69CHAPTER 4: Electrochemistry

Galvanic cells are usually classified into two groups, primary and secondary cells. Primary cells and batteries cannot be recharged. Once one of the reactants has been used up, these batteries no longer produce electricity—they become ‘dead’ and are discarded. Secondary cells and batteries, often called rechargeable or storage batteries, can be recharged many times by supplying electrical energy to reverse the cell reaction.

The dry cell and the lead–acid battery are two of the most widely used sources of portable electricity. Their impact on society has been significant and has led to the proliferation of portable devices we see around us. Nevertheless, new sources of portable power continue to be developed that provide more power for longer, are smaller, more lightweight and in general overcome some of the disadvantages of dry cells and lead–acid batteries. This unit investigates dry cells and lead–acid batteries in terms of their chemistry, cost and practicality, and their effect on society and the environment, and introduces some of the new generation of cells and batteries that may replace them.

Dry cells

Dry cells are so-called because in contrast to the galvanic cells discussed so far, they have electrolytes in the form of solids or pastes rather than liquids. Liquid electrolytes are prone to leakage and are therefore less suitable in many applications. Two types of dry cells are in everyday use: the ordinary acid dry cell (Leclanché cell) and the alkaline dry cell.Both acid and alkaline dry cells are relatively cheap compared to other types of portable power sources. Each type delivers approximately 1.5 volts and comes in a range of sizes to suit most applications, from small AAA cells through to large 9 V batteries. The impact of dry cells on society cannot be underestimated. A quick survey of your home will reveal many appliances, toys and devices that use dry cells as a source of power. Without dry cells we would be a much less mobile society. Devices requiring electrical power would need to be in close proximity to permanent power sources such as power points. Instead of torches we may still be using candles or kerosene lamps as a source of portable light, and many toys would be wind-up. The development of dry cells is directly responsible for the proliferation of the portable devices we have come to depend on. But these benefits come at some cost to the environment. Dry cells are primary cells and must be discarded once they are ‘dead’. Thus they end up in our landfills. Another problem is that the casing of dry cells is made from zinc, a heavy metal that is toxic to many animal species, particularly birds.

Leclanché cell

The ordinary dry cell or Leclanché cell commonly used in torches and other electrical appliances was developed in 1866 and has undergone very little modification since that time.In a dry cell, the outer zinc casing is the anode (–). The zinc oxidises according to the following half-equation:

Zn(s) Zn2+(aq) + 2e–

A graphite rod surrounded by a paste containing manganese dioxide is the cathode (+). The cathode reaction is as follows:

2MnO2(s) + 2H+(aq) + 2e– Mn2O3(s) + H2O(l)

In this half-reaction, the manganese is reduced from an oxidation state of (+4) to (+3).

+

–

Zinc anode

Graphitecathode

Moist pasteof MnO2,NH4Cland ZnCl2

(a) (b)

70 MODULE 1: Production of materials

The manganese dioxide paste also contains ammonium chloride and zinc chloride, which act as the electrolyte for the cell. The ammonium ions also provide the hydrogen ions needed for the cathode process.

NH4+(aq) L NH3(aq) + H+(aq)

The overall reaction in a dry cell can be represented in several ways, such as the following:

Zn(s) + 2NH4+(aq) + 2MnO2(s) ! Zn2+(aq) + Mn2O3(s) + H2O(l) + 2NH3(g)

or

Zn(s) + 2H+(aq) + 2MnO2(s) ! Zn2+(aq) + Mn2O3(s) + H2O(l)

The dry cell has a maximum cell potential (e.m.f.) of 1.48 V.

There are two main disadvantages associated with the dry cell. They have a relatively short shelf life because the zinc anode reacts with acidic NH4

+ ions. This may cause the cell to deteriorate and leak. Secondly, if the current is drawn rapidly from the cell, NH3(g) builds up, causing a drop in voltage. Despite these drawbacks, Leclanché cells are widely used in torches, portable radios, CD players and toys.

Alkaline dry cell

An alkaline version of the dry cell is also available and has somewhat superior performance to the acid form. The composition of the alkaline dry cell is similar to the Leclanché cell except that a powdered zinc anode is used and the electrolyte is 7 mol L–1 KOH instead of NH4Cl. The relevant half equations are:

Anode: Zn(s) + 2OH–(aq) ! Zn(OH)2(s) + 2e–

Cathode: MnO2(s) + 2H2O(l) + 2e– ! Mn(OH)2(s) + 2OH–(aq)

The alkaline dry cell has the same uses as a conventional dry cell but it has a longer working life and can supply current more rapidly than the conventional cell without a voltage drop. The alkaline dry cell is, however, more expensive to produce.

Figure 4.11 (a) A dry cell (b) The structure of an ordinary dry cell (Leclanché cell)

Cathode grid filled withlead dioxide

H2SO4electrolyte

Anode gridfilled withmetallic lead

–+

CHAPTER 4: Electrochemistry 71

The lead–acid battery

The lead–acid battery is commonly used in motor vehicles. In fact, its development revolutionised the automobile industry. Until batteries were used to power the starter motor, vehicles had to be cranked by hand. Much of the equipment we take for granted in modern vehicles, such as headlights, indicators, windscreen wipers, power windows, CD players and sophisticated computer systems, depends on the electrical energy stored in the vehicle’s battery. The main advantages of lead–acid batteries are that they can be recharged and are relatively cheap compared to alternative power sources.

The 12 volt storage battery, or accumulator, consists of six cells connected in series. Each cell has a cell potential (e.m.f.) of approximately 2 V. The cells are contained in a heavy-duty external casing. The electrodes in each cell consist of a bank of lead grids supporting a large surface area of the electrode material. The negative electrode (anode) grid is fi lled with spongy metallic lead and the positive electrode (cathode) grid is fi lled with brown lead dioxide. The electrolyte is sulfuric acid with a concentration of approximately 4.5 mol L–1.

The lead–acid accumulator is regarded as a storage battery, because electrical energy is stored as chemical energy in the recharging process. This stored chemical energy is released as electrical energy during discharge when the battery is used to start the vehicle, operate the lights and so on.

When discharging, the electrode reactions are as follows:

Anode: Pb(s) + SO42–(aq) ! PbSO4(s) + 2e–

Cathode: PbO2(s) + 4H+(aq) + SO42–(aq) + 2e– ! PbSO4(s) + 2H2O(l)

Overall: Pb(s) + PbO2(s) + 4H+(aq) + 2SO42–(aq) ! 2PbSO4(s) + 2H2O(l)

Figure 4.12 Lead–acid batteries are used in motor vehicles.

72 MODULE 1: Production of materials

As the battery runs down, the electrodes become coated with insoluble lead sulfate and the sulfuric acid is used up. Thus the density of the electrolyte decreases. A hydrometer is commonly used to check the degree of charge of a battery by testing the electrolyte density.

During recharging, the electrode reactions are reversed by connecting the terminals to another electrical source of higher voltage and reversing the direction of the electric current through the circuit. The overall equation in recharging is:

2PbSO4(s) + 2H2O(l) ! Pb(s) + PbO2(s) + 4H+(aq) + 2SO42–(aq)

In motor vehicles, recharging occurs during normal driving. This is achieved by a motor-driven generator, or alternator. If the battery runs fl at, as when the headlights are accidentally left on, an external recharger can be used.

There are two main disadvantages associated with lead–acid batteries. Because they are enclosed in a heavy casing, contain lead electrodes and several litres of sulfuric acid electrolyte, they are very heavy and this limits their use in many applications. The other problem concerns their disposal. Despite being rechargeable they have a limited life, usually about three years, and must then be replaced. Lead is a heavy metal and a cumulative poison. It builds up in the body over time and ultimately causes damage to the central nervous system, especially in younger children. Most car batteries are recycled and the lead electrodes within them are recovered. Nevertheless, many lead–acid batteries end up at the municipal tip, where they can contaminate the surrounding environment.

Alternative sources of portable power

Dry cells and lead–acid batteries are extremely useful and cost effective sources of portable power. Both have drawbacks, however, and therefore scientists continue their research into the development of more effi cient, lightweight cells and batteries. These use strong oxidants and reductants to produce maximum energy output. However, these materials are so highly reactive that new technologies such as solid polymer electrolytes must be employed. Some examples of this new generation of portable power sources are described in this section.

Button cells

The miniaturisation of electronic devices depends upon a corresponding decrease in the size of cells. Button cells, which are so named because of their compact size, have been developed for specifi c purposes where a small, lightweight power source is required. Two of the most commonly used button cells are the mercury(II) oxide–zinc and silver oxide–zinc cells.

Mercury(II) oxide cells

Contained in a stainless steel cylinder, mercury(II) oxide–zinc button cells use powdered zinc anodes and potassium hydroxide solution as the electrolyte. The anode process is:

Zn(s) + 2OH–(aq) ! Zn(OH)2(s) + 2e–

The reduction process at the cathode is:

HgO(s) + H2O(l) + 2e– ! Hg(l) + 2OH–(aq)

Mercury cells produce a cell voltage of 1.3 V and are ideal for applications where small size and light weight are important. They are commonly used in battery-powered watches and

Figure 4.13 Button cells are commonly used in calculators and watches.

Insulation

Steel (cathode)(+)

Paste of Ag2O on absorbentKOH and Zn(OH)2

Porousseparator

Zinc container(anode)

(–)

V

Anode Cathode

e– e–

Li

Li Li+ + e_ TiS2 + e_ TiS2–

TiS2

Solid electrolyte

Li+

Electrolyte

Water outlet

Porouscathode

Porousanode

Hydrogengas inlet

Oxygengas inlet

– +

73CHAPTER 4: Electrochemistry

scientifi c equipment such as calculators. The main disadvantages of the mercury(II) oxide button cell are that discarded cells can release toxic mercury into the environment and they are relatively expensive.

Silver oxide cell

The silver oxide cell is similar to the mercury(II) oxide cell. The anode process is the same:

Zn(s) + 2OH–(aq) ! Zn(OH)2(s) + 2e–

but the cathode process is the reduction of silver oxide to solid silver:

Ag2O(s) + H2O(l) + 2e– ! 2Ag(s) + 2OH–(aq)

Silver cells are used in cameras, heart pacemakers and hearing aids. They are very small and produce a very steady output of 1.5 V. In terms of disposal, silver oxide cells do not produce the highly toxic wastes associated with mercury(II) oxide cells but they are even more expensive.

Lithium cells

One of the exciting developments in battery research has been the production of solid-state lithium cells. These cells have an anode made of lithium, a lightweight metal that is a powerful reductant. The cathode is made of TiS2, a substance known as an ion-insertion compound. The electrode reactions are:

Anode: Li(s) ! Li+(s) + e–

Cathode: TiS2(s) + e– ! TiS2–(s)

While discharging, Li+ ions move through a solid polymer electrolyte which can conduct Li+ ions but not electrons. The electrons fl ow through an external circuit in (Figure 4.15).

Fuel cells

Fuel cells are part of a developing technology. They differ from primary and secondary galvanic cells in several ways. The major distinction is that fuel cells store neither the reactants nor the products. Their action is to supply electricity at a constant rate as reactants are fed into them and products are removed. They are direct energy converters. Another feature of fuel cells is that the reactants are usually gaseous, rather than the solid substances commonly used in primary and secondary cells.

The most successful fuel cell developed so far has been based on the combination of hydrogen and oxygen. The net reaction is the same as that for the combustion of hydrogen.

2H2(g) + O2(g) ! 2H2O(l)

Figure 4.16 shows one form of hydrogen–oxygen fuel cell. Hydrogen gas is supplied to the anode chamber and oxygen to the cathode chamber. The gases diffuse through the electrodes, which are porous metals such as platinum or nickel and act as catalysts.

Figure 4.14 A silver oxide ‘button’ cell

Figure 4.15 Schematic diagram of a solid-state lithium battery

Figure 4.16 A hydrogen–oxygen fuel cell

74 MODULE 1: Production of materials

The gases react with the electrolyte, which may be acidic or alkaline. For an acidic electrolyte, the reactions are given by the following equations:

Anode: H2(g) ! 2H+(aq) + 2e–

Cathode: O2(g) + 4H+(aq) + 4e– ! 2H2O(l)

If the electrolyte is alkaline, the reactions are:

Anode: H2(g) + 2OH–(aq) ! 2H2O(l) + 2e–

Cathode: O2(g) + 2H2O(l) + 4e– ! 4OH–(aq)

In both types of fuel cell the overall equation is the same:

2H2(g) + O2(g) ! 2H2O(l)

The theoretical e.m.f. of 1.23 volts is diffi cult to attain. One of the major problems is associated with the electrolyte. As the electrode reactions occur, the H+ ions in the acidic electrolyte must migrate from the anode to the cathode. In the alkaline electrolyte, OH– ions must move from the cathode to the anode. At normal temperatures, the rate at which these ions move is a limiting factor.

Another problem concerns the electrodes themselves. These must provide intimate contact between the gaseous reactants, the electrolyte and the catalyst. One development is to use graphite electrodes with large pores on the gas side and fi ne pores on the liquid electrolyte side. The fi nely divided metal catalyst adheres to the inside surface of the pores. Surface tension causes the liquid electrolyte to be retained in the fi ner pores while the gases diffuse into them from the larger pores.

Fuel cells have high fuel effi ciency and a relatively small mass. Another major advantage is their minimal impact on the environment. In a hydrogen–oxygen fuel cell, water and heat are the only products. Spacecraft including the Apollo and Space Shuttle programs have been using fuel cells for many years as a source of electrical energy and water. In spacecraft, the hydrogen and oxygen are stored as liquids until needed. The major disadvantages of fuel cells relate to cost. For example, the auxillary systems that feed the reactants into the fuel cell and remove products are quite expensive.

Figure 4.17 Space Shuttle fuel cell

+ –

Vanadium+ half-cell

storagetank

Vanadium– half-cell

storagetankCell

stack

Energy storage(kilowatts hours)

Energy storage(kilowatts hours)

Energy conversion(kilowatts)

Cell stack (2 cells)

Electrolyte tank

Electrolytetank

Current collector End plate

End electrode

Felt

Membrane

Felt Bipolar electrode

Catholyte V(V)/V(IV)

Anolyte V(II)/V(III)

Positive electrode: V(V) + e– V(IV) Negative electrode: V(II) V(III)+ e–

Overall: V(V) + V(II) V(IV) + V(III)

charge discharge

Pump

CHAPTER 4: Electrochemistry 75

Figure 4.18 Electrolytes are stored in two tanks and are pumped to the cell stack where energy conversion takes place.

Vanadium redox cells

The vanadium redox cell is a redox fl ow battery developed by the University of New South Wales. Redox fl ow batteries, also known as circulating-electrolyte cells, are rechargeable fuel cells. Because they offer greater effi ciency, less maintenance and a longer lifespan than other storage batteries such as lead–acid batteries, redox fl ow batteries are ideally suited to electrically powered vehicles. Used in conjunction with photovoltaic (solar) cells or other renewable energy systems, they could provide electrical power for residential and industrial applications. Vanadium redox batteries are currently undergoing trials in a golf cart and a solar demonstration house in Thailand.

A redox fl ow cell consists of two half-cells separated by a membrane. A number of redox cells are usually connected in series to form a cell ‘stack’. In a redox fl ow battery, two electrolyte solutions are pumped into a cell stack where their chemical energy is converted into electrical energy through a redox reaction. The basic concept of these batteries is shown in Figure 4.18.

In the vanadium redox battery, the electrolyte, which consists of vanadium ions in different oxidation states, can be charged or discharged by pumping it through the battery stack and either supplying electric power to the stack or taking power from the stack. Initially the positive and negative sides of the battery contain the same electrolyte. During charging, electrolysis within the battery stack changes the valency of the vanadium in the two electrolytes. Vanadium(IV) oxide (VO2) is changed to vanadium(V) oxide (V2O5) and vanadium(III) oxide (V2O3) is changed to vanadium(II) oxide (VO). During discharge when the battery is producing electricity, the process is reversed. Vanadium(V) oxide (V2O5) is reduced

Figure 4.19 A vanadium redox battery

76 MODULE 1: Production of materials

to vanadium(IV) oxide (VO2) in one half-cell, while in the other, vanadium(II) oxide (VO) is oxidised to vanadium(III) oxide (V2O3). The following half-equations represent the electrode reactions during discharge.

Cathode: V2O5(aq) + 2H+(aq) + 2e– L 2VO2(aq) + H2O(l)

Anode: 2VO(aq) + H2O(l) L V2O3(aq) + 2H+(aq) + 2e–

The overall equation for this reaction is:

V2O5 + 2VO L 2VO2 + V2O3

A porous, ion-selective membrane that only allows hydrogen ions to pass separates the electrolytes in the cell stack. The cell potential (e.m.f.) is 1.26 volts and its operating voltage is 1.4 volts.

The system can also be recharged by simply replacing the spent electrolytes. For an electric-powered vehicle this could involve offl oading the spent electrolytes and ‘fi lling up’ with fresh electrolytes at a vanadium refuelling station. At the refuelling station, the spent electrolytes could be recharged by connecting the battery to an external supply of electricity from the mains or from renewable energy sources.

Liquid junction photovoltaic cells: the Grätzel cell

The ability to convert light energy from the sun into electrical energy has long been the dream of scientists. In the 1960s, as part of the space program, scientists developed solar cells capable of converting light energy into electricity. These cells, known as photovoltaic cells, were manufactured from high-purity silicon. The problem with photovoltaic cells is that silicon is very expensive to process and the effi ciency of the cell in converting solar energy to electricity is low, typically 10–15%. Despite these limitations, solar cells have found a range of applications such as providing power for homes and outback installations, lighthouses, pumping water and emergency telephones.

New liquid junction photovoltaic cells, such as the Grätzel cell developed by Michael Grätzel and his colleagues in Lausanne, Switzerland, may revolutionise the solar energy industry. They have developed solar cells that are almost twice as effi cient at converting light energy into electrical energy as those currently available cells and, more importantly, are 80% cheaper to produce. Grätzel and colleagues have designed solar cells that mimic photosynthesis. Two very thin transparent electrodes enclose a thin layer of dye. This dye contains ruthenium ions coated onto tiny crystals of a semiconductor, titanium dioxide. The electrolyte surrounding the tiny crystals is a solution containing iodide ions. As the ruthenium dye absorbs visible light, photons of light cause electrons in the dye to become excited. These photoexcited electrons are pulled from the ruthenium dye by the titanium dioxide and are passed on to the electrode. The iodide ions provide a source of electrons to replace those knocked out from the dye by sunlight. When sunlight shines on this solar cell, a cell potential (e.m.f.) of 0.75 volts is produced. With the major obstacle of cost removed, these new-generation photovoltaic cells may signal the arrival of solar energy on a commercial scale.

Photon

Transparentelectrodes

Iodinesolution

Nanocrystals oftitania coated with

ruthenium dye

When light shines on the cell,electrons flow through thecircuit from one electrode

to another

e–

e–

e–

TiO2

O

O O

O

N N

Ru2+

77CHAPTER 4: Electrochemistry

1 Discuss how the development of the dry cell in 1866 has had an impact on society.

2 Draw a simple diagram of a conventional dry cell. Identify the anode, cathode and electrolyte on your diagram.

3 A dry cell is designed so that the MnO2 is concentrated around the carbon cathode and does not come into contact with the zinc.a Identify the reaction that would occur if the zinc and manganese dioxide were in direct contact.b Explain how this would this affect the usefulness of the cell.

4 Construct an overall equation for the redox reaction occurring in a Leclanché cell.

5 Use the separate half-equations for the lead–acid battery to identify the change in oxidation state of the lead in the two electrode processes.

6 Some small electric-start vehicles such as golf carts and ride-on mowers use a 6 V lead-acid battery. Predict the number of cells used to construct this battery.

7 Some solar-powered electric vehicles, such as those used in the World Solar Challenge held in Australia every three years, use lead–acid batteries to store energy. Many newer solar-powered cars now use alternative rechargeable batteries. Discuss the disadvantages of using lead–acid batteries in these vehicles.

! Review exercise 4.6

Figure 4.20 A section of a Grätzel cell