4(a) Introduction to the electrolysis of copper(II) sulfate solution

All copper sulfate electrolysis experiments are based on the principles illustrated in the diagram above.

Two electrodes, an electrolyte (conducting solution of copper ions and sulfate ions) and d.c. electricity supply.

More sophisticated apparatus for the electrolysis of copper sulfate solution are illustrated below.

Note the systematic name for this copper salt is copper(II) sulfate.

The electrolyte copper(II) sulfate, provides a high concentration of copper(II) ions Cu²⁺ and sulfate ions SO₄^2– to carry the current during the electrolysis process. There are tiny concentrations of hydrogen ions H⁺ and hydroxide ions (OH^–) from the self-ionisation of water itself, but these can be ignored in this experiment.

The electrolysis will only take place when electricity is passed through the copper ion solution.

The technical details of the electrolysis of copper sulfate solution with two different electrodes (a) inert graphite (carbon) electrodes and (b) copper electrodes are all explained below.

Electrolysis of a aqueous copper(II) sulfate solution CuSO₄(aq)

The products of electrolysing copper sulfate solution with inert electrodes (carbon-graphite or platinum) are copper metal and oxygen gas.

Using the simple apparatus (above left diagram) and inert carbon (graphite) or platinum electrodes, you can observe the products of the electrolysis of copper sulfate solution are (i) a copper deposit on the negative cathode electrode and (ii) oxygen gas at the positive anode electrode.

This anode reaction differs from when you use copper electrodes (see section (b) below). 

You have to fill the little test tubes with the electrolyte (dil. copper sulfate solution), hold the liquid in with your finger and carefully invert them over the nearly full electrolysis cell.

The very simple apparatus (above right) can be used with two inert wire electrodes.

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(b) The electrode products from the electrolysis of copper sulfate with inert graphite (carbon) electrodes

(or platinum electrodes if you can afford them!)

Note: The majority of liquid water consists of covalent H₂O molecules, but there are trace quantities of H⁺ and OH^– ions from the reversible self–ionisation of water:

H₂O(l) H⁺(aq) + OH^–(aq)

The half-equations for the electrolysis of copper(II) sulfate solution.

When the current is switched on, a copper deposit forms on the negative cathode and bubbles of the colourless oxygen come off the positive anode. The concept diagram below illustrates the process.

(2) The positive anode electrode reaction with graphite electrodes

Oxygen gas is formed at the positive electrode, an oxidation reaction (electron loss).

The negative sulfate ions (SO₄^2-) or the traces of hydroxide ions (OH^– from water) are attracted to the positive anode electrode.

But the sulfate ion is too stable and nothing happens, cannot be oxidised and discharged as oxygen.

Instead either hydroxide ions or water molecules are oxidised and discharged to form oxygen.

(i) 4OH^–_(aq) – 4e^– ===> 2H₂O_(l) + O_2(g)

The negative hydroxide ion is oxidised by electron loss

also written as:  4OH^–_(aq) ===> 2H₂O_(l) + O_2(g) + 4e^–

Use either of these equations

even though you can write an alternative equation involving oxidation of water

2H₂O_(l) – 4e^– ===> 4H⁺_(aq) + O_2(g)

The water molecule is oxidised by electron loss

also written as: 2H₂O_(l) ===> 4H⁺_(aq) + O_2(g) + 4e^–

BUT it is best to use the hydroxide ion oxidation equation because it fits in with the negative hydroxide ions migrating in the electrical field to the positive anode.

 

(ii) Test for the oxygen gas

The colourless gas should re-ignite a glowing splint - a simple test for oxygen.

 

AND in terms of the quantities of products (think moles)

(iii) The molar ratio of copper atoms to oxygen molecules is 2 : 1

To produce 1 molecule of oxygen requires the loss of 4 electrons, 1 from each of  hydroxide ions or water molecules.

These 4 electrons (i.e. the same current) can reduce 2 copper(II) ions to copper atoms.

Hence the ratio for the same current flow is 2Cu : O₂.

So

(iv) Overall balanced equation: 2Cu²⁺ +  4OH^–  ==>  2Cu  +  O₂  +  2H₂O

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4(c) The electrolysis of copper sulfate solution using copper electrodes

This is a method of purifying copper, copper plating and extracting other valuable metals from the anode sludge.

The products of electrolysing copper sulfate solution with copper electrodes are copper metal and copper ions (because the copper anode dissolves).

Using the simple apparatus (right diagram) and two copper electrodes the products of the electrolysis of copper sulfate solution are

(i) a copper deposit on the negative cathode electrode surface

(ii) copper dissolves from the positive anode electrode surface

This copper anode reaction differs from when you use an inert graphite electrode for the anode (see section (a) above).

When Copper(II) sulfate is electrolysed with a copper anode electrode (the cathode can be carbon or copper), the copper deposit on the cathode (–) equals the copper dissolves at the anode (+). Therefore the blue colour of the Cu²⁺ ions stays constant because Cu deposited = Cu dissolved.

Both half-reaction involve a two electron transfer (oxidation and a reduction) so it means mass of Cu deposited = mass of Cu dissolving for the same quantity of current flowing (flow of electrons).

You can check this out by weighing the dry electrodes before and after the electrolysis has taken place.

The experiment works with a carbon anode and you see the blackness of the graphite change to the orange-brown colour of the copper deposit and the anode becomes depleted in copper.

The electrode reactions and products of the electrolysis of copper sulfate solution with a copper anode are illustrated by the theory diagram above - it doesn't matter whether the cathode is carbon or copper - you get the same copper deposit and the copper anode is oxidised and dissolves to give the copper ion Cu²⁺(aq).

Electrode products from the electrolysis of copper sulfate with copper electrodes

Refer to the diagrams above when working through the reasoning of the half-reactions for the electrolysis of copper(II) sulfate solution explained below.

 

The half-equations

(1) The negative cathode reaction with copper electrodes or carbon electrodes

The negative cathode electrode attracts Cu²⁺ ions (from copper sulfate) and traces H⁺ ions (from water).

Only the copper ion is discharged, being reduced to a copper metal deposit on the cathode.

The less reactive a metal, the more readily its ion is reduced on the electrode surface. Copper ions are more readily reduced than hydrogen ions.

The reduction electrode reaction at the negative cathode

Cu²⁺_(aq) + 2e^– ===> Cu_(s)

A copper deposit forms, reduction of the copper ions to copper by electron gain, each Cu²⁺ ion gains 2 electrons.

Note on 'plating' - the formation of the copper deposit:

It doesn't matter what the cathode is made of, as long as it is a conducting material.

This is the basis of copper plating, and plating with any metal from a solution of its salt.

See section (d) The applications of electrolysis for electroplating in industry

 

(2) The positive anode reaction with a copper electrode

Its the copper anode that is the crucial difference than electrolysing copper sulfate solution with a inert carbon/graphite/platinum electrode.

The negative sulfate ions SO₄^2- (from copper sulfate) or the traces of hydroxide ions OH^– (from water) are attracted to the positive electrode.

But both the sulfate ion and hydroxide ion are too stable and nothing happens to them because the copper anode is preferentially oxidised to discharge Cu²⁺ copper ions into the electrolyte solution.

This is fairly unusual, because normally electrodes are 'inert', BUT, this technique is used in electroplating.

An oxidation electrode reaction at the positive anode

Cu_(s)  –  2e^– ===> Cu²⁺_(aq)

The copper dissolves after oxidation of the copper atoms,  each losing 2 electrons to form blue Cu²⁺ ions in solution - in this case the electrode is NOT inert.

also written as:  Cu_(s) ===> Cu²⁺_(aq)  +  2e^–

A balancing act !

copper atoms oxidised to copper(II) ions:

Dissolving of copper in its electrolytic purification or electroplating (must have positive copper anode).

The change involves two electrons per copper atom.

copper(II) ion reduced to copper atoms:

Deposition of copper in its electrolytic purification or electroplating using copper(II) sulfate solution, so the electrode can be copper or other metal to be plated OR any other conducting material.

The change involves two electrons per copper ion.

This means for every copper atom that gets oxidised, one copper ion is reduced, therefore ...

When copper electrodes are used in the electrolysis of copper sulfate solution, the mass loss of copper from the positive anode electrode should equal the mass of copper gained and deposited on the negative cathode electrode.

You can show this by weighing both electrodes at the start of the experiment.

After the current has passed for some time, carefully extract the electrodes from the solution, wash them, dry them and reweigh them.

The gain in mass of the cathode should be about the same as the loss of mass from the anode.

 

(3) In industry an anode sludge forms under the depleting impure block of copper

A deposit of dark material  gathers below the impure copper anode.

This is the residue left after the copper is oxidised, dissolves and transferred to the cathode.

In the electrolytic refining process, after the pure copper is deposited on the cathode plates insoluble impurities fall to the bottom of the cell as anode mud or anode sludge.

Anode sludge contains gold (Au) and other valuable metals like silver (Ag), platinum (Pt), and palladium (Pd).

These can be extracted from the anode sludge created by the electro-refining process.

In the formation of copper ore veins, copper concentrates often these precious metals and reclamation of these metals from anode slime is economically attractive and also it is environmentally friendly.

4(d) ELECTROPLATING - application of electrolysis in the chemical engineering industry

Metal electrodes dipped in aqueous salt solutions

For electroplating in general:

The negative cathode electrode is made the metal/conducting surface to be coated, and the positive anode electrode is made of the plating metal which dissolves and replaces any deposit formed on the cathode -which is the conducting article to be electroplated.

See also Extraction and purification of copper

INTRODUCTION TO ELECTROPLATING and its APPLICATIONS

- diagram and below it explanatory notes ...

This section below has some technical details e.g. the electrode equations, or go straight to the industrial applications of electroplating

As already described already the use of a copper positive anode electrode is the basis of the method of electroplating any conducting solid with a layer of copper which can be reproduced by electroplating other conducting materials with zinc (a way of galvanising steel), nickel, silver or chromium ('chromium plating'). Read on in conjunction with the theory diagram above describing the process of electroplating.

The CATHODE object to be electroplated must be a conducting material, usually a metal, and must be made the negative cathode electrode and completely immersed in the electrolyte solution.

The ANODE is usually a bar of the metal that is being electroplated onto the cathode object, giving a continuous supply of the coating metal and ensuring the concentration of electrolyte metal ion does not diminish as the electrolytic plating continues. The metal anode bar must be oxidised to provide a metal ion that can migrate across to the cathode and be discharged as the electrolysis takes place.

The electrolyte solution must contain ions of the metal that will form the electroplated deposit; and the ions come from an appropriate salt solution e.g. copper sulfate for copper, silver nitrate for silver, zinc sulfate for zinc or chromium chloride for chromium coatings.

The anode must be made of the metal that will form the electroplated coating on the positive anode object e.g. copper or silver.

As the metal is coated on the -ve cathode object, simultaneously the metal of the +ve anode is oxidised to refresh the solution of metal ions. so there is no depletion of the crucial ion concentration. These positive ions will migrate towards the negative electrode object to be coated.

The purification of copper by electrolysis amounts to copper plating so all you have to do is swap the pure negative copper cathode with the metal you want to coat (e.g. Ni, Ag or Au or any material with a conducting surface).

Swap the impure positive copper anode with any pure block of the metal you want to form the coating layer on the negative electrode object.

So any conducting (usually metal) object can be electroplated with copper, silver or gold for aesthetic reasons (decorative jewellery objects) or steel with zinc (galvanising) or a shiny chromium as anti-corrosion protective layer on steel. Any dull looking cheap metal can be made to look rather more shiny and attractive by electroplating. So cheap brass objects can be 'silver plated' and 'gold plated' to look more valuable that they really are!

Examples - half-reactions given, but read in conjunction with the general notes and diagram in the introduction.

 

(1) Copper electroplating (copper plating by electrolysis of a copper salt solution)

(-ve cathode electrode) Cu²⁺_(aq) + 2e^– ==> Cu_(s)

electron gain, reduction, copper deposited (electroplated) on the cathode object, dull object might look a lot prettier!

(+ve anode electrode) Cu_(s) – 2e^– ==> Cu²⁺_(aq)

supplies copper ions, electron loss, copper atoms oxidised

 

(2) Zinc electroplating (zinc plating by electrolysis)

a reduction electrode reaction at the negative cathode electrode in zinc salt solution

(– ve cathode electrode) Zn²⁺_(aq) + 2e^– ==> Zn_(s)

electron gain, zinc ion reduced, zinc deposit formed e.g. galvanising steel by electroplating

(+ve anode electrode) Zn_(s) ==> Zn²⁺_(aq) + 2e^–

zinc atoms of the positive zinc anode electrode are oxidised, electron loss, supplying more zinc ions

zinc ions reduced to zinc atoms: galvanising steel (the electrode) by electroplating from aqueous zinc sulfate solution, (or from molten zinc chloride?)

 

(3) Silver electroplating (silver plating by electrolysis)

a reduction electrode reaction at the negative cathode electrode in a silver salt solution

(– ve cathode electrode) Ag⁺_(aq) + e^–  ==> Ag_(s)

silver deposit as the silver ions are reduced to silver atoms, thereby electroplating the object, from cheaper metals like brass, to good looking silver ones and electroplated brass is much cheaper than pure silver and looks just as good!

(+ ve anode electrode) Ag_(s) ==> Ag⁺_(aq) + e^–

silver atoms oxidised on the surface of the silver anode, re-supplying the electrolyte with silver ions

You can do this using the electrolysis of silver nitrate solution.

Incidentally, as a school experiment, if you use lead nitrate solution you will get a coating of lead, despite lead being more reactive than hydrogen. BUT, who would want to coat anything with lead?!

(– ve cathode electrode) Pb²⁺_(aq) +  2e^–  ==> Pb_(s)

In both these cases in a school/college experiment you will get oxygen at the anode:

anode (+):   4OH^–_(aq) – 4e^– ==> 2H₂O_(l) + O_2(g)

However a solution of a gold salt is used to electroplate any other metal surface with a nice looking gold surface - but this is a bit costly for schools!

 

(4) Chromium electroplating (chromium plating by electrolysis)

a reduction electrode reaction at the negative cathode electrode in chromium(III) salt solution

(– ve cathode electrode) Cr³⁺_(aq) + 3e^–  ==> Cr_(s)

chromium deposit as the chromium ions from a chromium salt solution are reduced to chromium atoms, thereby electroplating the object, from cheaper metals like steel, to good looking shiny chromium plated ones!

 

(5) Tin electroplating (tin plating by electrolysis, 'tinning')

a reduction electrode reaction at the negative cathode electrode in a tin salt solution

(–ve cathode electrode) Sn²⁺_(aq) + 2e^– ==> Sn_(s)

electron gain, tin ion reduced, tin deposit formed.

 

(6) Nickel electroplating (nickel plating by electrolysis)

a reduction electrode reaction at the negative cathode electrode in a nickel salt solution e.g. nickel(II) sulfate

(–ve cathode electrode) Ni²⁺_(aq) + 2e^– ==> Ni_(s)

electron gain, nickel ion reduced, nickel deposit formed.

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4(e) Examples of APPLICATIONS of ELECTROPLATING

Please note that examples of electrode equations for plating are given in the previous sections on this page and note you can plate any conducting surface with a metal (less reactive) from its salt solution

Learning objectives for the electrolysis of copper sulfate solution with different electrodes and how this technique is applied to electroplating surfaces with other metals

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 : O₂) 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.