718.019 biomedical sensor systems (2018): electrochemistry ...€¦ · electrochemistry (exercises...

718.019 Biomedical Sensor Systems (2018):

Electrochemistry (exercises 5+6)

Christoph Aigner, Andreas Petrovic

March 5, 2018

Instiut für MedizintechnikKronesgasse 5/II

8010 Graz

Contents

0 Introduction 1

0.1 Why Electrochemistry? . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10.2 What Is a Potentiostat Good for? . . . . . . . . . . . . . . . . . . . . . . 20.3 How Does a Potentiostat Work (in a nutshell)? . . . . . . . . . . . . . . . 30.4 Electrodes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5

0.4.1 Working Electrode . . . . . . . . . . . . . . . . . . . . . . . . . . 50.4.2 Reference Electrode . . . . . . . . . . . . . . . . . . . . . . . . . . 70.4.3 Counter Electrode . . . . . . . . . . . . . . . . . . . . . . . . . . 80.4.4 Screen Printed and Film Electrodes . . . . . . . . . . . . . . . . . 9

1 Copper Deposition 11

1.1 Goals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11

1.2 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 111.2.1 Corrosion research and corrosion protection . . . . . . . . . . . . 111.2.2 Electrogravimetry and the Electrochemical Series . . . . . . . . . 14

1.3 The Experiments . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 151.3.1 Devices and Equipment . . . . . . . . . . . . . . . . . . . . . . . 151.3.2 Chemicals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 151.3.3 Procedure . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15

2 The Cottrell Experiment and Di�usion Limitation 17

2.1 Goals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 172.2 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17

2.2.1 The Cottrell Experiment . . . . . . . . . . . . . . . . . . . . . . . 172.2.2 Electrochemical Double Layer . . . . . . . . . . . . . . . . . . . . 22

2.3 The Experiments . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 252.3.1 Devices and Equipment . . . . . . . . . . . . . . . . . . . . . . . 252.3.2 Chemicals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 252.3.3 Procedure . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 252.3.4 Disposal of Chemicals . . . . . . . . . . . . . . . . . . . . . . . . 27

3 Cyclic Voltammetry - the Most Used Technique 28

3.1 Goals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 283.2 Instructions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 28

3.2.1 What is a Cyclic Voltammogram? . . . . . . . . . . . . . . . . . . 283.2.2 What Information Can a CV Provide? . . . . . . . . . . . . . . . 32

3.3 The Experiments . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 363.3.1 Devices and Equipment . . . . . . . . . . . . . . . . . . . . . . . 363.3.2 Chemicals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 363.3.3 Procedure . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 373.3.4 Disposal of Chemicals . . . . . . . . . . . . . . . . . . . . . . . . 40

0 Introduction

This script contains exercise 5 (Copper Deposition and The Cottrell experiment) andexercise 6 (Cyclic Voltammetry) of the Biomedical Sensor Systems laboratory course(718.019).

0.1 Why Electrochemistry?

Electrochemistry is a versatile discipline that has applications in various �elds of researchand industry. 1 Corrosion research develops methods to save millions of Euro that areneeded to remove the damage done by corrosion. Electrochemistry delivers the insightsabout the processes causing corrosion and is capable to investigate means against cor-rosion in hours instead of weeks. Although these experiments cannot replace the longterm test under real conditions, they can provide the knowledge for a fast pre-selectionof approaches that should be investigated more closely.Another big �eld of electrochemistry is energy conversion. Batteries, especially recharge-able batteries, have been important for portable technologies for quite some time, butsince the boom of the mobile internet the limitations of batteries have become the limi-tations of our electronic society. The electric car only makes sense if we have batterieswith enough capacity to drive hundreds of kilometers without recharging. Either thator we need another electrochemical device, the fuel cell, to power the cars and reducethe request for fossil oils. In the same context solar panels are important to gain energyfrom a sustainable source, the sun. The electrochemistry of photovoltaics is providingimportant knowledge to improve solar panels and brings the idea to satisfy our energydemand with solar power within our reach.In analytical chemistry electrochemistry makes miniaturized devices possible that delivera quantitative result in the ppb range. These devices are easy to handle point-of-caredevices. Electronic components are very suitable for mass production and as a conse-quence these devices are a�ordable to a broad part of the population. The most knownexample of an electrochemical point-of-care device is the blood sugar test. Easy to han-dle, inexpensive, and improving a diabetes patient's quality of life drastically it is themost successful electroanalytical chemistry product up to today. The blood sugar sensoris based on an enzymatic reaction, that is it is a biosensor.

Gaining an insight into biological electrochemistry is a powerful tool for modern tech-nology. Using enzymes in detection to create a very sensitive and selective sensor isjust one of the many possibilities biomolecules o�er to technology. Biofuel cells createenergy out of glucose or lactate and can thus power an implanted diagnostic device in-side a human body. Imitating the photosystem of plants leads to new developments inphotovoltaics. DNA chips or protein chips allow a high density of recognition elementson an electronic circuit turning it into a lab on chip that can detect quantitative various

1Whole theory acquired from [1]

viruses, bacteria, etc., with a single droplet of sample, e.g. human blood serum. In thelast 10 years electrochemical impedance spectroscopy has been opening new possibilitiesin corrosion and coating research as well as opened the path to more label free methods,that is no chemicals need to be added and the analyte needs no modi�cation to createan electrochemical signal.These are all modern electrochemical applications and additionally classic electrochem-istry applications are still essential like electrolysis or galvanization.This very brief list of technologies, available because of knowledge in electrochemistry,shows how important it is to teach the basic concepts and applications of electrochem-istry during the studies of chemistry, biochemistry, physics, biology, medicine, pharma,chemical engineering, and material sciences. For more detailed information please referto [2], [3] and [4]

0.2 What Is a Potentiostat Good for?

A voltmeter does measure the potential di�erence between two points. To do so thecircuit needs to be closed, but a very tiny current is usually �owing. This way you cansee the potential di�erence between two points with almost no disturbance for the inves-tigated system. Knowing the potential between two points is useful but manipulatingpotentials is even better.This is what a power source does. In every household batteries are the most commonexample of DC power supplies: a power supply having constant non-alternating po-tentials. They are used for �ashlights, mobile phones, clocks, etc. In electrochemicalapplications DC power supplies are used for galvanization (the deposition of metals onconducting materials, often other metals). Another well-known application is the elec-trolysis of compounds. A well-known industrial example is the chloralkali process wheresalt (NaCl) and water (H2O) are split into chlorine (Cl2), hydrogen (H2), and sodiumhydroxide (NaOH).The disadvantage is that you cannot investigate a single electrode and thus a singleevent. The current �ows through the anode (electrode where oxidation happens) andthe cathode (electrode where reduction takes place). So both these electrodes in�uencethe measured current and the current limiting process cannot be determined. This isespecially an issue in analytical chemistry. An electrochemical analysis that can be usedeasily with a power supply is the electrogravimetry. An electrode's weight is determinedand all the metal in a de�ned sample volume is reduced, which leads to the precipitationof the metal on the electrode. The weight is measured again and thus the amount ofmetal in the sample volume will be determined. Although this method works �ne, ithas some disadvantages. The process takes time, e.g. 30 min plus drying and weightingfor copper, nickel or lead oxide. In the given potential window only the analyte metalshould be reduced. Hazardous side products may be produced. The solution needs to beheated and stirred to decrease the concentration polarization and to make the conversionof the analyte complete. The metal layer must stick to the electrode properly, otherwise

a precise weighting of the electrode is not possible. To use other electroanalytical meth-ods a potentiostat is needed.A potentiostat uses three electrodes and a feedback loop to control the potential andmeasure the current �owing at just one of these electrodes, the working electrode. Thepotential will be measured to a �xed reference point and thus a lot of information aboutthe event happening at the working electrode can be gathered.

The potentiostat controls the potential of an electrode while measuring the current �ow-

ing through that electrode.

0.3 How Does a Potentiostat Work (in a nutshell)?

As mentioned before a potentiostat controls the potential of the working electrode andmeasures the current �owing through it.Why not just two electrodes? One of the reasons is that we cannot measure the potentialof the working electrode against a �xed point when we just have two electrodes. Imaginea two electrode system that consists of the already mentioned working electrode and theelectrode, which potential should be our �xed reference point, the reference electrode.We apply a certain potential between these electrodes and an electrochemical reactionhappens at the working electrode, but since the circuit needs to be closed and currentneeds to �ow, a reaction that is inverse to the reaction at the working electrode mustoccur, that is if an oxidation occurs at the working electrode, a reduction must take placeat the reference electrode. If a current �ows at a constant potential, an electrochemicalreaction must happen according to Faraday's law:

Q = n · z · F (0.1)

This equation says that the charge Q �owing through an electrode is proportional to theamount n of a species that took or gave z electrons at the electrode. F is the Faradayconstant and represents the charge of 1 mol electrons. The current I is the charge Q pertime t �owing through the electrode:

I =Q

t(0.2)

The equations' 3.1 and 3.2 combination shows that the current I �owing is connected tothe reaction happening at the electrode via the amount n:

I =n · z · F

t(0.3)

Imagine now that the current is �owing at the reference electrode. At this electrode aspecies' amount of n is converted. This conversion leads to a change of the surface orthe concentration of the solution surrounding the electrode. The Nernst equation shows

a clear correlation between the potential E of an electrode and its surrounding:

E = E0 +RT

zFlnaOx

aRed

(0.4)

E0 is the standard potential of the redox couple Red and Ox. R is the gas constant andT the temperature. The activity of the oxidized and reduced form of the species aOx

and aRed in the surrounding solution is not always easy to predict. This often leads toa simpli�cation of the equation:

E = E0 +RT

zFlncOxfOx

cRedfRed

= E0 +RT

zFlncOx

cRed

+RT

zFlnfOx

fRed

= E0′ +RT

zFlncOx

cRed

(0.5)

The two activity coe�cients fOx and fRed are included in the resulting potential E0′ ,which is called the formal potential. Since it contains parameters that depend on theenvironment, such as temperature and activity coe�cients, E0′ cannot be listed butneeds to be determined for each experiment, if necessary. Most experiments in analyticalchemistry are performed at room temperature (295 K). This makes another simpli�cationpossible. Out of convenience also the ln will be transferred to the log.

E = E0′ +RT

zF2.3log

cOx

cRed

= E0′ +0.059V

zlog

cOx

cRed

(0.6)

For practical application equation 0.6 is the most used form of the Nernst equation. Formany applications one can assume that E0 is roughly the same as E0′ , because bothof the activity coe�cients are close to one. In this form (equation 0.6) the correlationbetween the surrounding of an electrode and its potential is visible more easily.As mentioned before all the simpli�cations at equation 0.4 were performed: The changeof the solution surrounding the reference electrode, due to a �owing current, leads to achange of the potential that is supposed to be our �xed reference point. But we cannotlimit the current �ow through the reference electrode (RE), because all limitations shouldbe caused by the process that we want to investigate, that is the process at the workingelectrode (WE). The solution for this problem is a third electrode. At this counterelectrode (CE), also known as auxiliary electrode, the counter reaction to the workingelectrode's reactions takes place. The current is �owing between the working and thecounter electrode. The potential is controlled between working and reference electrode(see Figure 0.1). The potential between counter and reference electrode is adjusted insuch a way that the current �owing through the working electrode at a certain potentialbetween working and reference electrode is satis�ed. There are limits for the potentiala potentiostat can apply between RE and WE (DC potential range) and CE and WE(compliance voltage). Since you control the potential between RE and WE it is easy tostay within the limitation of the DC potential range. Because the compliance voltagecannot be controlled by the user the CE has to be bigger than the WE. A bigger surfaceat the same potential leads to a higher current . A rule of thumb suggests that the CEshould be 100 times bigger. For many experiments this may not be necessary, but for a

Figure 0.1: A schematic three electrode system

good praxis you should ensure that the CE is big enough so that it does not limit thecurrent �owing at the WE.Usually the distance between CE and WE is big enough so that the reactions do notin�uence each other, and the counter reaction can be ignored, but sometimes, in smallvolumes for example, it can be helpful to know which reaction happens at the counterelectrode.

The potential is applied between reference and working electrode, while the current

�ows through working and counter electrode. This way a constant reference point for

the potential is maintained.

0.4 Electrodes

While the potentiostat is the heart of an electrochemical experiment, modern poten-tiostats need little e�ort to set up and operate. The average electrochemist will spenta lot more time with his or her electrodes. They need to be polished maybe modi�ed,and sometimes it is rather challenging to �x them all in the right place. For this reasonwe will �rst talk about the three di�erent electrodes, before we take a closer look at thepotentiostat again.

0.4.1 Working Electrode

The working electrode is the place where the reaction happens that we want to controlor investigate. As a consequence this electrode should be carefully and reproducibleprepared. The most common working electrodes are disc electrodes. For example, ametal cylinder or a metal wire is surrounded by Te�on or PEEK and the cross section isexposed. The metal disc is connected to a wire at the other end of the coating, so it canbe connected (see Figure 0.2). Common materials for working electrodes are platinumand gold as well as a variety of carbon phases. Very popular due to its conductivityand reusability is glassy carbon. To ensure a clean and smooth surface the electrode

Figure 0.2: A gold disc working electrode in a Kel-F (Te�on) coat; photo and scheme

is usually polished before usage. Very �ne emery or sand paper as well as slurries ofaluminum oxide particles on cloth are used for mechanical polishing. Another way isthe electropolishing. This is used for surfaces that get contaminated very easily. Thehydrophobic character of gold makes it often necessary for gold electrodes to electropolishthem. When residues need to be removed by electropolishing a surface is cleaned byapplying high potentials.

Rough surfaces will lead to high currents due to the charging of the interface. Dirtysurfaces can show artifacts in the measurements or can be the cause that not the fullactive surface has access to the electrolyte. So cleaning an electrode is an importantstep in the electrode preparation.A blank electrode is often just the beginning of the preparation step. For many ap-plications an electrode's surface needs a modi�cation: To produce an ion selectiveelectrode polymer �lms are put on electrodes, to produce a biosensors enzymes orother biomolecules, for example DNA, are immobilized at the electrode, to produceimmunosensors parts of the immune system are immobilized at the electrode and alsoinorganic catalysts are used for electrode modi�cation.A very interesting working electrode is the mercury electrode. Due to its liquid statemercury is used as a droplet electrode. At the end of a tube a mercury droplet isformed and used for the corresponding experiment. To clean the electrode the dropletis enlarged until it drops and a new surface is prepared with clean mercury. This wayclean and easy to reproduce surfaces are obtained, but due to its toxicity in most labsmercury is not common. In general modern electrochemistry avoids mercury, but someexperiments deliver the best results with mercury, so that the mercury electrode is stillused nowadays.

The working electrode is the electrode where the investigated processes occur. It needs

to be carefully prepared, so that the surface is reproducible and known.

0.4.2 Reference Electrode

The reference electrode should deliver a constant potential. A current �owing throughan electrode leads to an electrochemical reaction that will change the composition of theelectrode's environment and thus the potential. As a consequence there should be aslittle current as possible �owing through the reference electrode.To establish and keep a constant potential an electrode of the second type is often chosen.Their potential usually depends indirectly on the concentration of a single anion. Theelectrodes of the �rst type are basically just metal surfaces in an electrolyte. Accordingto Nernst's law (see equation 0.4) the potential is depending directly on the surroundingsolution. An electrode of the second kind is usually a metal surrounded by a hardlysoluble salt of itself. This electrode is then immersed into a solution containing theanion of that salt. Usually the anion has a high concentration to make sure that themetal salt is not dissolving fast and that small changes in the concentration have a lowimpact on the potential. A common reference electrode is the silver / silver chlorideelectrode (see Figure 3.3). A silver wire is coated with a �lm of silver chloride. Thiswire is immersed in a potassium chloride solution.

Figure 0.3: Photo and scheme of a Ag/AgCl/3 M KCl reference electrode

The potential of the silver wire depends on the silver ions dissolved in the potassiumchloride solution.

E = E0′ +0.059V

zlog

[Ag+]

[Ag](0.7)

Due to the fact that it is a solid material the concentration of silver is 1 M. The con-centration of the silver ions depends on the concentration of chloride, as the solubility

product KL (equation 0.8) shows.

KL = [Ag+] · [Cl−] (0.8)

If 0.7 and 0.8 are combined and all constants are summarized in E0′(Ag/Ag+) the resultis:

E = E0′(Ag/Ag+) +0.059V

zlog[Cl−] (0.9)

If the chloride concentration is kept constant a good reference electrode is created. Thisis usually achieved by separating the measuring solution from the solution surroundingthe reference electrode using a porous frit. Even small concentration changes due todi�usion through the frit or evaporation have little impact on the concentration if theconcentration of chloride is high. The frit also prevents silver ions from di�using into themeasuring solution. The silver / silver chloride electrode is one of the two most popularchloride containing reference electrodes. The other one is the calomel electrode. Thisreference electrode contains mercury in contact with mercury(II) chloride in a chloridecontaining solution often saturated with chloride. You should always indicate clearlyversus which reference electrode the potentials were applied or measured. Often you�nd this information in the axis label e.g. �E / mV vs Ag/AgCl�. If the currents aresmall enough, that is in the nA range, it is possible to immerse the wire with the hardlysoluble salt coating in the measurement solution. This is called a pseudo referenceelectrode. The reference and counter electrode can also form a short circuit and you canwork with a two electrode system. If your currents are higher, you want to perform along term experiment, for example overnight, or need a very stable potential, a threeelectrode system with a proper reference electrode is necessary.

The reference electrode keeps the system stable. Fluctuations in the potential result

in noisy measurements. The reference electrode should be checked �rst if the system

behaves in an unexpected way.

0.4.3 Counter Electrode

The counter electrode usually needs the least maintenance. It is just some inert metalwith a large surface. For people who work with many di�erent materials platinum is agood counter electrode due to its inertness towards most solutions. A really big surfacearea is achieved by a mesh structure, but often a platinum wire has already a surfacearea several magnitudes bigger than the working electrode. Platinum is also very easyto clean. A hand torch is usually enough for making the platinum glow and to removeall stains.The counter electrode should have a large surface area, to enable a high current �oweven with low potentials. This way water splitting, gas production, or the production ofaggressive radicals is avoided. If this kind of reaction is happening in a small volume ofsolution, it will no longer be negligible. A large surface area will lead to high capacitive

currents. At the counter electrode capacitive currents provide a current �ow withoutchanging the chemical composition of the solution.

The counter electrode should not limit the current or in�uence the working electrode.

Popular electrodes are platinum wire or platinum mesh electrodes.

0.4.4 Screen Printed and Film Electrodes

Three properties make electrochemical detection very interesting for analysis in everydaylive, especially point-of-care measurements. The measuring devices as well as electrodescan be miniaturized, which makes them really portable. The results are quantitative.And produced in large quantities electronics get very cheap.However, a product that can be used by people who have no electrochemical educationmust skip all the demanding preparation and cleaning steps to reuse an electrode. Thisis why in recent years disposable test sensors have become more and more important.A gold disc in a Te�on coat is too expensive for a disposable sensor. Another groupof electrodes is more suitable for this objective. Flat electrodes that are deposited ona surface are the up-to-date commercial biosensors. These are again divided into twogroups: �lm electrodes and printed electrodes.Film electrodes are made by metal deposition on a surface that is covered with somekind of mask. The mask only exposes the parts of the metal that should be �lledwith metal afterwards. The metals are deposited by heating up the metal source in avacuum while the masked sample is present. If di�erent metals need to be on the samesensor, because a gold WE and a silver RE are requested, di�erent masks and cleaningprocesses need to applied to ensure that the gold and silver will not be deposited onthe same surface. As a consequence most sensors produced this way have WE, CE,and RE from the same metal. The advantage of these electrodes is that the masks arecreated by photolithography and very small structures can be created. For example,DNA chips with dozens of working electrodes in a 1mm2 or interdigitated electrodescan be produced. Interdigitated electrodes are always pairs. Each electrode looks like acomb and they are arranged in such a way that each tooth is between two teeth of theother electrode. With this setup species are cycling between the two electrodes and thisbehavior creates an ampli�cation of the signal.

The main di�erence with screen printed electrodes is the material the electrodes aremade of. While �lm electrodes are made from pure metals, screen printing uses pastesor inks to create a sensor. A mesh with a mask is placed above the empty support, apolyester sheet for example, without touching the support. The ink is put on the meshand a squeegee is moved across the mesh to press the paste through the template aswell as to create contact between the mesh and the support. The paste remains on thesupport. A drying step follows. The drying step can include heating, especially for waterbased inks. Each di�erent paste also needs a di�erent mask, same as with the thin �lmmethods, but since no vacuum is necessary the price for the needed equipment is lowerand depending on the drying step the procedure can be faster.

The inks usually contain the conducting material, for example the metal particles orcarbon compounds, some dispersion agents as well as other additives, polymer bindersand the solvents. For the results the ingredients and their ratios are very important.The results are usually rougher and have lower conductivity than classic electrodes, butthe production is cheaper, especially in mass production. Due to the fact that it is easierto use di�erent materials than while producing �lm electrodes, screen printed electrodeshave often WE, RE, and CE made from di�erent pastes.Counter electrodes are made of carbon pastes, because they are cheaper than metalpastes. Reference electrodes can be made of silver paste mixed with silver chloride, sothat a stable reference electrode is achieved.

The price per piece for a �lm or screen printed electrode decreases with the number ofproduced electrodes, due to the fact that the most expensive parts are the masks thatneed only to be produced once. The low price allows these electrodes to be producedas a one-time use or disposable electrode. As a consequence any steps for cleaning theelectrode after a measurement can be skipped, because the electrode is disposed. If theelectrode has a dedicated application, most preparations can already be done by theproducer. All this makes it possible to develop electrochemical applications that canbe operated by people with no electrochemistry background. The best example is theblood sugar measurement, which is used by diabetes patients.

These electrodes that are suitable for mass production are versatile, cheap, and easy to

handle. They make electrochemistry available to users without electrochemical knowl-

edge.

1 Copper Deposition

1.1 Goals

Most students already know the basics of electrochemical potentials and how they areapplied in batteries or electrolysis. This experiment is meant to facilitate a transitionfrom what is taught in most schools and the slightly more advanced level of potentiostaticexperiments. Because galvanization for depositing metals or other protective layers isa widely used technique in industry as well, some basic knowledge of this importantprocess should be gained. The experiment and its discussion should teach:

� A general idea about corrosion protection

� How can the di�erent standard potentials be exploited for metal deposition?

� How do electrodeposition paints work?

� How does electrogravimetry work?

1.2 Introduction

1.2.1 Corrosion research and corrosion protection

Corrosion is the destruction of materials by chemical reactions with the environment. Itis a big cost factor. Replacement of corroded materials in buildings, bridges, or monu-ments and painting or galvanization of building elements is a huge investment into main-tenance. The study "Corrosion Costs and Preventive Strategies in the United States"showed that in 1998 corrosion cost the U.S. 276 billion$. This is ca. 3.2 % of the USgross domestic product1. Painting the Ei�el tower costs ca. 4 million e every 7 years2.Without a proper and closed layer of paint this monument would become unstable andbreak down. As a consequence one of the most important �elds in electrochemical re-search focuses on corrosion. A lot of companies fund academic research for investigatingthe mechanism of corrosive processes and for creating new methods for corrosion protec-tion. Companies check paints and lubricants for their corrosion protection capabilitiesand develop new methods for creating coatings or creating non-corrosive environmentsfor machinery, for example �ooding the entire gears with oil. From a chemical pointof view protection by ordinary paints is easy to understand. The paint is usually amixture of a polymer, color pigments, and various additives for adjusting the physicalproperties of the paint. For example these additives create smooth or rough surfaces orgrant UV-light protection. The paint is usually dispersed or dissolved in some solvent

1Gerhardus H. Koch, Michiel P.H.Brongers, Neil G. Thompson, Y. Paul Virmani and Joe H. Payer.Corrosion Costs and Preventive Strategies in the United States - report by CC Technologies Labo-ratories, Inc. to Federal Highway Administration (FHWA), September 2001

2http://www.tourei�el.paris/en/everything-about-the-tower/themed-�les/97.html

(water or an organic solvent). The solvent evaporates after the paint has been appliedand a polymer with the additives and pigments stays on the painted surface. The resultis a physical barrier between the surface and its environment. This barrier protects thesurface from corrosion. A crack in the layer of paint will grant access to the surface andcorrosion may occur. Here already electrochemistry can help improve the protectionlayer. Electrodeposition paints are paints that are made insoluble by an electrochem-ical reaction. A conducting work piece is dipped into the paint bath and a potentialis applied. The created reaction close to the work piece's surface will make the paintinsoluble and thus a layer of paint precipitates on the metal surface. All conductingparts will be covered with the paint, even usually hard to reach places. If the paintforms an insulating layer, the reaction will stop itself by blocking the conducting surfacecompletely; only the needed amount of paint will be used. These paints have di�erentmechanisms. The most common one is neutralizing charges. The polymers in the paintshave some charged functional groups. These groups make them water soluble. Thesegroups can be deprotonated acids or protonated bases. By way of illustration we assumea polymer with deprotonated carboxylic groups (COO-) is used. Due to its negativecharge it will migrate towards the positive charged electrode, the anode. The potentialat the anode is high enough for water to split as follows:

2H2O − > 4H+ +O2 + 2e− (1.1)

The high concentration of H+, which implies a low pH value, close to the anode leadsto a protonation of the COO- groups. As a result the polymer's COOH groups haveno charge and the polymer is not soluble in water anymore. The polymer precipitateson the anode (see Figure 1.1). After completion of the deposition excess paint is rinsedaway and the paint layer is cured. This means that the paint is heated up, so it can �owand form a uniform surface.

Figure 1.1: electrodeposition paint principle shown for a gold disc electrode

Corrosion protection becomes easier if the metal itself forms a protective layer. Alu-minum oxide layers on aluminum are impenetrable for oxygen and thus protect thealuminum layer beneath. This has the advantage that damaged parts of the coating willbe replaced by the aluminum itself. The thickness of the layer can be increased elec-trochemically, leading to a better protection. Up to now the mentioned coatings werepassive coatings. Active coatings can be created by covering metals with metals. Thisis usually done by dipping the metal to protect in a solution with a salt of the coatingmetal. A potential is applied between the metal that needs protection and an electrodemade from the coating metal. A negative potential is applied to the core metal and allpositive metal ions (cations) will migrate to the core metal that acts as cathode, theelectrode where reduction takes place, and the ions are reduced to elemental metal atthe cathode. The other electrode is oxidizing its own metal and releasing it as ions intothe solution. This electrode is called anode, the electrode where oxidation happens (seeFigure 1.2).

Figure 1.2: the deposition of metals (galvanization)

Due to knowledge of the electrochemical series one can predict if a certain metal is noblerthan another one and thus predict its behavior in contact with it. Noble metals are usefulto create a coating that is inert to corrosion, but if the coating gets damaged and theless noble metal is exposed, a local element is created. The noble coating metal will takeelectrons from the less noble core metal as a result the speed of corrosion is increasedinstead of slowed down. An example of this would be chromium coated steel. Coatingwith a less noble metal that passivates itself is usually a better idea. In case the coatingis damaged and the core metal is exposed the coating will act as a sacri�cial anode.The noble core metal will drain electrons from the less noble coating and thus the corewill reduce itself, while the coating is oxidized. An example of this is zinc coated steel.There are more ways to protect a product from corrosion, but within the framework ofthis experiment the knowledge of these most important methods will su�ce.

Corrosion is an important research �eld due to the fact that corrosion causes big

amounts of damage to machines and monuments. Understanding and preventing cor-

rosion is the goal of corrosion research.

1.2.2 Electrogravimetry and the Electrochemical Series

The deposition of metal is not only used to modify a surface for its optical appearance,haptics, or corrosion protection; this technique can also be used for analytical purposes.An electrode with a known mass is immersed in a solution containing the metal ions forquanti�cation, the analyte. A potential is applied and the metal ions are deposited asmetal or metal oxide �lms. Which potential is applied and how the metal is deposited de-pends on the metal. Copper, for example, can be reduced at the electrode and depositedas elemental metal. It forms a properly attached �lm on platinum electrodes. This �lmcan be washed, dried, and scaled. Lead ions are oxidized for deposition. A �lm of leadoxide is formed that sticks well to a platinum electrode. This �lm can be washed, dried,and scaled as well. The composition of lead oxide (PbO2) is known and so the amount oflead in the �lm can be calculated. This technique for quanti�cation demands clean andcareful working. Damage of the �lm, remains of water or other solvents during scaling,un�nished deposition of the metal, or a wrong mass of the electrode without the �lm willlead to wrong results. To reduce the time needed for deposition the metal solution needsto be heated and stirred during the deposition. It seems easier to measure the chargethat passes through the electrode during the deposition, but often side reactions andcapacitive current appear during these processes and these will contribute to the chargepassing through the electrode. The electrochemical series show that di�erent processeshappen at di�erent potentials. For example, copper is deposited at less negative poten-tials than zinc, and cadmium is easier oxidized than lead. Of course these insights arealso true for other redox active species than metals: Chloride will be oxidized to chlorineat lower potentials than water will be oxidized to oxygen. These di�erences can be usedto determine di�erent molecules from the same solution, if the di�erence is large enoughfor these processes to be regarded as separate. A classic example is the determinationof copper and nickel from a mixed solution of copper and nickel salts.

Deposition of metals can also be used for analytical methods. The di�erences in the

standard potentials allow individual precipitation of di�erent metals in the same solu-

tion.

1.3 The Experiments

1.3.1 Devices and Equipment

� EmStat3 or PalmSens3 Sensor

� sensor cable

� computing unit (PC, Laptop, notebook, tablet PC (android))

� potentiostat software (PStrace4)

� calculation and plotting software (Excel, MatLab)

� counter electrode

� reference electrode

� working electrode

� a few metal paper clips or a metal wire

� retort stand

� retort clamp

� beaker (electrochemical cell)

� stirrer and magnetic stirring bar

1.3.2 Chemicals

� 18 mM CuSO4 in 0.5 M H2SO4 solution (prepared from solid CuSO4)

1.3.3 Procedure

Deposition of Copper from a Copper Sulfate Solution

1. Set up the retort stand including the clamp. Insert the white cell top with holesinto the clamp. Put the stirrer under the clamp and 50 mL beaker with a magneticstirring bar between them. Bend the paper clip into a straight piece of wire. Mostpaper clips have an insulating coating. Use some emery paper to remove it. Theactual shape of the paper clip wire is not important for the experiment, but astraight wire is easy to handle. Insert the counter electrode (platinum wire) andthe reference electrode, the silver / silver chloride electrode (the electrode with theliquid �lled glass body), into two of the cell top's holes. Connect the paper clipto the working electrode cable (red) with the crocodile clip. Connect the counterelectrode to the black cable and the reference electrode to the blue cable. Pleasebe aware that the part of the reference electrode containing the diaphragm alwaysneeds to be in the solution or otherwise protected from drying. During evaporationof solution little crystals form in the pores of the frit and the electrode will beblocked. This should be avoided.

2. Fill the beaker with the copper salt solution. Position the electrodes in such a waythat all three electrodes are in the solution. The paper clip wire should not beimmersed deeper than the platinum wire.

3. Start PSTrace.

4. Choose the scienti�c mode in the upper left corner.

5. Select your potentiostat from the drop down menu. If it is not in the list, pressthe refresh button (green arrows). Maybe your computer takes a while to installthe USB driver. Press the Connect button. Your device should be connected now.

6. Choose Amperometric Detection from the Technique list. In Sample and Sensoryou can add comments for your own data organization. Set t condition and t de-position to 0, because the electrode needs no treatment before the measurement.The t equilibration is the time during which the �rst potential of the measure-ment is already applied without recording the current. This is usually done toexclude capacitive current from the beginning of the graph, but since we are onlyinterested in the deposited metal and not in the measurement set this time to 0s. The potential applied during the measurement (E dc) has to be low enough toreduce copper but high enough to avoid zinc deposition. The electrochemical seriescontain the standard potentials for copper and nickel. They are E0(Cu2+/Cu) =0.345 V and E0(Zn2+/Zn) = -0.760 V. These values are the standard potential,which is relative to the potential of the standard hydrogen electrode (SHE) that isby convention 0 V. The reference electrode supplied with this kit is an Ag/AgClelectrode. It has a potential of E0(AgCl/Ag) = 0.205 V. So a potential of 0 V vsthe SHE is a potential of -0.205 V for the Ag/AgCl electrode. The other valuescan be treated analogously:E0(Cu2+/Cu) vs Ag/AgCl 0.140 V

E0(SHE) vs Ag/AgCl -0.205 V

E0(Zn2+/Zn) vs Ag/AgCl -0.965 V

So to deposit copper the potential has to be lower than 0.140 V but higher than-0.965 V. Two other species might interfere with our experiment. Potentials below-0.205 V might lead to hydrogen evolution depending on the overpotential of hydro-gen at your paper clip surface. Since experience tells that the large overpotentialsfor hydrogen evolution lead to an active evolution at around -1.2 V vs Ag/AgCl,hydrogen should be no issue. Most paper clips contain iron and E0(Fe2+/Fe) is-0.41 V (SHE) or -0.605 V vs Ag/AgCl. A potential above this value may lead tocorrosion of the paper clip itself. Assuming that no other species in our solutionwill react, an Edc of -0.5 V vs Ag/AgCl should su�ce and avoid gas bubbles fromhydrogen. The t interval is the time between two measurement points. Set it to 1s. The parameter t run is the time the experiment lasts or in this particular casehow long the potential is applied. Set it to 60 s.

7. Switch on the stirrer, adjust the speed to a gentle stir and start the measurement.Observe the paper clip.

8. After the measurement is �nished you can take out the paper clip and rinse it withdemineralized water to observe it more closely.

9. The solution can be reused for multiple paper clips. You can vary the depositiontime and observe the in�uence on the result. Or try to immerse an uncoated partof the paper clip without applying a potential.

10. If you have to dispose the solution keep the following in mind: Solutions containingcopper or zinc need to be collected in a container for liquid waste containing heavymetal. The pH of the collected heavy metal solution should be alkaline (pH > 9).The container should be disposed according to local laws. The electrodes and cellshould be cleaned with demineralized water. Otherwise the drying solution willleave salt stains.

2 The Cottrell Experiment and Di�usion Limitation

2.1 Goals

Electrochemically the Cottrell Experiment is very easy. It is easy to perform and theelectrochemical process is easy to understand. However, the physical observations duringthe experiment are important for many electrochemical techniques and experiments. Theimportance of mass transport, especially di�usion, as limitation for a �owing current andthe additional current �owing due to charging of the electrochemical double layer arebasic principles and in�uence almost every electrochemical experiment. The experimentand its discussion should teach:

� What is capacitive current and what is its in�uence?

� How does mass transport limit the current during electrochemical experiments?

� What is the Cottrell equation?

� How can a Cottrell experiment be used for analysis?

2.2 Introduction

2.2.1 The Cottrell Experiment

Simple experiments often lead to intensive theoretical observations, when performedunder real conditions. Suddenly temperature, pressure, etc. in�uence the outcome ofthe experiment. The Cottrell experiment is one of the most basic potential control

experiments and fortunately well understood. The basic idea is to keep an electrode ata potential that won't lead to any reaction at the electrode and when a stable state isreached a potential step is made that will lead to a chemical reaction. In detail thismeans the potential E of the electrode is before the potential step below the formalpotential E0′ of an electrochemical active species. This species is present in the solutionsurrounding the electrode. The potential before the potential step is named E1 in Figure2.1. In the following pictures and explanations it is assumed that the species is in itsreduced form (Red). The concentration of Red is c∗Red everywhere before the potentialstep, meanwhile the concentration of the oxidized form c∗Ox is 0. The potential step israising the potential so that

E2 > E0′ (2.1)

results. According to the Nernst equation (see equation 0.6) the ratio Ox/Red willbecome bigger. Red is consumed by the electrode by accepting an electron, that is Redis oxidized, and the oxidized form (Ox) is produced (see Figure 2.1). If E2 is signi�cantlyhigher than E0′ , c∗Red is 0 at the electrode surface and all Red close to the electrode isconsumed.

Figure 2.1: schematic potential step of a Cottrell experiment

The measured current after the potential step is a result of Red's oxidation to Ox. Thecurrent will �ow as long as Red is consumed. Before the potential step the lack ofpotential as driving force was limiting the reaction and by increasing the potential itis no longer the limiting factor for this reaction. Since all molecules of Red reachingthe electrode are consumed immediately, the mass transfer of Red towards the electrodebecomes the limiting factor. The species Red is completely depleted in front of theelectrode but in a distance x1 from the electrode surface the concentration is still c∗Red

(see Figure 2.2a).

This concentration gradient leads to di�usion of Red towards the electrode surface. Thedi�usion will also deplete all the Red in the distance x1, so Red from a distance x2 willdi�use to x1. This means that the layer of liquid in which di�usion occurs, the di�usion

Figure 2.2: a) concentration of Red cRed vs the distance from the electrode surface xduring di�erent times of the Cottrell experiment t0 < t1 < t2 < t3 < t4; b)same as a) but only t1 and t4 are shown with a schematic �ux

layer, is growing from the electrode into the bulk (see Figure 2.2a). There are limitsto the growth of the di�usion layer. For usual macroelectrodes, that is electrodes witha diameter above 25µm, the limit is reached when the di�usion layer is disturbed byconvection, that is movement of the liquid. This can be self-induced or due to stirring.In a stirred solution a di�usion layer will have a smaller �nal thickness than in a non-stirred one.

The current measured is always depending on the �ux J of molecules towards the elec-trodes. Actually the current I is proportional to the �ux J.

I ∝ J (2.2)

If the �ux J is known the current I can also be calculated. The �ux is described byFick's law

J = D∂c(x, t)

∂x(2.3)

which shows that the �ux J is proportional to the gradient or in a one dimensional caseslope of the concentration curve ∂c/∂x with the di�usion coe�cient D as proportionalfactor. D is a property and thus a �xed value for a molecule, ion, or complex in amedium. In literature one can �nd tables for D, but one should take care that this valueis also referring to the right medium.Directly after the potential jump in the Cottrell experiment the concentration of Redis 0 at the electrode surface only. With increasing time the di�usion layer grows intothe solution and the resulting gradient is less steep. As a result the �ux J is gettingsmaller and smaller with time (see Figure 2.2b). To describe the change of I with time,the change of J or c with time t is needed. The time di�erential of the concentration isdescribed by Fick's second law:

∂c

∂t= D

∂2c

∂x2(2.4)

Fick's laws are treating general di�usion; solving it for the particular case of the Cottrellexperiment requires some conditions for the initial state and the boundaries. Theseconditions are:

1. Before the potential step the concentration is c∗Red everywhere. cRed(x, 0) = c∗Red

2. After the potential step the concentration at the electrode surface is 0. cRed(0, t) = 0

3. In an in�nite distance from the electrode the concentration is always c∗Red. cRed(∞, t) = c∗Red

Using these conditions to solve Fick's law leads to the Cottrell equation:

I = zFAc∗√D

πt(2.5)

The z is the number of electrons transferred per molecule, F is the Faraday constant,and A the area of the electrode. This equation can be applied to many processes inelectrochemistry when depletion of observed species starts and the reaction gets di�usioncontrolled. As we now know equation 2.5, we expect a current responding to the potentialstep as seen in Figure 2.3 a. There will be a current step with the potential step andafterwards the currents decays with t−

12 . In most real life experiments at a certain time

Figure 2.3: a) current response to the potential step of Figure 2.1, b) current responsein the linear form of a)

the change of current becomes negligible and the current can be considered constant.A reason for this is that the di�usion layer cannot grow in�nitely. At some point thedisturbances and movement of the liquid will destroy the outer layers of the di�usionlayer.Although this seems to be a simple experiment with little practical value, it is actuallystill used to determine di�usion coe�cients or the electrode's surface area. This is doneby turning the I-t-curve into a linear form. From the Cottrell equation it is clear that Ivs t−

12 should lead to a linear plot with a slope

m = zFAc∗√D

π(2.6)

that can be used to calculate A, c∗, D, or z, if the other ones are known (see Figure2.3b). If the experiment is performed and the linear plot is made, it often does not showa perfect line, especially in the beginning of the curve and at the very end. The end ofthe curve is not perfectly linear since the di�usion layer does not expand perfectly intoin�nity as discussed before. The beginning of the curve is in�uenced by two factors.First, it is impossible to perform an ideal potential step. Potentiostats have rise timesin ns or even µs, that is they can reach a set potential in that time period, but this isstill not perfectly instantaneous. So the beginning of the curve deviates from the idealpotential step. The second factor, the capacitive current or capacitive charging current,will be explained in the next paragraph.

During the Cottrell experiment a potential step across a redox potential is applied and

the current caused by the electrochemical reaction is recorded. The current is controlled

by di�usion.

2.2.2 Electrochemical Double Layer

As soon as an electrode surface is charged, due to a potentiostat or its Nernst potential,an electric �eld is created. Charged particles will move in this �eld. Ions of the othercharge than the electrodes charge will accumulate directly at the electrode, forming alayer of ions. Assuming that the electrode is positive charged, anions will accumulate.Attracted by these anions kations will be attracted and form another loose layer on top ofthe �rst layer. The �rst layer is known as the outer Helmholtz plane. The charges accu-mulating in the metal surface are the inner Helmholtz plane (see Figure 2.4). The layerof ions and the electrode act like a capacitor and this has impact on most electrochem-ical techniques. This is only a rudimentary description of the electrochemical doublelayer; there are more sophisticated models, but for many electrochemical experimentsthe simple model will su�ce.

Figure 2.4: scheme of the electrochemical double layer

Up to now it was assumed that the electrochemical active species Red is transported bydi�usion or convection only, but there is one more way for mass transport: migration.Migration is mass transport due to an electric �eld. If Red is negatively charged, thepositive potential of the electrode will attract the Red-ions.Why are complete models and observations done without taking migration into consid-eration? Usually an electrochemical measurement is done in a well conducting solution.Since the current �ows from the working to the counter electrode through the solution, ahigh solution resistance will make the measured current smaller and increase the Ohmicdrop, the di�erence in the potential applied between the reference electrode and theworking electrode and the potential that the working electrode is feeling. To reduce theresistance of a solution an electrochemical inert support electrolyte is added. Often thebu�er itself in a pH bu�ered solution is su�cient or a salt with a high solubility, forexample KCl, NaCl, NaSO4, NaNO3, is added. If the support electrolyte has a highconcentration compared to the investigated species, in this example Red, the electric�eld will be compensated by the ions of the support electrolyte and almost only thesewill migrate. A rule of thumb is that the support electrolyte should have a hundred

times higher concentration. Since this e�ect is suppressed so easily, migration is oftennegligible.Another e�ect is the capacitive charging current or short capacitive current. As men-tioned the electrochemical double layers acts as capacitor. Capacitors store charge. Asimple capacitor is the plate capacitor. It comprises two conducting parallel plates thatare not in contact with each other. If a power source is connected to the plates, a cur-rent �ows that is exponentially decaying until it is insigni�cant. A current �ows becauseone plate is charged negative and the other positive. The separation of charges meanscurrent �ows. At some point the plates cannot store more charge and the current stops�owing. The current decays over time according to

I =Ec

Re−

tRC = I0e−

tRC (2.7)

Ec is the charging potential or voltage, I0 is the starting current, R is the resistance ofthe circuit around the capacitor, t the time and C the capacity of the capacitor. Thecapacity is a property of the capacitor and is de�ned as the charge Q that can be storedper applied potential E or as equation

C =Q

E(2.8)

Usually U is used for voltage, but since these equations need to be transferred to elec-trochemical experiments, it is useful to start with the potential E instead of the voltageU . These two are not synonymous but in this context it is �ne to exchange them. If itis assumed that the electrochemical double layer behaves exactly like a plate capacitor,the two equations 2.7 and 2.8 show three important facts:

1. The capacitive current decays exponentially with the time t. The higher resistanceR and capacity C are the slower it will be decaying. The product of resistance Rand capacity C is often called the time constant τ .

2. The charge Q that can be stored is proportional to the applied potential. Everytime the charge Q that can be stored changes a current I �ows until the chargeQ is adjusted. The charge Q that can be stored changes, if the potential E ischanging. This is expressed in the equation

I =∂Q

∂t= C

∂E

∂t(2.9)

3. In equation 2.8 it is shown implicitly and in 2.9 explicitly that the higher thecapacity C is the more capacitive current will �ow if the potential changes.

Usually electrochemists are interested in the Faraday current, that is the current causedby an electrochemical reaction; the capacitive current, caused by physics, is an unwantedside e�ect. What does this mean for measurements? If the potential of the electrode ischanged, for example during a potential step, a current will �ow that has no chemical

but only a physical meaning. This current decays exponential with t, while the Faradaycurrent decays with t−

12 . This means that the capacitive current decays much faster

than the Faraday current (see Figure 2.5). The higher the capacity C the higher thecapacitive current. The capacity C for a plate capacitor can be calculated with

C = ε0εrA

d(2.10)

where ε0 is the electric �eld constant, εr is the relative permittivity of the mediumbetween the plates, d is the distance between the two plates and A is the surface areaof the two plates. So most factors in�uencing the capacity cannot be altered in anelectrochemical experiment. The constant ε0 cannot be changed. The distance d andthe relative permittivity εr can only be changed by changing the solution, because d isde�ned by the distance between inner and outer Helmholtz plane (see Figure 2.4). Thearea A is in�uenced by the surface roughness. The rougher a surface the higher its area.If a reusable electrode is used, a proper polishing that leads to a smooth surface canreduce capacitive current drastically.

Figure 2.5: scheme of the capacitive and Faraday current through time

The electrochemical double layer acts as a capacitor and every change in the potential

of the electrode will induce a capacitive charging current that is caused by physics not

by a chemical reaction. This current decays exponentially.

2.3 The Experiments



� sensor cable







� retort stand

� retort clamp

� cell top


2.3.2 Chemicals

� 0.1 M KCl solution (prepared from solid KCl and demineralized water)

� 1 mM K3[Fe(CN)6] in 0.1 M KCl solution

2.3.3 Procedure

Capacitive Current

1. Set up the retort stand including the clamp. Insert the white cell top with holesinto the clamp. Insert the counter electrode (platinum wire), the working electrode(platinum disc in Te�on) and the reference electrode, the silver / silver chlorideelectrode (the electrode with the liquid �lled glass body), into the three holes ofthe cell top. Please be aware that the part with the diaphragm of the referenceelectrode always has to be in solution or otherwise protected from drying. Duringevaporation of solution little crystals form in the pores of the frit and the electrodewill be blocked. This should be avoided.

2. Fill the beaker with 0.1 M KCl solution. Position the electrodes in a way thatall three electrodes are immersed and �x the position with the help of the clamp.Connect the counter electrode to the black cable, the reference electrode to theblue cable, and the working electrode to the red cable.

3. Start PSTrace, choose the scienti�c mode in the upper left corner and select yourpotentiostat from the drop down menu. If it is not in the list, press the refreshbutton (green arrows). Maybe your computer takes a while to install the USBdriver. Press the Connect button. Your device should be connected now.

4. Go to the menu Method - Load method and choose the �le Exp1_1.psmethod.

5. Start the measurement. And save the curve by choosing Save curve ... from theCurve menu.

6. With this curve you have got the data in respect to the capacitive current of theelectrochemcial double layer.

The Cottrell Experiment

1. The setup is the same as in the previous experiment. Except for the solution inthe beaker which is now 1 mM K3[Fe(CN)6] in 0.1 M KCl solution.

2. Use the same parameters for the Multistep Amperometry as before. (Load �leExp1_2.psmethod)

3. Before you start the measurement choose in the drop down menu next to the Runbutton the option Overlay instead of New. If the curve of the capacitive currentyou recorded before is no longer displayed load it. Start the measurement. Thereaction at the electrode is: [Fe(CN)6]

3− + e− → [Fe(CN)6]4−

4. Save the new curve as well.

5. The capacitive current as well as the sum of the capacitive and Faraday currentcan be compared with each other. They should meet the expectations that wereset during chapter 2.2.2.

6. The next step should be a better time resolution. Look at the maximum currentthat was measured during step 3. Choose the smallest current range which valuedoubled is bigger than the maximum current. The part of the curve before thestep shows when a stable current is reached. For the equilibration time choose atime that is big enough to reach a constant current. Choose for the t1 5 s and t265 s. Reduce the tinterval to 0.001 s.

7. Choose New in the drop down menu next to the start button and start the mea-surement.

8. Save the curve.

9. To �nd the area of the electrode a plot of the absolute I vs t−11 and a linear

regression are needed. Export the data to software that can do this, for exampleExcel, MatLab, etc. If you decide to use Excel, you can use the one-click-to-Excel-export in the toolbar at the left side of the graph.

10. Set the time of the potential step to 0 s by subtracting 5 s from the time. After-wards calculate t−

12 and plot I vs t−

12 . According to equation 2.5 the resulting plot

should be a line (see Figure 2.3b).

11. Make a linear �t of the linear part of the curve. Calculate from the slope the areaof the electrode. Remember that the slope is described by equation 2.6. AlthoughD can be found in literature, it is not easy. We recommend to use D([Fe(CN)6]

3

in 0.1 M KCl) = (7.17± 0.18)10−6cm2/s found in [5].

12. Discuss the di�erence between expectations and measured values. What couldexplain a deviation from the expected values?

2.3.4 Disposal of Chemicals

� The potassium chloride solution is harmless and can be disposed in the sink. Theelectrodes and cell should be cleaned with demineralized water. Otherwise thedrying solution will leave salt stains.

� Solution containing iron need to be collected in a container for liquid waste con-taining heavy metal. The pH of the collected heavy metal solution should bealkaline (pH > 9). The container should be disposed according to local laws.Those containers will be provided from your laboratory supervisors.

3 Cyclic Voltammetry - the Most Used Technique

3.1 Goals

Moving from passive potentiometric experiments to potentiostatic experiments by con-trolling the potential was an important development. However, the step that followedtowards potentiodynamic experiments may have been even more important for mod-ern electrochemistry. Potentiodynamic experiments made it easy to collect all the dataneeded for a plot of current I versus potential E. These plots are called a voltammo-gram and the technique used for measuring is called voltammetry. In a short period oftime the cyclic voltammetry (CV) provides a lot of information and allows kinetic in-vestigations. It is by far the most used technique by PalmSens customers. Experiencedelectrochemists read quite some information from the shape of a CV. The experimentand its discussion should teach:

1. What is a cyclic voltammetry?

2. Why does a cyclic voltammogram show a speci�c shape?

3. What does the cyclic voltammogram say about reversibility?

4. How can a cyclic voltammogram help to distinguish adsorbed and free di�usingspecies?

5. How can a cyclic voltammogram be used to characterize catalysts?

3.2 Instructions

3.2.1 What is a Cyclic Voltammogram?

During a cyclic voltammogram the potential is controlled and the current is measured.The potential is linear increasing or decreasing. The change of the potential per time isthe scan rate v. As can be seen from the mathematical de�nition (see equation 3.1) thisis the slope of the linear potential.

v =∂E

∂t(3.1)

At the start the potential is usually in a region where no electrochemical reaction isoccurring. The linear sweep of the potential is usually chosen in such a way that thepotential crosses the formal potential of the investigated species (see Figure 3.1). Afterreaching a set potential the slope of the linear potential is inverted, that is a decreasingbecomes an increasing potential and vice versa. This potential is called the vertexpotential. One cycle is �nished when the potential reaches the starting potential again.It is possible to repeat this process several times. The intention behind multiple cycles isoften to observe the stability of a system. Modern software usually o�ers the option to

choose two vertex potentials and a start potential, that is the potential sweeps betweenthe two vertex potentials and starts at a potential between these two.

Figure 3.1: potential vs time during a cyclic voltammetry with indicated vertex poten-tials ( Evertex), formal potential of the investigated species ( E

′0), and scan

rate (v).

The prediction of the current signal has to take a few in�uences into consideration.Sweeping the potential will change the composition of the redox active species in frontof the working electrode according to the Nernst equation (equation 0.6). This changeis done by oxidation or reduction and leads to a current �owing. If there would beno di�usion limitation, the measured current would have a sigmoidal shape (see Figure3.2), but in real systems there will be di�usion limitation. If a species is consumed atan electrode, this species is depleted around the electrode and less species reaches theelectrode, thus less current �ows. If the potential is that high that the concentration ofthe consumed species is 0 at the electrode the depletion will lead to a current followingthe Cottrell equation (equation 2.5). As long as the potential is the limiting factor anincreasing potential leads to an increasing current. As soon as di�usion becomes thelimiting factor the current decreases. As a result the current during the sweep has amaximum current or peak current (see Figure 3.2).

When the scan rate is inversed, the same events happen for the opposite reaction, if thesystem is reversible. To directly read the potentials corresponding to the peak, usuallya voltammogram, a curve of I vs E, is plotted. This way many important parameterscan be determined faster than by plotting the E vs t and I vs t on top of each otheras in Figure 3.2. The I vs E curves are very compact and have characteristic shapes(see Figure 3.3). Symmetry is visible more easily. Very symmetric curves are hints toreversible systems, where both species have the same di�usion coe�cient.

As in chapter 2.2.1, only the contribution due to the faraday current was taken intoconsideration, but as before the capacitive current has a contribution to the total signal.

Figure 3.2: potential sweep (upper graph) and the corresponding current response(green, lower graph) including the contribution described by the Nernst equa-tion (pale violet) and Cottrell equation (pale blue)

Figure 3.3: E and I vs t curves in a voltammetric experiment (left) and the resultingvoltammogram (I vs E curve)

As seen from equation 2.9 the constant change of the potential described by the scanrate v will induce a constant capacitive current. This current will change positive ornegative depending on the scan's direction. If no species that is reduced or oxidizedin the potential range of the CV is present, the voltammogram will only show thecapacitive current (see Figure 3.4). A measurement like this allows the determinationof the electrode's capacity. The equation 2.9 and 3.1 combined show that the capacitive

current during a CV is

Icap(CV ) = v · C (3.2)

Another in�uence that could be visible is the resistance. If a high resistance is present inthe system, the current will be signi�cantly in�uenced by Ohm's law. If only a resistoris connected to the potentiostat, the resulting CV would be a diagonal line because thecurrent and potential will behave according to Ohm's law:

U = R · I (3.3)

If the resistance in an electrochemical cell is high, the CV looks like it is turned −45◦ (45◦ anti-clockwise). Clean contacts and a support electrolyte mostly keep the resistanceat a level that is barely noticed in a CV. If an electrode is smooth the contribution ofthe capacitive current is also rather small compared to the Faraday current, but still itcan be important when small peak currents are measured.

Figure 3.4: schematic CVs showing the contribution of a Faraday current (green) anda high capacitive current (red) as well as the resulting measured current(black)

Figure 3.5 shows how the peak data are read from a typical CV of a solution containingone free di�using reversible species. The peak data are the anodic peak current Ip,a andpeak potential Ep,a as well as their cathodic counterparts Ip,c and Ep,c. Much qualitativeinformation can be gathered from the shape of a CV. More important information canbe gathered either directly from the peak data or observing the peaks changing during aseries of measurements. Please notice that the baseline for the back scan is more di�cultto determine, because the current at the start of the back scan also includes the di�usion

limited current of the forward scan. In the instructions' part methods will be describedto take this into consideration. If the change of peak current is observed for quantitativeanalysis only, the o�set caused by ignoring the base lines is not relevant.

Figure 3.5: CV of a free di�using reversible species with indication of the peak currentsand peak potentials for the anodic and cathodic peak

3.2.2 What Information Can a CV Provide?

Quantitative analysis, that is determining how much of the species is in the solution, isusually done by observing the change of the peak current. The peak current is propor-tional to the concentration of the species inside (see equation 3.4). This makes it easyto perform a series of measurements to create a calibration curve or perform a standardaddition (see chapter ??).

Ip ∝ c (3.4)

Further information can be acquired from the shape of the CV. A reversible system,that is a system that allows the oxidation after reduction or the other way around,always shows a pair of peaks: one peak for the reduction at the cathodic sweep andone peak for the oxidation at the anodic sweep. If one of these peaks is missing, thesystem is not reversible under the given conditions. There are two di�erent reasons foran irreversible system. The species is simply not reversible, because the back reactionis blocked kinetically or thermodynamically or the converted species has a follow upreaction that consumes it. This can be due to another present reagent or due to decay.In a CV the time is a controlled parameter via the scan rate v. This allows investigatingthe cause of this behavior. If the scan rate is increased and the missing peak appears

at a certain velocity, the system was irreversible due to a follow up reaction. As soonas the scan rate is so fast that the back reaction is initiated before the reaction usuallyconsuming the converted species starts, the peak of the back reaction gets visible. Sothe back reaction is possible and the peak appears. If a back reaction is just very slow,reducing the scan rate v might lead to the appearance of a peak.Unfortunately an increase of the scan rate cannot be done unlimitedly. An obviousreason is the device. Each potentiostat, AD/DA converter, etc. has a limited amount ofvalues it can process per second, for example the EmStat3 has a maximum of 1000 pointsper second. This could be compensated by increasing the distance between each point ofthe measurement. Usually the potentiostat averages some values before returning themas one point. The higher the scan rate, the less time is used per step, the fewer pointscan be averaged. As a result random noise is not averaged out. Another limitation is therise time that is the minimum time needed for the potentiostat to adjust the potentialis not in�nitely small. The EmStat can still manage a scan rate of 5V/s with a steppotential of 5mV . Another limit is set by the capacitive current. As seen from equation3.2 the capacitive current will grow linear with the scan rate. The Faraday current willalso increase; since the same species is converted in a shorter time the current is higher.The di�usion layer is also thinner and the concentration gradient is steeper, thus the�ux towards the electrode is higher if the scan rate is increased. The correlation betweenthe peak current Ip of a free di�using species and the scan rate is:

Ip ∝√v (3.5)

This means that the capacitive current increases faster than the peak current withincreasing scan rate. The capacitive current will dominate the CV above a certainscan rate and the shape and peaks caused by the faraday current cannot be identi�edanymore. If experiments demand very high scan rates, ultramicroelectrodes are oftenused. These very small electrodes have very small capacities and thus high scan rates canbe used. Equation 3.5 can provide important information for an electrochemical systemindirectly. This equation is only valid for free di�using species. The peak shape andpeak current for adsorbed species are di�erent. Due to the fact that there is no di�usionfrom the bulk to the electrode of the species but depletion of the species at the surface,a symmetrical peak appears in the voltammogram (see Figure 3.6 a). Furthermore thepeak current of an adsorbed species Ip(ads) shows the following behavior:

Ip(ads) ∝ v (3.6)

A series of measurements using di�erent scan rates and plotting the measured peakcurrents versus v and v

12 will show if the active species is surface bound or free di�using.

If the plot versus v12 is linear, the peak belongs to a free di�using species and if the plot

versus v is linear, the species is adsorbed (see Figure 3.6 b and c). This is useful, becausemost systems have some backgrounds currents that will make it di�cult to recognize ifa peak is symmetrical or not.

Figure 3.6: a) schematic representation of the di�erent peak shapes in a voltammogramdue to adsorbed or free di�using species b) schematic plot of the peak cur-rents Ip of di�erent scan rates v versus the scan rate c) like b but versus

v12

Peak shapes can sometimes give information about surface and species surroundings. Acompletely homogeneous surface, for example the surface of a perfect single crystal, anda completely homogeneous species, for example a monolayer of ferrocene, should delivera very sharp peak. An experienced electrochemist can actually read a complete CV of awell-known system like others can read an infrared spectrum. Another parameter thatdelivers information about the system is the peak potential. A CV allows the fast deter-mination of the redox species' formal potential, although sometimes only approximately.Assuming that your reaction is controlled by di�usion and potential only and there is nokinetic inhibition, the reaction happen as soon as the potential for an oxidation or reduc-tion is reached. The reactions would be time independent and an increasing scan ratewould have no e�ect on the peak potential. The peak current would increase accordingto equation 3.5 or 3.6, due to the fact that current is passed charge per time. If thereduced and the oxidized species have the same di�usion coe�cient in the used solutionthe formal potential should be exactly between the two peak potentials of reduction andoxidation as equation 3.7 describes.

E0′ =(Ep,a − Ep,c)

2(3.7)

If the di�usion coe�cients are di�erent, the formal potential E0′ is still between Ep,a andEp,c, but not exactly in the middle. The described system is reversible in a classic way.The electrochemical reaction is happening so fast that in the time of the experiment thekinetics will never limit the reaction. In an ideal situation the minimum peak separation,so the di�erence between Ep,a and Ep,c is:

∆Ep =59mV

z(3.8)

As before z is the number of electrons transferred during the reaction. If there is a kineticinhibition, the peak separation ∆Ep can be larger and will increase with increasing v.The reason is that a kinetic inhibition means a reaction needs more time to happen. If theamount of time is similar to the time it takes to perform a scan, the kinetic inhibition canin�uence the measurement. When the potential for the reaction to happen is reached,the reaction starts slowly (due to the kinetic inhibition) but the potential is swept withconstant scan rate. As a result a slow reaction seems to happen at higher potentials.Imagine two people standing in front of a moving belt. On the moving belt there areyellow balls and a row of red balls behind the yellow ones. The �rst person has super-fastre�exes and you tell that person to grab the �rst red ball right in front of him/her. Thebelt starts moving and as soon as the �rst red ball reaches the person he or she grabsthe ball. You increase the speed of the belt several times, but due to the fast re�exes the�rst red ball is always grabbed. The second person is slower and has the same task. Heor she will take a second to realize that there is a red ball. The belt starts moving andthis person will grab the second red ball, because he or she was too slow for the �rst one.If the speed of the belt is increased, the person will grab the third, sixth or twentieth redball. Maybe you will have run out of balls before the second person was able to grab one.In the same way a kinetic inhibition leads to a peak separation that depends upon thescan rate. A system showing this behavior is considered quasi-reversible. An irreversiblesystem has, as mentioned before, either an anodic or cathodic peak. Examples of thisare catalytic processes which will be explained in the next chapter (??).

3.3 The Experiments



� sensor cable

� sensor conector







� ItalSens IS-Au

� retort stand

� retort clamp

� cell top


� stirrer

3.3.2 Chemicals

� 0.1 M KCl solution (prepared from solid KCl and demineralized water) for preparationof K3[Fe(CN)6] and K4[Fe(CN)6] solutions

� Mixed K3[Fe(CN)6]and K4[Fe(CN)6] in 0.1 M KCl solutions with ratios ofK3[Fe(CN)6]/K4[Fe(CN)6] = 100

1 ,101 ,

11 ,

110 ,

1100

1 2

� 0.1 M solutions of K3[Fe(CN)6] in 0.1 M KCl

� 0.1 M solutions of K4[Fe(CN)6] in 0.1 M KCl

� 0.1 M H2SO4

Note: If you use the 50 mL beaker, you will need 40 mL of solution for your measurement. We

recommend using a smaller vessel that allows the electrodes to be immersed in 5 mL solution

for the enzyme catalysis experiment.

1All these solutions should have the same sum of concentrations, that is c(K3[Fe(CN)6]) +c(K4[Fe(CN)6]) = constant

2A concentration for all solutions of (0.010 ± 0.001) M can be achieved by using 0.1 M solutionsof K3[Fe(CN)6] and K4[Fe(CN)6] in 0.1 M KCl with a pipette (20ml) and variable microliterpipettes (10− 100l and 100ı1000l)

3.3.3 Procedure

Determination of the Formal Potential

1. Make K3[Fe(CN)6] and K4[Fe(CN)6] in 0.1 M KCl solutions with ratios ofK3[Fe(CN)6]/K4[Fe(CN)6] = 100

1 ,101 ,

11 ,

110 ,

1100

2. Choose from the techniques drop down menu the Potentiometry / OCP. Choose 300sas trun (duration of the measurement) and 1s as tinterval (time between measurementpoints).

3. Immerse all three electrodes (reference, counter, and working electrode) into each one ofthe solutions. Press start and wait until the open circuit potential is stable. Note thesolution ratio and the corresponding open circuit potential. Stop the measurement.

4. Repeat step 3 for each of the solutions. Rinse the electrodes carefully between changingthe solutions.

5. Prepare a 5 mM K4[Fe(CN)6] in 0.1 M KCl solution, for example by adding 1mL of0.1M K4[Fe(CN)6] in 0.1 M KCl solution to 19 mL of 0.1 M KCl solution.

6. Fill the cell with the 5 mM solution. Immerse the three electrodes into the solution.Load the method �le named Exp2_1.psmethod.

7. Start the measurement. Save the curve afterwards under the menu Curve.

8. Determine the formal potential from both measurements.

� Potentiometry: The formal potential from the open circuit potentials can be deter-mined via the Nernst equation (equation 0.6).

E = E0′ + RTzF 2.3log cOx

cRed= E0′ + 0.059V

z log cOxcRed

When the ratio of cOx and cRed is 1/1 the log will be 0 and the potential E is equalto the formal potential E0. A good way to determine the formal potential is to plotE versus log cOx/cRed. A linear �t can be made and the E value at log cOx/cRed = 0is the formal potential.

� Voltammetry: Due to the fact that [Fe(CN)6]4−and[Fe(CN)6]3− have the same

di�usion coe�cient, determining the formal potential is simple. Read the peakpotentials for the anodic and cathodic peaks Ep,a and Ep,c from the CV. Theiraverage is the formal potential.

Identi�cation of a Reversible, Quasi-reversible, or Irreversible System

1. Prepare a 5 mM K4[Fe(CN)6] in 0.1 M KCl solution, for example by adding 2 mL of0.1 M K4[Fe(CN)6] in 0.1 M KCl solution to 38 mL of 0.1 M KCl solution.

2. Fill the cell with the 5 mM solution. Immerse the three electrodes into the solution.Load the method �le named Exp2_2.psmethod.

3. Start the measurement. Save the curve under the menu Curve.

4. Repeat the measurement with a Scan rate of 0.01 V/s, 0.025 V/s, 0.05 V/s, 0.1 V/s, 0.25V/s, and 0.5 V/s. Save each curve.

5. The goal is to use Table 3.1 to identify if [Fe(CN)6]4− and [Fe(CN)6]3− is a reversible,quasi-reversible, or irreversible system. Using the list of criteria di�erent parameters needto be extracted from the CV and compared.Extracting the Ip,c is not directly possible due to the fact that there is no direct baselineavailable. When the scan rate is inversed there is still current �owing due to conversionof the [Fe(CN)6]4−. The current overlaps with the baseline of the cathodic sweep, soeither a graphical or mathematical extrapolation has to be done. The graphical methodis rather di�cult, since parts of the curve has to be guessed. The mathematical way justneeds a single equation and some parameters from the curve. Determine Ip,c,0 and Ivertexaccording to Figure 3.7 and use equation3.9 to determine Ip,c.

Table 3.1: diagnostic criteria for reversible, quasi-reversible and electrochemical irre-versible systems

Reversible Quasi-reversible Electrochemical irreversible

∆Ep = (Ep,a0 − Ep,c) = 59mV/z ∆Ep = (Ep,a0 − Ep,c) = f(v) Ep shift about 30mV/αz (298)(298K) with tenfold hiegher v

(α = unknown factor)

E0′ = 0, 5(Ep,a + Ep,c) E0′ = 0, 5(Ep,a + Ep,c) -(DOx = DRed) (DOx = DRed)

Ip,a ∝ v12 - Ip,a ∝ v

12

Ip,a = const1 v12 Ip,a = const2 v

12

Ip,a/Ip,c = 1 for all v Ip,a/Ip,c = 1 for all v One peak current (Ip,c) missing

|Ip,c| = |Ip,c,0|+ 0, 485Ivertex + 0, 085|Ip,a| (3.9)

Figure 3.7: scheme of a CV indicating how to determine Ip,c,0 and Ivertex

Investigation of a Gold Surface

Gold is a special metal. The impact gold had in the history of mankind is well-known, butfrom a scienti�c point of view gold has quite a number of special properties. Under the rightconditions aurides can be made, that is salts where gold is present as the anion Au-. Workingwith gold electrodes is popular due to the fact that organic sulfur compounds like thiols (R-SH)easily bind to a gold surface. Working with gold electrodes has the disadvantage that gold issoft and hydrophobic. As a result a gold surface collects all organic residues in the air and getsdirty rather quickly. A popular method to clean electrodes is electropolishing. In this methodhigh anodic potentials are applied and adsorbed species will desorb. Furthermore gold formsa gold oxide monolayer, that is the gold is covered with a single layer of oxidized gold. Thereduction peak is very sharp and in sulfuric acid no other peak overlaps with it. The charge ofthe reduction peak can therefore be used to calculate the real surface of the gold electrode.

1. Take an ItalSens IS-Au electrode and insert it into the sensor holder. To attach theelectrode to the sensor connector take the strip of electrodes and hold it at the endwith the single letter (A to D) and a number (1 to 20). Since the electrodes are easilycontaminated or destroyed, please take care not to touch it with bare hands. Cut awaythe access plastic around one electrodes SPE and place it in the connector. You mustcut the sensor narrow enough for �tting it into the connector, but please take care notto damage the electrodes nor the lines during cutting.

2. Connect the sensor holder to the potentiostat and �x it in the retort stand.

3. Immerse the three electrodes on the stripe into a 0.5 M sulfuric acid.

4. Load the method named Exp2_3.psmethod.

5. Start the measurement. The electrode should be clean now.

6. Repeat with Scan rate = 0.005V/s, Number of Cycles = 1 and save the curve under themenu Curve.

7. Repeat the measurement with a Scan rate of 0.01 V/s, 0.025 V/s, 0.05 V/s, 0.1 V/s, 0.25V/s and 0.5 V/s. Save each curve.

8. The peak around 0.65V is the gold reduction peak. Since an adsorbed species is reducedequation 3.6 should be valid. Look at the reduction peak and plot Ip,c versus v to proofor falsify.

9. The next steps are used to calculate the real surface. Take the CV with v = 0.1 V/s tocalculate the surface. First take the area of the reduction peak Ap,c, which is also theintegral. To get the charge divide by the scan rate. This is also represented in equation3.10.

Q =Ap,c

v(3.10)

10. The real surface is determined by using the charge per surface area for a gold oxide layerq. This value can be found in the literature: 3.84C/m2. 3 The real surface Areal iscalculated according to

Areal =Q

q(3.11)

11. Usually the roughness factor r is used to characterize the gold electrode. It is the ratiobetween real and geometric area (see equation 3.12). It shows if electrodes are polishedin a reproducible way or if they are smooth enough to be used in following experiments.

r =Areal

Ageo(3.12)

3.3.4 Disposal of Chemicals

� The potassium chloride solution is harmless and can be disposed in the sink. The elec-trodes and cell should be cleaned with demineralized water. Otherwise the drying solutionwill leave salt stains.

� Solution containing iron need to be collected in a container for liquid waste containingheavy metal. The pH of the collected heavy metal solution should be alkaline (pH > 9).The container should be disposed according to local laws.

� Solution containing inorganic acids need to be collected in a container for liquid wastecontaining inorganic acids. The containers should be disposed according to local laws.Those mentioned containers will be provided from your laboratory supervisors.

References

[1] Palmsens Educational Kit. Teacher ' s Guide to the PalmSens Educational Kit. 2015.

[2] Allen J Bard, Larry R Faulkner, Johna Leddy, and Cynthia G Zoski. Electrochemical

methods: fundamentals and applications, volume 2. Wiley New York, 1980.

[3] Cynthia G Zoski. Handbook of electrochemistry. Elsevier, 2006.

[4] Joseph Wang. Analytical electrochemistry. John Wiley & Sons, 2006.

[5] John E. Baur and R. Mark Wightman. Di�usion coe�cients determined with microelec-trodes. Journal of Electroanalytical Chemistry, 305(1):73�81, 1991.

718.019 biomedical sensor systems (2018): electrochemistry ...€¦ · electrochemistry (exercises...

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