International System of Units

International System of Units

Purpose of the experiment:  The purpose of this experiment is to understand the International System of Units (measurement system), define a unit of measurement and demonstrate the ability to convert measurements.  The lab also aims to define length (meter), temperature (kelvin), time (second), volume (liter), mass (kilogram), density, and concentration, as well as to define significant figures and describe measurement techniques.  The goal is to perform measurements with a graduated cylinder, volumetric flask, graduated pipet, ruler, digital scale, beaker, and thermometer.  Additionally, the objective is to perform, compare, and contrast the water displacement and Archimedes' methods for measuring the volume of an irregularly shaped object.  Finally, the aim will be to calculate concentrations of created solutions, calculate experimental error, and practice basic math and graphing skills.

Procedure for the Experiments:

Exercise 1 Procedures

Part 1: Length Measurements

1. Gather the metric ruler, CD or DVD, key, spoon, and fork.
2. Look at the calibration marks on your ruler to determine the degree of uncertainty and number of significant figures that can be made when measuring objects with the ruler.

Note: Record every measurement you make with this ruler to the same decimal place. Remember to do this any time you use this ruler throughout the experiment.

1. Measure the length of each of the following objects (CD or DVD, Key, Spoon, Fork) with the ruler in centimeters (cm), to one degree of uncertainty and record in Data Table 1.
2. Measure the length of each of the following objects (CD or DVD, Key, Spoon, Fork) with the ruler in millimeters (mm), to one degree of uncertainty, and record in Data Table 1.
3. Convert the measurements for each of the objects from millimeters to meters and record in Data Table 1.

Part 2: Temperature Measurements

1. Gather the 100 mL glass beaker, cup (plastic or drinking), matches or lighter, burner stand, burner fuel, thermometer, 2 oz. aluminum cup, and aluminum pie pan.

Note: The thermometer is shipped in a protective cardboard tube, labeled "thermometer."

1. Look at the calibration marks on the thermometer to determine the degree of uncertainty and number of significant figures that can be made when measuring temperature.

Note: Record every measurement you make with this thermometer to the same decimal place. Remember to do this any time you use this measuring device throughout the experiment.

1. Turn on the tap water to hot. Let the water run as hot as possible for approximately 15 seconds.
2. Fill the 100 mL glass beaker with approximately 75 mL of hot tap water.
3. Measure the temperature of the hot tap water with the thermometer in degrees Celsius (°C) to one degree of uncertainty. Record the measurement in Data Table 2.

Note: When measuring the temperatures place the thermometer into the water so that the silver bulb is fully submerged, but not touching any sides of the glass beaker. The measurement is complete when the thermometer remains the same temperature without changing.

1. Put on safety glasses.
2. Assemble the burner setup and light the fuel, as shown in Figure 14.
   1. Place an aluminum pie plate on a solid work surface away from flammable objects.
   2. Set the burner stand towards the back of the pie plate.
   3. Place the beaker on the center of the stand.
   4. Uncap the burner fuel and set cap aside. Place the burner fuel on the pie plate just in front of the stand.
   5. Use matches or a lighter to ignite the fuel. BE CAREFUL- the flame may be nearly invisible.
   6. Gently slide the fuel under the stand without disturbing the beaker.
   7. The small, 2 oz. aluminum cup will be placed over the fuel to extinguish the flame. Set the aluminum cup next to the burner setup so you are ready to extinguish the flame at any point.

Figure 14.

Burner fuel setup.

1. Allow the water to heat unit it comes to a full boil. As soon as the water is boiling measure the temperature with the thermometer in degrees Celsius (°C), to one degree of uncertainty. Record the measurement in Data Table 2.
2. Allow the water to continue boiling for approximately 5 minutes. After 5 minutes, measure the temperature with the thermometer in degrees Celsius (°C) to one degree of uncertainty. Record the measurement in Data Table 2.
3. Use the small, 2 oz. aluminum cup to extinguish the burner fuel flame. See Figure 15.
   1. Do not touch the metal stand or the beaker; they may be hot.
   2. Carefully slide the burner fuel canister out from underneath the burner stand. The sides of the burner fuel canister will be warm, but not hot.
   3. Place the aluminum cup directly over the flame to smother it. The cup should rest on top of the fuel canister, with little or no smoke escaping. Do not disturb the burner stand and beaker; allow everything to cool completely.
   4. Once all equipment is completely cool, remove the aluminum cup and place the plastic cap back on the fuel. Ensure that the plastic cap “snaps” into place to prevent fuel leakage and evaporation. The aluminum cup, fuel, and all other materials may be used in future experiments.

Figure 15.

Using the aluminum cup to extinguish the flame.

1. Allow the 100 mL beaker to cool before touching it.
2. Turn on the tap water to cold. Let the water run as cold as possible for approximately 15 seconds..
3. Fill the cup (plastic or drinking) approximately half-full with cold tap water.
4. Measure the temperature of the cold tap water with the thermometer in degrees Celsius (°C) to one degree of uncertainty. Record the measurement in Data Table 2.
5. Add a handful of ice cubes to the cup of cold tap water and allow them to sit in the cold water for approximately 1 minute.
6. After 1 minute stir the ice water with the thermometer.
7. Measure the temperature of the ice water after 1 minute with the thermometer in degrees Celsius (°C) to one degree of uncertainty. Record the measurement in Data Table 2.
8. Allow the ice to remain in the water for an additional 4 minutes.
9. After the additional 4 minutes stir the ice water with the thermometer.
10. Measure the temperature of the ice water after 4 minutes with the thermometer in degrees Celsius (°C) to one degree of uncertainty. Record the measurement in Data Table 2.
11. Convert the temperature measurements for each of the 6 water samples from °C to °F and K. Record the converted temperatures in Data Table 2.

Part 3: Mass Measurements

1. Gather the pen or pencil, 5 pennies, 3 quarters, 4 dimes, and the key.
2. Read the instructions on how to use the digital scale. The lid of the scale must be opened to expose its weighing surface and make mass measurements.

Note: There may be a cardboard protector between the scale base and top. If so, remove the cardboard from the scale.

1. Turn the scale on by pressing the Φ/T button.
2. Make sure the scale is reading in grams by looking for the letter "g" in the upper right corner of the scale, if the "g" is not showing then press the "M" button until the scale is reading in grams.
3. Review the different object(s) listed in Data Table 3.
4. Estimate the masses for each of the object(s) in grams and record in Data Table 3. To help you with this process, a penny has a mass of approximately 2.5 grams.
5. Tare the scale by pressing the Φ/T button so that the scale reads 0.0 g.
6. Place the pen or pencil on the scale to measure the mass of the object. Record the mass in Data Table 3 under “Actual Mass (g).”
7. Repeat steps 33 and 34 for the remaining object(s) in Data Table 3.
8. For each object(s), convert the actual mass (in grams) to kilograms (kg). Record in Data Table 3.

Exercise 2 Volume and Density

In this exercise, you will make measurements using the SI system units for length, mass, and temperature.

Procedure

Part 1: Volume and Density Measurements (Liquid)

1. Gather the graduated cylinder, distilled water, short stem pipet, and isopropyl alcohol.
2. Place the clean, dry, 25 mL graduated cylinder on the tared scale. Record the mass of the graduated cylinder (g), in Data Table 4.
3. Fill the graduated cylinder with 5.0 mL of distilled water; use the short stem pipet to help measure exactly 5.0 mL of water. Record the volume in Data Table 4.
4. Place the 25 mL graduated cylinder with 5.0 mL distilled water on the tared scale. Record the mass of the graduated cylinder and the liquid in Data Table 4 .
5. Calculate the mass of the water by subtracting “Mass B” from “Mass A.” Record the mass of the water in Data Table 4.
6. Pour the water down the drain and fully dry the graduated cylinder.
7. Repeat steps 2 through 6 for the isopropyl alcohol.
8. Calculate the densities of both the water and the isopropyl alcohol and record in Data Table 4 .
9. The accepted value for the density of water is 1.00 g/mL and the accepted density for isopropyl alcohol is 0.786 g/mL. Determine the percent error between your calculated densities and the accepted values for both water and isopropyl alcohol. Record the percent error in Data Table 4.

Part 2: Volume and Density Measurements (Solid)

1. Gather the metal bolt, string, magnet, graduated cylinder, beaker, metric ruler, and scale.

Direct Measurement Method:

1. Tare the scale by pressing the Φ/T button so that the scale reads 0.0 g.
2. Place the magnet on the scale to measure the mass of the object. Record the mass in Data Table 5.
3. Use the ruler to measure the length, width, and height of the magnet in centimeters. Record the measurements in Data Table 5 .
4. Calculate the volume of the magnet by multiplying the length × width × height, record in Data Table 5 .
5. Calculate the density of the magnet by dividing the mass by the volume and record in Data Table 5.

Water Displacement Method:

1. Tare the scale by pressing the Φ/T button so that the scale reads 0.0 g.
2. Place the magnet on the scale to measure the mass of the object. Record the mass in Data Table 6.
3. Fill the graduated cylinder with 6 - 8 mL of distilled water. Record the volume, to the correct decimal place, in Data Table 6.
4. Carefully slide the magnet into the graduated cylinder so that the water doesn't splash, and read the volume of the graduated cylinder. Record the volume in Data Table 6, next to "Final volume of graduated cylinder."
5. Determine the volume of the object by calculating the difference in water displacement volumes (final – initial). Record in Data Table 6.
6. Calculate the density of the magnet and record in Data Table 6.
7. Carefully pour the water from the cylinder down the drain and collect the magnet.
8. Repeat steps 16 through 22 for the metal bolt.

Archimedes’ Method:

1. Tare the scale by pressing the Φ/T button so that the scale reads 0.0 g.
2. Place the metal bolt on the scale to measure the mass of the object. Record the mass, to the correct decimal place, in Data Table 7.
3. Attach the end of the string to the metal bolt, by tying the string to the bolt. See Figure 16.

Figure 16.

String attached to metal bolt.

1. Fill the glass beaker approximately ¾ full with distilled water and place the beaker on the scale.
2. Tare the scale by pressing the Φ/T button so that the scale reads 0.0 g. Refer to Figure 11 for questions.
3. Holding onto the string, submerge the bolt into the water so that the bolt is fully submerged, but not touching any part of the glass beaker. Record the mass reading from the scale and record in Data Table 7 under "Mass of Displaced Water."
4. Convert the mass of displaced water to volume of displaced water, assuming the density of water is 1.00 g/mL. Record in Data Table 7.
5. Calculate the density of the metal bolt and record in Data Table 7.
6. Repeat steps 24 through 31 for the magnet.

Exercise 3 Concentration, Solution, and Dilution

In this exercise, you will use laboratory equipment to create solutions of varying concentrations and densities by diluting a stock solution.

Note: View the following video for a demonstration of using the digital scale before you continue the procedures

Using The Digital Scale

Procedure

1. Gather the volumetric flask, distilled water, graduated pipet, pipet bulb, sugar, scale, glass beaker, cup (plastic or drinking), scissors, and a sheet of white paper.
2. Tare the digital scale by pressing the Φ/T button so that the scale reads 0.0 g.
3. Remove the stopper for the volumetric flask from the top of the flask.
4. Place the dry volumetric flask on the scale to determine its mass. Record the mass of the volumetric flask in Data Table 8.

Note: You will use the mass of the dry volumetric flask throughout the experiment.

1. Cut a small square of white paper, fold it in half and then in quarters, and place it on the scale.
2. Tare the scale by pressing the Φ/T button so that the scale reads 0.0 g with the paper on it. See Figure 17.

Figure 17.

Taring paper on a digital scale.

1. Place 8.0 grams of sugar on the paper. Record the mass in Data Table 8.
2. Carefully transfer the sugar from the white paper into the volumetric flask.
3. Add approximately 15 mL of distilled water to the volumetric flask containing the 8.0 g of sugar, and swirl the flask until the sugar is completely dissolved in the water.
4. Once the sugar is dissolved, use the short stem pipet to add additional water to the 25 mL mark. See Figure 18.

Figure 18.

Adding water to a volumetric flask.

1. Record the total volume in Data Table 8.
2. Determine the concentration of the sugar solution in the volumetric flask using the following equation:

% m/ in Volumetric Flask (g)Total Volume (mL)×100%% m/ in Volumetric Flask (g)Total Volume (mL)×100%

1. Record the concentration of the sugar solution in three places: In Data Table 8, for “Concentration,” in Data Table 9 for “Final Concentration” in the “Solution 0” row, and in Data Table 9 under “Initial Concentration” in the “Solution 1” row.
2. Tare the scale by pressing the Φ/T button so that the scale reads 0.0 g.
3. Place the volumetric flask containing the solution on the scale.
4. Subtract the mass of the dry, empty volumetric flask from the mass of the flask containing the solution. The difference is the mass of the sugar solution.
5. Record the mass of the sugar solution in Data Table 9 for “Solution 0.”
6. Calculate the density of the sugar solution and record it in Data Table 9 for “Solution 0.”
7. Pour the sugar solution (Solution 0) into the dry, empty 100 mL glass beaker.
8. Rinse the volumetric flask well with distilled water.
9. Use the serological (graduated) pipet to transfer 2.5 mL of Solution 0 into the volumetric flask.
10. Add approximately 15 mL of distilled water to the volumetric flask containing Solution 0 and swirl to thoroughly mix.
11. Using the short stemmed pipet, continue adding distilled water to the 25 mL mark of the flask.
12. Tare the scale by pressing the Φ/T button so that the scale reads 0.0 g and then place the volumetric flask containing the solution on the scale.
13. Subtract the mass of the dry, empty volumetric flask from the mass of the flask containing the solution.
14. Record the mass of the new sugar solution in Data Table 9 for “Solution 1.”
15. Calculate the density of the sugar solution and record in Data Table 9 for “Solution 1.”
16. Determine the Final Concentration (% m/V) for Solution 1 using the following equation:

C1××V2In this equation: Concentration, Transferred, mL, ConcentrationC1××V2In this equation: Concentration, Transferred, mL, Concentration

1. Record the Final Concentration of “Solution 1” in Data Table 9 under “Final Concentration.”
2. Repeat steps 21 through 29 for the remaining 3 dilutions (4.5 mL, 3 mL, and 6 mL) in Data Table 9.

Note: In step 21, transfer the appropriate volume of sugar solutions (4.5 mL, 3 mL, and 6 mL).

1. Create a graph displaying the relationship between Concentration and Density for the sugar solution. The x-axis on the chart will be Density and the y-axis is Concentration. Upload an image of the graph into Graph 1.

Cleanup:

• Clean all glassware and lab equipment with soap and water and rinse the equipment again with distilled water.
• Dry all items with paper towels and return to the lab kit for future use.

Experimental data: (Table or Graph)

Experiment 1

Data Table 1: Length Measurements

Experiment 1

Data Table 2: Temperature Measurements

Data Table 3: Mass Measurements

Experiment 2

Data Table 4: Liquid Measurements

Experiment 2

Data Table 5: Direct Measurement Method

Experiment 2

Data Table 6: Water Displacement Method

Experiment 2

Data Table 7: Archimedes' Method

Experiment 3

Data Table 8: Initial Concentration

Experiment 3

Data Table 9: Solutions

Experiment 3

Graph 1: Concentration versus Density

Observations and Conclusions:

Answers to the questions in the lab manual associated with the experiment:

Exercise 1 Questions

• Water boils at 100°C at sea level. If the water in this experiment did not boil at 100°C, what could be the reason?

The boiling point of water is achieved when the atmospheric pressure equals the vapor pressure of water, hence water boils at 100°C at sea level. As the pressure is decreased, however, the vapor pressure of water approaches the pressure of the surrounding air, that is the surrounding air that exists here in Loiza, the boiling temperature of water decreases with decreases in pressure.

When the air pressure is reduced enough, water begins to boil at room temperature and therefore water starts boiling as the water molecules at the surface gain the energy from the surrounding water molecules and escapes out, meaning, the water vapor takes heat of vaporization from water. Ergo, the temperature of water drops slightly.

2. While heating two different samples of water at sea level, one boils at 102°C and one boils at 99.2°C. Calculate the percent error for each sample from the theoretical 100.0°C.

100(|102-100|/100)= 100(2/100)= 2% error (for first sample that boils at 102°C)

100(|99.2-100|/100)= 100(0.8/100)(0.008)= 0.8% error (for second sample that boils at 99.2°C)

Exercise 2 Questions

• An unknown, rectangular substance measures 3.6 cm high, 4.21 cm long, and 1.17 cm wide. If the mass is 21.3 g, what is this substance’s density (in grams per milliliter)?

(3.6cm*4.21cm*1.17cm) = 17.7 cm^3
D = M/V = 21.3g/17.7cm^3 = Density = 1.20 g/cm^3 = 1.2 g/ml

• A sample of gold (Au) has a mass of 26.15 g. Given that the theoretical density is 19.30 g/mL, what is the volume of the gold sample?

26.15g Au

Theoretical density 19.30 g/ml

Therefore 26.15/19.30 = 1.35 ml

• What would happen if you dropped the object into the beaker while using the Archimedes’ Principle method instead of submerging the object?

Besides the potential splash, there may be no difference in measure or effect, except to say that this poor handling of the object could be a source of error. Still, depending on the object, beaker and liquid it could be the same.  That is to say, there should be no difference. Archimedes' Principle describes the net effect of gravity and pressure on the submerged object, and doesn't really matter precisely how the object came to be submerged.

4. How did the magnet’s density measurement using the Archimedes’ Principle compare to the density measurement using the calculated volume? Which method might be more accurate? Why?

Density measurement using the Archimedes' Principle might be less accurate if the magnet does not fully submerge into the water, thus causing the displaced volume to be slightly lesser than the actual calculated volume of the magnet. Still, using Archimedes' Principle has its advantage in that we can find the volume of irregularly shaped magnets or objects, that we cannot analytically calculate the volume for.

• You are given a small piece of gold colored material and want to determine if it is actually gold. Using the Archimedes Principle you find that the volume is 0.40 cm3 and the mass is 6.0 g. What conclusions can you reach from your simple density analysis?

density = mass / volume

= 6.0g / 0.40cm^3

density = 15g/cm^3

However, the actual density of gold is 19.30g/cm^3

We can conclude that the sample is not pure gold. It probably contains some combination of lead, silver or copper.

Exercise 3 - Questions

1. How would you prepare 10 mL of a 0.25% m/V HCl solution if 1% m/V HCl was available? How much 1% m/V HCl is needed? How much distilled water is used?

We use M1V1 = M2V2 (M= concentration or molarity, and V= volume)

Where M1 = 0.25 M

V1 = 10 mL

M2 = 1.0 M

So you need to solve for V2.

(0.25 M)(10 mL) = (1.0 M) V2

= 2.5 mL of 1.0 M HCl

So, we will need 2.5 mL of 1.0 M HCl to make a dilution of 0.25 M HCl and dilute to 10 mL, which means our solution will consist of 2.5 mL HCl and 7.5 mL of DI water.