CHEM 01A ExptA Determination of Densities F14 -...

CHEM-‐01A -‐ Experiment A: Determination of the Densities of Some Liquid and Solid Samples

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Name _______________________________________________ Date __________________Grade ____

Pre-‐Lab Questions: Determination of the Densities of Some Liquid and Solid Samples Show all work and be sure to include units. Express all answers to the correct number of significant figures. Q1. Which of the following properties is extensive or intensive?

a. Mass b. Volume c. Density

Q2. What are the metric units for the following properties of a liquid?

a. Mass b. Volume c. Density

Q3. Considering the definitions of extensive and intensive properties, which properties would be most useful identifying an unknown sample, such as an unknown metal sample? Why?

Q4. When you measure a liquid that has a meniscus that is concave down, do you read the bottom, top, or

middle of the meniscus? (You may want to sketch a diagram.)


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Q5. A student obtains a silvery piece of metal of an unknown identity. She weighs the metal and the mass is 56.58 g. When she places the metal in a graduated cylinder that initially contains 25.8 mL of water, the final volume reading on the graduated cylinder is 32.2 mL. a. Calculate the density of unknown metal sample. (Show your calculations with units.) b. Use the density information in your textbook or the CRC to determine the identity of the metal from

these possible unknowns: chromium, cadmium, molybdenum, nickel, or zinc. If you use the internet, cite your source.

c. Calculate the percent error of the measured value. (Show your calculations with units.) d. Is the density measurement accurate? Is the density measurement precise? Explain.


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Experiment A: Determination of the Densities of Some Liquid and Solid Samples Purpose

The purpose of this experiment is to learn a variety of techniques used to measure the density of samples in liquid or solid phase.

Introduction Matter is characterized by its physical and chemical properties. A physical property can be determined without chemically changing the material. Physical properties include color, odor, taste, mass, length, melting point, boiling point, conductivity, volume, and hardness. These properties are relatively easy to determine. Chemical properties tell us how the substance reacts with other substances and may include reaction with oxygen, chlorine, metals, etc. Determination of chemical properties results in the change of the identity of the substance.

Matter can also be categorized into two other properties:

1) Extensive properties: Properties that depend on the amount of matter present.

2) Intensive properties: Properties that do not depend on the amount of matter present. A) Measurements

In all sciences, measurements are essential, and the most fundamental properties that can be measured are length, mass, and time. In chemistry, temperature is also treated as a fundamental property. Other properties of matter -‐ such as volume, area, and density -‐ are ratios or products of the fundamental properties. For example,

(a) units of area are length x width, or (distance)2

(b) units of volume are length x width x height, or (distance)3

(c) units of density are mass/volume, or mass/(distance) 3

The metric system is used almost exclusively in all sciences. The meter, kilogram, and the second are the basic units in the International System of Units (SI), but the meter and the kilogram are generally too large for convenient use in the chemistry laboratory. Units in metric system are related to each other as power of ten and associated with prefixes described below:

Table 1: Prefixes to represent small or large units.

Large Units Small Units Prefix Symbol Multiple Prefix Symbol Multiple

tera-‐ T 1012 Unit 1

giga-‐ G 109 deci-‐ d 0.1

mega-‐ M 106 centi-‐ c 0.01

kilo-‐ K 1000 milli-‐ m 0.001

hecto-‐ H 100 micro-‐ µ 10–6

deka-‐ Da 10 nano-‐ n 10–9

Unit 1 pico-‐ p 10–12

This experiment has been designed to acquaint you with several types of measurements and measuring devices, and use the measurements to calculate density of liquid and solid substances.


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B) Density One of the fundamental properties of any sample of matter is its density. This property is dependent on the mass and the volume of the sample. The relationship between density, mass, and volume is:

Density = massvolume

𝑈𝑛𝑖𝑡𝑠 = 𝑔𝐿 𝑓𝑜𝑟 𝑔𝑎𝑠 ,

𝑔𝑚𝐿

𝑓𝑜𝑟 𝑙𝑖𝑞𝑢𝑖𝑑𝑠 ,𝑔𝑐𝑚3 (𝑠𝑜𝑙𝑖𝑑𝑠)

The density of a liquid or of a solution is usually reported in units of grams per milliliter (g/mL). The density of a solid is reported in units of grams per cubic centimeter (g/cm3). Because 1 mL is equivalent to 1 cm3, these units are interchangeable. The density of water is 1.00000 g/cm3 at 4°C, and 0.9970 g/cm3 at 25°C.

For any density determination, the mass and the volume of the substance must be determined.

C) Mass Determination The mass of a solid sample can easily be determined by weighing a sample on a calibrated laboratory balance. Generally, solids are placed on a tared weighing paper (or in a weighing boat) and never placed directly on the balance pan. The mass of a sample of liquid in a container can be found by taking the difference between the mass of the container with the liquid and the mass of the empty container:

𝑀𝑎𝑠𝑠!"#$"% = 𝑀𝑎𝑠𝑠!"#$"%!!"#$%&#'( − 𝑀𝑎𝑠𝑠!"#$%&#'( D) Volume Determination

The volume of a liquid can easily be determined by means of “graduated” containers such as graduated cylinders, pipettes, burets, or others for routine measurements. The volume of a solid can be determined by direct measurement if the solid has a regular geometric shape such as a cube, rectangle, or cylinder.

𝑉!"#$ = 𝑙!,𝑤ℎ𝑒𝑟𝑒 𝑙 is the length of one side of the cube.

𝑉!"#$%&'" = 𝑙.𝑤. ℎ,𝑤ℎ𝑒𝑟𝑒 𝑙,𝑤, 𝑎𝑛𝑑 ℎ 𝑎𝑟𝑒 𝑡ℎ𝑒 𝑙𝑒𝑛𝑔ℎ,𝑤𝑖𝑑𝑡ℎ 𝑎𝑛𝑑 ℎ𝑒𝑖𝑔ℎ𝑡 𝑜𝑓 𝑡ℎ𝑒 𝑟𝑒𝑐𝑡𝑎𝑛𝑔𝑙𝑒, 𝑟𝑒𝑠𝑝𝑒𝑐𝑡𝑖𝑣𝑒𝑙𝑦.𝑉!"#$%&'(= 𝜋𝑟!ℎ,𝑤ℎ𝑒𝑟𝑒 𝑟 𝑖𝑠 𝑡ℎ𝑒 𝑟𝑎𝑑𝑖𝑢𝑠 𝑜𝑓 𝑡ℎ𝑒 𝑏𝑎𝑠𝑒, 𝑎𝑛𝑑 𝑙 𝑖𝑠 𝑡ℎ𝑒 ℎ𝑒𝑖𝑔ℎ𝑡 𝑜𝑓 𝑡ℎ𝑒 𝑐𝑦𝑙𝑖𝑛𝑑𝑒𝑟.

A convenient way to determine the volume of an irregularly shaped solid such as a rock is to measure accurately the volume of liquid displaced (raised) when the solid is dropped (immersed) in the liquid. The volume of the solid will equal the volume of liquid that it displaces (see Figure 1).

𝑉!"#$% = 𝑉!"#$%!!"#$% − 𝑉!"#$%

Figure 1: Determination of volume by displacement


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E) Using a Graduated Cylinder Graduated cylinders are an important piece of equipment that we frequently use in a chemistry laboratory. Graduated cylinders come in various sizes such as 10-‐mL, 25-‐mL, 50-‐mL, 100-‐mL, 500-‐mL, and 1,000-‐mL. Similar to graduated cylinders, you may use other graduated glassware such as volumetric flasks and burets. When water or a clear aqueous solution is contained in a graduated cylinder, a meniscus forms as shown in Figure 2 below. The meniscus is the curve seen at the top of a liquid in response to its container. The meniscus can be either concave or convex. A concave meniscus as shown in the Figure 2a (e.g., water in glass) occurs when the molecules of the liquid are more strongly attracted to the container than to each other. A convex meniscus as shown in Figure 2b (e.g., mercury in glass) is produced when the molecules of the liquid are more strongly attracted to each other rather than to the container. In some cases, the meniscus appears flat (e.g., water in a narrow plastic cylinder).

When you read the scale on the side of a container with a meniscus, such as a graduated cylinder or volumetric flask, it is important that the measurement accounts for the meniscus. Measure so that the line you are reading is even with the center of the meniscus as shown in Figure 2c. For water and most liquids, this is the bottom of the meniscus. For mercury, take the measurement from the top of the meniscus. In either case, your measurement should be based on apex of the meniscus.

Most importantly, along with paying attention to meniscus reading, keep in mind the significant digit rules. For example, the graduated cylinder reading for Figure 2c should be recorded as 36.7mL. Note that bottom of the meniscus – indicated with the dotted line -‐ falls between 36-‐mL and 37-‐mL. Thus, you should estimate the volume between those minor increments. This “estimated digit” is always one decimal past the smallest increment. If the smallest increment is 10-‐mL, then you estimate the ones place. If the smallest increment is 1-‐mL, then you estimate to the tenths. If the smallest increment is 0.1-‐mL, then you estimate to the hundredths. And so on…

Figure 2: Formation of meniscus


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F) Percent error determination Analyzing possible errors is an important part of any experiment. Percent error is a useful formula to determine the accuracy of your experimental data. It is determined by taking the absolute value of the difference between theoretical and experimental values, divided by the theoretical value and multiplied by 100. Theoretical value is also referred to as the accepted value or true value. It is a value we depend on and is obtained from a reliable resource.

% 𝐸𝑟𝑟𝑜𝑟 = |𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑉𝑎𝑙𝑢𝑒 − 𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑉𝑎𝑙𝑢𝑒|

𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑉𝑎𝑙𝑢𝑒 The difference expressed within the “|” symbol indicates absolute value of the difference which means that the difference should be expressed as a positive quantity. Every determination is expected to have some error. Excluding the possibility of calculation error, a relatively high percent error usually points to poor lab techniques and/or errors in reading graduations. True value is often a value that is determined by more precise and more accurate technique. Resources Theoretical values or true values of a parameter, such as density, can be found in reliable references such as the CRC Handbook of Chemistry and Physics and The Merck Index, both of which are available in the laboratory. The internet can be searched for a true value, but the reliability of the source should be considered. For example, the ‘Sigma-‐Aldrich Chemical Catalog’ is a reliable resource, whereas sources like ‘Wikipedia’ and ‘Ask.com’ are often unreviewed and can be written by outside contributors. Regardless of their source(s), you should always cite your references.


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Experimental Procedure

Safety Considerations for this experiment:

1) If you break a piece of glassware during this experiment, notify your instructor immediately and dispose of the

broken pieces in the broken glass container.

2) Dispose of used alcohol in the waste container. Your instructor will point out the location of the waste

container.

3) Keep hot plates at a safe distance. The surface of the hot plate can get extremely hot and cause burns.

4) Do not pick up beakers containing hot water with bare hands.

5) If using digital thermometers, do not leave them in the container with boiling water. The plastic cover of the

thermometer might melt.

6) To avoid any damage from spillage, keep your data sheets at a safe distance from chemicals.

7) REMINDER -‐ eating or drinking is NEVER allowed in the laboratory.

Procedure

Part A. Mass Measurements: Be sure to record data on the following pages under appropriate section.

1) Using one of the top-‐loading balances on the countertop, take a coin (penny) and measure its mass to ± 0.01 g.

Make sure that the pan is clean and the balance was zeroed before taking the measurement.

2) Record the mass on the report sheet under ‘Balance 1’

3) Repeat step 1 with same penny on a different top-‐loading balance on the countertop. Record this mass under

‘Balance 2’

4) Using an analytical balance, measure the mass of the same penny to ± 0.0001 g. Record this mass under

‘Balance 3’

5) Repeat step 3 with same penny on a different analytical balance. Record this mass under ‘Balance 4’.

B. Volume Measurements: Be sure to record data on the following pages under appropriate section.

1) Read and record the volume of liquid in the graduated cylinder designated by the instructor.

2) Using a transfer pipette, measure and record the volume of 30 drops of tap water in a 10-‐mL graduated

cylinder.

3) Using a 10-‐mL Mohr (graduated) pipette, transfer 10 mL of tap water to a 50-‐mL graduated cylinder. Read and

record the volume of tap water in the graduated cylinder.

C. Temperature Measurements: Be sure to record data on the following pages under appropriate section.

Using a thermometer, determine the temperature of the following in oC. Then show conversion calculations to oF.

1) 50 mL of tap water in a 150-‐mL beaker.

2) Approximately 100 mL of tap water in a 250-‐mL beaker, with a few ice cubes added.

3) 100 mL of boiling deionized water in a 250-‐mL beaker. (For this step, you will need to boil water using a hot

plate.)


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D. Density Measurements: Be sure to record data on the following pages under appropriate section.

D.1 -‐ Method I for Solids

1) Weigh a given solid on a balance to the nearest 0.01g and record. Make sure you zero the balance before

proceeding with each measurement. Record the type of solid used on the report sheet.

2) Fill a 50 or 100-‐mL graduated cylinder about half full with tap water and record the volume as accurately as

possible, keeping in mind the significant figure rules.

3) Next, place the solid object in the graduated cylinder and record the water level after the solid is immersed.

4) Determine volume of the solid object and calculate density of the solid.

D.2 -‐ Method II for Solids

1) Using a ruler from your lab drawer or meter stick, measure and record the length, width, and height

(thickness) of a given wooden block, or the length and diameter of a metal cylinder.

2) From the dimensions in step 1, calculate the volume of the block.

3) Weigh and record the block on a balance to the nearest 0.01g. Make sure the balance is properly zeroed before

taking the measurement.

4) From the mass and volume data, calculate density of the solid.

D.3 -‐ Method for Liquids

1) Using the provided graduated cylinder or the volumetric pipette, measure 10 mL of rubbing alcohol (isopropyl

alcohol). Read and record the correct volume of alcohol.

2) Weigh a 50-‐mL Erlenmeyer flask to the nearest 0.01g and record the mass. Add the alcohol to the flask,

reweigh, and record under trial 1. Dispose of this liquid into the waste bottle.

3) Repeat steps 1 and 2. Record the data under trial 2.

4) From your data, calculate the average density of alcohol.


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Data Sheet Name _____________________________________________________________ Date _____________________________ Grade _______ A) Mass Measurements

Balance 1 (± 0.01 g) ________________ g ______________ mg

Balance 2 (± 0.01 g) ________________ g ______________ mg Balance 3 (± 0.0001 g) ________________ g ______________ mg

Balance 4 (± 0.0001 g) ________________ g ______________ mg

B) Volume Measurements 1) Volume of water in graduated cylinder (on the instructor’s desk) _________ mL 2) Volume of 30 drops of water _________ mL 3) Volume of water delivered by pipet in graduated cylinder _________ mL C) Temperature Measurements 1) Temperature of tap water _________ 0C _________ 0F (Show calculations for conversion of 0C to 0F): 2) Temperature of tap water + ice _________ 0C _________ 0F (Show calculations for conversion of 0C to 0F): 3) Temperature of boiling water _________ 0C _________ 0F (Show calculations for conversion of 0C to 0F):


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Density Measurements D.1) Method I: Solid Type of Solid: _________________ Mass of solid _________ g Initial volume of water (V1) _________ mL Final volume of water (V2) _________ mL Volume of solid (V2 – V1) _________ mL Density of solid _________ g/mL (show your calculations): D.2) Method II: Solid Type of Solid: Wood Block Mass of wooden block _________ g Dimensions: L = _________ cm

W = _________ cm H = _________ cm

Volume: __________ cm3

Density of solid _________ g/mL (show your calculations): D.3) Liquid Trial 1 Trial 2 Volume of alcohol _________ mL ________mL Mass of empty 50-‐mL flask _________ g _________ g Mass of flask and alcohol _________ g _________ g Mass of alcohol _________ g _________ g Density of alcohol _________ g/mL _________ g/mL (show your calculations): Average density of alcohol ___________ g/mL (Show your calculations):


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Post-‐Lab Questions:Determination of the Densities of Some Liquid and Solid Samples (All questions must be answered during the lab and must be submitted with your lab report at the end of the lab period) Please answer the following questions and show all work and units. Express all answers to the correct number of significant digits. Please show your calculations with units and significant figures: Q1. Define the following important terms pertaining to measurements. Give examples from this lab.

a. Accuracy

b) Precision

Q2. What is the difference between the mass and weight of a material? Q3. Calculate the density of a rectangular block, which has a mass of 25.71 g. The dimensions of the solid are

2.30 cm long, 2.01 cm wide, and 1.82 cm high. Q4. A rectangular wooden block, 22 cm x 13.2 cm x 4.4 cm, has a mass of 1562.0 g. What is the density of the

wood in kg/m3.


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Q5. A 24.20 g sample of an irregularly shaped metal piece displaces 1.26 mL of water in a graduated cylinder. Calculate the density of the metal.

Q6. The density of magnesium is 1.7 g/cm3, while that of iron is 7.9 g/cm3. A block of iron has a mass of 819 g.

What is the mass of a block of magnesium that has the same volume as the block of iron? Q7. A flask has a mass of 78.23 g when empty and 593.63 g when filled with water. When the same flask is filled

with concentrated sulfuric acid, H2SO4, the total mass is 1026.57 g. What is the density of concentrated sulfuric acid? (Assume water has a density of 1.00 g/cm3 at the temperature of the measurement.)

Q8. If 15 drops of ethanol from a medicine dropper weigh 0.60 grams, how many drops does it take from a

dropper to dispense 1.0 mL of ethanol? The density of ethanol is 0.80 g/mL. Write out complete dimensional analysis setup.


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Q9. A perfect cube of aluminum metal was found to weigh 20.00 g. The density of aluminum is 2.70g/mL. What are the dimensions of the cube? Hint: Calculate volume and then covert to the dimensions of a perfect cube.

Q10. Use the available resources in the laboratory and find the densities of the following substances at 25oC:

a. gold metal

b. methyl alcohol (methanol)

c. zinc chloride

Q11. Classify the following properties of hydrogen gas as either intensive or extensive.

d. The mass of the gas sample

e. The average speed of a molecule the gas sample

f. Temperature

g. Density

h. Number of molecules present


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Q12. Use the CRC Handbook of Chemistry and Physics and look up the density of water at 0oC, 10oC, 20.0oC, 50.0oC, 70oC, and 100.0 oC. Does temperature affect density? Why?

Q13. What is the difference between specific gravity and density? What are the units of specific gravity?

(You may need to research “specific gravity” in your textbook.) Q14. What is the specific gravity of alcohol having a density of 0.79 g/mL? Q15.* Talc is a mineral with low conductivity for heat and electricity, and it is not attacked by acid. It is used as

talcum powder and face powder. A sample of talc weighs 35.97 g in air and 13.65 g in mineral oil (density = 1.75 g/cm3). What is the density of talc? [This relates to buoyancy and Archimedes Principle. Hint: use the difference in weight and calculate the amount of mineral oil displaced. This volume of mineral oil is equal to the volume of the talc.]

CHEM 01A ExptA Determination of Densities F14 -...

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