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Molar Mass Determination by Freezing Point Depression

Formal Report

Lab #7

Molar Mass Determination by Freezing Point Depression

Introduction

Molar masses or formula masses are important quantities. Like atomic masses, these values are the foundation for the majority of numerical work when it comes to chemistry. The molar mass of a compound can be determined by different varieties of methods. It would merely depend on the nature of the compound. For example, when the sample is in the form of a liquid, it’s commonly treated as an ideal gas. The freezing point of a solution and the pure solvent is inclined by the quantity of moles of solute currently present; thus it correlates to molar mass (Russell, 2015).

In this experiment, the freezing point depression of a solution will be measured and used to conclude the molar mass of an unknown substance. Lauric acid, C12H24O2, will act as a solvent. The freezing point of lauric acid in the presence of a solute is decreased by a quantity determined by the following equation, ΔT = i * Kf * m. In this equation, ΔT is the change in temperature; the freezing point of the solution of the unknown in lauric acid is subtracted from the freezing point of lauric acid. A digital temperature probe will be used to explicitly monitor the temperature of a sample. The numeral of ions formed per formula unit of the compound are represented by “i”. In this particular experiment, “i” is equal to 1 (Russell, 2015). Kis the molal freezing point constant for the solvent. Since the solvent is lauric acid, the constant will be 3.9 °C/molal. The molality, m, is number of moles of solute divided by the mass of solvent in kilograms. This experiment comprises of placing the sample in a flask, melting, and then letting it cool. Supercooling occurs when a sample does not solidify or crystallize below the normal freezing point of 0 °C. This phenomenon was discovered by Daniel Gabriel Fahrenheit. He developed the design of the thermometer as well (Tusa, 2008).

A real life application of freezing point depression is the making of homemade ice cream. The ice cream mix is placed into a metal container surrounded by ice. Salt is then added to the ice to decrease its melting point. The melting of the solution has a tendency to lower the equilibrium temperature of the ice/water solution to the melting point of the solution. Temperature rises across the metal container into the saltwater-ice solution, which is lesser than 0 °C. The heat transfers out of the ice cream mix allowing it to freeze (Cook and Hartel, 2010).

Experimental

After the LoggerPro program started, a 500 mL water bath was set up in an 800 mL beaker and then heated to above 60 °C. A 125 mL Erlenmeyer flask’s weight was recorded. An approximation of 9 grams of lauric acid was added to the flask. The flask and lauric acid were weighed together and the mass was recorded. The flask was placed in the water bath and heated until the lauric acid melted entirely. Once it melted, the flask was removed from the water bath. The sample was cooled to about 60 °C before data collection. A temperature probe was placed in the acid to collect data. When the temperature remained constant for about two minutes, data collection stopped.

About 0.9 grams of the unknown sample was weighed and recorded to the nearest 0.001 grams. The unknown was added to the previous sample of the lauric acid. The flask was heated again until the contents melted. As previously completed, the freezing point data was collected with the temperature probe. Then the graph of both sets of concluding data was expanded and printed off.

Data and Results

Table 1

Contents

Masses

Dry 125 mL Erlenmeyer flask

92.5 grams

Flask and lauric acid

101.53 grams

Lauric acid

9.03 grams

Unknown #4

0.900 grams

Table 2

Contents

Temperature/Values

Freezing point of pure lauric acid

43.33 °C

Freezing point of lauric acid-unknown solution

44.69 °C

ΔT (decrease/change in freezing point)

1.36 °C

Kf (molal freezing point constant)

3.9 °C/molal

m (molality)

0.35 mol/kg

Number of moles of unknown

0.0032 mol

MM (molar mass: considering no significant figures)

281.25 g/mol

MM (molar mass: considering significant figures)

300 g/mol

Percent error

54.5%

Discussion

The first set of data recorded looked different from the original example shown in the lab manual. The data was collected before the temperature reached 60 °C, which may have been a mistake, skewing the results. Not measuring the solutions correctly could have added to different results as well. For example, the amount of lauric acid weighed out to be 9.03 grams instead of 9 grams (Table 1). After a few more attempts, the experiment finally worked as expected and the slope of data on the graph also appeared more precise; likewise, the original graph example.  Next time, it would be best to wait until the solution reached the required temperature instead of rushing the lab. Knowing that the experimental molar mass was actually 281.25 grams/mole but was about 300 grams/mole when considering significant figures (Table 2). The percent error, according to calculations, turned out to be 54.5% (Table 2).

Calculations

Calculations

Results

ΔT (decrease/change in freezing point):

= (44.69 °C) – (43.33 °C)

1.36 °C

m (molality):

= (1.36 °C) / (3.9 °C/molal)

0.35 molal (or mol/kg)

Number of moles of unknown:

= (0.35 mol/kg) * (0.00903 kg)

0.0032 mol

MM (molar mass: considering significant figures):

= (0.900 g) / (0.0032 mol)

300 g/mol

Percent error:

= (281.25 g/mol – 182 g/mol) / (182 g/mol)

= (99.25 g/mol) / (182 g/mol)

54.5%

Conclusion

A digital temperature probe was used to check the temperature of an unknown sample, lauric acid, and both compounds combined. The molar mass of the unknown substance was calculated by the data of the freezing points of the solutions. Due to the molal freezing point constant for the solvent, only one significant figure should be used to express the final result. If each molecule of unknown had dissociated into two ions in the lauric acid solution, the freezing point would be lower compared to an equal molality solution that had a solute that did not dissociate. The percent error actually turned out to be a surprising number as well. The results perceived from this experiment were only 54.5% accurate.

Citations

Cook, K. L. K.; Hartel, R. W. Mechanisms of Ice Crystallization in Ice Cream Production. Comprehensive Reviews in Food Science and Food Safety, http://onlinelibrary.wiley.com/doi/10.1111/j.1541-4337.2009.00101.x/full  (accessed Oct 1, 2016).

Tusa, S. Fahrenheit, Gabriel Daniel. - Enlightenment Revolution, http://enlightenment-revolution.org/index.php/fahrenheit,_gabriel_daniel  (accessed Oct 2, 2016).

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