USING THE PIPET
The purpose of this experiment is to calibrate a volumetric pipet and use it to measure the density of a liquid. A pipet is used to transfer or measure a liquid. When properly used, it is as accurate as a buret. A volumetric pipet is designed to deliver one set volume of liquid (called an aliquot), and has just one mark for that volume. Mohr pipets have graduated scales on the sides.
Procedure
1. Determine the uncertainty of the balances. Obtain a solid object and weigh it on each balance.
2. Now practice using the pipet. Place a rubber suction bulb onto the top of the pipet. The bulb does
NOT fit over the glass like a medicine dropper bulb; a soft conical fitting seals to the pipet when held against the glass. DO NOT force the bulb over the glass, because you will damage the fitting.
Use the bulb to draw water up into the pipet until the liquid is well above the mark. Try not to draw liquid into the bulb (keep the pipet tip in liquid when drawing it up); consult the instructor if you do draw liquid up into the bulb. NEVER (never, never) draw liquid into a pipet with your mouth. Never.
Remove the bulb and immediately seal the pipet top with your index finger, so the liquid won’t run out. Don’t use your thumb; it offers less control and looks dorky. Gently wipe liquid off the outside

of the pipet with a tissue. Use finger pressure to adjust the liquid level until the meniscus is at the mark. Release it and let it drain into the desired container. Do not blow out liquid in the pipet.
3. When you have practiced enough to feel comfortable with the pipet, calibrate it using this technique:
Obtain some deionized water in a clean beaker. Label 5 clean dry cups and record the mass of each cup. Record the temperature of the water, then pipet one aliquot of the water into each cup. Weigh the cups + water.
4. Obtain some of the unknown liquid. Rinse your pipet with the unknown liquid before doing any measurements, so the liquid will not be diluted with water left in the pipet: draw a small amount of the liquid into the pipet, roll it around to coat the inside of the pipet completely, then expel the liquid into the sink or waste beaker. Repeat two more times. After the pipet has been rinsed with the unknown liquid, use 5 new cups and the same procedure as above to measure the mass of several aliquots of the liquid.
Analysis
1. Calibration of the balance.
Calculate the average (mean) mass of your object and the average deviation from the mean.
For example, student X obtains 35.618 g, 35.615 g, and 35.619 g for her solid object.
Mean = (35.618 + 35.615 + 35.619) ÷ 3 = 35.617 g
Deviation from mean = absolute difference between measured value and mean:
| 35.618 – 35.617 | = 0.001
| 35.615 – 35.617 | = 0.002
| 35.619 – 35.617 | = 0.002 average deviation =
(0.001 + 0.002 + 0.002) ÷ 3 = 0.00167
The object weighs 35.617 g ± 0.002
2. Calibration of your pipet.
Calculate the mass of each aliquot of water, and convert that to volume of water delivered using the density of water at the temperature of your experiment (look up water density at various temperatures in the Handbook of Chemistry and Physics).
Calculate the average (mean) volume delivered and the average deviation from the mean.
For example, student X obtains 10.10 mL, 9.98 mL, and 10.00 mL for her 10 mL pipet.
Mean = (10.10 + 9.98 + 10.00) ÷ 3 = 10.03 mL
Deviation from mean = absolute difference between measured value and mean:
| 10.10 – 10.03 | = 0.07
| 9.98 – 10.03 | = 0.05
| 10.00 – 10.03 | = 0.03 average deviation = (0.07 + 0.05 + 0.03) ÷ 3 = 0.050
Pipet delivers 10.03 mL ± 0.05
The average deviation from the mean gives you an idea how precise your pipetting is. Be sure you round off appropriately: remember to show the beginning of uncertainty, but no more.
3. Density of an unknown liquid.
Calculate the average density of the liquid, using the average aliquot mass and the average aliquot volume delivered by your pipet, as calculated in your calibration. Round off your answer appropriately and be sure you include the correct units.
4. Calculate the % uncertainty in your density. The overall % uncertainty is the sum of the relative
(percent) uncertainties of the measurements. For example, student X obtains an average mass of
7.973 g ± 0.002 and an average volume of 10.03 mL ± 0.05 for her unknown liquid.

The density of her liquid is d

7.973
g
10.03
mL

0.794915254
g mL

0.7949
g mL
The percent uncertainty of each measurement is

0.002
g
7.973
g

100

0.0251%
0.05
mL
10.03
mL

100

0.499%
The overall uncertainty of her density is 0.02251% + 0.499% = 0.524%. The absolute uncertainty of
0.794915254
g
 mL

0.00524

0.004162
g mL
This result indicates that she should round her final answer to 0.795 g/mL ± 0.004 (3 sigfigs), even
When she compares her answer to the accepted value, a difference of 0.5% is within the uncertainty of her measuring instruments. Any difference beyond that is due to procedural errors.
Discussion
1. Obtain the accepted value of the density for that liquid and calculate your percent error. Discuss how accurate your result was, and if your answer was not accurate (within the uncertainty of the procedure), discuss reasons for the error.
2. Sigfig rules predict that student X should keep 4 sigfigs in her answer, but an analysis of the uncertainty shows that only 3 sigfigs are justified. Why didn’t the sigfig rules work in this example?
3. Examine your absolute and relative (percent) uncertainties. When designing an experiment, why is it best to use the largest sample that is practical for the procedure and the available instruments?