Chemistry

Key unit competence:

Predict the products of given electrolytes during electrolysis and work out quantitatively to determine how much is liberated at a given electrode using Faraday’s law.

Learning Objectives:

• Define electrolysis, cathode and anode.

• Explain the electrolysis of different substances.

• State Faraday’s laws and define the Faraday’s constant.

• Develop practical experimental skills related to electrolysis, interpret results, and draw valid conclusions.

• Carry out a practical activity to explain the phenomenon of electrolysis.

• Compare the electrolysis of dilute solutions and concentrated solutions.

• Calculate the masses and volumes of substances liberated during electrolysis.

• Relate the nature of electrode, reactivity of metal ion in solution to the products of electrolysis.

• Perform electroplating of graphite by copper metal

Introductory activity

Observe carefully the figure below and answer the following question. Record your answer and discuss them.

1. Label the set up and give the name of this Experiment.

2. Suggest how water can be decomposed into hydrogen and oxygen.

13.1. Definition of electrolysis and Description of electrolytic cells.

Activity 13.1:

A.1. In one case, you have a source of water at the top of a hill and you want to supply water to a community in the valley down the hill

2. In another case, you have a community at the top of a hill and you want to supply water to the community from a source located in the valley down. Students in groups discuss how they would proceed to supply water to the communities in the above two cases.

B. Why do we cook food by heating?

C. What is the difference between a spontaneous reaction and a non-spontaneous reaction?

D. Have you heard about electrolysis? If yes, can you say what it is about?

1.Definition of electrolysis.

A spontaneous reaction is a reaction that favors the formation of products without external energy. It is a process that will occur on its own. For example, a ball will roll down an incline, water will flow downhill, radioisotopes will decay, and iron will rust. No intervention is required because these processes are thermodynamically favorable.

A nonspontaneous reaction (also called an unfavorable reaction) is a chemical reaction that necessitates external energy to occur. For example, without an external energy source, water will remain water forever. Under the right conditions, using electricity (direct current) will help to produce hydrogen gas and oxygen gas from water. Cooking foods is not spontaneous reaction that is why heat is used.

Electrolyte: Sodium chloride is an ionic compound in which ions arrange themselves in a rigid cubic lattice when in solid state. In this state, it cannot allow electric current to pass through it. However, when it is melted, or dissolved in water, the rigid lattice is broken, ions are free to move and electric current can pass. Therefore, it is classified as an electrolyte.

Substances which cannot allow the flow of electriccurrent when in molten or in solution are referred to as non-electrolytes. When electric current (direct current) flows through an electrolyte, it decomposes it. This phenomenon is called electrolysis.

Thus electrolysis is the decomposition of an electrolyte by passage of an electric current through it. Therefore for electrolysis to take place, there must be a source of direct current. The direct current is conveyed from its source to the electrolyte by means of a metallic conductor and electrodes. The electrode connected to the positive terminal of the direct current is called the anode and the one connected to the negative terminal is the cathode. By convention, the electric current enters the electrolyte by the anode and leaves by the cathode.

When the current passes through an electrolytic solution, ions migrate and electrons are gained or lost by ions on the electrodes surface. Electrode that is positively charged has deficit of electrons is called anode and the other electrode negatively charged has excess of electrons and is called cathode. Chemical changes at the electrodes due to the passage of electric current are called electrolysis.

2. Description of electrolytic cells.

An electrolytic cell is an electrochemical cell that drives a non-spontaneous redox reaction through the application of externalelectrical energy. They are often used to decompose chemical compounds, in a process called electrolysis . The Greek word lysis means to break up.

Important examples of electrolysis are the decomposition of water into hydrogen and oxygen, production of sodium metal, Na, from molten NaCl, production of aluminium and other chemicals. Electroplating (e.g. of copper, silver, nickel or chromium) is done using an electrolytic cell.

An electrolytic cell has three component parts: an electrolyte and two electrodes (a cathode and an anode). The electrolyte is usually a solution of water or other solvents in which ions are dissolved. Molten salts such as sodium chloride are also electrolytes. When driven by an external voltage applied to the electrodes, the ions in the electrolyte are attracted to an electrode with the opposite charge, where charge-transferring (also called faradaic or redox) reactions can take place. Only with an external electrical potential (i.e. voltage) of correct polarity and sufficient magnitude can an electrolytic cell decompose a normally stable, or inert chemical compound in the solution. The electrical energy provided can produce a chemical reaction which would not occur spontaneously otherwise.

The main components required to achieve electrolysis are:

• An electrolyte is substance containing free ions which are the carriers of electric current in the electrolyte. If the ions are not mobile, as in a solid salt then electrolysis cannot occur.

• A direct current (DC) supply: provides the energy necessary to create or discharge the ions in the electrolyte. Electric current is carried by electronds in the external circuit.

• Electrolysis depends on controlling the voltage and current.

• Alternating current (AC) would not be appropriate for electrolysis. Because the “cathode” and “anode” are constantly switching places, AC produces explosive mixtures of hydrogen and oxygen.

• Two electrodes: an electrical conductor which provides the physical interface between the electrical circuit providing the energy and the electrolyte.

The key process of electrolysis is the interchange of atoms and ions by the removal or addition of electrons from the external circuit. The products of electrolysis are in some different physical states from the electrolyte and can be removed by some physical processes.

Electrodes of metal, graphite and semi-conductor material are widely used. Choice of suitable electrode depends on chemical reactivity between the electrode and electrolyte and the cost of manufacture.

Note:

The suitable electrode in electrolysis should be inert (Cu, Pt, etc.) therefore it will not participate in the chemical reaction.

It is very easy to be confused about the names CATHODE and ANODE and what their properties are, both with electrochemical cells and electrolytic cells.

(To help you to remember, Cathode is the site of reduction, or, if you prefer, CCC = Cathode Collects Cations. Anode is the site of oxidation, or, AAA = Anode Attracts Anions.)

Checking up 13.1:

Choose the correct answer from the options given below each of the following questions:

1) Which of the following substances is an electrolyte?

a) Mercury

b) Copper

c) Sodium sulphate

d) Aluminium

2) Which of the following substances is a weak electrolyte?

a) Dilute hydrochloric acid

b) Dilute sulphuric acid

c) A solution of potassium bromide.

d) Carbonic acid

3) Which of the following statements is true for the formation of sodiumchloride by the direct combination of sodium with chlorine?

a) Sodium is reduced

(b) Chlorine is oxidized.

c) Chlorine is the oxidising agent

(d) Sodium is the oxidizing agent.

4) Which of the following species will be deposited at the cathode on theelectrolysis of an aqueous solution of potassium bromide?

a) K

b) H₂

c) Br₂

d) O₂

5) If you want to electrolyse concentrated HCI, which of the following will youchoose for making the anode?

a) Graphite

b) Aluminium

c) Iron

d) Copper

13.2. Electrolysis of sodium chloride

Activity 13.2

Investigating the effect of concentration on the products formed during electrolysis of concentrated sodium chloride solution.

Materials: Carbon or graphite rods, connecting wire, U-tube, dry cell, glass syringes, concentrated sodium chloride, cork and switch.

Procedure:

2. Add 10g of sodium chloride to 100cm3 of distilled water.

3. Warm the mixture and continue adding sodium chloride until a saturated solution is formed.

4. Put the saturated solution in U-tube and fit it with carbon rods and glass syringes.

5. Level the brine solution in the two arms and switch on the circuit. Record any observations made after some time. Identify any gases collected in the syringe.

Questions:

a. Identify the gases formed by testing them using litmus papers.

b. Using ionic equations, explain how the products are formed.

Sodium chloride may be in different forms that can be electrolytes. It may be in its molten state, dilute solution or concentrated solution. In each case, the products of electrolysis differ because of different factors.

13.2.1. Electrolysis of molten sodium chloride

The molten salt is introduced in a container called electrolytic cell (or electrolysis cell) in which there are two inert electrodes (platinum or graphite). Electrodes are connected to a DC generator.

• Cations (Na⁺) move toward the cathode (negative electrode), where they take electrons and are reduced. On cathode metallic sodium is deposited:

• Anions (Cl^-) move towards the anode (positive electrode), where they give up electrons and are oxidized. On the anode Cl₂ is released:

The overall reaction is the addition of half reactions at electrodes:

The cathode provides electrons so it is a reducing site.

The anode takes electron, so it is an oxidizing site.

• Another important thing to note is that twice as much hydrogen is produced as oxygen. Thus the volume of hydrogen produced is twice that of oxygen. Refer to the equations above and note the number of electrons involved to help you

13.2.2. Electrolysis of Dilute Sodium Chloride Solution

An aqueous solution of sodium chloride contains four different types of ions. They are ions from sodium chloride: Na⁺ (aq) and Cl^- (aq) Ions from water: H⁺ (aq) and OH^- (aq)When dilute sodium chloride solution is electrolysed using inert electrodes, the Na⁺ and H⁺ ions are attracted to the cathode. The Cl^- and OH^- ions are attracted to the anode.

The table shown below is simply a table of standard reduction potentials in decreasing order. The species at the top have a greater likelihood of being reduced while the ones at the bottom have a greater likelihood of being oxidized. Therefore, when a species at the top is coupled with a species at the bottom, the one at the top will be easily reduced while the one at the bottom will be oxidized.

• At the cathode:

The H⁺ and Na+ ions are attracted to the platinum cathode. H⁺ ions gains electrons from the cathode to form hydrogen gas. (The hydrogen ions accept electrons more readily than the sodium ions. As a result, H⁺ ions are discharged as hydrogen gas, which bubbles off. Explanation why H+ ions are preferentially discharged will be given later.)

2H⁺(aq) + 2e^-→ H₂(g) ,Na⁺ ions remain in solution.

• At the anode:

OH^- and Cl^- are attracted to the platinum anode. OH^- ions give up electrons to the anode to form water and oxygen gas.

Note:

• Since water is being removed (by decomposition into hydrogen and oxygen), the concentration of sodium chloride solution increases gradually. The overall reaction shows that the electrolysis of dilute sodium chloride solution is equivalent to the electrolysis of water.

understand.

13.2.3. Electrolysis of Concentrated Sodium Chloride Solution

The only difference with dilute NaCl solution is that at the anode, Cl^- ions are more numerous than OH^- ions. Consequently, Cl^- ions are discharged as chlorine gas, which bubbles off. A half-equation shows you what happens at one of the electrodes during electrolysis.

Sodium ions Na⁺ and hydroxide OH^– are also present in the sodium chloride solution. They are not discharged at the electrodes. Instead, they make sodium hydroxide solution.

These products are reactive, so it is important to use inert (unreactive) materials for the electrodes.

One volume of hydrogen gas is given off at the cathode and one volume of chlorine gas is produced at the anode. The resulting solution becomes alkaline because there are more OH^- than H⁺ ions left in the solution.

Checking up 13.2:

With the help of equations of reactions which occur at each electrode, outline what happens during electrolysis of dilute aqueous sodium chloride. What happens to the pH of the solution as electrolysis continues?

13.3. Electrolysis of water

Activity 13.3: Investigate the products formed during the electrolysis of water

Materials:

• Distilled water

• Tap water

• 2 silver-colored thumb tacks

• 9V battery

• Small, clear plastic container

• 2 test tubes

• Stopwatch

• Baking soda

• Table salt

• Lemon

• Dish washing detergent

Procedure:

3. Insert the thumb tacks into the bottom of the plastic container so that the points push up into the container. Space them so that they’re the same distance apart as the two terminals of the 9V battery. Be careful not to harm yourself!

4. Place the plastic container with the terminals of the battery. If the cup is too large to balance on the battery, be sure thumb tacks are connect to positive and negative pushpins and do no touch each other.

5. Slowly fill the container with distilled water. If the tacks move, go ahead and use this opportunity to fix them before you proceed. Will distilled water conduct electricity on its own? Try it!

6. Add a pinch of baking soda.

7. Hold two test tubes above each push pin to collect the gas being formed. Record your observations. What happens? Does one tube have more gas than the other? What gases do you think are forming?

8. Discard the solution, and repeat the procedure with a different combination:

• Distilled water and lemon juice

• Distilled water and table salt

• Distilled water and dish detergent

• Distilled water (no additive)

• Tap water (Does tap water works? If so, why?)

Question: During the electrolysis of water, which electrolyte conducts electricity the best?

Water can be decomposed by passing an electric current through it. When this happens, the electrons from the electric current cause an oxidation-reduction reaction. At one electrode, called the cathode, electrons cause a reduction. At the other electrode, called the anode, electrons leave their ions completing the circuit, and cause an oxidation.In order to carry out electrolysis the solution must conduct electric current. Pure water is a very poor conductor.

To make the water conduct better we can add an electrolyte (NaCl) to the water. The electrolyte added must not be more electrolyzable than water. Many electrolytes that we add electrolyze more easily than water. Sulfate ions do not electrolyse as easily as water, so sulfates are often used to enhance the conductivity of the water

Water may be electrolyzed in the apparatus shown below. Pure water is however a very poor conductor of electricity and one has to add dilute sulphuric acid in order to have a significant current flow.

The electrodes consist of platinum foil. The electrolyte is dilute sulphuric acid.

Hydrogen gas is evolved at the cathode, and oxygen at the anode.

The ratio, by volume, of hydrogen to oxygen, is exactly 2:1.

Remember that electron flow in the circuit is opposite to the conventional current flow.

The reaction at the cathode (tube A) is the reduction of protons:

Oxidation takes place at the anode (tube B). There are two anions competing to give up their electrons, namely sulphate (SO4₂^-) from sulphuric acid, and hydroxide (OH^-) from the ionization of water. Here, the activity series is used to know the ion to be discharged.

The oxidation of OH^- according to the reaction:

In pure water at the negatively charged cathode, a reduction reaction takes place, with electrons (e⁻) from the cathode being given to hydrogen cations to form hydrogen gas. The half reaction, balanced with acid, is:

Combining either half reaction pair yields the same overall decomposition of water into oxygen and hydrogen:

The number of hydrogen molecules produced is thus twice the number of oxygen molecules. Assuming equal temperature and pressure for both gases, the produced hydrogen gas has therefore twice the volume of the produced oxygen gas.

Checking up 13.3

I understand the process of water electrolysis, that water as an electrolyte can be decomposed into hydrogen and oxygen via an external energy source (an electrical current). I know that the reduction of hydrogen takes place on the cathode and the oxidation reaction takes place on the anode. I also know that water, already partially split into H⁺ and OH^- (though there are very few of these ions in pure water).

1. Electric current (direct current) electrolyzes water. Discuss this statement.

2. Why alternative current are not used for the same process?

13.4. Electrolysis of concentrated copper (II) sulphate solution using inert electrode

Activity 13.4:

Investigating what happens when a solution of copper (II) sulphate is electrolysed using carbon and copper electrodes.

Apparatus and chemicals: Glass cell, Carbon rod, 2M copper (II) sulphate solution, connecting wires, dry cells, copper plates, propanone and litmus paper.

Procedure:

3. Determine the mass of the graphite rods and record it.

4. Put 0.5M of copper (II) sulphate solution in a glass cell with the carbon (graphite) rods and set up the apparatus. Carefully observe all the changes taking place at the electrodes and the solution. Test the resulting solution with blue litmus paper.

5. After some time, switch off the current, remove the electrodes, wash them in propanone, dry then and then weigh them.

6. Repeat the experiment using clean strips of copper metal as electrodes. Weigh them and then complete the circuit using freshly prepared copper (II) sulphate solution. Record your observations.

Questions:

1. Explain the changes observed during the electrolysis of copper (II) sulphate using :

a.Carbon electrodes

b.Copper electrodes

2. Outline the changes that occur in the solution from the beginning to the end of the experiment.

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

Using the simple apparatus (diagram above) and inert carbon (graphite) electrodes, you can observe the products of the electrolysis of copper sulfate solution are a copper deposit on the negative cathode electrode and oxygen gas at the positive anode electrode. This anode reaction differs if you use copper electrodes. You have to fill the little test tubes with the electrolyte (dilute copper sulphate solution), hold the liquid in with your finger and carefully invert them over the nearly full electrolysis cell. The simple apparatus (above) can be used with two inert wire electrodes.

The blue colour fades as more and more copper is deposited, depleting the concentration of blue copper ion Cu²⁺ in solution.

The electrode reactions and products of the electrolysis of the electrolyte copper sulphate solution (with inert carbon-graphite electrodes) are illustrated by the diagram above

(a) The electrode products from the electrolysis of copper sulphate with inert graphite (carbon) electrodes

The negative cathode electrode attracts Cu²⁺ ions (from copper sulphate) and H+ions (from water). Only the copper ion is discharged, being reduced into copper metal. The less reactive a metal, the more readily its ion is reduced on the electrode surface.

A copper deposit forms as the positive copper ions are attracted to the negative electrode (cathode)

The traces of hydrogen ions are not discharged, so you do not see any gas collected above the negative electrode. The blue colour of the copper ion will fade as the copper ions are converted into the copper deposit on the cathode

At the positive anode reaction with graphite electrodes

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

The negative sulphate ions (SO₄^2-) or the traces of hydroxide ions (OH^–) are attracted to the positive electrode. But the sulphate ion is too stable and nothing happens. Instead hydroxide ions are discharged and oxidised to form oxygen.

Checking up 13.4

1. Name the product at the cathode and anode during electrolysis of:

a. Molten lead bromide with inert electrode.

b. Acidified copper sulphate solution with inert electrodes.

c. Acidified water with inert electrode.

d. Dilute hydrochloric acid with inert electrode.

e. Concentrated hydrochloric acid with inert electrode.

2.Predict the products formed when the following molten compounds are electrolysed using carbon electrodes;

a. Lead(II) bromide

b. Magnesium oxide

13.5. Electrolysis of concentrated copper (II) sulphate solution using copper electrodes

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

The electrolysis of copper (II) sulphate solution using copper electrodes is shown below.

Using the simple apparatus and two copper electrodes the products of the electrolysis of copper sulphate solution are a copper deposit on the negative cathode electrode and copper dissolves, Cu⁺², at the positive anode electrode. This copper anode reaction differs from the one when you use an inert graphite electrode for the anode.

When Copper (II) sulphate 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 involve a two electron transfer 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 dryelectrodes 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.

The electrode reactions and products of the electrolysis of copper sulphate solution (with a copper anode) are illustrated by thediagram above.

(a) The electrode products from the electrolysis of copper sulphate with copper electrodes

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

Only the copper ion is discharged, being reduced to copper metal.A reduction electrode reaction at the negative cathode: (copper deposit, reduction 2 electrons gained) reduction by electron gain

The positive anode reaction with copper electrodes

Copper atoms oxidised to copper (II) ions: dissolving of copper in its electrolytic purification or electroplating (must have positive copper anode).

Copper (II) ionsreduced to copper atoms: deposition of copper in its electrolytic purification or electroplating using copper (II) sulphate solution, so the electrode can be copper or other metal to be plated or any other conducting material.

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 sulphate solution, the mass loss of copper from the positive anode electrode should equal tothe 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 the same as the loss of mass from the anode.

13.6. Faraday’s Laws

In an electrolysis where there are more than one species which can be discharged at the same electrode, only one of them is discharged at a time; for example, in an aqueous sodium chloride solution, we have four species that is,Na⁺ and Cl^- ions from sodium chloride and H⁺ and OH^- ions from water.

During electrolysis Na⁺ ions and H⁺ ions migrate to the cathode while Cl^-ions and OH^- migrate to the anode.

Now the question is, which species of ions will be discharged at the cathode and which ones will be discharged at the anode first?

The factors which decide the selective discharge of ions are:

• Nature of electrodes

• Position of the ion in electrochemical series

• Concentration

• The state of the electrolyte

13.7.2. Position of ion in Electrochemical Series

When solving for the standard cell potential, the species oxidized and the species reduced must be identified. This can be done using the activity series. The table 13.2 is simply a table of standard reduction potentials in decreasing order. The species at the top have a greater likelihood of being reduced while the ones at the bottom have a greater likelihood of being oxidized. Therefore, when a species at the top is coupled with a species at the bottom, the one at the top will become reduced while the one at the bottom will become oxidized.

During electrolysis of solution containing a mixture of ions, the ion lower in electrochemical series is discharged first in preference to the one high in the series.

13.7.3. Concentration of electrolyte solution

Increase of concentration of an ion tends to promote its discharge, for example in concentrated hydrochloric acid, containing OH^-(from water) and Cl^- as negative ions, the highly concentrated Cl^- is discharged in preference.

However, if the acid is very dilute, some discharge of OH^- will also occur. It is important to know that as the acid is diluted, there will not be a point at which chlorine ceases to be produced and oxygen replaces it. Instead a mixture of the two gases will come off, with the proportion of oxygen gradually increasing.

Another case in which the order of discharge according to the electrochemical series is reversed by a concentration effect is that of sodium chloride solution.

In concentrated solution of sodium chloride called brine, the following reactions occur.

Question 3

A university student set up three different electrolytic cells. The substances that were electrolysed were NaCl (l), 0.05 M NaCl (aq) and 5.0 M NaCl (aq). Which of the following statements correctly describes the results of the experiment?

a. The reactions occurring for the aqueous solutions will produce the same products at the anode and cathode.

b. Chlorine gas is the major product when molten NaCl (aq) and 0.05 M NaCl (aq) are electrolysed.

c. The pH at the cathode increases when solutions of NaCl are electrolysed.

d. The only means by which different products can be produced for varying concentrations of NaCl is to alter the voltage.

13.7.4. The state of the electrolyte

The half reactions taking place at the electrode depends on whether the electrolyte is in a molten or an aqueous state, and if in aqueous state its concentration. For example, the electrode reactions that take place during the electrolysis of molten potassium iodide are:

13.8. Application of electrolysis

Activity 13.8 Copper-Plated Key

Materials:

• 1.5-volt D batterywith battery holder

• Two alligator clip leads or insulated wire

• Beakeror glass

• Copper sulphate

• Copper electrode(or coil of copper wire)

• Brass key

• Safety equipment

Procedure:

5. Prepare the key for copper-plating by cleaning it with toothpaste or soap and water. Dry it off on a paper towel.

6. Stir copper sulphate into some hot water in a beaker until no more will dissolve. Your solution should be dark blue. Let it cool.

3. Use one alligator clip to attach the copper electrode to the positive terminal of the battery (this is now theanode) and the other to attach the key to the negative terminal (now called the cathode).

4. Partially suspend the key in the solution by wrapping the wire lead loosely around a pencil and placing the pencil across the mouth of the beaker. The alligator clip should not touch the solution.

5. Place the copper strip into the solution, making sure it doesn’t touch the key and the solution level is below the alligator clip. An electrical circuit has now formed and current is flowing.

6. Leave the circuit running for 20-30 minutes, or until you are happy with the amount of copper on the key.

Question: Observe carefully electrolysis process and records what happened during the electrolysis process.

Electrolysis has a number of important industrial applications. These include the extraction and purification of metals, electroplating and anodizing and the manufacture of other chemicals for example sodium hydroxide (NaOH).

Extraction of metals Metals in group I and II of the periodic table cannot be reduced by chemical reducing agents; they are extracted from their fused halides by electrolysis. Sodium is obtained by electrolysis of molten sodium chloride in the Dawncell.Magnesium is obtained by electrolysis of MgCl₂, generated from dolomite and sea water.

The electrolytic cell is an iron tank lined with carbon, which acts as the cathode. The anodes are blocks of carbon dipped into the electrolyte. The electrolyte is a solution of molten aluminum oxide in molten cryolite. Cryolite acts as a solvent to dissolve aluminium oxide and as an impurity to lower the melting point of aluminium oxide. The electrolytic cell is maintained at around 900°C