Purpose. The goal of this experiment is to introduce you to several classes of chemical reactions: oxidation/reduction, precipitation, decomposition and acid/base neutralization. The reactions will be performed in a specific order and, when completed, will regenerate the starting material, pure copper. Since you will not add copper to the reaction at any point after the initial step, it should be possible to recover all of the copper at the end of the cycle, thereby illustrating the Law of Conservation of Mass.

Historical note. These reactions are similar to a series of experiments performed by an Italian chemist named Angelo Sala in 1617. A widely held belief among alchemists was that minerals "fermented" over time and would ultimately "mature" into gold in a process known as transmutation. In other words, minerals were believed to have life cycles in much the same way that plants and animals do. These ideas had their roots in Aristotelean natural philosophy. Sala, a Calvinist and a follower of Paracelsus (see Carolyn Merchant's discussion in The Death of Nature), demonstrated that the conversion of metals to minerals, and vice versa, was a simple chemical process and that no transmutation (conversion of one metal into another) takes place. That is, the identity of the copper is unchanged in the metallic or mineral (copper sulfate, or blue vitriol). He proposed instead that the different properties of the new materials are due to different combinations of their component particles. This sequence of reactions clearly demonstrated the inadequacy of Aristotle's ideas to explain chemical phenomena.

Introduction. Of the thousands upon thousands of individual chemical reactions known, most of them can be classified into a very small number of reaction types - this is one way that chemists can gain understanding from what would otherwise be an enormous volume of individual observations. In this experiment, several different types of reactions will be performed to transform copper into a variety of its compounds. It is important to realize that each product contains copper and that the total number of copper atoms involved in each step is the same. Therefore, at the end of the cycle, the mass of copper recovered should equal the mass that was originally used. You will calculate the percent yield of copper at the end of the cycle to determine the "efficiency" of the process.

The cycle of reactions to be performed is shown in Figure 1. Beginning with pure copper at the top of the figure, these are:

Reaction 1: Oxidation of metallic copper with nitric acid (HNO₃). The charge on the copper changes from zero to +2, that is, the reaction produces copper ions with a charge of plus two. So where do the electrons go? This reaction is more complicated than initially appears. The balanced equation is:

To determine which species is reduced requires the use of formal oxidation states, a topic that we will not address in this introduction to chemistry. Suffice it to say that, in the balanced equation, three copper atoms lose a total of six electrons (two each). Four of these go to two oxygen atoms (two each) and the remaining two go to two nitrogen atoms (one each). Thus copper is oxidized and both oxygen and nitrogen are reduced.

Reaction 3. Thermal decomposition of copper (II) hydroxide to copper (II) oxide. Compounds that are stable under standard conditions often become unstable at elevated temperatures. Many times this can result in the loss of gases, for example, the loss of carbon dioxide from calcium carbonate (see page 226 in The Periodic Table by Primo Levi).

When copper (II) hydroxide is heated to roughly 100° C, it decomposes to copper (II) oxide according to the following reaction.

Reaction 4. Reaction of copper (II) oxide with sulfuric acid. Copper (II) oxide, like all metal oxides, is basic. That is, when placed in water, metal oxides will increase the hydroxide ion concentration of the solution. This is true even if the compound has a very low solubility. Because of this basic nature, metal oxides will undergo neutralization reactions with acids to produce water and a "salt". In this case the copper oxide is reacted with sulfuric acid, thus the salt is the ionic compound copper (II) sulfate.

Reaction 5. Reduction of copper (II) with zinc. In the final step, each copper (II) ion takes two electrons from a zinc atom, thereby regenerating the starting material, metallic copper..

After the reaction is complete, the excess zinc must be removed from the copper/zinc mixture in order to purify the copper and accurately measure its mass. To accomplish this, the mixture is treated with hydrochloric acid. The hydrogen ion (H⁺) will oxidize the zinc to form the soluble ZnCl₂ and hydrogen gas. At this point the solid, pure copper can be isolated, dried and weighed.
 
 

Experimental Procedure.

Safety Notes.

Reaction 2. While stirring with a glass stirring rod, add approximately 15 mL of 6 M NaOH. Record your observations in your notebook. Dilute the solution with deionized water to about 100 mL in preparation for step 3.

Reaction 3. Add a magnetic stir bar and place on a heatable stir plate. Boil gently while stirring for about 4 minutes. Record any changes that occur in your notebook. Remove the beaker from the heat and allow to cool. Add 40 mL of deionized water into a second clean beaker and begin heating.

Prepare a filter paper and funnel to filter the copper (II) oxide. Use a 250 mL beaker to collect the filtrate, supporting the funnel with a funnel support or iron ring on a ringstand. Filter the copper (II) oxide. The filtrate should be colorless and free of any solids. Transfer the last traces of solid material from the beaker using a stream of deionized water from a wash bottle. Use the deionized water that you have been heating to wash the solid collected in the filter paper. Pour about 5 mL of the hot water through the filter. Repeat three times. Leave the filtration apparatus in place for the next step. In your notebook, describe the appearance of the collected solid.

Reaction 4. Dissolve the CuO by carefully pouring about 15 mL of 3 M sulfuric acid directly through the residue on the filter into a 250 mL beaker. Record any changes that you see. If the solid is not completely dissolved the first time, replace the collection beaker with a clean new one and pour the acid in the first beaker through the filter again. Pour very carefully so as not to lose any of the liquid. Repeat this procedure as often as necessary to dissolve all the solid. It may take four or five times.

Once the solid is dissolved, you need to collect all of the copper (II) containing solutions in the same beaker. Rinse down the walls of the collection beaker that is not in position below the filter with deionized water from a wash bottle and pour the rinse water through the filter into the other beaker. Wash the empty filter paper with three or four 5 mL portions of cold deionized water and collect the washings in the beaker containing the acid solution.

Reaction 5. Add about 2 g of zinc metal to the copper (II) solution and stir rapidly. Hydrogen gas bubbles should appear – keep flames away! Record your observations in your notebook. The reaction between the zinc and copper (II) ion will be complete when the blue color of the copper solution is gone. If any blue color remains after the zinc has been consumed, add approximately 0.5g of additional zinc. Record this in your notebook.

When the copper has been completely reduced, decant most of the solution, even if it is still reacting with the zinc. Add 25 mL of 3 M HCl to speed up the rate of zinc oxidation. When no more bubbles are seen, proceed to the epilogue. Record your observations.

Epilogue. Allow the copper metal to settle to the bottom of the beaker. Carefully decant the supernatant. Do not try to pour off all of the liquid – it is better to leave a small amount of liquid than to lose any precipitate. Wash the copper metal precipitate with three 50 mL portions of water, removing each portion by decantation. Weigh a clean, dry evaporating dish and record the mass. Transfer all of the copper precipitate into the evaporating dish (use a rubber policeman – in good condition! – and a wash bottle to help with this). Let the solid settle in the evaporating dish and carefully decant most of the liquid. Set the evaporating dish in a 110° C oven for 30 minutes to evaporate the water. Remove from the oven and allow to cool to room temperature(place on a paper towel with your name on it). Weigh the evaporating dish and calculate the mass of copper recovered and your percent yield using the following equation:

Report Sheet Name: ______________________

Observations. Describe what happens physically in each of the following reactions. Be specific with regard to color changes, gases formed, etc.

Reaction 1. Oxidation of copper with nitric acid.
 
 
 
 
 
 
 
 

Reaction 2. Precipitation of copper (II) hydroxide with sodium hydroxide.
 
 
 
 
 
 
 
 

Reaction 3. Decomposition of copper (II) hydroxide to copper (II) oxide.
 
 
 
 
 
 
 
 

Reaction 4. Reaction of copper (II) oxide with sulfuric acid.
 
 
 
 
 
 
 
 

Reaction 5.

a) Reduction of copper (II) with zinc
 
 
 
 
 

b) Oxidation of excess zinc with hydrochloric acid.
 
 

Experimental Data:

Mass of copper foil ______________

Mass of copper recovered ______________

Percent Recovery ______________