Raymond Zmaczynski             Senn Metropolitan Academy
                               5900 N. Glenwood Avenue
                               Chicago, IL 60660
      The following properties of matter will be shown: density (two approaches);  
diffusion in a liquid; the effect of pressure on the boiling point of water; 
polarity of liquids; and the comparison of bond types and the physical properties 
of solids.  For each of these properties, a separate summary will be made.  

      To show the concept of density.

      Samples of metals of equal sizes (Al, Fe, Cu, Pb, Sn); unknown density 
samples (irregularly shaped pieces of sulfur and/or marble chips or assorted 
collection of fishing sinkers, available at Sears in a variety of metals and 
shapes); graduated cylinders (size depends on size and shape of unknowns); 
chemical balances; experiment sheets; graph paper. 

     The concept of density appears very early in most high school chemistry 
courses.  In many cases, the student already understands the concept or has been 
exposed to the concept.  Density is a useful and necessary property of matter and 
the student should have a good understanding of the concept.  It is assumed that 
the student has been taught to use a chemical balance and a graduated cylinder as 
tools of measurement. The idea of determining the volume of an irregularly shaped 
solid by the water displacement method may be taught previously or may be 
introduced as part of the experiment.  The student usually has the idea of 
identifying properties based on his or her general life experience. 

      Pass out for general classroom inspection or for group inspection, samples 
of some or all of the following metals: Al, Zn, Sn, Pb, Cu, and Fe (galvanized 
steel is all right to use).  Inquire of the students whether the samples are 
alike or different.  Ask whether any of the samples might  possibly be alike, 
especially those that look or feel alike. Suggest that certain pairs might be 
alike (Pb and Sn or Zn, Al, and Fe) and ask why the students don't think they are 
alike.  After some discussion, someone will usually say that they are not alike 
because one of the pieces is "heavier" than the other one.  From this, the 
discussion should lead to ways to determine which is "heavier" when the pieces are 
not the same size, have an irregular shape, or appear to have the same properties.  
The idea of density can be developed by having the student do a "density" 
experiment.  Two approaches are suggested and a possible experimental design for 
each is shown on the following pages. The first, adapted from the Chemical Guide 
and Laboratory Activities by McGill, Bradbury, and Sigler published by Lyons and 
Carnahan in 1966, introduces the concept of density and continues from there.  The 
second, adapted from the Interdisciplinary Approach to Chemistry (IAC), develops 
the concept of density by doing mass and volume measurements of different size 
samples of the same materials and then developing density by graphing the data 
collected from the entire class.  

               EXPERIMENT -- Determining the Density of a Solid

INTRODUCTION:  One of the important physical properties of a substance is 
its density, or what is its mass per unit volume.  The property of density 
may be used to help identify a substance.  Gold, for example, is much more 
dense (heavier) a substance than copper even though their colors may be 

Density is the mass in grams (g) for each milliliter (ml) or cubic 
centimeter (cm3) of a substance.  This relationship may be expressed as 
follows: Density = Mass / Volume  or D = M/V . 

PROCEDURE:  Obtain an unknown from your instructor and identify it by name 
or number in the data table below.  Determine the mass of the unknown and 
enter it in the data table. determine the volume of the unknown and then  
calculate the density of the unknown substance.  From your instructor 
obtain the accepted density of the unknown and determine the error and 
percentage of error in your experiment.  An error of only a few percent is 
considered good work.  Dry off your unknown and return it to the container 
your instructor supplies.  If you have time, you may wish to try another 

DATA:  Name or number of unknown   ____________________________________
(A)  Mass of unknown ------------------------------  ________________  g

(B)  Volume of water in graduate ------------------  ________________  ml

(C)  Volume of unknown and water in graduate ------  ________________  ml

(D)  Volume of unknown ----------------------------  ________________  ml

(E)  Density of unknown ---------------------------  ________________  g/ml

(F)  Accepted density of unknown ------------------  ________________  g/ml

(G)  Error = difference between (E) and (F) -------  ________________  g/ml

(H)  Percentage of error = (G)/(F) x 100 % --------  ________________  %

1.  What is the density of a solid whose mass is 19.57 g and whose volume 
is 3.6 ml ? 

2.  An irregular solid has a mass  of 12.77 grams.  When the solid is 
lowered into a graduated cylinder filled with 10.0 ml of water, the water 
level rises to 13.1 ml.  What is the density of the solid? 

3. The method used in this experiment would not be a satisfactory way to 
measure the volume of salt or sugar crystals.  Why? 

                 EXPERIMENT -- Mass and Volume 

Several different kinds of solid objects will be available in the 
laboratory.  Ask your instructor to identify each type and to tell you how 
many grams (g) of each to use.  Plan to keep a record of all the 
measurements you make.  Consult your instructor about how you should 
arrange your data in a table. 

Determine the mass of the specified number of grams of one of the solids.  
Then use the water-displacement method to determine the volume (in ml) of 
this quantity of material.  Be sure that you record on your data sheet the 
mass and the two measurements that you used to determine the volume of each 

Repeat these measurements with each of the other two solids that have been 
made available to you.  Record these measurements in your data table. 

Prepare a graph of the data that you and the other students have collected.  
Plot the mass (in g) of each sample along the vertical axis.  Plot the 
volume (in ml) of each sample along the horizontal axis.  Consult your 
teacher for suggestions in preparing your graph and for identifying the 
different kinds of solid materials plotted. 

Draw the best straight line you can through the points for each material.  
Study your graph.  What general pattern do you see? 

Could the graph be used to identify an unknown sample of one of the solid 
materials?              Explain.

What is your interpretation of the results?

     To show diffusion in liquids.

     Two large beakers or cylinders (usually 1000 ml, although somewhat smaller 
size probably will work too); potassium permanganate crystals or liquid food 

     While diffusion occurs in all states of matter, it is easier to observe in 
the liquid state.  Once observed, the concept can be related to all states of 
matter and to the idea of the energy and velocity of molecules.  Fill the one of 
the containers with cold water and the other container with hot water.  Allow the 
liquid in the containers to "settle."  Add a small crystal of potassium 
permanganate (or a drop of food coloring) to each container and have students make 
observations.  If desired, the containers may sit, undisturbed, until the end of 
the class period or until the next day.  

     The students should observe that the crystal of the permanganate dissolves 
faster and the color spreads out faster in the hot water than in the cold water.  
Discussion should center on why this is so.  Discussion should relate the rate of 
diffusion to the energy and velocity of the water molecules.  The discussion may be 
extended to the kinetic theory and to all the states of matter. 

     To show the effect of pressure on the boiling point of water.

     Florence flask (250 or 500 ml); rubber stopper (a stopper-thermometer 
assembly may also be used); clamp to hold the Florence flask; hot plate; ice or 
cold water. 

    [CAUTION: Make certain that the flask contains no cracks so that it will not 
explode.  An implosion hazard also exists due to the reduced pressure in the 
flask.  Goggles probably should be worn and this demonstration should be performed 
at a safe distance from the students.] 

     To many students, water boils at 100 C regardless of the atmospheric 
pressure.  This demonstration of the effect of atmospheric pressure on the boiling 
point of water may serve as an introduction to or as reinforcement of the standard 
definition of boiling point that "the boiling point is the temperature at which 
the vapor pressure is equal to the atmospheric pressure." 

     Fill the flask about one-third with water.  Heat the water to boiling in the 
unstoppered flask.  As the water is heated, ask the class about the  temperature 
change and its effect on the energy in the water and on the vapor pressure of the 
water.  As the water reaches its boiling point, relate the vapor pressure of the 
water to the atmospheric pressure as a definition of the boiling point.  Remove 
the flask from the heat and wait until the boiling stops.  Stopper the flask 
tightly.  Using the clamp, invert the flask. Rub the portion of the flask now 
containing the air (the bottom) with ice or put the inverted flask under running 
cold water.  The water will begin to boil again.  Ask the students to explain 
their observations.  After the demonstration turn the flask right side up and 
remove the stopper to equalize the pressure.  If the stopper that is inserted into 
the flask is fitted with a thermometer, the temperature at which the water 
boils can be determined.  The discussion should be related to the relationship 
between temperature and pressure of a gas (air, in this case) and the relationship 
among boiling behavior, air pressure, vapor pressure of a liquid, and temperature. 
The class should be asked to define boiling and boiling point. 
     To show the polarity of liquids.

     Burettes (the number depends on how many liquids you wish to show); 100 ml 
beakers, as needed; liquids to be tested (suggested liquids: water, carbon 
tetrachloride, ethanol, acetone, benzene);  source of static electricity 
(suggested sources or methods: plastic rod or comb and a piece of flannel cloth; 
plastic rod and animal fur; have students comb their hair vigorously with a 
plastic comb).  If only the polarity of water is to be shown, produce a fine 
stream of water from an available water faucet and have a source of static 

     Water is a common substance but one with special properties.  Much is usually 
said about the polarity or non-polarity of molecules in chemistry.  The polarity 
of the water molecule and the importance of this polarity in the ability of water 
to act as an "almost universal solvent" is usually stressed in the high school 
chemistry course.  What about other common liquids or solvents?  The following 
demonstration should help the student to "see" the polarity of some common 
liquids.  Following the demonstration, the shapes and structures of the molecules 
of the liquids can be related to their polarity.  From this starting point, 
depending on when in the course the demonstration is used, the teacher may develop 
further ideas on the structure of molecules and their shapes or deal with the 
properties of liquids and solutions. 

      Fill each burette with a different liquid.  Open the stopcock of the first 
burette and let the liquid flow out into a receiving beaker with a fairly fine 
stream.  Charge the plastic rod or comb and bring it near the stream of liquid as 
it flows from the burette.  Move the charged rod up and down near the stream.  If 
the liquid is polar (water, acetone, or ethanol), the stream will be deflected 
toward the charged rod.  If the liquid is non-polar (benzene or carbon 
tetrachloride), the stream flow will be unaffected by the charged rod.  When the 
burette is emptied, it may be refilled with the liquid from the collecting beaker.  
Interested students should be invited to test the liquids to see if individual 
differences in combs, rubbing techniques, and movements of the rod make a 
difference in the polarity of the liquids.  After all the students have satisfied 
themselves as to which liquids are polar and non-polar, class discussion should 
focus on which liquids are affected by the charged rod and why this is so.  

     Static electrical charges are more easily produced in dry weather (winter 
time), although this demonstration will work on a hot, humid day in the summer 
with the right materials and vigorous rubbing of the rod or comb.  Benzene and 
carbon tetrachloride are substances that are known carcinogens, but they are also 
the most common non-polar substances available.  The amounts used and the length 
of this demonstration, about 10 to 15 minutes, should not preclude their use.  
Nevertheless, other non-polar liquids that are apparently non-carcinogens , such 
as hexane or cyclohexane, may be used and should probably work just as well as 
benzene and carbon tetrachloride.  If one wishes to show only the polarity of 
water and does not wish to use a burette (or does not have a burette), an 
alternate procedure is available.  Adjust the water stream from a classroom faucet 
so that a fine stream of water is flowing.  This flow is approximately the same as 
that from a burette. The same effect can now be demonstrated using the water flow 
from the faucet when the charged rod is brought near the water flow. 

     To examine some of the physical properties of typical ionic and covalent 
molecular solids and to relate these properties to the type of bond found in each 

     sodium chloride, paradichlorobenzene, test tubes, bunsen burner, water, 
benzene ( or some other non-polar solvent).  The amounts of materials and 
equipment used will depend on whether the experiment is performed individually or 
in groups by the class or whether the instructor demonstrates the experiment for 
the class. 

     In a chemistry (or another science) course, it is helpful for the students to 
"see" the differences in properties between an ionic and molecular covalent solid.  
Sodium chloride and paradichlorobenzene are chosen as the typical solids because 
they are usually readily available and inexpensive for high school use.  
Additional typical ionic and covalent solids may be added to or substituted for 
these two substances depending on the purposes of the instructor.  This experiment 
may be used to introduce chemical bonding or the types of chemical bonds.  Since 
ionic and covalent molecular bonds are generally synonymous with inorganic and 
organic chemistry respectively, this experiment could serve as an introduction to 
the topic of organic chemistry.  In a biology, physical or general science course,   
the experiment could be demonstrated to summarize the general differences between 
ionic and covalent bonding compounds.  The experimental format that follows is 
based on an experiment in Laboratory Keys to Chemistry by Ledbetter and Young, 
published by Addison-Wesley in 1973.  

     Benzene is the best non-polar solvent for this experiment although carbon 
tetrachloride may also be used satisfactorily.  Unfortunately, both have been 
identified as known or probable carcinogens to humans.  Other non-polar solvents 
that are not probable carcinogens could probably be used instead of benzene or 
carbon tetrachloride but such an inexpensive, readily available substitute does 
not immediately come to mind.  This is a dilemma for the instructor, but not a 
readily soluble one. 


The bonds in ionic solids are formed by the transfer of electrons between 
atoms.  As a result, such solids consist of positive and negative 
ions arranged symmetrically.  The bonds in molecular solids are formed by 
sharing electrons; ions do not exist in such solids.  The forces between 
molecules in molecular solids are weak van der Waals forces.  In this 
experiment, you will investigate the relationship between bond types and 
the physical properties of solids.  

Procedure:  Obtain samples of sodium chloride (an ionic solid) and 
paradichlorobenzene (a molecular solid).

1.  Note the odor of each solid. An odor indicates that some of the solid's 
molecules can evaporate at room temperature. Record your observations.

2.  Rub a sample of each solid between your fingers.  Record the hardness 
of the solids by noting whether the "feel" is soft or hard and granular.

3.  Place a few crystals of sodium chloride in one test tube and a similar 
amount of paradichlorobenzene (PDCB) in a second test tube.  Heat each 
test tube separately. Heat cautiously at first and then more intensely if 
the solid does not melt readily.  (CAUTION.  If the solid does melt with 
gentle heating, do not heat it further.  It may also vaporize, producing 
fumes.  Be careful not to inhale such fumes since they may be poisonous.)  
Record your observations.

4. Place a few crystals of each solid in separate test tubes.  Add about 5 
ml of water to each test tube.  Shake or stir to mix. Compare the 
solubilities of the two solids in water.  Record your observations.

5.  Place a few crystals of each solid in separate test tubes.  Add about 5 
ml of benzene to each, shake or stir to mix.  Compare the solubilities of 
the two solids in benzene.  Record your observations. 

Conclusions. 1.  Prepare a chart in which you summarize the observations made 
on the ionic and covalent molecular solids.
2.  Compare the general properties of an ionic solid and of a covalent 
molecular solid.  
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