Experiment #2

Physical Properties of Substances

Background

Chemists frequently take advantage of the physical properties of substances to characterize, identify, and isolate them. Examples of physical properties are color, odor, density, melting point, boiling point, crystalline form, hardness, malleability, ductility, thermal conductivity, and electrical conductivity. In this experiment you will study the physical properties of a variety of substances. You also will learn how to use the Bunsen burners and to dispense reagents properly.

The basis of science is observation. Good scientists make clear and concise observations of the world around them. Scientists sometimes make inferences based upon their observations of substances. In this experiment you will make observations of density, physical state, volatility, color and odor. You should exercise care when smelling any substance because it may have an objectionable odor or be toxic. The proper technique for smelling substances is to waft the vapor towards your nose so that you may smell a very small quantity of it. You should never taste substances you find in the chemical laboratory as they may be toxic.

Experimental Procedure

Part I – Mechanics and Use of Bunsen Burners

Take apart the Bunsen burner as demonstrated by your instructor. Draw a Bunsen burner and identify each of the following parts of the Bunsen burner: barrel, stand, hose, and valve.

Reassemble the burner and light it. Adjust the flame so that you can see the lighter blue inner cone and draw a good flame on your lab sheet using colored pencils.

Test different parts of the flame for hotness. Insert a wooden splint at the top of the flame and observe the time it takes to ignite. Insert the other end of the splint just above the light blue cone and observe the ignition time. Compare to the previous measurement.

Part II – Proper Reagent Dispensing Technique

A chemist should always be confident of what is in a reagent bottle. To avoid contamination, nothing should ever be introduced into any reagent bottle. For this reason we have very specific rules for dispensing reagents.

1. For any reagent, the container cap should never be placed on the lab bench.

2. A reagent should be poured into a beaker or onto a piece of weighing paper, not into a test tube.

3. Spatulas, scoopulas, pipettes or other objects should never be introduced into the reagent bottle.

4. Excess sample should never be poured back into the reagent bottle. If you take more than you need, either offer the excess to another student or dispose of it properly.

The proper techniques will be demonstrated by your instructor.

Part III – Observations of Physical Properties

You will make two sets of observations

1. Observe pre-measured masses of the elements and compounds listed on your data sheet. Record the colors, physical states, and physical appearances of each. What can you infer regarding the relative densities, melting points or boiling points of these substances?
2. Observe the odors of the substances provided. Be sure to use the proper technique as demonstrated by your instructor.

Part IV – Observation of a Physical Changes and Crystal Structures

Some physical properties can only be observed by a physical change. Examples of these properties are melting point, boiling point, and sublimation point. (Look up the definition of sublimation.)

Sulfur

Pour some sulfur into a beaker. Observe the crystalline structure of the sulfur under a magnifying glass. Prepare a cone of filter paper and support it either in a funnel or a small beaker. Transfer enough of the sulfur to fill a 15 cm test tube to within three cm of the top. Holding the test tube with a test tube holder, heat it slowly and uniformly in order not to superheat any portion of it. The sulfur will darken if it is superheated; this can be avoided by moving the test tube in and out of the flame. When the sulfur is just melted, it should be an amber yellow, straw-colored, liquid. Pour it into the filter cone previously prepared, and with a wood splint in hand, watch the formation of long needle-shaped crystals. Just as the surface of the liquid begins to solidify, break it open with the wood splint, and quickly pour the remaining molten sulfur into a beaker of tap water. Briefly let the filter cone cool, then break it open and observe the crystals. Compare the newly formed crystals and the sulfur that falls into the water to each other and to the original sample of sulfur.

Iodine

Place a few small crystals of iodine in a dry 250-mL beaker. Cover the beaker with a watch glass and place ice on the watch glass. Support the beaker on a ring stand with wire guaze, and heat the iodine slowly until all the crystals vaporize and the vapor deposits on the bottom of the watch glass. Under a magnifying glass, observe the crystals that form.

Part V – Separation of Substances Based on Physical Properties

This part of the experiment involves separating a salt and sand from a mixture of the two. The separation is based on the fact that the salt is soluble in water and sand is not. The stockroom has made a mixture of sand and salt (sodium chloride) for each student. Each mixture has a different mass percent of salt. You will be required to separate this sample into its component substances and will be graded on how carefully you accomplish this separation. (Your grade is determined by comparing your experimental percent salt with the true value. Be careful!) Your mixture with your name on the label is on the instructor’s desk.

Weigh a clean, dry 150-mL beaker and a clean, dry evaporating dish, each to maximum precision of the balance (0.001 g or 0.0001 g). Also weigh the container with the unknown sample. Without loss, transfer the entire unknown to the weighed beaker. Weigh the empty sample container, and the beaker plus the unknown. (Note: This double weighing allows a check of the mass of the sample.)

Add about 10 mL of deionized water to the unknown in the beaker; gently warm this, while stirring, until the salt dissolves. Then let the mixture settle and carefully decant the clear solution (supernatant liquid) down the stirring rod into the weighed evaporating dish. Add a second 10-mL portion of deionized water to the sand mixture and repeat the extraction and decantation. Place the evaporating dish and solution on a wire gauze which is on a ring stand, and begin careful evaporation, using a cool flame. The complete extraction of any remaining salt from the sand will probably require that a third and possible a fourth 10-mL portion of deionized water be added to the evaporating solution. Finally, decant as much water as possible from the sand (without any loss of sand), and gently heat the beaker of moist sand on a hot plate until it is thoroughly dry. Let this cool completely and weigh it.

As the solution in the evaporating dish becomes concentrated and crystals form, be very careful to avoid splattering. When the salt is completely dry, let the dish cool completely and weigh it. Reheat the dry salt, cool, and reweigh to make sure the sample is completely dry. If the 1^st and 2^nd weights differ by more than 0.01 g, heat for a 3^rd time and reweigh. (This is called drying to a constant mass. Use the last mass in all of your calculations.) From these data you can calculate the percent salt in your unknown mixture.

Observe the sodium chloride that has crystallized in the evaporating dish and compare it to sodium chloride crystals found on the reagent shelf. Scrape some of the salt out of the evaporating dish and observe them under a microscope. Do the same with a few crystals of salt from the reagent bottle. Compare your observations.

Name

Date

Section

Prelaboratory Exercise

Experiment #2

Physical Properties of Substances

You are advised not to taste any chemicals in this experiment. Why not?

1. What is the conversion of a solid to a gas called?

2. A student took a white, crystalline solid and determined its melting point. The student was making an ___ ___ ___ ___ ___ ___ ___ ___ ___ ___ ___.

3. On the basis of what he saw and recorded, the same student decided the solid was sodium chloride. He was drawing an _I_ ___ ___ ___ ___ ___ ___ ___ ___.

4. To pour off a supernatant liquid is to ___ ___ ___ ___ ___ ___ it.

5. While doing Part V of this experiment, a student weighed the container with the unknown, poured it into a weighed beaker, then washed, dried and reweighed the original container. What did the student do that should not have been done?

6. A student obtained the following data. Fill in the blank spaces.

mass of sample container and sample	32.345 g
mass of sample container alone	29.972 g
mass of sample (answer box 1)
mass of evaporating dish and salt (1^st weighing)	45.676 g
mass of evaporating dish and salt (2^nd weighing)	45.671 g
mass of evaporating dish	44.780 g
mass of salt extracted (answer box 2)
percent of salt in sample (answer box 3)

7. Suppose the student in question #7 had weighed the empty sample container incorrectly, getting a lower value than it actually weighed. How would this affect the percent salt result (high or low)?

8. Iodine crystals may be easily ___ ___ ___ ___ ___ ___ ___ ___.

9. All data are to be recorded directly onto the data page in ___ ___ ___ !

Physical Properties of Substances Data Sheet

Part I Mechanics and Use of Bunsen Burners

Draw a Bunsen burner and identify the following parts of the Bunsen burner: barrel, stand, hose, and valve

Record your observations from your tests with the wooden splint, and draw a good flame. Use an arrow to identify the hottest portion of the flame.

Part II – Proper Reagent Dispensing Technique

Why is it important to dispense reagents properly?

Part III – Observations of Physical Properties

A. Record your observations regarding each of the following substances. Be sure to note the color, physical appearance, and physical state of each.

Mercury

Copper

Magnesium

Iodine

Bromine

Carbon

Copper carbonate

Copper nitrate

Potassium permanganate

Potassium dichromate

Can you make any inferences regarding the relative densities of copper, magnesium, carbon and mercury?

Can you make any inferences regarding the relative melting points of mercury and copper?

B. Record the odors of each of the following substances.

Acetone

Water

Ethanol

Para dichlorobenzene

Can you make any inferences regarding the relative boiling points of each of the substances above based on the odors you detected? Why or why not? Explain your reasoning.

Part IV – Observations of Physical Changes and Crystal Structures

Characterize and sketch crystals of the various salts you have prepared.

Sulfur

Crystals from reagent bottle

Crystals formed in filter cone

Amorphous sulfur in water (Look up this term.)

Iodine

Crystals from reagent bottle

Crystals from watch glass

Part V – Separation of Substances Based on Physical Properties

Mass of sample container and sample		Mass of beaker and sample
Mass of sample container		Mass of beaker and sand (extracted, dry)
Mass of sample (from sample container)		Mass of beaker
Mass of evaporating dish and salt	1^st weighing	Mass of sample (from beaker weighings)
	2^nd weighing
Mass of evaporating dish		Mass of sand
Mass of salt extracted		Mass of salt (by subtracting the beaker weighings)

Average mass of salt

Mass percent salt in sample

Sample Calculations

Draw the sodium chloride crystals from the reagent bottle and from your experiment.

Sodium chloride from reagent bottle

Sodium chloride from evaporating dish

How do the crystals differ?

Questions and Problems

1. What supporting evidence of Dalton’s atomic theory does the formation and growth of crystals suggest? (Base your answer on the observation you noted in part IV of the report.)

2. An experiment calls for 60.0 g of concentrated nitric acid, density 1.42 g/mL. Suppose no balance is available, and you decide to use a graduated cylinder. What volume should you use?

3. The dilute sulfuric acid on the laboratory desk has a density of 1.18 g/mL, and is 25.0% by mass sulfuric acid, the remainder being water.
1. How many grams does 15.0 mL of this acid weigh?

2. What is the weight of pure sulfuric acid in this 15.0 mL of solution?

4. Calculate the volume of magnesium, density 1.74 g/cm³, which would be equal in mass to 350.0 cm³ of lead, density 11.4 g/cm³.