Why Does The Ammonium Chloride Increase The Amount Of Solute?
The mass of urea was normalized so that it could become the equivalent of what it would’ve been in the larger solvent as numbered in Data Table 4. There wasn’t much difference between the corrected number of moles and grams of urea and the amount of urea used originally used in the experiment, so the ammonium chloride will be primarily compared to the corrected amount of urea to make sure that the amount of solvent is equal. Graph 1 and 5 and Data Table 2 show that the mass of solutes were overall very similar: 8.735 grams of ammonium chloride and 8.99 grams of urea. Overall, this experiment proved that the freezing point depression is a colligative property that depends on the number of moles of solute but not on the identity of the solute.
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The strong R2 values in graph 2, 4 and 6 show that a strong positive correlation exists between moles of solute and change in freezing temperature regardless of the solute identity.
As displayed on Graph 2 and Graph 6 and typed in Data Table 3 and 4, the moles of solute for ammonium chloride and the corrected moles of solute for urea were very close: .1627 moles of ammonium chloride and .1492 moles of urea. Even though the amount of moles was very similar, there is clearly a stark difference in the change in temperature when comparing Graph 2 and 6: When all of the moles of each solute were added to each beaker, the final change in temperature for ammonium chloride was nearly double that of the final change in temperature for urea. As shown in Data Table 3 and 4, the final change in temperature for ammonium chloride was 4.74 °C, and the final change in temperature for urea was 2.09 °C. To cause these results, ammonium chloride, which is a salt of a weak base, more than likely dissociated into ammonium ions and chlorine ions, resulting in twice as many more particles in solution than urea, which probably doesn’t dissociate much or at all, in the water/ice
solution. With twice as many solute particles in the ammonium chloride solution, the change in temperature should be double as well. As shown by Data Table 3 and 4 and Graphs 2 and 6, this is indeed the case. Clearly, the freezing point depression is a colligative property that depends only on the number of particles of solute and not on the identity of the solute. One possible source of error could come from the fact that some of the urea tended to stick to the inside of the test tube, which could’ve caused the change in temperature reading to be lower than it really should’ve been. Another slight source of error was that the temperature reading given by MeasureNet temperature probe would consistently vary by .01°C; this resulted in inaccurate readings that were higher or lower than what the true reading should be; this fluctuation could occur due to heat constantly coming into the beaker (high to low) to melt the ice and raise water temperature. Another possible source of error could’ve occurred because the solute didn’t fully dissolve into the solvent; this could result in a change in temperature reading being lower than it should be. Overall, the results that were found in this experiment make sense with what is already known about the colligative properties of the freezing point depression.
The strong R2 values in graph 2, 4 and 6 show that a strong positive correlation exists between moles of solute and change in freezing temperature regardless of the solute identity.
As displayed on Graph 2 and Graph 6 and typed in Data Table 3 and 4, the moles of solute for ammonium chloride and the corrected moles of solute for urea were very close: .1627 moles of ammonium chloride and .1492 moles of urea. Even though the amount of moles was very similar, there is clearly a stark difference in the change in temperature when comparing Graph 2 and 6: When all of the moles of each solute were added to each beaker, the final change in temperature for ammonium chloride was nearly double that of the final change in temperature for urea. As shown in Data Table 3 and 4, the final change in temperature for ammonium chloride was 4.74 °C, and the final change in temperature for urea was 2.09 °C. To cause these results, ammonium chloride, which is a salt of a weak base, more than likely dissociated into ammonium ions and chlorine ions, resulting in twice as many more particles in solution than urea, which probably doesn’t dissociate much or at all, in the water/ice
solution. With twice as many solute particles in the ammonium chloride solution, the change in temperature should be double as well. As shown by Data Table 3 and 4 and Graphs 2 and 6, this is indeed the case. Clearly, the freezing point depression is a colligative property that depends only on the number of particles of solute and not on the identity of the solute. One possible source of error could come from the fact that some of the urea tended to stick to the inside of the test tube, which could’ve caused the change in temperature reading to be lower than it really should’ve been. Another slight source of error was that the temperature reading given by MeasureNet temperature probe would consistently vary by .01°C; this resulted in inaccurate readings that were higher or lower than what the true reading should be; this fluctuation could occur due to heat constantly coming into the beaker (high to low) to melt the ice and raise water temperature. Another possible source of error could’ve occurred because the solute didn’t fully dissolve into the solvent; this could result in a change in temperature reading being lower than it should be. Overall, the results that were found in this experiment make sense with what is already known about the colligative properties of the freezing point depression.