Copper's Chemical Properties

Copper’s Chemical Properties

Topic Copper has unique properties that can be demonstrated in a series of chemical reactions.

Introduction Copper is a reddish-orange transition metal that is an excellent conductor of energy. Copper metal, which occurs commonly in nature, is widely used in electrical wiring, construction, plumbing, sculpture, and as currency. As element 29 on the Periodic Table, copper tends to form cations in chemical reactions. During oxidation, copper generally loses electrons and acquires a charge of 2, but it can also oxidize to form stable 1 compounds as well. Because of copper’s reactivity with other substances, it can undergo a variety of chemical changes. For instance, copper metal dissolves in nitric acid to form a liquid. The liquid can react with sodium hydroxide to yield copper (II) hydroxide. When the copper (II) hydroxide is heated, it forms copper (II) oxide, which can be dissolved in acid. Copper (II) ions can then be reduced with zinc metal to produce copper metal. In this experiment, you will perform a series of five reactions involving copper. At the conclusion of the reactions, you should end up with a yield of copper metal very similar to the mass of copper with which you started.

Time Required 40 minutes

Materials ✒ 250-milliliter (ml) beaker ✒ waste beaker

✒ stirring rod ✒ copper wire ✒ wire cutters

✒ 4 ml of concentrated (16M) HNO3 ✒ 20 ml of 3M NaOH

✒ 15 ml of 6M H2SO4 ✒ zinc metal (about 3 grams [g]) ✒ 20 ml of 6M HCl ✒ deionized water (about 200 ml) ✒ graduated cylinder ✒ electronic scale or triple-beam balance ✒ steel wool ✒ wire cutters ✒ hot plate ✒ rubber policeman ✒ watch glass ✒ pH paper ✒ science notebook

Safety Note Take care when working with Bunsen burners, open flames, and chemicals. Wear safety goggles and tie back long hair. Several steps of this experiment should be performed in a fume hood. Please review and follow the safety guidelines.

Procedure 1. Clean a piece of copper wire with steel wool. Cut a 4-to-5-inch (in.) piece of wire and coil it into a flattened spiral disc. 2. Use the electronic scale or triple-beam balance to find the mass of the copper metal. Record the mass on Data Table 2. 3. Place the copper metal into a clean, dry, 250-ml beaker.

4. Take the beaker of copper metal to the fume hood. Add about 4 ml of concentrated nitric acid to the beaker. Gently swirl the beaker until all of the copper dissolves. Record the appearance of the solution on Data Table 1. CAUTION: Take extreme care when working with concentrated acid. Wear goggles, apron, and protective gloves. Handle the beaker with tongs. 5. Add deionized water until there is 125 ml of liquid in the beaker. 6. Take the beaker back to your lab station. 7. Test the pH of the solution by touching a stirring rod from the solution to a piece of pH paper. Add 3M NaOH until the solution is basic when tested with pH paper, then add 2 more drops of NaOH. Record the appearance of the solution on Data Table 1. 8. Heat the beaker of solution on a hot plate. Stir the solution and allow it to boil gently until the precipitate formed turns a blackish brown color. Allow the beaker to cool to the touch 9. Carefully decant all of the liquid and pour into another beaker. Be careful not to lose any precipitate (CuO) (see Figure 1).

stirring rod

beakers

Decanting the liquid Figure 1

10. Rinse the solid precipitate with deionized water, allow the precipitate to settle, and decant the liquid. Note the appearance of the solid, and record it on Data Table 1.

11. Add 15 ml of 6M H2SO4 to the beaker containing the CuO precipitate. Stir with a stirring rod to mix the solid with the solution.

Record your observations on Data Table 1.

12. Take the beaker of CuO precipitate and H2SO4 to the fume hood. Add about 3.0 g of zinc metal to the solution in the beaker. In the fume hood, stir and heat (but do not boil) the solution until the zinc metal is dissolved. The liquid should be colorless. If it is blue, add a small amount of zinc to the solution until it turns clear. If silvery metal still exists, add a small amount of HCl to the solution until all of the zinc has reacted. 13. Take the beaker back to your lab station. Allow the copper to settle to the bottom of the beaker, then decant all of the liquid from the solution. 14. Rinse the copper with warm deionized water and decant again, then rinse a second time and decant the liquid. 15. Use the electronic scale or a triple-beam balance to find the mass of a watch glass. Record the mass on Data Table 2. 16. Using a rubber policeman, scrape all of the copper precipitate onto the watch glass that you just weighed. 17. Place the watch glass on top of a beaker of boiling water on a hot plate to dry the metal. Stop heating when the metal appears to be dry (see Figure 2).

copper precipitate

watch glass

beaker of boiling water hot plate

Figure 2

18. Find the mass of the copper metal and record it on Data Table 2.

Data Table 1

Procedure Product Observations

Nitric acid added

Sodium hydroxide added

Solution heated

Sulfuric acid added

Zinc metal added

Data Table 2

Beginning mass of copper

Mass of watch glass

Final mass of copper (minus watch glass)

Percentage yield (initial final/initial)

Analysis 1. Write the chemical equation for the process that occurred when the copper hydroxide was heated. 2. What was in the liquid that was decanted off of the solution in step 10? 3. Which gas is released when zinc was added to the solution of sulfuric acid and copper (II) oxide? 4. What was your percentage yield of copper after the series of reactions?

What’s Going On? In this experiment, copper was first dissolved in nitric acid, which caused the metal to ionize into copper (II) and release nitrogen dioxide. After the copper was ionized, sodium hydroxide was added to the solution. This step caused a neutralization reaction, until an excess of sodium hydroxide was added, binding the hydroxide ions to the copper ions to create copper (II) hydroxide. Heating this solution removed the water, and the copper (II) hydroxide was transformed into copper (II) oxide. When hydrochloric acid was added, copper (II) ions formed in solution once again; these ions bound to sulfate ions to create copper (II) sulfate. Zinc then reacted with copper (II) sulfate to form copper metal once again. Though ionized copper can react to form diverse products, the metal can be reduced back to its natural state. This is possible because of the unique properties of copper that enable it to be stable in its ionized state as well as in its form as a neutral metal. Because of these characteristics, copper does not corrode or degrade easily, making it an ideal metal not only for its conductive properties, but also for its durability.

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2.10 COPPER’S CHEMICAL PROPERTIES Suggestion for class discussion: Hold up a piece of copper metal and see how many students can identify it. Copper is well-known, and most students are familiar with its appearance. Ask them to name some of the uses of copper, then explain why copper has so many uses. On the Periodic Table, copper is in the family with silver and gold. Their electron structures make these elements stable, and all three metals have high thermal and electrical conductivity. Teacher notes: you may want to dispense the nitric acid instead of letting students pour it.

Analysis

1. Cu(OH)2(s) CuO(s) + H2O(l) 2. water 3. hydrogen gas 4. Answers will vary depending on individual results. High percent yields are expected.

Each experiment includes special safety precautions that are relevant to that particular project. These do not include all the basic safety precautions that are necessary whenever you are working on a scientific experiment. For this reason, it is absolutely necessary that you read and remain mindful of the General Safety Precautions that follow this note. Experimental science can be dangerous, and good laboratory procedure always includes following basic safety rules. Things can happen very quickly while you are performing an experiment. Materials can spill, break, or even catch fire. There will be no time after the fact to protect yourself. Always prepare for unexpected dangers by following the basic safety guidelines during the entire experiment, whether or not something seems dangerous to you at a given moment. We have been quite sparing in prescribing safety precautions for the individual experiments. For one reason, we want you to take very seriously every safety precaution that is printed in this book. If you see it written here, you can be sure that it is here because it is absolutely critical. Read the safety precautions here and at the beginning of each experiment before performing each lab activity. It is difficult to remember a long set of general rules. By rereading these general precautions every time you set up an experiment, you will be reminding yourself that lab safety is critically important. In addition, use your good judgment and pay close attention when performing potentially dangerous procedures. Just because the book does not say “Be careful with hot liquids” or “Don’t cut yourself with a knife” does not mean that you can be careless when boiling water or using knives. Notes in the text are special precautions to which you must pay special attention.

GENERAL SAFETY PRECAUTIONS Accidents caused by carelessness, haste, insufficient knowledge, or taking an unnecessary risk can be avoided by practicing safety procedures and being alert while conducting experiments. Be sure to check the individual experiments in this book for additional safety regulations and adult supervision requirements. Anytime you are working with an electrical current, it becomes possible to shock yourself on exposed wires. If you will be working in a lab, do not work alone. When

you are working off-site, keep in groups with a minimum of three students per group, and follow school rules and state legal requirements for the number of supervisors required. Ask an adult supervisor with basic training in first aid to carry a small first-aid kit. Make sure everyone knows where this person will be during the experiment.

PREPARING • Clear all surfaces before beginning experiments. • Read the instructions before you start. • Know the hazards of the experiments and anticipate dangers.

PROTECTING YOURSELF • Follow the directions step by step. • Do only one experiment at a time. • Locate exits, fire blanket and extinguisher, master gas and electricity shut-offs, eyewash, and first-aid kit. • Make sure there is adequate ventilation. • Do not horseplay. • Keep floor and workspace neat, clean, and dry. • Clean up spills immediately. • If glassware breaks, do not clean it up by yourself; ask for teacher assistance. • Tie back long hair. • Never eat, drink, or smoke in the laboratory or workspace. • Do not eat or drink any substances tested unless expressly permitted to do so by a knowledgeable adult.

USING EQUIPMENT WITH CARE • Set up apparatus far from the edge of the desk. • Use knives or other sharp, pointed instruments with care. • Pull plugs, not cards, when removing electrical plugs. • Clean glassware before and after use. • Check glassware for scratches, cracks, and sharp edges. • Let your teacher know about broken glassware immediately. • Do no use reflected sunlight to illuminate your microscope. • Do not touch metal conductors. • Use alcohol-filled thermometers, not mercury-filled thermometers.

USING CHEMICALS • Never taste or inhale chemicals • Label all bottles and apparatus containing chemicals • Read labels carefully. • Avoid chemical contact with skin and eyes (wear safety glasses, lab apron, and gloves). • Do not touch chemical solutions. • Wash hands before and after using solutions. • Wipe up spills thoroughly.

HEATING SUBSTANCES • Wear safety glasses, apron, and gloves when boiling water. • Keep your face away from test tubes and beakers. • Use test tubes, beakers, and other glassware made of Pyrex™ glass. • Never leave apparatus unattended. • Use safety tongs and heat-resistant gloves. • If your laboratory does not have heat-proof workbenches, put your Bunsen burner on a heat-proof mat before lighting it. • Take care when lighting your Bunsen burner; light it with the airhole closed, and use a Bunsen burner lighter rather than wooden matches. • Turn off hot plates, Bunsen burners, and gas when you are done. • Keep flammable substances away from flames and other sources of heat. • Have a fire extinguisher on hand.

FINISHING UP • Thoroughly clean your work area and any glassware used. • Wash your hands. • Be careful not to return chemicals or contaminated reagents to the wrong containers. • Do not dispose of materials in the sink unless instructed to do so. • Clean up all residues and put in proper containers for disposal. • Dispose of all chemicals according to all local, state, and federal laws.

BE SAFETY CONSCIOUS AT ALL TIMES!