The Сopper, or There and Back Again
Research a copper sulfate transformation chain!
- Put on protective gloves and eyewear.
- Conduct the experiment on the plastic tray.
- Observe safety precautions when working with boiling water.
- Do not allow chemicals to come into contact with the eyes or mouth.
- Keep young children, animals and those not wearing eye protection away from the experimental area.
- Store this experimental set out of reach of children under 12 years of age.
- Clean all equipment after use.
- Make sure that all containers are fully closed and properly stored after use.
- Ensure that all empty containers are disposed of properly.
- Do not use any equipment which has not been supplied with the set or recommended in the instructions for use.
- Do not replace foodstuffs in original container. Dispose of immediately.
- In case of eye contact: Wash out eye with plenty of water, holding eye open if necessary. Seek immediate medical advice.
- If swallowed: Wash out mouth with water, drink some fresh water. Do not induce vomiting. Seek immediate medical advice.
- In case of inhalation: Remove person to fresh air.
- In case of skin contact and burns: Wash affected area with plenty of water for at least 10 minutes.
- In case of doubt, seek medical advice without delay. Take the chemical and its container with you.
- In case of injury always seek medical advice.
- The incorrect use of chemicals can cause injury and damage to health. Only carry out those experiments which are listed in the instructions.
- This experimental set is for use only by children over 12 years.
- Because children’s abilities vary so much, even within age groups, supervising adults should exercise discretion as to which experiments are suitable and safe for them. The instructions should enable supervisors to assess any experiment to establish its suitability for a particular child.
- The supervising adult should discuss the warnings and safety information with the child or children before commencing the experiments. Particular attention should be paid to the safe handling of acids, alkalis and flammable liquids.
- The area surrounding the experiment should be kept clear of any obstructions and away from the storage of food. It should be well lit and ventilated and close to a water supply. A solid table with a heat resistant top should be provided
- Substances in non-reclosable packaging should be used up (completely) during the course of one experiment, i.e. after opening the package.
FAQ and troubleshooting
That’s a good question! You can learn about these formulas in the scientific description for this experiment (see the “Scientific description” section).
The hot water heats the liquid inside the vial and makes a precipitate (solid) inside the tube aggregate (clump together). This step makes it easier to filter and separate the precipitate.
Transfer all the solids and solutions from the filter paper, funnel, and flask into a clean beaker (use extra plastic cups if needed). Rinse the filter paper, funnel and flask with small volumes of water to make sure that you have recovered as much of the solid as possible.
Then, repeat steps 8–12.
The filtrate is no longer needed for this reaction chain and may be discarded down the sink. However, if you want an extra challenge, let some of the water evaporate from to obtain a little bit more copper(II) carbonate hydroxide.
If a generous amount of precipitate remains after you have added the ammonium carbonate solution to the vial, add just a little more ammonium carbonate solution dropwise (a maximum of 1 mL extra!).
Some precipitate may still remain after the addition of the extra ammonium carbonate solution, but this should not interfere with the outcome of this experiment.
When sodium hydrogen sulfate reacts with the copper tetra-amine complex gas evolves (is generated). If all of the sodium hydrogen sulfate were added at once, there would be a vigorous reaction and a lot of gas would be generated. This large volume of gas would probably cause the contents of the vial to bubble and spill out. By adding the sodium hydrogen sulfate dropwise, the rate of gas generation is controlled and the vial contents do not spill out.
- Take an empty vial and pour in two big spoons of 1M copper sulfate CuSO4 solution.
- Add two small spoons of sodium chloride NaCl.
- Pour in 2M sodium carbonate Na2CO3 solution from the bottle.
- Close the vial and shake it vigorously for 1 min.
- Secure the vial in the test tube holder.
- Place the vial in the beaker and pour in there freshly boiled water to completely cover the vial. Wait 15 min.
- Make a funnel out of a piece of filter paper, as shown. Fold the filter paper twice, insert your finger into the “pocket,” and open one layer of paper to the side to make a cone.
- Insert a plastic funnel into the flask and fit in the paper filter.
- Pour all the contents from the vial into the funnel.
- Wait for the liquid to drain through.
- Pour in some warm water into the funnel and let it filter through. Repeat this step two more times.
- Take a spoon and scoop the solids out of the funnel without ripping the filter. Collect these solids into a vial.
- Pour 3 big spoons of 2M ammonium carbonate (NH4)2CO3 solution into the vial.
- Close the vial with a cap and shake it vigorously.
- Now, slowly add all the 3M sodium bisulfate solution NaHSO4 from the bottle.
- You obtained initial copper sulfate solution!
Copper sulfate is transformed into a number of other copper compounds and back again during this experiment. The reaction starts with blue copper sulfate (CuSO4) then moves through green copper chloride ions ([CuCl4]2–), to turquoise copper(II) carbonate hydroxide (CuOH)2CO3, to black copper oxide (CuO), to a blue-black copper tetra-amine complex ([Cu(NH3)4]2+), and finally turns back into copper sulfate (CuSO4) again!
Dispose of solid waste together with household garbage. Pour solutions down the sink. Wash with an excess of water.
Why does CuSO4 solution changes its color from blue to green when NaCl is added?
In water, copper sulfate CuSO4 dissociates into ions:
CuSO4 ↔ Cu2+ + SO42-
Surrounded by water molecules, Cu2+ ions provide for blue color of solutions. If we add excess of NaCl into a solution, it dissociates into ions as well:
NaCl ↔ Na+ + Cl-
When copper ions react with chloride ions in water, they form [CuCl4]2– particles, which provide for the green color of the solution.
In this experiment, copper compounds undergo a whole chain of color changes: from blue to indigo-violet. Such transformations indicate that new chemical bonds are formed. In fact, ever-changing colors of these substances are a direct consequence of changes happening in quantity and nature of particles surrounding copper Cu2+ ions.
In an initial copper sulfate CuSO4 solution, each of copper Cu2+ ions has six neighboring water molecules. Even though there are plenty of H2O molecules around, only six of them can approach a copper ion closely enough to form weak chemical bonding. Such a stabilization of particles (in our case, Cu2+) by surrounding molecules of water or another solvent is called solvation, and the particles are regarded as solvated. In this experiment, copper ions are solvated by water. Though, in chemistry, water H2O is considered special, so chemists in this case would often say that particles are aquated (from the Latin aqua, meaning “water”).
At first, blue color is due to the presence of [Cu(H2O)6]2+ particles in the solution. Further, we add there an excess of NaCl, which dissociated into sodium Na+ and chloride Cl- ions. The latter can form relatively strong bonds with copper Cu2+ ions.
Soon enough, the solution becomes crowded with chloride ions, so that aquated copper ions start to bump into chloride ions. Most of these collisions lead to the formation of quite stable Cu–Cl bonding. Finally, copper and chloride ions form complex [CuCl4]2– particles:
[Cu(H2O)6]2+ + 4Cl– → [CuCl4]2– + 6H2O
These particles make the solution green.
By the way, the complex [CuCl4]2- can be obtained by other methods, for example, by oxidation of copper in presence of hydrochloric acid. For example, atmospheric oxygen or 30% hydrogen peroxide can be used as oxidizers, you can observe this in the following video. However, the reaction between copper sulfate and sodium chloride is the safest.
Why does adding sodium carbonate yield a precipitate?
Adding sodium carbonate Na2CO3 to a solution containing [CuCl4]2– results in formation of “basic” copper carbonate (CuOH)2CO3, a blue-green substance that is poorly soluble in water.
Carbonates are derivatives of carbonic acid H2CO3. Only few of them – such as sodium and potassium carbonates – easily dissolve in water. Most carbonates either decompose in water or do not dissolve in it at all.
The case in our experiment is not an exception. Turquoise-blue precipitate formed in our solution is one of copper carbonates. Interestingly, if we were to determine composition of this precipitate, we would have found not only carbonate CO32– and copper Cu2+ ions, which are responsible for the color of the solution, but also OH– ions that are usual for basic solutions (such as sodium hydroxide: NaOH ↔ Na+ + OH–).
These OH– ions are found in the precipitate because they are present in the initial sodium carbonate solution, which is characterized by an equilibrium:
Na2CO3 + H2O ↔ NaHCO3 + NaOH
NaOH ↔ Na+ + OH–
As a result, blue-green basic copper carbonate (CuOH)2CO3 is formed according to the reaction equation below:
2[CuCl4]2– + 3Na2CO3 + 2H2O → (CuOH)2CO3↓ + 4NaCl + 4Cl– + 2NaHCO3
You can see how this reaction is performed in the video below.
Why is it necessary to filter out and wash the precipitate?
A significant amount of copper that was previously in the solution now resides in the precipitate we obtained. However, besides copper, the solution also contains a lot of “unwanted” compounds left over after two reactions. And we have no intention for them to participate in further chemical transformations! In order to separate our precipitate from the solution, we filter it out.
Particles of the precipitate are large enough to stay on a filter, while molecules and ions from the solution easily pass through its pores. It is important to wash the precipitate several times to completely remove the solution residues.
Why does the precipitate dissolve in the solution of ammonium carbonate?
In a solution, ammonium carbonate dissociates, giving ammonium cations and carbonate anions:
(NH4)2CO3 → 2NH4+ + CO32-
Then, ammonium cation splits into a proton and ammonia:
NH4+ → NH3 + H+.
In fact, we obtain an aqueous ammonia solution. Though basic copper carbonate does not dissolve in water even upon persistent heating, the obtained solution allows to dissolve it and observe a nice color change.
The process yields formation of a compound with intense color that contains complex [Cu(NH3)4]2+ particles. Obviously, they are simply Cu2+ particles, each surrounded by four molecules of ammonia. Now, recall the two previous cases. When copper particles are each surrounded by six water molecules in [Cu(H2O)6]2+, they provide for a blue color of a solution. Furthermore, when copper particles are each surrounded by four chloride ions, a solution turns green. As we can see, the certain surrounding of an ion dramatically influences its properties (color, for instance!).
As we mentioned earlier, blue color of a copper sulfate solution is caused by the presence of [Cu(H2O)6]2+ particles, which intrinsically are copper cations Cu2+, each surrounded by six water molecules. Yet, we know that it is impossible to dissolve basic copper carbonate in water. Therefore, we may conclude that in our final dark violet solution, copper ions reside in a different surrounding. Indeed, in the final solution, copper is in form of cations Cu2+, each surrounded by four molecules of ammonia NH3:
(CuOH)2CO3 + 8NH3 → [Cu(NH3)4][OH]2 + [Cu(NH3)4]CO 3
Moreover, this reaction yields OH– ions, which we are already familiar with. These ions make aqueous solution of ammonia copper complex somewhat similar to NaOH solution:
NaOH ↔ Na+ + OH–
[Cu(NH3)4][OH]2 ↔ [Cu(NH3)4]2+ + 2OH–
And it is [Cu(NH3)4]2+ cations that provide such a beautiful color to the solution.
We can also obtain the same ammonia copper complex [Cu(NH3)4]2+ straight from the reaction of ammonia with copper sulfate. As soon as you mix the ammonia and copper sulfate solutions, you can observe the formation of blue precipitate of copper hydroxide Cu(OH)2, which dissolves after further addition of ammonia, giving intensive dark blue color. You can watch this process in slow motion in the following video.
Why does the solution turn blue when sodium hydrogen sulfate is added?
Influenced by sodium hydrogen sulfate solution, ammonia molecules “break loose” from copper Cu2+ ions surroundings, thus making room for water molecules. For that reason, both by color and composition, the final solution reminds the initial one – copper sulfate CuSO4 solution.
In the previous question, we noted that indigo-violet solution of [Cu(NH3)4][OH]2 resembles NaOH solution in a certain way. Particularly, it also contains OH– ions. Simultaneously, sodium hydrogen sulfate in water comes apart into ions (dissociates) according to the equation below:
NaHSO4 ↔ Na+ + H+ + SO42–
However, OH– and H+ ions cannot be present in water in huge quantities, because as they collide they immediately form a molecule of water themselves:
OH– + H+ → H2O
At first, this is exactly what happens:
[Cu(NH3)4][OH]2 + 2NaHSO4 → [Cu(NH3)4]SO4 + Na2SO4 + 2H2O
As we continue adding sodium hydrogen sulfate solution, complex [Cu(NH3)4]2+ particles decompose:
[Cu(NH3)4]SO4 + 4NaHSO4 → CuSO4 + 2Na2SO4 + 2(NH4)2SO4
And this is what leads to formation of copper Cu2+ ions, which after solvation by water molecules make the solution blue.
By the way, other chains of chemical reactions involving copper compounds are also possible. For example, in the video below you can observe how copper undergoes several transformations, turning back to… Copper!
What colors do copper compounds feature?
Why are some substances colored, while others are colorless? To begin with, we are able to see things around us only because they reflect impinging light. Moreover, white light may be split into a series of colored lights – a so-called spectrum. For instance, such decomposition may be observed when a rainbow appears in the sky or when light passes through water fountains or through a triangular prism. In these cases, white light is decomposed into its component colors: red, orange, yellow, green, blue, indigo, and violet. Objects around us absorb some of these colored rays and reflect the rest. Finally, it is these reflected colored rays that we see with our eyes and perceive as an object color.
Many copper compounds are colored. As we observed in the experiment “There and back again,” they feature vibrant colors: green, blue, and purple. In fact, copper cations Cu2+ normally absorb red, orange, and yellow light rays. And the remaining reflected colored rays determine color of copper compounds. Moreover, in each case, specific color depends on which colored light rays are absorbed best. As a result, the color heavily depends on environment surrounding copper cations Cu2+: a solution or a solid state matter. For instance, in most aqueous solutions, copper ions are surrounded by water molecules and effectively absorb red-orange light rays. It explains blue color of copper compounds.
A very similar environment for copper ions is in copper sulfate hydrate CuSO4*5H2O. Since concentration of copper ions in solid state is even higher than that in a saturated solution, copper sulfate hydrate features deep blue color.
A change of environment for copper ions leads to corresponding color change. For example, adding excess of sodium chloride to copper sulfate solution yields formation of complex particles [CuCl4]2, where each copper cation Cu2+ is surrounded by four chloride ions Cl–. In such an environment, absorption takes place in a red region of the visible spectrum, so the solution becomes green. Another process – formation of complex particles [Cu(NH3)4]2+ absorbing light in a yellow-green region – causes a change from blue to blue-violet color. You can observe the color changes in the video below.
However, not all copper compounds have color. For example, when water is removed from copper hydrate CuSO4*5H2O, copper ions are left with only sulfate ions SO42 around. Then copper cannot absorb visible light anymore and completely reflects it, thus providing for white color of anhydrous copper sulfate.
Under intense heat, many copper compounds yield black oxide CuO. In this substance, each copper atom is surrounded by four oxygen atoms, whereas each oxygen atom stays between two atoms of copper. This structure causes copper oxide CuO to absorb light in the entire visible spectrum, thus providing for its black color. (Watch the following video if you haven’t done the experiment “Anhydrous copper sulfate” yet).