Lab 1 - Laboratory Techniques and Measurements (Chemistry)

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Laboratory Techniques and Measurements Hands-On Labs, Inc. Version 42-0165-00-02

Review the safety materials and wear goggles when working with chemicals. Read the entire exercise before you begin. Take time to organize the materials you will need and set aside a safe work space in which to complete the exercise.

Experiment Summary:

You will use laboratory equipment to make measurements using the SI system units for length, mass, and temperature. You will prepare dilutions, and calculate the concentrations of the dilutions using the C1V1=C2V2 equation. You will use laboratory equipment to create solutions of varying concentrations and densities by diluting a stock solution.

EXPERIMENT

Learning Objectives Upon completion of this laboratory, you will be able to:

● Define the International System of Units (measurement system).

● Define a unit of measurement, and demonstrate the ability to convert measurements.

● Define length (meter), temperature (kelvin), time (second), volume (liter), mass (kilogram), density, and concentration.

● Define significant figures and describe measurement techniques.

● Perform measurements with a graduated cylinder, volumetric flask, graduated pipet, ruler, digital scale, beaker, and thermometer.

● Perform, compare, and contrast the water displacement and Archimedes methods for measuring the volume of an irregularly shaped object.

● Calculate concentrations of created solutions.

● Calculate experimental error.

● Practice basic math and graphing skills.

Time Allocation: 3.5 hours

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Materials Student Supplied Materials

Quantity Item Description 1 Aluminum pie pan 1 Bottle of distilled water 1 CD or DVD (to measure) 1 Cup, plastic or drinking 4 Dimes 1 Fork (to measure) 1 Ice cubes 1 Isopropyl (rubbing) alcohol, 10mL 1 Key (to measure) 1 Matches or lighter 1 Sheet of white paper 1 Pair of scissors 1 Pen or pencil 5 Pennies 3 Quarters 1 Spoon (to measure) 1 Sugar, 1 tablespoon 1 Source of tap water

HOL Supplied Materials

Quantity Item Description 1 Aluminum cup, 2 oz 1 Burner fuel 1 Burner stand 1 Digital scale 1 Glass beaker, 100 mL 1 Graduated cylinder, 25 mL 1 Magnet bar 1 Metal bolt 1 Pair of gloves 1 Pair of safety goggles 1 Ruler 1 Rubber bulb 1 Serological pipet, 2 mL 1 Short, thin-stem pipet 1 String, 1 m 1 Thermometer 1 Volumetric flask, 25 mL

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Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software, such as Microsoft Word® or PowerPoint®, to add labels, leader lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.

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Background International System of Units (SI)

Chemistry is the science of matter and its changes. In order to study matter and its changes, scientists make qualitative and quantitative observations. Quantitative observations, or measurements, always consist of a numerical value and a unit of measurement. Scientists use the International System of Units (SI), which is derived from the metric system, as the standard system of measurement. Table 1 shows the basic SI units to measure the five fundamental properties of length, mass, time, temperature, and amount of substance. The sixth property shown, volume, is considered a derived unit, as it is the cubic (3D) version of length. As shown in Table 1, the SI system and the metric system are closely related, differing only in the size of the fundamental unit. Table 1 also includes the units more likely to be used in the laboratory, where very small quantities are used.

Table 1. Comparison of units used to measure fundamental properties.

Measurement International System (SI) Metric System Common Laboratory Units

Length meter (m) meter (m) centimeter (cm) Mass kilogram (kg) gram (g) gram (g) Time second (s) second (s) second (s)

Temperature Kelvin (K) Celsius (oC) Celsius (oC)

Amount of Substance

mole (mol) mole (mol) mmole (mmol)

Volume cubic meter (m3) liter (L) milliliter (mL)

US customary units require conversion SI to metric units for scientific reporting. Table 2 illustrates the relationships (conversion factors) between US customary units and metric/SI units.

Table 2. Metric-US conversions.

Property Factor Factor Length 1 in = 2.54 cm 1 mi = 1.609 km Mass 1 lb = 454 g 1 kg = 2.206 lb

Volume 1 qt = 0.946 L 1 L = 1.06 qt Temperature oF = (oC x 1.8) + 32 K = oC + 273

Table 3 lists the meanings of the prefixes used in the metric and SI systems. Each prefix is a multiplication factor for the base unit. For example, the prefix kilo means 1000, so 1 kilogram is equivalent to 1000 grams.

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Table 3. Prefixes used in the SI and metric systems.

Prefix Symbol Meaning Exponential Notation mega M 1,000,000 106

kilo k 1,000 103

hecto h 100 102

deka da 10 101

deci d 0.1 10-1

centi c 0.01 10-2

milli m 0.001 10-3

micro µ 0.000001 10-6

Length

Length is defined as the distance (amount of space) of an object from end to end. The SI system unit of length is the meter (m) and was originally intended to represent one ten-millionth of the distance between the North Pole and the Equator. However, over time the definition of the meter changed. The current definition, which has been in place in 1983, is that a meter equals the distance that light travels in a vacuum in 1/299,792,458 seconds. There are many ways to measure length, including a caliper, a calibrated ruler, a tape measure, and even a laser. See Figure 1. Each of these different measuring devices measures length to a different degree of accuracy and precision.

Figure 1. Common equipment used to measure length. A. Vernier caliper © Paul Paladin B. Tape measure © Jiri Hera C. Calibrated ruler © Quang Ho

A ruler is a tool used to measure length. However, a

ruler (rule) is actually defined as an instrument used to rule (create) straight lines. A calibrated ruler is a ruler which contains measurements to measure length along a straight

line.

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Temperature and Time

Temperature is defined as a measure of the ability of a substance to transfer heat energy to another. More commonly, temperature is referred to as a measure of warmth or coldness of an object in reference to a standard value. The SI system unit of temperature is the kelvin (K), and the standard value (0 K) is defined as 1/273.16 of the temperature when water exists as a solid, liquid, and gas; at 1 atmosphere of pressure. This point, in which water co-exists as a solid, a liquid, and a gas is called the triple point. While the SI unit for temperature is the kelvin, the majority of thermometers are calibrated to degrees Celsius (°C or C) and/or degrees Fahrenheit (°F or F). Today, scientific measurements are commonly taken in the Celsius scale. A thermometer is used to measure temperature. Converting between Fahrenheit, Celsius, and Kelvin scales is common and can be performed with one of the following conversion formulas:

A comparison of common temperatures between the three different scales (K, °C, °F) is shown in Figure 2.

Figure 2. Comparison of common temperatures on the three different scales.

The second (s) is the basic SI unit of time and is defined as the duration of 9,192,631,770 cycles of radiation in an energy level change of the cesium atom. While smaller quantities of time than the second are described using the standard SI prefixes (Table 1), quantities of time larger than a second (minute, hour, day, week) are not SI units. Time is measured using watches (clocks), which are calibrated to atomic clocks. Atomic clocks are extremely precise clocks, which are regulated by the vibrations (resonance frequencies) of cesium atoms.

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Volume

Volume is defined as the amount of space occupied by a three-dimensional object or area of space. The SI unit for volume is the cubic meter (m3), which is equal to the volume of a cube measuring 1 meter on each side (1m x 1m x 1m). See Figure 3. The units for volume most commonly used in science laboratories include the liter (L) and the milliliter (mL). A liter is equal to the volume of a cube, with sides of 0.1 m (0.1m x 0.1m x 0.1m = 0.001m3 = 1L). A milliliter is 1/1000 of a liter and it is also equal to a cubic centimeter 1 mL = 1 cm3.

Figure 3. Volume of a cube. A cube that is 1 meter in length on all sides has a volume of 1 m3 or 1000 L.

1L = 1000mL = 1000cm3

In addition to the milliliter, the microliter ( ), which is one-millionth of a liter, is also a common unit of volume used by scientists. When measuring the volume of a liquid, common laboratory equipment options include a graduated cylinder, volumetric flask, and graduated pipet.

Graduated Cylinder, Meniscus, Volumetric Flask, and Graduated Pipet

A graduated cylinder is a slender container that is calibrated by specific volumetric amounts, such as milliliters or liters, and can measure a range of volumes depending on its capacity. See Figure 4.

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Figure 4. Graduated cylinder. A 50-mL graduated cylinder made of glass. Note the markings along the cylinder, which represent milliliters.

When reading a graduated cylinder made of glass, or any measuring device for volume that is made of glass, it is important to read the volume, at eye level, from the bottom of the meniscus. A meniscus is the curve that forms between the liquid and the surface of the container as the result of surface tension, cohesion, and adhesion. See Figure 5.

Figure 5. Meniscuses. A. Colored water in a graduated cylinder made of plastic (left). Note that the liquid forms a straight, non-curved, line at the 50-mL mark. Colored water in a graduated cylinder made of glass (right). Note that the liquid forms a curved line. B. When reading the

volume from a meniscus, the volume is read from the bottom of the curve.

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The majority of liquids in glass will form a concave shaped meniscus. However, some liquids,

such as mercury, will form a convex meniscus. In the case of a convex

meniscus, the volume is read from the top of the curve rather than

the bottom.

While graduated cylinders are designed to measure a variety of volumes, a volumetric flask is calibrated to measure only one volume, and is often used to prepare a specific volume of solu- tion. A volumetric flask has a bulb-shaped bottom and a very long slender neck. Each flask is calibrated with a mark on that slender neck, allowing for very careful and accurate measure- ments. See Figure 6.

Figure 6. Volumetric flask. A 25-mL glass volumetric flask. The line on the flask, as noted by the black arrow, marks 25-mL calibrated measurement.

For smaller volumes, scientists use graduated pipets (also referred to as serological pipets) to measure liquids. Just as for graduated cylinders, graduated pipets can be used to measure a range of volumes depending on the capacity of the pipet. Similar to volumetric flasks, the calibrated measurement area of graduated pipets is very narrow, allowing for very precise measurements. Graduated pipets use a suction mechanism to fill and release liquid. See Figure 7.

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Figure 7. Graduated pipet. Shown is a 2-mL graduated pipet, calibrated to 0.1 mL. The red bulb on the right end of the pipet is used to fill and release the liquid from the pipet.

The order of markings on a graduated pipet is opposite from the order of markings on a graduated cylinder. When a 50-mL graduated cylinder contains 50 mL of liquid, the liquid is at the 50-mL marking. However, when using a 2-mL graduated pipet, when liquid reaches the marking labeled “0”, it means that the pipet contains 2-mL of liquid. There are many different types of graduated pipets, and each has their own mechanism for dispensing liquid. For the graduated pipet that will be used in this experiment; to dispense the full 2 mL from the pipet the liquid is completely released from the pipet. Likewise, to dispense only 1 mL from the pipet the liquid is released from the “0” marking to the “1.0” marking. See Figure 8.

Figure 8. Measuring with a graduated pipet. Notice the “0” and “1.0” markings, both noted with black arrows.

Just as with graduated cylinders, graduated pipets are available in a variety of sizes, allowing for measurements as small as 0.1 mL.

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Measuring Volume of a Solid

There are multiple techniques to measure the volume of a solid, three of which are discussed here. For a solid with defined edges, such as a cube, a box, or a sphere, the length, width, and height or the radius of the object can be measured and used to calculate volume. See Figure 9.

Figure 9. Calculating the volume of solid objects with defined edges. © Image of sphere adapted from cobalt88.

When measuring the volume of an irregularly shaped object there are two common methods, the water displacement method and Archimedes’ method. In the water displacement method, the irregularly shaped object is placed into a known amount of water in a graduated cylinder. The increase in water volume, as measured by the graduated cylinder, is equal to the volume of the irregularly shaped object. See Figure 10.

Figure 10. A. Irregularly shaped object. B. Graduated cylinder with known amount of water. C. Graduated cylinder with irregularly shaped object. The difference (increase) in volume from 30 mL to

32.5 mL is equal to the volume of the irregularly shaped object.

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Archimedes’ method incorporates buoyancy into the water displacement method. Buoyancy is the upward force that a fluid places onto an object, which is equal to the weight of the displaced fluid. Put simply, when an object is submerged and suspended in water (1 mL of water = 1 gram of water), the change in mass, as the result of the object being submerged is equal to the volume of the object. See Figure 11.

Figure 11. Archimedes’ method. A. Water in a beaker on a tared scale. B. Irregularly shaped object is submerged in the water, creating a buoyant force equal to 2.5 grams. As 1 mL of water has a mass of 1 gram, the volume of displaced water is 2.5 grams, which is equal to the volume of the irregular shaped

object.

Archimedes, a Greek mathematician and scientist, was

asked by King Hiero II to determine if a crown made for him was composed of pure

gold, or contained a mixture of gold and silver. As Archimedes could not damage the crown, he had to find a way to determine its composition. It is from this challenge that he developed what is now known as Archimedes’ Principle. Knowing

that density is a physical property that does not change, Archimedes could determine if the density of the crown was equal to the density of

pure gold. As the densities of the crown and pure gold did not match, he concluded that the crown was not made from

pure gold!

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Mass and Density

Mass is defined as the measure of the amount of matter contained in a physical body. The SI unit of mass is the kilogram (kg) and is standardized (equal) to the mass of the International Prototype Kilogram (IPK). The IPK is composed of a platinum-iridium alloy, and is stored at the International Bureau of Weights and Measurements in Sevres, France. See Figure 12.

Figure 12. International Prototype Kilogram. The IPK is kept under two bell jars. Image courtesy of the U.S. Federal Government.

It is important to note that terms “mass” and “weight” are often used interchangeably, but are not the same thing. Mass is a quantifiable measure of matter, while weight refers to the gravitational force of attraction exerted upon an object. The SI system uses mass measurements, yet the verb “weigh” is the common verb used to describe or obtain the mass of an object. In the laboratory, scientists usually work with the gram (g) which represents one-thousandth of a kilogram, and the milligram (mg) which equals one-thousandth of a gram. Mass is measured using an instrument called a balance, which is commonly referred to as a scale.

Density is defined as mass per unit of measure, which is most often volume. It is a way to describe how heavy something is for its size. Like volume, density is not a fundamental SI unit of measure, but is derived from the SI units for mass and volume. The density of a liquid is usually reported as grams per milliliter (g/mL), while the density of a solid is usually reported as grams per cubic centimeter (g/cm3). The density of water is 1 g/mL. Substances with a density greater than 1 g/ mL will sink when placed in water, while objects with a density less than 1 g/mL will float when placed in water. The density of an object is determined with the following equation:

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Solution Concentration

In the laboratory, many substances are used in solution form, usually dissolved in water. There are many ways to describe the amount of a dissolved or concentrated substance. This laboratory introduces only one of those, the percent of mass/volume concentration, % m/V. The definition of % m/V is:

For example you would prepare a 5% solution of table salt, NaCl, in water by dissolving 5 grams of NaCl in enough water to make 100 mL of solution.

Chemists often make solutions more concentrated than the ones they use every day. Concentrated solutions are more stable than dilute ones, and it is easier to make accurate measurements of the masses needed to make concentrated solutions. Each day, the concentrated stock solutions can be diluted the desired concentration using the following relationship:

In this equation, C1 = the initial concentration of the solution, V1 = the initial volume of the solution, C2 = the final concentration of the solution, and V2 = the final volume of the solution.

For example, assume that you have 20 mL of a 5.6% m/V solution and want to make 40 mL of a 3.7% m/V solution. In this example: C1 = 5.6% m/V, C2 = 3.7% m/V, V2 = 40 mL, and V1 is the volume of a 5.6% m/V solution, which is needed to create 40 mL of 3.7% m/V solution. V1 is calculated with the equation as follows:

This means that to make the dilution solution, 26.43 mL of the 5.6% m/V solution is measured and then placed in a graduated cylinder or 40-mL volumetric flask and filled with water to a total volume of 40 mL.

Significant Figures and Percent Error

Significant figures are a combination of the certain and first uncertain digit of a measurement. The number of significant figures in a measurement is dependent upon the accuracy of the instrument. Measurements may be estimated for one decimal place, allowing for one uncertain digit. For example, if a ruler is calibrated to the millimeter length, a measurement may be taken to a tenth of a millimeter. Likewise if a ruler is calibrated to the centimeter length, a measurement may only be taken to the millimeter. While some may try to estimate the last number as 0, the number is only a guess and invites uncertainty in the value. See Figure 13.

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Figure 13. Significant figures in measurements.

The two rulers depicted in Figure 13 can be used to conduct measurements with different significant figures. The topmost, wooden ruler is calibrated to 0.5 cm. If a piece of string was measured, and it’s length was between 2 cm and 2.5 cm, the total length might be estimated as 2.3 cm. in this case, the tens place (2.3) is one significant digit, and the tenths place (2.3) is the estimated, uncertain digit. The bottom-most, metal ruler is calibrated to 0.1 cm (or 1 mm). The length of the same string could be reported as 2.25 cm. In this case, there are two significant figures: one in the tens place and one in the tenths place (2.25). The estimated, uncertain digit is in the hundredths place (2.25). In general, measurements are reported one decimal place beyond the instrument’s calibration.

Each time a particular measuring device is used, the scientist records the measurement to the same decimal place. That is, if a scientist is using a centimeter ruler similar to the bottom one in Figure 13, every measurement will end with a value in the hundredths column. The length of an item that is exactly 5 cm long must be recorded as 5.00 cm using that ruler. The zeros in the tenths and hundredths places indicate the accuracy of the measuring device. Remember, the last digit recorded is the uncertain one. If a student records a length of 5 cm, he/she is telling the instructor that the 5 is uncertain, and the actual length is somewhere between 4 and 6 centimeters. When a student records a length of 5.0 cm, the 0 is uncertain or estimated and the true length is somewhere between 4.9 and 5.1 centimeters. A measurement of 5.00 centimeters indicates that the 0 in the hundredths column is uncertain and the true length is between 4.99 and 5.01 centimeters.

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Calculations should be consistent with the values of measurements. The value of the least significant digit sets the standard for the degree of accuracy of measurements. When multiplying, the product cannot be greater than the number of significant figures in the factor having the fewest significant figures:

2.36 g x 3.1 g = 7.316 g2

Recorded as 7.3 g2

The product should be recorded as 7.3 g2 since there are only two significant figures in the least precise factor, 3.1 g.

When adding, the sum of the numbers cannot be more precise, have more decimal places, than the least precise number:

22.55 g + 75.3 g + 12 g = 109.85 g

Recorded as 110 g

The sum should be recorded as 110 g since the least precise number was precise to the full gram, 12 g.

Often measured results will differ from theoretical expectations, and there will be variations within repeated measurements. Percent error (the difference between the accepted/true value and the value that was measured) may be calculated using the following formula:

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Exercise 1: Length, Temperature, and Mass In this experiment, you will make measurements using the SI system units for length, mass, and temperature.

Procedure

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.

3. 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 of your Lab Report Assistant.

4. 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.

5. Convert the measurements for each of the objects from millimeters to meters and record in Data Table 1.

Temperature Measurements

6. 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.”

7. 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.

8. Turn on the tap water to hot. Let the water run as hot as possible for approximately 15 seconds.

9. Fill the 100 mL glass beaker with approximately 75 mL of hot tap water.

10. 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 of your Lab Report Assistant.

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Note: When measuring the temperature, place the thermometer into the water so that the silver bulb is fully submerged without touching any sides of the glass beaker. The measurement is complete when the thermometer remains the same temperature without changing.

11. Put on safety glasses.

12. Assemble the burner setup and light the fuel, as shown in Figure 14.

• Place an aluminum pie plate on a solid work surface away from flammable objects.

• Set the burner stand towards the back of the pie plate.

• Place the beaker on the center of the stand.

• Uncap the burner fuel and set cap aside. Place the burner fuel on the pie plate just in front of the stand.

• Use matches or a lighter to ignite the fuel. BE CAREFUL- the flame may be nearly invisible.

• Gently slide the fuel under the stand without disturbing the beaker.

• 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.

Note: When the burner is lit, the flame may be barely visible.

Figure 14. Burner fuel setup.

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13. 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.

14. 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.

15. Use the small, 2 oz. aluminum cup to extinguish the burner fuel flame. See Figure 15.

• Do not touch the metal stand or the beaker; they may be hot.

• 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.

• 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.

• 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.

16. Allow the 100 mL beaker to cool before touching it.

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17. Turn on the tap water to cold. Let the water run as cold as possible for approximately 15 seconds.

18. Fill the cup (plastic or drinking) approximately half-full with cold tap water.

19. 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.

20. 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.

21. After 1 minute stir the ice water with the thermometer.

22. 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.

23. Allow the ice to remain in the water for an additional 4 minutes.

24. After the additional 4 minutes stir the ice water with the thermometer.

25. 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.

26. 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.

Mass Measurements

27. Gather the pen or pencil, 5 pennies, 3 quarters, 4 dimes, and the key.

28. 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.

29. Turn the scale on by pressing the Φ/T button.

30. 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.

31. Review the different object(s) listed in Data Table 3 of your Lab Report Assistant.

32. 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.

33. Tare the scale by pressing the Φ/T button so that the scale reads 0.0 g.

34. 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).”

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35. Repeat steps 33 and 34 for the remaining object(s) in Data Table 3.

36. For each object(s), convert the actual mass (in grams) to kilograms (kg). Record in Data Table 3.

Questions A. Water boils at 100oC at sea level. If the water in this experiment did not boil at 100oC, what

could be the reason?

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

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Exercise 2: Volume and Density In this exercise, you will make measurements using the SI system units for length, mass, and temperature.

Procedure

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 of your Lab Report Assistant under (Mass A) column for water.

3. Fill the graduated cylinder with 5.0 mL of distilled water; use the short stem pipet to 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 + liquid (g) in Data Table 4 under (Mass B).

5. Calculate the mass of the water by subtracting “Mass A” from “Mass B.” 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.

Volume and Density Measurements (Solid)

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

Direct Measurement Method

11. Tare the scale by pressing the Φ/T button so that the scale reads 0.0 g.

12. Place the magnet on the scale to measure the mass of the object. Record the mass in Data Table 5 of your Lab Report Assistant.

13. Use the ruler to measure the length, width, and height of the magnet in centimeters. Record the measurements in Data Table 5.

Experiment Laboratory Techniques and Measurements

14. Calculate the volume of the magnet by multiplying the length × width × height, and record in Data Table 5.

15. Calculate the density of the magnet by dividing the mass by the volume and record in Data Table 5.

Water Displacement Method

16. Tare the scale by pressing the Φ/T button so that the scale reads 0.0 g.

17. Place the magnet on the scale to measure the mass of the object. Record the mass in Data Table 6 of your Lab Report Assistant.

18. Fill the graduated cylinder with 6 - 8 mL of distilled water. Record the volume, to the correct decimal place, in Data Table 6.

19. 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.”

20. Determine the volume of the object, by calculating the difference in water displacement volumes (final – initial). Record in Data Table 6.

21. Calculate the density of the magnet and record in Data Table 6.

22. Carefully pour the water from the cylinder down the drain and collect the magnet.

23. Repeat steps 16 through 22 for the metal bolt.

Archimedes’ Method

24. Tare the scale by pressing the Φ/T button so that the scale reads 0.0 g.

25. 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 of your Lab Report Assistant.

26. 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.

Experiment Laboratory Techniques and Measurements

27. Fill the glass beaker approximately ¾ full with distilled water and place the beaker on the scale.

28. Tare the scale by pressing the Φ/T button so that the scale reads 0.0 g. Refer to Figure 11 for questions.

29. 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.”

30. 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, under “Volume of Displaced Water.”

31. Calculate the density of the metal bolt and record in Data Table 7.

32. Repeat steps 24 through 31 for the magnet.

Questions A. 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)?

B. 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?

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

D. 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?

E. 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?

Experiment Laboratory Techniques and Measurements

Exercise 3: Concentration, Solution, and Dilution In this experiment, you will use laboratory equipment to create solutions of varying concentrations and densities by diluting a stock solution.

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 of your Laboratory Report Assistant under “Mass of Volumetric Flask.”

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

5. Cut a small square of white paper, fold it in half and then in quarters, and place it on the scale.

6. 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.

7. Place 8.0 grams of sugar on the paper. Record the mass in Data Table 8 under “Mass of Sugar.”

8. Carefully transfer the sugar from the white paper into the volumetric flask.

9. 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.

10. Once the sugar is dissolved, use the short stem pipet to add additional water to the 25 mL mark. See Figure 18.

Experiment Laboratory Techniques and Measurements

Figure 18. Adding water to a volumetric flask.

11. Record the total volume in Data Table 8.

12. Determine the concentration of the sugar solution in the volumetric flask using the following equation:

13. Record the concentration of the sugar solution in three places: In Data Table 8, for “Concentration,” in Data Table 9 of your Lab Report Assistant for “Final Concentration” in the “Solution 0” row, and in Data Table 9 under “Initial Concentration” in the “Solution 1” row.

14. Tare the scale by pressing the Φ/T button so that the scale reads 0.0 g.

15. Place the volumetric flask containing the solution on the scale.

16. 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.

Experiment Laboratory Techniques and Measurements

17. Record the mass of the sugar solution in Data Table 9 for “Solution 0.”

18. Calculate the density of the sugar solution and record it in Data Table 9 for “Solution 0.”

19. Pour the sugar solution (Solution 0) into the dry, empty 100 mL glass beaker.

20. Rinse the volumetric flask well with distilled water.

21. Use the serological (graduated) pipet to transfer 2.5 mL of Solution 0 into the volumetric flask.

22. Add approximately 15 mL of distilled water to the volumetric flask containing Solution 0 and swirl to thoroughly mix.

23. Using the short stemmed pipet, continue adding distilled water to the 25 mL mark of the flask.

24. 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.

25. Subtract the mass of the dry, empty volumetric flask from the mass of the flask containing the solution.

26. Record the mass of the new sugar solution in Data Table 9 for “Solution 1.”

27. Calculate the density of the sugar solution and record in Data Table 9 for “Solution 1.”

28. Determine the Final Concentration (% m/V) for Solution 1 using the following equation:

29. Record the Final Concentration of “Solution 1” in Data Table 9 under “Final Concentration.”

30. 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).

31. 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. Insert the graph into Data Table 10 of your Lab Report Assistant.

32. Clean all glassware and lab equipment with soap and water, rinse the equipment again with distilled water.

33. Dry all items with paper towels and return to the lab kit for future use.

Experiment Laboratory Techniques and Measurements

Questions A. 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?

B. From the graph of Concentration vs. Density, created in Data Table 10, what was the rela- tionship between the concentration of the sugar solution and the density of the sugar solu- tion?