exists negative deviations from ideality and the vapour pressure is lower. It should be expected that there are stronger forces in the mixture as compared to pure liquids. In these types of liquids one finds heat being evolved during mixing. When the new stronger bonds are made, more heat is given out as compared to breaking the weaker bonds. Typically one finds that strong ionic interactions are involved attributing the increase heat evolved and the enthalpy of mixing is exothermic. (Clark 2004)

(v) BOILING POINT / COMPOSITION DIAGRAMS FOR NON - IDEAL MIXTURES

POSITIVE DEVIATION

As mentioned earlier a large positive deviation will mean the production of a vapour pressure curve with a maximum value at some composition so a high vapour pressure is noted. Based on previous discussions, it leads to a lower boiling point at that specific composition. An advantage is that the molecules are escaping very easily and will not require excessive and incessant heating to overcome “intermolecular attractions”. (Clark 2004)

However the implication is that the boiling point - composition curve will have a low boiling point as compared to both A and B and even more so, the lowest of all. (Clark 2004)

NEGATIVE DEVIATION

There are mixtures in which particles escape the surface to form vapours with much more difficulty than in either of the pure liquids. It means that these mixtures tend have their boiling points “higher than the pre liquids” because they need more kinetic energy to break the strong attractions within the mixture. (Clark 2004)

The phase diagram looks like this for a nitric acid / water mixture:

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.

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).

THE BOILING POINTS OF THE LIQUIDS

The following table shows the boiling points of pure liquids:

Liquid

Boiling Point

1-Propanol

97 oC

2-Propanol

82 oC

Chloroform

61 oC

Acetone

56 oC

Boiling points of the liquids investigated. (Mortimer 2008)

ANTOINE'S EQUATION

The Antoine equation relates the vapour pressure of pure components with certain parameters, A, B and C. This relation is given as follows: (Assael et al 1996)

Where:

* are component specific constants based on the identity of the substance.

* is the temperature in oC.

* is the logarithm of the vapour pressure.

The parameters A, B and C are reported for the following substances as: (Dean 1999)

SUBSTANCE

A

B

C

1 - Propanol

7.84767

1499.210

204.640

2 - Propanol

8.11778

1580.920

219.610

Chloroform

6.49300

929.4400

196.030

Acetone

7.11714

1210.595

229.664

The above expression was used to find the pure vapour pressures of both liquids by assuming ideality (i.e. the mixture obeys Raoult's law without any deviations - positive or negative). When boiling occurs, the mixture's vapour pressure becomes equal to the external pressure (in this case atmosphere which is taken at 1.00 atm). Therefore:

By using Antoine's equation, the vapour pressure of the pure liquids were calculated, so it allowed to be found and since the mole fractions are known, the vapour pressure at each temperature can calculated quite easily.

INTERMOLECULAR FORCES

All molecules have the capability to form London forces (Wojciechowski and Cerpovicz 1998). These are solely dependent on the surface area and the polarizability of the surface of the molecule. They result from moving electrons that generate positive and negative charged regions in the molecule.

TREATMENT OF RESULTS

Mixture

Boiling Point

Vapour Pressure (atm)

Vol. of 1-CH3(CH2)2OH

Vol. of 2-CH3(CH2)2OH

Vol. % 1-CH3(CH2)2OH

Mole Fraction 1-CH3(CH2)2OH

Mole Fraction

2-CH3(CH2)2OH

° C

K

Total

(atm)

1-CH3(CH2)2OH

2-CH3(CH2)2OH

10.0 + 0.1

00.0 + 0.1

100.00

1.00

0.00

95.0 + 0.1

368.2 + 0.1

1.000

1.000

0.000

10.0 + 0.1

02.0 + 0.1

083.33

0.83

0.17

94.0 + 0.1

367.2 + 0.1

1.000

0.738

0.262

10.0 + 0.1

04.0 + 0.1

071.43

0.71

0.29

92.5 + 0.1

365.7 + 0.1

1.000

0.576

0.424

10.0 + 0.1

06.0 + 0.1

062.50

0.63

0.38

91.0 + 0.1

364.2 + 0.1

1.000

0.474

0.526

10.0 + 0.1

08.0 + 0.1

055.56

0.56

0.44

90.0 + 0.1

363.2 + 0.1

1.000

0.399

0.601

10.0 + 0.1

10.0 + 0.1

050.00

0.50

0.50

89.5 + 0.1

362.7 + 0.1

1.000

0.337

0.663

08.0 + 0.1

10.0 + 0.1

044.44

0.44

0.56

89.0 + 0.1

362.2 + 0.1

1.000

0.277

0.723

06.0 + 0.1

10.0 + 0.1

037.50

0.38

0.63

88.0 + 0.1

361.2 + 0.1

1.000

0.218

0.782

04.0 + 0.1

10.0 + 0.1

028.57

0.29

0.71

87.0 + 0.1

360.2 + 0.1

1.000

0.140

0.860

02.0 + 0.1

10.0 + 0.1

016.67

0.17

0.83

85.0 + 0.1

358.2 + 0.1

1.000

0.071

0.929

00.0 + 0.1

10.0 + 0.1

000.00

0.00

1.00

83.0 + 0.1

356.2 + 0.1

1.000

0.000

1.000

Table 1 - The manipulated results to show mole fractions, vapour pressures, etc for each liquid.

Mixture

Boiling Point

Vapour Pressure (atm)

Vol. of CH3COCH3

Vol. of CHCl3

Vol. % CH3COCH3

Mole Fraction CH3COCH3

Mole Fraction

CHCl3

° C

K

Total

(atm)

CH3COCH3

CHCl3

10.0 + 0.1

00.0 + 0.1

100.00

1.00

0.00

58.5 + 0.1

331.7 + 0.1

1.000

1.000

0.000

10.0 + 0.1

02.0 + 0.1

083.33

0.83

0.17

61.5 + 0.1

334.7 + 0.1

1.000

0.832

0.168

10.0 + 0.1

04.0 + 0.1

071.43

0.71

0.29

64.5 + 0.1

337.7 + 0.1

1.000

0.683

0.317

10.0 + 0.1

06.0 + 0.1

062.50

0.63

0.38

66.5 + 0.1

339.7 + 0.1

1.000

0.558

0.442

10.0 + 0.1

08.0 + 0.1

055.56

0.56

0.44

67.5 + 0.1

340.7 + 0.1

1.000

0.459

0.541

10.0 + 0.1

10.0 + 0.1

050.00

0.50

0.50

68.5 + 0.1

341.7 + 0.1

1.000

0.373

0.627

08.0 + 0.1

10.0 + 0.1

044.44

0.44

0.56

64.5 + 0.1

337.7 + 0.1

1.000

0.384

0.616

06.0 + 0.1

10.0 + 0.1

037.50

0.38

0.63

66.5 + 0.1

339.7 + 0.1

1.000

0.263

0.737

04.0 + 0.1

10.0 + 0.1

028.57

0.29

0.71

67.5 + 0.1

340.7 + 0.1

1.000

0.131

0.869

02.0 + 0.1

10.0 + 0.1

016.67

0.17

0.83

67.5 + 0.1

340.7 + 0.1

1.000

0.012

0.988

00.0 + 0.1

10.0 + 0.1

000.00

0.00

1.00

64.5 + 0.1

337.7 + 0.1

1.000

0.000

1.000

Table 2 - The manipulated results to show mole fractions, vapour pressures, etc for each liquid.

Liquid

Boiling Point (Lab)

Boiling Point (Literature)

Deviation

1-Propanol

95.0 oC

97.0 oC

2.0 oC

2-Propanol

83.0 oC

82.0 oC

-1.0 oC

Chloroform

64.5 oC

61.0 oC

-3.5 oC

Acetone

58.5 oC

56.0 oC

-2.5 oC

Table 3 - Table showing the boiling points of the liquids used, with their standard boiling points and deviations from normal.

TREATMENT OF RESULTS

TREATMENT OF RESULTS

TREATMENT OF RESULTS

TREATMENT OF RESULTS

DISCUSSION

The experiment investigates the variation of boiling point and vapour pressure with composition of a binary mixture of liquids. It is based on the extent to which intermolecular forces are affected by mixing of two miscible liquids.

* In ideal mixtures, intermolecular forces are identical.

* In non - ideal mixtures, intermolecular forces are not identical. They can be stronger or weaker compared to ideal mixtures.

(A) 1-PROPANOL / 2-PROPANOL

· VAPOUR PRESSURE COMPOSITION

1-Propanol and 2-Propanol according to Clark (2004) when mixed is considered an ideal mixture and thus can be treated as such. The intermolecular forces (van der Waals forces, dipole - dipole and hydrogen bonding) in such a mixture are considered to be equal. When a mixture is ideal (the liquids obey Raoult's law), the tendency of the molecules of these two liquids (considered as one) to escape is unchanged. When the proportion of the molecules is the same, only “half as many” molecules are able to escape in a “given time”. If one set of molecules have the “same tendency to escape” as opposed to the other, then “the intermolecular forces” between two molecules will be the same for both liquids.

Clark (2004) indicated that if you have an ideal mixture of two liquids A and B, then both A and B will contribute to the total vapour pressure of the mixture. So therefore, if there is an ideal mixture of 1-propanol and 2-proponal, then both of them will contribute to the total vapour pressure of the mixture. Since boiling occurs when the external pressure is equal to the vapour pressure, then the external pressure is taken to be atmospheric pressure (1 atm) and therefore, the total vapour pressure of 1-propanol and 2-propanol will be 1 atm as well however, it will differ based on the mole fractions of each liquid. According to Raoult's law, the partial vapour pressures of A and B at a particular temperature are proportional to the mole fraction.

In the pure 1-Propanol at its boiling point, the molecules with more energy were able to overcome intermolecular interactions and leave the surface forming a vapour in the process.

It should be noted that, when the intermolecular forces are small a lot more molecules are able to escape at a fixed temperature.

At any particular temperature molecules in a certain proportion will have enough energy to be able to exit the liquid's surface.

BOILING POINT AND COMPOSITION

A liquid with a high vapour pressure at a temperature T implies that the molecules are escaping quite easily from the surface of the said liquid.

Another liquid with a low vapour pressure at the same temperature T implies that the molecules of this liquid are NOT escaping at the same rate as opposed to the other.

The molecules of the liquid that will escape easily from the surface imply that the intermolecular forces are “relatively weak” and therefore do not require “much heat” to break the bonds (or interactions) to boil the liquid. Therefore, the liquid with the high vapour pressure will be the one with the low boiling point as compared to the other liquid with the lower vapour pressure.

Another analogy is based on that liquids will boil when the vapour pressures becomes equal to the external pressure, as was mentioned earlier. A liquid with a high vapour pressure means that a high temperature will not be required for this liquid to become equal with atmospheric pressure and thus boil quickly. On the other hand, if the vapour pressure of the other liquid is low, then a higher temperature will be required to raise the vapour pressure to the external pressure so that the liquid boils. (Clark 2004)

ANALYSIS OF THE GRAPHS

(i) VAPOUR PRESSURE AGAINST MOLE FRACTION

The partial vapour pressure of a gas is proportional to its mole fraction (from Raoult's law). The lines, within the limits of experimental error, seem to increase linearly for each liquid. Hence the vapour pressure is increasing almost linearly with increasing mole fraction. The vapour pressure of 2-propanol falls as its mole fraction decreases and the vapour pressure of 1-propanol increases as its mole fraction increase. The vapour pressure of 2-propanol is higher than that of 1-propanol. That means that molecules break away much more easily from the surface of 2-propanol than 1-propanol. Different concentrations of solutions have different vapour pressures. Solutions whose components have low vapour pressures will boil at higher temperatures than solutions in which the components have high vapour pressures. This is due to the fact that the external pressure is attainted at a lower temperature in case of volatile components. Thus the solution with a high vapour pressure starts boiling at a lower temperature.

(ii) BOILING POINT AGAINST % COMPOSITION (MOLE FRACTION & VOLUME %)

The graph shows a curve like increase of the boiling points with respect to the different compositions. One can garner that mixture behaves almost ideally and can therefore be assumed to be an ideal mixture. There are drops when the composition almost reaches 50:50 to about 62.50:37.50 and afterwards increases with the same rigor.

As the composition of 2-propanol increases, and the 1-propanol decreases, there is a rise in the boiling point of the mixture. The question is why? It means that the vapour pressure of the mixture is lowered since there is a reciprocate effect with the vapour pressure and the boiling point of a mixture. It means therefore, that the vapour pressure of the mixture is decreasing as the boiling points are increasing.

Increasing boiling points mean that the intermolecular forces of the two mixtures are increasing so they are harder to break. As a result of this, more energy needs to be introduced and the temperature increases. When the vapour pressure is then equal to the atmospheric pressure, it boils and therefore, the boiling points increases.

The liquid with the lower boiling point has the higher vapour pressure and from the graph, a pure mixture of 2-propanol has the lowest boiling point, as a result, 2-propanol has the higher vapour pressure. The liquid with the higher boiling point has the lower vapour pressure and since 1-propanol has the highest boiling point, it can be inferred to be the one with the lowest vapour pressure.

1-Propanol with the lowest vapour pressure means, it means that a greater amount of heat will need to be applied so that it reaches atmospheric pressure and thus be able to boil. 2-Propanol with the highest vapour pressure means that it will achieve equality with atmospheric pressure quite quickly and thus be able to boil sooner.

Boiling a binary liquid mixture; 1-propanol and 2-propano, it can be said that the more volatile liquid will escape quite easily as compared to the other. In this case, 1-propanol is the more volatile liquid. Therefore, one can expect to find a “higher proportion of 1-propanol”, which is more volatile in the vapour as opposed to it being in the liquid. And therefore one will find that less 2-propanol will be found in vapour as it is less volatile compared to 1-propanol.

Therefore, it can be said that: 1-propanol is the more volatile component where as 2-propanol is the more non-volatile component.

INTERMOLECULAR FORCES OF 1-PROPANOL AND 2-PROPANOL (EFFECT ON VAPOUR PRESSURE AND BOILING POINT)

(I) HYDROGEN BONDING

Organic liquids frequently have hydrogen bonding interactions especially the ones with oxygen atoms in it. The oxygen of one of the propanol molecule forms a hydrogen bonding interaction with the hydrogen atom of another propanol molecule. It can be of the same 1-propanol, 2-propanol or a 1-propanol and 2-propanol molecules. As a result, there is a framework of hydrogen bonding interactions within the mixture.

Starting off with a pure mixture of 1-propanol infers that there is hydrogen bonding with itself. However, when 2-propanol is added to the mixture, hydrogen bonding interactions are formed with the 2-propanol and the 1-propanol. There are still interactions with the same molecules such as 1-propanol and 1-propanol or 2-propanol and 2-propanol. In lieu of the interactions, the molecules are held more closely and bonds are harder to be broken. Hydrogen bonds are generally weak bonds but together pack a punch! This means that the addition of all the hydrogen bonding interactions considered as a will form a frame work and thus will be harder to break.

Now due to the increasing strength of the intermolecular hydrogen bonding interactions, more heat is required to break the bonds, which consequently leads to a higher boiling points of the mixture (with different compositions). It can be inferred from the graphs that by increasing the composition of 1-propanol increase the hydrogen bonding interactions, and consequently requires more energy to break the bonding leading to higher boiling points and lower vapour pressures.

(ii) DIPOLE - DIPOLE INTERACTIONS

As the different molecules of 1-propanol and 2-propanol come into contact with each other by varying the composition) there are dipole - dipole interactions created. Dipole - dipole interactions occur between molecules that have permanent dipoles (those that are polar). The partial positive charges on one molecule are electrostatically attractive to the partial negative charge on a neighbouring molecule. It should be noted that H bonding is a special type of dipole - dipole interactions. But, regardless of H - bonding, there are other dipole - dipole interactions when you consider the molecule as a whole.

As the dipole - dipole interactions increase, they become harder to break, as in H - bonding. As a result, a higher temperature is required to make the vapour pressure of the mixture equal to the external pressure on the mixture and therefore causing boiling to occur.

1-propanol being a primary alcohol is more polar compared to the secondary alcohol 2-propanol. The more polar the molecule is, the stronger the dipole - interactions between itself and a less polar molecule. The result is stronger intermolecular interactions and consequently requiring more energy to pry them apart.

(iii) LONDON FORCES (VAN DER WAALS)

All molecules have the capability to form London forces (Wojciechowski and Cerpovicz 1998). These are solely dependent on the surface area and the polarizability of the surface of the molecule. They result from moving electrons that generate positive and negative charged regions in the molecule. This increases when two different polar are mixed together to form binary liquid. As a result, the interactions are stronger between the 1-propanol and 2-propanol which will require higher temperature to break the bonds and therefore cause the liquid to boil by increasing the vapour pressure to atmospheric pressure.

(B) CHLOROFORM / ACETONE

ANALYSIS OF THE GRAPHS

(iii) VAPOUR PRESSURE AGAINST MOLE FRACTION

The partial vapour pressure of a gas is proportional to its mole fraction (from Raoult's law). The lines, within the limits of experimental error, seem to increase linearly for each liquid. Hence the vapour pressure is increasing almost linearly with increasing mole fraction. The vapour pressure of chloroform falls, as its mole fraction decreases and the vapour pressure of acetone increases as its mole fraction increases. The vapour pressure of acetone is higher than that of chloroform. That means that molecules break away much more easily from the surface of acetone than chloroform. Different concentrations of solutions have different vapour pressures. Solutions whose components have low vapour pressures will boil at higher temperatures than solutions in which the components have high vapour pressures. This is due to the fact that the external pressure is attainted at a lower temperature in case of volatile components. Thus the solution with a high vapour pressure starts boiling at a lower temperature.

(iv) BOILING POINT AGAINST % COMPOSITION (MOLE FRACTION & VOLUME %)

The graph shows a curve like decrease of the boiling points with respect to the different compositions as it moves from chloroform to acetone. One can garner that mixture behaves almost ideally and can therefore be assumed to be an ideal mixture except at the 44:44 to 55:56 fraction. The drop indicates that the boiling point is lower at that interval. This means that the vapour pressure is higher at that composition meaning that there is a greater likelihood of more molecules escaping the liquid's surface. It could have also been inconsistencies in data collection i.e. reading the thermometer incorrectly, or messing up the composition of the mixture. This conclusion can be drawn because, afterwards increases with the same rigor. However, in some cases, this is not unexpected (such as in azeotropes). It just means at that point, the intermolecular forces are weaker and therefore easier for molecules to escape into the vapour phase at a much lower temperature.

As the composition of chloroform increases, and the acetone decreases, there is a fall in the boiling point of the mixture. The question is why? It means that the vapour pressure of the mixture is increased since there is a reciprocate effect with the vapour pressure and the boiling point of a mixture. It means therefore, that the vapour pressure of the mixture is increasing as the mole fractions are increasing.

Decreasing boiling points mean that the intermolecular forces of the two mixtures are decreasing so they are easier to break. As a result of this, less energy needs to be introduced and the temperature decreases. When the vapour pressure is then equal to the atmospheric pressure, it boils and therefore, the boiling points decreases.

The liquid with the lower boiling point has the higher vapour pressure and from the graph, a pure mixture of acetone has the lowest boiling point, as a result, acetone has the higher vapour pressure. The liquid with the higher boiling point has the lower vapour pressure and since chloroform has the highest boiling point, it can be inferred to be the one with the lowest vapour pressure.

Chloroform with the lowest vapour pressure means that a greater amount of heat will need to be applied so that it reaches atmospheric pressure and thus be able to boil. Acetone with the highest vapour pressure means that it will achieve equality with atmospheric pressure quite quickly and thus be able to boil sooner.

Boiling a binary liquid mixture; chloroform and acetone, it can be said that the more volatile liquid will escape quite easily as compared to the other. In this case, acetone is the more volatile liquid. Therefore, one can expect to find a “higher proportion of acetone”, which is more volatile in the vapour as opposed to it being in the liquid. And therefore one will find that less chloroform will be found in vapour as it is less volatile compared to acetone.

Therefore, it can be said that: acetone is the more volatile component where as chloroform is the more non-volatile component.

COMPARISON WITH LITERATURE VALUES

The boiling points recorded for the pure mixtures are follows

THE METHODOLOGY

(i) Critiquing Each Step

The boiling point of the liquids were obtained (either by combining in fractions or separately) so the mixture had to be heated in order to reach the boiling point (vapour pressure = external pressure). When the liquid mixture is heated, it eventually boils, and if this operation were to be carried out in a beaker (or open flask) there would be several undesirable consequences of this. The vapours from the boiling liquids escaping from the apparatus would in all probability be toxic or, at the very least, harmful to health. These vapours would also be likely to be flammable, so there would be the additional danger of fire. Finally - the least of the problems - the escape of vapours from the boiling mixture would reduce the liquid contents in the beaker or open flask, so that by the time the reaction had completed, there would be no liquid product left!

The way to get round all these problems is to use “reflux conditions”. In these, whilst the mixture is heated, the rising vapour comes into contact with the cold glass of the condenser, this vapour then condenses back into liquid and this is subsequently returned to the flask for further “processing”. The reflux system prevents the reactants and products from escaping into the open laboratory and ensures that all the components remain in the flask until the reaction is complete.

The liquid was allowed to cool briefly before adding the second liquid. This ensured that there were no unnecessary accidents. Often due to rapid heating and cooling cracks can occur in the apparatus in question, consequently leading to instrument being unusable or worse, causing injury.

Further additions of the other liquid were done to ensure that the mole fractions of the mixtures changed. This allowed obtaining boiling points and different mole fractions and thus obtaining a boiling point - composition graph and by extension a vapour pressure - composition graph. This was done until 10 cm3 of the other liquid was added and allowed a 50:50 composition of the first liquid to the other. In other words, the composition of the 1st liquid was ½ and so was the 2nd. This step was reversed so that the boiling points would be ascertained for the next set of mole fractions so that it will go from 0.00 to 1.00 and vice versa.

(ii) Problems During The Experiment

As it relates to accuracy of data, there could have been

(iii) Possibilities Of Improving The Experiment And Results

In order to plot the vapour pressure - composition curve, the antoine's equation was used which also assumed that the mixture was ideal and therefore most, if not all the intermolecular forces within the mixture were equal. It would have been better if the vapour pressure could have been calculated directly, or even indirectly with less approximations as these. The result of the approximation perhaps, gave a graph that was in no way close to the correct vapour pressure graph considering that chloroform and acetone formed a non - ideal mixture. It perhaps would work for a mixture of 1-propanol and 2-propanol which is considered to be ideal. The methodology therefore needs to be re constructed to allow for this.

The experiment could have been more done in a more useful way instead of just getting boiling point composition graphs and to a lesser extent vapour pressure composition graphs. It would have been more useful, if accurate tests were done to give more meaningful data that would be reliable enough to be published in a journal (not to be published but the quality and accuracy needs to be adhered to). When dealing with non - ideal mixtures, approximations are not enough as it most often leads to inaccurate data.

A better approach in finding the vapour pressure, if the partial pressures are not easy to be found would be to measure the total pressure, then use a refractometer or another method to ascertain the composition of the vapour and then calculate the vapour pressure from that. It is more of an indirect method, but it will give better results and will not be based on so much on approximations.

REFERENCES

1. Anand, A and Negi, S. A Textbook of Physical Chemistry. USA: John Wiley & Sons, 1985.

2. Atkins, Peter and De Paula, Julio. 2001. Physical Chemistry 6th Edition. USA : W. H Freeman & Company, 2006.

3. Castellan and Gilbert. 1983. Physical Chemistry 3rd Edition. Massachusetts: Addison Wesley Publishing Company, 1983.

4. Flowers and James. 2004. Cracking the MCAT with CD-ROM. USA: Princeton Review, 2004.

5. Lide, David. 1993. Handbook of Chemistry and Physics 85th Edition. USA: CRC, 1993.

6. Mortimer, Roger. 2008. Physical Chemistry 3rd Edition. Canada: Elsevier Academic Press, 2008.

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