Q1. One major disadvantage of the dry cell was that the zinc casing would begin to disintegrate if it was left operating in an appliance for several weeks, potentially causing damage to the appliance. Explain why this occurs.
A1. In dry cells the zinc forms the casing. Since zinc is oxidised during the cell reaction, holes may develop in the case and the contents leak out.
Q2. Although very small alkaline cells can be made, they are not used in watches. Why are button cells preferred for this application?
A2. The electronic circuitry in watches requires an almost constant voltage for its operation. Alkaline cells do not provide a constant voltage, unlike silver–zinc and mercury–zinc button cells.
Q3. Write overall equations for the cell reactions in the dry cell, alkaline cell, and lithium and silver–zinc button cells using the anode and cathode reactions given in the text.
Q5. A galvanic cell is constructed by connecting a Cl2/Cl– half cell to a Zn2+/Zn half cell using a salt bridge at standard conditions. You will need to refer to Table 26.2 (p. 418). a Write an overall equation for the cell. b What is the polarity (positive or negative) of the zinc electrode? c Calculate the potential difference of the cell. d What happens to the concentrations of Cl– and Zn2+ ions when electric current is
drawn from the cell? e Explain the purpose of the salt bridge in the cell.
A5.
a Zn(s) + Cl2(g) → Zn2+(aq) + 2Cl–(aq) b Zinc is oxidised, generating negatively charged electrons. Hence the zinc
electrode is negative. c cell potential difference = higher half cell Eo – lower half cell Eo
= Cl2/Cl– half cell Eo – Zn2+/Zn half cell Eo = 1.36 – (–0.76) V = 2.12 V
d Both increase as the reaction proceeds in the forward direction until equilibrium is reached.
e Movement of ions through the salt bridge balances charges that are formed in the half cells during the cell reaction. This movement of charge completes the flow of current in the circuit.
E1. Write equations for the reactions that occur at the anode and the cathode of the sodium–sulfur cell as it discharges.
E2. Write an overall equation for the cell reaction during: a discharge b recharge
AE2. The overall reaction that occurs when the cell discharges is obtained by adding the anode and cathode reactions. In this case the anode reaction must first be doubled so that there equal number of electrons on both sides of the equation
2Na(l) + S(l) + 2e– → S2–(l) + 2Na+(l) + 2e–
This equation simplifies to 2Na(l) + S(l) → S2–(l) + 2Na+(l)
The reverse reaction occurs when the cell is recharged. S2–(l) + 2Na+(l) → 2Na(l) + S(l)
Worked solutions to student book questions Chapter 27 Cells and batteries
E3. Why do you think it is necessary for the cells to be heated in order for them to produce a voltage?
AE3. The operation of the cells requires that sodium anode, sulfur cathode and electrolyte are all in the molten state.
E4. Provided the cells are discharged and charged every day, they remain at working temperatures. Suggest the reason for this.
AE4. During recharging, some of the electrical energy entering the cell is transformed into chemical energy and some is transformed into heat. Heat energy is also released in the cell when it discharges. This heat is sufficient to maintain the cell at working temperatures if a discharge–recharge cycle is carried out every day.
E5. According to the article, what are the disadvantages of electric vehicles compared with petrol-driven vehicles at present?
AE5. Electric vehicles do not generate carbon dioxide, which is a greenhouse gas.
E6. Suggest the meaning of the term ‘energy density’. Why is energy density an important factor in the construction of electric vehicles?
AE6. The energy density of a cell is the energy available per unit mass of the cell. The greater the energy density of a cell, the greater the range of the electric vehicle that is using the cell.
E7. ‘Electric vehicles cause no pollution.’ Comment on the accuracy of this statement.
AE7. While electric vehicles cause almost no pollution at the site at which they are used, the electrical energy that is used to charge their batteries is produced by a series of energy transformations that involve the production of pollutants. Typically, fossil fuels will be burnt in order to produce the electrical energy.
E8. Suppose your family operated an electric car that used lead–acid batteries. In what ways would this affect your family’s lifestyle?
Worked solutions to student book questions Chapter 27 Cells and batteries
AE8. Owning an electric car, rather than a petrol-driven car, might restrict your range and frequency of travel. You would save money by not having to buy petrol, and maintenance costs (e.g. oil changes, tune-ups) could also decrease, but cost would be involved in replacing your battery every few years.
Q6. The pH of the electrolyte in a lead–acid battery varies according to the state of charge. Describe how the pH would change as a battery discharged and was recharged.
A6. As a lead–acid battery is discharged, H+ ions are consumed. As the concentration of H+ drops, the pH of the electrolyte will increase. The opposite occurs when a lead–acid battery is recharged. As the concentration of H+ increases, the pH will decrease.
Q7. A lead–acid battery uses 50 g of lead to provide the electrical energy needed to start a typical car. Calculate the mass of lead oxide consumed at the positive electrode and the total mass of lead sulfate produced in the battery while starting the car.
2–(aq) + 4H+(aq) → 2PbSO4(s) + 2H2O(l) Step 2 From the equation, 1 mol of PbO2 reacts with 1 mol Pb. Use stoichiometry to
calculate the amount of PbO2.
(Pb)
)PbO( 2
nn =
11
n(PbO2) = n(Pb)
= 1mol g 207.19g 50
−
= 0.24 mol Step 3 Calculate the mass of PbO2. m(PbO2) = 0.24 mol × 239.19 g mol–1 = 57.51 g = 58 g Step 4 From the equation, 2 mol PbSO4 is produced by 1 mol Pb. Use stoichiometry
to calculate the amount of PbSO4.
(Pb)
)PbSO( 4
nn =
12
n(PbSO4) = 2 × n(Pb)
= 2 × 1mol g 207.19g 50
−
= 0.48 mol
Worked solutions to student book questions Chapter 27 Cells and batteries
Step 5 Calculate the mass of PbSO4. m(PbSO4) = 0.48 mol × 303.19 g mol–1 = 145.5 g = 1.5 × 102 g
Q8. A mechanic wishes to recharge a car battery using a battery charger. a Which battery terminals should the positive and negative terminals of the battery
charger be connected to? b What would be the effect on the battery if the charger were connected
incorrectly?
A8. a The positive terminal of the charger is connected to the battery’s positive terminal
and the negative terminal of the charger is connected to the battery’s negative terminal.
b Connecting the battery charger to the battery so that terminals of opposite polarity are linked will force the spontaneous cell reaction to occur at a faster rate, further flattening the battery and possibly damaging it irreparably.
Q9. Why is large-scale production of electricity using fuel cells suggested as part of a solution to the greenhouse problem?
A9. Fuel cells are about twice as efficient as coal-fired power stations. Consequently, fuel cells produce the same quantity of energy using about half as much fuel. Less fuel means less carbon dioxide gas is produced. As carbon dioxide is a major greenhouse gas, the use of fuel cells has the potential to reduce the greenhouse effect.
Q10. The use of hydrogen-powered fuel cells in public transport is being considered in many cities around the world. What are the advantages and disadvantages of using fuel cells for this purpose?
A10. Hydrogen powered fuel cells generate water as the only product and do not generate noise when in operation. Wide spread use of fuel cells requires extensive new infrastructure such a generation plant, storage facilities, specialised transport and distribution outlets. E1. Write equations for the reactions that occur at the anode and cathode of a NaS cell as it discharges. AE1. Anode: Na(l) → Na+(l) + e– Cathode: S(l) + 2e– → S2–(l)
Worked solutions to student book questions Chapter 27 Cells and batteries
E2. Write and overall equation for the cell reaction during: a discharge b recharge. AE2. a 2Na(l) + S(l) → 2Na+(l) + S2–(l) b 2Na+(l) + S2–(l) → 2Na(l) + S(l) E3. To operate efficiently, the cells require a temperature above 270°C. Why do you think this is necessary? AE3. To operate efficiently, the cells require that both the sodium and sulfur be in the liquid state. As both are solid at room temperature, a temperature above 270°C is necessary. E4. Provided the cells are discharged and recharged every day, they remain at working temperatures. Suggest the reason for this. AE4. The recharging of most cells and batteries is less than 100% efficient and some heat energy is usually created. Similarly, the discharge of most cells and batteries also releases a small amount of heat. In NaS batteries, this heat is sufficient to maintain a working temperature provided the batteries are discharged and recharged every day. E5. Suggest a meaning for the term ‘energy density’. Why is energy density an important factor in the construction of most types of cells and batteries? AE5. The term ‘energy density’ is the ratio of the energy available from a cell to its mass. The greater the energy density, the more energy is available for the same mass of cells or the less mass of cells is required for the same amount of energy. The higher the energy density, the smaller and lighter the appliance that uses them can be. This is a crucial factor for electric cars, mobile phones and hearing aids. E6. Explain how the use of NaS batteries enables an intermittent source such as wind power to be used for baseload power. AE6. The use of NaS batteries enables an intermittent source such as wind power to be used for baseload power because electrical energy from the wind turbines can be collected and stored for later release. Thus, when there is little or no wind, the power station can continue to provide energy by using the NaS store. This enables the power station to provide a constant supply of electricity regardless of wind conditions.
Worked solutions to student book questions Chapter 27 Cells and batteries
Q11. When constructing a galvanic cell in the laboratory, two separate containers are generally used for the half cells. However, in commercial cells such as the dry cell, all the reactants are placed in the one container. a How are commercial cells constructed so that the use of two containers is
avoided? b Design a cell that could be constructed in one container in the laboratory and
which uses the same cell reaction as in the Daniell cell (p. 415).
A11. a Most commercial cells contain an oxidant and reductant in the form of a solid or
moist powder; each is unable to migrate and react directly with the other. The electrolyte is usually a paste, allowing ions to move.
b
Q12. Explain the difference between the terms: a positive electrode and negative electrode b cell and battery c primary cell and secondary cell d discharge and recharge
A12.
a The negative electrode anode in a galvanic cell or is the electrode at which oxidation occurs. The positive electrode or cathode is the electrode at which reduction occurs.
b A battery consists of a number of cells connected together in series. c A primary cell cannot be recharged and is discarded once it no longer supplies
electrical energy. A secondary cell can be connected to a battery charger to reverse the cell reaction. Once the cell has been recharged, it can again be used to supply electrical energy.
d When a cell discharges chemical energy is converted to electrical energy. When it is recharged, electrical energy is converted to chemical energy.
Worked solutions to student book questions Chapter 27 Cells and batteries
Q13. What feature enables secondary cells such as lithium ion cells to be recharged?
A13. In secondary or rechargeable cells such as lithium ion cells or a lead acid battery, the products of cell reaction remain in contact with the electrodes and are in forms that can be converted back to reactants.
Q14. Explain what happens to the oxidant, reductant, anode and cathode when a secondary cell is switched from discharge to recharge.
A14. When a cell discharges, oxidation occurs at the anode which has a negative polarity. Reduction occurs at the cathode which has positive polarity. The oxidation and reduction reactions in the cell generate electricity. When the cell is recharged the direction of the electron flow is reversed. An external supply of electricity is required to reverse these reactions. Oxidation now occurs at the positive electrode while reduction occurs at the negative electrode.
Q15. Both dry and alkaline cells use powdered graphite mixed with a manganese dioxide oxidant. From your knowledge of the properties of graphite, suggest what its role might be in these cells.
A15. Graphite is a good conductor of electricity because it contains delocalised electrons between the layers of carbon atoms. Graphite is incorporated in the cells to conduct electrons to the manganese dioxide oxidant.
Q16.
Describe an application where each of the following would be used: a a nickel–metal hydride cell b a lead–acid accumulator c a lithium button cell d a dry cell e an alkaline cell
A16.
a an appliance that draws a moderate current, e.g. camera flash, portable cassette player, remote controlled car
b starting internal combustion engines in motor bikes, cars, trucks, and buses c an appliance with a small current drain that requires a constant voltage, e.g.
watch, camera, hearing aid, calculator d an infrequently used device that requires a cheap power source, e.g. glove box
torch, transistor radio e an appliance with a moderate to high current drain that is not used extensively,
e.g. electric toothbrush, portable cassette player
Worked solutions to student book questions Chapter 27 Cells and batteries
Q17. A manual car with a flat battery can often be started by pushing, but this is not done with an automatic car. Find out why a manual car can be started in this way but an automatic cannot.
A17. The role of the battery is to turn the starter motor which, in turn, cranks the engine. Pushing a manual car that has its gears engaged will cause the motor to turn and, subsequently, cause the alternator to produce an electric current to fire the spark plugs and start the engine. An automatic car has no direct coupling between the engine and transmission; a fluid coupling system is used to drive the transmission when the engine is running.
Q18. Mains electricity costs about 5 cents per MJ of energy, or less, depending on the tariff. The cost of the same amount of electrical energy produced by a cell is far more—about $1300 for a dry cell and even more for a button cell. Why are people prepared to use cells and pay such relatively high prices for electricity?
A18. While the cost of electrical energy purchased in the form of a dry cell or button cell is far higher than the cost of mains electricity, people are prepared to pay the higher price for the convenience and flexibility of the portable equipment powered by these cells. Furthermore, the price of individual cells is regarded as relatively low.
Q19. A car has been powered using an experimental ten-cell aluminium–air battery. Each of the ten cells contains aluminium plates and sodium hydroxide solution, and air is fed into the battery. New aluminium plates and water must be added regularly to the battery. a Write half equations for the anode and cathode reactions in the battery. b Draw a diagram to show how each cell might be constructed. c Give two reasons why this battery weighs less than a similar lead–acid battery. d Why can this battery be described as a fuel cell?
A19.
a anode: Al(s) → Al3+(aq) + 3e–; cathode: O2(g) + 2H2O(l) + 4e– → 4OH–(aq) b
Worked solutions to student book questions Chapter 27 Cells and batteries
c One reason why an aluminium–air battery weighs less than a similar lead–acid battery is that the density of aluminium is less than lead. Secondly, the reaction at the cathode of an aluminium–air battery uses oxygen obtained directly from the air as it is required, whereas a lead–acid cell must incorporate lead(IV) oxide around the cathode.
d A fuel cell is an electrochemical cell in which the oxidant and reductant are supplied continuously so that the cell does not flatten. By replacing the aluminium and adding water and oxygen from the air, an aluminium–air cell will function as a fuel cell.
Q20. Preliminary roadside breath tests are often performed by measuring the voltage produced when a motorist breathes into a small fuel cell. The cell operates like a hydrogen–oxygen fuel cell, except that ethanol replaces hydrogen. Carbon dioxide is one of the products formed when the ethanol is oxidised at the anode. a Predict the anode, cathode and overall reactions in the cell. (Assume the cell
electrolyte is acidic.) b Would you expect the fuel cell to generate a voltage if natural gas were blown
b Yes. The net cell reaction in the fuel cell is the oxidation of ethanol to form carbon dioxide and water. Like ethanol, methane in natural gas can be oxidised to carbon dioxide and water. Provided the electrode materials used in the cell act as effective catalysts for both reactions, it is likely that the use of natural gas would produce a voltage.
Q21. An experimental fuel cell that uses methanol as the fuel has the half equations:
a Write the equation for the overall cell reaction. b Which reaction occurs at the positive electrode of the cell? c Suggest a suitable electrolyte for the cell. d When the cell begins to produce electricity, the pH of the electrolyte near the
cathode increases and eventually reaches a constant value. Explain why this occurs.
e Electricity could be obtained from thermal energy produced by combustion of methanol. What is the main advantage of using a fuel cell to produce electricity?
Worked solutions to student book questions Chapter 27 Cells and batteries
a 2CH3OH(g) + 3O2(g) → 2CO2(g) + 4H2O(l) b reduction of O2(g) c NaOH(aq) or KOH(aq) d Oxygen is reduced, forming OH– ions at the cathode of the fuel cell, so an
increase in pH occurs at this electrode initially. The OH– ions migrate to the anode where they are consumed. Once the rate of production of OH– ions at the cathode becomes equal to the rate at which they depart, the pH near the electrode will be constant.
e A fuel cell converts chemical energy into electrical energy directly, with relatively little energy being converted into thermal energy. If electrical energy were obtained by burning methanol (in a process similar to that used to obtain electrical energy from coal in coal-fired power stations), the energy ‘losses’ of the various energy transformations involved would be greater. In particular, large losses occur when thermal energy is converted into mechanical energy.
Q22. Many of the ‘alkaline cells’ on the market contain zinc electrodes in contact with an electrolyte containing hydroxide ions. The half-cell reaction might be represented as:
Zn(s) + 4OH–(aq) → Zn(OH)42–(aq) + 2e–
To investigate if the standard electrode potential (E°) of these half cells is the same as that of a zinc electrode in contact with a zinc nitrate electrolyte, for which the electrode reaction would be
Zn(s) → Zn2+(aq) + 2e– a student is provided with the two half cells shown in Figure 27.20, as well as a Cu2+/Cu half cell.
Figure 27.20 The three half cells.
a Write a net equation for the reaction in a galvanic cell in which the ‘alkaline zinc half cell’ is connected to the Cu2+/Cu half cell.
b Carefully explain how the student could use the half cells that were provided to determine whether the two different half cells containing zinc had the same E° value. Include fully labelled diagrams with your answer and explain how the results could be interpreted.
Worked solutions to student book questions Chapter 27 Cells and batteries
a Zn(s) + 4OH–(aq) + Cu2+(aq) → Zn(OH)42–(aq) + 2Cu(s)
b The student should make two electrochemical cells consisting of I the Cu2+/Cu half cell and the ‘alkaline zinc half cell’ II the Cu2+/Cu half cell and the Zn2+/Zn half cell as shown in the diagrams. The cell voltages should be measured and if they are identical then the two half
cells have the same E0 values (they would not be expected to have the same E0 values).
Worked solutions to student book questions Chapter 27 Cells and batteries
Q23. In the quest for power plants with lower emissions of carbon dioxide, a Victorian company is developing solid oxide fuel cells (SOFC). Also known as ceramic fuel cells, the units have a ceramic electrolyte through which oxide ions, O2–, can move. On the top and bottom of the electrolyte is a thin layer of electrode material. Fuel cells are designed so that the fuel and air are directed separately onto the appropriate electrode. The units are arranged in stacks so that high voltages can be generated. The diagram represents a fuel cell that is using methane gas, CH4, as the energy source. The direction of electron drift through the external circuit is shown in Figure 27.21. The electrode reactions are:
O2 + 4e– → 2O2– CH4 + 4O2– → CO2 + 2H2O + 8e–
Figure 27.21 Methane fuel cell.
a Write an overall equation for the reaction that occurs in this SOFC. b On the diagram, clearly label the positive and negative electrodes. c Explain the claim that this method of generating electricity will reduce carbon
dioxide emissions, by comparison with a power plant that burns natural gas (containing methane) and uses the heat released to boil water and then generate electricity.
d Why is it necessary to aim for reduced emissions of carbon dioxide from power-generating facilities?
Worked solutions to student book questions Chapter 27 Cells and batteries
Q24. The current generation of hybrid cars such as the Honda Insight and the Toyota Prius use nickel–metal hydride batteries. Nickel–metal hydride technology has replaced lead–acid batteries previously used in many developmental hybrid and electric cars. Yet the ‘holy grail’ for both hybrid and electric vehicles remains lithium ion technology. a What aspects of lead–acid battery technology have resulted in their replacement
in hybrid cars with nickel–metal hydride batteries? b What features of lead–acid batteries make them an attractive option for future
hybrid and electric vehicles? c What is currently the main impediment to the use of lithium ion batteries in
hybrid and electric vehicles?
A24. a Lead–acid batteries are heavy and have a lower energy density, which limited
fuel economy gains of the early model hybrid cars b Lithium ion batteries are attractive for future hybrid and electric vehicles because
they have an even higher energy density than nickel–metal hydride batteries. c The main impediment is its limited lifespan in comparison to nickel–metal
hydride batteries. At around three years, lithium ion batteries have about half the lifespan of the nickel–metal hydride batteries used in current model hybrid cars.
A23.
a CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) b
c Since fuel cells are more efficient at converting chemical energy to electrical energy, there will be less CO2 emitted to the atmosphere per joule of energy generated.
d Increased levels of CO2 in the atmosphere are thought to be causing global warming.
