Natural Sciences Tripos Part II

MATERIALS SCIENCE

II

MATERIALS SCIENCE

C2 : Electrochemistry

Dr R. V. Kumar

Michaelmas Term 2014 15

Name............................. College..........................

Michaelmas Term 2014-15

1

C2 - ELECTROCHEMISTRY

Dr R.V. KUMAR

COURSE CONTENT (6 Lectures + 1 Examples Class)

Electrolytes

Aqueous solutions; Molten salts; Ionic liquids; Solid electrolytes Electrochemical Reactions

Electrode potentials & Nernst equation; Reference electrodes; Concentration cells; Redox potentials; Kinetics of electrode reactions

Applications

Chemical sensors; Batteries; Fuel Cells; Electrolysis; Photocatalytic redox reactions

Examples Class (one) / Question Sheet (one) Resources: Chemical Metallurgy, JJ Moore, Butterworth-Heinmann (PG14) Modern Batteries, C Vincent and B Scrosati, Butterworth-Heinmann (Pn118) Fuel Cell Systems Explained, J Larmire and A Dick, John-Wiley & Sons Ltd (Pn123) DoITPoMS Teaching & Learning Packages http://www.msm.cam.ac.uk/doitpoms/tlplib/index.php The Nernst equation & Pourbaix diagrams Ellingham diagrams Kinetics of aqueous corrosion Batteries Fuel cells

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Introduction

A basic electrochemical cell comprises two electrodes and one electrolyte. Reactions in electrochemical cells occur when electrons are transferred from one species to another and are known as redox reactions. Any given redox reaction can be expressed such as to depict actual gain or loss of electrons. If electrons are donated, the reaction is an oxidation, while if electrons are accepted, the reaction is a reduction.

+ + eaqZnsZn 2)()( 2 Oxidation )(2)(2 sCueaqCu + + Reduction

The processes of oxidation and reduction are spatially separated by the electrolyte with reduction occurring at one electrode, the cathode, and oxidation at the other, the anode. In the electrochemical cell, both oxidation and reduction reactions take place at the two electrodes (as 2 half-reactions) so as not to violate the laws of conservation of matter and energy. Each electrode is in contact with the electrolyte and at the interface there is an interchange of ionic to electronic current (or vice versa). Electrolyte is by definition an ionic conductor, while electrodes are electronic conductors. Sketch of an electrochemical cell

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Galvanic Cell The electrodes are linked externally through electrically conducting wires, high impedance meters, electrical shunts, electrical loads, devices, batteries and power supplies or systems. When electrode reactions take place by the virtue of a thermodynamic driving force, the cell is termed a galvanic cell. Cathode is the positive pole and anode is the negative pole of a galvanic cell. In galvanic cells, chemical energy embodied in the reactants is converted to electrical energy. Technological examples of galvanic cells are sensors, fuel cells, batteries and supercapacitors. Corrosion of materials is also a result of galvanic action.

4

Electrolytic Cell Electrical energy can be supplied using an external power source in order to drive the electrode reactions in a direction opposing the driving force. In this cell the anode is made the positive electrode via which current enters the electrolyte and the cathode is made the negative electrode from where the current exits out of the cell. Electrolytic cells are used in charging batteries and supercapacitors and in electroplating, metal production and electrophoresis.

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Electrolytes Aqueous electrolytes: These are based on salts, acids, bases dissolved in water and they conduct electricity via a mixture of ions such as H+, OH-, Mz+, Xz- . All ions present in the aqueous solution contribute to the current but not necessarily by the same amount. Aqueous electrolytes are extensively used in fuel cells, batteries, sensors, supercapacitors, electroplating and electrolysis. Molten Salts: When salts undergo melting they are capable of conducting ions and carry current. Large scale application of NaF-AlF3 molten salt saturated with alumina at 1000oC is used industrially in electrolytic cells to produce over 20 million tonnes per year of metallic aluminium. Other halides based on NaCl, KCl, LiCl and MgCl2, CaCl2 are also used in order to produce the very reactive alkali and alkaline earth metals. Fuel cells using molten salts of carbonates have achieved a high state of development for commercial applications. All ions present in the molten salt system contribute to the current but not necessarily by the same amount. Ionic Liquids: These are special class of electrolytes based on organic salts (typically made up of inorganic anions and large organic cations) that are liquid at or near room temperature and thus can effectively offer the advantages of molten salts at low temperatures without the need for water as the solvent as in aqueous electrolytes. Electrochemical applications in solar cells, metal production, and synthesis of nano-structures and in batteries and supercapacitors are being researched extensively. Typically, ionic liquids have low vapour pressure and low flammability in comparison with many organic solvents, and are intrinsically conductive due to their ionic nature. Ionic liquids can be constituted by combining a cation which can be metal cation (e.g. Zn2+) or organic (e.g. Tetraalkyl ammonium) with an anion which can be inorganic (e.g. halides) or organic (e.g. methylsulfates)

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Solid Electrolytes: Ion conductors in the solid state exist as an intermediate between conventional ionic solids and ionic liquids. They have a rigid immobile structure formed in one sub-lattice through which a component can move relatively freely within another sub-lattice. Many of the ionic conductors have the potential to be employed in electrochemical cells, ranging from sensors, fuel cells, supercapacitors and electrolysers to batteries, electrochemical capacitors and electrochromics, as well as devices with ion-sensitive electrodes including pH meters.

Oxygen ion conductors represent the most important class of solid state ionically conducting materials. Of these, the most widely investigated, and used, solid electrolytes are based upon ZrO2. The interest in the zirconias, both from the scientific and the technological viewpoint, has been truly overwhelming over the last five decades, and there are now countless practical applications relying on this material. ZrO2 when doped with aliovalent oxides, such as Y2O3, CaO and MgO, acquires ionic conduction for oxygen ions over a wide range of temperature and partial pressure of oxygen. Pure zirconia undergoes two structural transformations on heating (or the reverse on cooling): cubicl tetragona monoclinic

CC 23401170 ,

Melting takes place at approximately 2680 C. (See Phase Diagram in the next page.) During transformation from the tetragonal to the monoclinic phase there is a large specific volume change of 5 %, resulting in cracking of a pure zirconia monolith rendering the material useless for most applications. The high-temperature cubic fluorite phase of zirconia can be partially or fully stabilised to lower temperatures by the substitution of a portion of the Zr4+ cations with lower valent cations like Mg2+, Ca2+, Y3+, Sc3+ or trivalent rare earth (RE) metal cations. This substitution is possible because ZrO2 can form solid solutions with the respective divalent or trivalent oxides. These aliovalent dopants stabilise the fluorite structure and thereby avoid the destructive tetragonal-to-monoclinic transformation. In addition, the dopants introduce vacancies in the oxygen ion sites due to their lower valences. When yttria is used to stabilize zirconia, it is called yttria-stabilized zirconia (YSZ).

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Phase Diagram

The introduction of yttria into zirconia and the formation of a solid solution can be described as: ++ OO/Zr VOYOY 3232 Oxygen ion conduction in stabilised zirconia occurs via these vacancies, because oxygen ions, xOO , are able to exchange sites with adjacent oxygen vacancies,

..OV . Thus the

oxygen ion motion can be considered to be equal and opposite to the vacancy motion. As the dopant concentration is increased, the number of oxygen vacancies will increase according to the stoichiometric relationship: [ ] [ ]/ZrO YV =2 As can be seen from the figure below, the conductivity in stabilised zirconia first increases with increasing dopant content, before reaching a maximum, beyond which the conductivity then decreases. The reason for this maximum is that above a critical vacancy concentration defect clusters are prevalent, decreasing the mobility of the oxygen vacancies and thus decreasing the conductivity. The composition that displays the conductivity maximum in

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stabilised zirconia corresponds to the minimum amount of dopant required to fully stabilise the cubic phase. At a temperature of 1000 C, which represents a typical fuel cell operating temperature, the conductivity maximum occurs near 8 mol% Y2O3.

Conductivity vs mol % dopant

The tetragonal phase within a partially stabilised zirconia (PSZ) can transform to the monoclinic phase under mechanical stress. Energy absorbed through the corresponding increase in volume of 3 to 5 % arrests crack growth resulting in toughening of the ceramic. Thus, a partially stabilised zirconia containing the cubic and the tetragonal phases at typically less than 8 mol% yttria combines high ionic conductivity with high toughness and thermal shock resistance. A fully tetragonal phase is formed in the 2 to 3 mol% Y2O3 system in the temperature range from 1200 to 1500 C, commonly known as tetragonal zirconia polycrystals (TZP). The transition to monoclinic can be completely suppressed at temperatures below 500 C, provided the grain size is less than 300 nm, which results in a strong and tough material. At less than 50 nm the material is extremely tough while retaining hardness and strength. Within the solid solution range, the vacancy concentration in stabilised zirconia is determined by the dopant concentration. At higher temperatures and low partial pressure of oxygen, electrons may appear as a result of the non-stoichiometric reactions between lattice and gaseous oxygen:

/OO e(g)OVO 221

2 ++

9

with equilibrium constant K expressed as: 212

2OO/ ]p[V][eK =

where 2Op is the partial pressure of oxygen. At high partial pressure of oxygen, electron

holes may appear:

++ hO(g)OV OO 221

2

with equilibrium constant given by:

212

2OO]p[V

][hK

=

The conductivity is given by the equation: = charge concentration of charge carrier mobility Since electronic mobility is usually much higher than ionic mobility, typically by nearly about three orders of magnitude, the ionic conductivity is dominant only when the ionic defect concentration is very large relative to the electronic defects. The ionic transference number is defined as the ratio of ionic conductivity to total conductivity. In the case of an oxygen ion conductor, such as YSZ, the oxygen ion transference number 2Ot is defined as:

++= heO

OO

t

/2

22

For a given temperature there is a range of oxygen partial pressure values within which ionic defects dominate. This renders oxygen ions the majority charge carrier and oxygen ion conductivity independent of oxygen pressure. constO =2 The ionic (or electrolytic) domain is defined by the range of oxygen partial pressures for which 2Ot 0.99. This is represented in Figure shown below which has been calculated

on the basis of numerical values typical for stabilised zirconia. Oxygen ion, electron and hole conduction are all activated processes with typical activation energies in 8 mol% YSZ of 0.85, 3.88 and 1.67 eV, respectively. Due to their higher activation energies, the electron and hole conductivities increase more than the ionic conductivity as temperature rises. This causes the width of the electrolytic domain to diminish, as is shown schematically. Knowledge of the electrolytic domain is important in order to be able to use the oxygen ion conductor as a solid electrolyte in sensors and fuel cells.

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Ionic DomainLog pO2(atm.)

1/T, (per K)

Electronicconduction

Hole conductiontion = 0.99

tion = 0.99

T, K

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Separator: In addition to the electrodes, the two other constituents that are required for electrochemical reactions to be usefully harnessed are the electrolyte and the separator. The electrolyte is an ion conducting material, which can be in the form of an aqueous system, an ionic liquid, an organic solvent containing a soluble salt, a molten salt, or a solid electrolyte, while the separator is a membrane that physically prevents direct contact between the two electrodes and allows ions to pass through; it therefore ensures electrical insulation for charge neutralization in both the anode and cathode once the reaction is completed. A separator is soaked in the electrolyte when liquid electrolyte is used.

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Electrochemical Reactions

The general expression for a reduction reaction is Ox + Z e- = Red

Both Ox and Red may be positive, neutral or negative species. The standard reduction potential Eo is a measure of how easily a species can gain (or loose) electrons. Table Standard Electrode Potentials in Aqueous Electrolyte at 298K (written as reduction reactions by convention)

Reaction E0 / V Li+ + e Li -3.10 Na+ + e Na -2.71 Mg2+ + 2e Mg -2.36 Mn2+ + 2e Mn -1.18 MnO2 + 2H2O + 4e Mn + 4OH -0.98 2H2O + 2e H2 + 2OH -0.83 Cd(OH)2 + 2e Cd + 2OH -0.82 Zn2+ + 2e Zn -0.76 Ni(OH)2 + 2e Ni + 2OH -0.72 Fe2+ + 2e Fe -0.44 Cd2+ + 2e Cd -0.40 PbSO4 + 2e Pb + SO42 -0.35 Ni2+ + 2e Ni -0.26 MnO2 + 2H2O + 4e Mn(OH)2 + 2OH -0.05 2H+ + 2e H2 0.00 Ag2O + H2O + 2e 2Ag + 2OH +0.34 Cu2+ + 2e Cu +0.34 O2 + 2H2O + 4e 4OH +0.40 2NiOOH + 2H2O + 2e 2Ni(OH)2 + 2OH +0.48 NiO2 + 2H2O + 2e Ni(OH)2 + 2OH +0.49 MnO42 + 2H2O + 2e MnO2 + 4OH +0.62 Ag+ + e Ag +0.80 O2 + 4H+ + 4e 2H2O +1.23 PbO2 + 4H+ + 2e Pb2+ + 2H2O +1.47 PbO2 + SO42 + 4H+ + 2e PbSO4 + 2H2O +1.70 F2 + 2e 2F +2.87

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The Standard Hydrogen Electrode (SHE) H2 of pressure 1 atm is bubbled into an HCl (aq) solution at a concentration of 1M onto a high surface area Pt. By definition, Eo(SHE) = 0 V The very electropositive metals, such as Li, Mg, Mn will directly react with water decomposing it and produce hydrogen and thus thermodynamically they are unstable as electrodes in aqueous electrolytes. The standard potential of a half-reaction is reported with respect to the SHE in an aqueous solution, whereby the species are in their standard states defined by 1M for soluble species Saturation for slightly soluble species 1 atm. Pressure (101325 Pa) for any gas Pure substance in stable form for a metal Pure substance in stable form in contact with the metal for other solids The overall cell reaction (redox reaction) can be expressed by combining 2 half reactions with the more +ve potential undergoing reduction; the one with the less +ve potential will be reversed and proceed as oxidation. Nernst Equation When the species are not in their standard state, we can calculate the electrode potential by using the Nernst Equation. The Nernst Equation for the reaction:

Ox + Z e- = Red is given by:

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][][Relog303.2

Oxd

zFRTEE o =

Reference Electrode Since the standard hydrogen electrode may not be a convenient reference electrode in practical usage, other reference electrode systems have been developed for practical use in aqueous systems. Ag/AgCl electrode: Cl-(of defined concentration) | AgCl(s) | Ag(s) The Cl- concentration is fixed using a saturated aqueous solution of KCl The electrode reaction for this reference electrode is: AgCl + e- = Ag + Cl- Eo (298K) = +0.23 V Using the Nernst equation: ]log[3.3.2 = Cl

FRTEE o

The electrode potential with a solution of KCl (usually saturated) is given by E = +0.197 V at 298K. Saturated calomel electrode (SCE)

)(|)(|)( 22 lHgsClHgdefinedCl

The Cl- concentration is fixed using a saturated aqueous solution of KCl The electrode potential a solution of KCl (usually saturated) is given by E = +0.244 V at 298K.

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Concentration Cells: In an electrochemical cell, the cell potential can arise simply by using 2 different concentration of a selected substance at the 2 electrodes. If the electrodes are connected in the external circuit a current would flow through the cell system in order to try and eliminate the concentration difference. [In reality, it is activity difference rather than concentration difference that will drive the net electrochemical reaction in a concentration cell] Concentration cells are very commonly employed in practice. Electrorefining of impure Cu to pure Cu in an aqueous electrolyte is a concentration cell. Galvanic cells using a YSZ solid electrolyte separated by different oxygen partial pressures at the anode and the cathode is a concentration cell as shown below.

Ni, NiO | YSZ | Fe, FeO

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Electrode Kinetics Interface When a metal electrode is in an electrolyte, the charge on the metal will attract ions of opposite charge in the electrolyte, and the dipoles in the solvent will align. This forms a layer of charge in both the metal and the electrolyte, called the electrical double layer, as shown in Fig below. The electrochemical reactions take place in this layer, and all atoms or ions that are reduced or oxidized must pass through this layer. Thus, the ability of ions to pass through this layer controls the kinetics, and is therefore the limiting factor in controlling the electrode reaction. The energy barrier towards the electrode reaction, described as the activation energy of the electrochemical reaction lies across this double layer.

Electrode Kinetics When an electrode is not at equilibrium an overpotential exists, given by = E Eo, where is the overpotential, E is the actual potential, and E0 is the equilibrium potential. Overpotential is used synonymously with polarization. Electrode kinetics is described by Tafel equation at each of the electrode: Anode:

oAA EE =

AAAA iba ln+=

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or in terms of exchange current density oi (which reflects the rate at which equilibrium is achieved for a given electrode/electrolyte system):

)/ln( oAAA iib= and similarly for the cathode, substitute A with C. Space for a typical Tafel plot: Thus for an applied potential, the current density, i, can be found from the Tafel plot in an electrolytic cell when the battery is being charged or vice versa in a galvanic cell as in a battery being discharged. At very high currents a limiting current may be reached as a result of concentration overpotential, C, restricting mass transfer rates to the diffusion rate of the electroactive species. A limiting current arises which can be derived from Ficks first law of diffusion under the condition that the electrode surface is depleted of the ion and the recovery of the ion concentration is limited by ion-transport through the electrolyte diffusion boundary layer. The limiting current is diffusion limited, and can be determined by Ficks law of diffusion as, where iL is the limiting current density over a boundary layer:

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zFDCiL =

where iL is the limiting current density over a boundary layer, D is the diffusion coefficient of metal cations in the electrolyte (or chemical component in the electrode in some other situation), C is the concentration of metal cations in the bulk electrolyte, and is the thickness of the boundary layer. Typical values for deposition for Cu from Cu2+ for example would be: D = 2x 10-9 m2s-1, C = 0.05x 104 kg m-3, = 6x 10-4 m, which gives iL = 3.2 x 102 A m-2 The concentration overpotential thus represents the difference between the cell potential at the electrolyte concentration and the cell potential at the surface concentration because of depletion (or accumulation) at high current densities, given by

)1ln(303.2.)(L

C ii

zFRTconc =

Space for a Tafel curve showing this diffusion limiting of the current:

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Electrochemical Sensors

Electrochemical principles form the basis of applying solid state ionic conductors in chemical sensors for monitoring and/or controlling chemical components in gaseous, liquid and solid systems. Solid electrolytes interfaced with suitable electrodes can directly and quantitatively transduce the chemical activity of the species to be sensed into a readily measurable electrical quantity. Electrochemical sensors are used either in the potentiometric or in the amperometric mode. In potentiometric sensors, the open circuit potential difference between the anode and the cathode, also termed electro-motive force, emf, is measured at zero current with a high impedance voltmeter. One of the electrodes contains the species to be sensed, the other is the reference electrode, and the emf can be related directly to the unknown activity of the target component. In amperometric sensors, the current is measured under an applied voltage between anode and cathode. This current can be related to the activity of the target component. Oxygen Sensor

The commercialized oxygen sensor is composed of a partially stabilised zirconia thimble as the oxygen ion conducting solid electrolyte and porous coatings of platinum - rhodium on both sides as the electrodes. As described earlier, 3 mol% YSZ, corresponding to a mixture of tetragonal and cubic phases, provides optimum mechanical and electrical properties. The electrodes are applied to the solid electrolyte by thermal evaporation, sputtering or chemical deposition, followed by thermal treatment. During sensor operation both electrodes are hermetically sealed from each other. One electrode is exposed to the test gas, the other electrode is exposed to air serving as the reference. A heater brings the sensor to its operating temperature, thus called Heated Exhaust Gas Oxygen Sensor (HEGO). A porous layer of thickness 50 to 300 m around the sensing electrode protects the sensor from mechanical damage. The layer also acts as a diffusion barrier for the gas and ensures thermal equilibration of the gas before it reaches the electrode. The layer is made from MgAl2O4, Al2O3, ZrO2 or a combination thereof, and is produced with organic pore forming substances. A schematic diagram is shown below: The oxygen sensor (also referred to as a probe) operates as a galvanic oxygen concentration cell: (+) Pt, O2 (ref) | YSZ | O2 (test), Pt (-). The test chamber contains the sample gas of oxygen partial pressure testOp 2 , and the

reference chamber contains air, rendering the reference oxygen partial pressure refOp 2

equal to 0.21 atm. The emf generated by the difference in oxygen partial pressure is given by the Nernst equation:

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refO

testO

p

p

FRT.emf

2

2log4

3032=

where R is the universal gas constant, T is the absolute temperature, and F is the Faraday constant. The high catalytic activity of the platinum and rhodium helps establish thermodynamic equilibrium at both electrodes. The oxygen partial pressure difference between the exhaust gas and the air causes opposite reactions at both electrodes. Under operating conditions, the working electrode is the anode and the electrode reaction is:

( ) ( ) += etestOYSZO 42 22

The reference electrode is the cathode and the electrode reaction is:

( ) ( )YSZOereferenceO =+ 22 24 The stoichiometric point of an automotive internal combustion engine is at an air-to-fuel mass ratio of ~14.7, ensuring complete reaction of both the fuel and the oxygen. The air/fuel ratio, A/F, is commonly expressed in terms of the normalised quantity lambda, :

tric A/FstoichiomeA/Foperating

=

21

The stoichiometric point is at = 1, and here the equilibrium partial pressure of oxygen in the exhaust gas varies between 10-20 and 10-2 atm, i.e., from slightly fuel-rich to slightly lean-burn. This sharp change in oxygen pressure produces a large change in the emf. Figure (below) shows the resulting emf curves relative to the normalised air/fuel ratio for a fixed temperature. This response characteristics makes the sensor ideal for closed-loop control of the automotive internal combustion engine centred at = 1.

0.90 0.95 1.00 1.05 1.100

0.2

0.4

0.6

0.8

1.0

10-6

10-1

10-11

10-16

10-21

Calculatedsensor voltage

pO2

pO2

(bar

)

V N

erns

t (V

)

Normalized air fuel ratio Lowering the concentrations of pollutant gases CO, hydrocarbons (HCs) and NOx from automobile emissions is achieved by the use of three-way catalysts (TWC) placed upstream in the exhaust gas. When operated at stoichiometric air/fuel ratio, the TWC can simultaneously oxidise CO and HCs to CO2 and H2O and reduce NOx to N2, and in this way convert pollutant gases into innocuous gases. The concentration of gases as a function of air/fuel ratio is shown below. The catalyst is deposited on a CeO2 layer with oxygen storage capacity such that oxygen can be stored when the exhaust gas is slightly lean and released when it is slightly rich. Keeping inside a narrow window around the stoichiometric point is therefore critically important in the engine management.

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CO2

CO

H2

O2

NO

Hydrocarbons

Stoichiometricmixture =1

6

5

4

3

2

1

12

10

8

6

4

2

10:1 12:1 14:1 16:1 18:1Air to fuel ratio (wt)

CO

, CO

2, O

2 or H

2 (v o

l %)

NO

, hy d

roca

rbo n

s (x

1000

pp m

)

1

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Batteries Introduction There are two main types of batteries: primary and secondary batteries. In primary batteries the chemical energy stored in the cell is such that it can be used only once to generate electricity, i.e. once the cell is fully discharged it cannot be of further use. In secondary batteries the reverse redox reaction (also referred to as electrolysis and charging) can occur when the current is applied at a potential higher than the cell potential (Ecell) and the battery can be used reversibly many times. During charging, electrons flow to the anode through the external circuit and cations from the cathode diffuse through the electrolyte. Primary Batteries Primary batteries are not easily or safely rechargeable, and consequently are discharged and then disposed of. Many of these are dry cells cells in which the electrolyte is not a liquid but a paste or similar. The cell electrochemical reactions are not easily reversible and the cell is operated until the active component in one or both the electrodes are exhausted. Any attempt for reversing the reaction via recharging in a primary cell is dangerous and can cause the battery to explode.

The electrical resistance in primary cells is usually high, thus, even if charging was possible it would be a slow process; at normal practical charging rates, a large proportion of the current would have been likely dissipated as heat, causing further safety hazards. Primary batteries are therefore designed to operate at low currents and have a long lifetime. Generally primary batteries have a higher capacity (Ah/Kg), a higher specific energy (Wh/Kg) and a higher initial voltage than secondary (rechargeable) batteries of comparable chemistries. They are used in portable devices, toys, watches, hearing aids, medical implants. The commercially used primary batteries are shown below.

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Primary battery chemistries in use

System Cathode/Anode

Nominal Cell

Voltage(V)

Specific Energy (Wh/Kg)

Advantages Disadvantages Applications

Carbon/Zinc 1.50 65 Lowest cost;

variety of shapes and sizes

Low energy density; poor low-

temperature performance

Torches; radios; electronic toys and

games;

Mg/MnO2 1.60 105 Higher capacity than C/Zn; good

shelf life

High gassing on discharge;

delayed voltage

Military and aircraft receiver-transmitters

Zn/Alk/MnO2 1.50 95 Higher capacity than C/Zn; good low-temperature

performance

Moderate cost Personal stereos; calculators; radio;

T.V

Zn/HgO 1.35 105 High Energy density; flat

discharge; stable voltage

Expensive; energy density only moderate

Hearing aids; pacemakers;

photography; military sensors/detectors

Cd/HgO 0.90 45 Good high and

low-temperature performance; good shelf life

Expensive; low energy density

Zn/Ag2O 1.50 130 High Energy density, good

high rate performance

Expensive (but cost effective for

niche applications)

Watches; photography;

missiles; Larger space applications

Zn/Air 1.50 290 High Energy density; long

shelf life

Dependent on environment; limited power

output

Watches; hearing aids; railway signals;

electric fences

Li/SOCl2 3.60 300 High Energy density; long

shelf life

Only low to moderate rate applications

Memory devices; standby electrical power devices; medical devices

Li/SO2 3.00 280

High energy density; best

low-temperature performance; long shelf life

High-cost pressurized

system

Military and special industrial needs

Li/MnO2 3.00 200

High energy density; good

low-temperature performance; cost effective

Small in size, only low-drain

applications

Electrical medical devices; memory

circuits; fusing

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Secondary Batteries In the secondary battery market, several new cells have been introduced just within the last two decades, for example, the Ni-MH cell (1990), Li-ion cell (1991), rechargeable alkaline cells (1992), Li polymer (1999), and also the concept of mechanically rechargeable Zn-air batteries (2001). The new batteries, especially those based on the Li chemistries have not only met some of the existing demand, but can be said to have revolutionized the battery market by accelerating demand for laptops, mobile phones and cordless hand held tool products.

An inspection of Fig. above clearly reveals that with time, the specific energy (Wh/Kg) and the energy density (Wh/litre) have continued to grow as batteries advance. It is this combination of high values of Wh/Kg and Wh/litre that have been the key factors heralding this rapid growth. The number of batteries produced each year is staggering. The relatively newcomers, Li-ion batteries, alone are produced in quantities exceeding 1.2 billion cells per year. The global market anticipated for the year 2015 is shown below: Table Global Battery Market Forecast for 2015 Type of Battery Global Demand in US$, 150 billion

Primary Primary Total : 40% Carbon Zinc 8% Alkaline 22% Others 12%

Secondary Secondary Total: 60% Lead Acid 28% Ni-Cd / Ni-MH and others 12% Li 14%

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Li-ion Batteries

In a Li-ion cell, both electrodes are based on insertion materials which can insert and de-insert lithium reversibly over many cycles. The most efficient secondary battery, though, is the lithium-ion battery which was commercially introduced by Sony in 1991. This cell chemistry included a LiCoO2 cathode and a carbon anode, whose potential after being fully charged is 4.2 V. The reversible reaction occurs by the intercalation reaction of lithium ions between the LiCoO2 and carbon frameworks in a non aqueous liquid organic electrolyte (1M lithium hexafluorophosphate (LiPF6) in 1:1 weight mixture of EC:DEC (ethylene carbonate : diethyl carbonate)). The face-centered cubic structure of LiCoO2 met the requirement of small volume changes (about 2%) during the Li-insertion and de-insertion that took place upon electrochemical cycling. The delithiated state is believed to maintain its layered structure, which is similar to a hexagonal close-packed structure.

The cell potential arising from the voltage difference between delithiatedLiCoO2

and lithiatedcarbon is large making it a high energy density battery. Graphite is a suitable anode material since its layered structure with hexagonal vacant sites allows lithium ions to be inserted with a very small volume change taking place (graphite interlayer expansion, which is about 11%) and can return back to its initial volume after the de-lithiation process. Graphite and other carbon anode materials can readily achieve this at a potential which is still relatively close to that of pure Li. While the rate of charging, cyclability and safety are significantly improved with a carbon anode, in comparison with a pure Li, the self-discharge rate is much higher.

The gravimetric capacity of this cell is 372 mAh/g and the corresponding

volumetric capacity is 800 mAh/dm3. However, a safety concern arises from the high reactivity between the lithium and the organic electrolyte in the deep charging state and may result in fire initiation. Use of a gel polymer electrolyte can prevent such risks. This gel electrolyte is used as a very thin film in order to overcome the resistance arising from the low conductivity of the polymer. A typical Li-ion battery can be schematically represented as shown in Figure below, while the working principle is depicted in the next Figure. In addition to the Co based compounds, the Li cathode can comprise of other chemistries based on Mn, P, Ni and Fe and various combinations.

27

28

Fuel Cells

A fuel cell is an electrochemical device that is able to continuously convert the chemical energy embodied in a fuel and an oxidant directly into electrical energy. The conversion to electrical energy takes place silently without combustion. Fuel cells are classified according to the type of electrolyte employed. The most common types are described in the Table below: Type Electrolyte Operating Temperature, oC Alkaline Fuel Cell (AFC) KOH (aq) 0 - 150 Phosphoric Acid Fuel Cell (PAFC)

H3PO4(aq) 200

Polymer Electrolyte Membrane Fuel Cell (PEM)

Nafion polymer 0 - 120

Molten Carbonate Fuel Cells (MCFC)

(K,Li)2CO3 550-700

Solid Oxide Fuel Cell (SOFC)

YSZ 900-1000

Polymer Electrolyte Fuel Cell The most widely used polymer membrane in PEM fuel cells is Nafion, produced by DuPont Company and readily available in the form of a membrane. It is a semi-crystalline co-polymer tetrafluorethylene and sulphonyl fluoride vinyl ether in a molar ratio of 7:1, with typical equivalent mass of around 1200 g. The hydrophobic regions dominate the structure, but hydrophilic sulphonate groups are placed in close proximity to the tetrafluoroethylene backbone and act as cross-links in the matrix, endowing the material with rigidity, inertness and insolubility. Nafion can absorb a large quantity of water and its conductivity for protons is dependent upon the water content, and thus the relative humidity, surrounding the membrane. Thinner membranes have higher conductivity but lower physical strength and higher permeability for the fuel. Lower strength can result in faster degradation and poor life. Higher permeability is responsible for fuel crossover and loss of electrochemical efficiency. Electrochemical Reactions: Considering hydrogen as the fuel and oxygen gas as the oxidant, the following reactions take place: Anode:

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H2 (g) 2 H+ + 2 e- Cathode: O2 + 2 H+ + 2 e- H2O Overall Cell Reaction: H2 (g) + O2 H2O (l) At 298 K, the standard free energy change for the overall reaction is given by: Go = -237,000 J = - 2 F Eo Thus the maximum cell potential or the open circuit potential (at zero current) can be calculated as 1.23 V. In a fuel cell the useful energy output is electrical energy produced and the energy input is the enthalpy of hydrogen (286,000 J/mol). If all the free energy can be converted to electrical energy, then the maximum fuel cell efficiency at 298 K is given as: = Go / Ho = 237,000 / 286,000 = 0.83 = 1.23/ 1.48 where the thermoneutral potential is 1.48V. Efficiency is reduced by many factors, such as crossover of fuel from anode to cathode via the electrolyte; unreacted fuel passing out of the anode zone; parasitic reaction of the fuel with contaminants such as oxygen. The mechanism for oxidation of hydrogen takes place by adsorption on a suitable electrocatalyst such as Pt on which the overpotentials are low and exchange current density high with Nafion (0.1 A cm-2). Presence of CO (in ppm levels) in the fuel, can poison the catalyst by adsorbing strongly and blocking the active surface. CO tolerance is improved when Pt is alloyed oxophilic elements such as Sn, Mo, W or Ru. The oxygen reduction reaction is a lot more sluggish with low exchange current density even on the best electrocatalyst Pt ( 6x10-6 A cm-2) and high overpotentials. This reaction is the major cause of low energy conversion in both PEM and Aqueous Electrolyts-based fuel cells. Problem arises from incomplete conversion to water by the 4-electron pathway per mole of oxygen (equation 25) rather losses from conversion to H2O2 by 2-electrons pathway and formation of other intermediates such as oxygen superoxide ions O2- and O22-. Despite major research on Pt, the oxygen reduction reaction is still the major limiting factor in PEMFC. In any event Pt is scarce even at the high price it commands; focus of research is shifting towards developing non-noble electrocatlaysts for both anode and

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cathode reactions. Use of metal particles (such as Pt, Pd, Au, Cu, Ni, Ta) and metal carbides (WC) on nanoparticles of oxides or carbon have shown to be promising avenues. Carbides also display CO tolerance. Fuel cells electrode reactions also follow polarization curves similar to that of batteries. The actual cell voltage is always less than the open circuit voltage (ocv). The overpotential losses are dependent upon the operating current density. The anode potential will move to a more positive value, thus polarizing towards the cathode while the cathode potential will move towards the anode to a more negative value. In general, as described previously for batteries Ecell = Erev a c I R When Pt is used, it is normally supported on C black. C black is made of nanoparticles of amorphous (and semi-crystalline) spheres aggregated into a network of highly porous structure. Aggregates can also form stable agglomerates. The structure permits flow of electrolyte and provides diffusion path for the gaseous fuel and the oxidant. CNTs are also being investigated as possible support materials and also for holding hydrogen within the tubes as well as intercalated in the graphitic structure. A schematic Diagram of Carbon black Structure (each C Nanopaticle is 5 -10 nm in size) Surface areas > 1000 m2/g are readily available. PEM fuel cells are also developed using liquid fuels such as methanol, ethanol or formic acid. The key performance measure of a fuel cell is the voltage output as a function of electrical current density drawn, or the polarisation curve shown below in Fig. below.

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Current density [A cm-2]

Schematic fuel cell polarisation and power density curves Solid Oxide Fuel Cells (SOFC) The most commonly used electrolyte in SOFC is yttria-stabilized zirconia (YSZ). SOFC electrodes are typically made of mixed electronic and ionic conducting materials, thus both electrons and ions can flow through them. In the figure below for anode as an example, ions coming from the electrolyte flow through the anode thus combining with H2 and releasing electrons at the so triple-phase-boundary (TPB). Consequently electrons flow towards the anodic current collector. The TPB concept holds that the H2 oxidation reaction and the oxygen reduction reaction can only occur at confined spatial sites, called triple phase boundaries, where electrolyte, gas and electrically connected catalyst regions contact.

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Diagram showing the triple phase boundary

The most commonly used anode for hydrogen oxidation for YSZ electrolytes is a porous cermet structure consisting of two interpenetrating and interconnecting networks of Ni-metal and Yttrium stabilized Zirconia particles. The gas phase in the pores account for more than 50 % of the volume. The large internal surface in the porous structure makes it possible to obtain large external current densities while keeping the local internal current densities and over potentials low. At the same time the pores enable the H2 molecules to penetrate all the way into the cermet and the H2O molecules produced to escape out of the electrode. The cathode is required to act as an electrocatalyst i.e. promoting the electrocatalytic reduction of O2 to O2-: O2 (g) + 4e- 2O2- Although several metals, such as Pt and Ag can also act as electrocatalysts for the above reaction, the most commonly used materials for the cathode are taken from the perovskites. The most common of all is the perovskite La1-xSrxMnO3- (LSM). This material demonstrates both ionic and electronic conductivity and thus the electrocatalytic sites are not restricted to the geometric TPB, but can also exist on the gas-exposed electrode surface as well, forming an electrochemically active zone, which can in principle, extend over the entire gas-exposed electrode surface. The electrical energy converted from chemical energy in the cell, is carried by electrons to the external circuit, by what are referred to as current collectors. Depending on the cell configuration, current collectors might also be situated between series cells, which are then referred to as interconnects. In either case they both serve the same purpose. In the past platinum has been used due to its high electrical conductivity and excellent adhesion, however its relative cost has not permitted further consideration for future designs.

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Until recently, the leading candidate material for the SOFC interconnect was electronically conducting doped lanthanum chromite, LaCrO3 a ceramic which could easily withstand the 1000C operating temperature of an electrolyte or air electrode supported SOFC design. It is typically doped with Sr and Ca in order to improve electrical conductivity and decrease sintering temperature to help facilitate dense impermeable ceramic monolith. It is still quite expensive, fragile and can be partially reduced by the fuel with time. It is desirable to be able to decrease the operating temperatures of SOFCs to < 700oC, so that high temperature metallic alloys (Cr and/or Ni based alloys) can be used as interconnects. Typical Designs:

Tubular Oxide Fuel Cell

Planar Oxide Fuel Cell

ELECTROLYSIS

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ELECTROWINNING Extraction of metals from aqueous solutions or molten salts in electrochemical cells - the metal is plated or deposited on to the cathode (-ve), while the anode is an insoluble conductor (e.g. a Pb alloy), where a gas such as oxygen or chlorine is evolved From Aqueous Solutions Copper sulphate in aq. solution is decomposed in an electrochemical cell: At cathode: Cu2+ +2e = Cu EO = + 0.34 V + RT/2F ln [Cu++] At anode: H2O = 1/2 O2 + 2H+ + 2e E

O = + 1.23 V - RT/F pH Net reaction for electrowinning copper is:

CuSO4(aq) + H2O = Cu + 1/2 O2 + H2SO4(aq)

G = 221.6kJ; E = -1.2V 1.2V is the minimum potential required to bring about the electrochemical reaction as shown above. Due to polarisation processes the operating voltages are nearly 2.5 V at a typical current density of 200 A/m2. The actual potential required is = 1.2 + A + C + IR Note: Electrolytic extraction from aqueous solutions is possible for Ni, Co, Cd, Fe, Zn and Mn, by decomposing their salts in aq solutions without the danger of decomposing H2O to form hydrogen gas at the cathode If the reversible electrode potential required to deposit a metal is < - 1 V (or Total potential > 3 V), the possibility of evolving hydrogen at the cathode increases

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ELECTROREFINING: Electrorefining of copper anode consists of applying an electrical potential between the copper anode (+ve) and a starter sheet (pure Cu, stainless steel or Ti) as the cathode, both immersed in a cell containing an acidified CuSO4 solution Overall Cell Reaction: Cu (anode, 99.4 %) = Cu (cathode, 99.999 %) G = 0 and Emin = 0 V E(actual) = 0.3V to overcome the resistance of the electrolyte and the external circuit + overpotentials to carry out the reactions at reasonable rates Typically, I = 200 A/m2 and Energy consumption = 300 KWH per tonne of Cu deposited Au, Ag, Pt and other precious metals do not dissolve in the electrolyte and are collected in the anode slimes S, Se, Te, Pb and Sn are also collected in the anode slimes As, Bi, Co, Fe, Ni and Sb dissolve in the electrolyte, but do not deposit in the cathode and are removed by purification of the electrolyte

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ELECTROWINNING OF ALUMINIUM: Reactive metals such as Al, Mg, Li, Ca, Na, K etc cannot be deposited from their salts in aqueous solutions but are deposited from their fused salts Bauxite ore is purified in Bayer's process by leaching with NaOH(aq) in order to produce pure Alumina for electrowinning Al Alumina can be dissolved up to 8% in a fused salt of cryolite NaF-AlF3 at 1273K and electrochemically decomposed: Al2O3(l) = 2Al(l) + 3/2 O2 (g) G (1273K) = 1279 kJ E(min) = -2.2V The molten electrolyte is covered with a crust of alumina which acts as a reservoir of feed as well as insulation Both the anode and the cathode is made up of graphite - the graphite anode consumed during the process as a result of oxidation with oxygen evolved at the anode, which leads to a reduction in the min. potential required: Al2O3(l) + 3/2 C = 2Al(l) + 3/2 CO2 (g) for which G (1273K) = 686 kJ E(min) = -1.18 V; E(total) = 6 V (due to and IR losses) Primary Al production is a power intensive process (17 KWH/Kg of Al)

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Photocatalytic Redox Reactions

In photocatalysis, an electron in the valence band is excited to the conduction band (see Figure below) if light of energy greater than the band gap is provided to the semi conductor. This excitation creates charge carriers, electron and holes in CB and VB respectively. The charge carriers migrate to the surface of the photocatalyst where the redox reaction takes place with adsorbates. In an aqueous environment, water is oxidized by the hole which generates powerful oxidizing agents OH, called hydroxyl free radicals (OH) which can readily oxidize organic compounds accompanied with mineralization, and evolution of CO2 and H2O. The standard electrode potential values, which determine the strength of oxidants, are given for OH relative to the other common oxidant in the table below.

The excited electrons in the conduction band are trapped by dissolved oxygen in water which results in the formation of super oxide anion (O2) by electrochemical reduction. The different type of redox reactions which can take place are shown in reactions below. The super oxide anion (O2) can further react with H+ which form hydroperoxyl radical (OOH) which can further react to produce H2O2.

Photochemical and Redox Reactions

TiO2 + hv +VBh + CBe

H2O + +VBh OH + H+

O2 + CBe 2O

OH + *pollutant H2O + CO2

2O + H+ OOH

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OOH + OOH H2O2 + O2

2O + *pollutant CO2 + H2O

OOH +* pollutant CO2 + H2O *Pollutant here refers to organic compounds.

Oxidant Reduction potential (Eo )/V

OH (Hydroxyl radical) 2.8

O3 (Ozone) 2.07

H2O2 (Hydrogen peroxide) 1.77

HClO (Hypochlorous acid) 1.49

Cl2 (Chlorine) 1.36

Standard electrochemical reduction potentials of common oxidants vs SHE at 298K

During photocatalysis, the positions of both conduction band (CB) and valence band (VB) of a photocatalyst play important roles, denoting the relative power of electrons for photo-reduction and that of holes for photo-oxidation. Using photolysis of water to H2 and O2 as the example to understand the reduction and oxidization power respectively, the relative position of the CB and VB are considered. The CB position must be above the H+ ion/ H2 redox couple level in the electrochemical series for reduction to take place and in similarly the VB position must be below the O2/ H2O redox level in the electrochemical series for oxidation reaction to take place. A schematic showing the position of CB and VB for different semiconductors with respect to H+/H2 and O2/H2O is shown in Figure below. Some metal oxides are unsuitable for producing H2 in an aqueous system, due to the mismatch of their CB and VB with respect to the position of H+ / H2 and O2/H2O from the electrochemical series.

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Valence and conduction bands for a number of semiconducting materials on a potential scale (V) versus the standard hydrogen electrode (SHE). Redox potentials for the water-splitting half reactions are indicated by the dotted red lines. In order for a semiconductor to be an effective catalyst its conduction band energy should be higher than the H2 producing reaction potential and its valence band energy should be lower than the O2 producing reaction potential

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C2 - Examples Class

1. Pure NiO is a p-type semiconductor which forms vacancies on the cation sub-lattice in a metal deficient Ni1-xO (where x = 0.001 at 600K). Show that electrical conductivity of nickel oxide will vary as 6/1 2Op . The holes are 100x more mobile than the Ni vacancy at 600K. What is the ratio of ionic to electronic conductivity? What will be the effect of dissolving (a) 0.01 mol % of Li2O; (b) 0.01 mol % of Cr2O3 in NiO (c) Over oxidising NiO such that some of the Ni appear to be in Ni3+ state ? In a NiO-0.01 mol % Cr2O3, polycrystalline sample with an average grain size of 10 nm, the effective Ni vacancy concentration was found to increase by an order of magnitude. Discuss the implication of this finding. 2. Electrochemical cells are divided into electrolytic cells and galvanic cells. What is the characteristic feature of an electrolytic cell, and what is the characteristic feature of a galvanic cell? For each type of cell, name two examples that are of practical relevance. Tin is used as a coating on iron to improve corrosion resistance of the iron component. The electrochemical series gives the following standard reduction potentials for the Sn2+/Sn and the Fe2+/Fe redox couples: 2

0/

0.136 VSn Sn

E + = ; 20 / 0.440 VFe FeE + = . When combining both half-cells, what is the overall cell reaction, what is the standard cell potential, and what polarity do the two metal phases adopt? What is the molar Gibbs free energy of the cell reaction under standard conditions, and what is the equilibrium constant? Which concentration ratio of the two types of ions would be required in order to generate a cell potential of 0 V? (assume unit activity coefficients) [T = 298 K , R = 8.314 J mol-1 K-1 , F = 96500 As mol-1] 3. Impure Cu (99.4 wt%) is electrorefined at 298 K to a pure Cu (99.999 wt%) in an aqueous acidified electrorefining cell containing 1M H2SO4 (aq) and 1M CuSO4 (aq) electrolytic solution at a current density of 200 A m-2. The surface area of the electrode is 1 m2. Answer the following using the Data given below:

(a) Calculate the reversible cell potential

(b) Calculate the applied cell potential required for sustaining a current density of 200 A

m-2.

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(c) Discuss the fate of the following impurities present in the impure anode: Ni, Pb, Ag and Au and explain why electroefining can help produce very pure copper [40%]

(d) Can hydrogen or oxygen be evolved at the electrodes?

[Tripos 2014] Data: (all data specified at 298 K) Standard Electrode Potentials

AueAu + + 33 +1.50V

OHeHgO 22 44)( ++ + +1.23V

AgeAg + + +0.80V

CueCu + + 22 +0.34V

)(22 2 gHeH + + 0.00V

PbePb + + 22 -0.13V

NieNi + + 22 -0.25V The value of the limiting current density for reduction of Cu2+(aq) , Li = 205 A m

-2 Tafel slope for oxidation of Cu, b = 0.295 V per decade The exchange current density for CueCu + + 22 is, oi = 1 A m-2 The ohmic resistance, R of the electrolyte is given by 6 x 10-3

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C2 Question Sheet 1. The following solid-state oxygen concentration cell was set up using yttria stabilized zirconia (YSZ) as the electrolyte: Pt, Ni, NiO | YSZ | Cu-Ni, NiO, Pt At 1200K, the potential of the cell (electro motive force) was 0.062V, for 30 mol % Ni in the Cu-Ni alloy. What is the activity of nickel in the alloy? What can you say about the copper-nickel solid solution? 2. A fuel cell consists of hydrogen/Pt and an oxygen/Pt electrode with both gases at 1 atmosphere. Neglecting ohmic losses and concentration overpotentials, calculate the power density (W m-2) of the fuel cell at a current density of 100 A m-2, given: At 298 K, the standard free energy change for the reaction: H2 (g) + O2 H2O (l) is Go = -237,000 J [The Tafel constant, b for the hydrogen and the oxygen reactions are 0.0295 and 0.09 V per decade of current density respectively, while the exchange current densities for the condition in the cells are 10-4 and 10-7 A m-2 respectively; Faradays Constant = 96540 J mol-1 K-1] 3. In the context of a discharging a battery, distinguish between activation polarisation and concentration polarisation at a zinc anode plate. Calculate the reversible electrode potential for pure Zn in contact with an electrolyte with effective concentration of Zn2+ at 0.2M, given the standard electrode potential is -0.76V. The exchange current density for the above reaction is 0.01 A m-2 and the anodic Tafel slope b is 0.11V. Calculate the operating electrode potential for a current density of 1 A m-2. 4. Using Faradays laws, calculate the theoretical capacity in Ah per Kg of pure Li as an anode in a Li-ion cell. (Molecular wt of Li: 6.94). In practice, for safety reasons, carbon is used as the anode into which Li can be intercalated up to a maximum of 1 atom of Li per 6 atoms of C. Calculate the theoretical capacity (Ah) per Kg of the C anode. In the cathode, 1 Li atom can be intercalated per CoO2 molecule. Calculate the theoretical capacity (Ah per Kg) of the cathode. It is found that both the anode and the cathode can practically and safely utilize only 50% of the available intercalation sites. In a practical Li-ion cell, 20% of weight is accounted

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for by the electrolyte, separator, current collector, binders and conductive additives in the cathode. Calculate the capacity of a practical Li-ion cell (i) per Kg of Li and (ii) per Kg of the cell weight. The open circuit cell voltage decreases from a maximum initial value of 4.4 V for a fully charged Li-ion cell [corresponding to LiC6 anode and Li0.5CoO2] to 2.5 V corresponding to Li0.5C6 anode and LiCoO2] when it is discharged for an average cell potential of 3.5 V. Calculate the average specific energy density (Wh/Kg) of the cell with respect to 1 Kg of the cell weight. Write the anodic, the cathodic and the overall cell reactions during discharge and during charging.

C2 coverC2 Course Handout