Creating Solutions of Standard Molarity - Lab Report PDF

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This is the first lab report required for the first general chemistry lab at UNCC. The title gives the description perfectly. I made a perfect grade on this assignment and it is a rich source of information for chemistry students at the beginner-level....

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Creating Solutions of Standard Molarity Joshua Farley Partner: Keegan Bradley CHEM 1251L-923 October 8, 2014 Introduction In this experiment, the main principles of spectrometry were to be studied by creating solutions of a specified molarity and measuring them using the spectrophotometer, or Spec 20, to discover their levels absorbance and transmittance by measuring the amount of light passing through the solution. Because all solutions absorb certain wavelengths of light, while transmitting other complementary colors through them, they tend to take on the colors that are not absorbed by the sample for us to see and observe. The higher the concentration of the solute in the solvent, the darker the color will appear. Absorbance can be calculated using the equation A=-log(T) where T represents the transmittance recorded as a decimal value. Using a correct calibration curve, the actual concentration, or molarity, of the solution could be identified. Molarity is defined as the number of moles of solute per one liter of solvent, which was water in this experiment. Additionally, the serial dilution technique was utilized when making new solutions of a specific molarity. Using a standard solution and the following equation, MC x VC = MD x VD, a conversion between the molarity and volumes of the original solution and those of the new, diluted solution could be made to determine how much solute and solvent must be added to

new solution to give it a certain molarity. Therefore, the more concentrated a solution is with a particular solute, the more wavelengths of light it will absorb. Procedure Initially, ten small test tubes were thoroughly cleaned, dried, wiped, and label with a piece of tape. Test tube #1 was filled about ¾ full with distilled water to be used as a control on the spec 20 instrument. The Spec 20 was to be warmed up for approximately 15 minutes before use and was then calibrated to 0% transmittance and maximum absorbance for Cu(II) solutions, which is 600 nm. After wiping the control tube with a Kim-wipe to remove fingerprints, the tube was placed in the well of the Spec 20 and transmittance was set to 100% to ensure that the instrument was correctly calibrated and ready for use. With the Spec 20 calibrated, the standard solutions were to be poured into their respective test tubes, labeled 0.50 M, 0.20 M, 0.10 M, and 0.05 M CuSO4. These test tubes were also wiped with Kim-wipes because fingerprints can interfere with the detection of wavelengths of light, which can disrupt the ultimate results. The spectrophotometer was then utilized to find and record the percent transmittance, as well as the absorbance. Because of the design of the Spec 20, the absorbance values could have been read slightly inaccurately compared to the transmittance values. Therefore, the actual absorbance of each standard solution of CuSO4 was calculated using the equation, A=-log(T), where the percent transmittance was converted into a decimal value. Using the absorbance and the concentration (or molarity) of each standard solution, a calibration curve was created and included a line of best fit, an equation that can be used to find the absorbance of a solution with a certain concentration (or vice-versa), and an R2 value, which is a percentage that determines how close the data is fitted to the regression line.

After recording all of the data from the previous steps and constructing a calibration curve, new Copper(II) solutions were to be prepared from the solute, Copper(II) sulfate pentahydrate, and the solvent, distilled water. In order to make a 20.0 mL solution with a molarity of 0.50 M, the moles of solute needed were determined using the formula for molarity. Since 0.50 M = moles of solute/.02 L of water, 0.50 M was multiplied by .02 to result in .01 moles of CuSO4, which when multiplied by the molar mass of Copper(II) sulfate pentahydrate (249.96 g/mol), results in about 2.5 grams of the solute. A 100 mL beaker was tapered on a scale and then 2.5 grams of the solid were added. This was then added to a graduated cylinder contained about 15.0 mL of water and mixed until all of the solute had completely dissolved in the cylinder. Afterwards, water was slowly added to the mixture with a Pasteur pipette until the meniscus read exactly 20.0 mL. Some of the mixture was then poured into one of the test tubes and labeled as 0.50 M CuSO4; this was the first made solution. In order to create a solution of 20.0 mL at 0.20 M CuSO4, the serial dilution technique was to be used. Since 0.50 M x V1 must equal 0.20 M x 20.0 mL, simple algebra was performed to find that the volume that needed to be removed from the first solution was 8.0 mL. Therefore, 8.0 mL of the solution of 0.50 M was extracted and placed into the graduated cylinder. About 12.0 mL of water was then added using a Pasteur pipette until the meniscus read 20.0 mL again. This became the second solution and was placed in a test tube labeled 0.20 M CuSO4. Using the serial dilution technique again, it must be true that 0.20 M x V1 = 0.10 M x 20.0 mL. The equation was solved and the volume that needed to be removed from the 0.20 M solution was 10.0 mL. 10.0 mL of the solution was removed and placed into the graduated cylinder where 10.0 mL of water were then carefully added to ultimately create a new solution that was added to the third test tube and labeled as 0.10 M CuSO4. Finally, the final solution was created by finding

the volume that must be removed from the 0.10 M solution by solving the equation: 0.10 M x V1 = 0.05 M x 20.0 mL. The volume that was to be extracted came out as 10.0 mL. Therefore, 10.0 mL of the 0.10 M solution were removed and placed into the graduated cylinder where another 10.0 mL of water would be slowly added until the reading was at 20.0 mL. This resulted in a new, diluted solution that was labeled 0.05 M CuSO4. Now that the final solutions had all been prepared, they were taken to the Spec 20 where their percent transmittance and absorbance values could be viewed and recorded. The calibration curve that was constructed for the standard solutions was now being used to determine the actual concentrations of the made solutions. The equation of the line of best fit, y = 0.8033x + 0.0193, where y is the absorbance and x is the concentration, had the absorbance values plugged into the y-value so that the actual concentrations could be found. A final, unknown Copper(II) solution was given so that it could be placed in the Spec 20 in order to find the percent transmittance and the absorbance which were the utilized to determine the concentration of the solution. Results Table 1: Shows the recorded data from the 4 standard solutions. Concentration of

Percent

Transmittance

Absorbance

Standard Solutionss 0.05 M 0.10 M 0.20 M 0.50 M

Transmittance 87.5% 79.7% 65.6% 38.0%

0.875 0.797 0.656 0.380

0.058 0.099 0.183 0.420

Figure 1: Calibration curve of the standard Copper(II) solutions.

Absorbance

Absorbtion of Standard Copper(I I) Solutions 0.45 0.4 0.35 0.3 0.25 0.2 0.15 0.1 0.05 0

f(x) = 0.8 x + 0.02 R² = 1

0

0.1

0.2

0.3

0.4

0.5

Concentration (mol/L)

Table 2: Shows the data from the 4 made solutions. Concentration of

Percent

Absorbance

Actual Concentration

Made Solutions 0.05 M 0.10 M 0.20 M 0.50 M

Transmittance 86.0% 76.8% 66.0% 37.0%

0.067 0.118 0.175 0.434

0.06 M 0.12 M 0.19 M 0.52 M

Actual concentrations of the made solutions were found by plugging the absorbance into the equation for the calibration curve as the y-value; y = 0.8033x + 0.0193. Table 3: Shows the data from the unknown solution. Solution

Percent

Absorbance

Concentration

Unknown

Transmittance 54.2%

0.265

0.31 M

Table 4: Compares concentrations of standard and made solutions. Concentration of Standard

Concentration of Made

Percent Error

0.6

Solution 0.05 M 0.10 M 0.20 M 0.50 M

Solution 0.06 M 0.12 M 0.19 M 0.52 M

20% 20% 5% 4%

Percent error was calculated using the following equation:

|measured value−actual value| actual value

x 100 % .

Table 5: Compares absorbance values of standard and made solutions. Concentration

Absorbance of

Absorbance of Made

Percent Error

0.05 M 0.10 M 0.20 M 0.50 M

Standard Solutions 0.058 0.099 0.183 0.420

Solutions 0.067 0.118 0.175 0.434

15.52% 19.19% 4.37% 3.33%

Discussion Table 1 illustrates the data recorded from the original four standard solutions of a definite concentration. The percent transmittance, transmittance, and absorbance were all found using the Spec 20 with the solutions. However, the absorbance values observed on the instrument were not very precise. Because of this, further calculations were done using the equation A=-log(T) to find the actual absorbance of the solution. This calculated value is the one present on all of the tables above. Finding the absorbance is crucial to this experiment because it directly relates to concentration according to the Beer-Lambert Law. Therefore, it establishes a relationship between the two subjects so that any other absorbance value plugged into the equation of the linear trendline in Figure 1 could be used to find the actual concentrations of the made solutions.

As seen on all of the tables, as the concentration of solute in a solution increased, the transmittance decreased while the absorbance increased. This shows a direct correlation between all three values. The R2 value in Figure 1 is a percentage that describes how close the data is fitted to the line of best fit. In this calibration curve, the R2 value was 0.9998, or 99.98%. This simply shows that the concentrations shown on the graph are nearly perfectly aligned with how the data should be represented in a linear form. It also serves as a percent error value for the data points in comparison to the linear trendline. The calibration curve is relevant because, as seen in Table 2, the concentrations of the made solutions are not exactly what they were expected to be. Using the equation from the graph, y = 0.8033x + 0.0193, the actual concentrations could be found by using the calculated absorbance values. Table 3 simply shows the recorded data of the unknown solution of CuSO4 with the calculated absorbance, as well as the calculated concentration of 0.31 M. It was already known that the concentration would be between 0.20 M and 0.50 M because the color of the unknown sample was darker than the 0.20 M solution and lighter than the 0.50 M solution. Table 4 shows the relationship between the concentrations of standard solutions and the concentrations of the made solutions. Additionally, it shows the percent error of the results. The concentrations of the made solutions were mostly a bit higher than those of the standard solutions. However, as the concentrations increased in both the standard and made solutions, the percent error decreased dramatically. Therefore, the techniques and instruments used in the experiment show a higher level of accuracy as the molarity of the solutions increase. There are also several other factors that could play into a high percent error. One could have put too much or too little solute in the container and not mixed it up well enough. There may have been an excess of water in the mixture which would result in a more diluted solution

than expected. The Spec 20 could have been misread since its absorbance values are typically slightly inaccurate. There may have been too many fingerprints on the test tube which would disrupt the detection of wavelengths of light or some excess water molecules may have still been stuck to the side, resulting in a less concentrated solution. All of these factors, plus many more, could affect the outcome of the made solutions, leading to a higher percent error. Similar to Table 4, Table 5 compares the absorbance of the standard solutions with the absorbance of the made solutions, along with the calculated percent error. This table shows how as the concentration of a solute increases, so does the amount of wavelengths of light it absorbs compared to the wavelengths of light transmitted through it. As the absorbance of the solutions increased, the solution turned into a darker shade of blue due to the complementary colors being transmitted throughout the solution. The made solutions in this experiment were different from the standard solutions in that the standard solutions had a definite molarity, or concentration, while the made solutions had to be created by those in the lab. Therefore, several of the factors that can result in a higher percent area, like those listed above, ultimately caused the made solutions to have a slightly different molarity. This gave off more inaccurate data than what was expected for the experiment. Water was used to dilute samples of the original solution into new solutions that were to have a certain molarity based on measurements of water put into the new sample. A large amount of the percent error resulted from inaccurate measurements of water mixed into the solution, which would result in a lower or higher concentration than what was expected. Conclusion

This experiment showed how different solutions of a specific molarity could be created using the serial dilution technique and the formula for finding molarity. These conversions were crucial in determining the amount of solute and solvent to put in a mixture in order to create a solution of a certain concentration. The Spec 20 device showed the amount of wavelengths of light transmitted through and absorbed by the solution so that the actual concentration could be determined. Because the colors of light absorbed by the solution are contained within in, the complementary, or opposite, colors are the observable ones that humans see. The higher the concentration of solute in a solution, the darker those colors become. The amount of light absorbed by a sample is related to the concentration of that sample because as the molarity of a substance in a solution increases, more wavelengths of light are absorbed by the sample. Therefore, there are more complementary colors transmitting through the sample, giving it a darker color and hence, a higher concentration and absorbance. Based on the values seen in Table 4 and Table 5, it’s obvious that the techniques used for this experiment are not completely accurate when working with solutions of a lower concentration. A digital spectrophotometer would probably yield more accurate results, which would lead to more precise data. Despite this, there are several other common errors that could have led to this inaccuracy, such as amount of solute/solvent put in a solution, the possibility of contamination in a solution, fingerprints on the test tubes, faulty calibration on the Spec 20, and not allowing the crystals to dissolve completely in the solution. All of these errors could result in a higher percent error by giving off misreadings on the meter.

Sample Calculations

1) Converting percent transmittance into a decimal value:

Percent Transmittance 100 %

87.5% / 100% = 0.875

2) Finding absorbance using the transmittance values: Absorbance=-log(Transmittance) A = -log(0.875) A = 0.058

3) Find moles of solute needed in a 20.0 mL solution: Molarity =

0.50 M = moles of solute / 0.2 L 0.50 M x 0.2 L = moles of solute Moles of solute = 0.1 mol

4) Convert moles to grams: moles x molar mass (g/mol) 0.1 mol x 249.69 g/mol = x grams x = 2.5 grams

5) Serial Dilution Technique: M1V1 = M2V2 0.50 M x V1 = 0.20 M x 20.0 mL

molesof solute liters of solution

0.50 M x V1 = 4 moles 4 moles / 0.50 M = V1 V1 = 8 mL

6) Finding absorbance using percent transmittance: A = 2 – log(%T) A = 2 – log(86.0%) A = 0.066

7) Finding actual concentration of made solutions: y = 0.8033x + 0.0193 Y is absorbance, X is concentration 0.067 = 0.8033x + 0.0193 0.8033x = 0.0987 X = 0.06 moles

8) Finding percent error:

|measured value−actual value|

|0.06 M – 0.05 M| / 0.05 M |0.01 M| / 0.05 M = .20 x 100% = 20%

actual value

x 100 %...