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Name: Erica
Status: educator
Age: 6-8
Location: CA

Country: N/A
Date: 11/30/2005


Question:
I was going to teach my class about precipitates. While explaining the definition, I wanted to do a demonstration so they can actually see the precipitate. These are the chemicals my new school has in stock: Potassium nitrate, Strontium chloride, potassium chloride, calcium chloride, barium chloride,and calcium carbonate. Are any of them good for a precipitate demonstration? Can you advise as to how to go about this?


Replies:
There are two issues here. The first is safety. Strontium, and Barium can be toxic if ingested so cautious handling and supervision are necessary. Nitrates can also be toxic but to a lesser extent. An interesting precipitation reaction is to take a solution of freshly prepared calcium chloride and have the students exhale, that is, blow bubbles through a soda straw immersed in the calcium chloride solution. Carbon dioxide exhaled from the lungs forms calcium carbonate which causes the solution to turn cloudy. Keep the stock solution of CaCl2 protected from the atmosphere because it will absorb CO2 from the atmosphere. Strontium and barium chloride will undergo the same precipitation reaction, but I would be cautious about handling these salts by 6-8th graders without careful supervision. You could compare the results of exhaling through a solution of Ca(Cl)2 and Mg(SO4) [Epsom's salt] which is available at any pharmacy and most grocery stores. The solubility product constant of Mg(CO3) and Ca(CO3) is 6x10^-6 and 3x10^-9, so the difference in the amount of breath necessary to precipitate the two solutions should be observable, although I have not done the experiment.

Another demonstration could be done with Ca(CO3). While it is only very slightly soluble in water, you could start with a slurry of the solid and add HCl. This will cause the precipitate to dissolve, and the solution to become clear. Then, raise the pH with KOH (or NaOH). As the solution becomes more alkaline, Ca(OH)2 will precipitate because its solubility product constant is about 5x10^-6.

Vince Calder


Erica,

Unfortunately, all the combinations of the chemicals you listed are soluble in water (and therefore do not react to form a precipitate). The easiest, brightest precipitation reaction I can think of that makes for a brilliant demonstration is: KCl(aq) + Pb(NO3)2(aq) --> KNO3(aq) + PbCl2 (s).

Hope this helps.
--Michelle W.


Erica, The only way you can form a precipitate is if the cation and anion combination have a very low solubility product. Unfortunately, most nitrates and carbonates are soluble, and of the chlorides only silver chloride and lead chloride are insoluble. You may have to get access to some silver nitrate or lead(II) nitrate - which when combined with the chlorides should form a precipitate. Alternatively, barium hydroxide and calcium hydroxide are only sparingly soluble - so if you can get a hold of sodium or potassium hydroxide - that would work too.

A possible alternative is to use commercial products. I'm not an expert on products that you can get at your local supermarket . . . hopefully some other respondent can point you in that direction.

Greg (Roberto Gregorius)


Sodium and Potassium (column 1 on the periodic table), and chloride and nitrate, are all renowned as highly soluble ions. They are unlikely to make any quick precipitates for you.

You have 3 metals from column 2 in the periodic table: calcium, strontium, and barium. Column 2 metals make many more precipitates and insoluble salts. Calcium carbonate is such an insoluble salt. But that is the end product, not your starting chemical. Your demonstration needs to start from two soluble salts. When mixed, one member from each soluble salt will find each other and form an insoluble salt which precipitates itself.

Strontium and barium are too probably toxic to be using in bulk in a class that young. Or by yourself, if you are this unfamiliar with chemistry. That leaves you calcium, which is fine. You need to start with your calcium chloride, which is moderately soluble. Now you need an anion which precipitates with calcium. As I implied above, carbonate is good. So I would try buying some baking soda , make a solution of that and a separate solution of calcium chloride, and see if they make calcium carbonate precipitate when mixed. (Please read to the end here before doing this.)

The reaction formula might be expressed this way:

calcium chloride, dissolved, + sodium bicarbonate, dissolved -> calcium carbonate solid + sodium chloride, dissolved + hydrochloric acid, dissolved.

[Ca(2+) & Cl-] + [Na+ & H+ & CO3(2-)] -> CaCO3 & [Na+ & Cl-] & [H+ & Cl-]

With that hydrochloric acid the resulting water could be a fairly strong acid, so rinse within a minute if you get a drop on your hands. The acid will also react with some baking soda and make a very strong CO2 fizz or erupting froth, so be careful, practice before class with small amounts in a sink, and dilute your starting solutions until the fizz is manageably weaker but a lot of white powder still forms. You need extra baking soda to use up the HCl acid by fizzing:

[H+ & Cl-] + [Na+ & H+ & CO3(2-)] --> CO2(gas) & H20(liquid) & [Na+ & Cl-]

so dilute your calcium chloride solution in preference to the baking soda solution.

Separating the precipitate from the liquid, to show it off, is a larger part of the experiment than you might think. The powder which forms can be so fine that it almost never settles to the bottom. When the precipitation happens fast, in a panic, zillions of new crystals are formed and they need not get very big to absorb all the precipitating substance. Tiny particles can stay suspended in water a long time. When the precipitation happens slowly, fewer new crystals form and those simply grow bigger to absorb all the precipitate which is forming. Bigger crystals can sink much faster, are easier to filter out with coarse paper, and easier to drain water from. So pouring your starting solutions together slowly and steadily is more likely to give you a manageably coarse precipitate. Heating or lightly boiling the liquid with precipitate can help change milky ultra-fine silt into crystallites which settle to the bottom well. Days of waiting-time at room temperature also help, though perhaps a little less.

Sorry I have not figured out the concentrations for you yet. You might want to look up the molecular weights of CaCl2 and NaHCO3. Personally I tend to make saturated solutions of each salt (with leftover powder undissolved on the bottom), then pour out some top-liquid and dilute it by a factor of 2 or 10 or whatever. Have a marker to label such bottles. Then you never forget what they were. My first try would be with baking soda 1/2 of saturated, and calcium chloride 1/10th of saturated. That should be reasonably tame, but you might wear goggles in case it splashes upwards a little.

I think the calcium chloride solution always tries to be slightly milky, not quite clear. It can grab CO2 from the atmosphere to make a fine haze of calcium carbonate precipitate within itself. You will probably want to keep a cap on the bottle as much as possible. Occasionally adding a drop of HCl solution might help.

Jim Swenson



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