Physical Chemistry Experiment

PARTIALLY MISCIBLE LIQUIDS:

DETERMINATION OF MUTUAL SOLUBILITY

Abstract

A phenol - water solution was used to determine the solubility of two partially miscible liquids.  The group calculated the volume of water required to prepare the following mixtures with volume percentage ranging from 5% to 95% sample at 5% increment using 10mL phenol sample.  The different volume ratios of mixtures prepared were subjected to constant heating and cooling in order to gather the needed temperature necessary for the construction of the mutual solubility curve of Phenol- Water solution.  The critical solution temperatures were determined at 30% Phenol - 70% Water ratio, 64˚C (for single phased region) and 61.8˚C (for double - phased region)…

I. Introduction

Oil and water don’t mix. Pouring 10 mL of olive oil into 10 mL of water results in two distinct layers, clearly separated by a curved meniscus. Each layer has the same volume and essentially the same composition as the original liquids. Because very little mixing has apparently occurred, the liquids are called “immiscible” or unmixable.

Pouring grain alcohol into water results in a single liquid phase. No meniscus forms between the alcohol and the water, and the two liquids are considered “miscible”. Nearly any pair of liquids is miscible if only a trace amount of one of the liquids is present.

Many liquid mixtures fall between these two extremes. Two liquids are “partially miscible” if shaking equal volumes of the liquids together results in a meniscus visible between two layers of liquid, but the volumes of the layers are not identical to the volumes of the liquids originally mixed. For example, shaking water with certain organic acids results in two clearly separate layers, but each layer contains water and acid (with one layer mostly water and the other, rich in acid.)  Liquids tend to be immiscible when attractions between like molecules are much stronger than attractions between mixed pairs. (Logan, 1998)

The objectives of this experiment are 1) to determine the solubility of two partially liquids (phenol - water solution), 2) to construct a mutual solubility for the pair, and 3) to determine their critical solution temperature.

II. Review of Related Literature

Mutual solubility of polymers and properties of their mixtures

“Heats of mixing of polymers with each other have been measured, the behavior of the mixtures of solutions of various polymers has been studied, and the dependence of mechanical properties of polymer mixtures on the ratio of components has been investigated. It has been shown that mixing of polymers with each other is usually an endothermic process and, therefore, leads to formation of macroscopically homogenous, but actually microheterogenous, systems with an extremely high degree of dispersion. These microheterogenous polymer mixtures are formed because of the enormous viscosity of polymer mixtures, which prevents macroscopic separation into phases but does not hinder the considerable mobility of the segments of flexible chain molecules. It has been shown that the dependence of mechanical properties of microheterogenous polymer mixtures on the ratio of polymers in the mixture have sharp maxima or minima which cannot be found in the case of true polymers in polymer solutions. It has been found that the behavior of some polymer pairs is anomalous, in that exothermal mixing is supplemented by separation of the solution mixture into phases and by the appearance of maxima or minima in the dependences of the properties of polymer mixtures on the ratio of polymers in the mixture. This anomaly has been attributed to the effect of loose packing of the molecules of the polymers which show anomalous behavior. It has been shown that, in these systems, there necessarily exists a lower critical temperature of mixing whose value can be decreased by adding low-molecular solvents to the loosely packed polymer. Attention has been drawn to the fact that, although mixing of amorphous polymers should be considered on a thermodynamic basis to be a mutual solution of two liquid phases, the large dimensions and the flexibility of polymer chain molecules require a critical revision of the possibility of formal application of the basis thermodynamic concepts and relations to a theoretical analysis of the behavior of polymer mixtures.” (Slonimski, 1998)

Equations of state for the calculation of fluid-phase equilibria

“Progress in developing equations of state for the calculation of fluid-phase equilibria is reviewed. There are many alternative equations of state capable of calculating the phase equilibria of a divers (Sadus and Song Wei)e range of fluids. A wide range of equations of state from cubic equations for simple molecules to theoretically-based equations for molecular chains is considered. An overview is also given of work on mixing rules that are used to apply equations of state to mixtures. Historically, the development of equations of state has been largely empirical. However, equations of state are being formulated increasingly with the benefit of greater theoretical insights. It is now quite common to use molecular simulation data to test the theoretical basis of equations of state. Many of these theoretically-based equations are capable of providing reliable calculations, particularly for large molecules.” (Sadus and Ya, 2000)

Mutual solubility study for 94.2:5.8 of ethanol to octane with supercritical carbon dioxide solvent

“Solubility data of a mixture containing 94.2% ethanol and 5.8% octane was measured in carbon dioxide solvent using a high-pressure type phase equilibrium apparatus at pressures up to 103.5 bar and at temperature of 75 °C. The results showed that considerable separation was not achieved in this ethanol and octane ratio. However, the experimental data were then compared with the theoretical data which were obtained from two models which are regular solution theory and Redlich-Kwong equation of state. Regular solution theory is employed to each phase by applying activity coefficient expressions. Redlich-Kwong equation of state is employed to the vapor phase and then with applying fugacity coefficient, liquid phase data is obtained. The regular solution theory as a novel model approach has been found to be encouraging for the prediction of phase equilibria solubilities. It concluded that the regular solution theory model could predict two phases equilibrium data better than Redlich-Kwong equation of state.” (Davarnejad et al, 2008)

Solubility, miscibility and their relation to interfacial tension in ternary liquid systems

“The terms, miscibility and solubility, are widely used in phase behavior stud (Ayiralam and Rao)ies of multicomponent hydrocarbon systems. The distinction between these two terms appears to be still hazy, leading to their synonymous use in some quarters. Also, the relation of these two thermodynamic properties with interfacial tension has largely remained unexplored. However, recently a new experimental technique of vanishing interfacial tension (VIT) has been reported relating miscibility with interfacial tension in gas-oil systems. Therefore, the objectives of this study are to correlate miscibility and solubility with interfacial tension and to investigate the applicability of the new VIT technique to determine miscibility conditions in ternary fluid systems. For this purpose, a standard ternary liquid system of benzene, ethanol and water was chosen since their phase behavior and solubility data were readily available. The interfacial tensions of benzene in aqueous ethanol at various ethanol enrichments were measured using the drop shape analysis (DSA) and capillary rise techniques.

The experimental results indicate the applicability of VIT technique to determine miscibility conditions for ternary liquid systems as well. Comparison of IFT measurements with solubility data showed a strong mutual relationship between these two properties, in addition to demonstrating a clear distinction between solubility and miscibility. The interfacial tension appears to be independent of solvent-oil ratio in feed, provided that complete equilibration of fluid phases is allowed to incorporate all the mass transfer effects during experimentation. All these experimental observations have immense application in fluid-fluid phase equilibria studies and to determine the miscibility conditions of gas injection improved oil recovery projects.” (Ayiralam et al, 2006)

Heats of mixing of the partially miscible liquid system cyclohexane + methanol

“The molar excess enthalpies of cyclohexane + methanol were systematically measured with a Picker flow microcalorimeter operated in the discontinuous mode at 298.15, 303.15, 308.15, 313.15, 318.15 and 323.15K. Our measurements are higher than the literature data. This work shows that molar excess enthalpies increase with temperatures, an the straight segments of the excess enthalpy curves in the mid region become shorter with increasing temperature. At 323.15K the curve become one of a miscible liquid system, and the position of the maximum value is at X=0.6. In addition, the calorimetric measurement can be used to determine the compositions of two immiscible phases for the binary mixture.” (Dai and Chao, 1985)

III. Methodology

Apparatus and Materials

Phenol Sample

Distilled water

Stirring Rod

Hot plate

200ml beaker

Thermometer (0.1 deg calibration)

1 L Beaker (2pcs)

A.   Preliminaries

Before the experiment, the group calculated the volume of water required to prepare the following mixtures with volume percentage ranging from 5% to 95% sample at 5% increment, using 10mL sample in each proportion.  The calculations should be approved before proceeding.

B.   Experiment Proper

After the preparation of a 95% sample - 5% water volume to volume mixture based on 10mL of the sample, [Caution: All the samples are corrosive while triethylamine is flammable, a lachrymator and readily forms explosive in air] the mixtures was heated in a water bath with mild stirring until the cloudiness in it disappears. Its temperature was noted. It was cooled in a second water bath with mild stirring until the cloudiness appears. Once again, the temperature was noted. This process was repeated until a fairly constant reading was observed for a specific volume ratio mentioned in the preliminaries. Constant temperature was recorded.

IV. Data and Discussion

After the group prepared a 95% phenol -5% water volume to volume mixtures based on 10mL of the sample (see Table I), the mixture was heated in a water bath with mild stirring and recorded its constant temperature until the cloudiness of the solution disappeared and cooled instantly until the cloudiness appears. As shown in Table I, it shows that on 95% phenol - 85% phenol and 5% phenol, there is no significant changes appeared both for heating and cooling of the mixtures due to the concentration of phenol in the solution.  Cloudiness of the solution started to appear at 80% phenol.

Table I: The prepared amount of water needed to add at the given

percentage of Phenol - Water Solution based on 10ml sample

% Phenol by Volume	Volume of Added Water / mL
95	0.53
90	0.59
85	0.65
80	0.74
75	0.83
70	0.95
65	1.10
60	1.28
55	1.52
50	1.81
45	2.22
40	2.78
35	3.57
30	4.76
25	6.67
20	10.00
15	16.67
10	33.33
5	100.00

Based on the data in Table II, the group constructed the mutual solubility curve for Phenol - Water solution which is important in the determination of the critical temperature of the mixture.  Critical solution temperature is the temperature at which a mixture of two liquids (Phenol and Water for this experiment), immiscible at ordinary temperatures, cease to separate into two distinct phases.  The black line symbolizes the temperature reading of phenol in the hot water bath while the gray line is the temperature reading of phenol in the cold water bath.

On the other hand, the red curve is the polynomial trend line of mixture in the hot water bath in which the solution tends to become in single phased (no cloudiness appeared), while the blue curve is the polynomial trend line of the cooled mixture that tends to start the double phased region (appearance of cloudiness).  The critical temperatures of the solution was located at 30% phenol - 70% water,   64˚C (heating) and 61.8˚C (cooling).

Table II : Constant temperature reading of Phenol - Water Solution at heating and cooling process on 10ml sample.

% Phenol by Volume	Constant Temperature Reading /˚C
	HEATING	COOLING
95	no change appeared	no change appeared
90	no change appeared	no change appeared
85	no change appeared	no change appeared
80	36.3	32.8
75	41.2	35.2
70	43.7	39.7
65	46.8	42.9
60	52.5	47
55	53	50.9
50	54.2	52.4
45	58.5	56.7
40	65.2	62.2
35	68.7	66.1
30	64	61.8
25	60.4	56.2
20	54.9	50.7
15	51.3	47.2
10	48.5	43.6
5	no change appeared	no change appeared

Figure 1. Mutual Solubility Curve of Phenol - Water Solution

V. Conclusion and Recommendation

Throughout the experiment, the critical solution temperature of the solution was 64˚C (heating) and 61.8˚C (cooling) at 30% phenol - 70% water.  There are factors that affect the solubility of the mixtures, the nature of solute and solvent, the temperature and the pressure.

a)      Nature of Solute and Solvent

• Molecular Size - The larger the molecule o the bigger its molecular weight, the less soluble the substance will be.
• Polarity - Polar solutes will dissolve polar solvents; Non - polar solute molecules will dissolve non- polar solvents

b)      Temperature

If the solution process absorbs energy, then the solubility will be Increased as the temperature is increased.  If the solution releases energy, then the solubility will Decreased with increasing temperature.

c)       Pressure

If solid and liquid, there is no change in solubility if pressure changes, likewise, in gas, as pressure increased, solubility also increases.

Cloudiness is significant in this experiment for immiscible liquids.  Through cloudiness, we cansay that the substance is still unmix due to the presence of stable emulsion, but when a completely clear solution with no trace of cloudiness, we can assume that the substance is mixed

Appendix

Ayiralam, Subhash C. and Dandina N. Rao. “Solubility, miscibility and their relation to interfacial tension in ternary liquid systems.” Fluid Phase Equilibria 249.1 (2006): 82-91.

Dai, Ming and Jian-Ping Chao. “Heats of mixing of the partially miscible liquid system cyclohexane + methanol.” Fliuid Phase Equilibria 23.2 (1985): 315-319.

Davarnejad, R, K.M Kassim and A Zainal. “Mutual solubility study for 94.2:5.8 of ethanol to octane with supercritical carbon dioxide solvent.” Journal of the Chinese Institute of Chemical Engineers 39.4 (2008): 343-352.

Logan, R.S. “The Behavior of a Pair of Partially Miscible Liquids.” Chemical Education 75.339 (1998): 206-208.

Sadus, Richard J and Ya Song Wei. “Equations of state for the calculation of fluid-phase equilibria.” AIChE 46.1 (2000): 169-296.

Slonimski, G.L. “Mutual solubility of polymers and properties of their mixtures.” Journal of Polymer Science 30.121 (1958): 625 - 637.

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Experiment
SOLUTIONS AND SOLUBILITIES

INTRODUCTION

What is a solution? Solutions are homogeneous mixtures of two or more pure substances. A homogeneous mixture is a physical combination of two or more pure substances whose distribution throughout the mixture is uniform. What this means is that if we were to make a solution and take only a portion of the solution at random called an aliquot, the proportion of each pure substance in the aliquot would be the same as the proportion of that pure substance in the whole solution. We call these proportions to the whole solution the concentration.
Within any solution there are electrical interactions called intermolecular forces between the solute and solvent molecules. In aqueous solutions the polar water molecules will be attracted to other molecules that are also polar.

The water molecules will surround each solute molecule and form a “solvent cage” isolating each solute molecule. We say the solute molecules have been solvated. If the bonds holding the solute molecule together are weak enough, the weakest bond within the solute molecule could be broken by the electrical force of attraction the water molecules have for the solute molecules. If this should take place the solute molecule will undergo what is called ionization forming ions (fragments of the solute molecule) within the solution. In some cases the ionization process can be quite complete while in other cases the ionization process will occur in a very limited manner. Those solutions that do not undergo ionization have the solute molecules solvated within the solution and are called non-electrolytes. Such solutions will not conduct electrical current because of the absence of ions within the solution. Those solutions whose molecules undergo ionization during the solution process are called electrolytes and those solutions will conduct electrical current because of the presence of ions.

Energy Transfers In Solution Formation

When solutions form the solvent molecules form solvent cages around the solute molecules. This solvation process involves a certain amount of thermal energy to be exchanged between the solution system and its immediate surrounding environment. Some solutions absorb energy as they are formed. The solution is said to have an endothermic Heat of Solution.

Solute + Solvent + Thermal Energy —-> Solution

Most aqueous solutions involving liquid or solid solutes will have endothermic Heats of Solution. However, a few are exceptions to this statement.
Other solutions involving gaseous solutes in water will release thermal energy during the solution formation process. These solutions are said to have an exothermic Heat of Solution.

Solute(g) + Solvent —–> Solution + Thermal Energy

Solutions can be unsaturated, saturated, or supersaturated. Unsaturated solutions are those that are below the solubility limits of the solute in that solvent. Saturated solutions are those that are at the solubility limits. Supersaturated solutions are those solutions that are above the solubility limits. Supersaturated solutions are meta stable. Such solutions will have the excess solute crystalize out with any disturbance of the supersaturated solution. Providing a tiny crystal of the solute or scratching the sides of the container which introduce micro chips of glass into the supersaturated solution so the excess solute can crystalize upon its surface, will illicit the dramatic crystallization of the excess solute out of the solution restablishing a saturated solution.

Solubility of Solutions

The solubility of a particular solute in a solvent is the maximum amount of solute that will dissolve in a specified amount of solution or solvent. It represents the saturated level of the solution where no more solute will dissolve within the solution. This saturated condition creates a dynamic physical equilibrium between the solute and solvent and the solution:

solute + solvent = solution

Such dynamic equilibrium involve two processes, a forward process and a reverse process. When the rate of the forward process is equal to the rate of the reverse process, then the system is said to be in a dynamic equilibrium. Dynamic Equilibria will involve two opposing processes occuring simultaneously. We can prove this by taking a salt crystal and chipping off one end. Then we suspend the deformed cystal in a saturated solution of the salt and allow the saturated solution to be in contact with the deformed cystal for several weeks. When we come back we will have discovered that the deformed crystal will have been reformed. This could only happen if the solution process and its reverse process, the dissolution process, were occuring simultaneously.

After this experiment, we would be able to relate the effect of particle size on the rate of solubility and we should be able to determine the effect of temperature and nature of solute and solvent on solubility

METHODOLOGY

Materials and Apparatus

Coarse and finely ground NaCl
Distilled water
Hexane
Iodine crystals
Ethyl alcohol
KClO3
Thermometer
Stirrer
Ignition tube
Hot plate
Test tube holder
600.0 mL beaker
Erlenmeyer flask

Procedure

A. Rate of Solution and Particle Size

1. Place about 0.5 gram coarse sodium chloride (NaCl) crystals in a test tube. Add approximately 5.0 mL of water to the crystals. Shake the test tube contents. Record the time it takes for the crystals to dissolve.

2. Repeat step 1, using a finely ground sodium chloride (NaCl). Record the time it takes for the crystals to dissolve

B. Solubility of Various Substances in Different Solvents

B.1 Solid in Liquid

1. Prepare two clean, dry test tubes. To each of the test tube, place a pinch of sodium chloride (NaCl). Use finely ground sodium chloride.

2. To the first test tube, add 1.0 mL of distilled water. Shake to mix contents.

3. To the second, add 3.0 mL of hexane and mix the contents. Then, add 6.0 mL of distilled water to the mixture and shake well.

4. Prepare three more test tubes. To each of the test tube, place a pinch of iodine crystals

5. To the first test tube, add 3.0 mL of hexane and shake well.

6. To the second test tube, add 10.0 mL of distilled water and mix the contents well. Then, add 3.0 mL of hexane to the mixture and shake well.

7. To the third test tube, add 3.0 mL of ethyl alcohol. Record all observations. Indicate solubility of the mixtures as follows:

• soluble – when all of the solid dissolve
• insoluble – when none of the solid dissolves
• partially soluble – when not all of the solid dissolve

B.2 Liquid in Liquid

1. Prepare two dry test tubes and to each of the test tube, place 3.0 mL of ethyl alcohol.

2. To the first test tube, add 3.0 mL of distilled water.

3. To the second test tube, add 3.0 mL of hexane.

4. Record all observations and indicate miscibility of the mixtures as follows:

• miscible – when only one phase is formed
• immiscible –when two distinct phases are formed.

C. Effect of Temperature on Solubility

1. Heat the water bath to near boiling.

2. Weigh approximately 8.0 g of KNO3 . Place it in an ignition tube.

3. Measure the assigned amount of volume of the distilled water for each group. (The volume of distilled water assigned for our group was 14.0 mL)

4. Run down the distilled water along the side of the ignition tube.

5. Immerse the test tube in the heated water bath. Continue the heating process and constantly stir the mixture until all the solid KNO3 have completely dissolved.

6. Remove the test tube from the water bath. Transfer it into an ice –water bath and wait until crystals start to appear.

7. Record the temperature of crystallization to solubility temperature.

8. Share experimental results gathered by the other groups in order to plot the temperature against solubility

TREATMENT OF RESULTS

A. Rate of Solution and Particle Size

Through this experiment we had determined that the fine ground particles will dissolve faster than the coarse particle. In just 1min 16.4 secs, fine particles already dissolve unlike coarse particle dissolves for about 1 min 45.1 seconds. It is 29 seconds difference between their time.

B. Solubility in Different Solvents

Through the data we had gathered, we determined that solute may be soluble, miscible or not soluble or immiscible or maybe partially soluble or partially miscible in other solvents. For example water, H2O, and ethyl alcohol are completely “miscible”. Both water and ethanol are polar molecules with hydrogen bonding between molecules too. Also water, H2O and NaCl are completely “soluble” while water and I2 is somewhat “partially soluble” with one another.

C. Effect of Temperature on Solubility

After we gathered the saturation temperature of KNO3 in different volume of H2O, we find out that the solubility of KNO3 is some- what directly proportional to temperature. It means that when the temperature of the solution increases, the solubility of the solid also increases.

(solubility ∞ temperature)

DISCUSSION OF RESULTS

Questions

1. Why do finely ground particles dissolve faster than the coarse particles?

Finely ground particles dissolve faster than the coarse particles it is because the H2O molecules can penetrate better in finely ground particle because of its area unlike the coarse particle.

2. What may happen as a solvent is added to a solute? How would you qualitatively compare the behavior of one solvent on a solute?
When a solvent is added to solute, solubility may take place or not. It depends whether the solute is soluble or miscible to the solvent. The solute cannot be dissolved when the volume of the solvent is not proportion to the weight of the solute. It means that the correct volume of the solvent is important in order to dissolve the solute.

3. What significant factors can you give to account for the interaction between solute and solvent?
There are many significant factors that can give account for the interaction between a solute and a solvent. First factor is the solute/solvent interactions and their identities.

• Identity of solvent
• Like dissolves like rule

polar solvents such as water are best solvent for ionic and polar compounds such as common salt

non-polar solvents such as benzene and tetrachloroethene are best for non-polar compounds such as wax, grease

• Identity of solvent

• For ionic compounds - solubility rules - larger anions with low charge are more soluble

• lack of nitrates in mineral deposits - high solubility

• bones of aquatic animals not dissolved - low solubility of phosphates in water

• hard water due to the presence of carbonates and bicarbonates of calcium and magnesium

Second factor is the temperature. Generally speaking the water solubility of a liquid or solid will increase with increasing temperature. However, there are some exceptions to this. Some solutes like solid Ce2(SO4)3 will have a decreasing water solubility with increasing temperature. This depends upon the thermodynamics of the solution process. Since a solution that has reached the solubility limit is in equilibrium, the various laws of equilibrium apply to that system. There is a principle that will be more completely discussed in a future lessons called Le Chatlier’s Principle. The principle says that when a stress is applied externally to an equilibrium, the equilibrium is disrupted temporarily and will shift in such a way as to undo the stress that had been applied. One such stress that can be applied is temperature change.

According to the Principle, increasing the temperature of an equilibrium will always favor the endothermic process of an equilibrium since it is the endothermic process that can absorb the added energy resulting from the increase in temperature. That effectively counteracts the temperature increase. Since most liquid and solid solutes dissolved in water have the solution formation process endothermic, that would be favored when the temperature was increased resulting in an increase in the solubility limit. However some solutes like Ce2(SO4)3 have an exothermic Heat of Solution. Since increasing the temperature will always favor the endothermic process the dissolution (solution breakdown) process endothermic) will be favored and the solubility will decrease. Gaseous solutes always have an exothermic Heat of Solution. Consequently, the solubility of all gases in water decrease with increasing temperature. That is why carbonated drinks that have Carbon Dioxide gas dissolved in them will become “flat” tasting when heated. The sparkle of the drink will have disappeared along with the Carbon Dioxide gas.
And the last significant factor is the pressure. Pressure changes above the solution do not affect the solubility limits of solids or liquids dissolved in water. However gaseous solutes are affected. If the pressure of the gas is increased above the gaseous solution then the solubility will be increased in a linear fashion. This was investigated by Henry and resulted in Henry’s Law.

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