Cells and Batteries Chapter 27
Portable Power • A mobile phone, a laptop, an MP 3 player and a hearing aid all depend on small portable sources of electricity: cells and batteries. • Here we examine a range of common commercial galvanic cells to see how they use chemical reactions to produce electricity. • Read the case study on page 428
Primary Cells • Cells that can not be recharged are called primary cells. • They go flat when the cell reaction reaches equilibrium and you have to buy a replacement. • The products slowly migrate away from the electrodes and are consumed by side reactions occurring in the cell.
The zinc-carbon dry cell • The first mass-produced and widely used small-scale source of electrical energy. • Has changed very little in the last 100 years. • An electrolyte composed of a moist paste of zinc chloride and ammonium chloride plays the same role as the salt bridge.
The zinc-carbon dry cell • At the anode (-) oxidation of the zinc case produces electrons: Zn(s) → Zn 2+(aq) + 2 e- • At the cathode (+) Manganese dioxide is reduced in a complicated reaction that is thought to be: 2 Mn. O 2(s) + 2 NH 4+(aq) + 2 e- → Mn 2 O 3(s) + 2 NH 3(aq) + H 2 O(l)
The zinc-carbon dry cell • A new cell produces about 1. 5 volts, but this diminishes significantly during use. • To maintain a net forward reaction, the soluble reaction products must migrate away from the electrodes. • During use the build up of products around the electrodes slows and can even stop the forward reaction. • This is known as polarisation. If the cell is allowed to rest, the products migrate away from the electrodes and the cell can recover.
Alkaline cells • The alkaline cell is optimised for performance and longevity. • The alkaline cell is designed for more high capacity use appliances than the zinc- carbon dry cell.
Alkaline cells • At the anode Zinc powder around the central metal rod is oxidised: Zn(s) → Zn 2+(aq) + 2 e • Once formed, Zn 2+ reacts immediately with OHions in the electrolyte to form zinc hydroxide. The overall reaction at the anode is therefore written as: Zn(s) + 2 OH-(aq) → Zn(OH)2(s) + 2 e • At the cathode Manganese dioxide is reduced: 2 Mn. O 2(s) + H 2 O(l) + 2 e- → Mn 2 O 3(s) + 2 OH-(aq)
Alkaline cells • They have about 5 times the life of a zinccarbon dry cell. • There is no build up of electrolyte so no ‘rest’ time is needed. • They are slightly more expensive but offer better value for money.
Button Cells • Used in very small devices such as watches, some calculators and remote car locks. • There are two main types: silver-zinc cells and lithium cells. • Both are relatively expensive because of their small size and the materials used to make them. • Lithium cells produce about 3 volts during discharge and silver-zinc cells give an almost constant 1. 6 volts.
Examples • • Look at the two worked examples on page 431. Your Turn Page 432 Questions 1, 2, 3 and 5
Rechargeable cells and batteries • Rechargeable cells, such as lithium ion cells, are known as secondary cells or accumulators. • To recharge a cell, the products of the reaction must be converted back into the original reactants: the cell reaction must occur in reverse.
Rechargeable cells and batteries • This is done by connecting the cell to a ‘charger, ’ a source of electrical energy, which has a potential difference greater than the potential difference of a cell. • Electrical energy supplied by the charger is converted into chemical energy in the cell. • In order for it to be possible to regenerate the reactants, the products formed in the cell during discharge must remain in contact with the electrodes in a convertible form.
Car Batteries • Lead-acid batteries are the most widely used type of secondary cell. • They are relatively cheap and reliable, provide high currents, and have a long lifetime. • These are commonly known as car batteries, these are used to start a car’s engine and operate the car’s electrical accessories when the engine is not running. • An alternator provides electrical energy when the car is running.
Lead-acid batteries • Comprised of six separate cells connected together in series. • In a typical car battery each cell contains three positive electrodes sandwiched between four negative electrodes. • Contact between the electrodes is prevented by the presence of a porous separator. • The positive electrodes consist of a lead grid packed with Pb. O 2, while the negative electrodes consist of a lead grid packed with powdered lead. • A solution of sulfuric acid (about 4 M) acts as the electrolyte.
Lead-acid batteries • Each cell has a potential difference of just over 2 volts. • A car battery has six of these cells connected in series giving a total potential difference of about 12 volts.
Lead-acid batteries • At the anodes (-) Pb(s) + SO 42 -(aq) → Pb. SO 4(s) + 2 e- – These ions combine with sulfate ions from the electrolyte to form a coating of lead(II) sulfate on the electrodes (the white stuff you can sometimes see on the battery terminals) • At the cathodes (+) Pb. O 2(s) + SO 42 -(aq) + 4 H+(aq) + 2 e- → Pb. SO 4(s) + 2 H 2 O(l) – This reaction forms a coating of lead(II) sulfate on the electrodes
Lead-acid batteries • The two half equations can be combined as an overall equations: Pb(s) + Pb. O 2(s) + 2 SO 42 -(aq) + 4 H+(aq) → 2 Pb. SO 4(s) + 2 H 2 O(l) • The product of both electrode reactions, lead(II) sulfate, forms as a solid on the surface of the electrodes. This enables the battery to be recharged.
Recharging • To recharge the battery, the electrode reactions are reversed. • The alternator, with a potential difference of about 14 V, is used to force the electrons into the batteries negative terminal and draw them out at the positive terminal. • In effect driving a spontaneous reaction backwards. • The recharging process converts electrical energy into chemical energy. • What is the overall equation fro the reaction as it recharges?
Nickel-based cells and batteries • • Read about these cells on page 434. Your Turn Page 435 Questions 6 and 7
Fuel Cells • The major limitations of the cells we have examined so far is that they contain relativley small amounts of reactants. • When the reaction reaches equilibrium, they must either be recharged or discarded. • Cells can be constructed in which the reactants are supplied continuously, allowing constant production of electrical energy. • These devices are called fuel cells.
Fuel cells • They transform chemical energy directly into electrical energy. • This enables efficient use to be made of the energy released by spontaneous redox reactions. • Energy losses such as those that occur in a coal-fired power station are avoided. • They are up to 80% efficient compared with 30 -40% for thermal power stations. • In addition modern designs for fuel cells employ the waste heat they produce to make steam. • This steam can be sued for heating or to operate a turbine.
Fuel cells • The fuel cell used in the Apollo program used pure oxygen and hydrogen gas as reactants. • Potassium hydroxide solution was used as the electrolyte and the cell operated at about 250°C. • This is commonly referred to as an alkaline fuel cell.
Alkaline fuel cells • At the anode (-): – Hydrogen gas reacts with hydroxide ions from the electrolyte H 2(g) + 2 OH-(aq) → 2 H 2 O(l) + 2 e- • At the cathode (+): O 2(g) + 2 H 2 O(l) + 4 e- → 4 OH-(aq) • The overall equation is: 2 H 2(g) + O 2(g) → 2 H 2 O(l)
Fuel Cells • Each cell produces about one volt. • Higher voltages are obtained by connecting a number of fuel cells in series to form a battery. • The only by-products are water and heat. • The nature of the electrodes is crucial to successful operation of the cell as the function as catalysts for the reaction, and the size of the current depends on their surface area. • Scientists are striving to reduce the overall costs of cells and to improve the current that can be drawn by increasing the rate of reaction at the electrodes.
Transport • Most major vehicle manufacturers are investigating the use of fuel cells as an alternative to the internal combustion engine. • There a number of buses that use hydrogen-powered fuel cells. • The use of fuel cells improves fuel efficiency and reduces greenhouse gases and other emissions as well as our reliance on oil. • Look at page 438 for some other fuel cells
Advantages and disadvantages of fuel cells • • • read page 439 Your turn Page 439 Questions 9 and 10 Page 442 Questions 12, 13 and 21
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