Full Download Electrochemical Methods: Fundamentals and Applications, 2nd Edition - Ebook PDF Version File PDF All Chapter On 2024
Full Download Electrochemical Methods: Fundamentals and Applications, 2nd Edition - Ebook PDF Version File PDF All Chapter On 2024
Full Download Electrochemical Methods: Fundamentals and Applications, 2nd Edition - Ebook PDF Version File PDF All Chapter On 2024
https://ebookmass.com/product/molecular-diagnostics-fundamentals-
methods-and-clinical-applications-2nd-edition-ebook-pdf/
https://ebookmass.com/product/quantitative-methods-and-socio-
economic-applications-in-gis-2nd-edition-ebook-pdf-version/
https://ebookmass.com/product/fluid-mechanics-fundamentals-and-
applications-ebook-pdf-version/
https://ebookmass.com/product/research-methods-statistics-and-
applications-second-edition-ebook-pdf-version/
Fundamentals of Building Construction: Materials and
Methods 6th Edition – Ebook PDF Version
https://ebookmass.com/product/fundamentals-of-building-
construction-materials-and-methods-6th-edition-ebook-pdf-version/
https://ebookmass.com/product/stage-lighting-second-edition-
ebook-pdf-version-the-fundamentals-2nd-edition-ebook-pdf-version/
https://ebookmass.com/product/heat-and-mass-transfer-
fundamentals-and-applications-5th-edition-ebook-pdf-version/
https://ebookmass.com/product/electrochemical-power-sources-
fundamentals-systems-and-applications-tom-smolinka/
https://ebookmass.com/product/fracture-mechanics-fundamentals-
and-applications-fourth-edition-4th-edition-ebook-pdf-version/
SECOND EDITION
ELECTROCHEMICAL
METHODS
Fundamentals and
Applications
Allen J. Bard
Larry R. Faulkner
Department of Chemistry and Biochemistry
University of Texas at Austin
Section
Symbol Meaning Usual Units References
Section
Symbol Meaning Usual Units References
Section
Symbol Meaning Usual Units References
of species /
k° true standard heterogeneous rate cm/s 13.7.1
constant
xiv Major Symbols
Section
Symbol Meaning Usual Units References
Section
Symbol Meaning Usual Units References
R (a) gas constant Jmol^K"
1
Section
Symbol Meaning Usual Units References
2
*>mt rate of mass transfer to a surface mol cm s ' 1.4.1
1
Wj(A,E) probability density function for species j eV" 3.6.3
w width of a band electrode cm 5.3
Wj work term for reactant j in electron eV 3.6.2
transfer
*c capacitive reactance n 10.1.2
x mole fraction of species j none 13.1.2
>
X distance, often from a planar electrode cm
X\ distance of the IHP from the electrode cm 1.2.3, 13.3.3
surface
x2 distance of the OHP from the electrode cm 1.2.3, 13.3.3
surface
Y admittance rr1 1
10.1.2
Y admittance vector ft" 10.1.2
У distance from an RDE or RRDE cm 9.3.1
z (a) impedance n 10.1.2
(b) dimensionless current parameter in none B.1.6
simulation
z impedance vector ft 10.1.2
faradaic impedance ft 10.1.3
Z\m imaginary part of impedance a 10.1.2
^Re real part of impedance ft 10.1.2
7 Warburg impedance ft 10.1.3
z (a) distance normal to the surface of a cm 5.3
disk electrode or along a cylindrical
electrode
(b) charge magnitude of each ion in a none 13.3.2
z: z electrolyte
Z
j
charge on species j in signed units of none 2.3
electronic charge
GREEK SYMBOLS
Section
Symbol Meaning Usual Units References
Section
Symbol Meaning Usual Units References
2
surface excess of species j at saturation mol/cm 13.5.2
(a) surface tension dyne/cm
(b) dimensionless parameter used to define none 5.4.2, 5.5.2
frequency (time) regimes in step
experiments at spherical electrodes
П activity coefficient for species у none 2.1.5
Д ellipsometric parameter none 17.1.2
l/2
8 r0(s/Do) , used to define diffusional none 5.5.2
regimes at a spherical electrode
"diffusion" layer thickness for species у at 1.4.2,9.3.2
an electrode fed by convective transfer
(a) dielectric constant none 13.3.1
(b) optical-frequency dielectric constant none 17.1.2
(c) porosity none 11.6.2
complex optical-frequency dielectric none 17.1.2
constant
molar absorptivity of species у M" 1 cm " 1 17.1.1
]
permittivity of free space m"2 13.3.1
zeta potential mV 9.8.1
overpotential, E — Eeq V 1.3.2,3.4.2
charge-transfer overpotential V 1.3.3, 3.4.6
1
viscosity of fluid у gem' " V = poise 9.2.2
mass-transfer overpotential V 1.3.3, 3.4.6
none 5.4.1
s 1/2 5.8.2
fractional coverage of an interface by none 13.5.2
species у
(a) conductivity of a solution = fl" 1 -i
Section
Symbol Meaning Usual Units References
Section
Symbol Meaning Usual Units References
standard Galvani potential of ion transfer V 6.8
for species j from phase a to phase /3
Фо total potential drop across the solution side mV 13.3.2
of the double layer
ф2 potential at the OHP with respect to bulk V 1.2.3, 13.3.3
solution
X (12/7)1/2£fT1/2/Do/2 none 7.2.2
XU) dimensionless distance of box; in a none B.1.5
simulation
x
(bt) normalized current for a totally irreversible none 6.3.1
system in LSV and CV
x
(at) normalized current for a reversible system in none 6.2.1
LSV and CV
Xf rate constant for permeation of the primary cm/s 14.4.2
reactant into the film at a modified
electrode
Ф (a) ellipsometric parameter none 17.1.2
(b) dimensionless rate parameter in CV none 6.5.2
(a) angular frequency of rotation; s" 1 9.3
2тг X rotation rate
(b) angular frequency of a sinusoidal s" 1 10.1.2
oscillation; 2rrf
STANDARD ABBREVIATIONS
Section
Abbreviation Meaning Reference
Section
Abbreviation Meaning Reference
1
Letters may be subscripted /, q, or r to indicate irreversible, quasi-reversible, or reversible reactions.
Major Symbols xxi
Section
Abbreviation Meaning Reference
QRE quasi-reference electrode 2.1.7
RDE rotating disk electrode 9.3
RDS rate-determining step 3.5
RPP reverse pulse polarography 7.3.4
RPV reverse pulse voltammetry 7.3.4
RRDE rotating ring-disk electrode 9.4.2
SAM self-assembled monolayer 14.2.2
SCE saturated calomel electrode 1.1.1
SECM scanning electrochemical microscopy 16.4
SERS surface enhanced Raman spectroscopy 17.2.2
SHE standard hydrogen electrode = NHE 1.1.1
SHG second harmonic generation 17.1.5
SMDE static mercury drop electrode 7.1.1
SNIFTIRS subtractively normalized interfacial Fourier transform infrared 17.2.1
spectroscopy
SPE solid polymer electrolyte 14.2.6
SPR surface plasmon resonance 17.1.3
SSCE sodium saturated calomel electrode, Hg/Hg2Cl2/NaCl (sat'd)
STM scanning tunneling microscopy 16.2
swv square wave voltammetry 7.3.5
TBABF4 tetra-/2-butylammonium fluoborate
TBAI tetra-ft-butylammonium iodide
TBAP tetra-w-butylammoniumperchlorate
TEAP tetraethylammonium perchlorate
THF tetrahydrofuran
UHV ultrahigh vacuum 17.3
UME ultramicroelectrode 5.3
UPD underpotential deposition 11.2.1
XPS X-ray photoelectron spectrometry 17.3.2
VB valence band 18.2.2
CHAPTER
1
INTRODUCTION
AND OVERVIEW
OF ELECTRODE
PROCESSES
1.1 INTRODUCTION
Electrochemistry is the branch of chemistry concerned with the interrelation of electri-
cal and chemical effects. A large part of this field deals with the study of chemical
changes caused by the passage of an electric current and the production of electrical en-
ergy by chemical reactions. In fact, the field of electrochemistry encompasses a huge
array of different phenomena (e.g., electrophoresis and corrosion), devices (elec-
trochromic displays, electro analytical sensors, batteries, and fuel cells), and technolo-
gies (the electroplating of metals and the large-scale production of aluminum and
chlorine). While the basic principles of electrochemistry discussed in this text apply to
all of these, the main emphasis here is on the application of electrochemical methods to
the study of chemical systems.
Scientists make electrochemical measurements on chemical systems for a variety of
reasons. They may be interested in obtaining thermodynamic data about a reaction. They
may want to generate an unstable intermediate such as a radical ion and study its rate of
decay or its spectroscopic properties. They may seek to analyze a solution for trace
amounts of metal ions or organic species. In these examples, electrochemical methods are
employed as tools in the study of chemical systems in just the way that spectroscopic
methods are frequently applied. There are also investigations in which the electrochemi-
cal properties of the systems themselves are of primary interest, for example, in the design
of a new power source or for the electrosynthesis of some product. Many electrochemical
methods have been devised. Their application requires an understanding of the fundamen-
tal principles of electrode reactions and the electrical properties of electrode-solution in-
terfaces.
In this chapter, the terms and concepts employed in describing electrode reactions
are introduced. In addition, before embarking on a detailed consideration of methods
for studying electrode processes and the rigorous solutions of the mathematical equa-
tions that govern them, we will consider approximate treatments of several different
types of electrode reactions to illustrate their main features. The concepts and treat-
ments described here will be considered in a more complete and rigorous way in later
chapters.
2 • Chapter 1. Introduction and Overview of Electrode Processes
In this notation, a slash represents a phase boundary, and a comma separates two compo-
nents in the same phase. A double slash, not yet used here, represents a phase boundary
whose potential is regarded as a negligible component of the overall cell potential. When
a gaseous phase is involved, it is written adjacent to its corresponding conducting ele-
ment. For example, the cell in Figure 1.1.1ft is written schematically as
The overall chemical reaction taking place in a cell is made up of two independent
half-reactions, which describe the real chemical changes at the two electrodes. Each half-
reaction (and, consequently, the chemical composition of the system near the electrodes)
1.1 Introduction 3
Pt H2
Zn Ag
СГ СГ
j
Excess Excess
AgCI AgCI
(а) (Ь)
Figure l.l.l Typical electrochemical cells, (a) Zn metal and Ag wire covered with AgCI immersed
in a ZnCl2 solution, (b) Pt wire in a stream of H2 and Ag wire covered with AgCI in HC1 solution.
responds to the interfacial potential difference at the corresponding electrode. Most of the
time, one is interested in only one of these reactions, and the electrode at which it occurs
is called the working (or indicator) electrode. To focus on it, one standardizes the other
half of the cell by using an electrode (called a reference electrode) made up of phases
having essentially constant composition.
The internationally accepted primary reference is the standard hydrogen electrode
(SHE), or normal hydrogen electrode (NHE), which has all components at unit activity:
Pt/H2(a - l)/H + (a = 1, aqueous) (1.1.3)
Potentials are often measured and quoted with respect to reference electrodes other than
the NHE, which is not very convenient from an experimental standpoint. A common ref-
erence is the saturated calomel electrode (SCE), which is
Hg/Hg2Cl2/KCl (saturated in water) (1.1.4)
Its potential is 0.242 V vs. NHE. Another is the silver-silver chloride electrode,
Ag/AgCl/KCl (saturated in water) (1.1.5)
with a potential of 0.197 V vs. NHE. It is common to see potentials identified in the litera-
ture as "vs. Ag/AgQ" when this electrode is used.
Since the reference electrode has a constant makeup, its potential is fixed. Therefore,
any changes in the cell are ascribable to the working electrode. We say that we observe or
control the potential of the working electrode with respect to the reference, and that is
equivalent to observing or controlling the energy of the electrons within the working elec-
trode (1, 2). By driving the electrode to more negative potentials (e.g., by connecting a
battery or power supply to the cell with its negative side attached to the working elec-
trode), the energy of the electrons is raised. They can reach a level high enough to transfer
into vacant electronic states on species in the electrolyte. In that case, a flow of electrons
from electrode to solution (a reduction current) occurs (Figure 1.1.2a). Similarly, the en-
ergy of the electrons can be lowered by imposing a more positive potential, and at some
point electrons on solutes in the electrolyte will find a more favorable energy on the elec-
trode and will transfer there. Their flow, from solution to electrode, is an oxidation cur-
rent (Figure 1.1.2b). The critical potentials at which these processes occur are related to
the standard potentials, E°, for the specific chemical substances in the system.
4 Chapter 1. Introduction and Overview of Electrode Processes
Vacant
0 MO
Potential
Energy level
of electrons
0j Occupied
MO
A + e —> A
(a)
Vacant
0 Energy level
MO
of electrons
Potential
0l Occupied
MO
A - e -^ A+
(b)
Figure 1.1.2 Representation of (a) reduction and (b) oxidation process of a species, A, in
solution. The molecular orbitals (MO) of species A shown are the highest occupied MO and the
lowest vacant MO. These correspond in an approximate way to the E°s of the A/A~ and A + /A
couples, respectively. The illustrated system could represent an aromatic hydrocarbon (e.g.,
9,10-diphenylanthracene) in an aprotic solvent (e.g., acetonitrile) at a platinum electrode.
Power
supply
-Ag
Pt
-AgBr
Figure 1.1.3 Schematic diagram of the
electrochemical cell Pt/HBr(l M)/AgBr/Ag attached
to power supply and meters for obtaining a current-
1МНВГ
potential (i-E) curve.
Let us now consider the particular cell in Figure 1.1.3 and discuss in a qualitative
way the current-potential curve that might be obtained with it. In Section 1.4 and in later
chapters, we will be more quantitative. We first might consider simply the potential we
would measure when a high impedance voltmeter (i.e., a voltmeter whose internal resis-
tance is so high that no appreciable current flows through it during a measurement) is
placed across the cell. This is called the open-circuit potential of the cell.1
For some electrochemical cells, like those in Figure 1.1.1, it is possible to calculate
the open-circuit potential from thermodynamic data, that is, from the standard potentials
of the half-reactions involved at both electrodes via the Nernst equation (see Chapter 2).
The key point is that a true equilibrium is established, because a pair of redox forms
linked by a given half-reaction (i.e., a redox couple) is present at each electrode. In Figure
1.1.1/?, for example, we have H + and H 2 at one electrode and Ag and AgCl at the other.2
The cell in Figure 1.1.3 is different, because an overall equilibrium cannot be estab-
lished. At the Ag/AgBr electrode, a couple is present and the half-reaction is
AgBr + e ±± Ag + Br = 0.0713 Vvs. NHE (1.1.6)
Since AgBr and Ag are both solids, their activities are unity. The activity of Br can be
found from the concentration in solution; hence the potential of this electrode (with re-
spect to NHE) could be calculated from the Nernst equation. This electrode is at equilib-
rium. However, we cannot calculate a thermodynamic potential for the Pt/H+,Br~
electrode, because we cannot identify a pair of chemical species coupled by a given half-
reaction. The controlling pair clearly is not the H2,H+ couple, since no H 2 has been intro-
duced into the cell. Similarly, it is not the O 2 ,H 2 O couple, because by leaving O 2 out of
the cell formulation we imply that the solutions in the cell have been deaerated. Thus, the
Pt electrode and the cell as a whole are not at equilibrium, and an equilibrium potential
*In the electrochemical literature, the open-circuit potential is also called the zero-current potential or the rest
potential.
2
When a redox couple is present at each electrode and there are no contributions from liquid junctions (yet to be
discussed), the open-circuit potential is also the equilibrium potential. This is the situation for each cell in
Figure 1.1.1.
Chapter 1. Introduction and Overview of Electrode Processes
does not exist. Even though the open-circuit potential of the cell is not available from
thermodynamic data, we can place it within a potential range, as shown below.
Let us now consider what occurs when a power supply (e.g., a battery) and a mi-
croammeter are connected across the cell, and the potential of the Pt electrode is made
more negative with respect to the Ag/AgBr reference electrode. The first electrode reac-
tion that occurs at the Pt is the reduction of protons,
+
2H + 2 e - * H 2 (1.1.7)
The direction of electron flow is from the electrode to protons in solution, as in Figure
1.12a, so a reduction (cathodic) current flows. In the convention used in this book, ca-
3
thodic currents are taken as positive, and negative potentials are plotted to the right. As
shown in Figure 1.1.4, the onset of current flow occurs when the potential of the Pt elec-
+
trode is near E° for the H /H 2 reaction (0 V vs. NHE or -0.07 V vs. the Ag/AgBr elec-
trode). While this is occurring, the reaction at the Ag/AgBr (which we consider the
reference electrode) is the oxidation of Ag in the presence of Br~ in solution to form
AgBr. The concentration of Br~ in the solution near the electrode surface is not changed
appreciably with respect to the original concentration (1 M), therefore the potential of the
Ag/AgBr electrode will be almost the same as at open circuit. The conservation of charge
requires that the rate of oxidation at the Ag electrode be equal to the rate of reduction at
the Pt electrode.
When the potential of the Pt electrode is made sufficiently positive with respect to the
reference electrode, electrons cross from the solution phase into the electrode, and the ox-
1 : /
Onset of H +
reduction on Pt.
1 I i
\
\y
1
1.5 0 -0.5
L
/ \
1 \
/ Onset of Br"
/ oxidation on Pt
Figure 1.1.4 Schematic current-potential curve for the cell Pt/H + , Br~(l M)/AgBr/Ag, showing
the limiting proton reduction and bromide oxidation processes. The cell potential is given for the Pt
electrode with respect to the Ag electrode, so it is equivalent to £ P t (V vs. AgBr). Since ^Ag/AgBr =
0.07 V vs. NHE, the potential axis could be converted to EPt (V vs. NHE) by adding 0.07 V to each
value of potential.
3
The convention of taking / positive for a cathodic current stems from the early polarograhic studies, where
reduction reactions were usually studied. This convention has continued among many analytical chemists and
electrochemists, even though oxidation reactions are now studied with equal frequency. Other
electrochemists prefer to take an anodic current as positive. When looking over a derivation in the literature
or examining a published i-E curve, it is important to decide, first, which convention is being used (i.e.,
"Which way is up?").
1.1 Introduction 7
idation of Br~ to Br2 (and Br^~) occurs. An oxidation current, or anodic current, flows at
potentials near the E° of the half-reaction,
Br2 + 2 e ^ 2 B r ~ (1.1.8)
which is +1.09 V vs. NHE or +1.02 V vs. Ag/AgBr. While this reaction occurs (right-
to-left) at the Pt electrode, AgBr in the reference electrode is reduced to Ag and Br~ is
liberated into solution. Again, because the composition of the Ag/AgBr/Br~ interface
(i.e., the activities of AgBr, Ag, and Br~) is almost unchanged with the passage of modest
currents, the potential of the reference electrode is essentially constant. Indeed, the essen-
tial characteristic of a reference electrode is that its potential remains practically constant
with the passage of small currents. When a potential is applied between Pt and Ag/AgBr,
nearly all of the potential change occurs at the Pt/solution interface.
The background limits are the potentials where the cathodic and anodic currents start
to flow at a working electrode when it is immersed in a solution containing only an elec-
trolyte added to decrease the solution resistance (a supporting electrolyte). Moving the
potential to more extreme values than the background limits (i.e., more negative than the
limit for H2 evolution or more positive than that for Br2 generation in the example above)
simply causes the current to increase sharply with no additional electrode reactions, be-
cause the reactants are present at high concentrations. This discussion implies that one can
often estimate the background limits of a given electrode-solution interface by consider-
ing the thermodynamics of the system (i.e., the standard potentials of the appropriate half-
reactions). This is frequently, but not always, true, as we shall see in the next example.
From Figure 1.1.4, one can see that the open-circuit potential is not well defined in
the system under discussion. One can say only that the open-circuit potential lies some-
where between the background limits. The value found experimentally will depend
upon trace impurities in the solution (e.g., oxygen) and the previous history of the Pt
electrode.
Let us now consider the same cell, but with the Pt replaced with a mercury electrode:
Hg/H + ,Br-(l M)/AgBr/Ag (1.1.9)
We still cannot calculate an open-circuit potential for the cell, because we cannot define a
redox couple for the Hg electrode. In examining the behavior of this cell with an applied
external potential, we find that the electrode reactions and the observed current-potential
behavior are very different from the earlier case. When the potential of the Hg is made
negative, there is essentially no current flow in the region where thermodynamics predict
that H2 evolution should occur. Indeed, the potential must be brought to considerably
more negative values, as shown in Figure 1.1.5, before this reaction takes place. The ther-
modynamics have not changed, since the equilibrium potential of half-reaction 1.1.7 is in-
dependent of the metal electrode (see Section 2.2.4). However, when mercury serves as
the locale for the hydrogen evolution reaction, the rate (characterized by a heterogeneous
rate constant) is much lower than at Pt. Under these circumstances, the reaction does not
occur at values one would predict from thermodynamics. Instead considerably higher
electron energies (more negative potentials) must be applied to make the reaction occur at
a measurable rate. The rate constant for a heterogeneous electron-transfer reaction is a
function of applied potential, unlike one for a homogeneous reaction, which is a constant
at a given temperature. The additional potential (beyond the thermodynamic requirement)
needed to drive a reaction at a certain rate is called the overpotential. Thus, it is said that
mercury shows "a high overpotential for the hydrogen evolution reaction."
When the mercury is brought to more positive values, the anodic reaction and the po-
tential for current flow also differ from those observed when Pt is used as the electrode.
8 • Chapter 1. Introduction and Overview of Electrode Processes
Onset of H +
reduction ,
Anodic
Figure 1.1.5 Schematic current-potential curve for the Hg electrode in the cell Hg/H + , Br (1
M)/AgBr/Ag, showing the limiting processes: proton reduction with a large negative overpotential
and mercury oxidation. The potential axis is defined through the process outlined in the caption to
Figure 1.1.4.
With Hg, the anodic background limit occurs when Hg is oxidized to Hg2Br2 at a poten-
tial near 0.14 V vs. NHE (0.07 V vs. Ag/AgBr), characteristic of the half-reaction
Hg 2 Br 2 + 2e«±2Hg 2Br" (1.1.10)
In general, the background limits depend upon both the electrode material and the solu-
tion employed in the electrochemical cell.
Finally let us consider the same cell with the addition of a small amount of Cd 2 + to
the solution,
Hg/H + ,Br"(l M), Cd 2+ (10" 3 M)/AgBr/Ag (1.1.11)
The qualitative current-potential curve for this cell is shown in Figure 1.1.6. Note the
appearance of the reduction wave at about -0.4 V vs. NHE arising from the reduction
reaction
CdBr|~ + 2e S Cd(Hg) + 4Br~ (1.1.12)
where Cd(Hg) denotes cadmium amalgam. The shape and size of such waves will be cov-
ered in Section 1.4.2. If Cd 2 + were added to the cell in Figure 1.1.3 and a current-poten-
tial curve taken, it would resemble that in Figure 1.1.4, found in the absence of Cd 2 + . At a
Pt electrode, proton reduction occurs at less positive potentials than are required for the
reduction of Cd(II), so the cathodic background limit occurs in 1 M HBr before the cad-
mium reduction wave can be seen.
In general, when the potential of an electrode is moved from its open-circuit value to-
ward more negative potentials, the substance that will be reduced first (assuming all possi-
ble electrode reactions are rapid) is the oxidant in the couple with the least negative (or
most positive) E®. For example, for a platinum electrode immersed in an aqueous solution
containing 0.01 M each of F e 3 + , Sn 4 + , and N i 2 + in 1 M HC1, the first substance reduced
will be F e 3 + , since the E° of this couple is most positive (Figure 1.1.7a). When the poten-
Another random document with
no related content on Scribd:
The Project Gutenberg eBook of Perseus; or, of
dragons
This ebook is for the use of anyone anywhere in the United
States and most other parts of the world at no cost and with
almost no restrictions whatsoever. You may copy it, give it away
or re-use it under the terms of the Project Gutenberg License
included with this ebook or online at www.gutenberg.org. If you
are not located in the United States, you will have to check the
laws of the country where you are located before using this
eBook.
Language: English
WIRELESS POSSIBILITIES
By Prof. A. M. Low. With four Diagrams
PERSEUS, or Of Dragons
By H. F. Scott Stokes, M.A.
BY
New York
E. P. DUTTON & COMPANY
681 Fifth Avenue
Copyright 1925
By E. P. Dutton & Company
OF DRAGONS
OF DRAGONS
PREFACE
I must first ask your pardon for troubling you with this digression,
but the facts and the theories are so many and curious and the
whole subject is of such vital importance, that I dared not trust to my
memory alone. It must be admitted that most of the facts recorded
are obviously untrue and that most of the theories are unsupported
by evidence and highly improbable; though the facts are guaranteed
by the highest authorities of the most ancient religions and the
theories upheld by the most eminent modern scientists. Whence it is
evident that I, like Sir John Mandeville, “have taken pains to
ascertain the exact truth”; and yet I think that from this jumble of
superstitions and fables and conjectures and absurdities there does
emerge something that may repay you for the weariness of an hour,
and throw some light upon the hopes and fears with which the
unconquerable spirit of man has progressed through the ages to the
crowning triumphs of this twentieth century. Indeed, as Sir Thomas
Browne finely says in his Preface: “A work of this nature is not to be
performed upon one legg: and should smell of oyl, if duly and
deservedly handled.” And, on second thoughts, I fully and
unreservedly withdraw my apology.
I shall begin with some attempt to define our subject, and then
take you through the dragons of classical antiquity and early
Christendom down to the twilight uncertainty of the sixteenth and
seventeenth centuries and the blank incredulity of our modern age. I
shall then go back to the earliest legends of all—the Egyptian, and
close with an attempt to harmonize the whole, and with a critical
estimate of the place of the dragon in human thought and
experience.
I
OF DRAGONS IN GENERAL
In Glastonbury, where St. Dunstan took the devil by the nose and
where the Pilgrim’s Inn is dedicated to St. George and the Dragon,
the dragon will always be an object of peculiar interest, not to say
veneration. (So true is this that on the 16th October, 1906, the
Somerset County Council, on the advice of its Chairman, adopted as
its sole device “Gules, a dragon rampant, or,” though the recognition
of its increasing importance has since led it to add—15th October,
1912—the mace of office, “at a cost not exceeding £20.”)
For Milton writes in one of his most harmonious numbers:
and St. John (Revelation, xx, 2) speaks of “the Dragon, the old
Serpent, which is the Devil and Satan,” while the Serpent that
tempted Eve in Paradise has been familiar to us all since our earliest
childhood; though commentators differ as to whether it appeared
with a virgin’s head (as some say) and how it was enabled to speak,
and in Eve’s own language; and why the event excited no surprise in
her. (Milton tells us—Paradise Lost, ix, 550, and what follows—that it
did, and that “not unamazed” she took up the matter with the
Serpent, which explained that it had been elevated above all the
“other beasts that graze” by tasting of the tree of knowledge. This
answer at once satisfied Eve and lured her on to her fall. The whole
account is circumstantial but undocumented.) Some say that Eve
was inexperienced with animals, not having been present when
Adam named them; Eugubinus suggests that the Serpent was a