22
The Electrolytic Behaviour of Thin Films. Part I. Hydrogen Author(s): F. P. Bowden and E. K. Rideal Source: Proceedings of the Royal Society of London. Series A, Containing Papers of a Mathematical and Physical Character, Vol. 120, No. 784 (Aug. 1, 1928), pp. 59-79 Published by: The Royal Society Stable URL: http://www.jstor.org/stable/95030 . Accessed: 07/05/2014 15:29 Your use of the JSTOR archive indicates your acceptance of the Terms & Conditions of Use, available at . http://www.jstor.org/page/info/about/policies/terms.jsp . JSTOR is a not-for-profit service that helps scholars, researchers, and students discover, use, and build upon a wide range of content in a trusted digital archive. We use information technology and tools to increase productivity and facilitate new forms of scholarship. For more information about JSTOR, please contact [email protected]. . The Royal Society is collaborating with JSTOR to digitize, preserve and extend access to Proceedings of the Royal Society of London. Series A, Containing Papers of a Mathematical and Physical Character. http://www.jstor.org This content downloaded from 169.229.32.136 on Wed, 7 May 2014 15:29:51 PM All use subject to JSTOR Terms and Conditions

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Page 1: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

The Electrolytic Behaviour of Thin Films. Part I. HydrogenAuthor(s): F. P. Bowden and E. K. RidealSource: Proceedings of the Royal Society of London. Series A, Containing Papers of aMathematical and Physical Character, Vol. 120, No. 784 (Aug. 1, 1928), pp. 59-79Published by: The Royal SocietyStable URL: http://www.jstor.org/stable/95030 .

Accessed: 07/05/2014 15:29

Your use of the JSTOR archive indicates your acceptance of the Terms & Conditions of Use, available at .http://www.jstor.org/page/info/about/policies/terms.jsp

.JSTOR is a not-for-profit service that helps scholars, researchers, and students discover, use, and build upon a wide range ofcontent in a trusted digital archive. We use information technology and tools to increase productivity and facilitate new formsof scholarship. For more information about JSTOR, please contact [email protected].

.

The Royal Society is collaborating with JSTOR to digitize, preserve and extend access to Proceedings of theRoyal Society of London. Series A, Containing Papers of a Mathematical and Physical Character.

http://www.jstor.org

This content downloaded from 169.229.32.136 on Wed, 7 May 2014 15:29:51 PMAll use subject to JSTOR Terms and Conditions

Page 2: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

59

The Electrolytic Behaviour of Thin Films. Part I.-Hydrogen.

By F. P. BOWDEN and E. K. RIDEAL.

(Communicated by T. M. Lowry, F.R.S.-Received May 9, 1928.)

Introduction.

Few investigations have been made on the electrolytic behaviour of very thin metallic films. Whilst the work of Oberbeck* and of Pringt reveals the fact that a deposited layer of metal but a few atoms thick will produce an electrode possessing all the electromotive properties of the massive metal, yet information is lacking on the alteration of the electromotive force as these

layers are built up. It might be anticipated that the behaviour of the electrode

during the deposition of the first few layers would lead to interesting results,

giving some insight into the mechanism of electrode processes and the range of action of forces of adhesion. In view of this it was considered a matter of some interest to investigate how far it might be possible to obtain data on the

:electrode potential and the rate of solution of the deposited atoms during the

building up of the first atomic layer. The problem of metal ion deposition from aqueous solutions is complicated

by the presence of other ions, such as the hydrogen ion, which can deposit :simultaneously with the metal and affect the potential. For this reason the

deposition of the hydrogen ion was first studied. It is well known that, in

general, in order to bring about the continuous deposition of hydrogen ions at a metallic cathode, the potential must be maintained at a value considerably more negative than that of a reversible hydrogen electrode in the same electro-

lyte. The view most generally accepted is that this overpotential is due to an accumulation of electromotively active material on the electrode, and it has been suggested by various workers that it may consist of metallic hydrides, hydrogen atoms or negative hydrogen ions.: With the exception of a paper by Knobel,? little work has been done in determining the actual quantity of

hydrogen accumulated on the cathode during the establishment of over-

potential. Knobel, making the assumptions that the material was atomic

hydrogen, that the relation between the solution pressure of the hydrogen P and the surface concentration of atoms C. is given by the relation P = kC,

* Wied. Ann.,' vol. 31, p. 337 (1887). t ' Z. Elektrochem.,' vol. 19, p. 255 (1913). : 'H eyrovsky, 'Trans. Faraday Soc.,' vol. 19, p. 785 (1924). ? ' J. Amer. Chem. Soc.,' vol. 46, p. 2613 (1924).

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Page 3: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

F. P. Bowden and E. K. Rideal.

and that the potential is related to the pressure by the Nernst expression, calculated this quantity from the rate of growth of overpotential and found that for most metals it was considerably less than an atomic layer.

It is generally assumed that the relation between the potential and the

concentration of active hydrogen is given by the Nernst expression

--E RT

log fCH-

where C.a is the concentration of active hydrogen on the cathode surface and

fC1;+ is the activity of hydrogen ions in the solution. This assumption is open to criticism, for there appears little doubt that the electromotive behaviour

of a substance in the form of a thin film at an interface differs in many respects from its behaviour in the bulk phase. The work of Guyot* and of Frumkint on the change of potential at an air-liquid interface on the addition of a uni-

molecular film to the surface has shown that, over a considerable range, the

alteration in potential is to a first approximation directly proportional to the

surface concentration of the film forming material. Again, Langmuirj has

shown that for the emission of electrons from a tungsten surface coated with a

single layer of caesium atoms, the logarithm of the saturation current is directly

proportional to the surface concentration of caesium atoms. In this investi-

gation it is found that the relation between the electrode potential and the

surface concentration of active material is a linear one, viz.,

- E = F'r + const.,

where E is the electrode potential and r is the true surface concentration of

active hydrogen on the cathode surface. Moreover, the constant P is the same

for all metals; thus the overpotential depends only on the surface concentra-

tion of the added hydrogen and is independent of the nature of the underlying metal. This quantity is very small indeed, the deposition of sufficient hydrogen to form only 1/3000th of an atomic layer raising the potential of the cathode

100 millivolts. Also it is shown that the rate of decay, - dl/dt, of the active

material is not proportional to r2 as usually assumed, nor to r as supposed

by Heyrovsky,? but is an exponential function of the potential, viz.,

dF dt -- kle-k2E

* ' Ann. d. Physique,' vol. 11, p. 506 (1924). t ' Z. Phys. Chem.,' vol. 109, p. 34 (1924). + ' Science,' vol. 57, p. 57 (1923). ? 'Trans. Faraday Soc.,' vol. 18, p. 785 (1924); 'Ree. Trav. Chim.,' vol. 46, p. 582

(1927).

60

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Electrolytic Behaviour of Thin Films. 61

and thus also of the surface concentration

dr - ke-k3r. dtP

This relation holds for all the metals investigated. It is shown that the behaviour of the electrode potential and the magnitude of the quantities involved is compatible with the assumption that the potential of the electrode is due to the presence of electric doublets on its surface, the electric moment of these doublets being equal to that given by a proton and a negative hydrogen ion separated from each other by a distance equal to the diameter of a

hydrogen atom.

Apparatus.

The experimental method involved the measurement of the rapidly changing electrode potential during the initial deposition of active material on a cathode and during its subsequent decay. The usual commutator method was rejected as unsuitable since it measures the average of a large number of cycles and not the initial growth and decay; a cathode-ray oscillograph was unsatisfactory for the same reason. An oscillograph of the Duddell type could not be

employed since no appreciable current must be drawn from the test electrode. The instrument finally adopted and which gave satisfactory results was an Einthoven string galvanometer supplied by the Cambridge Instrument Co. The sensitivity could be varied by altering the tension on the fibre, and as

genezally used was greater than 100 mm. per micro-ampere. The period could be decreased to 1/300th of a second. Electrode potentials were usually measured to the nearest millivolt. The general arrangement of the apparatus is shown in fig. 1.

Light from a 500-c.p. pointolite lamp A was focussed by a large condensing lens B on to the galvanometer, where it was further condensed by C on to the fibre D. The shadow of the fibre was magnified by the microscope E and fell on the camera H ; with the camera at a distance of 1 metre the magnification effected was 600. The cylindrical lens F increased the intensity of illumination

by condensing the light into a narrow band without affecting the horizontal

magnification of the fibre. Another cylindrical lens I on the camera focussed this on to the sensitive paper. This lens had gradation marks scratched on it so that lines 1 mm. apart were reproduced on the photographic paper. The

camera, as constructed, consisted of a light tight box and within this another

compartment containing a 50-yard roll, J, of rapid bromide paper 6 cm. wide. The paper was fed through a spring slit K which maintained a constant tension

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Page 5: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

62 F. P. Bowden and E. K. Rideal.

on it, over the roller L on to the driving roller M. This roller was directly

coupled to a clockwork (not shown) the speed of which could be varied by

H

A JD

100

+-o 0

H2 3

ELECTROL,YTE 9 FIG. 1.

screwing fans of different size on to the fly wheel so tha paper speeds ying

screwing fans of different size on to the fly wheel so that paper speeds varying from 0 5 to 60 cm. per second could be obtained. The timing was given by an electrically driven tuning fork N of period 1/64 second vibrating in front

of the lens I and throwing a shadow on the paper. At low paper speeds a

rotating disc driven by a Wilberforce synchronous motor was employed instead. The galvanometer and the camera were mounted on concrete columns

to diminish mechanical vibration. The circuit for exciting the magnetic field

of the galvanometer is not shown, but the current was supplied from a

constant e.m.f. of 220 volts. The potential of the test electrode P was measured against a saturated calomel

electrode Q, a resistance of 106 Q being included in the circuit so that the current

drawn from the electrode was always very small. The galvanometer was

calibrated by a potentiometer. The polarising current for the cell was supplied

by a 50-volt battery of small accumulators with a variable resistance of several

million ohms in the circuit. The larger currents were measured by a milli-

ammeter R and the smaller by a uni-pivot micro-ammeter S (1 div. = 5 X 10-8

amps.). The apparatus was insulated by blocks of paraffin wax which rested

on an equi-potential metal shield.

Preliminary Experiments.

It is necessary to consider the factors which determine the potential of a

metal plate in a solution containing none or very few of its ions. For example,

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Page 6: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

Electrolytic Behaviour of Thin Films.

if a silver plate be immersed in N/5 . H2S04, the potential being defined by the

expression -E - log PAg

fcAg+

should be infinitely negative at the moment of immersion, since fCAg+ = 0. There will, however, be rapid'attainment of equilibrium between the electrode and all the ions in the solution, so that hydrogen ions will be deposited and the

potential should remain constant at a value which is governed by the solution

pressure of hydrogen on the electrode and the activity of the hydrogen ions in solution. Smits* has shown that if a silver or a copper electrode be immersed in pure water in the absence of oxygen or hydrogen, the potential attains the same value as would a hydrogen electrode in the same solution, indicating that the hydrogen pressure on the plate comes into equilibrium with the external

pressure, viz., one atmosphere. In N/5 sulphuric acid, the hydrogen electrode

potential on the saturated calomel scale is ca. - 0 3 V.; if at this potential the silver is in equilibrium with its ions in the solution, the concentration of the latter required by the above relation is 10-15 N. If for any reason the number of silver ions present in solution is greater than this, the potential will be more positive and will be determined by the value of this silver-ion concentration.

Experimentally it was found that for a silver plate immersed in N/5. H2S04 which had not been freed from oxygen, the potential was + 0 25 V. on the saturated calomel scalet and gradually became more positive on standing. If this potential were due to equilibrium between the silver plate and silver ions in the electrolyte, the necessary concentration of the latter would be 10-5 N. It was considered probable that silver ions were present in the electrolyte in the vicinity of the cathode in sufficient quantity to cause this potential, their solution being brought about by the oxygen dissolved in the electrolyte. To test this point N/5 . 12S04 was freed from oxygen by prolonged boiling of the solution under reduced pressure, and cooling in a hydrogen atmosphere. The silver electrode (area 4 cm.2) was cleaned in dilute nitric acid, washed, made cathodic in dilute sulphuric acid to reduce any oxide, removed, rapidly washed in distilled water, placed in a hydrogen atmosphere and oxygen-free electrolyte then admitted. The potential was now constant at a considerably

*' Rec. Trav. Chim.,' vol. 44, p. 638 (1925). t Unless otherwise stated, all potentials are given on the saturated calomel scale (hydrogen

electrode = - 0- 25 V. on sat. cal.), and the cell is at room temperature ca. 15? C. Nega- tive potentials are spoken of as " high," positive as " low."

63

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Page 7: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

F. P. Bowden and E. K. Rideal.

more negative value, viz., + 0 03. This would correspond to a silver-ion concentration of less than 10-9 N at the electrode. It is probable that silver ions could still be present in this amount due to a trace of oxide on the electrode or to residual dissolved oxygen. If this were so, passage of a small current should deposit this silver and cause the potential to rise, i.e., become more

negative. A current density of 25 X 10-5 amps. cm.2 was passed and the poten- tial rose in 0 35 second to - 0-76 V. This corresponds approximately to

9 X 105 coulombs per cm.2, or enough to deposit ca. 9 X 10-10 gm. ions of

silver. On opening the circuit the potential did not fall rapidly to that of the

hydrogen electrode, as might be expected at first sight, but fell very slowly,

taking 25 minutes to fall from - 0 76 to 0. This behaviour is very different from that observed in solutions from which

oxygen has not been rigorously excluded. In fig. 2, curve I shows the growth

ua 0

o-l.o ui -J LJ .J w - 'I E .': .--- oo 0 J TO -0'69V IN 30 SEC.

5-6 4

-2

F o z 0

+.2 - 5 l.O 1r5 2.0 215 3-0 3.5 4-'o 4'5

TIMlE IN SECONDS

FIG. 2.

of potential at the silver cathode in oxygen free N/5 sulphuric acid on passing a current density of 25 X 10-5 amps. cm.2. Curve II shows the growth when the electrolyte has not been freed from oxygen, all other experimental con- ditions being identical. In fig. 3 is shown the decay of potential on open circuit, curve I for oxygen-free electrolyte, curve II for ordinary electrolyte.

64

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Page 8: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

Electrolytic Behaviour of Thin Films.

The marked effect which even a small amount of oxygen has on the apparent rate of growth and decay of hydrogen overpotential should be emphasised,

Li

LU CI RC IT OPE NED

u4

0r . - r

oLT I T., 20 NOrINUTES

Vu - -..

. 6 ^ ^ - '

.

t3 _ .4 . .7t k ' " ' -" .

Z C

-2

O2 c

,0 ; o5

:.o . . . . . t I

F I .

solutions which are open to the air or which contain dissolved oxygen. The

fall of potential is in this case determined by the rate of supply of oxygen to

the cathode and is not a measure of the natural decay of hydrogen over-

potential. In ordinary solutions, photographs similar to those of Newbery*

have been obtained showing a very rapid fall in the back electromotive force,

the initial drop being too rapid to be recorded on the film. If, however, the

solution be freed from oxygen, the decay rate is very much slower and can be

followed from the moment of opening the circut. The bearing of this on the work Oberbeck should also be considered. For example, Oberbek in one5 0

TIME.N .SECON'DS FIz. 3.

of hsince a very considerabced le numbe of silver in copperiments sulphae beensolution carried out in

solutions which are open tohe number of coulombs which contain dissed oxygen. The

*'Roy. Soc. Proc.,' A, vol. 111, p. 182 (1926). t An air solution saturated with hydrogen behaves in an identical manner to an ordinary

fall of potential is in that the observed effects are due to the ratbsee of dissolved oxygen

the cathode and is no a measre of the natural decay of hydrogen over-

potential. In ordinary solutions, photographs similar to those of Newbery*

have been obtained showing a very rapid fall in the back electromotive force,

the initial drop being too rapid to be recorded on the film. If, however, the

solution be freed from oxygen, the decay rate is very much slower and oan be

followed from the moment of opening the circuit.t The bearing of this on the

work of Oberbeck; should also be considered. For example, Oberbeck in one

of his experiments placed a piece of silver in copper sulphate solution, made it

cathodic, and determined the number of coulombs which passed before the

* ' Roy. Soc. Proc.,' A, vol. 111, p. 182 (1926). t An air solution saturated with hydrogen behaves in an identical manner to an ordinary

air solution, showing that the observed effects are due to the absence of dissolved oxygen and not to the presence of hydrogen.

' Ann. der Physik,' vol. 31, p. 337 (1887).

VOL. CXX.-A. F

65

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Page 9: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

F. P. Bowden and E. K. Rideal.

silver cathode rose to the potential of copper. This was taken as a measure of the number of copper atoms deposited. It is apparent that this conclusion is unjustified: the initial potential of the silver plate in the copper sulphate solutions is that corresponding to equilibrium between the silver plate and the silver ions present in the solution, and before the potential can rise to that characteristic of copper these silver ions have to be deposited. If precautions were taken to prevent the occurrence of these silver ions in solution, the characteristic potential of the silver would initially be more negative than that of copper, and copper ions would deposit on the plate, causing the potential to fall without passage of any external current.

Measurement of Quantity oJ Material deposited on Cathode Surface during Establishment of Hydrogen Overpotential.

Two experimental methods were employed. Method (a).-If the test electrode be made cathodic and a constant current

I amps./cm.2 passed for a time T, the amount of material m gm. ions which accumulates on the cathode surface during the time T is given by

m I T -dt IT- dt dJt.

where i denotes the current which is equivalent to the loss of material at each instant due to natural decay. This rate of loss is a function of the potential. E is the electrode potential and R the resistance of the galvanometer circuit.

Experimentally i can readily be determined for each value of E by measure- ment of the current density necessary to maintain the potential at that value. The last term represents the current drawn off through the galvanometer circuit and is in general negligible. It is not included in the second term, since

during the determination of i there is no deflection of the galvanometer. By plotting the observed rate of growth of potential and correcting for the

natural loss during time of growth, the amount of material actually on the cathode can be determined.

A modification of this may be used. Since i falls off very rapidly with E

(see later), by making I large, the last two terms become negligible if the

potential change is measured for low values of E, so that over this range the current passed is a direct measure of the amount of material on the cathode.

Method (b).-In general, if the test cathode be polarised and the circuit then

opened, the potential, in the absence of traces of oxygen, falls quite slowly. If now a current be passed in the reverse direction, the deposited material is

66

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Page 10: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

Electrolytic Behaviour of Thin Films.

removed from the cathode surface by electrolysis and the potential falls rapidly, the current passed being a direct measure of the amount of material removed.

These two methods, (a) that of depositing material on the cathode and

measuring the growth of potential and (b) of forcing the material off the cathode and measuring the fall of potential, have been carried out on a number of metals. It was found that both methods gave the same result if the solution, was free from dissolved oxygen. If, however, any oxygen were present, there was a large discrepancy between the two measurements, method (a) giving a value for the amount of material on the cathode which may be several hundred times greater than that given by method (b).

During the fall of potential on open circuit, arrests are frequently observed in the decay curve. These have been attributed to the formation of metallic

hydrides,* but the work of one of ust has shown that they are in realitydue to

very small quantities of negative metal impurities which deposit on the cathode

during electrolysis. Even in carefully purified solutions, if the cathode be

'subjected to prolonged electrolysis, this deposition can occur and affect the

potential. Arrests in the decay curve due to this cause were occasionally observed in this work. The behaviour of the cathode under these conditions was erratic and methods (a) and (b) gave different values. In general, the excellent reproducibility of the results under widely varying experimental -conditions and the agreement between methods (a) and (b) indibate that both oxygen and negative metal impurities are absent and that the growth of the

potential of the electrode and its back electromotive force on open circuit is actually caused by the deposition of hydrogen.

Experimental. The cell (see fig. 1) consisted of two glass bulbs each of 200 c.c. capacity

connected at the base by a U-tube which could be closed by a tap. In the top -orifice of each bulb was fitted a stopper, one carrying a platinum anode and inlet for hydrogen and the other the test cathode, a glass tube leading from the immediate vicinity of the cathode to the calomel electrode, and an outlet for

hydrogen. The cell could be evacuated and filled with hydrogen and the oxygen-free

electrolyte blown in under hydrogen pressure through another inlet in the base of the cell, without coming into contact with air. Hydrogen from a

,cylinder was purified by passing through caustic potash and a long train of * Newbery, 'Roy. Soc. Proc.,' A, vol. 111, p. 182 (1926). t Bowden, 'Trans. Faraday Soc.,' vol. 23, p. 571 (1927).

67

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Page 11: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

F. P. Bowden and E. K. Rideal.

heated palladinised copper to remove oxygen. The water was purified by redistillation in a silica still and the sulphuric acid by redistillation in pyrex glass. The glassware was cleaned with chromic acid and distilled water before use.

Mercury. The mercury electrode was contained in a glass cup and the procedure was

either to run the purified mercury (withdrawn from the interior of a mass of

mercury so as to be free from oxide) into the cell in an atmosphere of hydrogen and then admit the electrolyte, or to allow the liquid metal to flow in under the

electrolyte. In either case the initial potential of the electrode depended on the freedom

or otherwise from traces of oxygen. It was generally about 0.0 volt but has been obtained as negative as - 015. Under similar conditions in ordinary solutions not freed from oxygen it is + 0 * 3 volt. The passage of a very small

current is sufficient to deposit this mercury and establish a high hydrogen overpotential, e.g., with a current density of 5 x 10-7 amps. cm.2 the potential is maintained at - 08 volt and on open circuit it falls only very slowly from this value.

Method (a).-In fig. 4 is shown the observed rise of potential on now passing

I . , . a.. . .. =.. .. . . ... .

O 0O04 O'OS 0o12 0-16 0-20 0-24 - 0-28 0

TIME IN SECONDS

FG. 4.

68

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Page 12: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

Electrolytic Behaviour of Thin Films.

a current density of 4 X 10-5 amps. cm.2 The dotted line shows this cor rected for loss during the period of growth. The current density/potential curves used for this correction will be considered later.

Method (b).-Fig. 5 shows the behaviour of the potential on opening the

polarising circuit and passing the current in the reverse direction, making the

0-20 0'25 0-30 0-35 TIME IN SECONDS

Fia. 5.

test electrode anodic: the current in this case is 4 5 X 10-5 amps. cm.2. At the

point A the circuit is opened and the slow fall in potential is due to the natural

decay of the material from the surface, at B the circuit is closed in the reverse direction so that the material is removed electrolytically from the surface.

Both methods show clearly that if rF be the amount of material added to the surface and AE the increase in potential, then

where K is constant, or - AE/AP = K,

- E = KrP + const.,

where P is the amount of active material on the cathode surface at the potential E. Thus the relation between the increase of potential and the amount of hydrogen added is a linear one and not logarithmic as might be expected if the deposited hydrogen behaved as an amalgam electrode of finite bulk

69

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Page 13: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

F. P. Bowden and E. K. Rideal.

concentration. Moreover, the amount of material deposited per square centimetre for every 100 millivolt rise in potential is very small indeed. Method (a) gives 6.7 x10-7 coulombs/cm.2/100 m.v. Method (b) gives 5.9 X 10-5 coulombs/cm.2/100 m.v., or approximately 6 X 1-12gm. ions. Taking the diameter of the hydrogen atom as 1 X 10-I cm.2, the number of coulombs

required to deposit a monatomic layer is ca. 1 7 x 10-3 coulombs/cm.2. If, then, the deposited material is atomic hydrogen, the potential is raised 100

m.v. for every 1/3000th of an atomic layer added. If the hydrogen is regarded as being adsorbed by the metal atoms, then (taking the diameter of a metal atom as 3 X 10-8 cm.) approximately 1/300th of the metal atoms in the

surface is covered for a potential rise of 100 m.v. This linear relation is found

to hold accurately over a potential range of - 0 2 volt to - 10 volt, and there

is no evidence of deviation at more negative potentials. In the following table are given the results of a number of determinations of this quantity

using both methods (a) and (b) under widely varying conditions of cathode

area, current density, time of electrolysis and strength of solution. Con-

sidering the magnitude of the quantities involved the agreement is good, and within the limits of experimental error is independent of the above factors.

Area of Curt 'r in coulombs/cm.2 cathode Slin amps. et 10' to change

in square . e . potential centimetres. x 100 millivolts.

0.5 N/5. H2SO\ 8 (a) 6-7 05 ,, 8 (b) 5. 0 0.5 8 (b) 6 0-Mean 5.9

0 33 .N/500. H,2SO'4 8 (a) 7.9 i 1 (a) 5.0

-?. b 3 (b) 6-2-Mean 6-3

314 N/500. H2804 24 (b) Few minutes elec- 5-3 ...,, 24 (b) trolysis 5-4 ; ? ,,? ' 54 (b) 3 hours electrolysis 5 2

24 (b) 45 hours electro- 5-3 126 (b)J lysis 6-1--Mean 5.8

3 14 iN/ . H2SO 24-8 (b) . 5.1 3-14 , 24-8 (b) 5.2

14-4 ,, 18 (a) 7.1 . ,, 580 (a) 6-7 -,, 640 (b) 5.*9

- : ,, 300 (b) 6-9-Mean 6-2

a':;~~~~~~ . . . . Mean value 6.0 X 10-7

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Page 14: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

Electrolytic Behaviour of Thin Films. 71

Silver.

The cathode consisted of a sheet of pure silver, area 4 3 cm.2, and the electro-

lyte was N/5 sulphuric acid. The behaviour was in all respects similar to that of mercury, the same linear relation being found between AF and AE. The

magnitude of the quantities was, however, different and depended on the nature of the silver surface. This is shown in the following table; the values in the last column are the mean of several determinations using different current

densities, and the maximum deviation from this mean is given.

Potential at CD Xr in coulombs/cm. Nature of surface. of 1 X 10-s amps. Method. 10 to hange

cm.2. potential 100 milli- volts.

Volts. (i) Freshlyetched with dilute 0-465. 2 by (a) 310 ?40

HN03. Fine crystalline struc- 4 by (b) ture.

(ii) Same as (i) but 20 hours 0-525 1 by (a) 225?13 old. 2by (b)

(iii) Polished with "flour" sand- 0 60 2 by (a) 100 ?10 paper. Shows very scratched sur- 2 by (b) face under the microscope.

It will be seen that the amount of hydrogen added per apparent square centimetre of an etched silver surface to raise its potential 100 m.v. is about 50 times as great as that found for mercury. We must conclude that either the free energy of hydrogen deposited or adsorbed on silver is very much less

than that of hydrogen on mercury, or else the accessible area of the silver cathode is very much greater than the apparent area.

Amalgamated Silver.

To investigate this point the surface of the sand-papered silver cathode was

amalgamated with a very small amount of mercury, the surface being rubbed with another piece of silver for 30 minutes, giving a bright surface resembling

mercury. The results were as follows:-

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Page 15: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

72 F.P. Bowden and E. K. Rideal.

Potential for apparent Ar in coulombs/cm.2 Nature of surface. CD of 1 x 10-5 amps. x 107 to change

cm.2 potential 100 millivolts.

Amalgam 1 hour old. Surface bright -0 73 7 T-30 .8 resembling mercury.

Amalgam 20 hours old. Surface uni- --1 06 7. 9 0 7 form grey and solid.

Amalgam 150 hours old. Surface -0.9 10-5 40-8 solid. Dull grey and under micro- scope shows depressions and scratches I of original sand papering.

The electrolyte was N/5 . H2SO4 as before, and the values given are the mean of several determinations using both methods (a) and (b) and with varying current densities. Since the amalgam is solid, the number of silver atoms in the surface must be high, yet it is seen that AF is practically the same as that found for pure mercury (6 X 10-7), and only increases slightly as the surface becomes obviously rougher.

This indicates that the quantity of hydrogen which must be deposited on a metal plate to raise its potential 100 milli-volts is not a specific property of the metal atom, but depends only on the physical structure of the surface. The large differences found in the value of A F for etched, for polished and for

amalgamated silver is due to the fact that the accessible area of the etched or polished surface (i.e., the total area of the metal which is accessible to

hydrogen atoms and on which they can deposit or be adsorbed) is very much

greater than its apparent area. If A cm.2 is the accessible area of a metal cathode per apparent square centi-

metre, we may write equation (1), p. 73,

- E = K/A + const. or

- E -= p + const., (2)

where { = IK/A and is a constant, having the same value for all metals. It appears that the magnitude of the hydrogen overpotential depends

only on the true surface concentration of the added hydrogen and is inde-

pendent of the underlying metal. The rate of removal of this added hydrogen (as measured by the current necessary to keep the surface concentration con-

stant) does, however, vary greatly with the nature of the underlying metal. The behaviour of platinum lends strong support to these conclusions.

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Page 16: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

Electrolytic Behaviour of Thin Filrms.

Platinum. The following table shows the behaviour of platinum and platinised sur-

faces. The electrolyte was N/5 . H2SO4, and the values given are the mean of several determinations.

The platinum was deposited on the mercury electrolytically by adding known amounts of a platinum salt to the electrolyte.

It is seen that the accessible area of platinum foil is about twice, while that of platinum black is about 2000 times, its apparent area. If, however, the platinum be deposited not on to solid platinum but on to the liquid mercury, the atoms cannot form a sponge having a large area but lie in a plane surface

Potential for an Ar in coulombs /cm.2 Nature of surface. apparent C.D. of 10-3 X 107 to change

amps. cm.2. potential 100 millivolts.

Bright platinum foil ................................ -060 13 (?1) Platinised platinum " platinum black" -0 31 1]1,000 (?2000)

as used for H2 electrode. Platinised mercury-

(i) Pure mercury ................................ -119 6 (?1) (ii) Mercury surface covered with

about 1/100 of an atomic layer of platinum. -1-09 5.1 (? 0-5)

(iii) Mercury surface covered with several atomic layers of platinum. -0.74 5 0 (5?0.5)

having practically the same accessible area as the liquid mercury and giving the same value for A P. The rate at which such a surface can catalyse the evolution of hydrogen is, however, many thousand times greater than a pure mercury surface.

Relation between Surface Concentration of Hydrogen and its Rate of Decay.

In fig. 6 the logarithm of the apparent current density in amps. per cm.2 10-7 is plotted against the observed electrode potential for the metals investigated. The electrolyte is N/5 sulphuric acid, so that the

potential of the reversible hydrogen electrode is - 0 3 volt on this scale. It is seen that over a wide range of current density the points lie accurately

on a straight line. In fig. 7 the logarithm of the true current density, i.e., the current per unit of accessible area, is plotted against the potential. The curves are marked with the same numbers as in fig. 6.

If E be the electrode potential and i the true current density, then

- E = b log i + const., (1) where b is a constant.

73

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Page 17: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

F P. PBowden and E. K. Rideal.

-1.4 LLI C)

0

Id

-J ?0

~ -.4 . -4? '- --

- 0 *5 i.o .i5 2-0 2-5 3.0 3-5 4.0 4-

LOGARTHM OF (APPARENT CURRENT DENSITY IN AMPS. Cmrn XlOd

FIG. 6.

Curve I.-Mercury. Curve V.-Etched silver, old. Curve II.-Platinised mercury, 1/100 covered. Curve VI.-Polished silver. Curve III.-Platinised mercury, thin film of Pt. Curve VII.-Bright platinum. Curve IV.--Etched silver, new. Curve VIII.-Spongy platinum.

(In Curve VIII the current density is expressed in amps. X 10-6 in order to bring it on the figure. It should be displaced two divisions to the right.)

Also it has been shown on p. 72 that

- E U= pr + const., (2)

where p is a constant and where r is the true surface concentration of active

hydrogen at the potential E.

'Combining (1) and (2) af- -1 n *- I-s *I /4?\ o log = p I' +- consr.

But i is a direct measure of the velocity of the reaction by which the active

hydrogen disappears from the plate or - ddt = i. Therefore equation (3)

becomes log (- dr/dt) rFib + const.

or

-drldt = ke^,

74

to

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Page 18: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

Electrolytic Behaviour of Thin Films. 75

'4 0 0

L)

-J W -j uJ

0 -j

Q v- (4

(j:

I-

z

-j

2

0' 0.

42 ~ ~~~~ J

oi . ,..!?

;s I E\d

-r,

.?o ,.. I....

t-0~~~~~~~~~~~~~~~~~~~t .4 --o- <

iii__

-5 -4- -3 -2 -. 0 1 2 3 4 LOqARITHM OF(TRUE CURKRENT DE-NSITY IN AMPS CM2- X 0-7)

FIG. 7.,

so that the rate of removal of the active hydrogen is an exponential function of its surface concentration.

The logarithmic relation between the current i and the potential holds

accurately for all metals over a wide range of current density. In the case of

high overvoltage metals such as mercury, deviation from this only occurs at

very low current densities (less than 5 X 107 amps. cm.2), and it is probable that this break is due to a trace of dissolved oxygen still present in the electro-

lyte, its rate of diffusion up to the cathode at this point becoming comparable with the current density, so causing the potential to fall below its equilibrium value. The supply of oxygen necessary for this is less than 5 X 10-12 gm. ions per second, and if any oxygen be admitted, the break occurs at much

higher current densities. With low overvoltage metals such as platinum,

hydrogen can be evolved at an appreciable rate without any overpotential, and it is only when this is exceeded that overpotential occurs.

In fig. 7 the graphs for mercury, polished silver and smooth platinum are

extrapolated until they intersect a line drawn through the reversible hydrogen

potential. The points of intersection are a measure of the maximum rate at

which hydrogen can deposit reversibly on the respective metals.i On mercury

.

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Page 19: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

F. P. Bowden and E. K. Rideal.

it is approximately 3 X 10-12 coulombs/sec., on polished silver 2 x 10-9 and on smooth platinum 2 X 10-6. If for each metal the current be reduced below this value, no further decrease will occur in the cathode potential. If this maximum reversible current be denoted by i0, then the overpotential - is given by

- = b (log z - log io), where b is a constant.

This may be written T = a + B log i,

where a = - b log i,.

This last relation between overpotential and the current density was first noted by Tafel* and has been confirmed by other workers.t These investiga- tors found it to hold accurately for mercury but only approximately for other

metals. In fig. 7 the value of b is practically the same for mercury, for silver, and for smooth platinum, viz., 0-120. The value obtained by Tafel for

mercury was 0 110. Since i represents the rate at which hydrogen is removed from the cathode

surface per unit of accessible area, it is a measure of the specific catalytic activity of the surface for this reaction. It is seen that the curves in fig. 7 for the pure metals fall into three groups characteristic of mercury, silver and

platinum, indicating the specfic activity of a particular atom; for example,

platinum is not very different whether it exists in the form of platinum black or bright platinum, the apparent increase in activity being due mainly to the increase in the accessible area. For the respective metals, however, it differs

very greatly, the rate at which platinum catalyses the deposition of hydrogen being 106 times greater than mercury per unit of accessible area of the surface, the surface concentration of hydrogen being the same in each case.

It is apparent that this experimental procedure gives an interesting method

of investigating the areas of metallic surfaces and correlating it with their

catalytic activity, and this will be considered in more detail later.

Mechanism of the Electrode Potential.

If the potential across the interface be due to a number of electric doublets, each possessing a finite electric moment, then the interfacial potential is given

by - E = 47cnpL, (1)

*' Z. Phys. Chem.,' vol. 50, p. 641 (1905). t Lewis and Jackson, 'Z. Phys. Chem.,' vol. 56, p. 193 (1906); Harker and Adams,

' J. Phys. Chem.,' vol. 29, p. 205 (1925).

76

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Page 20: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

Electrolytic Behaviour of Thin Films. 77

where E is the absolute potential difference, n is the number of doublets per square centimetre, and , is the electric moment contributed by each doublet. If the number of doublets on the surface is increased, then the increase in

potential is given by -- AE = 4An{x. (2)

Experimentally this linear relation is found, the cathode potential becoming 100 millivolts more negative for every 6 X 107 X 6-06 X 1023 - 96540 =

3-7 X 1012 hydrogen ions deposited per square centimetre of accessible surface and independent of the nature of the cathode. Substituting these values in

(2) we obtain

i x 01 X 10-2 - 47 3-7 X 1012 sz, whence P = 7-2 X 10-18 E.S.U.

If the electric moment be produced by a doublet consisting of an electron

separated from a proton by a distance a, we obtain

. ==,3, whence 8 = 7 2 X 10-18 /(4 77 X 10-10) = 1.5 X 10-8 cm.

This value is in agreement with that given for the diameter of the hydrogen atom, namely, 1 X 10- cm.

It is also significant that the rate of discharge of hydrogen is an exponential function of the potential

- d ldt -= k1e-E,

a relationship which is similar to that governing the emission of electrons across the double layer at a metallic surface.

It is found experimentally that if the deposited hydrogen be removed by electrolysis to a value below that corresponding to the reversible hydrogen electrode, the potential continues to fall with exactly the same gradient until well below the reversible potential, and only deviates from this when the number of metallic ions which go into solution becomes appreciable. For example, on polarising a silver electrode and then making it anodic, the linear relation between the amount of material removed and the fall of potential held over a range of - 0 9 to + 0 20; beyond this it fell more slowly and exponentially to the silver value.

We are led to infer that the mechanism of the reversible hydrogen electrode is similar to the case of overpotential just considered, the number of doublets on the surface being here dependent only on the pressure of the gas. Whilst the absolute value of the hydrogen electrode potential is not known, the dropping electrode gives a value of - 0 28 volt, which is equivalent to 1 X 1013 doublets

per square centimetre, thus covering approximately 1 per cent. of the metal

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Page 21: The Electrolytic Behaviour of Thin Films. Part I. Hydrogen

F. P. Bowden and E. K. Rideal.

atoms in the surface. On deposition of hydrogen ions, doublets are formed

which in turn must be discharged to form molecular hydrogen. The rate at

which the discharge of these doublets can occur at the concentration corre-

sponding to the reversible hydrogen potential varies from metal to metal.

On mercury it is 2 X 107, on silver 1 X 1010 and on platinum 1 X 1013 doublets

per second. Unless this rate of discharge be exceeded there is no overpotential. If this rate is exceeded, the surface concentration of the doublets increases and

the electrode potential becomes more negative, the overpotential being directly determined by the number of doublets added to the surface. Since the equili- brium concentration of doublets is now exceeded, changes in the gas pressure will not affect appreciably the electrode potential.*

A possible mechanism by which these doublets are brought into existence

may be suggested. The linkage between the metal and the deposited hydrogen atom is of such a nature that the hydrogen atom (possessing an electron

affinityt) behaves as a negative ion which attracts to it a hydrogen ion

from the solution, thus forming a doublet of a positive and negative hydrogen ion orientated on the surface. Discharge of this doublet involves the loss of

an electron from the metal surface, and the rate at which it occurs is a function

of the potential, viz., - dl/dt = -ke-7"eE, which is the same relation as that

governing the emission of electrons from metals. When the doublet is dis-

charged, the two hydrogen atoms unite to form a neutral hydrogen molecule.4

Heyrovsky? on very different grounds has proposed a mechanism for

hydrogen overpotential which is somewhat similar to the above in that it

involves the formation of negative hydrogen ions on the electrode surface.

It, however, involves the assumptions that the potential is proportional to the

logarithm of the surface concentration of negative hydrogen ions and that the

rate of removal of these ions is directly proportional to their surface concentra-

tion. The above experimental work does not support these assumptions.

* Knobel, 'J. Amer. Chem. Soc.,' vol. 46, p. 2751 (1924); Bircher and Harkins, ibid., vol. 45, p. 2890 (1923).

t Joos and Huttig, 'Z. f. Physik,' vol. 39, p. 473 (1926). : Alternatively the potential change which occurs may be regarded as due simply to a

flux of electricity across the boundary. Taking the dielectric capacity across the interface as unity, the thickness of the double layer is 1 5 A.U., which is approximately the diameter of a hydrogen atom.

? ' Rec. Trav. Chim.,' vol. 44, p. 499 (1925), and vol. 46, p. 582 (1927).

78

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Electrolytic Behaviour of Thin Films.

Summary.

A quantitative investigation is made on the changes of electrode potential at the surfaces of metallic cathodes during the electrolytic deposition and removal of very small quantities of hydrogen. It is found that the electrode

potential is a linear function of the surface concentration of the hydrogen; moreover, the potential of the polarised cathode depends only on the true surface concentration of the added hydrogen and is independent of the nature of the underlying metal. Apparent differences observed are due to differences in the real areas of the cathodes, and this gives a method of measuring the accessible area of metallic surfaces. The amount of hydrogen deposited is

small, sufficient to form 1/3000th of an atomic layer, causing a change of 100 millivolts in the electrode potential. The rate of decay of the electromotively active hydrogen from the cathode surface is an exponential function of the

potential. The general behaviour and the magnitude of the quantities involved can be explained on the assumption that the electrode potential is due to the

presence of electric doublets on its surface, the electric moment of these doublets

being that given by a proton and an electron separated from each other by a distance equal to the diameter of a hydrogen atom.

Out best thanks are due to the Commissioners of the 1851 Exhibition for a

Scholarship granted to one of us (F. P. B.) which made it possible to carry out this work, to Prof. T. M. Lowry, F.R.S., who communicated this paper, and to the Imperial Chemical Industries and the Cambridge Instrument Co. for assistance in obtaining the necessary apparatus.

79

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