Chapter 1_Basic Concepts(4)

1.1-1.5 CHBE 477/577: Fuel Cells and Electrochemical Engineering-Dr. Eld Gyenge 1

Chapter 1: Basic Concepts

Electrochemical processes can be classified as either electric power consuming (i.e.,

power sinks) or electric power generating (i.e., power sources or galvanic cells). The

first basic requirement for electrochemical power sources (batteries, fuel cells) is that

the electrochemical transformations taking place in the system must be

thermodynamically spontaneous, i.e., the Gibbs free energy of the power source

(Gcell) must be negative. In case of power sinks on the other hand, (e.g.,

electrosynthesis (inorganic and organic), electrodeposition of metals) the cell reaction

thermodynamically is not spontaneous, the Gibbs free energy is positive, hence, an

external dc electricity source is required to carry out the electrochemical reactions.

Fig. 1.1 Conceptual representation of electrochemical systems.

R1

O1 +

ne

I

Ecell

Anode

(+) ()

Cathode

Porous

Separator

R2

O2 +

ne

(DC power supply)

e-

Power Sink Thermodynamically

not spontaneous system

Gcell > 0

() (+)

Power Source Thermodynamically

spontaneous system

Gcell < 0

R1

O1 +

ne

I

Ecell

Anode

() (+)

Cathode

R2

O2 +

ne

Load

e-

Anode: electrode where oxidation occurs Cathode: electrode where reduction occurs occurs

Batteries, Fuel Cells

Electrolyte Electrolyte

Electrolysis Electrodeposition

Electrosynthesis


1.1 Thermodynamic considerations using Voltas pile as an example

Chemical species: Zn, Ag, O2, N2, H2O, Na+, Cl

Fig. 1.2 Schematic diagram of Voltas pile comprised of two cells connected in series.

Features of Voltas pile:

1. Zn anode

2. O2 (or air) cathode

3. Separator (electronic insulator, ionic conductor)

4. Electrolyte (salt water)

5. Bipolar plate stacking to increase the cell voltage (observe how the two cells are

connected in series in Fig. 1.2).

Table 1.1 Potential electrochemical reactions involving chemical species present in

Voltas pile and the corresponding standard Gibbs free energies of reactions.

Nr. Potential electrochemical reactions G0 (kJ) [at 298 K]

1 Zn Zn2+ + 2e 147.1 2 1/2O2 + H2O + 2e

2OH 77.4 3 Ag Ag+ + e 77.1

4 Na+ + e Na 261.9

5 2Cl Cl2 + 2e 262.6

In Table 1.1 the standard Gibbs free energies for the reactions were calculated from

the tabulated standard Gibbs free energies of formation (fG0) of the chemical species

(Langes Handbook 15th edition Table 6.3). Elements in their most basic, natural, form

have fG0 = 0.

Aerated salt water soaked in a porous material (Note in Voltas days a cheese cloth or wood bark was used. Now various porous polymeric materials are employed as separators).

Ag (or Cu): + electrode

Zn: electrode

Loadad

A

V

e


Based on Table 1.1 there is only one combination of two electrochemical reactions

(oxidation and reduction) that could lead to a thermodynamically spontaneous Volta pile

characterized by a very negative standard Gibbs free energy:

Anode (): Zn Zn2+ + 2e oxidation (1.1)

Cathode (+): 1/2O2 + H2O + 2e 2OH reduction (1.2)

Volta cell: Zn + 1/2O2 + H2O Zn(OH)2 (1.3)

00 ]2.1[0

]1.1[

000 GGGGG redoxidcell . (1.4)

G0cell = -147.1-77.4 = 224.5 kJ molZn

-1.

Since G0cell < 0, it means thermodynamically SPONTANEOUS system. Thus,

theoretically, the combination of reactions (1.1) and (1.2) could function as an

electrochemical power source. Whether it is a practical system, it is a more complex

question and it depends on a number of factors such as rate of electrochemical reactions,

electrolyte employed, cell design and cost.

Based on thermodynamic considerations it can be shown that the Gibbs free energy of

the cell (for either standard G0cell or non-standard (actual) states Gcell) is equal to the

useful non pressure-volume work exerted or received by the system at equilibrium.

thermocellcell wG (1.5)

For an electrochemical power source (Gcell < 0), eq. (1.5) means that the Gibbs free

energy of the cell is equal to the thermodynamic maximum electric work the system

could produce, wcell-thermo. The latter can never be achieved in practice since eq. (1.5)

implies an ideal reversible systems, where there is no net flow of electrons in either

direction, hence there is no net consumption of reactants. In other words eq. (1.5)

describes the electrochemical equilibrium case and in practice, when the cell is operating,

the generated electric work is always smaller than the thermodynamically determined

maximum value due to various losses that will be discussed in detail in upcoming

chapters. Thus, for a practical electrochemical power source the actual electric work

produced wcell is:

.thermocellcell ww (1.6)


Note: for electric energy and work is common to use [Wh] or [kWh] as unit. 1 kWh =

3.6 MJ =3.6x106 J

For an electric power consuming cell Gcell > 0 (e.g., electrolysis reactor), the

reaction is NOT spontaneous, hence, electric work in the form of electric charge flow

needs to be provided to the reactor to carry out the reaction. In this case, eq. (1.5)

expresses the minimum electric work that has to be provided by the external power

supply in order to drive the electrochemical reactions. In practice, more work has to be

provided than that expressed by (1.5) in order to overcome the various losses in the

system:

.thermocellcell ww (1.7)

1.2 The electrochemical equilibrium

Consider the electrode interface where an electronic conductor (or in some cases a

semi-conductor) is in contact with an ionic conductor phase (e.g., liquid or solid

electrolyte). The chemical species from the ionic conductor phase can be in equilibrium

with the electrons or vacancies from the conduction band of the electronic conductor.

Thermodynamics dictates that at equilibrium the Gibbs free energy change across the

electrode interface (better referred to as the electrochemical Gibbs free energy change

across the electrode interface, G

) must be zero.

For any electrochemical (half-cell) reaction (see for example Table 1.1), separating

conceptually the electrons and the chemical species at the electrode interface, the general

stoichiometry can be written as:

jz

j

jjMsne ; (1.8)

) phase te,(electrolyconductor ionic ) (phaseconductor electronic .

Where: zj = 0, and/or, +, and/or

nszj

jj . (1.9)

n- nr. of electrons exchanged in the half-cell reaction, sj stoichiometric coefficient, zj

oxidation state. Mj symbolizes a chemical species. The electrode material itself might or

might not be a chemically active participant in the electrochemical reaction.


Note: eq. (1.8) implicitly is written in the reduction direction, with the electrons on

the left-hand side. This means that sj is taken as negative if species j in the reaction

stoichiometry was on the electron side (i.e., reactant oxidized species) and positive if

species j in the reaction stoichiometry is situated on the opposite side vs. electrons

(product reduced species).

Under electrochemical equilibrium conditions there is no net flow of electrons at any

temperature. In other words, the rates of the half-cell oxidation and reduction reactions

are equal. Hence, the net electrode reaction rate is zero. Thermodynamically the

electrochemical equilibrium is defined as the Gibbs free energy change between phases

(electrolyte ionic conductor) and (electronic conductor) is zero:

efj

jfj GnGsG 0

, (1.10)

where ef G is the Gibbs free energy of formation per mole of electron in the conduction

band, jf G is the Gibbs free energy of formation of the chemical species Mj involved in

the half-cell electrode reaction.

The Gibbs free energy of formation of the electrons (J mole--1

) can be expressed in

terms of the equilibrium electrode potential Ee multiplied by the charges carried:

eefFEG (1.11)

where Ee is the equilibrium electrode potential at any temperature the system is at, and F

is Faradays constant, expresses the charge carried by one mole of electrons (~96,500 C

mole--1

). The negative sign in eq. (1.11) corresponds to the electron charge.

Note the units: [Joule] = [Coulomb] x [Volt].

Combining eqns. (1.10) and (1.11) for electrochemical equilibrium at any temperature

we obtain:

0 ej

jfj nFEGsG

, (1.12)

or

.ej

fj nFEGsG j (1.13)


Eq. (1.13) is fundamental for electrochemical thermodynamics because it allows the

calculation of the equilibrium electrode potential from thermodynamic data with respect

to the free energy change of the reaction (G).

For the particular case when the activities of the chemical species are all equal to one,

(i.e. for ideal cases this means molar concentrations of 1 M and/or pressure = 1 bar (

1atm), see also further Chapter 2), eq. (1.13) becomes:

.000 nFEGsG

j

fj j (1.14)

E0 is the standard electrode potential at any temperature. In Chapter 2 the temperature

dependence of E0 is further discussed in detail.

Furthermore, it is of interest to note the thermodynamic relationship between the

equilibrium constant K and the standard potential E0 at any temperature, using the well-

known thermodynamic equation:

KRTG ln0 (1.15 I)

.ln0 KnF

RTE (1.15 II)

Combining now two half-cell electrode reactions of general form (1.8) into a

stoichiometrically balanced cell reaction (i.e., n is the same for both half-cell reactions):

cathode jz

j

cjcj Msne ,, , cec nFEG , (1.16 I)

anode jz

j

ajaj Msne ,, aea nFEG , (1.16 II)

-----------------------------------

[Cell] = [Cathode] [Anode]

jj z

j

ajaj

z

j

cjcj MsMs ,,,,0 .

At any temperature, at equilibrium:

celleaccell nFEGGG , (1.16 III)

or similarly, under standard conditions (i.e. activities equal to one):

00

cellaccell nFEGGG . (1.16 IV)


From eqns. (1.16 I) (1.16 IV) the cell potential at equilibrium under non-standard

(actual) conditions (Ee,cell) and under standard equilibrium conditions (E0cell) is expressed

as:

,,,, aececelle EEE (1.17 I)

.000 accell EEE (1.17 II)

At this point is important to note the IUPAC (International Union of Pure and

Applied Chemistry) electrochemical convention. Accordingly, the electrode reactions are

always considered in the reduction direction (i.e., with the electrons as reactants on the

left hand side of the stoichiometric equation), irrespective of the actual direction in an

electrochemical system. Thus, all the electrode potentials are expressed for reductions.

ccathodered EEE , (cathode is the electrode where reduction occurs) (1.18 I)

aanodeoxid EEE , (anode is the electrode where oxidation takes place). (1.18 II)

From (1.16 III and 1.6 IV):

- for electrochemical power sources (and galvanic corrosion): Ee,cell and E0cell > 0.

- for electrolysis (power consuming systems): Ee,cell and E0cell < 0.

Standard electrode potentials E0 for inorganic and organic electrochemical reactions

are listed at 298 K in many databases such as Langes and CRC Handbook of Chemistry.

According to convention, E0

is given for the reduction direction (i.e., electrochemical

reactions are written with electrons on the left hand side).

For the H+/H2 redox couple E

0 = 0 V at 298 K because G0 = 0 at 298 K, due to zero

standard Gibbs energies of formation for both H+ and H2 (see Langes Handbook 15

th ed.

p. 6.93). Hence, the standard hydrogen electrode (SHE) is used as primary reference

electrode for measurement of the electrode potentials (see Chapter 2). Table 1.2 lists

selected standard electrode potentials for inorganic and organic reactions.


Table 1.2 Selected standard electrode potentials at 298 K (source CRC Handbook

of Chemistry and Physics 75th

ed. and Langes Handbook 15th ed.)

Electrochemical Reaction E0 at 298 K (V vs. SHE)

Co3+

+ 1e- Co2+ 1.92

Ce4+

+ 1e- Ce3+ 1.72

N2O + 2H+ +2e

- N2 + H2O 1.77

PbO2 +SO42-

+4H+

+ 2e- PbSO4 + 2H2O 1.69

O2 + 4H+ + 4e

- 2H2O 1.23

2H+ + 2e

- H2 0

O2 + H2O + 2e- HO2

- + OH

- -0.076

CO2 + 2H+ +2e

- HCOOH -0.20

Ni(OH)2 + 2e- Ni + 2OH- -0.72

Zn2+

+ 2e- Zn -0.76

Li+ + e

- Li -3.04

Example:

A) for O2 electroreduction under neutral or alkaline conditions, written under the

form of eq. (1.8):

2e 2OH - 1/2O2 - H2O. (1.19)

Using eq. (1.14):

.2/122 0000/

0

222OHfOfOHfOHO

GGGFEG (1.20)

Substituting the Gibbs free energies of formation at 298 K in J mol-1

(e.g., data from

Langes Handbook 15th ed., Table 6.3):

).1014.237(02/1)1028.157(22 330/2

xxxFEOHO

K). 298(at SHE vs.V 4.00/2

OHOE

B) The standard cell potential for Voltas battery at 298 K is calculated using eq.

(1.16 I) (1.16 IV).

() [anode]: Zn2+ + 2e Zn o

ZnZnG

/2 147.1 kJ; E

0a = 0.76 V vs. SHE


(+) [cathode]: 1/2O2 + H2O + 2e 2OH

o

OHOG

/277.4 kJ ; E0c = 0.40 V vs. SHE

Cell = [cathode] [anode]

Standard cell potential at 298 K:

E0

cell = E0

c E0a = 0.40-(-0.76) = 1.16 V. (1.21)

1.3 The rate of electrochemical transformation and Faradays law

A fundamental question for electrochemical processes is the correlation between the

electric current and the consumption/production of chemical species. Faradays constant

expresses that 1 mol of e carries a charge of F = 96,500 Coulombs = 26.8 Ah.

488,96106.110023.6 1923 xxeNF A C, (1.22)

where NA is Avogadros number, and e is the elementary charge.

In other words, Faradays law states that 1 equivalent-gram of any species if undergoes

electrochemical transformation with 100% current efficiency, involves F = 96,500

Coulombs of charge.

Consider the general electrochemical stoichiometry given by eq. (1.8). The rate of

consumption or production of a species Mj in the electrode reaction normalized per

electrode area A, can be related to the current density across the electrode interface i, by:

nF

i

Adt

dn

sr

j

j

j 1

, (1.23)

where nj is the number of moles of species Mj.

Note: the units for the reaction rate

[rj] = [mol m-2

s-1

], just like for any heterogeneous reaction.

The current density i, in SI units is expressed in [A m-2

] (where A Amperes).

However, units such as [mA cm-2

] are also very common in the literature. The current

density is defined in terms of the current I, as:

A

Ii . (1.24)

If:

A is the geometric electrode area, than i is referred to as the superficial current

density;


A is the electrocatalytically active area, than i is referred to as the real (or

actual or effective) current density. The electrocatalytically (or electrochemically)

active area can be determined using special techniques such as H2 underpotential

deposition and stripping, COad voltammetric stripping or Cu underpotential

deposition and stripping

By convention the following sign rules apply:

at the cathode, i and I < 0,

at the anode, i and I > 0.

For the cell (composed of the anode and cathode):

cacell III . (1.25)

Example:

For a 5 cm2 geometric area anode of the Volta pile, how much Zn is consumed to produce

100 mA cm-2

current density for 20 minutes?

Writing the anode reaction in terms of eq. (1.8) (see also Table 1.1):

2ZnZn2e , (1.26)

sZn = +1, sZn(2+) = -1, n = 2.

From eq. (1.24):

F

i

Adt

dnr ZnZn

2)1(

. (1.27)

Since it is an anode: i > 0, Zn is consumed thus, dnZn < 0.

Substituting the numerical values, gives a Zn consumption of: 0.2 g.

Example:

For a 5 cm2 geometric Ag cathode area of the Volta pile, how much O2 is consumed and

how much hydroxide is produced when the cell operates at 100 mA cm-2

current density

for 20 minutes?

Based on eq. (1.22): sO2 = -1/2, sOH(-) = 2, n = 2

F

i

Adt

dnr

O

O2)2/1(

2

2

. (1.28)

Since it is a cathode: i < 0, and O2 is consumed thus, dnO2 < 0.

F

i

Adt

dnr OHOH 2)2(

. (1.29)


At the cathode: i < 0, and OH is produced thus, dnOH(-) >0.

Substituting the numerical values gives: 1.55x10-3

mole O2 consumed and 6.22x10-3

mole OH produced.

In cases when more than one electrochemical reaction can occur at the same

electrode, the current efficiency must be taken into account. The current efficiency

indicates the fraction of the total current consumed by a particular electrochemical

reaction.

,ii kk (1.30)

where ik is the current density corresponding to reaction k occurring at the electrode

surface, k is the current efficiency for reaction k (0 k 1) and I is the total current

density at a particular electrode (anode or cathode).

k

kii , (1.31)

Therefore,

.1k

k (1.32)

For virtually all practical electrolysis systems the current efficiency for the main (or

desired) electrode reaction is below 100%. During electrolysis in aqueous solutions at the

cathode, the electrochemical generation of H2 is a ubiquitous secondary reaction

competing with the main (or principal) reaction. Similarly, during electrolysis in aqueous

solutions at the anode oxygen evolution is a common secondary reaction.

As an example, during electrodeposition of Cu from aqueous acidic solutions, at the

cathode (negative electrode in this case) the following reactions can take place: 1.

electroreduction of Cu2+

, 2. evolution of H2,(g) from the hydronium ions in the acidic

solution, and 3. reduction of dissolved O2. Thus, k in eq. (1.31), (1.32) is k =1, 2, 3.

1. Cu2+ + 2e Cu

2. 2H+ + 2e H2,(g)

3. O2 + 2e +2H

+ H2O

In case of efficient Cu electrodeposition plants the current efficiency for Cu deposition

(option 1, main reaction) is typically between 90-95% (i.e., 1 = 0.9 0.95). The current

efficiency is a function of a large number of factors that must be engineered for optimal


performance such as electrolyte composition, hydrodynamics, electrode and cell design.

The total current density is expressed as the sum of the partial current densities for each

reaction. In case of Cu electrodeposition:

321 iiii . (1.33)

Including the current efficiency the general rate equation is expressed as:

nF

i

Adt

dn

sr

kj

j

j

1. (1.34)

Integrating eq. (1.34) with I = const. (galvanostatic condition):

t

k

j

jjj dtiAnF

snnn

0

0, . (1.35)

Assuming k = const. over time (which is not always the case because due to electrode

fouling and under factors, the current efficiency is likely to change over time), eq. (1.35)

can be easily integrated, generating in essence the mathematical form of Faradays law:

.tiAnF

sn k

j

j (1.36)

Note index k refers to a particular reaction on the surface which involves a number of

species j.

If the concentration is expressed instead of moles,

,tainF

st

V

Ai

nF

sC sk

j

k

j

j (1.37)

where as - specific surface area of the electrode, it is the area per volume of the electrodic

compartment (either anodic or cathodic compartment) or total cell volume, depending

where the concentration of species j is defined:

V

Aas . (1.38)

In order to increase the conversion of species j, it follows from eq. (1.37) that as

should be increased. Therefore, porous electrodes have been introduced with specific

surface areas between 100 to 20,000 m2 m

3. Porous electrodes can have various

configurations such as bed of particles, metal meshes, fibrous materials. Fig. 1.3 presents

the diagram of an electrochemical cell equipped with a porous cathode.


Fig. 1.3 Schematic of a divided electrochemical cell

(reactor) equipped with a porous cathode. O indicates oxidized species which are reactants at the

cathode; R indicates reduced species which are the products at the cathode. The anode could have

either a separate line of reactants or products or it

could be fed from the cathode side with the same

species. The anode and cathode are electronically

separated but ionically connected using a porous

separator material which is filled with electrolyte.

1.4 Polarization curves, power and energy density

Directly related to the rate of the electrode processes is the polarization phenomena,

which in essence is responsible for the deviation of the electrode and cell potentials under

operating current density conditions from the values given by thermodynamics. In other

words, the dependence of the electrode or cell potentials on the current density is referred

to as polarization.

At the cathode: cec EiE ,)( ; at the anode: aea EiE ,)( .

The cell potential - current density profile for an electrochemical power source is

expressed by polarization curves such as those shown in Fig. 1.4. The open circuit cell

potential Eoc (defined at i = 0) is typically somewhat lower than the equilibrium

(reversible) cell potential Ee,cell. The dependence of the operating cell potential on current

density depends on a number of phenomena such as electrode kinetics, ionic conductivity

and mass transfer.

anode

Ecell I

O

R

cathode separator


Fig. 1.4 Cell potential as a function of superficial current density for an electrochemical

power source (polarization curve). A and B represent two different performance

characteristics, such as beginning of life and end of life performance, respectively. Ee,cell is the equilibrium cell potential, Eoc is the open circuit cell potential (at i = 0).

The electric power output for a power source (Ecell > 0) or the power consumption for

an electrolysis unit (Ecell < 0) per cell (or reactor) volume Vcell:

cell

cellcell

VV

EIPD . (1.39)

where PDV is the volume based power density (W m3

).

If the anode and cathode geometric areas are equal (i.e., Ac = Aa = A), than it is common

to use the current density ( ca iii ) and to calculate an area-specific power density:

cellcellcell

A EiA

EIPD , (1.40)

where PDA is the area-specific power density (W m2

).

A major issue in the design and operation of electrochemical power sources is to

increase the power density, while for electrolysis the goal is to lower the power

consumption.

The polarization curves of the type shown by Fig. 1.4, when converted to power

density generate a parabolic function with respect to current density (Fig. 1.5).

Current density, i (A m2)

0

Cell

po

ten

tia

l, E

cell

(V)

A

B

Ee,cell

Eoc


Fig. 1.5 Power density for an electrochemical power source as a function of

superficial current density.

1.5 Review and classification of fuel cell systems

Fig. 1.6 Basic components of a generic fuel cell [modified after Larminie and Dicks,

2001].

Fu

el

(e.g

. H

2 )

and Separator

() (+)

Ox

ida

nt

(e.g

. O

2 )

(contains a catalyst layer, gas diffusion region and flow field-current collector plate)

(contains a catalyst layer, gas diffusion region and flow field-current collector plate)

~ 3 mm

Current density, i (A m-2) 0

PD

A (

W m

-2)


Table 1.3 Examples of oxidant choices for fuel cells and the standard cathode potentials

at 298 K

Oxidant Half-cell cathode reaction Eoc (VSHE)

Oxygen

In acid: O2 + 4H+ + 4e 2H2O

In alkali: O2 + 2H2O +4e 4OH

1.23 0.40

Hydrogen peroxide H2O2 + 2H+ + 2e 2H2O 1.78

Ozone O3 + 2H

+ + 2e O2 + H2O 2.08

Table 1.4 Possible fuel choices for fuel cells and the respective standard anode potentials

at 298 K.

Fuel Half-cell anode reaction (Note: indicates the actual direction of the

reaction in the fuel cell)

Eoa (VSHE)

H2

2H+ + 2e H2 0

Methanol CO2 + 6H+ + 6e CH3OH + H2O

0.04

Ethanol Formic Acid Hydrocarbons

2CO2 + 12H+ + 12e C2H5OH + 3H2O

CO2 + 2H+ + 2e HCOOH

nCO2+ (6n+2)H+ + (6n+2)e CnH2n+2+ 2nH2O(g)

0.08

0.20 0.17 (for n=1 CH4)

Coal CO2 + 4H+ + 4e C + 2H2O(g) 0.21

Carbon monoxide CO2 + 2H

+ +2e CO + H2O(g) 0.11

Hydrazine N2 + 4H+ + 4e N2H4 0.33

Biomass rich in sugars (e.g. glucose, maltose)

Borohydride

S + 2H+ + 2e SH2 (gluconate) enzyme (glucose)

BO2 + 6H2O + 8e

BH4 + 8OH

0.47

1.24


Table 1.5 Classification of fuel cells based on the type of ionic conductor (electrolyte)

No. Fuel Cell Typical Electrolyte and separator

Symbol

1

Solid polymer electrolyte (or polymer electrolyte-membrane) < 90 oC (occasionally up to ~130 oC)

Proton-exchange membrane (fluorinated organic polymer scaffold with acidic (-SO3H) groups) functions as both electrolyte and separator

SPE PEM

2 Alkaline 80 200 oC

KOH (or NaOH) asbestos matrix

AFC

3 Phosphoric acid 200 oC

H3PO4 silicon carbide matrix

PAFC

4 Molten carbonate (or molten salt) 600 800 oC

mix. of LiCO3 K2CO3 ceramic matrix of LiAlO2

MCFC

5 Solid oxide 600 1000 oC

ZrO2 stabilized with Y2O3 functions as both electrolyte and separator

SOFC


Fig. 1.7 Schematic representation of the ion-conduction through electrolyte for solid

oxide (SOFC), proton exchange membrane (PEM), phosphoric acid (PAFC) and molten

carbonate fuel cells (MCFC) [Kinoshita, 1992].

PEM and


Schematic representation of a fuel cell gas diffusion electrode

Fig. 1.8 Schematic of the electrolyte-electrode interface for fuel cells.

Note: in the case of PEM fuel cells a polymer electrolyte membrane is used and the

electrode/electrolyte system is referred to as membrane-electrode assembly (MEA)

[modified after Larminie and Dicks, 2001].

Cb

Cs

Reactant conc.

Profile (e.g. O2)

Macro-porous material, e.g. carbon cloth or carbon paper Pore diameter:

~ 3 50 m

Catalyst

layer,

Reaction

sites

H+; H2O

~ 100-200 m ~50 - 200 m

and Electronic Separator

Meso-pores filled with electrolyte between the catalyst particles and particle agglomerates Pore diameter: 2- 50 nm ~ 5-20

m


Enlarged view of the catalyst layer: The importance of ionic conductivity of the catalyst layer

~ 0.1 2 m

Fig. 1.9 The interaction between the electrolyte and electrode in a low-temperature fuel

cell (e.g. PEM) [modified after Larminie and Dicks, 2001].

Gas O2 or H2

~ 2-10 nm


1.6 The Specific Power and Energy Relationship: The Ragone Plot

battery)(or cell fuel theof weight Total

Power ElectricSP Power, Specific ; [W kg1] (1.41)

Electric Energy tIEcell ; [W h] (1.42)

;battery)(or cell fuel theof weight Total

Energy ElectricSE Energy, Specific [Wh kg1] (1.43)

Rule of thumb: a 50% increase of specific energy for an electrochemical power source device provides approximately a 74% increase in driving range for a vehicle.

Figure 1.7 Ragone plot: Specific power vs. specific energy for various electrochemical

power sources.

Specific energy (Wh/kg)

Sp

ec

ific

Po

wer

(W/k

g)

Typical transportation PEM fuel cell

battery

AgO - Zn

Sodium sulfur

battery


battery discharge reaction; battery charge reaction

Thermal battery (molten salt):

4LiAl + FeS2 2Li2S + 4Al + Fe; T: 400-700

0C,

Electrolyte: LiCl-Li2SO4-KCl (molten salt);

Ecell (i.e. operating cell potential range) = 2 1.6 V

Sodium-sulfur battery:

2Na(l) + 4S(l) Na2S4; T: 200 - 400

0C,

Electrolyte: sodium beta-alumina (solid porous electrolyte)

Ecell (i.e. operating cell potential range) = 2 1.75 V

Lead-acid battery:

PbO2 + Pb + 2H2SO4 2PbSO4 + 2H2O;

T: between 40 and + 60 0C, Electrolyte: H2SO4 ~ 38% by weight

Ecell = 2 1.7 V

Nickel-Cadmium battery:

2NiOOH + Cd + 2H2O 2Ni(OH)2 + Cd(OH)2;

T: between 40 and + 85 oC, Electrolyte: KOH ~24% LiOH ~5% by weight Ecell = 1.3 0.9 V

Zinc-Silver Oxide battery: (satellites, button cell for photocamera),

2AgO + Zn + H2O Ag2O + Zn(OH)2;

T: between 20 and + 55 oC, Electrolyte: KOH conc. immobilized

Ecell = 1.5 1.2 V.

Primary lithium-thionyl chloride battery: (button cell for electronic devices, military)

4Li + 2SOCl2 4LiCl + SO2 + S;

T: between 20 and + 55 oC, Electrolyte: propylene carbonate-LiAlCl4;

Ecell = 3.6 2.9 V.

Chapter 1_Basic Concepts(4)

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