WB01542_.gif (729 bytes) Demonstration 1: Growing Crystals in Gels
WB01542_.gif (729 bytes) Demonstration 2: An Inorganic Polymerization Demonstration
WB01542_.gif (729 bytes) Demonstration 3: Supersaturated Solution
WB01542_.gif (729 bytes) Demonstration 4: Precipitation and Redissolution of Calcium Carbonate
WB01542_.gif (729 bytes) Demonstration 5: Pot-O-Gold

CAUTION: Use appropriate safety guidelines in performing demonstrations.










Demonstration 1: Growing Crystals in Gels


In nature minerals are often found crystallized in bands. Large, nearly perfect specimens are occasionally found. These phenomena can be simulated by preparing a gel from water glass (a solution of Na2 SiO3 ) to represent the silicate rock. Small quantities of other compounds are introduced, and crystals are formed by precipitation or oxidation reduction reactions. A solution of one of the ions of the desired crystal is added to an acidified solution of water glass. This mixture is allowed to stand. When the other reactant is added to the top of the solution the ions diffuse downward, and the reaction takes place. The longer these crystals are permitted to grow, the more beautiful they become.



Dispose of solutions according to local codes.


  1. Preparation of acidified gel.
  2. Prepare a hot water bath by half filling a 400-mL beaker and placing it on a hot plate.
  3. Put 25 mL 1.0 M acetic acid in a test tube. (NOTE: The concentration of acetic acid and the sodium silicate solution are critical for gelling.)
  4. Add the specified amount of the first reactant (see following table) to the acetic acid in the test tube, and mix well. This may be done by stoppering the test tube and gently turning it upside down several times.
  5. Add 25mL sodium silicate solution to the acid mixture. Mix well as above, and cover with Parafilm (with a pinhole) or a stopper.
  6. The gel will set in minutes if put into very hot water.
  7. Add the metal or solution listed for specific systems (see following table), stopper, and display.
  8. Observe daily. 

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Demonstration 2: An Inorganic Polymerization Demonstration


Silicates polymerize in acid solution according to the following equation:

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Although the product is represented by the formula Si(OH)4 , it consists of a complex mixture of polymeric acids formed by condensation.

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The Si(OH)4 cannot be isolated from the mixture. If the gelatinous mixture is heated, the reaction is reversible. If the NaCl(aq) is removed by washing, the reaction is not reversible. On heating, the unwashed silica becomes a solid, and the washed product is a powder that can absorb a large amount of water.



  1. Add 100 mL 1 M HCl to 100 mL 17% Na2SiO3 in a weighed 250-mL beaker. Decant supernatant liquid.
  2. Find the mass of the gel. Divide the gel into 2 equal parts.
  3. Place half in a weighed Petri dish, and weigh again.
  4. Wash the remaining gel in the beaker with distilled water until no cloudiness appears when a few drops of AgNO3 is added to the decanted wash water.
  5. Heat both gels in a drying oven at 105 C for 1.5 hr or in a microwave oven for 15 min.
  6. Reweigh the samples.

Data Analysis

  1. Describe the differences in the two samples.
  2. Students can calculate the number of grams of water absorbed by one gram of gel for each sample.
  3. Explain the differences in structure that can account for the results in Question 2.
  4. Which of the samples would be most likely to be packed with cameras and electronic equipment for use as a drying agent?

Demonstration 3: Supersaturated Solution


The action of heat and pressure on water beneath the surface of the earth sometimes creates supersaturated solutions. This demonstration shows how such solutions could produce crystals.



  1. Weigh 160 g of sodium acetate trihydrate in a 500-mL Erlenmeyer flask.
  2. Add 30 mL distilled water to the flask.
  3. Heat the mixture in a hot water bath stirring occasionally until all of the solid is dissolved. (This may take 15 min or so.)
  4. Remove all crystals from the sides of the flask by rinsing them down with small squirts of water from the wash bottle.
  5. Cover the flask with Parafilm or the inverted 100-mL beaker.
  6. Allow the solution to cool to room temperature undisturbed, or to speed up the cooling process, run cold water over the sides of the flask making sure no tap water gets into the flask and contaminates the solution.
  7. While holding a single crystal of sodium acetate over the open mouth of the flask, snap your fingers, and allow the crystal to drop into the flask.   The single crystal should start crystallization.
  8. After crystallization is complete turn the flask upside down, and nothing should fall out.
  9. The sides of the flask should be warm since this is an exothermic process.
  10. The solution may be used over and over again by reheating it to redissolve the sodium acetate.


The addition of too much water will result in leftover liquid after recrystallization.  Variations of this demonstration include placing a single crystal of sodium acetate in a shallow container and pouring the solution described above on the crystal to produce a "stalagmite" of sodium acetate. A buret may also be used to add the solution to the crystal.

Demonstration 4: Precipitation and Redissolution of Calcium Carbonate


CaCO3 is the major chemical component of many minerals and natural products. These minerals include marble, calcite, aragonite, coral, sea shells, stalactites, stalagmites. Precipitation of CaCO3 forms these substances.  Dissolution of CaCO3 allows the raw materials to be transported.

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This demonstration illustrates how sensitive the direction of this equilibrium expression is to reaction conditions (pH, carbonate concentration, temperature).



The Ca(OH)2 solution is basic with a pH = 12.4. It can irritate eyes and skin.   In case of ingestion, give copious drinks of water to dilute the limewater in the stomach, and seek medical advice.


  1. Place 150-250 mL of clear, saturated Ca(OH)2 in a 500-mL beaker. Have a volunteer blow gently through a straw or glass tube into the solution.  Caution: because Ca(OH)2 is a strong base, warn the volunteer to not get any of the liquid into the mouth (see Safety above). A white precipitate of CaCO3 will form.
  2. Place a second sample of Ca(OH)2 in a second beaker. Bubble pure CO2 through the solution by dropping in a piece of dry ice or by using the compressed gas. Initially, a white precipitate of CaCO3 will form. However, the precipitate will slowly dissolve eventually giving a clear solution.
  3. Have your volunteer return to try to redissolve the first sample of CaCO3 by blowing again into the sample. No amount of blowing will redissolve the precipitate.
  4. Put the solution from Step 2 on a hot plate with stirring. The carbon dioxide gas is less soluble in hot water and is removed. Near the boiling temperature, the solid begins to reform. When the solution cools, the appearance of suspended solid will be quite pronounced. (Adding pure CO2 to this final solution will not redissolve the CaCO3 .)


Step 1.  Although H2CO3 is not stable, it is useful to view CO2 dissolving inwater to form carbonic acid.

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The acid is immediately neutralized by hydroxide forming carbonate ion that in turn reacts with the calcium ion forming insoluble CaCO3 .

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Step 2. When pure CO2 is being added, its concentration increases until all the hydroxide originally in the solution is consumed. Thus, any excess dissolved CO2 is unneutralized carbonic acid. Some hydrogen ions dissociate and dissolve the precipitate converting the carbonate to bicarbonate. (CaHCO3 is soluble.)

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Step 3. Exhaled breath is only 4% CO 2 . By contrast, the gas obtained from dry ice is 100% CO 2. Blowing into the solution never provides a concentration of carbonic acid high enough to exceed the available hydroxide. Think of all the reactions as reversible. Thus, in each, the reactants and products are in a state of balance. When the pressure of CO 2 (g) in reaction (a) cannot get very high, reaction (b) cannot be forced to consume all the OH . Thus the effect of the last two reactions cannot be seen. That is, the equilibrium in reaction (d) and (e) lies far to the left.

Step 4. Heating the solution decreases the solubility of CO 2 in water. Thus reaction (a) is forced to the left. This in turn forces reactions (d) then (e) to the left and CaCO 3 reforms.

Demonstration 5: Pot-O-Gold


This demonstration should precede a discussion on solubility. Solubility and precipitation are important in understanding the occurrence of minerals. This demonstration catches the attention of students who watch the platelets swirling and catching the light.




Both lead and iodide ions are listed as hazards. Students should not handle these chemicals. Normal precautions should be followed. Excess solid lead iodide should be stored in a lead-waste container. The filtrate is used as part of the second solution.


  1. Mix the lead nitrate and sodium iodide solutions. A yellow precipitate of lead iodide will form.
  2. Filter the supernatant liquid into 1-L beaker and add enough distilled water for a total volume of 300 mL.
  3. Heat solution containing the filtrate to boiling.
  4. Add enough solid lead iodide to make a saturated solution at 100C.
  5. Pour the saturated solution, while hot, into a 500-mL Florence flask.
  6. Let the solution cool slowly to room temperature. The lead iodide will precipitate out as shiny platelets that look like gold.
  7. Stopper the flask securely. The mixture can be kept in a stoppered flask for several years.