What happens when copper reacts with nitric acid in this lab?

Solid copper is oxidized by nitric acid to form copper nitrate, water, and NO2 gas. The copper nitrate dissociates into copper(2+) ions and nitrate anions in water. The reddish-brown fumes confirm NO2 gas production, while a greenish-blue tint indicates the presence of copper ions in solution.

How is copper oxide formed from copper(II) hydroxide?

Copper(II) hydroxide precipitate forms when a base is added to the copper solution, creating a blue color. When heated to 85°C, the copper(II) hydroxide undergoes decomposition, removing water and producing black copper oxide precipitate. This thermal decomposition converts the unstable hydroxide into the more stable oxide form.

What role does zinc play in recovering solid copper from copper oxide?

Zinc acts as a reducing agent in a single displacement reaction with copper sulfate solution. When zinc is added, copper(2+) ions are reduced to solid copper while zinc is oxidized to zinc(2+) ions. The solution becomes colorless once all copper ions are reduced, indicating complete copper recovery.

Why is hydrochloric acid added after the zinc displacement reaction?

Hydrochloric acid dissolves excess zinc remaining in the solid copper precipitate through a neutralization reaction that produces bubbling. Adding acid in increments ensures complete removal of unreacted zinc without over-dissolving the copper product, leaving pure copper for final collection and weighing.

How do you determine if the final copper product is pure?

Test for methanol using an indicating desiccant that changes from blue to turquoise if present. Drop 6 M HCl on the product; bubbling indicates residual NaOH contamination. Pure copper appears as dark-red granules. Black color suggests incomplete conversion and the presence of copper oxide instead.

What does percent yield calculation reveal about the copper recovery experiment?

Percent yield is calculated by dividing the final copper mass by the initial copper mass and multiplying by 100. Yields exceeding 100% indicate calculation errors or residual moisture on the product. This metric shows how efficiently copper was recovered through the oxidation and reduction sequence.

Why is methanol used in the final drying step instead of water?

Methanol is used because it evaporates more readily than water at room temperature, ensuring thorough drying of the copper precipitate without heat damage. Multiple methanol washes followed by extended vacuum application remove residual moisture and impurities, preparing the copper for accurate final mass measurement.

Redox Reactions

Concepts Instructor Prep Student Protocol

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Redox, which is shorthand for reduction-oxidation, is a type of chemical reaction classified by the net transfer of electrons. In this reaction, one molecule loses electrons, called oxidation, and the other molecule gains electrons, called reduction.

To help you differentiate between the two, remember the phrase 'OIL-RIG', which stands for 'oxidation is losing, reduction is gaining'. The molecule that is oxidized is called the reducing agent because it reduces the other reactant. Similarly, the molecule that is reduced is called the oxidizing agent because it oxidizes the other molecule.

Now that we have the terminology sorted out, let's look at an example of a redox reaction, the formation of the mineral magnesium oxide. During the reaction, each magnesium atom loses two electrons. Thus, magnesium is oxidized. Each oxygen atom gains two electrons; thus, oxygen is reduced.

However, not all reactions are redox reactions. For example, the reaction of calcium oxide with carbon dioxide to form calcium carbonate is not a redox reaction. So, how can we identify a redox reaction?

To do this, we track the oxidation number of each element as it goes from reactant to product. The oxidation number is the hypothetical charge that an atom would have if its bonds to different elements were ionic, meaning that the electrons are assigned to the more electronegative atom. The sum of the oxidation numbers in a molecule equals its overall charge.

Let's look back at magnesium oxide. It is a neutral compound, so the sum of the oxidation numbers for magnesium and oxygen equals zero. Magnesium can give two electrons, so its oxidation number is plus two. Oxygen can accept two electrons, so its oxidation number is minus two.

How about the reaction? Pure neutral elemental compounds have an oxidation number of zero. Thus, both magnesium and oxygen start with oxidation numbers of zero. Both magnesium and oxygen's oxidation numbers changed during the reaction, so this is a redox reaction.

Now, let's look at the calcium carbonate reaction we saw earlier. Both reactants are neutral, so the sum of the oxidation numbers for both compounds is zero. As we saw with magnesium oxide, calcium has an oxidation number of plus two and oxygen minus two. Then, the carbon in the carbon dioxide molecule has an oxidation number of plus four and each oxygen minus two.

How about the product? Calcium is plus two and carbon plus four, just like in the reactants. Each oxygen is minus two, totaling minus six, with the net oxidation number zero. Since none of the oxidation numbers have changed, this is not a redox reaction.

Now let's introduce the four types of redox reactions. The first is a single displacement reaction, where one atom displaces another. You'll see this in a thermite reaction, where one metal is reduced, and the other metal is oxidized.

The next type is a combustion reaction, which occurs between a fuel and oxidant to form oxidized products and heat. You see this in the lab during the combustion of methane with oxygen when using a Bunsen burner.

The third is a synthesis reaction, where two reactants combine to form one product, like in the synthesis of ammonia, where nitrogen is combined with hydrogen to form ammonia.

Finally, the fourth type is a decomposition reaction, where a reactant absorbs enough energy to break its bonds to form smaller compounds. This is what happens with fireworks, where potassium chlorate decomposes to potassium chloride and oxygen after heating.

In this lab, you'll perform and identify various types of redox reactions that transform solid copper to copper oxide, and then back again to solid copper.

Print Procedure Steps

Source: Smaa Koraym at Johns Hopkins University, MD, USA

Making Solid Copper
In the first half of this lab, you will make solid copper. In this reaction, the solid copper is first oxidized by the strong oxidizing agent, nitric acid, to form water and the salt copper nitrate. The copper nitrate quickly dissociates into the ions, copper(2+) and nitrate, in water. The nitrate anions are reduced to NO₂ gas, which is confirmed by the presence of reddish-brown fumes. The presence of copper ions is indicated by a greenish-blue tint. Adding a base to the mixture changes the color to blue as the copper(II) hydroxide precipitate is formed. The copper(II) hydroxide is relatively unstable; thus, when heat is introduced, a decomposition reaction occurs. In this reaction, water is removed to produce copper oxide, which appears as a black precipitate.
- Before starting the experiment, put on personal protective equipment, including a lab coat, gloves, and chemical splash goggles.
- Weigh between 0.45 and 0.55 g of copper shot and record the exact mass. This experiment must be conducted in the hood, so bring the copper shot back to your workstation.
  Table 1. Determining percent yield
  
  Mass_{Cu initial}(g)
  Mass_{Cu final} (g)
  Percent yield
  Click Here to download Table 1
- Label a 10-mL graduated cylinder as ‘8 M nitric acid’ and measure 8 mL of the solution provided by your instructor. Cover the cylinder with a watch glass when transporting the solution back to your hood.
- To begin the experiment, place the copper in a 50-mL beaker.
- Slowly add the 8 mL of 8 M nitric acid while stirring with a stirring rod. Note: If the copper does not fully dissolve, place the beaker on a hotplate set to 30 °C for approximately 2 min, while continually stirring. Here, the solid copper is reacting with the nitric acid to form copper nitrate, water, and NO₂ gas. Do not overheat the solution or the nitric acid will boil off.
- Once the copper is fully dissolved, remove the beaker from the hotplate and set the temperature on the hotplate to 85 °C.
- Measure 150 mL of deionized water using a 50-mL graduated cylinder and transfer the water into a 400-mL beaker.
- Slowly, and while stirring, pour the copper solution into the 400-mL beaker containing water.
  Note: If there is copper solution left in the beaker, use a pipette to rinse the rest of the solution into the larger beaker.
- Label a 50-mL graduated cylinder with '3 M NaOH’ and obtain 20 - 30 mL of 3 M NaOH from your instructor. Cover the graduated cylinder with a watch glass while carefully carrying it back to your hood.
- Add 10 mL of the NaOH to the 400-mL beaker while stirring.
- Test the pH of the solution using pH paper. Continue adding NaOH until the solution is basic or about pH 10. Here, the copper nitrate reacts with NaOH to form the blue precipitate, copper(II) hydroxide, and sodium nitrate.
- Place the 400-mL beaker on the hotplate and use a glass thermometer to monitor the temperature. Once the solution reaches 85 °C, let it remain there for 10 min. Stir slowly with the stirring rod, as this will evenly distribute heat while not disrupting crystal formation.
- After 10 min, turn off the hotplate and remove the beaker. Allow the beaker to cool to room temperature. As the solution cools, a dark precipitate will appear at the bottom of the beaker, while a clear and colorless liquid, the supernatant, will separate towards the top. The crystals are the copper oxide.
- Once the solution has cooled completely, label a new beaker as 'aqueous waste', and remove the supernatant by carefully pouring it into the waste container. Try not to lose the precipitate at the bottom of the beaker.
- Set up the filtration system by attaching the silicone tubing between the house vacuum and the vacuum flask. Place a Büchner funnel into the vacuum flask using a rubber adapter. Now, place one filter paper in the Büchner funnel and wet it with deionized water.
- Turn on the vacuum. Pour the precipitate into the Büchner funnel and direct it into the center with the stirring rod. Use deionized water to rinse all the precipitate into the funnel. When all the precipitate is out of the beaker, turn off the vacuum.
- Remove the filter paper with the precipitate and place it into the beaker. The dark precipitant on the filter paper is the copper oxide, which should be left to dry for a day or two before it is used in the next reaction.
- Clean up from the experiment. First, empty the liquid in the filter flask into the aqueous waste beaker. Then, test the aqueous waste with pH paper. Note: Any waste that is higher than pH 10 should be neutralized with citric acid. Any waste below pH 4 should be neutralized with baking soda.
- Neutralize the unused NaOH with citric acid and then pour it down the drain with copious amounts of water. Anything containing black particles indicates the presence of copper oxide and should be disposed of in the heavy metal drum container.
Making Copper Oxide
The copper oxide produced in the previous lab section is insoluble in water but can be dissolved in sulfuric acid. When this occurs, the copper(2+) cations and the sulfate anions are reintroduced through a double displacement reaction. When zinc is added in the final step, the copper(2+) ions are reduced to solid copper through a single displacement reaction.
- Put on the appropriate protective equipment, such as a lab coat, gloves, and chemical splash goggles.
- Obtain the dried copper oxide that was prepared in the last lab exercise and remove as much copper oxide from the filter paper as possible using a rubber policeman. Place the solid and filter paper back into the same 400-mL beaker it was stored in.
- Label a 10-mL graduated cylinder as '5.93 M sulfuric acid' and obtain 8 mL of sulfuric acid from the instructor. Cover the graduated cylinder with a watch glass, and carefully carry the sulfuric acid to your hood. Note: Perform the next step of the experiment in the hood.
- Use a plastic pipette to add about 5 mL of sulfuric acid directly onto the filter paper in the beaker.
- Then, hold the filter paper with tweezers and use a metal spatula to scrape off as much solid as possible without ripping the filter paper. Use as little sulfuric acid as possible to rinse the copper oxide, as excess sulfuric acid will require more zinc later.
- Rinse copper oxide from the filter paper with 2 mL of deionized water and add a little more deionized water to the beaker. Once the copper oxide has been removed from the filter paper, dispose of the filter paper as solid waste.
- Stir the solution until the copper oxide is fully dissolved. If the solid does not dissolve, place the beaker on the hot plate and heat it to about 60 °C.
- Add 3 - 5 mL of water and stir to dissolve. Be careful not to go higher than this temperature, as the liquid could evaporate leaving a precipitate of copper sulfate.
- Weigh 1.2 g of zinc and record the exact mass in your lab notebook.
- Slowly add half of the zinc to the solution. Crush and stir the zinc using the stirring rod to fully dissolve all particles. The solution should be colorless. If color is still present, add more zinc until the color disappears.
- Once the color fully disappears, weigh the remaining zinc in the weigh pan, and subtract it from the original mass to obtain the mass of the zinc used. Here, the zinc has reacted with the copper sulfate to form zinc sulfate and solid copper.
- Decant the supernatant from your solution into the container labeled 'aqueous waste'. Then, preheat the hot plate to 40 °C. Obtain 10 mL of 6 M hydrochloric acid from your instructor in a labeled graduated cylinder. Cover the opening of the cylinder with a watch glass and bring the acid back to your hood.
- Add 5 mL of the hydrochloric acid to the remaining precipitate to dissolve any zinc present within the solid copper.
- Place the beaker on the preheated hot plate and warm the solution for 2 min. The bubbling observed here is the neutralization reaction taking place to dissolve the zinc. Make sure that the reaction does not boil. Then, turn the heat setting on the hot plate off.
- Once the bubbling subsides, add 2 - 3 mL of additional hydrochloric acid. Continue adding acid in 1-mL increments until the bubbling completely stops.
- Decant the supernatant, wash the precipitate with 5 - 10 mL of water, and decant any remaining supernatant. Note: If the copper particles are too small to decant the liquid without losing them, set up a Büchner funnel apparatus with filter paper and no vacuum. Pour the solution through the funnel so the particles are retained on the filter paper, then return the particles to the beaker. Repeat the wash and decant steps two more times.
- Replace the filter paper in the Büchner funnel with a fresh one and wet it slightly with water. Then, turn on the vacuum for a moment to help the filter settle into place.
- Add approximately 5 mL of water to the copper precipitate, and then pour it into the funnel in the center of the filter paper. Stir the copper gently to remove the impurities while taking care not to rip the filter paper.
- Turn the vacuum on, and let it run for 2 min.
- Close the vacuum and disconnect the tubing from the flask. Remove the funnel and dispose of the liquid in the aqueous waste container.
- Reconnect the Büchner funnel setup and reattach the vacuum tubing.
- Obtain about 10 mL of methanol and pour 3 mL onto the copper in the funnel. Then, turn on the vacuum and let the copper dry for 2 min.
- Dry the precipitate with methanol two more times, running the vacuum after each wash. During the last drying phase, place a large rubber stopper on the Büchner funnel, and let the vacuum run for 5 min. Then, close the vacuum and disconnect the vacuum tube.
- Remove the large stopper and let the copper air dry for 30 min.
- Label an empty beaker as 'organic waste' and pour the liquid from the filter flask into the organic waste container.
- Tare an empty weighing pan and transfer as much of the copper as possible from the filter paper to it. Weigh the copper and record its mass.
- Clean up from the lab exercise. Dispose of the organic waste and excess methanol in the organic waste container. Test the aqueous waste with pH paper and neutralize it with either citric acid for highly basic pH or baking soda for acidic pH.
- Dispose of the supernatant liquid in the heavy metal drum waste container because it contains zinc. The unused sulfuric acid and hydrochloric acid should be neutralized with baking soda and then washed with plenty of water down the sink.
Results
- Determine if the product is pure. Spread some of the product on an indicating desiccant to test for the presence of methanol. The desiccant will change from blue to turquoise if methanol is present.
- Drop a few drops of 6 M HCl onto the product to see if it will react with any NaOH that may be left over. If NaOH is present, you will see bubbling.
- If the product is copper, it should look like dark-red granules. If the product is black, you have failed to make copper and instead made copper oxide.
- Calculate the percent yield using the initial and final mass of copper. The percent yield is the ratio of the actual yield, in this case, the final mass of copper, to the theoretical yield, or the initial mass of the copper, multiplied by 100. It is impossible to obtain a percent yield greater than 100. If this occurs, there was an error in calculation — or the product was still wet — which added more weight.