Question

outline a lab experiment which could be done in order to determine delta H, delta G and dellta S of dissolution of Urea

Answer 1

Background theory to determine the asked thermodynamic variables:

Upon dissolution of solids in liquids the entropy of the system increases. Therefore, the entropy change for the first step upon dissolution"process" is positive (ΔS₁ = +). In the system, the solvent molecules must move about to make space for the solute molecule smoemnt. This causes in disruptions of the solvent-solvent forces and more freedom of movement for the solvent molecules, thus additional pathways for energy dispersal. So ΔS₂ is typically positive. In the final step solvent and solute interactions decrease the free movement of both the species to certain extent and the entropy will decrease.

It is possible to find the heat of the solution in the lab but determining the entropy of solution (ΔS_soln) is very cumbersome. In this experiment the thermodynamic requirements for the equilibrium condition can be explored to determine ΔS_soln from the determined heat of solution and the free energy change. In a saturated solution of urea the following equilibrium can exist:

urea (s)                    urea(aq)

Therefore, the equilibrium constant expression for this reaction is:

K _c = [urea(s)] / [urea(aq)]

Because the solid has unit activity (and a constant concentration) it is generally not given as part of the constant so this expression may be written as:

Kc = [urea(aq)]

The relationship between the equilibrium constant and the standard free energy change is given in the following realtion:

ΔG_o = -RT ln Kc

where R = 8.31 J/mol·K and T is the temperature in kelvins. If the concentration of urea in a saturated solution can be determined then Kc is known and it is possible to calculate the free energy change for the dissolving of urea (or at least an approximation of ΔG_o ). Combined with the heat of solution from a calorimetry experiment this allows the calculation of ΔS_soln since:

ΔG_soln = ΔH_soln - TΔS_soln

The measurement of urea concentration can be done by a variety of methods. In this experiment we will use an adaptation of a colorimetric method. This assay is so sensitive that the saturated urea solution has to be diluted 20,000-fold to give a suitable concentration.

The diluted urea solution is combined with a test reagent which consists of an acidic solution of diacetyl monoxime and some additional minor constituents. In acidic solution the monoxime hydrolyzes to give 2,3-butanedione

Scheme is given separately as an attachment

The resulting five-membered ring compound is pink in color. The intensity of the color is proportional to the concentration of urea so that a comparison can be made with standards in accordance with Beer's Law.

The Experiment

There are three parts to this experiment:

• determination of the heat of aqueous solution of urea

• determination of the specific heat of the urea solution

• determination of the urea concentration in a saturated solution

The following chemicals and apparatus is needed to carry out the carry experiment:

• solid urea

• expanded polystyrene calorimeter

• CBL w/thermometer probes

• matched pair of resistance-heater calorimeters w/power supply and cable

• saturated urea solution [record temperature from label]

• 100 mL volumetric flasks • micropipettor w/tips

• 4 urea solution standards [record concentrations from labels]

• six 13 x 100 mm test tubes in rack

• Fisher Model 415 Spectrophotometer w/cuvettes

• BUN reagent

The first two parts of the experiment involve a series of calorimetry measurements.

About 4 g of urea in 50 mL of water should give a spectacular change in temperature. Despite the fact that the urea is soluble, it is good to use the magnetic stirrer. The determination should be done in triplicate with the first trial reserved for the next step. Masses for the solution components are necessary. To be able to calculate accurately the heat involved in the temperature change of the urea solution, it is necessary to know the specific heat of the mixture. Typically the addition of solute lowers the specific heat of water. One way to determine the specific heat of the mixture is to place a sample in an electrically heated calorimeter which is connected in series with an identical calorimeter containing water. As electricity flows through the circuit both solutions receive the same amount of heat energy but their temperature changes will be different owing to their different specific heats. The only unknown in this situation is the specific heat of the urea solution. Using the first solution from the three trials allows time for the mixture to come back to room temperature. This is not absolutely necessary but it tends to give better results since heat exchange with the surroundings is less. The mass of the solution added to the calorimeter and the mass of water in the other calorimeter (about 50 mL) should be recorded. Filling the calorimeters before wiring them together helps to stabilize them (in any case the electricity should not be allowed to flow if the calorimeters are empty; the heaters will burn up). With the power supply set on 12 v the liquids should be heated until the water temperature has changed by about 10 degrees. Initial and final temperatures should be recorded. The following procedure will be done by the instructor and temperature data will be given along with the prepared solution: About 35 g of urea should be mixed with 25 mL of 0.1% benzoic acid water solution (the benzoic acid acts as a preservative to retard bacterial growth). After vigorous stirring, the mixture is allowed to stand at least overnight. After recording the temperature of the solution, the mixture is then gravity filtered. The saturated solution must be diluted 1:20,000 using special glassware. The dilution of 1:20,000 can be made in the following fashion: remove 1000 μL (1 mL) of the saturated urea with the micropipette (which is shown in the following figure) and place it in a clean, dry 100 mL volumetric flask.

Figure. Micropippet (Care should be taken to not press harder so that the graduation will not be disturbed)

The volumetric flask has a line etched on the neck somewhere. You need to fill the flask with distilled water so that the meniscus sits on this line (at eye-level). Stopper firmly and invert at least 20 times to insure through mixing.!! Using a clean tip, 500 μL of the new solution is withdrawn and transferred to another 100 mL volumetric flask and diluted as before. 170 The two BUN reagents must be mixed the same day you will make the absorption measurements. It will be dispensed from burets. 2.5 mL of the BUN solution mixture is dispensed into each of six clean, dry test tubes. The "blank" has 0.150 mL (150 μL) of water added. The remaining five samples have an identical amount of the four standards and the diluted urea solution added. Because the test is extremely sensitive it is imperative that separate tips be used for each new solution. The six tubes should then be immersed in already boiling water for exactly 10 minutes, and then cooled for 5 minutes under cold tap water. Transfer the "blank" mixture to the cuvette provided. Use this solution to zero the spectrophotometer. Empty this solution back into its original test tube and rinse the cuvette with a little of the solution from the next tube. Discard the rinse and then fill the cuvette and measure its Absorbance at 520 nm. Continue in this way to record the Absorbance of each sample.

The Report Your initial calculations should include:

[you may assume the various calorimeter constants all = 0].

1.The specific heat of the urea solution

2. The heat absorbed or released per mole of urea (i.e., ΔHsoln) [best value/standard deviation, 95% confidence, relative error]

3. A calibration graph for urea at 520 nm with the BUN reagents, Absorbance vs. concentration (blank = zero concentration)

4. The concentration (M) of the diluted saturated solution (as read from the calibration graph)

5. The concentration (M) of the original saturated solution of urea

6. The free energy change, ΔG, for the dissolution of urea [relative error]

7. The entropy change, ΔS, for the dissolution of urea at the appropriate temperature [relative error]

Your conclusion to this experiment should include a brief discussion of discrepant data. The following values for the thermodynamics of the dissolution of urea are taken from Basic Tables in Chemistry by R. Keller: ΔG = -6.86 kJ/mol ΔH = +13.8 kJ/mol ΔS = +69.4 J/mol K A brief summary of the probable signs and magnitudes of the enthalpy and entropy changes in the hypothetical steps for the formation of the solution as described in the Background section should also be included.