EXOTHERMIC AND ENDOTHERMIC REACTIONS   
                 
HERBERT TARNOR                 DU SABLE HIGH SCHOOL
                               4934 S. WABASH AV.    
                               Chicago, IL 
                               1-312-536-8600

  OBJECTIVES: 
     (1) Students are to observe several exothermic and endothermic reactions.
     (2) Students are to determine the changes that take place in a chemical 
reaction; to observe how fast a reaction is taking place and what factors 
influence a chemical reaction. 
              
MATERIALS:  DEMONSTRATIONS.
    (1) A picture puzzle.
    (2) Two burettes, ring stand and clamp, plastic rod, fur, water, alcohol. 
    (3) Model of a crystal, nails, wooden blocks, and magnets.
    (4) Dry ice (CO2), two 1000 ml graduated cylinders, red and blue dyes, 
pneumatic trough, gas bottle, rubber tubing, fish tank, candles, steps.  
    (5) Nitrogen triiodide (NI3) 2-3 g, ring stand, 3 iron rings, three sheets of 
filter paper which are larger than the rings, feather, 2 m pole.                    
    (6) Potassium chlorate (KClO3) 6 g granulated sugar 2 g, 1 drop H2SO4, ring 
stand with metal base. 
    (7) One 50 ml beaker with  20 g barium hydroxide (Ba(OH)2.8H2O).  One 50 
ml.beaker with  10 g ammonium thiocyanate (NH4SCN), wooden block. 

LABORATORY PROCEDURE. 
     Safety glasses, 3 test tubes with stoppers, marking pencil, test tube holder, 
10 ml graduated cylinder, small spatula, stirring rod, thermometer.   Reagents: 
Sodium thiosulfate crystals, 18 M (concentrated) sulfuric acid (H2SO4), ammonium 
chloride (NH4Cl).               
                
STRATEGIES:  DEMONSTRATIONS.
    (1)  Drop the picture puzzle and allow the pieces to scatter.  Discuss entropy 
and the Second Law of Thermodynamics (basically that no process can occur unless 
there is an increase in the disorder in the universe when it happens). 
    (2)  Allow the water to stream from the burette, rub the plastic rod with fur 
and hold it close to the stream. Explain that water is polar and is being 
attracted by the electrons on the plastic rod. Repeat using alcohol and carbon 
tetrachloride.          
    (3)  Display the model of a crystal. Use a magnet to pull the ions apart in 
the crystal lattice, and show the point of attack of solvent molecules (the exposed 
corners), pointing out that the solvent molecules envelop the ions and actually 
tear the crystal apart. 
    (4)  Collect the CO2(g) and show the effect when it is poured in to the fish 
tank where the candles are burning on steps. 
                          
  LABORATORY PROCEDURE.                                     

   (1)  Label a test tube  #1.  Fill a test tube nearly full a sodium thiosulfate 
crystals.  Record the temperature. Add just enough drops of water to make the 
crystals look wet half-way up their height. Now warm the tube gently, with steady 
swirling, over a burner flame.  The wet crystals will melt rather readily. When 
the tube contents are fully liquid, and uniformly stirred so no concentration 
waves show, record the temperature. Set the tube aside to cool (if care is used, 
it can even be cooled in cold water, though in this case premature crystallization 
may occur). We now have a supercooled liquid. Even though the liquid is far below 
2
the freezing point (or saturation temperature), the sodium thiosulfate will not 
readily crystallize unless some kind of nucleus is provided. When the tube is 
cooled, drop a single crystal of sodium thiosulfate into the tube-it makes no 
difference how tiny.  Immediately crystals will start spreading from this nucleus. 
When the reaction is completed the tube will be hot. Record the temperature. 

 (2) Label a test tube #2.  Place 2 ml of tap water in the test tube. Add 10 drops 
of concentrated sulfuric acid (H2SO4) to the water.  Feel the tube and record the 
temperature. Carefully wash off the thermometer with tap water. 

  (3) Label a test tube #3. Place 1 spatula of ammonium chloride (NH4Cl) in th 
test tube. Add 2 ml of water. Stopper and shake the test tube. Record any change 
in the temperature of the solution and discard the materials. 

REFERENCES:
BOOKS:
Kerkut, G.A.        Implications of Evolution (1960).
Morris,Daniel Luzon.The Dynamic Equilibrium Approach to Teaching Chem-
                    istry (West Nyack, New York, Parker Publishing
                    Company,Inc.: 1974), p. 72.
Olmsted, Michael P. Chemistry Teacher's Guide (West Nyack, New York, 
                    Parker Publishing Company,Inc.:1972),pp.116f.
Shakhashiri, Bassam Z. Chemical Demonstrations: A Handbook for Teachers
                    of Chemistry, Vol.1 (Madison, Wisconsin, The University
                    of Wisconsin Press: 1983), pp. 10f.,96-98. 

ARTICLES
(NO AUTHOR)        "Life on Earth: From Chemicals in Space?" Chemical
                   and Engineering News (19), Nov. 19,1973, pp.21-22.
Brinkman, R. T.   "Dissociation of Water Vapor and Evolution of Oxygen
                  in the Terrestrial Atmosphere,"   Journal of Geophysical
                  Research (74), Oct.20, 1969, pp.5355-5368.
Carruthers,George R. "Apollo 16 Far-Ultraviolet Camera/Spectograph: Earth
                  Observations."Science (177), 1 Sept. 1972, pp. 788-
                  791.