Electrochemical Deposition of Nanostructured Metal/Metal ...170636/FULLTEXT01.pdf · Journal of Electroanalytical Chemistry 594, 35 (2006) II. Pulsed galvanostatic and potentiostatic

ACTAUNIVERSITATIS

UPSALIENSISUPPSALA

2007

Digital Comprehensive Summaries of Uppsala Dissertationsfrom the Faculty of Science and Technology 336

Electrochemical Deposition ofNanostructuredMetal/Metal-Oxide Coatings

JONAS ESKHULT

ISSN 1651-6214ISBN 978-91-554-6956-6urn:nbn:se:uu:diva-8186

�� !"� !##$ �� %&## '�� ' �� ' (��) *�� +�� ,��)

��

,�-�� .) !##$) ,�� ' /�� 0��10��23�� 4��)5�� ) �� 667) 8% ��) ��) 9 :/ ;$"2;�288%27;8727)

,�� '�� 2 �� ) *�� <�� '� �� ' �� '��)=+�� ' �� '�� ' �� '�� ) 9 �� +�� + �� '� �� ' ��1��2�� +�� )0��1�� +�� '�� <��

�� ' �� ) *�� ' ��+�-�� +�� +�� <��> �� ?,@40A ��<�� +�� '�� ' 4�

!3)

*�� +�� '�� >�� +�� BC�� B( � ,0� *,0� �� C��) ,�� ' �1 �

!3

6� � ��

�� +�� ' �� )/�� ' 4�14�

!3 +��

�� -�� 4�?99A2�� ) *�� +�� <�� +�� 4� ��4�!D) E��2��'�� ' 4� �� 4�

!3 ��

��<�� +�� -�� ' �� )4�� 2�� ' �

!3

6�+�� -�� ' �� !##

� +�� '�� ''�� ?999A2�� ) F�� +�� 2� ��'� �� 8# �� +�� ' 77# �5�1�)*�� + �� ' ��

�� '�� +�� ) *�� ''��+ �� ' �� )

� �� 4�!3� �!36� �� ,@40� ��

�� !��" � �� #�� " $% &'(" �� " � �)&*+*��" ��

G .�� ,�-�� !##$

9 / �78�27!�%9 :/ ;$"2;�288%27;8727��&�&��&��&��2"�"7 ?��&11��)-�)��1��H��I��&�&��&��&��2"�"7A

Till Mamma och Pappa

”När man är en björn med en myck-et liten hjärna och tänker ut saker, upptäcker man ibland att en idé som verkade vara riktigt idéaktig inne i hjärnan, är helt annorlunda när den kommer ut i det fria och andra människor ser på”

Nalle Puh

List of papers

This thesis is based on the following papers, referred to in the text by their roman numerals.

I. On the origin of the spontaneous potential oscillations ob-served during galvanostatic deposition of layers of Cu and Cu2O in alkaline citrate solutions J. Eskhult, M. Herranen, L. Nyholm Journal of Electroanalytical Chemistry 594, 35 (2006)

II. Pulsed galvanostatic and potentiostatic electrodeposition of Cu and Cu2O nanolayers from alkaline Cu(II)-citrate solu-tions J. Eskhult and L. Nyholm Submitted to Journal of the Electrochemical Society

III. Current oscillations during chronoamperometric and cyclic voltammetric measurements in alkaline Cu(II)-citrate solu-tions J. Eskhult, C. Ulrich, F. Björefors, L. Nyholm Submitted to Electrochimica Acta

IV. Electrodeposited Sb and Sb/Sb2O3 nanoparticle coatings as anode materials for Li-ion batteries H. Bryngelsson, J. Eskhult, L. Nyholm, M. Herranen, O. Alm, K. Edström Chemistry of Materials 19, 1170 (2007)

V. Electrodeposition and electrochemical characterisation of thick and thin coatings of Sb and Sb/Sb2O3 particles for Li-ion battery anodes H. Bryngelsson, J. Eskhult, K. Edström, L. Nyholm Electrochimica Acta, Published on the web 13 February 2007

Reprints were made with permission of the publishers

Comments on my contribution to the papers in this thesis I have planned all the experiments and performed all the experimental work in papers [I-III], except for the ellipsometry measurements in paper [III], which were done together with Christian Ulrich. The data analysis in [I-III] was done together with Leif Nyholm. In [I-III], I also made the literature studies and wrote the first version of the manuscripts.

In papers [IV] and [V], which were done in co-operation with the battery research group at our department, I performed the initial electrochemical deposition of the coatings, which then was reproduced and improved by Hanna Bryngelsson. I performed the XRD and Raman spectroscopy meas-urements in paper [IV] and the EQCM measurements in paper [V]. I also contributed significantly to the data analysis in papers [IV] and [V].

Contents

Introduction.....................................................................................................9

1 Electrochemical deposition..................................................................11 1.1 Oxidation and reduction .............................................................11 1.2 Mass transfer versus kinetic control ...........................................12 1.3 Deposition from solutions containing metal ion complexes ......13 1.4 Deposition of metals...................................................................14 1.5 Deposition of metal oxides.........................................................14 1.6 Deposition of nanostructured materials......................................15

2 Experimental........................................................................................17 2.1 Preparation of electrolytes and experimental setup....................17 2.2 Characterization of the electrodeposits ......................................18

3 Electrochemical deposition from alkaline Cu(II)-citrate solutions......22 3.1 Complex speciation and electrode reactions ..............................22 3.2 Deposition of copper ..................................................................24 3.3 Deposition of Cu2O ....................................................................25 3.4 Electrochemical stability of Cu2O ..............................................27 3.5 Deposition of nano-composites of Cu/Cu2O ..............................29 3.6 Deposition of well-defined layers of Cu and Cu2O....................32

4 Electrochemical deposition from Sb(III)-tartrate solutions .................34 4.1 Complex speciation and electrode reactions ..............................34 4.2 Deposition of nano-composites of Sb and Sb2O3 .......................35 4.3 Electrochemical characterization of the deposited material .......38 4.4 Deposition of thick Sb2O3-containing films ...............................40 4.5 Deposition of pure antimony films.............................................41

Concluding remarks ......................................................................................43

Acknowledgements.......................................................................................45

Summary in Swedish/Sammanfattning på svenska ......................................46

Bibliography .................................................................................................51

List of abbreviations

EQCM Electrochemical quartz crystal microbalance XRD X-ray diffraction SEM Scanning electron microscopy TEM Transmission electron microscopy XPS X-ray photoelectron spectroscopy SCE Saturated calomel electrode RDE Rotating disc electrode RRDE Rotating ring disc electrodeAFM Atomic force microscopy STM Scanning tunneling microscopy ESTM Electrochemical STM SECM Scanning electrochemical microscopy CCD Charged coupled device

9

Introduction

Electrochemical deposition deals with the synthesis of solid films from dis-solved species by alteration of their oxidation states using electricity. Gold plating is an example where dissolved gold ions are reduced by donation of electrons from an external current source, thus forming a solid gold metal film. Not only metals can be prepared by electrochemical deposition but also chemical compounds such as oxides and phosphides. Important applications within the electronics industry are the deposition of copper interconnects in integrated circuits [1] and the deposition of thin film magnetic materials, e.g. CoNiFe alloys [2]. When the electrons are supplied by a chemical reducing agent added to the solution, the process is called electroless deposition. The latter method is often used for deposition of protective coatings such as nickel-phosphorous alloys [3]. Electrochemical deposition has a widespread use in nanotechnology since it can be used to fill three-dimensional features at room temperature with good control of thickness and morphology.

In the present thesis, nanostructured metal/metal-oxide coatings have been deposited under reductive conditions from solutions containing metal complexes. Two different systems have been studied: The alkaline Cu(II)-citrate solution where the intermediate oxidation state Cu(I) is stable as the oxide Cu2O and the Sb(III)-tartrate solution where Sb(III) is stable as Sb2O3 above pH 6. In both systems, electrochemically induced local production of hydroxide ions gives rise to precipitation of oxide at potentials where the metal otherwise is stable. Deposition from the alkaline Cu(II)-citrate solution yields Cu/Cu2O coatings and deposition from the Sb(III)-tartrate solution gives rise to Sb/Sb2O3 coatings, respectively.

The deposition processes has been studied primarily with the electro-chemical quartz crystal microbalance technique. Recordings of the changes of the electrode mass have revealed the presence of deposition of Cu2O dur-ing conditions of open circuit as well as precipitation of Sb2O3. The phase content and microstructure of the deposited materials have been character-ized using spectroscopy and diffraction techniques. Ellipsometry was also used to detect the formation of Cu2O on Cu. The Sb/Sb2O3 deposits were further electrochemically characterized by galvanostatic cycling in a Li+-containing organic electrolyte in co-operation with the battery research group at our department.

The aim of this thesis is to give further insights into the electrochemical deposition of nanostructured materials. In-situ studies of the spontaneous

10

potential oscillations obtained during constant current electrolysis in alkaline Cu(II)-citrate solutions have paved the way for the deposition of well-defined layers of Cu and Cu2O. Local pH changes during the deposition from solutions of metal complexes with organic ligands have been found to be essential for the deposition of Cu2O as well as the co-deposition of Sb2O3 during Sb deposition. This thesis contains a summary of the most important results concerning the deposition of alternating layers of Cu/Cu2O and coat-ings of Sb/Sb2O3 obtained by electrochemical deposition

11

1 Electrochemical deposition

Electrochemical deposition has typically been used for the plating of metals and the art has historically been developed industrially, mainly by experi-ence. Research has deepened the understanding of the deposition processes and electrochemical deposition today provides many exciting routes for the synthesis of metal oxides, alloys and nanostructured materials. Fundamental aspects of electrochemical deposition include the heterogeneous electron-transfer step between the electrode and the electroactive species present in the solution as well as the transition of the discharged metal atoms into the crystalline state [4]. The deposition is inexpensive and fast and the technique finds many important applications in semi-conductor industry, fabrication of energy storage materials and corrosion protection coatings. This thesis deals with both the deposition of oxides and the generation of nanostructured met-als.

1.1 Oxidation and reduction Electrochemical deposition of metals and metal oxides typically proceeds by oxidation or reduction of species in a solution. The standard electrode poten-tial for an electrochemical reaction is the potential where the rate of the re-duction and the oxidation reactions are equal at standard conditions of con-centrations, pressure and temperature. The Nernst equation relates the stan-dard electrode potential E0 to the electrode potential E:

� �� red

oxnFRTEE ln0 �� (1.1)

where R denotes the standard gas constant (8.314510 J·K–1·mol–1), T the absolute temperature in Kelvin, n the number of electrons transferred and F the Faraday’s constant (96485.309 C·mol–1). The potential also depends on the ratio of the natural logarithm of the activities of the oxidized and reduced species.

12

The mass deposited on the working electrode during the electrolysis is proportional to the number of moles of electrons transferred, as is stated by Faraday’s law of electrolysis:

nFQMm � (1.2)

where m is the mass of the deposit produced at the electrode (in grams) and Q is the total electric charge (in coulombs) required for the process.

The electrolysis of species can be performed using a constant current forced trough the electrochemical cell, while the electric potential is moni-tored. Alternatively, a desired potential can be chosen, which is then main-tained by the instrument while the necessary current used to maintain that potential is monitored.

The electrode potential is directly connected to the energy change of the electrode process through the relationship: �G0 = – nFE0. When the elec-trode potential is made more negative in relation to the standard reduction potential for an electrochemical reaction, the reduction current increases because the rate of electron transfer of the reduction increases. In the elec-tron transfer controlled potential region, there is a linear relationship be-tween the potential and the logarithm of the deposition current known as the Tafel linearity. However, the current can also be limited by other factors such as mass transfer, preceding chemical steps and crystallization proc-esses.

1.2 Mass transfer versus kinetic control In an electrochemical cell the current forced through the solution is carried by ions. The consumption of species during the reduction at the working electrode gives rise to a concentration gradient in the solution outside the electrode surface. The random motion of the species by which this concen-tration difference is equalized is called diffusion. In an unstirred solution the diffusion layer thickness extends into the solution until vibration and thermal movement start to contribute to the mass transport.

The mass transport of ions also takes place by migration, which is the movement of charged particles in an electric field. The contribution of mi-gration to the mass transport of electroactive species is usually diminished by the addition of inert ions that carry the current but do not participate in the electrode reactions.

13

The rate of the electrode reactions depends on chemical reactions preced-ing or following the electron transfer, the electron transfer kinetics and sur-face reactions, such as adsorption and crystallization steps [5].

The current is generally limited either by the mass transport of species towards the electrode or by the kinetics of the electrode reaction. If one wants to study the kinetics of an electrode reaction, care must be taken to prevent problems with the mass transport of the species. This can be achieved by forced convection of the solution using the rotating disc elec-trode (RDE) technique. On the other hand, determinations of the diffusion coefficient of species in the solution by electrochemical methods require that the kinetics of the electrode reaction does not limit the current.

When an electrochemical experiment is performed using a constant cur-rent, the potential shifts to the value required to maintain that current. Under conditions of mass transfer control (i.e. when the kinetics of the electrode reaction does not limit the current), the potential shifts when the concentra-tion of electroactive species at the electrode surface is reduced to zero. The time needed for this depletion of the species, the transition time �, is given by the Sand equation:

� �I

CDnF2

*2/12/1 � � (1.3)

where I is the current (in mA), C* is the bulk concentration (in mol·cm–3), and D is the diffusion coefficient expressed in cm2·s–1 [5]. For Cu2+ in aque-ous solution, D = 3.6·10–6 cm2·s–1 [6].

1.3 Deposition from solutions containing metal ion complexes

Positively charged metal ions are Lewis acids and many organic molecules form stable coordination compounds with metal ions in which oxygen or nitrogen are the electron donor atoms [7]. The formation of stable complexes enables deposition under kinetic (or activation) control, which generally results in smooth deposits [8]. Complexes are also often used for alloy depo-sition since the reduction potentials of two metal ions then can be brought together by the proper choice of complexes, as for example in the deposition of W-Ni alloys from ammonia-citrate electrolytes [9]. Citrate is a very com-mon complex agent that has been used for the deposition of Cu-Zn alloys [10] and Co-Ni-Cu alloys [11].

Deposition from alkaline solution is generally made possible by ligands forming soluble complexes with metal ions. Hydroxycarboxylic acids form very stable complexes with metal ions where both the carboxylic and the

14

alcoholic hydroxyl groups take part in complex formation [12]. When these acids are liberated from the complexes with the metal ion upon electro-chemical deposition, the hydroxyl groups are reprotonated again because of the small dissociation constants of these acids. This consumption of H+ (or generation of OH–) leads to a local rise in the pH value when the buffer ca-pacity of the solution has been consumed. Electrochemical deposition from complexes of hydroxycarboxylic may therefore induce local precipitation of hydroxides or oxides.

1.4 Deposition of metals The reductive electrochemical deposition of metals from aqueous solutions is a well-established field [4]. The optimum current density for the deposi-tion of compact coatings is generally that corresponding to the end of the Tafel linearity range, as the nucleation rate increases with more negative potentials but mass-transport limitations causes irregular growth [13].

Additives are commonly used in plating baths to improve the adhesion and material properties of the metal deposits. Leveling agents are organic molecules that adsorb on the surface and accelerate the rate of the deposition within trenches, thus enabling depositions of smooth and bright deposits [14, 15]. Tartrate has for example been used in tin and copper deposition [16].

Deposition of copper from electrolyte baths containing citric acid has been described for acidic solutions [17-19] as well as alkaline solutions [20, 21]. Examples include studies of the growth behavior of copper on TiN and TaN using Cu(II)-citrate solutions of pH 3.5. The effect of citrate on copper crystal size during pulsed deposition has also been studied [17]. A high nu-cleation density was found for the deposition of copper on Ta at pH 12 [21].

1.5 Deposition of metal oxides The deposition of metal oxides from aqueous solutions is mainly performed in alkaline solutions containing metal complexes. Electrochemical deposi-tion of metal oxides can be carried out under oxidizing conditions as well as reducing conditions from alkaline solutions. In both cases, the metal ion dissociates from the complex and precipitates on the electrode as the oxide. What controls the ability to deposit an oxide is the stability of that oxide under the experimental conditions, i.e. the potential, temperature and pH.

Deposition under oxidizing conditions includes the deposition of CuO from Cu(II)-tartrate [22] and Cu(I) cyanide [23], Bi2O3 from Bi(III)-tartrate [24], AgO from Ag(II)-acetate [25] and CeO2 from Ce(III)-acetate [26]. Further examples are the deposition of NiOx and MnOx from Ni(II) and Mn(II) ammine complexes, respectively [27, 28] and the deposition of Co3O4

15

from Co(II)-glycine in alkaline solutions [29]. In the latter cases, the oxida-tion state of the metal ion in the oxide is the same or higher than for the dis-solved ions. The oxide film is not stripped even though the current is oxida-tive. Instead species in the solution are oxidized, as is the case for the Ce(III) ions slowly released from Ce(III)-ligand complexes [30]. In the case of CuO deposition from Cu(II)-tartrate, it has been suggested that the oxidation of the ligand destabilizes the metal complex, thus enabling the anodic deposi-tion [22].

Deposition of oxides under reductive conditions includes the reduction of metal ions that form stable oxides. The Cu(I) oxide, Cu2O, has thus been demonstrated to be deposited from alkaline Cu(II) solutions containing vari-ous ligands [31, 32]. Otherwise cathodic deposition of metal oxides relies on precipitations due to electrochemically induced local pH alterations. Reduc-tion of dissolved oxygen or nitrate ions [33] results in local production of hydroxide ions, which for example enables the deposition of CdO and ZnO [34, 35]. However, hydroxide ions are also produced upon the dissociation of complexes with ligands like lactic, tartaric and citric acid containing �-hydroxyl groups [36], [I].

1.6 Deposition of nanostructured materials Nanostructured materials have been extensively studied during the last dec-ade due to their interesting electronic, magnetic, electrochemical and optical properties and potential use as catalytic- and electrode materials in various devices. Nanocrystalline metals and alloys has been achieved using pulses of high current density [37], while the effect of organic additives on nanocrys-talline copper has been specifically studied [38]. Nanocrystalline metal ox-ides have been prepared using oxidizing conditions in a non-aqueous me-dium [39].

Manganese oxide nanocomposites intercalated with polyelectrolytes have been deposited using oxidation of Mn2+ [40], while nanostructured compos-ites of Sn/SnO2 have been deposited by the method of local generation of base through the reduction of NO3

– ions [41]. Nanocomposites of Cu/Cu2O have been deposited from alkaline Cu(II)-lactate solution under conditions resulting in spontaneous potential oscillations [42]. Deposition of nanostruc-tured compounds based on spontaneous current oscillations has also been described for Cu-Sn [43]. Layered nanostructures of iron-group alloys have likewise been deposited under current oscillations arising from adsorp-tion/desorption of electrolyte components [44]. Copper deposits of different growth direction have been deposited using the oscillatory behaviour of ad-sorption of Cu(II)-complexes in the presence of o-phenanthroline [45].

16

Co-deposition of metal oxides during electrochemical deposition has been described for the mixed oxide of WO3-TiO2 [46] and for Al2O3 and TiO2 during nickel deposition [47].

Multi-layered structures of transition metals are quite common, as for ex-ample Co/Cu multilayers [48] and Ni/Co multilayers [49]. Multilayers of MoS2 have been deposited using an electrochemical/chemical synthesis method [50]. Magnetic multilayers of metal/metal-oxide have been produced electrochemically by deposition of the metal followed by anodic oxidation, as for example for Fe/Fe-oxide [51].

17

2 Experimental

2.1 Preparation of electrolytes and experimental setup Alkaline Cu(II)-citrate electrolytes, typically containing 0.4 M Cu(II) and 1.2 M citrate, were prepared by dissolving copper sulphate and citric acid in a small amount of deionised water. Concentrated, about 5 M, sodium hy-droxide was then added to the electrolyte until the pH was about 11. Final adjustment of the total volume of the solution was typically done after evaporation during one night in the laboratory fume hood. Controlled tem-peratures were maintained by a water bath with a double-walled glass beaker. The solutions were carefully purged with nitrogen to remove dis-solved oxygen and a flow a nitrogen gas was kept over the surface of the solution during experiments.

Antimony tartrate electrolytes were prepared by dissolving potassium an-timony tartrate trihydrate in deionised water. The pH of the resulting electro-lyte was about 4.1. Various buffer solutions were also prepared with the intention to keep the local pH low for the deposition of pure antimony films. Solutions with higher ligand to antimony ratios were made by additions of tartrate as potassium tartrate.

The electrochemical experiments were performed using a conventional three-electrode cell setup consisting of a working electrode, a counter elec-trode and a reference electrode. The working electrode in the experiments performed with forced convection was a polished platinum rod sealed in epoxy. When the electrochemical quartz crystal balance (EQCM) technique was used, the gold-coated crystals were made part of the electrochemical cell as the working electrode.

A laboratory designed combined sample holder and electrochemical cell was invented to minimize laboratory work when several samples of Sb and Sb/Sb2O3 coatings were needed in the electrochemical characterizations of the material as anode material in lithium ion batteries. The latter coatings were deposited at 23C in quiescent solutions.

18

2.2 Characterization of the electrodeposits Commonly used techniques for the study of electrochemical deposition are various electrochemical methods: cyclic voltammetry, potential step meth-ods and controlled-current techniques.

In cyclic voltammetry, the current is recorded as the potential is swept be-tween two potentials causing reduction and oxidation of species. The shape of the voltammogram gives information about the redox potential, and the occurrence of diffusion limitation or surface confined reactions. Varying the potential scan rate, which set the time-scale of the experiment, can test the reversibility of the electrochemical reaction.

When a constant potential is applied, the current is measured as a function of time. Potential step (Chronoamperometric) methods are used to determine diffusion coefficients and concentrations of electroactive species. The shape of the current-time curves can also give information about the type of nu-cleation taking place during the electrochemical deposition [52].

When a constant current is forced through an electrochemical cell, the po-tential takes on the value necessary to enable an electrochemical reaction to keep up with the current. Controlled current techniques are used among other things to study multicomponent systems and multistep reactions be-cause the measured potential changes as the surface concentration of electro-active species decreases sufficiently. The galvanostatic double pulse method can be used to determine rate constants of very rapid electron transfer reac-tions [5]. The thickness of a metal film can be determined by oxidation using a constant current.

In this thesis, the shape of potential-time curves during constant current electrolysis and the shape of voltammograms as well as that of potentiostatic current transients have been the matter of investigation. Electrochemical methods have likewise been used for the determination of the reduction po-tential of Cu2O [I] and Sb2O3. The Sb/Sb2O3 deposits were also electro-chemically characterized using constant current techniques in an organic electrolyte [IV, V].

The rotating disc electrode (RDE) is used for measurements under condi-tions of controlled mass transport of the reacting species. The rotating ring-disc electrode (RRDE) is often used for the detection of reaction intermedi-ates [5]. However, the latter technique requires more sophisticated and ex-pensive equipment in the form of a bi-potentiostat. RDE and RRDE methods were recently applied to develop an understanding of the reaction mecha-nism for the deposition of manganese oxide [53].

In this thesis, the RDE technique has been used to construct plots of cur-rent versus rotation rate using a constant potential [III]. The technique was also used to study the influence of mass transport during spontaneous poten-tial oscillations [I].

19

A powerful and very valuable tool in electrodeposition studies is the elec-trochemical quartz crystal microbalance (EQCM) technique. A gold-coated quartz crystal is then used as the working electrode in a conventional three-electrode setup. The quartz crystal frequency is continuously measured, al-lowing in-situ detection of mass changes of the working electrode. The de-tection of mass changes enables the determination of charge efficiencies during electrochemical depositions.

The EQCM technique relies on the converse piezoelectric effect where an applied electric field induces shear deformation of a thin quartz wafer, pre-pared from a special angle of cut with respect to the crystallographic axis of the quartz crystal [54]. An alternating potential across the quartz crystal causes a vibrational motion with amplitude parallel to the surface of the crystal that results in a transverse acoustic wave. Deposition of a layer on top of the quartz crystal increases the thickness of the resonator and results in a longer acoustic wavelength [54]. The measured frequency of the wave thus decreases according to the Sauerbrey equation:

� � 2/12

0 /2 qqAmff �� (2.1)

where f0 is the frequency of the bare quartz crystal, �m is the mass change, A is the electrode area, �q is the shear modulus and �q is the density of quartz. The Sauerbrey relationship requires that the deposited material is a rigid and very thin film on the surface of the shear wave antinode [5]. Fac-tors that lead to nonideal behavior include viscoelastic effects, high mass loadings, surface roughness, surface stress, interfacial slippage and nonuni-form mass distribution [54].

The EQCM technique has been used for studies of the deposition of e.g. PbTe [55], ZnO [56], Ag and Cu [57] and Cu2-xSe [58] and the effect of polyethylene glycol and chloride ions on copper electrodeposition [59]. An EQCM experimental setup including hydrodynamic control has been used for studies of phase composition quantification of electrodeposited ternary chalcogenide compounds [60].

The EQCM experiments in this thesis were performed with gold-coated quartz crystals (9 MHz, AT-cut) as working electrodes. The area of these electrodes was 0.196 cm2. Under the experimental conditions employed, a frequency change of 1 Hz corresponded to a mass change of 1.068 ng. The technique was used for the interpretation of potential- and current curves [I, III] the determination of layer thickness [II] and the detection of co-deposition of oxide [IV, V].

Structural information concerning materials can be obtained by scanning probe techniques such as transmission electron microscopy (TEM), scanning electron microscopy (SEM), scanning tunneling microscopy (STM) and atomic force microscopy (AFM). The latter techniques allow examinations

20

of electrode surfaces in air, as well as in solution during experiments. STM measures the tunneling current between the surface and a tip and has, for example, been used for the structural characterization of bright copper sur-faces [61]. STM has also been used for studies of nanoscale electrodeposited superlattices [62]. Atomic force microscopy (AFM) is used to create a map of the surface topography on a micrometer scale by scanning a small tip over the surface. AFM has been used for the study of the deposition of copper on platinum [63]. The morphology and nucleation kinetics of copper islands during electrodeposition was recently studied using in-situ TEM [64].

A very valuable technique is the electrochemical scanning tunneling mi-croscopy (ESTM), which measures the tunneling current during an electro-chemical experiment. Several studies on the initial stages of electrodeposi-tion on highly ordered surfaces have been performed with ESTM, for exam-ple Cu on Ag(100) [65] and Sb on Au(111) [66]. Mapping of the electro-chemical activity can be achieved by measurements of the current caused by an electrochemical reaction, using a technique called scanning electrochemi-cal microscopy (SECM) [67]. The latter technique can also be used for local-ized electrochemical deposition, as was demonstrated for copper microstruc-tures [68].

In this thesis, an electrochemical cell for AFM measurements was de-signed to study morphological changes during spontaneous potential oscilla-tions. However, the concentration of CuSO4 must not exceed 10 mM as the intensity of the laser light reaching the detector otherwise is too small. Due to the relatively low Cu(II) concentration, it was then difficult to obtain po-tential oscillations. Another drawback was that several minutes were re-quired to obtain an image, while the phenomenon to be studied took place on a far shorter timescale.

The technique of scanning electron microscopy (SEM) uses the detection of secondary electrons generated in the sample by an electron beam of high intensity. An image of the sample is created by the topographic contrast that arises due to variations of the intensity of electrons that reach the detector from different points of the sample surface [69]. SEM is a nowadays a stan-dard analysis technique and was recently used for studies of copper crystalli-zation [70]. In this thesis, the morphology of the electrodeposited films was characterized employing a LEO 1550 (Zeiss, Germany) instrument [I-V].

Spectroscopic techniques commonly used to study electrochemical inter-faces include ultraviolet and visible spectroscopy, vibrational spectroscopy, electron- and ion spectrometry and X-ray methods [5]. A powerful technique for the in-situ study of electrode surfaces is synchrotron X-ray diffraction [71].

X-ray diffraction (XRD) is based on the diffraction phenomenon that arises when electromagnetic radiation (light) with a wavelength of the order of interatomic distances interacts with crystalline (ordered) material [72]. In

21

this thesis, the phase analysis experiments were performed with a T2T Sie-mens D5000 diffractometer (Siemens, Germany) with CuK� radiation [I-V].

In X-ray photoelectron spectroscopy (XPS), the kinetic energy of elec-trons leaving the sample is measured, providing elemental identification and information about the bonding character by detection of chemical shifts in the binding energy of the electrons [73]. XPS is commonly used for surface analysis and was recently used for studies of the deposition of Cu [74]. In this thesis, XPS was used in paper [I] and [IV]. The experimental data were obtained with a Perkin Elmer PHI 5500 instrument with monochromatic Al K� radiation using pass energy of 58.75 eV.

The technique of Raman spectroscopy uses a monochromatic laser beam to create vibration and rotation movement of molecules in solution or vibra-tions in solid matter [75]. A small fraction of the scattered light emerges with a change in frequency and is detected by a CCD camera [75]. The tech-nique was recently used for the study of levelers for Cu electrodeposition from acidic sulphate solution [76]. In this thesis, the technique was used in [I] to detect surface layers of Cu2O and in [IV] to detect Sb2O3. The experi-mental data were obtained with a Renishaw 2000 Raman spectrometer equipped with a 20 mW near-infrared diode laser with a wavelength of 783 nm.

Ellipsometry is based on the rotation of plane-polarized light during inter-action of matter [77]. Since the electrode surface changes during an electro-chemical experiment, the technique is well-suited for in-situ studies of sur-faces [5]. The technique was recently used for the study of adsorption of surfactants on copper [78]. In this thesis, the ellipsometry technique was used to detect Cu2O during the cyclic voltammetry experiments in paper [III]. The off-null ellipsometry experiments were performed on an EP3 im-aging null-ellipsometer equipped with a laser with a wavelength of 532 nm and a CCD camera as detector.

22

3 Electrochemical deposition from alkaline Cu(II)-citrate solutions

Electrochemical deposition of nanostructured coatings of Cu/Cu2O as well as well-defined multilayers of Cu and Cu2O could be performed from solutions containing copper complexes. In paper [I] and [III] it is shown that nanos-tructured coatings of Cu/Cu2O were formed as a result of oscillations in the potential or current, respectively. Cu2O was also shown to precipitate due to the local pH increase during electrochemical deposition from the copper citrate complex. The electrochemical deposition of well-defined multilayers of Cu and Cu2O by a galvanostatic pulsing technique is described in paper [II], as Cu2O can be electrochemically deposited from alkaline Cu(II)-citrate solutions.

In papers [I-III], EQCM and ellipsometry data show that Cu2O is not thermodynamically stable at the potentials where Cu is being deposited. In paper [II] it is shown that the ability to deposit a layer of Cu on top of a Cu2O layer, without entirely reducing the latter, relies on the preferential reduction of Cu2+ in highly concentrated Cu(II)-citrate solutions. In paper [I], EQCM and Raman spectroscopy results show that copper deposition from the present solutions proceeds by a two-step electron reduction reac-tion, in which Cu2O can be formed as an intermediate species.

3.1 Complex speciation and electrode reactions Citric acid is an organic acid with the molecular formula COOH-CH2-COH(COOH)-CH2-COOH that contains three carboxylic groups and one �-hydroxyl group. The basic form of citric acid is called citrate, Cit3–. When citrate forms complexes with copper, the �-hydroxyl proton is detached [79]. The latter form of citrate will be denoted H–1Cit4– in this thesis. Titration curves (see figure 3.1) were recorded showing that an equivalent of 3.5 and 4 protons had been detached at the equivalence points in solutions with Cu(II) to citrate ratios of 0.5:1 and 1:1, respectively. This result verifies that the hydroxyl proton is detached due to complex formation with Cu2+ ions in aqueous solutions.

23

Fig. 3.1 Titration of 50 ml 0.1 M citric acid + 0.05 M CuSO4 (left curve) and of 0.1 M citric acid + 0.1 M CuSO4 (right curve) with 5 M NaOH.

The distribution of Cu(II)-citrate species as a function of pH and total citrate concentration has been calculated [80] and the dimer [Cu2H–2Cit2]

4– has been found to be the predominant species in Cu(II)-citrate solutions at pH values greater than 7 in the excess of citrate [79-81]. The electrochemical deposition of copper thus proceeds according to reaction (3.1).

In the absence of copper ions, Cit3� dominates above pH 7 [80]. The H�1Cit4� species for which the �-hydroxyl group also has been detached is only stable in very alkaline solutions [81]. The �-hydroxyl proton is there-fore restored to form the Cit3� species according to reaction (3.2) when the Cu(II)-citrate complex dissociates. As the dissociation of the complex takes place at the working electrode this results in a local production of hydroxide ions during the electrochemical deposition of copper according to reaction (3.3).

[Cu2H–2Cit2]

4– + 4e– = 2Cu + 2H–1Cit4– (3.1) 2H–1Cit4– + 2H2O = 2Cit3– + 2OH– (3.2) [Cu2H–2Cit2]

4– + 4e– + 2H2O = 2Cu + 2Cit3– + 2OH– (3.3)

The logarithm of the stability constant for the reaction 2Cu2+ + 2Cit3– = [Cu2H–2Cit2]

4– + 2H+ has been determined to 5.2 ± 0.1 [82] by potentiometric titration with copper ion-selective electrode and to 5.8 ± 0.07 [79] by means of combined pH and pCu measurements. The strength of the Cu2+-citrate complex should therefore be high enough to keep the concentration of free Cu2+ (aq) ions below the limit of precipitation of Cu(OH)2 (s) in alkaline solutions. The solubility product of Cu(OH)2 (s) is Ksp = 10-20 [83], which

24

means that the concentration of free Cu2+ must not exceed 10-14 M. This re-quirement is met at pH 11 due to the presence of the citrate complex.

The logarithm of the stability constant for the reaction Cu + Cu2+ + 2OH– = Cu2O + H2O is 23.5 [83]. Thermodynamically, Cu2O is formed by com-proportionation between Cu and Cu2+ if log [Cu2+] > 4.5 – 2pH. At pH 11, Cu(II)-citrate solutions are not thermodynamically stable in the presence of copper with respect to formation of Cu2O and comproportionation occurs according to reaction (3.4).

[Cu2H–2Cit2]

4– + 2Cu + 2OH– = 2Cu2O + 2Cit3– (3.4)

Electrochemical deposition of Cu2O from alkaline Cu(II)-citrate solutions, as is described in chapter 3.2 and 3.3, proceeds by a continuous precipitation of Cu+ ions (see reaction 3.6) generated by reduction of the Cu(II)-citrate com-plex (see reaction 3.5). Since the OH– ions consumed in reaction (3.6) are restored due to reaction (3.2), there is no net production of OH– during the deposition of Cu2O according to reaction (3.7).

[Cu2H–2Cit2]

4– + 2e– = 2Cu+ + 2H–1Cit4– (3.5) 2Cu+ + 2OH– = Cu2O + H2O (3.6) [Cu2H–2Cit2]

4– + 2e– + H2O = Cu2O + 2Cit3– (3.7)

As will be described in chapter 3.4, Cu2O can be reduced to copper accord-ing to reaction (3.8). Cu2O + 2e– + H2O = 2Cu + OH– (3.8)

The dissociation rate constant of the [Cu2H–2Cit2]4– complex has been deter-

mined to 3·10-2 s–1 by a chromatographic technique using 64Cu [84], employ-ing a solution adjusted to pH 6. Since the currents used in the experiments are not limited by mass transport [III], it is reasonable to assume that the electrochemical deposition of copper proceeds by reduction of Cu2+, dissoci-ated from the [Cu2H–2Cit2]

4– complex, according to reaction (3.9) and (3.10). [Cu2H–2Cit2]

4– = 2Cu2+ + 2H–1Cit4– (3.9) 2Cu2+ + 4e– = 2Cu (3.10)

3.2 Deposition of copper The charge efficiency for the deposition of copper was about 90 %, as calculated from the EQCM frequency data [I]. The electrochemical deposi-tion of copper in Cu(II)-citrate solutions generally proceeds with charge efficiencies lower than 100%, as previously observed [80]. EQCM data

25

show that the frequency continues to decline after the current has been turned off [I], as has also been observed during copper deposition in acidic solutions [85].

Precipitation of Cu2O occurs during open circuit conditions after copper has been deposited, as verified by Raman measurements [I]. The accumulation of Cu+ during the reduction of Cu2+ has previously been stud-ied in 5 M HClO4 [86]. The deposition of copper proceeds by a two-step electron transfer to Cu2+, while a small part of the generated Cu+ ions, stabilised as Cu2O, escapes into the solution. A loss of Cu(I) into the solution is also in accordance with previous measurements with the rotating ring-disc electrode (RRDE) technique [87, 88]. Figure (3.2) depicts the open circuit mass increase detected by EQCM after copper has been deposited for differ-ent times. The graph shows that the amount of Cu2O that precipitates in-creases with increasing copper deposition charge. The positive intercept provides evidence for reaction (3.4).

Fig 3.2 The open circuit mass increase as a function of the duration of a preceding galvanostatic copper deposition step using a current density of –1.0 mA/cm2 and a temperature of 50 ºC. The solution contained 0.4 M CuSO4 and 1.2 M citrate, pH 11.

3.3 Deposition of Cu2O Electrochemical deposition of Cu2O from alkaline Cu(II)-citrate solutions has been achieved in a solution containing 0.4 M Cu(II) and 1.2 M citrate adjusted to about pH 12.5. The deposition was performed at 50 ºC employ-ing a constant potential of -0.30 V. There is also a number of reports on the deposition of Cu2O from alkaline Cu(II)-lactate solutions [89-92].

The chronopotentiogram obtained with the EQCM (see fig. 3.3) shows that the current stabilizes at about –1.0 mA/cm2 after a long nucleation pe-

26

riod. The X-ray diffractogram (figure 3.4) shows that the (111) diffraction peak dominates, indicating that [111] is the kinetically preferred orientation of the crystal growth. The influence of the deposition potential on the orien-tation of Cu2O films is further discussed in paper [II].

The flux of Cu2+ to the electrode surface depends on the rate of dissocia-tion of the complex, which in turn is a function of the nature of the complex, pH and temperature. Since EDTA is a very strong complexing agent, pre-liminary experiments showed that the nucleation rate of Cu2O became very low even when the pH was raised to 13.

Fig 3.3 Potential– (thick line) and EQCM frequency– (thin line) versus time curves during a potentiostatic deposition step for a solution containing 0.4 M CuSO4 and 1.2 M citrate at pH 12.5. The potential was E = – 0.3 V vs. SCE and the temperature was 50 ºC. The charge efficiency for the deposition of Cu2O was 88 %.

Fig. 3.4 XRD diffractogram recorded for the deposit resulting from the experiment depicted in figure (3.3). The peak denoted by * stems from the gold substrate.

27

Deposition of Cu2O has also been performed in polycarbonate templates (see figure 3.6). XRD measurements showed marked (111) texture when the temperature was 40 ºC and the potential was –0.2 V vs. SCE.

Fig 3.5 Transmission electron microscopy image (TEM) showing cylinders of Cu2O, about 20 �m in length and 300 nm in diameter. The electrodeposition was performed using a polycarbonate template with around 8 holes per �m2 at a temperature of 40 ºC. The electrolyte contained 0.86 M CuSO4 and 0.86 M citrate, pH 10.6. A total charge of 0.78 C had passed after 12 hours using a constant potential of E = – 0.2 V vs. SCE.

3.4 Electrochemical stability of Cu2O The electrochemical stability of electrochemically deposited Cu2O films has been studied with the EQCM technique, as has been described previously [88, 93]. Figure 3.6 shows that the reduction of a 150 nm thick Cu2O coating in 0.5 M KNO3 starts at a potential of about –0.6 V, as the current starts to increase according to reaction (3.8) and the EQCM frequency simultane-ously starts to increase (i.e. mass is lost). The more negative potentials needed to reduce the remaining part of the Cu2O is most probably due to the formation of a passivating layer of Cu, as is described in paper [I] and [II].

In-situ ellipsometry data show that the potential at which Cu deposition replaces that of Cu2O is about –0.5 V vs. SCE in a Cu(II)-citrate solution, as is seen in figure 3.7. This transition is not detected as clearly in the EQCM–CV experiments, as will be shown in chapter 3.5 (figure 3.12). The small anodic peak on the cathodic scan, together with the ellipsometry intensity results, show that Cu2O also is formed by oxidation of Cu in alkaline solu-tion, as is further described in paper [III].

28

Fig 3.6 Reduction of a 150 nm thick coating of Cu2O (prepared in alkaline Cu(II)-citrate solution at pH 12 using E = -0.3 V) during a potentiodynamic sweep (thick line) from –0.3 to –1.5 V vs. SCE at a scan rate of 1 mV/s. The electrolyte contained 1.2 M citrate adjusted to pH 11 and was kept at 50 ºC. The thin line denotes the corresponding EQCM frequency, which is seen to increase (i.e. the mass decreases).

Fig 3.7 The second cycle voltammogram (thick line) and in situ ellipsometric inten-sities (thin line) as a function of the potential during a 10 mV/s voltammetric scan in a solution containing 0.4 M Cu(II) and 1.2 M citrate adjusted to pH 11. The working electrode was a gold electrode with a Cu2O coating of a thickness of about 9 nm (deposited at -0.45 V). The temperature was about 50 ºC at the start of the first scan.

29

3.5 Deposition of nano-composites of Cu/Cu2O Nanostructured coatings of Cu/Cu2O were obtained during oscillations in the potential or current. The oscillations were obtained at 50 ºC employing a constant current density of –1.0 mA/cm2 or a constant potential of -0.90 V, respectively, in a solution containing 0.4 M Cu(II) and 1.2 M citrate adjusted to about pH 11 (See figures 3.8 and 3.11).

The potential oscillations resulted in a coating consisting of both Cu and Cu2O, as is seen in figure 3.9. The XRD analysis showed that the particles of Cu and Cu2O were small, which is also seen in the scanning electron micro-graph in figure 3.10. The nanostructured morphology is supposed to stem from the local precipitation of Cu2O during the oscillations.

While similar oscillations in the potential has been described for the alkaline Cu(II) lactate– [94, 95] and tartrate– [96] systems, oscillations in the current have not yet been reported. In the present system the oscillations are due to local changes in the pH, as described in [I] and [III].

The proposed oscillation mechanism is based on a change of the deposition rate of Cu2O according to reaction (3.7) that is dependent on the local pH. When the latter rate is sufficient to maintain the constant current, the potential is less negative (i.e within the potential peak region). The pH increases as a result of the local production of hydroxide ions according to reaction (3.3) when the buffer capacity has been consumed. The pH decreases due to the comproportionation reaction (3.4) and diffusion of OH– away from the electrode that is expected at potentials where Cu2O is stable on the electrode surface.

Fig 3.8 Potential oscillations (red line) obtained during constant current deposition using a current density of –1.0 mA/cm2 and a temperature of 50 ºC. The deposition bath contained 0.4 M CuSO4 and 1.2 M citric acid, adjusted to pH 11.2. Also seen in the figure is the EQCM frequency curve (black line). (See paper [I] for details.)

30

Fig 3.9 X-ray diffractogram of the Cu/Cu2O coating obtained during the spontane-ous potential oscillations depicted in fig. (3.8). The peak denoted by * stems from the gold substrate.

Fig 3.10 Scanning electron micrograph of the Cu/Cu2O coating obtained during the spontaneous potential oscillations depicted in fig. (3.8).

Oscillations in the current was observed during potentiostatic deposition from alkaline Cu(II)-citrate solutions at elevated temperatures, as is seen in figure 3.11. These oscillations are analogous to the spontaneous potential oscillations during controlled current conditions as they result from changes in the Cu(II)-reduction rate due to local pH changes [I, III]. The Cu2O depo-sition is preceded and facilitated by saturation of the solution in the vicinity of the electrode with respect to Cu2O, as is indicated by detailed analysis of the mass changes in the region of the current peak [III]. An XRD analysis showed that the material deposited under the current oscillations likewise contained both Cu2O and Cu of small particle sizes [II].

31

Fig 3.11 Current oscillations (red line) and EQCM curve (black line) obtained at 50 ºC employing a constant potential of -0.90 V in a solution containing 0.4 M Cu(II) and 1.2 M citrate adjusted to about pH 11 showing the deposition experiment with 30 peaks during 3000 seconds. (See paper [III] for details.)

Cyclic voltammetry (see fig. 3.12) can be used to test if a solution of Cu(II) containing hydroxycarboxylic acids is likely to give rise to spontaneous os-cillations in the potential or current [III]. For an oscillating system, a ca-thodic peak should be seen during the anodic scan [III].

The sharp cathodic peak is most likely a result the local pH increase, as deduced from data obtained with the EQCM and the RDE [III]. The pres-ence of a surface confined reaction is excluded because the amount of Cu2O generated in the diffusion layer before the current peak according to reaction (3.6) is not sufficient to explain the peak charge [III].

32

Fig. 3.12 Cyclic voltammogram (thick line) and EQCM frequency (thin line) versus potential plot obtained at 50 ºC with a EQCM gold electrode in a solution containing 0.4 M Cu(II) and 1.2 M citrate adjusted to about pH 11. The scan rate was 10 mV/s. The voltammogram and EQCM curve denote the third cycle responses.

3.6 Deposition of well-defined layers of Cu and Cu2O Thick layers of Cu and Cu2O have been deposited from alkaline Cu(II)-citrate solutions, as is shown in figure 3.13. The method is based on pulsing the constant current density between a relatively high current density (i = –1.0 mA/cm2), where Cu is deposited according to reaction (3.3), and a lower current density (i = –0.1 mA/cm2), where Cu2O is deposited according to reaction (3.7), see figure 3.14 [II]. The layer thicknesses can then be independently controlled by setting the deposition time to the desired value. This is clearly not possible for the spontaneous oscillations described in chapter 3.5. The Cu layer thickness is limited by the onset of oscillations, whereas micrometer thick Cu2O layers can be deposited.

If a layer of Cu first is deposited, the nucleation of Cu2O is generally straightforward due to reaction (3.4) and a thick layer of Cu2O may be deposited. The problem arises when the next layer of Cu is to be deposited on top of the previously deposited Cu2O layer, as the latter is not stable at the potentials where Cu is deposited and therefore will be partially reduced according to reaction (3.8). However, because of the high concentration of Cu(II) used in the deposition experiments, the current is still carried by reaction (3.3) and a layer of Cu protects the underlying Cu2O layer from

33

further reduction according to reaction (3.8). The nucleation of Cu on a layer of Cu2O has previously been reported to be difficult [97]. However, the method presented in this thesis enables deposition of Cu on top of Cu2O without addition of compounds that facilitate the Cu nucleation. This provides a way of depositing well-defined multilayers of Cu and Cu2O.

Fig. 3.13 SEM micrograph of a cross section of the EQCM electrode used in the experiment described in Figure (3.14) showing three thick Cu2O layers intervened by thin Cu layers. (Result described in [II].)

Fig 3.14 Potential (thick line) and frequency (thin line) versus time curves during pulses of a cathodic constant current density of –1.0 mA/cm2 for 300 seconds fol-lowed by –0.1 mAcm2 for 900 seconds. The deposition bath was composed of 0.4 M CuSO4 and 1.2 M citric acid, adjusted to pH 11 and kept at 50 ºC. (Results described in [II].

34

4 Electrochemical deposition from Sb(III)-tartrate solutions

Electrochemical deposition of nanostructured anode materials for Li-ion batteries could be performed from solutions containing antimony complexes. In paper [IV] and [V] it is shown that nanoporous coatings of Sb and Sb2O3 were formed, as Sb2O3 precipitated due to the local pH increase during the electrochemical deposition from the antimony tartrate complex. The difficul-ties associated with the deposition of thick films of Sb/Sb2O3 with similar morphologies as that for thin films are discussed at the end of this chapter.

The cycling performance of the Sb/Sb2O3 coatings used as anodes in lith-ium ion batteries were very good due to the presence of small Sb particles that became embedded in a matrix of nanoporous Li2O, upon the reduction of the co-deposited Sb2O3. Depositions of pure antimony films were per-formed to compare the shape of the galvanostatic potential-time cycling curves with those for coatings also containing Sb2O3.

4.1 Complex speciation and electrode reactions Tartaric acid, –OOC-CH(OH)-CH(OH)-COO–, is a dicarboxylic acid for which the pKa values for the carboxyl groups are 2.98 and 3.34, respectively [98]. The molecule also has two �-hydroxyl groups with pKa values of 15.5±0.5 and 17.5±0.5 [99]. The protons of the hydroxyl groups are lost when this organic ligand forms complexes with Sb(III) in which both the two carboxylic groups and the two hydroxyl groups coordinate to the central atom. If H�2Tar4� denotes the fully deprotonated form of tartaric acid, �OOC–CH(O�)–CH(O�)–COO�, the complex can be written as [SbH�2Tar]�. Although we have not been able to find the relevant stability constants for the antimony-tartrate system, electrospray mass spectrometry data [100] indicate that the dimer [Sb2H–4Tar2]

2– is the dominating species in the pH 4.1 solution used in this work. The electrochemical deposition should thus proceed according to reaction (4.1). The hydroxyl protons are restored upon the liberation of the tartrate ligand from the complex with the Sb(III) ion according to reaction (4.2), yielding the total reaction (4.3).

35

[Sb2H–4Tar2]2– + 6e– = 2Sb + 2H–2Tar4– (4.1)

2H�2Tar4� + 4H2O = 2Tar2� + 4OH� (4.2) [Sb2H–4Tar2]

2– + 4H2O + 6e– = 2Sb + 2Tar2– + 4OH– (4.3)

As the pH of the solution near the working electrode increases, Sb2O3 can precipitate according to reaction (4.4). [Sb2H–4Tar2]

2– + 2 OH– + H2O = Sb2O3 + 2Tar2– (4.4)

According to the Pourbaix diagram for the antimony-water system [101], the stable form of Sb(III) is SbO+ below pH 1, HSbO2 at pH values between 1 and 11, and Sb(OH)4

– for pH > 11. In acidified antimony tartrate solutions with pH values significantly lower than the dissociation constants of the carboxyl groups of tartrate, i.e. at pH < 2, the predominating Sb(III) species should be HSbO2 and reduction should occur according to (4.5). When chlo-ride ions are present, the antimony is instead likely to be present as SbCl2+and SbCl2

+ [102] and the reactions should be written as (4.6) or (4.7) depending on the concentration of Cl–.

HSbO2 + 3 H+ + 3 e- � Sb + 2 H2O (4.5) SbCl2+ + 3 e- � Sb + Cl- (4.6) SbCl2

+ + 3 e- � Sb + 2 Cl- (4.7)

According to thermodynamic considerations, the Sb2O3 particles that reach the electrode surface should be reduced to antimony according to reaction (4.8).

Sb2O3 + 6H+ + 6e– = 2Sb + 3H2O (4.8)

4.2 Deposition of nano-composites of Sb and Sb2O3 The solution used in the experiments for the deposition of nano-composites of Sb and Sb2O3 was prepared by dissolving potassium antimony(III)tartrate in deionized water to a concentration of 0.15 M. The pH was about 4 and the buffer capacity was low. The charge efficiency was equal to 120 %, indicating that precipitation of Sb2O3 occurred during the deposition according to reaction (4.4) due to the increase in the local pH during the deposition of Sb according to reaction (4.3).

36

Fig. 4.1 Chronopotentiogram and EQCM frequency vs. time curves during the depo-sition of a thin Sb/Sb2O3 coating. The solution was 0.15 M potassium antimony tartrate at pH 4.1 and the current density 1.9 mA/cm2.

The presence of precipitated Sb2O3 in the antimony coating has also been verified by XPS and Raman measurements [IV]. As is also seen in figure 4.1, the constant slope of the frequency curve indicates that Sb2O3 was con-tinuously co-deposited during the deposition. While reduction of Sb2O3 is expected at the potentials needed for the deposition according to reaction (4.8), the constant co-deposition of Sb2O3 most likely proceeds as a result of a sufficiently slow reduction of the oxide.

Fig. 4.2 Scanning electron micrograph for an Sb2O3-containing coating deposited using a constant current of 1.92 mA/cm2 for 120 s.

37

Fig. 4.3 TEM micrographs showing small (<2 nm) Sb2O3 nanoparticles agglomer-ated onto the surface of larger Sb particles.

The precipitation of Sb2O3 provides nucleation sites, which facilitate the deposition of antimony. It is therefore easy to obtain a coating of even distri-bution over a large electrode area without pre-treatment of the substrate. Scanning electrode microscopy shows cauliflower-like morphologies of the deposits, as is seen in figure 4.2. TEM analysis showed that the deposit con-sisted of Sb and Sb2O3 nanoparticles (see figure 4.3). The co-deposition of Sb2O3 influenced the size of the Sb grains [IV]. X-ray diffraction shows a very broad (012) Sb peak at 28.7 ºC in 2, from which a crystallite size of 16 nm was obtained using Scherrer’s formula [IV].

Fig. 4.4 X-ray diffraction pattern obtained for an Sb/Sb2O3 coating. The * symbols denote nickel substrate peaks, while all other peaks are due to Sb.

38

Electrochemical deposition from solutions of lower pH, as well as solutions with higher buffer capacities, resulted in lower concentrations of Sb2O3 in the deposits and larger particles as a result of less precipitation of Sb2O3 in the vicinity of the working electrode [V].

4.3 Electrochemical characterization of the deposited material

Reduction of co-deposited Sb2O3 was performed in aqueous solutions using constant current electrolysis in 0.5 M KNO3. The reduction of the precipi-tated Sb2O3 particles started at potentials below –1 V vs. SCE. The mass loss during the reduction was determined from recordings of the frequency changes with the EQCM, as is shown in figure 4.5. In the figure, it is seen that there was some dissolution of Sb2O3 in the acidic solution but also that the onset of the current induced a significant change in the slope of the fre-quency versus time curve. The results indicate that about 90 % of the co-deposited Sb2O3 was reduced.

Figure 4.5 Potential (thick line) and EQCM frequency (thin line) during reduction of a Sb/Sb2O3 coating in acidified 0.5 M KNO3 using a current density of –2.0 mA/cm2. The frequency curve shows that the mass was decreasing during the ex-periment.

Reduction of Sb2O3 and Sb in organic solutions was performed as the depos-ited material was tested as anode material for lithium ion-batteries. In the chronopotentiometric cycling method used, the alloying/de-alloying of the Sb/Sb2O3 deposits took place in organic solutions containing Li+ ions.

As is seen in figure 4.6, the specific capacity was about 930 mA·h·g�1, which is much higher than the theoretical value of 660 mA·h·g�1 expected

39

for the formation of Li3Sb according to reaction (4.9). This discrepancy can be explained by the much higher theoretical capacity of 1103 mA·h·g�1 for the reduction Sb2O3 according to reaction (4.10), as the coatings contained 20-25 % (w/w) Sb2O3.

The reduction potential versus Li+/Li in an organic electrolyte depends on the lithiation reaction, particle size and texture of the electrodeposited coat-ing. During the first cycle for the Sb/Sb2O3 deposit, a plateau is seen at a potential of around 1.6 V vs. Li+/Li. The constant current forced through the electrochemical cell is then maintained by the reduction of Sb2O3 yielding Sb and Li2O according to reaction (4.11). The charge used for reaction (4.11) as deduced from the potential curve in fact provides an alternative way to calculate the amount of Sb2O3 in the deposits.

Sb + 3Li+ + 3e– = Li3Sb (4.9) Sb2O3 + 12Li+ + 12e– = 2Li3Sb + 3Li2O (4.10) Sb2O3 + 6Li+ + 6e� = 3Li2O +2Sb (4.11)

Fig. 4.6 Chronopotentiograms showing the first cycle for a pure Sb (solid line) and an Sb/Sb2O3 (dotted line) coating, respectively. The potential scale refers to the potential of the working electrode versus the Li+/Li couple. The cycling rate was C/10 (i.e. 10 h for a complete lithiation of the electrode material).

The cycling stability of the Sb/Sb2O3 coatings as Li-ion anode materials was tested by repeated galvanostatic cycling between 1.5 and 0.01 V vs. Li+/Li using a constant current density. The reversibility of the lithiation-delithiation process of the Sb/Sb2O3 coatings was good as the charge used for the lithiation of the Sb/Sb2O3 material up to the turning point of 0.01 V vs. Li+/Li was found to be practically constant for more than 50 cycles. The

40

excellent cycling stability of this anode material was ascribed to the small particle size and the formation of an amorphous Li2O matrix, as is well-known from other metal oxide materials [103-106].

The experimental specific capacity obtained from figure 4.7 was 640 mA·h·g�1. The theoretical capacity for a coating containing 25% Sb2O3 is 771 mA·h·g�1, assuming reversible reduction of Sb2O3 according to reaction (4.11). The fact that the cycling is stable at a level of 83% of the latter value suggests that the Sb2O3 contributes to the capacity also during repetitive cycling of the coating.

The loss of capacity for the pure Sb coating also seen in figure 4.7 is most likely due to cracking of the material as a result of the volume changes asso-ciated with the formation of Li3Sb during the cycling [107, 108].

Fig. 4.7 Specific capacities as a function of the cycle number during galvanostatic cyling between 1.5 and 0.01 V vs. Li+/Li for a pure Sb (squares) and a Sb/Sb2O3 (stars) coating, respectively. The cycling rate was C/10 (i.e. 10 h for a complete lithiation of the electrode material).

4.4 Deposition of thick Sb2O3-containing films For longer deposition times than that used in figure 4.1, the slope of the frequency curve was found to flatten out and the amount of co-deposited antimony oxide decreased. During the deposition, the concentration of free tartrate increases locally as the complex dissociates. The composition of the diffusion layer therefore changes during the deposition and at higher tartrate to Sb(III) ratios the extent of precipitation of Sb2O3 decreases. This effect makes it difficult to deposit coatings of Sb and Sb2O3 that are thicker than about 200 nm.

41

One way to circumvent this problem could be to restore the equilibrium conditions using pulsed deposition, allowing for precipitation of Sb2O3 according to reaction (4.4) during reduction of Sb according to reaction (4.3). However, as detected by the EQCM, the mass decreases under periods of open circuit, probably because of dissolution of the Sb2O3 as the more stable Sb(III)-tartrate complex is formed.

To counterbalance the effect of the higher local tartrate concentration, deposition was performed at a potential negative enough for the onset of reduction of water, which produces hydroxide ions. About 50 �m thick, oxide-containing films was deposited in this way (unpublished results). The XRD data indicated the presence of nano-crystalline Sb. However, the coating rapidly lost capacity upon galvanostatic cycling.

To carefully investigate the influence of the temperature on the deposition conditions, deposition using constant potentials was carried out in quiescent Sb(III)-tartrate electrolytes at temperatures ranging from 10 ºC to 70 ºC. The current increased when the temperature was raised, indicating that the rate-limiting step is the temperature dependent dissociation of the Sb(III)-tartrate complex, in analogy with the case for the Cu(II)-citrate complex [III]. EQCM data showed that the mass increased during conditions of open cir-cuit after the deposition step at elevated temperatures, probably due to pre-cipitation of Sb2O3 according to reaction (4.4) (unpublished data). Recent results show that thick oxide-containing films can be deposited using higher temperatures if higher current densities than those used for the previously deposited successful films are used. A coating deposited at 70 ºC proved to contain large amounts of Sb2O3 judging from the charge efficiency. How-ever, this coating would probably not be a good Li-ion anode material as SEM showed large crystals, probably due to rapid grain growth.

4.5 Deposition of pure antimony films The electrochemical deposition of pure antimony films can be performed [IV, V] from acidified Sb(III)-tartrate electrolytes. At pH 1.3, the buffer capacity is high enough to cope with the consumption of H+ during the depo-sition according to reaction (4.5). Thick films of antimony was readily be prepared from acidified solutions also containing chloride, where reaction (4.6) and (4.7) also takes part in the deposition of Sb. An even distribution of the film over the large (31.5 cm2) electrode area was achieved when the nu-cleation of antimony was facilitated by etching the nickel foil with 4 M HNO3 at 40 ºC to remove surface oxide.

42

Fig 4.8. Scanning electron micrographs for crystalline films of antimony showing a thin film (left) and a thick film (right) with larger crystallites. The films were depos-ited from an acidified chloride containing solution (150 mM potassium antimony tartrate , 91 mM HCl, 110 mM KCl) using a current density of –4.0 mA/cm2. (See paper [V] for details.)

The scanning electron micrograph in figure 4.8 shows large well-defined Sb crystals and the X-ray diffraction pattern showed no peak broadening due to a small Sb crystal size [V]. The Raman spectrum showed peaks characteris-tic for antimony and no peaks due to Sb2O3 [IV]. An XPS depth profile analysis showed that Sb2O3 was present only on the surface since the inten-sity for the Sb(III) peaks decreased dramatically after 60 s of sputtering [IV].

43

Concluding remarks

This thesis deals with electrochemical deposition from metal complexes with organic molecules containing �-hydroxyl protons in aqueous solutions. The results indicate that deposition from these complexes is accompanied by a local production of hydroxide ions. The results have been discussed in the context of finding suitable conditions for depositions of the metal and the metal oxide, as well as to reach an understanding of fundamental aspects during the deposition including the origin of spontaneous oscillations in the potential or current. Two different systems have been studied where an in-crement in the local pH is induced by the protonation of the hydroxyl groups of the released organic ligands. In the alkaline Cu(II)-citrate system, Cu2O is stable as an intermediate oxidation state. In the Sb(III)-tartrate system, Sb2O3 precipitates in deposition baths with low tartrate to antimony ion ratios.

Deposition of copper from alkaline Cu(II)-citrate electrolytes proceeds as a two-step reduction reaction of Cu2+, in which the intermediate Cu+ species can be lost in the solution as Cu2O. The deposition rate of Cu2O on the elec-trode surface depends strongly on the local pH. The deposition of Cu2O films from alkaline Cu(II)-citrate electrolytes is possible within a limited range of current densities (or potentials). Deposition of Cu takes over when the constant current forced trough the cell is set higher than the nucleation rate of Cu2O. Although Cu becomes the thermodynamically stable phase when the electrochemical potential shifts negatively, reduction of Cu2O is hindered by the deposition of a protecting layer of Cu. Well-defined layers of Cu and Cu2O have been deposited using this approach.

Spontaneous potential oscillations arise in the Cu(II)-citrate system due to variations of the local pH, which govern the deposition rate of Cu2O. The model presented in this thesis predicts that oscillations should be possible to obtain also for other Cu(II)-complexes provided that the ligand undergoes protonation upon its liberation from the complex and that a decrease in the pH can be obtained as a result of comproportionation. The present approach, which provides new exciting possibilities for the design of nanostructured layers of copper and cuprous oxide, can most likely also be used to gain a better understanding of the behavior of other Cu(II)-complex based electro-lytes commonly used for the electroplating of copper.

Nanostructured coatings of Sb and Sb/Sb2O3 can be obtained by electro-chemical deposition from poorly buffered antimony tartrate solutions in the absence of an excess of ligand. These coatings exhibit excellent properties as

44

anode material for Li-ion batteries, mainly due to the small Sb particle size. The stability of the electrodeposited Sb/Sb2O3 coatings, when used as anodes in Li-ion batteries, depends on the concentration of Sb2O3 in the coatings, the size of the Sb and Sb2O3 particles, as well as the texture of the grains. The preparation of these coatings is complicated by the local accumulation of tartrate ligands during the deposition, which hinders co-deposition of Sb2O3. However, deposition of thick oxide-containing films are possible by per-forming the electrolysis at a constant potential sufficiently negative to enable reduction of water where the extensive local production of hydroxide ions counteracts the influence of the excess of the tartrate. Electrochemical depo-sition of pure antimony coatings from Sb(III)-tartrate electrolytes is possible at conditions where precipitation of Sb2O3 is prevented.

In summary, it can be concluded that electrochemical deposition readily can be used to make metal coatings as well as oxide- containing coatings. A protective layer of metal can be deposited on top of the oxide in solutions of high concentration of metal ions. Metal oxides with stable intermediate oxi-dation states can be deposited in a controlled way. Nanometer-sized metal particles can be manufactured by reduction of locally precipitated metal ox-ide particles.

45

Acknowledgements

First of all I would like to express my gratitude to Professor Leif Nyholm for the excellent education in fundamental electrochemistry that I have received.

I am indebted to Professor Jan-Otto Carlsson, Research Coordinator of Inorganic Chemistry and Vice-rector for the faculty of Science and Technol-ogy, for providing me the opportunity to perform my work at the Department of materials chemistry and for the materials analysis facilities put at my dis-posal.

I wish to express my appreciation of Professor Kristina Edström, Dean of Chemistry, who introduced me to the group of Professor Josh Thomas at our department. It has been instructive to take part of measurements at the Swed-ish National Synchrotron Radioation Laboratory in Lund (MAXLAB) and workshops within the European network for research on energy storage ma-terials (Alistore).

I am indeed very grateful to Dr. Merja Herranen for the initial support with computer software and for the acquisition of the electrochemical quartz crystal microbalance instrument (EQCM), which has been invaluable for my research. The means provided by the Göran Gustafsson Foundation is also gratefully acknowledged.

I am very thankful for the co-operation with Ph. D. student Hanna Bryn-gelsson, concerning the antimony deposits that were used as lithium ion anode material.

I wish to thank all persons who have taught me chemistry (mentioned without titles): Krister Åkervall (gymnasium), Marit Andersson, Maja Elm-gren, Kerstin Sunnerheim, Ove Bohman, Erik Dahl, Anders Wahlberg and Anders Eriksson to mention a few.

Particularly, I wish to record my appreciation of Professor Rolf Berger for his skilful teaching in inorganic and solid-state chemistry. The exam work regarding high-temperature synthesis of two-dimensional magnetic materials was rewarding for my perception of chemistry.

Finally, I wish to express my thanks towards Ph. D. student Daniel Petrini for many interesting discussions about undergraduate laboratory experi-ments.

46

Summary in Swedish/Sammanfattning på svenska

Inledning Elektrokemisk deponering innebär att en yta beläggs med ett fast material genom att höja eller sänka energinivån hos laddade partiklar, joner, upplösta i vatten eller något annat lösningsmedel med hjälp av elektricitet. Överföring av elektroner från en yttre strömkrets sker från en elektriskt ledande yta, arbetselektroden, till jonerna i lösningen vars oxidationstal då minskar (detta kallas reduktion). Från en vattenlösning är det möjligt att reducera många joner i det periodiska systemet, exempelvis kobolt-, nickel-, koppar- och zinkjoner. För joner som kräver en högre energinivå för elektronöverföring-en än den för reduktion av vatten används organiska lösningsmedel eller jonsmältor. Det finns två typer av fundamentala problem vid elektrodepone-ring: (i) kinetik och mekanism för deponeringsprocessen, och (ii) kärnbild-ning och tillväxt av den på elektroden stabila fasen.

Elektrokemisk deponering används industriellt för att göra ledare i elek-troniska kretskort, magnetiska material i datorminnen, beläggningar på hård-diskar, nötningståliga och korrosionsbeständiga beläggningar, metalliska kompositmaterial, elektrokroma material, dekorativa beläggningar, oxider och organiska polymerer.

Mitt arbete har gått ut på att förstå ett mycket speciellt fenomen som upp-står vid elektrodeponering från kopparjoner bunda i komplex med vissa or-ganiska molekyler. Spontana och regelbunda svängningar, oscillationer, i den elektrokemiska potentialen uppstår när en konstant ström drivs igenom den elektrokemiska cellen. Det material som då bildas är tunna lager av kop-paroxid emellan tjockare lager av koppar. Under arbetets gång framgick det hur väldefinierade multilager kunde deponeras genom att istället växla mel-lan en låg och hög ström för deponering av kopparoxid respektive koppar. De inbördes tjocklekarna kunde då varieras oberoende av varandra med god kontroll av tjockleken genom att variera tiden för strömpulsen. Kopparoxid är en halvledare och koppar är en god elektrisk ledare där väldefinierade multilager representerar ett nytt material (se figur 1).

Multilager av Cu/Cu2O uppvisar flera intressanta fysikaliska egenskaper. Vid mätning av materialets elektriska egenskaper visar det sig att strömmen minskar då spänningen ökar, ett fenomen som beror på tunnelemission och kallas negativ differentiell resistans. Det är en egenskap som utnyttjas i halv-

47

ledarkomponenter som till exempel dioder. Dessutom gör kopparoxidens halvledaregenskaper att tunna filmer är genomskinliga för synligt ljus. Has-tigheten för ljuset beror av mediet och eftersom ljuset bryts vid flera gräns-snitt uppstår interferens hos det reflekterade ljuset. Färgen ändras i ordning-en blå–gul–grön då tjockleken ökar hos tunna filmer av Cu2O. Lagerstruktu-rer av Cu/Cu2O har intressanta optiska egenskaper och skulle också kunna användas som stomme för magnetiska lagerstrukturer. Genom elektroke-misk deponering kan lagerstrukturen och därmed de fysikaliska egenskaper-na styras med god precision.

Fig. 1 Modell av koppar (underst), kopparoxid (mellan) och lösning med kopparjo-ner bundna till citrat (överst). Figuren genererades med programmet Material Stu-dio, utvecklat av Accelrys Inc., San Diego.

Arbetet med elektrokemisk deponering från lösningar med komplexbundna joner i basisk lösning har också gett en god förståelse för elektrodeponering-en av ett annat material som består av en porös blandning av nanometerstora partiklar av antimon och antimonoxid. Detta material fungerar utmärkt som anodmaterial i ett litiumjon-batteri, då det har nästan dubbelt så hög ladd-ningskapacitet som dagens anodmaterial, grafit. Elektrokemisk deponering från lösningar innehållande me-tall-komplex Ett elektriskt ledande föremål kan beläggas med koppar genom att doppa ned det i en lösning av kopparjoner. När ström går genom cellen sker elektron-överföring från den elektriskt ledande ytan till de positivt laddade kopparjo-nerna i lösning.

48

Organiska molekyler tillsätts ofta vid elektrodeponering av metaller i vat-tenlösningar för att det ska bli en slät och jämn beläggning med god vidhäft-ning. Så används till exempel citronsyra vid elektrodeponering av koppar. Molekylerna binder till jonerna i komplex, ofta med flera funktionella grup-per, och kallas i sammanhanget flertandade ligander. Vid överskott av ligand nedbringas den fria halten av joner i lösningen till ett minimum. Det medför att deponeringshastigheten blir beroende av hastigheten för dissociationen av komplexet.

När jonerna är skyddade av organiska ligander går det att höja lösningens pH utan att det blir fällning av hydroxid i lösningen. Det går att belägga ett föremål med kopparoxid eftersom citronsyra bildar lösliga komplex med kopparjoner i basisk lösning.

Basisk koppar(II)-citratlösning Vanlig citronsyra bildar ett komplex med jonen Cu2+ som kan skrivas [Cu2H�2Cit2]

4�, där Cit står för citrat som är basformen till citronsyra där molekylens karboxylsyragrupper är dissocierade i formen �COO�. I kom-plex med kopparjoner är dessutom hydroxogruppens väteatom dissocierad (avlägsnad) och formen �O� binder till kopparjonen. När jonerna är bundna i ett komplex måste de först frigöras innan elektronöverföring kan ske. I basisk vattenlösning sker elektronöverföring till Cu2+ i två steg: Cu2+ + e� � Cu+, följt av Cu+ + e� � Cu (s). Då kan det hända att hydroxidjoner, OH�, som det finns gott om vid högt pH reagerar med mellanprodukten Cu+ så att kopparoxiden Cu2O bildas. Spontana potentialoscillationer Potentialen talar om hur mycket elektrisk energi som behövs för att upprätt-hålla den konstanta strömmen. Potentialen hamnar oftast på ett konstant värde när flödet av reducerbara joner fram till arbetselektroden är tillräckligt snabb. I en basisk lösning av kopparjoner bundna i komplex till citrat kan potentialen dock börja oscillera med en potentialskillnad på ca 300 mV, se figur 2. En nanoporös beläggning bestående av både koppar (Cu) och kop-paroxid (Cu2O) bildas under dessa potentialoscillationer.

49

Fig. 2. Potentialoscillationer vid konstant ström (tunn linje) och frekvensen (tjock linje) mätt med kvartskristallmikrovågen.

Som en följd av att citrat frigörs från kopparjonen under elektrodeponering-en produceras hydroxidjoner. Den koncentrerade citratlösningen har en stor förmåga att motstå förändringar i pH vid tillsats av bas eller syra. I det här fallet titreras lösningen invid elektroden med bas till dess att buffertkapaci-teten är slut och pH snabbt ökar.

När en konstant ström tvingas genom en elektrokemisk cell beror potenti-alen av den reaktion som hinner med att uppehålla strömmen. Deponerings-hastigheten för kopparoxid ökar med pH och det gör att den konstanta strömmen kan uppehållas genom deponering av kopparoxid en kort stund. Potentialen blir då mindre negativ och hamnar vid en sådan nivå att koppar-oxid är stabil. Vad som händer då är att den elektrodeponerade kopparen reagerar med Cu2+ och bildar Cu2O i en så kallad komproportioneringsreak-tion, där en kemisk förening med ett mellanliggande oxidationstal bildas genom en reaktion mellan två närliggande oxidationstal. Reaktionen som sker spontant gör att pH minskar. Det sker ingen nettoproduktion av hyd-roxidjoner vid deponering av kopparoxid. Deponeringshastigheten för kop-paroxid minskar efterhand då pH invid elektroden minskar. Strömmen uppe-hålls genom deponering av koppar när potentialen åter blir mer negativ. Antimon(III)-tartratlösning Antimon med beteckningen Sb är ett grundämne och dess joner bildar med vinsyra ett kaliumsalt som har det traditionella namnet kräkvinsten. Saltet kan lösas upp i vatten och antimonjonerna är då bundna till vinsyran i ett komplex som kan skrivas [Sb2H�4Tar2]

2–, där Tar står för tartrat som är bas-

50

formen av syran som på engelska heter tartaric acid. I en sådan lösning är förhållandet mellan antalet tartratmolekyler och antalet antimonjoner lika med 1:1. I likhet med citratkomplexen produceras hydroxidjoner när hyd-roxylgruppen protoneras efter det att tartratmolekylen frigjorts från antimon-jonen. Eftersom förmågan att motstå förändringar i pH vid tillsats av syra eller bas är liten i den här lösningen kommer pH att öka snabbt vid depone-ring av antimon. Då bildas oxiden Sb2O3 och deponeras samtidigt med me-tallen.

Utfällning av antimonoxid i lösningar utan överskott av ligand har an-vänts för att deponera svarta, tunna filmer som består av nanopartiklar av antimonmetall och antimonoxid som fungerar bra som anodmaterial i litium-jonbatterier. Lokal utfällning av oxid Metalloxid kan deponeras under betingelser när oxiden är stabil vid arbets-elektroden samtidigt som metalljonerna är stabila i lösningen varifrån depo-neringen sker. Det finns i huvudsak två olika sätt att gå tillväga på. Det för-sta är att bilda lösliga komplex av metalljonerna i en basisk lösning. Lokal utfällning av metalloxid sker då komplexet spjälkas upp i ligand och fri me-talljon vid arbetselektroden. Det kan göras både under anodiska och katodis-ka förhållanden. Det andra sättet att erhålla en lokal utfällning av oxid är att utnyttja en elektrokemisk eller kemisk reaktion för att lokalt höja pH. Re-duktion av vatten till syrgas är möjlig om potentialen på arbetselektroden höjs tillräckligt och ger då hydroxidjoner som biprodukt så att pH ökar.

Nu är det så finurligt att i båda fallen som omtalats ovan � kopparjon-komplexen med citronsyran och antimonjonkomplexen som löses i vatten från kräkvinstenen � är liganderna sådana att när de spjälkas upp från sina respektive metalljoner kommer deras hydroxylgrupper att automatiskt rycka till sig en vätejon. När så sker ökar pH lokalt eftersom spjälkningen av kom-plexen sker vid arbetselektroden där jonerna reduceras. Det gör det möjligt att från en svagt sur eller neutral lösning erhålla en lokal utfällning av ox-iden, tack vare den automatiska pH-ökningen som sker under elektrodepone-ringen. Vad som avgör om det blir utfällning av oxid är förutom pH, förhål-landet mellan koncentrationen av ligand och metalljon respektive om metall-jonen vid reduktion bildar intermediat som lätt bildar oxid.

51

Bibliography

[1] P. C. Andricacos, C. Uzoh, J. O. Dukovic, J. Horkans, and H. Deligianni, IBM

Journal of Research and Development 42:567 (1998). [2] E. I. Cooper, C. Bonhote, J. Heidmann, Y. Hsu, P. Kern, J. W. Lam, M. Rama-

subramanian, N. Robertson, L. T. Romankiw, and H. Xu, IBM Journal of Re-search and Development 49:103 (2005).

[3] D. D. N. Singh and R. Ghosh, Surface and Coatings Technology 201:90 (2006). [4] E. Raub, Müller, K., Fundamentals of metal deposition, Elsevier, Amsterdam,

1967. [5] F. L. R. Bard A. J., Electrochemical methods, John Wiley & Sons, Texas, 2001. [6] E. N. Drakr and J. L. Jones, Analytica Chimica Acta 54:483 (1971). [7] D. F. Shriver, Atkins, P.W., Inorganic chemistry, Oxford university press, Ox-

ford, 1990. [8] K. I. Popov, S. S. Djokic, and B. N. Grgur, Fundamental aspects of electrome-

tallurgy, Kluwer Academic/Plenum Publishers, New York, 2002. [9] M. D. Obradovic, R. M. Stevanovic, and A. R. Despic, Journal of Electroana-

lytical Chemistry 552:185 (2003). [10] F. B. A. Ferreira, F. L. G. Silva, A. S. Luna, D. C. B. Lago, and L. F. Senna,

Journal of Applied Electrochemistry 37:473 (2007). [11] E. Gomez, S. Pane, and E. Valles, Electrochimica Acta 51:146 (2005). [12] M. T. Beck, B. Csiszar, and P. Szarvas, Nature 188:846 (1960). [13] K. I. Popov, Djokic S.S, Grgur B.N, Fundamental aspects of electrometallurgy,

Kluwer Academic/Plenum Publishers, New York, 2002. [14] P. M. Vereecken, R. A. Binstead, H. Deligianni, and P. C. Andricacos, IBM

Journal of Research and Development 49:3 (2005). [15] T. P. Moffat, D. Wheeler, and D. Josell, Journal of the Electrochemical Society

151:262 (2004). [16] I. A. Carlos, C. A. C. Souza, E. M. J. A. Pallone, R. H. P. Francisco, V. Car-

doso, and B. S. Lima-Neto, Journal of Applied Electrochemistry 30:987 (2000). [17] H. Natter and R. Hempelmann, 100:19525 (1996). [18] S. Kim and D. J. Duquette, Journal of the Electrochemical Society 154:195

(2007). [19] S. Kim and D. J. Duquette, Journal of the Electrochemical Society 153:417

(2006). [20] S. M. Mayanna and B. N. Maruthi, Metal Finishing 94:42 (1996). [21] M. Zheng, J. J. Kelly, and H. Deligianni, Journal of the Electrochemical Society

154:400 (2007). [22] P. Poizot, C.-J. Hung, M. P. Nikiforov, E. W. Bohannan, and J. A. Switzer,

Electrochemical and Solid-State Letters 6:21 (2003). [23] I. G. Casella and M. Gatta, Journal of Electroanalytical Chemistry 494:12

(2000).

52

[24] E. W. Bohannan, C. C. Jaynes, M. G. Shumsky, J. K. Barton, and J. A. Switzer, Solid State Ionics 131:97 (2000).

[25] B. E. Breyfogle, C.-J. Hung, M. G. Shumsky, and J. A. Switzer, Journal of the Electrochemical Society 143:2741 (1996).

[26] A. Q. Wang and T. D. Golden, Journal of the Electrochemical Society 150:616 (2003).

[27] M. Chigane and M. Ishikawa, Electrochimica Acta 42:1515 (1997). [28] M. Chigane and M. Ishikawa, Journal of the Electrochemical Society 147:2246

(2000). [29] K. Nakaoka, M. Nakayama, and K. Ogura, Journal of the Electrochemical Soci-

ety 149:159 (2002). [30] T. D. Golden and A. Q. Wang, Journal of the Electrochemical Society 150:621

(2003). [31] K. Nakaoka and K. Ogura, Journal of the Electrochemical Society 149:579

(2002). [32] J. A. Switzer, H. M. Kothari, and E. W. Bohannan, Journal of Physical Chemis-

try B 106:4027 (2002). [33] M. Nobial, O. Devos, O. R. Mattos, and B. Tribollet, Journal of Electroanalyti-

cal Chemistry 600:87 (2007). [34] A. Seshadri, N. R. De Tacconi, C. R. Chenthamarakshan, and K. Rajeshwar,

Electrochemical and Solid-State Letters 9:1 (2006). [35] M. Izaki and T. Omi, Journal of the Electrochemical Society 143:53 (1996). [36] Y. Zhou and J. A. Switzer, Journal of Alloys and Compounds 237:1 (1996). [37] H. Natter, M. Schmelzer, S. Janssen, and R. Hempelmann, Berichte der Bun-

sengesellschaft/Physical Chemistry Chemical Physics 101:1706 (1997). [38] H. Natter and R. Hempelmann, Journal of Physical Chemistry 100:19525

(1996). [39] A. Dierstein, H. Natter, F. Meyer, H. O. Stephan, C. Kropf, and R. Hempel-

mann, Scripta Materialia 44:2209 (2001). [40] M. Nakayama and H. Tagashira, Langmuir 22:3864 (2006). [41] S. T. Chang, I. C. Leu, and M. H. Hon, Journal of Alloys and Compounds

403:335 (2005). [42] J. A. Switzer, E. W. Bohannan, T. D. Golden, C.-J. Hung, L.-y. Huang, and M.

Shumsky, Vol. 451, Materials Research Society, Pittsburgh, PA, USA, Boston, MA, USA, 1997, p. 283.

[43] A. Finke, P. Poizot, C. Guery, and J. M. Tarascon, Journal of the Electrochemi-cal Society 152:2364 (2005).

[44] S. Sakai, S. Nakanishi, and Y. Nakato, Journal of Physical Chemistry B 110:11944 (2006).

[45] S. Nakanishi, S. Sakai, K. Nishimura, and Y. Nakato, Journal of Physical Chem-istry B 109:18846 (2005).

[46] T. Pauporte, A. Goux, A. Kahn-Harari, N. De Tacconi, C. R. Chenthamarak-shan, K. Rajeshwar, and D. Lincot, Journal of Physics and Chemistry of Solids 64:1737 (2003).

[47] F. Erler, C. Jakob, H. Romanus, L. Spiess, B. Wielage, T. Lampke, and S. Steinhauser, Electrochimica Acta 48:3063 (2003).

[48] E. Gomez, A. Labarta, A. Llorente, and E. Valles, Surface and Coatings Tech-nology 153:261 (2002).

[49] C. Bonhote and D. Landolt, Electrochimica Acta 42:2407 (1997). [50] Q. Li, E. C. Walter, W. E. Van Der Veer, B. J. Murray, J. T. Newberg, E. W.

Bohannan, J. A. Switzer, J. C. Hemminger, and R. M. Penner, Journal of Physi-cal Chemistry B 109:3169 (2005).

53

[51] M. Lakatos-Varsanyi, A. Miko, L. K. Varga, and E. Kalman, Corrosion Science 47:681 (2005).

[52] M. Palomar-Pardave, M. Miranda-Hernandez, I. Gonzalez, and N. Batina, Sur-face Science 399:80 (1998).

[53] C. J. Clarke, G. J. Browning, and S. W. Donne, Electrochimica Acta 51:5773 (2006).

[54] D. A. Buttry and M. D. Ward, Chem. Rev. 92:1355 (1992). [55] H. Saloniemi, M. Kemell, M. Ritala, and M. Leskela, Journal of Electroanalyti-

cal Chemistry 482:139 (2000). [56] S. Otani, J. Katayama, H. Umemoto, and M. Matsuoka, Journal of the Electro-

chemical Society 153:551 (2006). [57] E. Gileadi and V. Tsionsky, Journal of the Electrochemical Society 147:567

(2000). [58] M. Kemell, H. Saloniemi, M. Ritala, and M. Leskela, Electrochimica Acta

45:3737 (2000). [59] M. Eyraud, S. Kologo, L. Bonou, and Y. Massiani, Journal of Electroceramics

16:55 (2006). [60] M. T. Neshkova, V. Nikolova, and V. Petrov, Analytica Chimica Acta 573-

574:34 (2006). [61] N. D. Nikolic, Z. Rakocevic, and K. I. Popov, Journal of Electroanalytical

Chemistry 514:56 (2001). [62] R. P. Raffaelle, T. D. Golden, R. J. Phillips, and J. A. Switzer, Nanostructured

Materials 2:175 (1993). [63] R. M. Rynders and R. C. Alkire, Journal of the Electrochemical Society

141:1166 (1994). [64] A. Radisic, P. M. Vereecken, P. C. Searson, and F. M. Ross, Surface Science

600:1817 (2006). [65] M. Dietterle, T. Will, and D. M. Kolb, Surface Science 396:189 (1998). [66] Q. Wu, W.-H. Shang, J.-W. Yan, and B.-W. Mao, Journal of Molecular Cataly-

sis A: Chemical 199:49 (2003). [67] O. E. Husser, D. H. Craston, and A. J. Bard, Journal of the Electrochemical

Society 136:3222 (1989). [68] E. M. El-Giar, R. A. Said, G. E. Bridges, and D. J. Thomson, Journal of the

Electrochemical Society 147:586 (2000). [69] S. Hogmark, S. Jacobsson, and Å. Kassman-Rudolphi, Materialvetenskap, Upp-

sala, 1998. [70] K. Kremmer, O. Yezerska, G. Schreiber, M. Masimov, V. Klemm, M. Schnei-

der, and D. Rafaja, Materialwissenschaft und Werkstofftechnik 38:121 (2007). [71] M. Imafuku, M. Kurosaki, and K. Kawasaki, Journal of Synchrotron Radiation

5:935 (1998). [72] G. Hägg, Allmänn och oorganisk kemi, Almqvist & Wiksell, Uppsala, 1989. [73] A. R. West, Basic solid state chemistry, John Wiley & Sons, Chichester, 1984. [74] N. Tantavichet and M. D. Pritzker, Transactions of the Institute of Metal Finish-

ing 84:36 (2006). [75] J. R. Ferraro, Nakamoto, K. , Introductory Raman Spectroscopy, Academic

Press, San Diego, 1994. [76] B. Bozzini, L. D'Urzo, and C. Mele, Electrochimica Acta 52:4767 (2007). [77] R. M. A. Azzam and N. M. Bashara, Ellipsometry and Polarized Light, North-

Holland, New York, 1987. [78] M. L. Walker, L. J. Richter, and T. P. Moffat, Journal of the Electrochemical

Society 153:557 (2006). [79] E. R. Still and P. Wikberg, Inorganica Chimica Acta 46:147 (1980).

54

[80] S. Rode, C. Henninot, C. Vallieres, and M. Matlosz, Journal of the Electro-chemical Society 151:405 (2004).

[81] P. G. Daniele, G. Ostacoli, C. Rigano, and S. Sammartano, Transition Metal Chemistry 9:385 (1984).

[82] A. Olin and B. Wallien, Analytica Chimica Acta 151:65 (1983). [83] G. Aylward and T. Findley, SI Chemical data, Wiley, New York, 2002. [84] J. van Doornmalen, J. T. van Elteren, and J. J. M. de Goeij, Anal. Chem.

72:3043 (2000). [85] V. Tsionsky and E. Gileadi, Materials Science and Engineering A 302:120

(2001). [86] A. I. Molodov, G. N. Markosyan, and V. V. Losev, Electrochimica Acta 17:701

(1972). [87] A. Survila and V. Uksiene, Electrochimica Acta 37:745 (1992). [88] A. Survila, A. Surviliene, S. Kanapeckaite, J. Budiene, P. Kalinauskas, G. Stal-

nionis, and A. Sudavicius, Journal of Electroanalytical Chemistry 582:221 (2005).

[89] T. D. Golden, M. G. Shumsky, Y. Zhou, and R. A. VanderWerf, Chemistry of Materials 8:2499 (1996).

[90] A. E. Rakhshani, Solid-State Electronics 29:7 (1986). [91] T. Mahalingam, J. S. P. Chitra, J. P. Chu, and P. J. Sebastian, Materials Letters

58:1802 (2004). [92] R. P. Wijesundera, M. Hidaka, K. Koga, M. Sakai, and W. Siripala, Thin Solid

Films 500:241 (2006). [93] A. Survila, S. Kanapeckaite, and A. Surviliene, Journal of Electroanalytical

Chemistry 501:151 (2001). [94] E. W. Bohannan, L.-Y. Huang, F. S. Miller, M. G. Shumsky, and J. A. Switzer,

Langmuir 15:813 (1999). [95] S. Leopold, M. Herranen, J. O. Carlsson, and L. Nyholm, Journal of Electroana-

lytical Chemistry 547:45 (2003). [96] S. Leopold, J. C. Arrayet, J. L. Bruneel, M. Herranen, J. O. Carlsson, F. Argoul,

and L. Servant, Journal of the Electrochemical Society 150:472 (2003). [97] M. Y. Fujita M., Sogoshi N., Nakabayashi S., Kondo A, Ishii M, Trans. Mater.

Res. Soc. Japan 30:1041 (2005). [98] G. Aylward, Findley, T., SI Chemical data, Wiley, New York, 2002. [99] C. B. Beck M. T, Szarvas P., Magyar Kemiai Folyoirat 70:217 (1964). [100]J. Zheng, A. Takeda, and N. Furuta, Journal of Analytical Atomic Spectrome-

try 16:62 (2001). [101]A. L. Pitman, M. Pourbaix, and N. de Zoubov, Journal of The Electrochemical

Society 104:594 (1957). [102]A. M. Bond and A. B. Waugh, Electrochimica Acta 15:1471 (1970). [103]P. Poizot, S. Laruelle, S. Grugeon, L. Dupont, and J. M. Tarascon, Nature

407:496 (2000). [104]I. A. Courtney and J. R. Dahn, Journal of the Electrochemical Society

144:2045 (1997). [105]M. Winter, J. O. Besenhard, M. S. Sphar, and P. Novak, Advanced Materials

10:725 (1998). [106]J. Yang, M. Winter, and J. O. Besenhard, Solid State Ionics 90:281 (1996). [107]H. Li, X. Huang, and L. Chen, Solid State Ionics 123:189 (1999). [108]J. O. Besenhard, M. Hess, and P. Komenda, Solid State Ionics 40-41:525

(1990).

Acta Universitatis UpsaliensisDigital Comprehensive Summaries of Uppsala Dissertationsfrom the Faculty of Science and Technology 336

Editor: The Dean of the Faculty of Science and Technology

A doctoral dissertation from the Faculty of Science andTechnology, Uppsala University, is usually a summary of anumber of papers. A few copies of the complete dissertationare kept at major Swedish research libraries, while thesummary alone is distributed internationally through theseries Digital Comprehensive Summaries of UppsalaDissertations from the Faculty of Science and Technology.(Prior to January, 2005, the series was published under thetitle “Comprehensive Summaries of Uppsala Dissertationsfrom the Faculty of Science and Technology”.)

Distribution: publications.uu.seurn:nbn:se:uu:diva-8186

ACTAUNIVERSITATIS

UPSALIENSISUPPSALA

2007

Documents

Electrochemical Deposition of Nanostructured Metal/Metal ...170636/FULLTEXT01.pdf · Journal of Electroanalytical Chemistry 594, 35 (2006) II. Pulsed galvanostatic and potentiostatic