This was by far the worsr practical probably because our teacher 'modelled' the entire thing, wrote up some equations and barely explained anything. And probably its being the last practical of the year we didnt give two craps at the time either. Big mistake come exam time though *groan* anyway, I did some research and here's basically what I came up with:
Risk Assesment: Ammonia is toxic by all routes of exposure. Concentrated ammonia is to be handled in the fume cupboard by the teacher only.
Procedure:
1. Dissolve 3.5-4.5g NaCl in 10mL distilled or deionised water in a small beaker. This is the brine solution.
2. The teacher will dissolve 5-6g ammonium chloride in 10mL concentrated ammonia in a small beaker. This is the ammonium saturator.
3. Transfer the brine solution into the saturated ammonia beaker and stir. If the solution is clear, add a small amount of ammonium chloride until some is left undissolved.
4. Decant the clear solution into a small conical flask. Bubble carbon dioxide gas through the solution from a gas cylinder. A white ppt. forms in 5-7 minutes.
5. Filter the ppt. through a Hirsh funnel using a vacuum filtration or if unavailable, decant the liquid into another beaker and scoop up the ppt.
6. Place some of the dry ppt. into a clean test tube. Add a few drops of water to dissolve and test 0.1M barium nitrate sltn for the presence of carbonate ions. A faint white ppt. confirms this.
7. The rest of the solid ppt. is placed into a Pyrex test tube and heated.
8. To regenerate ammonia from the filtrate, mix equal volumes of saturated Ca (OH)2 in a conical flask and heat. Hold a moist litmus paper to the mouth of the flask that should change from red to blue.
Difficulties:
1. Reaction is Lime kiln: CaCO3(s) >< CaO + CO2; Most schools do not have strong enough heat from Bunsen burners to decompose calcium carbonate. Without pure dry CO2 gas students will not be able to saturate brine to produce NaHCO3. In industry, such a process produces CO2 gas under steady pressure above 150kPa. The brine solution was not purified, a step that could not be done in the lab.
2. Reaction in slaker: CaO(s) + H2O >< Ca(OH)2(aq) + E; this be easily demonstrated by using very dry CaO which has been in the oven and cooled in a dedicator. Addition of water releases a lot of heat that can be conveyed using a thermometer.
3. Reaction in saturator: NaCl(aq) + NH3(g) >< Na+(aq) + Cl-(aq) + NH3(aq); Saturated ammonia cannot be produced in the lab due to regulations. Instead, saturating with NH4+ under pressure is used. While brine is added to the reaction vessel, the equilibrium established between all ions and the ammonia vapor.
4. Reaction is carbonator: NaCl(aq) + CO2(g) + NH3(g) +H2O >< NaHCO3(s) + NH2Cl (aq); By bubbling CO2 through the liquid in the conical flask, the equilibrium mixture is not like that of the industrial process as it is mostly NH4Cl that forms, demonstrated by heating the solid that represents the converter.
5. Reaction in converter: 2NaHCO3(s) >< Na2CO3(s) + CO2(g); in the lab, heating produces only traces of Na2CO3. In industry however, sodium carbonate is pure.
6. Ammonia recovery: 2NH4CL(aq) + Ca(OH)2(aq) ><CaCl2(aq) + 2NH3(g) + 2H2O(l); regeneration of ammonia can be done qualitatively in the lab. In industry, it is performed in a closed system.
Overall reaction:
2NaCl(aq) + CaCO3(s) ><Na2CO3(s) + CaCl2(aq)
Hope that helps. I think you should make notes in a tabular form, like column one: Process and how it was modelled, column 2: difficulties.
