Electrolysis: driving chemical reactions by electricity Chapter 28

Chemical energy from electrical energy We have seen that spontaneous redox reactions can be used as a source of electrical energy. We can cause redox reactions to occur by the passage of electrical energy from a power supply through a conducting liquid. This process is called electrolysis.

Electrolysis Electrical energy is converted into chemical energy. The reactions that occur in electrolytic cells are essentially the opposite to those occurring in galvanic cells. Reactions in electrolytic cells would not normally happen without the application of electrical energy, and so they are called non-spontaneous reactions. Chemicals formed by electrolysis are often difficult to obtain by other means.

Important applications of electrolysis Plating a thin film of metal on surfaces of other metals to improve the appearance or prevent corrosion. Extraction of reactive metals such as sodium and aluminium from their ores. Industrial production of sodium hydroxide, copper, chlorine and hydrogen Recharging of car batteries and other rechargeable cells, such as lithium batteries Increasing the thickness of the surface oxide layer on aluminium metal to improve its resistance to corrosion.

Electroplating We tend to take tin cans for granted. They are not really tin cans, but are mainly composed of steel with a thin surface coating of tin only a fraction of a millimetre thick that has been applied by electrolysis. Tin corrodes very slowly and prevents contact between the iron, moisture and air. It therefore prevents rusting.

Electroplating The deposition of a layer of metal on the surface of another metal by electrolysis is called electroplating. It is performed in an electrolytic cell such as the simplified one used for tin plating as shown. The object to be plated is connected by a wire to the negative terminal of a power supply. This is the negative electrode of the cell. It is immersed in a solution, such as tin nitrate solution, which contains ions of the metal that is to form the plating. This is the electrolyte An electrode of tin metal is connected to the positive terminal of the power supply.

Electroplating When the cell is in operation the power supply acts as an ‘electron pump’ pushing electrons onto the negative electrode and removing electrons from the positive electrode.

Electroplating At the negative electrode: The negative terminal of the power supply pushes electrons along the wire to the electrode. Tin ions are attracted to the electrode, accept electrons and are converted to tin metal: Sn2+(aq) + 2e- → Sn(s) A coating of tin forms on the object. Since reduction of the tin ions has occurred, the negative electrode is acting as a cathode.

Electroplating At the positive electrode: The positive terminal of the power supply withdraws electrons from the tin electrode. A reaction that released electrons must occur. Tin metal slowly dissolves as Sn2+ ions are formed: Sn(s) → Sn2+(aq) + 2e This reaction replaces the Sn2+ ions in solution that were consumed by the reaction at the negative electrode. Since an oxidation reaction has occurred the positive electrode is the cathode.

Electroplating The excess of positive charge caused by the production of Sn2+ ions at the anode will be balanced by movement of NO3- ions in the electrolyte to that region and by movement of Sn2+ ions away from it. At the cathode, loss of positive charge causes cations to migrate towards it and anions away from it. It is still cations towards the cathode and anions away from it.

Other examples of electroplating Iron is widely used because it is cheap and strong. Electroplating with iron improves their resistance to corrosion. Electroplating jewellery makes an object more attractive. Chromium is used for decoration, non-tarnishing and hard-wearing.

Your Turn Page 448 Question 1 and 2

Faraday’s Laws – The first law of electrolysis When electroplating an object we need to know: How can I determine how much metal is being plated? How long should I leave the object being plated in the electroplating cell? What size current should I use We can determine the electric charge needed to be used by using this law.

Faraday’s Laws – The first law of electrolysis Electric charge, is given the symbol Q and is measured in coulomb. The relationship is: charge(coulombs) = current(amps) x time(sec) Q = I x t Look at the table on page 449 for the relationship between charge and mass

Faraday’s Laws – The first law of electrolysis We can link this to mass as the mass of metal produced at the cathode is directly proportional to the quantity of electricity passed through the cell. This relationship is known as Faraday’s first law of electrolysis. It may be written as: m α Q

The second law of electrolysis In order to deposit one mole of silver from a solution of Ag+ ions, just one more of electrons is required: Ag+(aq) + e- → Ag(s) The charge on one mole of electrons is 96 500C The charge on one mole of electrons is known as a faraday, and given the symbol F. 1 Faraday = 96 500 coulomb mol-1.

The second law of electrolysis The charge on a given number of moles of electrons, n(e-), may be calculated by: Q = n(e-) x F In order to produce one mole of a metal, one, two, three, or another whole number of moles of electrons must be consumed. This is Faraday’s second law of electrolysis.

Worked Examples 28.3a A silver-plating cell operates with a steady current of 30.0 A for 20.0 minutes. What mass of silver is plated on the object at the cathode? 28.3b How long would it take to deposit 50.0 g of copper at the cathode of the copper-plating cell operating at a current of 8.00 A?

Your Turn Page 451 Question 3 - 6

Competition at electrodes Knowing that oxidation occurs at the anode and reduction occurs at the cathode it is possible to predict the electrode reactions in a particular electrolytic cell. Usually there are several chemicals present at each electrode and you must decide which of the possible reactions have the greatest tendency to occur.

During Electrolysis The highest reaction in the electrochemical series that can occur in the forward direction is likely to occur at the cathode. That is the strongest oxidant will usually react at the cathode. The lowest reaction in the electrochemical series that can occur in reverse is likely to occur at the anode. That is the strongest reductant will usually react at the anode. It is the opposite to the Galvanic Cell

Worked Example 28.4 Predict the products of electrolysis of 1 M nickel sulfate solution using copper electrodes

Your Turn Page 452 Question 7

Useful Chemicals From Electrolysis To be commercially viable, electrolysis processes must: Be located in places with access to cheap, abundant and reliable electrical power. Use electrolytes that are cheap and readily available Where possible use electrodes made from inexpensive materials such as carbon or iron Be sited in locations that allow relatively inexpensive transport of materials and products Operate continuously – it should not be necessary to stop the cells in order to add electrolyte or remove products Involve low electrolyte temperatures to reduce heating costs and minimise wear on equipment Consider the environmental impact of the materials used in the construction and operation of the electrolytic cell.

Commercial cells using molten electrolytes Electrolysis is a useful method for the production of highly reactive elements. The electrolytes used in the production of these metals are molten salts rather than solutions of the salt in water. The electrochemical series shows that water is a stronger oxidant than the metal ions for Na, K, Ca and Al. If water were present it would be reduced preferentially at the cathode and no metal would form The use of a molten electrolyte usually involves greater energy expenditure and greater wear on the cell.

Sodium Sodium is extracted from molten sodium chloride by electrolysis. A simple cell of molten salt is shown above. The electrodes are made of an unreactive conducting material such as graphite or platinum. The reaction occurring here is the reverse of the reaction that occurs spontaneously between sodium and chlorine to form sodium chloride. Electrical energy from the power supply has caused a non- spontaneous reaction to occur

Sodium At the cathode (-) The power supply pushes electrons towards this electrode. Positive sodium ions accept electrons and become sodium atoms: Na+(l) + e- → Na(l) Sodium is solid at room temp, but is liquid at the temperatures required to melt NaCl. It is less dense than molten NaCl and floats to the top of the cell.

Sodium At the anode (+): Electrons are withdrawn from this electrode by the power supply. Since the electrode is made of unreactive material, the electrode itself will not participate in an oxidation reaction to supply electrons. Instead Cl- ions in the electrolyte give up electrons and form chlorine atoms. These atoms quickly form molecules of Cl2 and bubbles of chlorine gas form at the electrode. 2Cl-(l) → Cl2(g) + 2e-

Sodium 2Na+(l) + 2Cl-(l) → 2Na(l) + Cl2(g) Since Cl2 is a strong oxidant and Na is a strong reductant, there must be no contact between them; otherwise they will reform sodium chloride and the products from the electrolysis will be lost. This will certainly occur in the simple cell.

Sodium In order to produce sodium and chlorine commercially, a modified cell, known as the Downs cell has been developed to minimise contact between the two products. A screen is used to prevent contact between the chlorine gas formed at the anode and the sodium formed at the cathode. Read about aluminium on page 455

Anodising We have seen that the oxide layer on the surface of aluminium metal restricts further reaction of the metal with other chemicals. The corrosion resistance and hardness of this layer can be further increased by a process known as anodising. In this process, the oxide layer is made about 1000 times thicker by connecting the metal to the anode in an electrolytic cell with sulfuric acid as the electrolyte. The thick oxide film that forms is easily dyed to give an attractive appearance.

Commercial cells using aqueous electrolytes Compared with cells using molten electrolyte such as used to produce sodium and aluminium, electrolytic cells using aqueous electrolytes tend to use less energy and have lower operating costs. Have a read of pages 457 – 459 about the production of chlorine and sodium hydroxide and the electrorefining cell used for purifying copper metal.

Your Turn Page 459 Question 8 Page 460

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