The Solubility and Thermodynamics of Borax Lab Report PDF

Preview

Summary

Lab Report...

Description

The Solubility and Thermodynamics of Borax Rachel Blasczyk, Carter Pendergrass, Doug Macarthur, Alyssa Wallenta

Abstract The purpose of this experiment was to determine the enthalpy (△H°), entropy (△S°), the value of Gibbs free energy (△G°), and determine the KSP of borax. The experimental plan consisted of generating a KSP versus temperature graph that measured enthalpy, entropy, and the value of Gibbs free energy via the relationship to the equilibrium constant to thermodynamics. The hypothesis was if temperature increased, then the Ksp for borax increased. The equilibrium constant is the free energy change in the reaction: △G°= -RT ln K. The free energy change is also written as: △G°= △H°-T△S° to account for the enthalpy and entropy changes during the reaction. Both equations are combined and divided by the negative rate times temperature, and the relationship to the equilibrium constant becomes: ln K = -△H°/RT + △S°/R where T is in Kelvin and R = 8.314 J/Kmol. The enthalpy and entropy is determined by linear regression of the y=mx+b equation through the Ksp versus temperature graph. Through plotting each five of the solutions on the graph on page five, it was determined through the equation: ln K = -△H°/RT + △S°/R that enthalpy(△H) = -1682.64 kJ/mol and entropy(△S) = 5.2555 J/Kmol. It was also determined that as temperature increased, so did Ksp for borax.

Introduction The objectives of this lab were to determine the values of enthalpy, entropy and Gibbs free energy for the dissolution of Borax in water, and then use these values to determine the equilibrium constant (Ksp) of Borax. The value of enthalpy (△H°), is the change in temperature. Entropy (△S°) is the change in disorder. The value of Gibbs free energy (△G°) is the energy associated with a chemical reaction that can be used to do work, and represents the relation between entropy, temperature and enthalpy.

Experimental Method A hot plate was set to a heating level of 7, while the hot plate was warming up a solution of 25.00 grams of borax was mixed with 80.00 mL of water, in a 100. mL beaker. This beaker was then placed on the hot plate and stirred continuously, while the temperature was monitored by a digital thermometer. The thermometer was not resting on the bottom of the beaker, to insure an accurate reading. The beaker was removed from the hot plate after reaching a temperature of 57.0°C. After the initial heating period, the hotplate was tuned down to a heating level of 4, and a 100.0 mL beaker half filled with water was placed on it. Then a graduated cylinder was placed into the beaker for later use. The beaker containing saturated borax was stirred until the temperature reaches 55.0°C and then left to settle. When the solid settled out, 7.09.0 mL of the borax solution was poured into the graduated cylinder, the volume was then recorded. After this the solution was poured into an erlenmeyer flask, then the graduated cylinder was rinsed once, with the excess water being added to the flask. This procedure was repeated four more times, cooling the borax to gradually decreasing temperatures of 45.0°C, 35.0°C, 25.0°C, and 15.0°C. For the lower temperatures an ice bath was used to cool down the borax solution. After all five of the borax samples were taken, five drops of methyl-red indicator was added to each flask. Then 200.0 mL of 0.20 M hydrochloric acid was obtained and poured into a clean buret. The samples were then titrated until the borax samples reached a salmon-pink color. The volume of HCl used in each titration was recorded. Then the moles of the tetraborate ion was determined using the molarity of HCl used in titration, then Ksp was calculated for borax at each temperature was recorded. Using the Ksp and temperature data, entropy, enthalpy, and Gibbs free energy for the dissolution of borax at 25.0°C was calculated.

Results Sample

Amount of Borax Solution (mL)

Drops of Methyl Red

HCl Added (mL)

Moles of Tetraborat e Ion (M Borax)

ln(KSP)

Temperature (K)

#1

9.0

5

72.3

0.803

0.729

328

#2

8.7

5

49.1

0.564

-0.3299

318

#3

8.1

5

34.4

0.425

-1.1809

306

#4

8.5

5

15.0

0.176

-3.826

297

#5

8.9

5

9.8

0.11

-5.23

287

Line of Best Fit: ln(KSP) =(ΔH/R)(1/T) + (ΔS/R) ΔH = -1682.64 kJ/mol ΔS = 5.2555 J/Kmol

Discussion The solubility of Borax increases as temperature increases because it takes more acid at higher temperatures to get a reaction with Borax. In the first sample, the temperature was 328 K and it took more acid to cause a reaction than when the temperature was lower at 287 K. The solubility is higher in the higher temperatures because the Borax dissolved easier into the water, causing you to need more acid to react with the Borax. As the temperature decreases, it takes less and less acid to get a reaction with the Borax meaning as temperature increases, solubility increases. Dissolving of Borax into water is endothermic because as temperature increases, solubility increases. This means that heat is a reactant, which means that heat is absorbed. Dissolving Borax into water is endothermic. ΔS would increase as temperature decreases because it would be harder to predict how much of the Borax would dissolve into the water. ΔS would decrease as temperature increases because it would be easier to predict how much of the Borax would dissolve into the water. The dissolution of Borax into water is dependent on temperature because there is a trend in the solubility as temperature changes. As temperature increases, solubility increases. As temperature decreases, solubility decreases. One source of error that we may have had was with the amount of acid added to the Borax solution. It was somewhat difficult to keep track of how much acid was in the buret before we started to drip the acid into the Borax solution. We fixed this during our experiment by making sure that the buret was full at 50 mL so that it was easy to see how much we had put in. Some of our earlier numbers may have been off a little bit but it didn’t seem to affect the outcome of the experiment by too much.

Conclusion In conclusion, the purpose of this experiment was to determine the enthalpy (△H°), entropy (△S°), the value of Gibbs free energy (△G°), and determine the KSP of borax through generating a KSP versus temperature graph. △H° and △S° were -1682.64 kJ/mol and 5.2555 J/Kmol respectively. The enthalpy and entropy were found by the equation: lnK=(ΔH/R)(1/T) + (ΔS/R) as this accounted for the line of best fit on the graph using y=mx+b. It was found that as temperature increased so did the Ksp for Borax. There is an uncertainty on when the borax was supposed to be cooled to 45.0°C, 35.0°C, 25.0°C, and 15.0°C. In the experiment, solutions 2,3,4, and 5 were cooled to 45.0°C, 33.0°C, 23.8°C, and 14.3°C respectively. To expand on this experiment, the experiment should include more trials of other solutions cooled at different temperatures. Also, the experiment should be re-done more than once and the enthalpy and entropy found in both experiments can be averaged out. This will create a more accurate enthalpy and entropy throughout the experiment....