What happens when CUSO4 is electrolyzed

If different metal rods are placed in a copper (II) sulfate solution, a copper mirror forms on the iron rod, while no copper deposition takes place on the silver rod. On all metals that are more oxidizable than copper, copper is deposited in a copper (II) salt solution. In the case of metals that are more difficult to oxidize than copper, for example the noble metals silver or gold, there is no copper deposition.



Iron rod and silver rod in copper (II) sulfate solution
  

The phenomenon can be explained using the electrochemical potential with the help of a redox reaction. If a zinc rod is immersed in the copper (II) salt solution, the zinc atoms become zinc with the release of electrons2+-Ions are oxidized, which go into solution, while the copper2+-Ions in the solution are reduced with the help of the released electrons and are deposited on the zinc sheet.
   
Oxidation (electron donation)Zn Zn2+ + 2 e
Reduction (electron uptake) Cu2+ + 2 e  Cu
Overall responseZn + Cu2+   Cu + Zn2+
 


Zinc rod in copper (II) sulfate solution, right after 90 minutes


Oxidation and reduction each form in this electrochemical reaction Redox couple. The zinc rod will gradually dissolve, so the copper layer on the rod will not last for long. The reaction continues until all of the copper (II) sulfate solution has been consumed. The products are copper powder and the (colorless) zinc (II) sulphate solution, the zinc rod appears to be “pitted”. If you had enough copper (II) sulfate solution available, the zinc rod would dissolve completely.

If a copper sheet is placed in a silver (I) nitrate solution, elemental silver is deposited on the copper sheet, since the electrode potential of the silver is positive in relation to the copper. Compared to silver, copper is therefore a less noble metal.



Copper sheet in silver (I) nitrate solution

Film available on> DVD  


Oxidation (donation of electrons)Cu Cu2+ + 2 e
Reduction (electron uptake) 2 Ag+ + 2 e  2 Ag
Overall responseCu + 2 Ag+   2 Ag + Cu2+
   

The ions of the non-metals also form redox pairs. If chlorine gas is passed into a potassium bromide solution, elemental bromine is formed, and chloride ions are formed in the solution. The formation of bromine can be recognized by the brown fumes. If, on the other hand, bromine is passed into a potassium chloride solution, no reaction takes place. Chlorine atoms take up electrons more easily than bromine atoms, so they like to react with the electrons of the bromide ions.
   
Oxidation (electron donation)2 Br  Br2 + 2 e
Reduction (electron uptake) Cl2 + 2 e  2 cl 
Overall responseCl2 + 2 Br  2 cl + Br2
 


Introducing chlorine into a bromide solution

Film available on> DVD    
   

If one conducts appropriate experiments with all redox pairs of metals or non-metals, one can do a Redox series or one electrochemical series put up. In the left column are all atoms or ions that can donate electrons and are oxidizable, i.e. all reducing agents such as metal atoms and non-metal anions. In the right column are all oxidizing agents, for example the non-metal atoms and the metal cations. The redox series can be used to determine whether a reaction is taking place at all. In principle, no reactions in ascending order from left to right are possible. A reducing agent can only react with an oxidizing agent if it is stronger and can "snatch" the electrons from the reducing agent.




  
The ions of the noble metals are more likely to accept electrons than the ions of the base metals. The strength of the ability to absorb electrons can be achieved with the Redox potential or measure the normal potential in relation to a standard hydrogen electrode. The redox potential of the redox couple Cu / Cu2+ + 2e is +0.34 volts, that of the redox couple Zn / Zn2+ + 2e is given as −0.76 volts. Building on this knowledge, one can calculate the voltage if one builds a galvanic cell from the respective redox pairs.

A galvanic cell consists of two half-cells with different electrodes and electrolytes, it converts chemical energy into electrical energy. At the Daniell element the half-cells consist of a copper electrode that is immersed in copper (II) sulfate solution and a zinc electrode that is immersed in zinc sulfate solution. So that charges can be exchanged between the two cells, a pipe with a conductive salt solution is required as a connecting bridge for the electrolytes, as well as a conductive connection between the electrodes. This is how you get a chemical battery. In high-performance Daniell elements, the connection of the electrolyte solutions consists of a porous diaphragm made of clay or earthenware. In the Daniell element shown below, a sheet of zinc sits in a porous crucible that is filled with a zinc sulphate solution. The copper electrode sits in a copper (II) sulfate solution in the outer glass container:



Daniell element with zinc and copper electrodes

Film available on> DVD

  
The resulting voltage of 1.10 volts can be calculated using the normal potential of the redox couples. The normal potential of the redox couple Zn / Zn2+ + 2e is −0.76 volts lower than that of the pair Cu / Cu2+ + 2ewith +0.34 volts. The difference is 1.10 volts. The Solution pressure is larger in the less noble zinc than in the copper, there are many more zinc ions in solution than copper ions in the other half-cell. The released electrons migrate to the copper electrode and react with the copper ions in the electrolyte solution to form elemental copper. The charge exchange between the electrolytes takes place through the reaction of the Zn2+-Ions with the SO42−-Ions take place across the diaphragm.



 

Batteries, accumulators or fuel cells can be built from galvanic cells. If you connect them in series, the tensions add up. If they are connected in parallel, the available current is added. With the Daniell element shown above, a current of about 0.1 amperes is obtained at 1.10 volts.

  

Electrolysis of a zinc iodide solution:
Formation of a zinc tree and elemental iodine

Film (with detailed shots) available on> DVD  


The processes that take place in a galvanic cell can also be reversed: in the electrolysis chemical reactions are forced with the help of electrical energy. In this way, chemical elements can be obtained from electrically conductive salt solutions in a material breakdown (analysis). When a zinc iodide solution is electrolyzed, elemental zinc and elemental iodine are obtained. When zinc bromide is melted electrolysis, the salt is heated until it melts. In the melt there are freely moving ions that enable conductivity. When a DC voltage is applied, bromine and zinc are generated at the electrodes. Electrolysis is an important process in the chemical industry for the extraction of metals, for example in chlor-alkali electrolysis or in melt-flow electrolysis according to Downs and Hall-Hérout. In copper refining, electrolysis is used to purify copper. The electrolysis of water to hydrogen and oxygen plays an important role in hydrogen technology.
  

  
 

In the Electroplating Metals are electrolytically coated in a bath with other metals such as chromium, zinc or copper to protect them from corrosion. The workpiece to be coated is hung on the negative pole of a direct current source as a cathode in the bath. The metal to be applied hangs on the positive pole as a consumption anode in the electrolyte bath. A salt solution of the anode metal is always used as the electrolyte solution. In the case of galvanic copper plating, this would be a copper (II) sulfate solution. As soon as a direct voltage is applied, the excess electrons on the workpiece (cathode) reduce the copper (II) ions in the solution to elemental copper. At the consumption anode, copper atoms are oxidized to copper ions with the release of electrons.

   


   
Rusting is an attack by oxidizing substances such as atmospheric oxygen, water or acids on metals. In the corrosion materials react with substances from the environment, they are attacked and decompose. Corrosion processes can also occur on a small scale at the points of contact between two different metals. In building construction, for example, questions arise such as: Can a copper rain gutter be built in together with a drainage pipe made of zinc? An experiment serves as an answer: If you place a copper rod and a zinc rod isolated from each other in a vessel with dilute hydrochloric acid, gas is only generated on the zinc rod. The hydrochloric acid does not initially attack the copper. As soon as the two metal rods touch, the gas development can be seen on the copper rod, while it decreases on the zinc rod:



As soon as the zinc rod and the copper rod are in the
Touching hydrochloric acid solution creates a local element.

Film available on> DVD  
  

When the two metals touch, a galvanic cell is created. Electrons flow from the less noble zinc to the more noble copper. The electrons reduce the H.+-Ions from the acidic solution on the surface of the copper to hydrogen. Such a short-circuited galvanic cell is called Local element. In the case of the local element, the oxidation and thus the decomposition of the zinc are accelerated.
  
oxidationZn Zn2+ + 2 e
reduction2 H+ + 2 e  H2   
Overall responseZn + 2H+   Zn2+ + H2   
 
To prevent local elements, one must pay attention to soldering points, weld seams or screw connections as soon as moisture is added.
  
  
Alloys such as stainless steel or chrome argan are considerably more rust-resistant than pure metals. If you want to use pure metals, you have to cover the surface with a protective layer to protect against corrosion, which prevents the entry of oxidising substances. Oiling, painting or covering with an oxide layer are suitable. Metals such as aluminum react immediately with the oxygen in the air and form a self-protecting layer of aluminum oxide. Aluminum objects are used to reinforce this oxide layer Anodizing protected:

 


 
The aluminum workpiece to be anodized is hung in an electrolysis bath with dilute sulfuric acid as the anode and an electrode made of carbon, lead or aluminum as the cathode. If a direct voltage is applied, the H3O+-Ions of the acid at the cathode are reduced to hydrogen and water while accepting electrons. On the aluminum workpiece, the aluminum atoms oxidize to Al, releasing electrons3+-Ions. These react with the water and form aluminum oxide Al2O3. At the same time, H3O+-Ions free:
     
   
To prevent corrosion, metals are also coated with a more noble metal. The gold plating of electronic contacts or the tin coating on tin cans made of sheet iron are examples of this. However, the layer only protects as long as the coating is intact. If the tin layer on the cans is damaged, a local element can form under the influence of water or acid.

    

Hot-dip galvanizing of steel girders  
  

But you can also apply a less noble metal as corrosion protection. This is the case with chrome plating. Chromium is considered to be extremely corrosion-resistant, as it immediately forms a protective oxide layer with the oxygen in the air and forms very corrosion-resistant alloys with other metals. Another example is galvanizing. The zinc coating can be carried out galvanically or through Hot dipping in a liquid zinc bath at the Hot-dip galvanizing. A corrosion-resistant alloy is formed at the transition between the steel and the zinc. At the same time, the zinc serves as a Sacrificial anode, since the less noble zinc corrodes first before the iron oxidizes. The magnesium sacrificial anodes serve the same purpose on large ships or on pipelines.
    

Related pages 
Redox reaction
Normal potential
Batteries and accumulators
Copper refining