copper sulfate electricity experiment

Edexcel GCSE Chemistry

Practical: electrolysis of copper sulphate solution.

Copper sulphate is an inorganic salt used as a disinfectant and antiseptic. It appears blue when hydrated due to water molecules in its crystalline structure. This compound is also used in electrolysis to produce pure copper.

When dissolved in water, copper sulphate is highly soluble, separating into Cu 2+ and SO 4 2- ions. This creates an electrolytic solution.

During electrolysis, Cu 2+ ions, which are cations, are reduced by gaining two electrons to form neutral copper (Cu) atoms. This process allows us to obtain pure metals from metal salts. We’ll look at how it can be carried out in a laboratory.

Part 1: Electrolysis of Copper Sulphate Using Inert Electrodes

• Safety glasses
• 250 cm 3 beaker
• Graphite electrodes
• Clamps for holding electrodes
• 6-volt direct current (DC) electricity source (e.g., a battery)
• Small 6-volt light bulb to verify current flow
• Wires and alligator clips
• 200 cm 3 of aqueous copper (II) sulphate (pentahydrate) concentration 0.5 M
• Read the safety sheet carefully and understand the precautions needed when working with copper sulphate and sulphuric acid. Sulphuric acid may form when copper sulfate is dissolved in water.
• Pour the copper sulphate solution into the beaker.
• Place the graphite electrodes in the solution, ensuring they do not touch each other.
• Secure the electrodes with clamps.
• Connect the electrodes above the solution to the alligator clips.
• Attach the wires to a 6-volt DC electricity source.
• Turn on the electricity source.
• Observe the changes at each electrode for 20 minutes and record your observations.
• Turn off the DC supply.
• Cathode : An uneven brown/pink solid can be seen forming
• Anode : Bubbles can be observed

Copper ions are more likely to gain electrons than hydrogen ions, which is why copper reduces at the cathode:

Cu 2+ (aq) + 2e –  → Cu 4 (s)

This reaction explains the solid copper observed forming at the cathode during the experiment.

In contrast, because SO 4 2- ions are not halides, they do not get oxidised at the anode. Instead, hydroxide ions present in the solution are oxidised, leading to this reaction:

4OH – (aq) → O 2 (g) + 2H 2 O (l) + 4e –

The bubbles seen in the experiment are oxygen gas being released.

The remaining ions in the solution, H + and SO 4 2- , do not typically react under these conditions. Therefore, no sulphuric acid is formed during this process. The beaker contains the original copper sulphate solution with some water and oxygen.

Part 2: Electrolysis of Copper Sulphate Using Copper Electrodes

• Use the same materials as in Part 1, but replace graphite electrodes with copper foil.
• Carefully read the safety sheet and take note of precautions when working with copper sulphate, sulphuric acid and propanone.
• Measure and record the mass of a piece of copper foil. Connect this to the negative terminal of a DC supply and immerse it in the copper sulphate solution as the negative electrode (cathode).
• Take another piece of copper foil, connect it to the positive terminal, and immerse it in the solution as the positive electrode (anode). Ensure the electrodes do not touch.
• Use clamps to hold the electrodes in place.
• Connect the parts of the electrodes above the solution to the alligator clips.
• Turn on the electricity source and observe the changes at each electrode for 20 minutes, recording your observations.
• Turn off the DC supply and carefully remove the electrodes.
• Gently wash the electrodes with distilled water, then soak them in propanone, and allow to air dry. Do not wipe the electrodes.
• Measure and record the mass of each electrode post-experiment, noting which was the anode and which was the cathode.
• Repeat the experiment five times with fresh electrodes and increasing currents.

Here’s an example of a table you can use to track your results at the cathode:

For example:

Observations: The final masses of the cathode consistently increase as the current increases. This indicates that a material, likely copper, is being deposited on the cathode.

This is when a material (like copper in your experiment) reacts with oxygen or another substance, leading to a loss of electrons. In the context of your electrolysis experiment, oxidation at the anode means copper is reacting and changing chemically.

Here’s an example of a table you can use to track your results at the anode:

Observations: The final masses of the anode consistently decrease as the current increases. This shows us that copper, which is the main material at the anode, is being lost. It happens because the copper reacts (oxidation) and starts mixing into the water around it (dissolving).

You should note that the cathode gained mass while the anode lost mass.

The mass lost at the anode is roughly equivalent to the mass gained at the cathode. This balance occurs because the anode, made of copper, contributes to the reaction by forming Cu 2+ ions that dissolve into the CuSO 4 solution (electrolyte). These copper ions are then reduced and deposited on the cathode, maintaining a near-equivalent mass transfer. For every copper ion discharged at the cathode, one copper ion is generated at the anode.

This relationship explains why the mass gained by the cathode matches the mass lost by the anode.

With higher current, more ions are moved between the electrodes. This leads to a faster and greater copper mass transfer due to the accelerated redox reactions. The increased electron flow results in quicker reactions at both the anode and cathode.

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Principle of Electrolysis of Copper Sulfate Electrolyte

Electrolysis

Before understanding the principle of electrolysis , we should know what is electrolyte or definition of electrolyte

Definition of Electrolyte

Principle of electrolysis.

Next, we immerse two metal rods in the solution and apply an electrical potential difference between them using a battery .

These partly immersed rods are technically referred as electrodes. The electrode connected with negative terminal of the battery is known as cathode and the electrode connected with positive terminal of the battery is known as anode. The freely moving positively charged cations are attracted by cathode and negatively charged anions are attracted by anode. In cathode, the positive cations take electrons from negative cathode and in anode, negative anions give electrons to the positive anode. For continually taking and giving electrons in cathode and anode respectively, there must be flow of electrons in the external circuit of the electrolytic. That means, current continues to circulate around the closed loop created by battery, electrolytic and electrodes. This is the most basic principle of electrolysis .

Electrolysis of Copper Sulfate

Video presentation of basic principle of electrolysis.

Core Practical: Electrolysis of Copper(II)Sulfate ( Edexcel GCSE Chemistry )

Revision note.

Chemistry Lead

Core Practical: Electrolysis of Copper(II)Sulfate

Part 1- electrolysis with passive electrodes.

To electrolyse copper(II) sulfate solution using inert(graphite) electrodes

Apparatus for the electrolysis of copper(II)sulfate using passive(inert) electrodes

Method: (Graphite electrodes)

• Pour copper sulfate solution into a beaker
• Place two  graphite  rods into the copper sulfate solution. Attach one electrode to the  negative  terminal of a DC supply, and the other electrode to the  positive terminal
• Completely fill two small test tubes with copper sulfate solution and position a test tube over each electrode as shown in the diagram
• Turn on the power supply and observe what happens at each electrode
• Test any gas produced with a glowing splint and a burning splint
• Record your observations and the results of your tests

Analysis of results:

• Record observations of what happens at each electrode, including the results of the gas tests

Conclusion:

• Copper  metal is formed at the negative electrode and  oxygen gas is formed at the positive electrode

Part 2: Electrolysis with Active Electrodes

Apparatus for the electrolysis of copper(II)sulfate using active electrodes

Method: (copper electrodes)

• Measure and record the mass of a piece of copper foil. Attach it to the  negative terminal of a DC supply, and dip the copper foil into the copper sulfate solution
• Repeat with another piece of copper foil, but this time attach it to the  positive terminal
• Make sure the electrodes do  not   touch   each   other , then turn on the power supply
• Adjust the power supply to achieve a constant current and leave for 20 minutes
• Remove one of the electrodes and wash it with distilled water, then dip it into propanone
• Lift the electrode out and allow all the liquid to evaporate.  Do   not   wipe the electrodes clean. Measure and record the mass of the electrode
• Repeat with the other electrode making sure you can identify which electrode is which
• Repeat the experiment with fresh electrodes and different currents.
• Record the currents used and the masses of each electrode in suitable table format
• Calculate the change in mass of each electrode
• The  cathode increases  in mass while the  anode decreases
• This occurs as copper atoms are  oxidised  at the anode and form ions while copper ions are  reduced at the cathode, forming copper atoms
• The gain in mass by the negative electrode is the  same as the loss in mass by the positive electrode
• Therefore the copper deposited on the negative electrode must be the  same copper ions that are lost from the positive electrode
• That implies that the concentration of the Cu 2+  ions in the solution remains  constant

Hazards, risks and precautions

Hazard symbols to show substances that are corrosive, harmful to health and flammable

• Copper(II) sulfate solution is corrosive and harmful to health as it is a skin irritant and can cause serious eye damage 
• Propanone, which is often used to clean the electrodes, is flammable
• Avoid contact with the skin and use safety goggles when handling copper(II) sulfate solution
• Propanone should be kept away from naked flames, e.g. a Bunsen burner

Explaining the Electrolysis of Copper(II)Sulfate

Copper refining.

• The electrolysis of CuSO 4 using graphite rods produces oxygen and copper
• By changing the electrodes from graphite to pure and impure copper, the products can be changed at each electrode
• Electrolysis can be used to purify metals by separating them from their impurities
• In the set-up, the impure metal is always the anode , in this case the impure copper
• The cathode is a thin sheet of pure copper
• The electrolyte used is an aqueous solution of a soluble salt of the pure metal at the anode, e.g. CuSO 4
• Copper atoms at the anode lose electrons, go into solution as ions:
• The anode thus becomes thinner due to loss of atoms and the impurities fall to the bottom of the cell as sludge
• The copper(II) ions are attracted to the cathode where they gain electrons and form now purified copper atoms
• The cathode gradually becomes thicker

Cu 2+   +  2e -   ⟶ Cu

• The anode sludge is a highly valuable material and is further refined as it often contains small quantities of precious metals like silver which are found as impurities in the unrefined copper

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Electroplating: Copper-Plated Key

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Electroplating uses a form of electrolysis in which the electrodes play a bigger role than just conducting the current.

Using electricity, you can coat the metal of one electrode with the metal of the other in an electroplating process, also known as electrochemistry.

Electroplating, also known as electrodeposition, is essentially a chemical reaction that helps to make various items we see and use every day. There are also specific types of electroplating such as copper plating, silver plating, and chromium plating. Jewelry and silverware can be silver- or gold-plated, while zinc is often used to coat iron to protect against corrosion. Professional electroplating requires specialized chemicals and equipment to make a high-quality coat, but in this project, you can try your hand at a simple procedure that will transfer copper (a versatile, naturally occurring metal ) to a brass key.

Being red in color, copper is known for its high electrical conductivity, malleability, and corrosion resistance. In copper electroplating, a metal substrate is placed in an electrolytic bath and an electric current is used to cause copper ions to adhere to the base material's surface.

(Adult supervision and chemical safety equipment required.)

Watch us use electricity to copper-plate a brass key (copper electroplating) in this Home Science Tools video. See this project in action!

What You Need:

• 1.5-volt D battery with the battery holder (for power supply)
• Two alligator clip leads or insulated wire
• Beaker or glass (250-ml beaker is recommended or similar glass size)
• Copper strip (pure copper)
• Copper sulfate  
• Copper electrode (or coil of copper wire)
• Safety equipment

How to Electroplate Copper:

• Prepare the key for the DIY copper plating by cleaning it with a thin layer of toothpaste or soap and water. Dry it off on a paper towel.
• Stir copper sulfate into some hot water in a beaker until no more will dissolve. Your solution should be dark blue. Let it cool.
• Use one alligator clip and attach the copper electrode to the positive terminal of the battery (this is now the anode ) and then attach the key to the negative terminal (now called the cathode ).
• Partially suspend the key in the solution by wrapping the wire lead loosely around a pencil and place the pencil across the mouth of the beaker. The alligator clip should not touch the solution.
• Place the copper strip/mass of copper into the solution, making sure it doesn't touch the key. The plating solution level is now below the alligator clip. The copper strip will produce a path for conductivity. An electrical circuit has now formed with the positive electrodes & negative electrodes and an electrical current is flowing.
• Leave the circuit running for 20-30 minutes, or until you are happy with the amount of copper on the key.

What Happened During the Plating Process:

The copper sulfate solution is an electrolyte solution that conducts electricity from one electrode to the other, creating an electrical current.

When the current is flowing, oxidation (loss of electrons) happens at the copper anode, adding copper ions to the solution.

Those ions travel on the electric current to the cathode, where reduction (gain of electrons) happens, plating the copper ions onto the key.

There were already copper ions present in the copper sulfate solution before you started, but the oxidation reaction at the anode kept replacing them in the solution as they were plated with a thin layer onto the key, keeping the reaction going.

This project has many variables, including the cleanness and smoothness of the key, the strength of the copper sulfate solution, and the strength of the current.

If a black soot-like substance starts forming on the key, your solution is not strong enough for the current. Take the electrodes out and add more copper sulfate. When you put them back in, make sure the anode and cathode are as far apart as possible. Be sure to take notes for your science experiment to ensure you have great data collection.

There are lots of projects you can do with electroplating!

One fun idea is to use a flat piece of brass as your cathode and draw a design on it with an oil-based marker. The copper will not bond where the marker is.

After you're done plating it, you can use acetone (or nail polish remover) to wipe off the marker, leaving a design of the brass showing through the copper. Copper is relatively dull in color, which means other additives may be needed if a brighter finish is required. You can use a little metal polish to make the copper shiny if you desire.

You may want to try this simple copper-plating experiment that doesn't use electrolysis and requires only household materials.

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Electrolysis of Copper Sulphate

Related Topics: More Lessons for IGCSE Chemistry Math Worksheets

A series of free IGCSE Chemistry Activities and Experiments (Cambridge IGCSE Chemistry).

Electrolysis of Copper(II) Sulphate Solution This experiment is designed to demonstrate the different products obtained when the electrolysis of copper(II) sulfate solution is carried out first with inert graphite electrodes and then with copper electrodes. The use of copper electrodes illustrates how copper is refined industrially.

Using Graphite Rods

• Set up an electrolysis cell using graphite rods as electrodes and copper(II) sulphate solution as electrolyte.
• At the anode: Bubbles of gas (oxygen) are formed at the anode. Anode reaction: 4OH - (aq) → O 2 (g) + 2H 2 O(l) + 4e -
• At the cathode: A deposit of copper forms on the cathode; this will often be powdery and uneven. Cathode reaction: Cu 2+ (aq) + 2e - → Cu(s)

Using Copper Plates (Refining Copper)

• Replacing the graphite rods with clean copper plates produces a different anode reaction.
• At the anode: No oxygen is produced, rather the copper anode dissolves. Anode reaction: Cu(s) → Cu 2+ (aq) + 2e -
• At the cathode: A deposit of copper forms on the cathode. Cathode reaction: Cu 2+ (aq) + 2e - → Cu(s) During this electrolysis, the mass gained of copper at the cathode is equal to the mass lost at the anode.

Electrolysis in the Lab

• Using a table, highlight the similarities and differences between using graphite electrodes and copper electrode for the electrolysis of copper sulphate.

Check out the following link to find out how electrolysis can be used for electroplating. Electroplating

We welcome your feedback, comments and questions about this site or page. Please submit your feedback or enquiries via our Feedback page.

Electrolysis of Copper sulfate (CuSO 4 ) Solution

Copper sulfate is soluble in water and gives a blue colour aqueous solution though anhydrous copper sulfate solid is white. In an aqueous solution, Copper sulfate completely dissociates to copper 2+ (Cu 2+ ) cation and sulfate (SO 4 2- ) anion. Therefore, CuSO 4 is a strong electrolyte. We can perform electrolysis of copper sulfate in different methodologies to perform copper plating and will discuss those stuff in detail in this tutorial.

In an electrolysis process, you should have the capability of deciding which chemical will be oxidized and which one will be reduced. Follow below steps to decide products of electrolysis and to observe the behavior of electrolysis.

Electrolysis of copper sulfate solution with inert and active anode and for copper plating

Electrolysis of copper sulfate solution with graphite / platinum anode and iron cathode.

In this experiment, we are going to deposit metallic copper layer in the surface of a iron piece. We use graphite or platinum electrode because they do not react easily with water and is not oxidized or reduced too.

Anode reaction / Oxidation reaction

4OH - (aq) → O 2(g) + 2H 2 O (aq) + 4e

Cathode reaction / Reduction reaction

There are two cations (Cu 2+ , H + ) around cathode. To decide the, which cation will be reduced, electrochemical series with standard potential is used as below. Let's consider standard potential values of Cu 2+ and H + .

Cu 2+ (aq) + 2e → Cu (s)

Physical changes such as colour changes, gas formations, precipitating, electrolysis of copper sulfate solution with copper anode (active) and iron cathode.

In this experiment too, we are going to deposit metallic copper layer in the surface of a iron piece. But, We use a Copper electrode electrode here. So, there is a change in the overall reactions and we will see how products are given.

Anode reaction: Cu (s) → Cu 2+ (aq) + 2e

So, copper atoms in the copper piece are oxidized and Cu 2+ ions are released to the aqueous solution. Released electrons due to the oxidation process are flown towards cathode through the DC power supply. As well as, due to the released of Cu 2+ cations to the solution, Cu 2+ cation concentration will be kept at a constant value.

There are two cations (Cu 2+ , H + ) around the cathode. To decide, which cation will be reduced, electrochemical series with standard potential is used as below. Let's consider standard potential values of Cu 2+ and H + .

Cathode reaction: Cu 2+ (aq) + 2e → Cu (s)

Example problems for calculating current requirement in cuso 4 electrolysis.

Let's try few questions to understand this.

1.27 g of Copper should be plated on a Iron spoon by using a Copper bar as an anode. If electrolysis process should be completed within 10 minutes, Calculate the current.

A copper sulphate solution (100 cm 3 ) is electrolyzed using graphite electrodes. initially, cuso 4 concentration was 1.0 mol dm -3 . if constant 1a current is supplied, what is the concentration of cu 2+ ion after 1 minute, related tutorials to electrochemistry, related tutorials to copper.

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D escribing and explaining electrolysis of copper sulfate solution and electroplating

Know that electrolysis requires a conducting solution of ions (electrolyte of copper sulfate) and two solid conducting electrodes e.g. graphite, platinum or copper. Know that the electrolyte here is copper(II) sulfate solution containing high concentrations of copper ions and sulfate ions. Know that electrolysis will only happen if a d.c. electrical current is passed through the copper sulfate solution. Be able to describe the apparatus required to electrolyse copper sulfate solution and be able to explain and understand the formation of the electrolysis products by: knowing that the positive copper ion is reduced by electron gain and discharged at the negative cathode as copper atoms and the blue colour intensity decreases (unless the anode is made of copper), knowing that the negative sulfate ion is NOT oxidised by electron loss and so NOT discharged at the positive anode, know that with inert electrodes, oxygen gas is formed at the positive anode by the oxidation of the hydroxide ion or water molecule, and be able to write out the electrode equations for the formation of copper, oxygen or copper(II) ions depending on the nature of the electrodes.

From the electrode equations, be able to explain why the mole ratio of copper atoms to oxygen molecules (Cu : O 2 ) is theoretically 2 : 1

Know how to test for oxygen gas from the electrolysis of copper sulfate solution with inert electrodes.

Know that when the anode is made of copper, the copper is oxidised to copper ions which dissolve in the solution.

Be able to explain that if both electrodes are made of copper, the intensity of the blue colour due to the Cu2+ ions, will remain constant.

Know and be able to explain how an electrolysis cell can be used to plate any electrically conducting surface with a metal coating when the electrolyte contains ions of the same metal and this process is called electroplating.

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keywords and phrases: revision study notes for AQA Edexcel OCR IGCSE/GCSE chemistry topics modules on explaining the electrolysis copper sulfate solution with copper electrodes carbon graphite electrodes platinum electrodes electroplating half-equations products at the positive anode products at the negative cathode description of apparatus for doing electrolysis experimental investigation electrolyte diagrams of electrolysis cell experiments anode sludge the purification of copper by electrolysis, applications of electroplating in industry preventing corrosion improving appearance electrical conduction nickel chromium tin silver zinc gold plating by electrolysis

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How to Make Copper Sulfate (Copper Sulphate)

It’s easy to make copper sulfate using readily available materials. Copper(II) sulfate is also known as copper sulphate, blue vitriol, or bluestone. Usually, it is a vibrant blue salt encountered as copper sulfate pentahydrate (CuSO 4 ·5H 2 O). The chemical has several uses, but most people use it for growing blue blue copper sulfate crystals .

Here is how you make copper sulfate yourself, using a battery, copper wire, and dilute sulfuric acid.

Materials for Making Copper Sulfate

The easiest and safest method of making copper sulfate uses electrochemistry.

• Copper wire (which is high-purity copper )
• Sulfuric acid (H 2 SO 4 or battery acid )
• 6-volt Battery

Concentrated sulfuric acid (like the kind in a lab) is 98% sulfuric acid and 2% water, with a concentration of 18.4 M. That is too strong for this project, so you need to dilute it. If you have dilute sulfuric acid, go ahead and use that. If you’re making copper sulfate at home, you’re probably using battery acid, which averages around 37% acid in water or about 4 M. You don’t need to dilute it much in this project.

While the instructions call for a 6-volt battery, a 9-volt battery or a constant power supply work fine.

Make Copper Sulfate

• Pour 30 ml of water and then 5 ml concentrated sulfuric acid into a small glass jar or beaker. Always add acid to water , not the other way around. This minimizes the chance of the acid splashing. For battery acid, use less water. The concentration of acid is not critical, so either 30 milliliters of acid in 40 milliliters of water or mixing half battery acid and half water is fine.
• Inspect your wires. If they are insulated, strip enough insulation that you have bare copper ends to put in the liquid. Attach a copper wire to each battery terminal and immerse the exposed ends in the solution to that the wires are not touching each other.
• The liquid turns blue as copper sulfate is produced.

Concentrate the Copper Sulfate

The reaction between sulfuric acid and copper yields a dilute copper(II) sulfate solution. If left undisturbed, copper sulfate crystals form as the water evaporates. However, the solution still contains some sulfuric acid, so use care when removing the crystals (which are your solid product).

Alternatively, concentrate the solution by boiling it. After evaporating the liquid, blue copper sulfate powder remains. Any remaining liquid that does boil away is concentrated sulfuric acid. Pour this liquid off and save it for future science experiments.

Once you have copper sulfate, dissolve it in water and grow copper sulfate crystals.

Tips for Success

When you run electricity through the copper electrodes, expect bubbling from the anode (negative electrode) in the liquid. These bubbles contain hydrogen gas. Meanwhile, the copper at the cathode (positive electrode) dissolves. Some of the dissolved copper ions make their way to the anode and are reduced. When this happens, it reduces the copper sulfate yield. But, a little care with the set-up reduces the loss.

If you have enough wire, coil the copper for the cathode (connected to the “+”) and place it on the bottom of the jar or beaker. Either leave the wire insulation in place above the coil or else slide a section of plastic tubing (such as aquarium tubing) over exposed wire just above the coil. This minimizes the reaction between the cathode and anode. Place the anode (connected to the “-“) higher in the liquid and distant from the coil. Ideally, hydrogen bubbles only form from the anode. If both electrodes bubble, move the copper wires further apart. With this set-up, copper sulfate forms at the bottom of the container, near the cathode.

Make Copper Sulfate Using Sulfuric Acid and Nitric Acid

While the electrochemical method is the safest way of making copper sulfate, there are other synthesis routes. Another method uses sulfuric acid, nitric acid, and copper (either as a chunk or wire). The disadvantage is that nitric acid and concentrated sulfuric acid are not common home chemicals. They come from a chemical supply store. The acid mixture is highly corrosive and produces toxic vapor, so the procedure is best done in a fume hood. This reaction is popular as a chemistry demonstration because of the color changes. Note that the product includes both copper(II) sulfate and copper(II) nitrate.

• 70% nitric acid
• concentrated (98%) sulfuric acid
• Place 30 milliliters of water in a beaker.
• Add 5 milliliters of nitric acid and 3 milliliters of concentrated sulfuric acid.
• Gently drop about 6 grams of copper into the acid solution. The reaction releases a brown gas and the solution turns blue.
• Within the fume hood, let the acid evaporate. Collect the copper sulfate crystals.

Make Copper Sulfate Using Sulfuric Acid and Hydrogen Peroxide

You can make copper sulfate from copper in a mixture of sulfuric acid and hydrogen peroxide called piranha solution . This is not a recommended synthesis method. It is not very efficient and the acid and peroxide often boil during mixing and may overflow or break a glass container. While 30% hydrogen peroxide is available from a beauty supply store, the concentrated sulfuric acid comes from a chemical supply store.

• 30% hydrogen peroxide (H 2 O 2 )
• concentrated (98%) sulfuric acid (H 2 SO 4 )
• Pour 10 milliliters of 30% hydrogen peroxide into a borosilicate glass beaker.
• Add 3 milliliters of concentrated sulfuric acid. This reaction is exothermic , so expect heat!
• Carefully add about 3 grams of copper. The copper bubbles and the clear liquid turns blue.
• Pour the liquid onto a shallow glass dish. Leave any remaining copper in the original container. Copper sulfate crystals form as the liquid evaporates.

Copper Sulfate Safety and Disposal

• Wear gloves and eye protection. Sulfuric acid is corrosive and causes burns upon contact. Do not touch or inhale the acid. In the event of a splash, immediately rinse the affected are with lots of water. Neutralize a spill using a weak acid, such as baking soda. Then, rinse with plenty of water.
• Avoid skin contact with the copper sulfate solution. Copper sulfate is a skin irritant. It is only mildly toxic, but please don’t drink the liquid. It still contains some acid and may be corrosive. In case of accidental contact, rinse affected skin with water.
• While municipal water treatment can handle copper just fine, copper sulfate is toxic to invertebrates, so don’t dump copper sulfate outdoors. Rinse unused product down the drain with plenty of water.
• Clayton, G. D.; Clayton, F. E. (eds.) (1981). Patty’s Industrial Hygiene and Toxicology (3rd ed.). Vol. 2, Part 6 Toxicology. NY: John Wiley and Sons. ISBN 0-471-01280-7.
• Copper Development Association Inc. “ Uses of Copper Compounds: Copper Sulphate .”
• Wiberg, Egon; Nils Wiberg; Arnold Frederick Holleman (2001). Inorganic Chemistry . Academic Press. ISBN 978-0-12-352651-9.
• Zumdahl, Steven; DeCoste, Donald (2013). Chemical Principles . Cengage Learning. ISBN 978-1-285-13370-6.

Related Posts

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• Class 9 Practical Experiment on Reaction of Iron With Copper Sulphate Solution in Water

Experiment on Reaction of Iron with Copper Sulphate Solution in Water

A physical change occurs when there is no change in the composition of a substance and no change in the chemical nature of the substance.

The interconversion of state occurs during physical change.

SOLID ⇄ LIQUID ⇄ GAS

A chemical change is a change that causes a change in the chemical properties of matter, resulting in the formation of a new substance. As an example, consider the burning of oil or fuel.

Heat is evolved or taken in, the formation of bubbles, gas, and fumes, as well as a change in the colour of the reactants, can take place when they form a product.

Reactants → Products

A + B → C (Chemical reaction)

Table of Contents

Materials required.

• Observation

Precautions

• Frequently Asked Questions – FAQs

To carry out the reaction between Copper sulphate solution and water and Classify it as physical change and chemical changes.

Iron nails, Copper Sulphate solution, Test Tube, Clamp Stand, Sandpaper.

The colour of pure iron is greyish. Pure copper is a reddish-brown metal. The presence of Cu2+ ions causes the aqueous C solution of copper sulphate to be blue. The presence of Fe2+ ions causes the aqueous solution of ferrous sulphate to be pale green.

Since iron is more reactive than copper, it removes copper from its salt solution.

2. Separate two test tubes and label them A and B. Add 10 mL of freshly prepared copper sulphate solution to each test tube and secure these test tubes in two separate clamp stands.

3. Thread the nail and hang it in test tube B. It is important to ensure that the iron nail is completely immersed in CuS0 4 solution. Tie the other end of the thread to the stand.

4. Keep the other iron nail on a piece of white paper.

5. Leave the setup alone for a while.

6. Take the nail out of the solution and place it along the side of the second iron nail on the sheet of paper.

7. Record your observations.

Observations

1. The brown coating on the iron nail indicates that copper is deposited on the iron nail as a result of iron displacement.

2. The colour of the blue colour copper sulphate solution changes to green.

3. The greenish colour of the solution in the test tube indicates the presence of Fe 2+ ions in the solution.

4. This is a single displacement reaction in which copper is displaced by iron from copper sulphate solution, resulting in the formation of a new compound, ferrous sulphate.

5. A chemical change occurs as a result of the reaction.

1. Clean iron nails by rubbing with sandpaper.

2. Copper sulphate solution is poisonous, so use caution when handling it.

3. The test tubes should not be touched or disturbed during the experiment.

4. After completing the experiment, the copper-coated iron nail should not be touched.

Frequently Asked Questions on Reaction of Iron with Copper Sulphate Solution in Water

What is the colour of copper sulphate solution.

The colour of the copper sulphate solution is blue.

Why are iron nails rubbed with sandpaper?

Iron nails are rubbed with sandpaper so as to remove any impurities present like rust, dust or greasy surface. Iron nails are rubbed with sandpaper so as to remove any impurities present like rust, dust or greasy surface.

Does the colour of the copper sulphate solution change?

Yes, the colour of the copper sulphate solution changes from blue to light greenish.

What does the greenish colour of the solution show?

The greenish colour of the solution shows that Fe 2+ Ions are present in the solution.

What does the brown coating on the iron nails show?

The brown coating in the iron nail shows that copper is deposited in it by displacing iron.

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Researchers observe “locked” electron pairs in a superconductor cuprate

The finding could help future efforts to design superconductors that work at higher temperatures.

By David Krause

For the past century since their discovery, superconductors and their mysterious atomic properties have left researchers in awe. These special materials allow electricity to flow through them without any energy loss. They even allow trains to levitate.

But superconductors typically only work at extremely cold temperatures. When these materials are heated, they become ordinary conductors, which allow electricity to flow but with some energy lost, or insulators, which don’t conduct electricity at all.

Researchers have been hard at work looking for superconductor materials that can perform their magic at higher temperatures – perhaps even room temperature someday. Finding or building such a material could change modern technology, from computers and cell phones to the electric grid and transportation. Furthermore, the unique quantum state of superconductors also makes them excellent building blocks for quantum computers.

Now, researchers have observed that a necessary characteristic of a superconductor – called electron pairing – occurs at much higher temperatures than previously thought, and in a material where one least expects it – an antiferromagnetic insulator. Although the material did not have zero resistance, this finding suggests researchers might be able to find ways to engineer similar materials into superconductors that operate at higher temperatures. The research team from SLAC National Accelerator Laboratory, Stanford University, and other institutions published their results August 15 in Science .

“The electron pairs are telling us that they are ready to be superconducting, but something is stopping them,” said Ke-Jun Xu, a Stanford graduate student in applied physics and paper co-author. “If we can find a new method to synchronize the pairs, we could apply that to possibly building higher temperature superconductors.”

Out-of-sync electrons

Over the past 100 years, researchers have learned a lot about how exactly superconductors work. We know, for instance, that for a material to superconduct, electrons have to pair off, and these pairs must be coherent – i.e., their movements must be synchronized. If electrons are paired but incoherent, the material might end up being an insulator.

In superconductors, the electrons act like two reticent people at a dance party. At first, neither person wants to dance with the other. But then the DJ plays a song that both people like, allowing them to relax. They notice one another enjoying the song and become attracted from afar – they have paired but have not yet become coherent.

Then the DJ plays a new song, one that both people absolutely love. Suddenly, the two people pair and start to dance. Soon everyone at the dance party follows their lead: They all come together and start dancing to the same new tune. At this point, the party becomes coherent; it is in a superconducting state.

In the new study, the researchers observed electrons in a middle stage, where the electrons had locked eyes, but were not getting up to dance.

Cuprates acting oddly

Not long after superconductors were first discovered, researchers found that the thing that got electrons paired up and dancing was vibrations in the underlying material itself. This kind of electron pairing happens in a class of materials known as conventional superconductors, which are well understood, said Zhi-Xun Shen, a Stanford professor and investigator with the Stanford Institute for Materials and Energy Sciences (SIMES) at SLAC who supervised the research. Conventional superconductors work at temperatures typically close to absolute zero, below 25 Kelvin, in ambient pressure.

Unconventional superconductors – such as the copper oxide material, or cuprate, in the current study – work at significantly higher temperatures, sometimes up to 130 Kelvin. In cuprates, it is widely believed something beyond lattice vibrations helps pair up electrons. Although researchers aren't sure exactly what's behind it, the leading candidate is fluctuating electron spins, which cause the electrons to pair and dance with a higher angular momentum. This phenomenon is known as a wave channel – and early indications of such a novel state were seen in an experiment at SSRL about three decades ago.  Understanding what drives electron pairing in cuprates could help design superconductors that work at higher temperatures.

In this project, scientists chose a cuprate family that had not been studied in depth because its maximum superconducting temperature was relatively low – 25 Kelvin – compared to other cuprates. Even worse, most members of this family are good insulators. To see the atomic details of the cuprate, researchers shined ultraviolet light onto material samples, which eject electrons from the material. When the electrons are bound, they are slightly more resistant to being ejected, resulting in an “energy gap." This energy gap persists up to 150 Kelvin, suggesting that electrons are paired at much higher temperatures than the zero resistance state at about 25 Kelvin. The most unusual finding of this study is that the pairing is the strongest in the most insulating samples.

The cuprate in the study might not be the material to reach superconductivity at room temperature, around 300 Kelvin, Shen said. "But maybe in another superconductor material family, we can use this knowledge for hints to get closer to room temperature," he said.

“Our findings open a potentially rich new path forward,” Shen said. “We plan to study this pairing gap in the future to help engineer superconductors using new methods. On the one hand, we plan to use similar experimental approaches at SSRL to gain further insight into this incoherent pairing state. On the other hand, we want to find ways to manipulate these materials to perhaps coerce these incoherent pairs into synchronization.”

This project was supported in part by the DOE’s Office of Science. SSRL is a DOE Office of Science user facility.

Citation: Ke-Jun Xu et al., Science, 15 August 2024 ( 10.1126/science.adk4792 )

SLAC National Accelerator Laboratory explores how the universe works at the biggest, smallest and fastest scales and invents powerful tools used by researchers around the globe. As world leaders in ultrafast science and bold explorers of the physics of the universe, we forge new ground in understanding our origins and building a healthier and more sustainable future. Our discovery and innovation help develop new materials and chemical processes and open unprecedented views of the cosmos and life’s most delicate machinery. Building on more than 60 years of visionary research, we help shape the future by advancing areas such as quantum technology, scientific computing and the development of next-generation accelerators.

SLAC is operated by Stanford University for the U.S. Department of Energy’s Office of Science . The Office of Science is the single largest supporter of basic research in the physical sciences in the United States and is working to address some of the most pressing challenges of our time.

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Exothermic metal displacement reactions

In association with Nuffield Foundation

Try this class experiment to explore what happens when different metals are added to a copper(II) sulfate solution

In this practical, students add powdered or finely-divided metals to a copper(II) sulfate solution and measure the temperature rises. The experiment reinforces ideas about energy changes during reactions, the reactivity series of the metals and the chemical behaviour of metals.

The finely-divided metals should be distributed on labelled plastic weighing dishes or watch glasses to avoid cross-contamination and wastage.

• Eye protection
• Expanded polystyrene cups standing in glass beakers for stability, x4 (see note 5 below)
• Measuring cylinder, 25 cm 3
• Thermometer, –10 °C to +110 °C or similar
• Large bowl or bucket for collecting the residues from experiments (do not let students put residues into sinks; see note 6 below)
• Copper(II) sulfate(VI) solution, 1.0 M (HARMFUL, DANGEROUS FOR THE ENVIRONMENT), 100 cm 3
• Iron (HIGHLY FLAMMABLE)
• Magnesium (HIGHLY FLAMMABLE)
• Zinc (HIGHLY FLAMMABLE)

Health, safety and technical notes

• Read our standard health and safety guidance.
• Wear eye protection throughout.
• Copper(II) sulfate(VI) solution, CuSO 4 (aq), (HARMFUL, DANGEROUS FOR THE ENVIRONMENT) – see CLEAPSS Hazcard  HC027c  and CLEAPSS Recipe Book RB031.
• Powdered or finely divided metals: iron, Fe(s), magnesium, Mg(s), zinc, Zn(s), (all HIGHLY FLAMMABLE) and tin, Sn(s) – see CLEAPSS Hazcards HC055A , HC059A , HC107 and HC102A . Iron filings tend to be greasy and should be degreased using propanone (HIGHLY FLAMMABLE, IRRITANT, see CLEAPSS Hazcard HC085A ) and dried before being used. Carry out degreasing in a fume cupboard.
• A desirable, but not essential, addition are lids for the polystyrene cups. A lid can be made by cutting a suitably-sized piece from a polystyrene ceiling tile and making a hole for the thermometer.
• Provide a bowl or bucket for discarding the residues. Metal residues in sinks are almost impossible to remove. Iron particles rust and cause unsightly stains.
• Measure 20 cm 3  of the copper(II) sulfate(VI) solution into a polystyrene cup.
• Put the cup into a beaker so that it does not fall over.
• Measure and record the temperature of the solution.
• Add the first of the powdered metals and stir the mixture with the thermometer. 
• Observe the temperature over the next few minutes until a maximum temperature is reached.
• Record the temperature rise.
• Repeat the procedure with fresh polystyrene cups using each of the other metals.

Teaching notes

The temperature rises should be approximately:

• Magnesium: 39°C

The results are approximately in line with the reactivity series of the metals. The ‘wrong’ order of magnesium and zinc might be due to oxidation on the surface of the magnesium being more extensive than for zinc.

Additional information

This is a resource from the  Practical Chemistry project , developed by the Nuffield Foundation and the Royal Society of Chemistry. This collection of over 200 practical activities demonstrates a wide range of chemical concepts and processes. Each activity contains comprehensive information for teachers and technicians, including full technical notes and step-by-step procedures. Practical Chemistry activities accompany  Practical Physics  and  Practical Biology .

© Nuffield Foundation and the Royal Society of Chemistry

• 14-16 years
• Practical experiments
• Thermodynamics
• Reactions and synthesis

Specification

• Deduce an order of reactivity of metals based on experimental results.
• C3.2.1 deduce an order of reactivity of metals based on experimental results including reactions with water, dilute acid and displacement reactions with other metals
• C4.1f deduce an order of reactivity of metals based on experimental results
• C4.1e deduce an order of reactivity of metals based on experimental results
• A reaction or process that releases heat energy is described as exothermic.
• (c) the relative reactivities of metals as demonstrated by displacement (e.g. iron nail in copper(II) chloride solution) and competition reactions (e.g. thermit reaction)
• 2.1.5 collect and/or analyse experimental data to predict where an unfamiliar element should be placed in the reactivity series or make predictions about how it will react;

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