Round-Trip Copper Reactions Lab Report
AP Chemistry
12/13/11
Round-Trip Copper Reactions Lab
The purpose of this lab was to evaluate our skills of decanting a supernatant liquid without losing the solid and successful completion of a series of reactions. This was done through five chemical reactions involving copper. In this lab, elemental copper was put through five different chemical reactions in order to convert it into different compounds. By the end of the fifth reaction, the copper was back to its elemental state. In the first reaction, 0.95 g of pure copper was reacted with 4.0 mL of concentrated nitric acid under the fume hood. The solution was swirled until all of the copper had dissolved. The balanced equation for this reaction is as follows:
Cu (s) + 4HNO3 (aq) à Cu(NO3)2 (aq) + 2NO2 (g) + 2H2O (l)
When the nitric acid came into contact with the solid copper, a brown gas was immediately released into the fume hood. This brown gas was nitrogen dioxide as the balanced equation above shows. As the copper dissolved, the solution slowly turned blue because of the copper ions. In the second reaction, distilled water was added to the copper (II) nitrate solution until the beaker was half full. Then, 30 mL of 3.0 M sodium hydroxide was added to the solution. The balanced equation for this reaction is as follows:
Cu(NO3)2 (aq) + 2NaOH (aq) à Cu(OH)2 (s) + 2NaNO3 (aq)
When sodium hydroxide was added to the copper (II) nitrate solution, a bright blue gel-like precipitate was formed instantly. This precipitate was copper (II) hydroxide. The precipitate was blue because of the copper ions. In the third reaction, the copper (II) hydroxide solution was heated above a Bunsen burner. Continual stirring was required to reduce the “bumping”, or formation of bubbles that release gas very quickly that have the potential to cause injury. The balanced equation for this reaction is as follows:
Cu(OH)2 (s) à CuO (s) + H2O (l)
When the solution was heated, the bright blue precipitate
12/13/11
Round-Trip Copper Reactions Lab
The purpose of this lab was to evaluate our skills of decanting a supernatant liquid without losing the solid and successful completion of a series of reactions. This was done through five chemical reactions involving copper. In this lab, elemental copper was put through five different chemical reactions in order to convert it into different compounds. By the end of the fifth reaction, the copper was back to its elemental state. In the first reaction, 0.95 g of pure copper was reacted with 4.0 mL of concentrated nitric acid under the fume hood. The solution was swirled until all of the copper had dissolved. The balanced equation for this reaction is as follows:
Cu (s) + 4HNO3 (aq) à Cu(NO3)2 (aq) + 2NO2 (g) + 2H2O (l)
When the nitric acid came into contact with the solid copper, a brown gas was immediately released into the fume hood. This brown gas was nitrogen dioxide as the balanced equation above shows. As the copper dissolved, the solution slowly turned blue because of the copper ions. In the second reaction, distilled water was added to the copper (II) nitrate solution until the beaker was half full. Then, 30 mL of 3.0 M sodium hydroxide was added to the solution. The balanced equation for this reaction is as follows:
Cu(NO3)2 (aq) + 2NaOH (aq) à Cu(OH)2 (s) + 2NaNO3 (aq)
When sodium hydroxide was added to the copper (II) nitrate solution, a bright blue gel-like precipitate was formed instantly. This precipitate was copper (II) hydroxide. The precipitate was blue because of the copper ions. In the third reaction, the copper (II) hydroxide solution was heated above a Bunsen burner. Continual stirring was required to reduce the “bumping”, or formation of bubbles that release gas very quickly that have the potential to cause injury. The balanced equation for this reaction is as follows:
Cu(OH)2 (s) à CuO (s) + H2O (l)
When the solution was heated, the bright blue precipitate