Labreport Solubility

Topics: Thermodynamics, Gibbs free energy, Entropy Pages: 7 (2285 words) Published: April 30, 2013
Solubility and Thermodynamics

The purpose of the lab was to determine the thermodynamics variables of ∆H, ∆S, and ∆G for the dissolution reaction of potassium nitrate in water. The solubility of potassium nitrate in mol/L was measured over a range of various temperatures by finding out at what temperature crystallization began for solutions of different molarities. Then, the equilibrium constant was calculated and a graphical relationship between the natural logarithm of the equilibrium constant and the inverse of the temperature gave a linear plot that allowed the determination of Gibbs free energy and enthalpy changes associated with the reaction. From there, the entropy change associated with the reaction was determined.

1. Weigh out 20.0g of potassium nitrate and transfer it to a 25x200mL test tube. Do not ingest the potassium nitrate 2. Add 15 mL of water and heat the test tube in a boiling water bath, stirring until all of the potassium nitrate dissolved 3. Determine and record the volume of the potassium nitrate solution. Use a second large 25x200mm test tube and fill it with an equal volume of water. Then, measure that volume in a graduated cylinder 4. Remove the test tube with the potassium nitrate from the water bath and allow it to cool while slowly stirring the solution with a thermometer 5. Record the temperature when crystals first appear. It is assumed that at this temperature, the system is at equilibrium and it is possible to calculate the concentration of the ions 6. Add 5mL of water to the test tube and warm mixture in a hot water bath until all solid dissolves. Determine the solution volume as before and record it 7. Cool the solution slowly and record the temperature at which crystals first appear 8. Repeat the cycle of adding water/heating/cooling/recording temperature until crystals appear near room temperature. Add 5mL of water each time

1. 20.00g of potassium nitrate taken, white solid spheres (fairly large sized spheres) 2. Initially, add 10.01mL + 4.98mL of water to potassium nitrate 3. First time, heated to 89.9 degrees Celsius

4. Test tube feels cool to touch as potassium nitrate dissolves 5. Crystals form rapidly after some point when temperature drops, easy to see

|Trial |Volume |Temperature | |1 |23.9mL |67.5˚C | |2 |29.9mL |55.2˚C | |3 |35.0mL |46.3˚C | |4 |38.9mL |40.6˚C | |5 |43.9mL |36.2˚C | |6 |48.5mL |32.5˚C |

Moles of KNO3
20.00g * (1mol/(39.10g + 14.01g + 16.00g * 3)) = 0.1978mol

Kelvin Temperatures:
67.5˚C + 273.2 = 340.7K; 1/T = 2.935 * 10-3K-1
55.2˚C + 273.2 = 328.4K; 1/T = 3.045 * 10-3K-1
46.3˚C + 273.2 = 319.5K; 1/T = 3.130 * 10-3K-1
40.6˚C + 273.2 = 313.8K; 1/T = 3.187 * 10-3K-1
36.2˚C + 273.2 = 309.4K; 1/T = 3.232 * 10-3K-1
32.5˚C + 273.2 = 305.7K; 1/T = 3.271 * 10-3K-1

Concentration at Equilibrium:
At 340.7K: [KNO3] = 0.1978mol / 23.9mL * (1000mL/1L) = 8.28M At 328.4K: [KNO3] = 0.1978mol / 29.9mL * (1000mL/1L) = 6.62M At 319.5K: [KNO3] = 0.1978mol / 35.0mL * (1000mL/1L) = 5.65M At 313.8K: [KNO3] = 0.1978mol / 38.9mL * (1000mL/1L) = 5.08M At 309.4K: [KNO3] = 0.1978mol / 43.9mL * (1000mL/1L) = 4.51M At 305.7K: [KNO3] = 0.1978mol / 48.5mL * (1000mL/1L) = 4.08M

Equilibrium Constant:
K = [NO3-][K+] and [NO3-] = [K+] = concentration of KNO3 at equilibrium. K = [KNO3]2 At 340.7K: K = 8.282 = 68.6; lnK = 4.23
At 328.4K: K = 6.622 = 43.8; lnK = 3.78
At 319.5K: K = 5.652 = 31.9; lnK = 3.46
At 313.8K: K = 5.082 =...
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