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CHM-201 General Chemistry and Laboratory I Laboratory 1 ­ Introduction to Experimental Measurement


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CHM-201 General Chemistry and Laboratory I Laboratory 1 ­ Introduction to Experimental Measurement Purpose This laboratory introduces you to basic laboratory measurement techniques with a focus on the proper use of significant figures. You will become familiar with measuring mass, volume and length. The proper recording of scientific raw data using non-eraseable ink to record data directly on the data sheets provided is expected in every laboratory session. Calculations will be performed which require the correct use of significant figures. Introduction It is extremely important when performing laboratory work to be able to measure quantities properly. The accuracy (closeness of a measurement to its true value) and precision (variability of identical measurements) of a measurement are determined by the instrument(s) used to obtain this data as well as the skill and training of the operator ­ you. All instruments have an inherent limit to the accuracy and precision of a measurement. Put another way, an instrument will normally provide data with a set number of significant figures. When recording data in the laboratory, you should ask yourself "How many of the numbers I am recording are meaningful?" Sometimes you don't know the answer to this question until you perform the same measurement more than once (called replication) and look at the variability of the data. In certain cases, the instrument will have a limit to the number of significant figures it will display. In general, if an instrument displays a value with several digits, we assume that all are reliable, but the last digit is uncertain. If we are unable to reproduce results to the same level of accuracy and precision, we may reassess the accuracy of the instrument. In all cases replication of the measurement is the only true way to assess the precision of the measurement. This is why laboratory data is often taken more than once. When evaluating significant figures, the first digit that has some uncertainty is reported and all further digits are discarded. Graphical Analysis of Data Graphical analysis is a powerful method for presenting scientific data. There are seveal methods for representing data but one of the most common is an x-y plot of the data. You will generate computer or calculated generated plots of data in later experiments in CHM-201 and 202. Before going to this level of sophisication, it is important to understand how to generate and interperet a graph manually. Read the section in the Laboratory Handbook (pages 83-85) carefully. In addition keep the following points in mind when you generate the graph of mass vs. volume in this experiment (and for any futher linear plots in later laboratories). 1. Be sure your graph is large enough. The graph paper provided is 8.5 x 11 inches and any graph should cover more than half of the page. If this is not the case you have not selected appropriate values on the axes for the range of data you have obtained. Use values such that each line is a simple integer in order to make the data easier to plot and read. Note that it is not neccessary to start all graphs at x = 0, y = 0. 2. Draw the best straight line possible through the data. There should be one line (not a series of connected line segments ­ do not connect each data point in a dot to dot fashion) which goes through the data you plot. Since some data points will reside above and below the line use your best approximation for your line in which the positive deviations from the line equal the negative

deviations from the line (for purely random error you would expect to find an equal number of data points above and below the line). 3. Draw your best fit line only through the data range for which you have experimental values. You do not know what happens to the data for x values smaller than what you measured or for x values greater than what you measured. If you project the line you obtained into regions for which you have no x,y data this is called extrapolation. 4. Slope determinations should be performed from x,y pairs which originate from the line (NOT from individual data points). You have graphed the data in order to reduce the experimental error and the x,y values should be obtained from reading new values from the line you generated. Exercises Note: Record all of your laboratory data in non-eraseable ink on the sheets provided. 1 Determining the conversion factor between centimeters and inches (10 points) 1.1 Using a 12 inch ruler, determine the number of centimeters that correspond to line segments of 1/4, 1/2, 1, 5 1/2, and 11 inches. Measure the value for each length in inches (remember that you are interpreting a scale here ­ what are the finest markings on the scale? how accurately can your measure in inches using this ruler?) and in centimeters (ask the same questions about the cm scale) and record these values using the correct number of significant figures. Calculate the conversion factor from inches to centimeters using this data. Be careful to follow the rules for significant figures. Do not use the known value of the conversion factor 2.54 cm = 1.000 inch to convert between inches and centimeters since that is the quantity you should be calculating in this exercise.


Inches 1/4 1/2 1 5 1/2 11

Decimal Value (in inches)

Centimeters (measured)

Calculated Conversion Factor (cm/in)

2. Determining the volume of a balloon. (10 points) 2.1 Blow up a balloon and measure its circumference three times. Compute the volume of the balloon using these measurements. If you don't recall the formula for the volume of perfect sphere, look it up. Be sure to use significant figures and units correctly. Provide a statement which indicates how close the balloon shape is to a perfect sphere. How, if at all, does this effect your calculated volume? Trial 1 Circumference Diameter Radius Calculated Volume





Statement regarding spherical shape of balloon:

General: You will use a microcentrifuge tube (commonly called an Eppendorf tube) for these measurements although a small test tube can be substituted. Since isopropanol is a volatile liquid (readily evaporates to become a gas) you should keep the top of the tube closed after each transfer to avoid loss of the isopropanol. You will perfom the transfers with an Eppendorf automatic transfer pipette. The instructor will explain how to use this piece of equipment but these are relatively expensive pipettes so please be careful and be sure to use a disposable tip on the end of the pipette. The instructor may provide you with an alternative method of transferring the required amounts such as using a glass 0.5 or 1.0 mL graduated pipette. 3. Determining the densities of liquids (isopropanol and salt water solution) (50 points) 3.1 Obtain the mass of an empty microcentrifuge tube on a balance that can read to 1 mg (0.001 g) precision. If the balance you are using does not show three places to the right of the decimal point contact the instructor as to how to reset the balance. Transfer approximately 1.0 mL of isopropanol (20 drops or as measured by the markings on the tube) into a second microcentrifuge tube (not the one used in the previous step). Transfer 100 L of isopropanol (isopropyl alcohol or rubbing alcohol) using the Eppendorf pipette as instructed below: 3.3.1 Be sure that the window at the top of the pipette reads 100 and that a plasic pipette tip is firmly attached to the bottom of the pipette. If the pipette does not read 100 (for 100 microliters) please contact the instructor. Depress the plunger of the pipette until you reach a `resistance point' (not all the way down) and then place the tip into the solution. Release the plunger of the pipette to the fully released position (liquid should be drawn into the plastic tip) which transfers 100.0 L of isopropanol into the tip. Place the tip of the pipette over the microcentrifuge tube which you placed on the balance and depress the plunger of the pipette all the way down and cap the tube. This should transfer the desired volume of isopropanol into the microcentrifuge tube. Once you have completed all the transfers for a given solution push the lever on the top of the pipette to discard the tip and dispose of the tip in the trash.



3.3.2 3.3.3 3.3.4



Record the mass of the microcentrifuge tube containing the isopropanol to a precision of 0.001 grams (1 mg). Without removing the isopropanol from the microcentrifuge tube, transfer an additional 100 L of isopropanol to the microcentrifuge tube. Again seal the tube and record the mass of the microcentrifuge tube containing both transferred volumes of isopropanol. Repeat 3.5 once to transfer a third 100 L aliquot of isopropanol and again record the mass after the transfer. You should have three masses of the microcentrifuge tube containing 100, 200 and 300 L of isopropanol. Repeat this procedure (3.1-3.6) using salt solution in place of the isopropanol. Since salt water is not as volatile as isopropanol it is much less important to seal the cap of the microcentrifuge tube. As above your should have a tare mass (mass of the empty tube, as well as three masses of the tube containing 100, 200 and 300 L of salt water solution.





There should be a graduated cylinder at the instructor's desk which contains the salt water solution and a hydrometer (a weighted and sealed glass tube with a calibration scale in the narrow upright tube which rests above the liquid level). Record the specific gravity for this salt water solution by observing where the scale intersects with the level of the solution. This should provide a reading with at least 4 places to the right of the decimal point. Determine the density of each liquid measured above by graphing the mass of the liquid transferred on the y-axis (Net weight = Gross weight - Tare weight) by the volume transferred on the x-axis. You can plot both sets of data on the same piece of graph paper. Calculate the average density for isopropanol and salt solution by determining the slope of each line on the graph (mark which data is ethanol and which is salt solution). You will need to submit a copy of this graphical analysis performed manually. For this exercise you cannot use graphing software (cricket graph, excel, graphical analysis, etc.) nor simply perform a linear regression analysis on a scientific calculator.


3.10 The specific gravity as measured by the hydrometer is a density relative to water. Since the density of water at room temperature is 0.998 g/mL we can assume that the specific gravity and density can be compared directly. Calculate the percent difference between the value you obtained using the hydrometer and the value you obtained for the slope of the graph for salt water solution. The percent difference can be calculated by dividing the difference between the values by the average of the values and mutiplying by 100%. Do not be concerned if this is a negative value. Determining the densities of liquids (isopropanol and salt water solution) Isopropanol 2 Salt Water Solution 2 3

Trial Tare Mass

(same for all trials)




Gross Mass Net Mass Volume Density (from graph) Specific Gravity Percent Difference

Graphs of Net Mass versus Volume for Isopropanol and Salt Water Solution.

Questions 1. In exercise #1, what conclusions can you come to regarding the accuracy in determining the conversion factor? Include a statement about the significant figures in this exercise (5 points).

2. Convert the volume of the balloon calculated in part #2 into cubic feet (ft3). Is there any way to correct for the non-spherical nature of the balloon? Is so how would you do so and what would be the expected result on your measured volume. Do you think that the number of significant figures that you report for your volume is appropriate? Comment on how many significant figures you think this corrected volume should have and if that is the same number of significant figures that you used in your calculation. (10 points)

3. In your graphical analysis which variable (mass or volume) is the dependent variable? Why do you think this is the case when you examine the experiment performed? Can you describe a different experiment which would make the other quantity the dependent variable? If so describe what you would measure in this experiment and what you would then graph. (10 points)

4. If your laboratory partner were to leave the cap off the microcentrifuge tube during the isopropanol density determination then what error would you predict for your density value? Would you expect a density determined from the slope of the graphical data to be larger or smaller than the actual density? Explain your reasoning. (5 points)


Introduction to Experimental Measurement

10 pages

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