718.019 Biomedical Sensor Systems (2018): Electrochemistry (Exercises 5+6)

718.019 Biomedical Sensor Systems (2018): Electrochemistry (exercises 5+6)

Christoph Aigner, Andreas Petrovic March 5, 2018

Instiut für Medizintechnik Kronesgasse 5/II 8010 Graz

Contents

0 Introduction1 0.1 Why Electrochemistry? ...... 1 0.2 What Is a Potentiostat Good for? ...... 2 0.3 How Does a Potentiostat Work (in a nutshell)? ...... 3 0.4 Electrodes ...... 5 0.4.1 Working Electrode ...... 5 0.4.2 Reference Electrode ...... 7 0.4.3 Counter Electrode ...... 8 0.4.4 Screen Printed and Film Electrodes ...... 9

1 Copper Deposition 11 1.1 Goals ...... 11 1.2 Introduction ...... 11 1.2.1 Corrosion research and corrosion protection ...... 11 1.2.2 Electrogravimetry and the Electrochemical Series ...... 14 1.3 The Experiments ...... 15 1.3.1 Devices and Equipment ...... 15 1.3.2 Chemicals ...... 15 1.3.3 Procedure ...... 15

2 The Cottrell Experiment and Diusion Limitation 17 2.1 Goals ...... 17 2.2 Introduction ...... 17 2.2.1 The Cottrell Experiment ...... 17 2.2.2 Electrochemical Double Layer ...... 22 2.3 The Experiments ...... 25 2.3.1 Devices and Equipment ...... 25 2.3.2 Chemicals ...... 25 2.3.3 Procedure ...... 25 2.3.4 Disposal of Chemicals ...... 27

3 Cyclic Voltammetry - the Most Used Technique 28 3.1 Goals ...... 28 3.2 Instructions ...... 28 3.2.1 What is a Cyclic Voltammogram? ...... 28 3.2.2 What Information Can a CV Provide? ...... 32 3.3 The Experiments ...... 36 3.3.1 Devices and Equipment ...... 36 3.3.2 Chemicals ...... 36 3.3.3 Procedure ...... 37 3.3.4 Disposal of Chemicals ...... 40 0 Introduction

This script contains exercise 5 (Copper Deposition and The Cottrell experiment) and exercise 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 research and industry. 1 Corrosion research develops methods to save millions of Euro that are needed to remove the damage done by corrosion. Electrochemistry delivers the insights about the processes causing corrosion and is capable to investigate means against corrosion in hours instead of weeks. Although these experiments cannot replace the long term test under real conditions, they can provide the knowledge for a fast pre-selection of 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, but since the boom of the mobile internet the limitations of batteries have become the limitations of our electronic society. The electric car only makes sense if we have batteries with enough capacity to drive hundreds of kilometers without recharging. Either that or we need another electrochemical device, the fuel cell, to power the cars and reduce the request for fossil oils. In the same context solar panels are important to gain energy from a sustainable source, the sun. The electrochemistry of photovoltaics is providing important knowledge to improve solar panels and brings the idea to satisfy our energy demand with solar power within our reach. In analytical chemistry electrochemistry makes miniaturized devices possible that deliver a quantitative result in the ppb range. These devices are easy to handle point-of-care devices. Electronic components are very suitable for mass production and as a consequence these devices are aordable to a broad part of the population. The most known example of an electrochemical point-of-care device is the blood sugar test. Easy to handle, inexpensive, and improving a diabetes patient's quality of life drastically it is the most successful electroanalytical chemistry product up to today. The blood sugar sensor is based on an enzymatic reaction, that is it is a biosensor. Gaining an insight into biological electrochemistry is a powerful tool for modern technology. Using enzymes in detection to create a very sensitive and selective sensor is just one of the many possibilities biomolecules oer to technology. Biofuel cells create energy out of glucose or lactate and can thus power an implanted diagnostic device inside a human body. Imitating the photosystem of plants leads to new developments in photovoltaics. DNA chips or protein chips allow a high density of recognition elements on 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 the last 10 years electrochemical impedance spectroscopy has been opening new possibilities in 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 modication to create an electrochemical signal. These are all modern electrochemical applications and additionally classic electrochemistry 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 electrochemistry during the studies of chemistry, biochemistry, physics, biology, medicine, pharma, chemical engineering, and material sciences. For more detailed information please refer to [2], [3] and [4]

0.2 What Is a Potentiostat Good for?

A voltmeter does measure the potential dierence between two points. To do so the circuit needs to be closed, but a very tiny current is usually owing. This way you can see the potential dierence between two points with almost no disturbance for the investigated system. Knowing the potential between two points is useful but manipulating potentials is even better. This is what a power source does. In every household batteries are the most common example of DC power supplies: a power supply having constant non-alternating potentials. They are used for ashlights, mobile phones, clocks, etc. In electrochemical applications DC power supplies are used for galvanization (the deposition of metals on conducting materials, often other metals). Another well-known application is the electrolysis of compounds. A well-known industrial example is the chloralkali process where salt (NaCl) and water (H2O) are split into chlorine (Cl2), hydrogen (H2), and sodium hydroxide (NaOH). The disadvantage is that you cannot investigate a single electrode and thus a single event. The current ows through the anode (electrode where oxidation happens) and the cathode (electrode where reduction takes place). So both these electrodes inuence the measured current and the current limiting process cannot be determined. This is especially an issue in analytical chemistry. An electrochemical analysis that can be used easily with a power supply is the electrogravimetry. An electrode's weight is determined and all the metal in a dened sample volume is reduced, which leads to the precipitation of the metal on the electrode. The weight is measured again and thus the amount of metal in the sample volume will be determined. Although this method works ne, it has some disadvantages. The process takes time, e.g. 30 min plus drying and weighting for copper, nickel or lead oxide. In the given potential window only the analyte metal should be reduced. Hazardous side products may be produced. The solution needs to be heated and stirred to decrease the concentration polarization and to make the conversion of 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 methods a potentiostat is needed. A potentiostat uses three electrodes and a feedback loop to control the potential and measure the current owing at just one of these electrodes, the working electrode. The potential will be measured to a xed reference point and thus a lot of information about the event happening at the working electrode can be gathered. The potentiostat controls the potential of an electrode while measuring the current owing 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 and measures the current owing through it. Why not just two electrodes? One of the reasons is that we cannot measure the potential of the working electrode against a xed point when we just have two electrodes. Imagine a two electrode system that consists of the already mentioned working electrode and the electrode, which potential should be our xed reference point, the reference electrode. We apply a certain potential between these electrodes and an electrochemical reaction happens at the working electrode, but since the circuit needs to be closed and current needs to ow, a reaction that is inverse to the reaction at the working electrode must occur, that is if an oxidation occurs at the working electrode, a reduction must take place at the reference electrode. If a current ows at a constant potential, an electrochemical reaction 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 the amount n of a species that took or gave z electrons at the electrode. F is the Faraday constant and represents the charge of 1 mol electrons. The current I is the charge Q per time t owing through the electrode:

Q I = (0.2) t The equations' 3.1 and 3.2 combination shows that the current I owing is connected to the reaction happening at the electrode via the amount n:

n · z · F I = (0.3) t Imagine now that the current is owing at the reference electrode. At this electrode a species' amount of n is converted. This conversion leads to a change of the surface or the concentration of the solution surrounding the electrode. The Nernst equation shows a clear correlation between the potential E of an electrode and its surrounding:

RT a E = E0 + ln Ox (0.4) zF aRed E0 is the standard potential of the redox couple Red and Ox. R is the gas constant and T 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 to a simplication of the equation:

RT c f RT c RT f 0 RT c E = E0 + ln Ox Ox = E0 + ln Ox + ln Ox = E0 + ln Ox (0.5) zF cRedfRed zF cRed zF fRed zF cRed

00 The two activity coecients fOx and fRed are included in the resulting potential E , which is called the formal potential. Since it contains parameters that depend on the environment, such as temperature and activity coecients, E00 cannot be listed but needs to be determined for each experiment, if necessary. Most experiments in analytical chemistry are performed at room temperature (295 K). This makes another simplication possible. Out of convenience also the ln will be transferred to the log.

0 RT c 0 0.059V c E = E0 + 2.3log Ox = E0 + log Ox (0.6) zF cRed z cRed For practical application equation 0.6 is the most used form of the Nernst equation. For many applications one can assume that E0 is roughly the same as E00 , because both of the activity coecients are close to one. In this form (equation 0.6) the correlation between the surrounding of an electrode and its potential is visible more easily. As mentioned before all the simplications at equation 0.4 were performed: The change of the solution surrounding the reference electrode, due to a owing current, leads to a change of the potential that is supposed to be our xed reference point. But we cannot limit the current ow through the reference electrode (RE), because all limitations should be caused by the process that we want to investigate, that is the process at the working electrode (WE). The solution for this problem is a third electrode. At this counter electrode (CE), also known as auxiliary electrode, the counter reaction to the working electrode's reactions takes place. The current is owing between the working and the counter electrode. The potential is controlled between working and reference electrode (see Figure 0.1). The potential between counter and reference electrode is adjusted in such a way that the current owing through the working electrode at a certain potential between working and reference electrode is satised. There are limits for the potential a 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 to stay within the limitation of the DC potential range. Because the compliance voltage cannot be controlled by the user the CE has to be bigger than the WE. A bigger surface at the same potential leads to a higher current . A rule of thumb suggests that the CE should 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 the current owing at the WE. Usually the distance between CE and WE is big enough so that the reactions do not inuence each other, and the counter reaction can be ignored, but sometimes, in small volumes for example, it can be helpful to know which reaction happens at the counter electrode. 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 potentiostats need little eort to set up and operate. The average electrochemist will spent a lot more time with his or her electrodes. They need to be polished maybe modied, and sometimes it is rather challenging to x them all in the right place. For this reason we will rst talk about the three dierent electrodes, before we take a closer look at the potentiostat again.

0.4.1 Working Electrode

The working electrode is the place where the reaction happens that we want to control or investigate. As a consequence this electrode should be carefully and reproducible prepared. The most common working electrodes are disc electrodes. For example, a metal cylinder or a metal wire is surrounded by Teon or PEEK and the cross section is exposed. The metal disc is connected to a wire at the other end of the coating, so it can be connected (see Figure 0.2). Common materials for working electrodes are platinum and gold as well as a variety of carbon phases. Very popular due to its conductivity and 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 (Teon) coat; photo and scheme

is usually polished before usage. Very ne emery or sand paper as well as slurries of aluminum oxide particles on cloth are used for mechanical polishing. Another way is the electropolishing. This is used for surfaces that get contaminated very easily. The hydrophobic character of gold makes it often necessary for gold electrodes to electropolish them. When residues need to be removed by electropolishing a surface is cleaned by applying high potentials. Rough surfaces will lead to high currents due to the charging of the interface. Dirty surfaces can show artifacts in the measurements or can be the cause that not the full active surface has access to the electrolyte. So cleaning an electrode is an important step in the electrode preparation. A blank electrode is often just the beginning of the preparation step. For many applications an electrode's surface needs a modication: To produce an ion selective electrode polymer lms are put on electrodes, to produce a biosensors enzymes or other biomolecules, for example DNA, are immobilized at the electrode, to produce immunosensors parts of the immune system are immobilized at the electrode and also inorganic catalysts are used for electrode modication. A very interesting working electrode is the mercury electrode. Due to its liquid state mercury is used as a droplet electrode. At the end of a tube a mercury droplet is formed and used for the corresponding experiment. To clean the electrode the droplet is enlarged until it drops and a new surface is prepared with clean mercury. This way clean and easy to reproduce surfaces are obtained, but due to its toxicity in most labs mercury is not common. In general modern electrochemistry avoids mercury, but some experiments deliver the best results with mercury, so that the mercury electrode is still used 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 through an electrode leads to an electrochemical reaction that will change the composition of the electrode's environment and thus the potential. As a consequence there should be as little 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. The electrodes of the rst type are basically just metal surfaces in an electrolyte. According to Nernst's law (see equation 0.4) the potential is depending directly on the surrounding solution. An electrode of the second kind is usually a metal surrounded by a hardly soluble salt of itself. This electrode is then immersed into a solution containing the anion of that salt. Usually the anion has a high concentration to make sure that the metal salt is not dissolving fast and that small changes in the concentration have a low impact on the potential. A common reference electrode is the silver / silver chloride electrode (see Figure 3.3). A silver wire is coated with a lm of silver chloride. This wire 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 potassium chloride solution. + 0 0.059V [Ag ] E = E0 + log (0.7) z [Ag] Due to the fact that it is a solid material the concentration of silver is 1 M. The concentration 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 E00 (Ag/Ag+) the result is: 0 0.059V E = E0 (Ag/Ag+) + log[Cl−] (0.9) z If the chloride concentration is kept constant a good reference electrode is created. This is usually achieved by separating the measuring solution from the solution surrounding the reference electrode using a porous frit. Even small concentration changes due to diusion through the frit or evaporation have little impact on the concentration if the concentration of chloride is high. The frit also prevents silver ions from diusing into the measuring solution. The silver / silver chloride electrode is one of the two most popular chloride containing reference electrodes. The other one is the calomel electrode. This reference electrode contains mercury in contact with mercury(II) chloride in a chloride containing solution often saturated with chloride. You should always indicate clearly versus 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 are small enough, that is in the nA range, it is possible to immerse the wire with the hardly soluble salt coating in the measurement solution. This is called a pseudo reference electrode. The reference and counter electrode can also form a short circuit and you can work with a two electrode system. If your currents are higher, you want to perform a long term experiment, for example overnight, or need a very stable potential, a three electrode 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 metal with a large surface. For people who work with many dierent materials platinum is a good counter electrode due to its inertness towards most solutions. A really big surface area is achieved by a mesh structure, but often a platinum wire has already a surface area several magnitudes bigger than the working electrode. Platinum is also very easy to clean. A hand torch is usually enough for making the platinum glow and to remove all stains. The counter electrode should have a large surface area, to enable a high current ow even with low potentials. This way water splitting, gas production, or the production of aggressive radicals is avoided. If this kind of reaction is happening in a small volume of solution, 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 without changing the chemical composition of the solution. The counter electrode should not limit the current or inuence 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 everyday live, especially point-of-care measurements. The measuring devices as well as electrodes can 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 education must skip all the demanding preparation and cleaning steps to reuse an electrode. This is why in recent years disposable test sensors have become more and more important. A gold disc in a Teon coat is too expensive for a disposable sensor. Another group of electrodes is more suitable for this objective. Flat electrodes that are deposited on a surface are the up-to-date commercial biosensors. These are again divided into two groups: lm electrodes and printed electrodes. Film electrodes are made by metal deposition on a surface that is covered with some kind of mask. The mask only exposes the parts of the metal that should be lled with metal afterwards. The metals are deposited by heating up the metal source in a vacuum while the masked sample is present. If dierent metals need to be on the same sensor, because a gold WE and a silver RE are requested, dierent masks and cleaning processes need to applied to ensure that the gold and silver will not be deposited on the 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 are created by photolithography and very small structures can be created. For example, DNA chips with dozens of working electrodes in a 1mm2 or interdigitated electrodes can be produced. Interdigitated electrodes are always pairs. Each electrode looks like a comb and they are arranged in such a way that each tooth is between two teeth of the other electrode. With this setup species are cycling between the two electrodes and this behavior creates an amplication of the signal. The main dierence with screen printed electrodes is the material the electrodes are made of. While lm electrodes are made from pure metals, screen printing uses pastes or inks to create a sensor. A mesh with a mask is placed above the empty support, a polyester sheet for example, without touching the support. The ink is put on the mesh and a squeegee is moved across the mesh to press the paste through the template as well as to create contact between the mesh and the support. The paste remains on the support. A drying step follows. The drying step can include heating, especially for water based inks. Each dierent paste also needs a dierent mask, same as with the thin lm methods, but since no vacuum is necessary the price for the needed equipment is lower and depending on the drying step the procedure can be faster. The inks usually contain the conducting material, for example the metal particles or carbon compounds, some dispersion agents as well as other additives, polymer binders and 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, but the production is cheaper, especially in mass production. Due to the fact that it is easier to use dierent materials than while producing lm electrodes, screen printed electrodes have often WE, RE, and CE made from dierent pastes. Counter electrodes are made of carbon pastes, because they are cheaper than metal pastes. Reference electrodes can be made of silver paste mixed with silver chloride, so that a stable reference electrode is achieved. The price per piece for a lm or screen printed electrode decreases with the number of produced electrodes, due to the fact that the most expensive parts are the masks that need only to be produced once. The low price allows these electrodes to be produced as a one-time use or disposable electrode. As a consequence any steps for cleaning the electrode after a measurement can be skipped, because the electrode is disposed. If the electrode has a dedicated application, most preparations can already be done by the producer. All this makes it possible to develop electrochemical applications that can be operated by people with no electrochemistry background. The best example is the blood 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 knowledge. 1 Copper Deposition

1.1 Goals

Most students already know the basics of electrochemical potentials and how they are applied in batteries or electrolysis. This experiment is meant to facilitate a transition from what is taught in most schools and the slightly more advanced level of potentiostatic experiments. Because galvanization for depositing metals or other protective layers is a widely used technique in industry as well, some basic knowledge of this important process should be gained. The experiment and its discussion should teach: A general idea about corrosion protection How can the dierent 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. It is a big cost factor. Replacement of corroded materials in buildings, bridges, or monuments and painting or galvanization of building elements is a huge investment into maintenance. 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 US gross domestic product1. Painting the Eiel tower costs ca. 4 million e every 7 years2. Without a proper and closed layer of paint this monument would become unstable and break down. As a consequence one of the most important elds in electrochemical research focuses on corrosion. A lot of companies fund academic research for investigating the mechanism of corrosive processes and for creating new methods for corrosion protection. Companies check paints and lubricants for their corrosion protection capabilities and develop new methods for creating coatings or creating non-corrosive environments for machinery, for example ooding the entire gears with oil. From a chemical point of view protection by ordinary paints is easy to understand. The paint is usually a mixture of a polymer, color pigments, and various additives for adjusting the physical properties of the paint. For example these additives create smooth or rough surfaces or grant 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.toureiel.paris/en/everything-about-the-tower/themed-les/97.html (water or an organic solvent). The solvent evaporates after the paint has been applied and a polymer with the additives and pigments stays on the painted surface. The result is a physical barrier between the surface and its environment. This barrier protects the surface from corrosion. A crack in the layer of paint will grant access to the surface and corrosion may occur. Here already electrochemistry can help improve the protection layer. Electrodeposition paints are paints that are made insoluble by an electrochemical reaction. A conducting work piece is dipped into the paint bath and a potential is applied. The created reaction close to the work piece's surface will make the paint insoluble and thus a layer of paint precipitates on the metal surface. All conducting parts will be covered with the paint, even usually hard to reach places. If the paint forms an insulating layer, the reaction will stop itself by blocking the conducting surface completely; only the needed amount of paint will be used. These paints have dierent mechanisms. The most common one is neutralizing charges. The polymers in the paints have some charged functional groups. These groups make them water soluble. These groups can be deprotonated acids or protonated bases. By way of illustration we assume a polymer with deprotonated carboxylic groups (COO-) is used. Due to its negative charge it will migrate towards the positive charged electrode, the anode. The potential at 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 leads to a protonation of the COO- groups. As a result the polymer's COOH groups have no charge and the polymer is not soluble in water anymore. The polymer precipitates on the anode (see Figure 1.1). After completion of the deposition excess paint is rinsed away and the paint layer is cured. This means that the paint is heated up, so it can ow and 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 the aluminum layer beneath. This has the advantage that damaged parts of the coating will be replaced by the aluminum itself. The thickness of the layer can be increased electrochemically, leading to a better protection. Up to now the mentioned coatings were passive coatings. Active coatings can be created by covering metals with metals. This is usually done by dipping the metal to protect in a solution with a salt of the coating metal. A potential is applied between the metal that needs protection and an electrode made from the coating metal. A negative potential is applied to the core metal and all positive metal ions (cations) will migrate to the core metal that acts as cathode, the electrode where reduction takes place, and the ions are reduced to elemental metal at the cathode. The other electrode is oxidizing its own metal and releasing it as ions into the solution. This electrode is called anode, the electrode where oxidation happens (see Figure 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 nobler than another one and thus predict its behavior in contact with it. Noble metals are useful to create a coating that is inert to corrosion, but if the coating gets damaged and the less noble metal is exposed, a local element is created. The noble coating metal will take electrons from the less noble core metal as a result the speed of corrosion is increased instead of slowed down. An example of this would be chromium coated steel. Coating with a less noble metal that passivates itself is usually a better idea. In case the coating is damaged and the core metal is exposed the coating will act as a sacricial anode. The noble core metal will drain electrons from the less noble coating and thus the core will 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 of this experiment the knowledge of these most important methods will suce. Corrosion is an important research eld due to the fact that corrosion causes big amounts of damage to machines and monuments. Understanding and preventing corrosion 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 for quantication, the analyte. A potential is applied and the metal ions are deposited as metal or metal oxide lms. Which potential is applied and how the metal is deposited depends on the metal. Copper, for example, can be reduced at the electrode and deposited as elemental metal. It forms a properly attached lm on platinum electrodes. This lm can be washed, dried, and scaled. Lead ions are oxidized for deposition. A lm of lead oxide 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 of lead in the lm can be calculated. This technique for quantication demands clean and careful working. Damage of the lm, remains of water or other solvents during scaling, unnished deposition of the metal, or a wrong mass of the electrode without the lm will lead to wrong results. To reduce the time needed for deposition the metal solution needs to be heated and stirred during the deposition. It seems easier to measure the charge that passes through the electrode during the deposition, but often side reactions and capacitive current appear during these processes and these will contribute to the charge passing through the electrode. The electrochemical series show that dierent processes happen at dierent potentials. For example, copper is deposited at less negative potentials than zinc, and cadmium is easier oxidized than lead. Of course these insights are also true for other redox active species than metals: Chloride will be oxidized to chlorine at lower potentials than water will be oxidized to oxygen. These dierences can be used to determine dierent molecules from the same solution, if the dierence is large enough for these processes to be regarded as separate. A classic example is the determination of copper and nickel from a mixed solution of copper and nickel salts. Deposition of metals can also be used for analytical methods. The dierences in the standard potentials allow individual precipitation of dierent metals in the same solution. 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 holes into the clamp. Put the stirrer under the clamp and 50 mL beaker with a magnetic stirring bar between them. Bend the paper clip into a straight piece of wire. Most paper clips have an insulating coating. Use some emery paper to remove it. The actual shape of the paper clip wire is not important for the experiment, but a straight wire is easy to handle. Insert the counter electrode (platinum wire) and the reference electrode, the silver / silver chloride electrode (the electrode with the liquid lled glass body), into two of the cell top's holes. Connect the paper clip to the working electrode cable (red) with the crocodile clip. Connect the counter electrode to the black cable and the reference electrode to the blue cable. Please be aware that the part of the reference electrode containing the diaphragm always needs to be in the solution or otherwise protected from drying. During evaporation of solution little crystals form in the pores of the frit and the electrode will be blocked. This should be avoided. 2. Fill the beaker with the copper salt solution. Position the electrodes in such a way that all three electrodes are in the solution. The paper clip wire should not be immersed deeper than the platinum wire. 3. Start PSTrace. 4. Choose the scientic mode in the upper left corner. 5. Select your potentiostat from the drop down menu. If it is not in the list, press the refresh button (green arrows). Maybe your computer takes a while to install the USB driver. Press the Connect button. Your device should be connected now. 6. Choose Amperometric Detection from the Technique list. In Sample and Sensor you can add comments for your own data organization. Set t condition and t deposition to 0, because the electrode needs no treatment before the measurement. The t equilibration is the time during which the rst potential of the measurement is already applied without recording the current. This is usually done to exclude capacitive current from the beginning of the graph, but since we are only interested in the deposited metal and not in the measurement set this time to 0 s. The potential applied during the measurement (E dc) has to be low enough to reduce copper but high enough to avoid zinc deposition. The electrochemical series contain 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 is by convention 0 V. The reference electrode supplied with this kit is an Ag/AgCl electrode. It has a potential of E0(AgCl/Ag) = 0.205 V. So a potential of 0 V vs the SHE is a potential of -0.205 V for the Ag/AgCl electrode. The other values can 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 hydrogen at your paper clip surface. Since experience tells that the large overpotentials for 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 to corrosion of the paper clip itself. Assuming that no other species in our solution

will react, an Edc of -0.5 V vs Ag/AgCl should suce and avoid gas bubbles from hydrogen. The t interval is the time between two measurement points. Set it to 1 s. The parameter t run is the time the experiment lasts or in this particular case how 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 with demineralized water to observe it more closely. 9. The solution can be reused for multiple paper clips. You can vary the deposition time and observe the inuence on the result. Or try to immerse an uncoated part of the paper clip without applying a potential. 10. If you have to dispose the solution keep the following in mind: Solutions containing copper or zinc need to be collected in a container for liquid waste containing heavy metal. 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 cell should be cleaned with demineralized water. Otherwise the drying solution will leave salt stains.

2 The Cottrell Experiment and Diusion Limitation

2.1 Goals

Electrochemically the Cottrell Experiment is very easy. It is easy to perform and the electrochemical process is easy to understand. However, the physical observations during the experiment are important for many electrochemical techniques and experiments. The importance of mass transport, especially diusion, as limitation for a owing current and the additional current owing due to charging of the electrochemical double layer are basic principles and inuence almost every electrochemical experiment. The experiment and its discussion should teach: What is capacitive current and what is its inuence? 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 performed under real conditions. Suddenly temperature, pressure, etc. inuence the outcome of the 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 at a potential that won't lead to any reaction at the electrode and when a stable state is reached a potential step is made that will lead to a chemical reaction. In detail this means the potential E of the electrode is before the potential step below the formal potential E00 of an electrochemical active species. This species is present in the solution surrounding the electrode. The potential before the potential step is named E1 in Figure 2.1. In the following pictures and explanations it is assumed that the species is in its reduced form (Red). The concentration of Red is ∗ everywhere before the potential cRed step, meanwhile the concentration of the oxidized form ∗ is 0. The potential step is cOx raising the potential so that 00 E2 > E (2.1) results. According to the Nernst equation (see equation 0.6) the ratio Ox/Red will become bigger. Red is consumed by the electrode by accepting an electron, that is Red is oxidized, and the oxidized form (Ox) is produced (see Figure 2.1). If E2 is signicantly higher than 00 , ∗ is 0 at the electrode surface and all Red close to the electrode is E cRed consumed.

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. The current will ow as long as Red is consumed. Before the potential step the lack of potential as driving force was limiting the reaction and by increasing the potential it is no longer the limiting factor for this reaction. Since all molecules of Red reaching the electrode are consumed immediately, the mass transfer of Red towards the electrode becomes the limiting factor. The species Red is completely depleted in front of the electrode but in a distance from the electrode surface the concentration is still ∗ x1 cRed (see Figure 2.2a). This concentration gradient leads to diusion of Red towards the electrode surface. The diusion will also deplete all the Red in the distance x1, so Red from a distance x2 will diuse to x1. This means that the layer of liquid in which diusion occurs, the diusion Figure 2.2: a) concentration of Red cRed vs the distance from the electrode surface x during dierent 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 limits to the growth of the diusion layer. For usual macroelectrodes, that is electrodes with a diameter above 25µm, the limit is reached when the diusion layer is disturbed by convection, that is movement of the liquid. This can be self-induced or due to stirring. In a stirred solution a diusion 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 electrodes. 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 by Fick's law ∂c(x, t) J = D (2.3) ∂x which shows that the ux J is proportional to the gradient or in a one dimensional case slope of the concentration curve ∂c/∂x with the diusion coecient D as proportional factor. D is a property and thus a xed value for a molecule, ion, or complex in a medium. In literature one can nd tables for D, but one should take care that this value is also referring to the right medium. Directly after the potential jump in the Cottrell experiment the concentration of Red is 0 at the electrode surface only. With increasing time the diusion layer grows into the solution and the resulting gradient is less steep. As a result the ux J is getting smaller 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 dierential of the concentration is described by Fick's second law: ∂c ∂2c = D (2.4) ∂t ∂x2 Fick's laws are treating general diusion; solving it for the particular case of the Cottrell experiment requires some conditions for the initial state and the boundaries. These conditions are:

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

2. After the potential step the concentration at the electrode surface is 0. cRed(0, t) = 0 3. In an innite distance from the electrode the concentration is always ∗ . ∗ cRed cRed(∞, t) = cRed Using these conditions to solve Fick's law leads to the Cottrell equation: r D I = zF Ac∗ (2.5) πt 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 in electrochemistry when depletion of observed species starts and the reaction gets diusion controlled. As we now know equation 2.5, we expect a current responding to the potential step as seen in Figure 2.3 a. There will be a current step with the potential step and − 1 afterwards the currents decays with t 2 . 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 response in 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 diusion layer cannot grow innitely. At some point the disturbances and movement of the liquid will destroy the outer layers of the diusion layer. Although this seems to be a simple experiment with little practical value, it is actually still used to determine diusion coecients or the electrode's surface area. This is done by turning the I-t-curve into a linear form. From the Cottrell equation it is clear that I − 1 vs t 2 should lead to a linear plot with a slope r D m = zF Ac∗ (2.6) π that can be used to calculate A, c∗, D, or z, if the other ones are known (see Figure 2.3b). If the experiment is performed and the linear plot is made, it often does not show a perfect line, especially in the beginning of the curve and at the very end. The end of the curve is not perfectly linear since the diusion layer does not expand perfectly into innity as discussed before. The beginning of the curve is inuenced by two factors. First, it is impossible to perform an ideal potential step. Potentiostats have rise times in ns or even µs, that is they can reach a set potential in that time period, but this is still not perfectly instantaneous. So the beginning of the curve deviates from the ideal potential 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 diusion. 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 other charge than the electrodes charge will accumulate directly at the electrode, forming a layer 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 of the 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 layer of ions and the electrode act like a capacitor and this has impact on most electrochemical techniques. This is only a rudimentary description of the electrochemical double layer; there are more sophisticated models, but for many electrochemical experiments the simple model will suce.

Figure 2.4: scheme of the electrochemical double layer

Up to now it was assumed that the electrochemical active species Red is transported by diusion 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, the positive potential of the electrode will attract the Red-ions. Why are complete models and observations done without taking migration into consideration? 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, a high solution resistance will make the measured current smaller and increase the Ohmic drop, the dierence in the potential applied between the reference electrode and the working electrode and the potential that the working electrode is feeling. To reduce the resistance of a solution an electrochemical inert support electrolyte is added. Often the buer itself in a pH buered solution is sucient or a salt with a high solubility, for example KCl, NaCl, NaSO4, NaNO3, is added. If the support electrolyte has a high concentration 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 these will migrate. A rule of thumb is that the support electrolyte should have a hundred times higher concentration. Since this eect is suppressed so easily, migration is often negligible. Another eect is the capacitive charging current or short capacitive current. As mentioned the electrochemical double layers acts as capacitor. Capacitors store charge. A simple capacitor is the plate capacitor. It comprises two conducting parallel plates that are not in contact with each other. If a power source is connected to the plates, a current ows that is exponentially decaying until it is insignicant. A current ows because one plate is charged negative and the other positive. The separation of charges means current ows. At some point the plates cannot store more charge and the current stops owing. The current decays over time according to

c E − t 0 − t I = e RC = I e RC (2.7) R Ec is the charging potential or voltage, I0 is the starting current, R is the resistance of the circuit around the capacitor, t the time and C the capacity of the capacitor. The capacity is a property of the capacitor and is dened as the charge Q that can be stored per applied potential E or as equation Q C = (2.8) E Usually U is used for voltage, but since these equations need to be transferred to electrochemical experiments, it is useful to start with the potential E instead of the voltage U. These two are not synonymous but in this context it is ne to exchange them. If it is 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 resistance R and capacity C are the slower it will be decaying. The product of resistance R and capacity C is often called the time constant τ. 2. The charge Q that can be stored is proportional to the applied potential. Every time the charge Q that can be stored changes a current I ows until the charge Q is adjusted. The charge Q that can be stored changes, if the potential E is changing. This is expressed in the equation

∂Q ∂E I = = C (2.9) ∂t ∂t

3. In equation 2.8 it is shown implicitly and in 2.9 explicitly that the higher the capacity C is the more capacitive current will ow if the potential changes. Usually electrochemists are interested in the Faraday current, that is the current caused by an electrochemical reaction; the capacitive current, caused by physics, is an unwanted side eect. What does this mean for measurements? If the potential of the electrode is changed, 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 Faraday − 1 current decays with t 2 . This means that the capacitive current decays much faster than the Faraday current (see Figure 2.5). The higher the capacity C the higher the capacitive current. The capacity C for a plate capacitor can be calculated with A C = ε ε (2.10) 0 r d where ε0 is the electric eld constant, εr is the relative permittivity of the medium between the plates, d is the distance between the two plates and A is the surface area of the two plates. So most factors inuencing the capacity cannot be altered in an electrochemical experiment. The constant ε0 cannot be changed. The distance d and the relative permittivity εr can only be changed by changing the solution, because d is dened by the distance between inner and outer Helmholtz plane (see Figure 2.4). The area A is inuenced 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 can reduce 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

2.3.1 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 retort stand retort clamp cell top beaker (electrochemical cell)

2.3.2 Chemicals

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

1 mM K3[F e(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 holes into the clamp. Insert the counter electrode (platinum wire), the working electrode (platinum disc in Teon) and the reference electrode, the silver / silver chloride electrode (the electrode with the liquid lled glass body), into the three holes of the cell top. Please be aware that the part with the diaphragm of the reference electrode always has to be in solution or otherwise protected from drying. During evaporation of solution little crystals form in the pores of the frit and the electrode will be blocked. This should be avoided.

2. Fill the beaker with 0.1 M KCl solution. Position the electrodes in a way that all 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 the blue cable, and the working electrode to the red cable. 3. Start PSTrace, choose the scientic mode in the upper left corner and select your potentiostat from the drop down menu. If it is not in the list, press the refresh button (green arrows). Maybe your computer takes a while to install the USB driver. 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 the Curve menu. 6. With this curve you have got the data in respect to the capacitive current of the electrochemcial double layer.

The Cottrell Experiment

1. The setup is the same as in the previous experiment. Except for the solution in

the beaker which is now 1 mM K3[F e(CN)6] in 0.1 M KCl solution. 2. Use the same parameters for the Multistep Amperometry as before. (Load le Exp1_2.psmethod) 3. Before you start the measurement choose in the drop down menu next to the Run button the option Overlay instead of New. If the curve of the capacitive current you recorded before is no longer displayed load it. Start the measurement. The 3− − 4− reaction at the electrode is: [F e(CN)6] + e → [F e(CN)6] 4. Save the new curve as well. 5. The capacitive current as well as the sum of the capacitive and Faraday current can be compared with each other. They should meet the expectations that were set during chapter 2.2.2. 6. The next step should be a better time resolution. Look at the maximum current that was measured during step 3. Choose the smallest current range which value doubled is bigger than the maximum current. The part of the curve before the step shows when a stable current is reached. For the equilibration time choose a

time that is big enough to reach a constant current. Choose for the t1 5 s and t2 65 s. Reduce the tinterval to 0.001 s. 7. Choose New in the drop down menu next to the start button and start the measurement. 8. Save the curve.

− 1 9. To nd the area of the electrode a plot of the absolute I vs t 1 and a linear regression are needed. Export the data to software that can do this, for example Excel, 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- − 1 − 1 wards calculate t 2 and plot I vs t 2 . 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 area of the electrode. Remember that the slope is described by equation 2.6. Although 3 D can be found in literature, it is not easy. We recommend to use D([F e(CN)6] in 0.1 M KCl) = (7.17 ± 0.18)10−6cm2/s found in [5]. 12. Discuss the dierence between expectations and measured values. What could explain 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. The electrodes and cell should be cleaned with demineralized water. Otherwise the drying solution will leave salt stains. Solution containing iron need to be collected in a container for liquid waste containing heavy metal. The pH of the collected heavy metal solution should be alkaline (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 followed towards potentiodynamic experiments may have been even more important for modern electrochemistry. Potentiodynamic experiments made it easy to collect all the data needed for a plot of current I versus potential E. These plots are called a voltammogram and the technique used for measuring is called voltammetry. In a short period of time 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. Experienced electrochemists read quite some information from the shape of a CV. The experiment and its discussion should teach: 1. What is a cyclic voltammetry? 2. Why does a cyclic voltammogram show a specic shape? 3. What does the cyclic voltammogram say about reversibility? 4. How can a cyclic voltammogram help to distinguish adsorbed and free diusing species? 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 is the scan rate v. As can be seen from the mathematical denition (see equation 3.1) this is the slope of the linear potential.

∂E v = (3.1) ∂t At the start the potential is usually in a region where no electrochemical reaction is occurring. The linear sweep of the potential is usually chosen in such a way that the potential crosses the formal potential of the investigated species (see Figure 3.1). After reaching a set potential the slope of the linear potential is inverted, that is a decreasing becomes an increasing potential and vice versa. This potential is called the vertex potential. 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 is often to observe the stability of a system. Modern software usually oers the option to choose two vertex potentials and a start potential, that is the potential sweeps between the two vertex potentials and starts at a potential between these two.

Figure 3.1: potential vs time during a cyclic voltammetry with indicated vertex potentials ( ), formal potential of the investigated species ( 0 ), and scan Evertex E0 rate (v).

The prediction of the current signal has to take a few inuences into consideration. Sweeping the potential will change the composition of the redox active species in front of the working electrode according to the Nernst equation (equation 0.6). This change is done by oxidation or reduction and leads to a current owing. If there would be no diusion limitation, the measured current would have a sigmoidal shape (see Figure 3.2), but in real systems there will be diusion limitation. If a species is consumed at an electrode, this species is depleted around the electrode and less species reaches the electrode, thus less current ows. If the potential is that high that the concentration of the consumed species is 0 at the electrode the depletion will lead to a current following the Cottrell equation (equation 2.5). As long as the potential is the limiting factor an increasing potential leads to an increasing current. As soon as diusion becomes the limiting factor the current decreases. As a result the current during the sweep has a maximum current or peak current (see Figure 3.2). When the scan rate is inversed, the same events happen for the opposite reaction, if the system is reversible. To directly read the potentials corresponding to the peak, usually a voltammogram, a curve of I vs E, is plotted. This way many important parameters can be determined faster than by plotting the E vs t and I vs t on top of each other as 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 to reversible systems, where both species have the same diusion coecient. As in chapter 2.2.1, only the contribution due to the faraday current was taken into consideration, 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 equation (pale violet) and Cottrell equation (pale blue)

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

As seen from equation 2.9 the constant change of the potential described by the scan rate v will induce a constant capacitive current. This current will change positive or negative depending on the scan's direction. If no species that is reduced or oxidized in the potential range of the CV is present, the voltammogram will only show the capacitive current (see Figure 3.4). A measurement like this allows the determination of 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 inuence that could be visible is the resistance. If a high resistance is present in the system, the current will be signicantly inuenced by Ohm's law. If only a resistor is connected to the potentiostat, the resulting CV would be a diagonal line because the current 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 resistance at a level that is barely noticed in a CV. If an electrode is smooth the contribution of the capacitive current is also rather small compared to the Faraday current, but still it can be important when small peak currents are measured.

Figure 3.4: schematic CVs showing the contribution of a Faraday current (green) and a 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 containing one free diusing reversible species. The peak data are the anodic peak current Ip,a and peak potential Ep,a as well as their cathodic counterparts Ip,c and Ep,c. Much qualitative information can be gathered from the shape of a CV. More important information can be gathered either directly from the peak data or observing the peaks changing during a series of measurements. Please notice that the baseline for the back scan is more dicult to determine, because the current at the start of the back scan also includes the diusion limited current of the forward scan. In the instructions' part methods will be described to take this into consideration. If the change of peak current is observed for quantitative analysis only, the oset caused by ignoring the base lines is not relevant.

Figure 3.5: CV of a free diusing reversible species with indication of the peak currents and 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, is usually done by observing the change of the peak current. The peak current is proportional to the concentration of the species inside (see equation 3.4). This makes it easy to perform a series of measurements to create a calibration curve or perform a standard addition (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 and one peak for the oxidation at the anodic sweep. If one of these peaks is missing, the system is not reversible under the given conditions. There are two dierent reasons for an irreversible system. The species is simply not reversible, because the back reaction is blocked kinetically or thermodynamically or the converted species has a follow up reaction 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 investigating the 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 soon as the scan rate is so fast that the back reaction is initiated before the reaction usually consuming the converted species starts, the peak of the back reaction gets visible. So the 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 obvious reason is the device. Each potentiostat, AD/DA converter, etc. has a limited amount of values it can process per second, for example the EmStat3 has a maximum of 1000 points per second. This could be compensated by increasing the distance between each point of the measurement. Usually the potentiostat averages some values before returning them as one point. The higher the scan rate, the less time is used per step, the fewer points can be averaged. As a result random noise is not averaged out. Another limitation is the rise time that is the minimum time needed for the potentiostat to adjust the potential is not innitely small. The EmStat can still manage a scan rate of 5V/s with a step potential of 5mV . Another limit is set by the capacitive current. As seen from equation 3.2 the capacitive current will grow linear with the scan rate. The Faraday current will also increase; since the same species is converted in a shorter time the current is higher. The diusion 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 between the peak current Ip of a free diusing species and the scan rate is:

√ Ip ∝ v (3.5) This means that the capacitive current increases faster than the peak current with increasing scan rate. The capacitive current will dominate the CV above a certain scan rate and the shape and peaks caused by the faraday current cannot be identied anymore. If experiments demand very high scan rates, ultramicroelectrodes are often used. These very small electrodes have very small capacities and thus high scan rates can be used. Equation 3.5 can provide important information for an electrochemical system indirectly. This equation is only valid for free diusing species. The peak shape and peak current for adsorbed species are dierent. Due to the fact that there is no diusion from 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 the peak current of an adsorbed species Ip(ads) shows the following behavior:

Ip(ads) ∝ v (3.6) A series of measurements using dierent scan rates and plotting the measured peak 1 currents versus v and v 2 will show if the active species is surface bound or free diusing. 1 If the plot versus v 2 is linear, the peak belongs to a free diusing species and if the plot versus v is linear, the species is adsorbed (see Figure 3.6 b and c). This is useful, because most systems have some backgrounds currents that will make it dicult to recognize if a peak is symmetrical or not. Figure 3.6: a) schematic representation of the dierent peak shapes in a voltammogram due to adsorbed or free diusing species b) schematic plot of the peak cur-

rents Ip of dierent scan rates v versus the scan rate c) like b but versus 1 v 2

Peak shapes can sometimes give information about surface and species surroundings. A completely homogeneous surface, for example the surface of a perfect single crystal, and a completely homogeneous species, for example a monolayer of ferrocene, should deliver a very sharp peak. An experienced electrochemist can actually read a complete CV of a well-known system like others can read an infrared spectrum. Another parameter that delivers information about the system is the peak potential. A CV allows the fast determination of the redox species' formal potential, although sometimes only approximately. Assuming that your reaction is controlled by diusion and potential only and there is no kinetic inhibition, the reaction happen as soon as the potential for an oxidation or reduction is reached. The reactions would be time independent and an increasing scan rate would have no eect on the peak potential. The peak current would increase according to equation 3.5 or 3.6, due to the fact that current is passed charge per time. If the reduced and the oxidized species have the same diusion coecient in the used solution the formal potential should be exactly between the two peak potentials of reduction and oxidation as equation 3.7 describes.

0 (E − E ) E0 = p,a p,c (3.7) 2 00 If the diusion coecients are dierent, the formal potential E is still between Ep,a and Ep,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 the kinetics will never limit the reaction. In an ideal situation the minimum peak separation, so the dierence between Ep,a and Ep,c is:

59mV ∆E = (3.8) p z As before z is the number of electrons transferred during the reaction. If there is a kinetic inhibition, 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 the amount of time is similar to the time it takes to perform a scan, the kinetic inhibition can inuence 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 with constant 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 are yellow balls and a row of red balls behind the yellow ones. The rst person has super-fast reexes and you tell that person to grab the rst red ball right in front of him/her. The belt starts moving and as soon as the rst red ball reaches the person he or she grabs the ball. You increase the speed of the belt several times, but due to the fast reexes the rst red ball is always grabbed. The second person is slower and has the same task. He or she will take a second to realize that there is a red ball. The belt starts moving and this 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 red ball. 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 the scan rate. A system showing this behavior is considered quasi-reversible. An irreversible system has, as mentioned before, either an anodic or cathodic peak. Examples of this are catalytic processes which will be explained in the next chapter (??). 3.3 The Experiments

3.3.1 Devices and Equipment

EmStat3 or PalmSens3 Sensor sensor cable sensor conector

computing unit (PC, Laptop, notebook, tablet PC (android)) potentiostat software (PStrace4) calculation and plotting software (Excel, MatLab) counter electrode reference electrode working electrode ItalSens IS-Au retort stand retort clamp cell top beaker (electrochemical cell) stirrer

3.3.2 Chemicals

0.1 M KCl solution (prepared from solid KCl and demineralized water) for preparation of K3[F e(CN)6] and K4[F e(CN)6] solutions

Mixed K3[F e(CN)6]and K4[F e(CN)6] in 0.1 M KCl solutions with ratios of 100 10 1 1 1 1 2 K3[F e(CN)6]/K4[F e(CN)6] = 1 , 1 , 1 , 10 , 100 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.

1 All these solutions should have the same sum of concentrations, that is c(K3[F e(CN)6]) + c(K4[F e(CN)6]) = constant 2A concentration for all solutions of (0.010 ± 0.001) M can be achieved by using 0.1 M solutions of K3[F e(CN)6] and K4[F e(CN)6] in 0.1 M KCl with a pipette (20ml) and variable microliter pipettes (10 − 100l and 100ı1000l) 3.3.3 Procedure

Determination of the Formal Potential

1. Make K3[F e(CN)6] and K4[F e(CN)6] in 0.1 M KCl solutions with ratios of 100 10 1 1 1 K3[F e(CN)6]/K4[F e(CN)6] = 1 , 1 , 1 , 10 , 100 2. Choose from the techniques drop down menu the Potentiometry / OCP. Choose 300s as trun (duration of the measurement) and 1s as tinterval (time between measurement points). 3. Immerse all three electrodes (reference, counter, and working electrode) into each one of the solutions. Press start and wait until the open circuit potential is stable. Note the solution ratio and the corresponding open circuit potential. Stop the measurement. 4. Repeat step 3 for each of the solutions. Rinse the electrodes carefully between changing the solutions.

5. Prepare a 5 mM K4[F e(CN)6] in 0.1 M KCl solution, for example by adding 1mL of 0.1M K4[F e(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 determined via the Nernst equation (equation 0.6).

0 0 E = E0 + RT 2.3log cOx = E0 + 0.059V log cOx zF cRed z cRed

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

4− 3− Voltammetry: Due to the fact that [F e(CN)6] and[F e(CN)6] have the same diusion coecient, determining the formal potential is simple. Read the peak

potentials for the anodic and cathodic peaks Ep,a and Ep,c from the CV. Their average is the formal potential. Identication of a Reversible, Quasi-reversible, or Irreversible System

1. Prepare a 5 mM K4[F e(CN)6] in 0.1 M KCl solution, for example by adding 2 mL of 0.1 M K4[F e(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, 0.25 V/s, and 0.5 V/s. Save each curve. 5. The goal is to use Table 3.1 to identify if [F e(CN)6]4− and [F e(CN)6]3− is a reversible, quasi-reversible, or irreversible system. Using the list of criteria dierent parameters need to be extracted from the CV and compared.

Extracting the Ip,c is not directly possible due to the fact that there is no direct baseline available. When the scan rate is inversed there is still current owing due to conversion of the [F e(CN)6]4−. The current overlaps with the baseline of the cathodic sweep, so either a graphical or mathematical extrapolation has to be done. The graphical method is rather dicult, since parts of the curve has to be guessed. The mathematical way just

needs a single equation and some parameters from the curve. Determine Ip,c,0 and Ivertex according to Figure 3.7 and use equation3.9 to determine Ip,c.

Table 3.1: diagnostic criteria for reversible, quasi-reversible and electrochemical irreversible 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)

00 00 E = 0, 5(Ep,a + Ep,c) E = 0, 5(Ep,a + Ep,c) - (DOx = DRed)(DOx = DRed)

1 1 Ip,a ∝ v 2 - Ip,a ∝ v 2 1 1 Ip,a = const1 v 2 Ip,a = const2 v 2

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, but from a scientic point of view gold has quite a number of special properties. Under the right conditions aurides can be made, that is salts where gold is present as the anion Au-. Working with 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 is soft and hydrophobic. As a result a gold surface collects all organic residues in the air and gets dirty rather quickly. A popular method to clean electrodes is electropolishing. In this method high anodic potentials are applied and adsorbed species will desorb. Furthermore gold forms a gold oxide monolayer, that is the gold is covered with a single layer of oxidized gold. The reduction peak is very sharp and in sulfuric acid no other peak overlaps with it. The charge of the 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 the electrode to the sensor connector take the strip of electrodes and hold it at the end with the single letter (A to D) and a number (1 to 20). Since the electrodes are easily contaminated or destroyed, please take care not to touch it with bare hands. Cut away the access plastic around one electrodes SPE and place it in the connector. You must cut the sensor narrow enough for tting it into the connector, but please take care not to 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 the menu 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, 0.25 V/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 reduced equation 3.6 should be valid. Look at the reduction peak and plot Ip,c versus v to proof or falsify.

9. The next steps are used to calculate the real surface. Take the CV with v = 0.1 V/s to calculate the surface. First take the area of the reduction peak Ap,c, which is also the integral. To get the charge divide by the scan rate. This is also represented in equation 3.10. A Q = p,c (3.10) v 10. The real surface is determined by using the charge per surface area for a gold oxide layer q. This value can be found in the literature: 3.84C/m2. 3 The real surface Areal is calculated according to Q A = (3.11) real q

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

A r = real (3.12) Ageo

3.3.4 Disposal of Chemicals

The potassium chloride solution is harmless and can be disposed in the sink. The electrodes and cell should be cleaned with demineralized water. Otherwise the drying solution will leave salt stains. Solution containing iron need to be collected in a container for liquid waste containing heavy 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 waste containing 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. Diusion coecients determined with microelec- trodes. Journal of Electroanalytical Chemistry, 305(1):7381, 1991.