Undergraduate Instrumental Analysis, Sixth Edition Simpo PDF Merge and Split Unregistered Version hopdf com 15 Electroanalytical Chemistry Electrochemistry is the area of chemistry that.Undergraduate Instrumental Analysis, Sixth Edition Simpo PDF Merge and Split Unregistered Version hopdf com 15 Electroanalytical Chemistry Electrochemistry is the area of chemistry that.
Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com 15 Electroanalytical Chemistry Electrochemistry is the area of chemistry that studies the interconversion of chemical energy and electrical energy Electroanalytical chemistry is the use of electrochemical techniques to characterize a sample The original analytical applications of electrochemistry, electrogravimetry and polarography, were for the quantitative determination of trace metals in aqueous solutions The latter method was reliable and sensitive enough to detect concentrations as low as ppm of many metals Since that time, many different types of electrochemical techniques have evolved, each useful for particular applications in organic, inorganic, and biochemical analyses A species that undergoes reduction or oxidation is known as an electroactive species Electroactive species in general may be solvated or complexed, ions or molecules, in aqueous or nonaqueous solvents Electrochemical methods are now used not only for trace metal ion analyses, but also for the analysis of organic compounds, for continuous process analysis, and for studying the chemical reactions within a single living cell Applications have been developed that are suited for quality control of product streams in industry, in vivo monitoring, materials characterization, and pharmaceutical and biochemical studies, to mention a few of the myriad applications Under normal conditions, concentrations as low as ppm can be determined without much difficulty By using electrodeposition and then reversing the current, it is possible to extend the sensitivity limits for many electroactive species by three or four orders of magnitude, thus providing a means of analysis at the ppb level In practice, electrochemistry not only provides a means of elemental and molecular analysis, but also can be used to acquire information about equilibria, kinetics, and reaction mechanisms from research using polarography, amperometry, conductometric analysis, and potentiometry The analytical calculation is usually based on the determination of current or voltage or on the resistance developed in a cell under conditions such that these are dependent on the concentration of the species under study Electrochemical measurements are easy to automate because they are electrical signals The equipment is often far less expensive than spectroscopy instrumentation Electrochemical techniques are also commonly used as detectors for LC, as discussed in Chapter 13 Coauthor: R.J Gale, Department of Chemistry, Louisiana State University, Baton Rouge, LA 919 920 Chapter 15 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com 15.1 FUNDAMENTALS OF ELECTROCHEMISTRY Electrochemistry is the study of reduction –oxidation reactions (called redox reactions) in which electrons are transferred from one reactant to another A chemical species that loses electrons in a redox reaction is oxidized A species that gains electrons is reduced A species that oxidizes is also called a reducing agent because it causes the other species to be reduced; likewise, an oxidizing agent is a species that is itself reduced in a reaction An oxidation – reduction reaction requires that one reactant gain electrons (be reduced) from the reactant which is oxidized We can write the reduction and the oxidation reactions separately, as half-reactions; the sum of the half-reactions equals the net oxidation reduction reaction Examples of oxidation half-reactions include: Fe2ỵ ! Fe3ỵ þ eÀ Cu(s) À! Cu2þ þ 2eÀ AsH3 (g) À! As(s) ỵ 3Hỵ ỵ 3e H2 C2 O4 ! 2CO2 (g) ỵ 2Hỵ ỵ 2e Examples of reduction half-reactions include: Co3ỵ ỵ e ! Co2ỵ (IO3 ) ỵ 6Hỵ þ 5eÀ À! I2 (s) þ 3H2 O Cl2 (g) ỵ 2e ! 2Cl Agỵ ỵ e ! Ag(s) If the direction of an oxidation reaction is reversed, it becomes a reduction reaction; that is, if Al3ỵ accepts electrons, it is reduced to Al(s) All of the reduction reactions are oxidation reactions if they are written in the opposite direction Many of these reactions are reversible in practice, as we shall see A net oxidation – reduction reaction is the sum of the appropriate reduction and oxidation half-reactions If necessary, the half-reactions must be multiplied by a factor so that no electrons appear in the net reaction For example, the reaction between Cu(s), Cu2ỵ, Ag(s), and Agỵ is: Cu(s) ỵ 2Agỵ ! Cu2ỵ ỵ 2Ag(s) We shall see why the reaction proceeds in this direction shortly The net reaction is obtained from the half-reactions as follows: Oxidation reaction: Cu(s) ! Cu2ỵ ỵ 2e Reduction reaction: Agỵ ỵ e À! Ag(s) Each mole of copper gives up moles of electrons, while each mole of silver ion accepts only mole of electrons Therefore the entire reduction reaction must be multiplied by 2, so that there are no electrons in the net reaction after summing the half-reactions: Oxidation reaction: Cu(s) ! Cu2ỵ ỵ 2e Reduction reaction: 2(Agỵ ỵ e ! Ag(s)) Net reaction: Cu(s) ỵ 2Agỵ ! Cu2ỵ 2Ag(s) The equal numbers of electrons on both sides of the arrow cancel out Electroanalytical Chemistry 921 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com Electrochemical redox reactions can be carried out in an electrochemical cell as part of an electrical circuit so that we can measure the electrons transferred, the current, and the voltage Each of these parameters provides us with information about the redox reaction, so it is important to understand the relationship between charge, voltage, and current The absolute value of the charge of one electron is 1.602  10219 coulombs (C); this is the fundamental unit of electric charge Since 1.602  10219 C is the charge of one electron, the charge of one mole of electrons is: (1:602  10À19 C/eÀ )(6:022  1023 eÀ =mol) ¼ 96,485 C/mol (15:1) This value 96,485 C/mol is called the Faraday constant (F), and provides the relationship between the total charge, q, transferred in a redox reaction and the number of moles, n, involved in the reaction q¼nÂF (15:2) In an electric circuit, the quantity of charge flowing per second is called the current, i The unit of current is the ampere, A; A equals C/s The potential difference, E, between two points in the cell is the amount of energy required to move the charged electrons between the two points If the electrons are attracted from the first point to the second point, the electrons can work If the second point repels the electrons, work must be done to force them to move Work is expressed in joules, J, and the potential difference, E, is measured in volts The relationship between work and potential difference is: w (in joules) ¼ E (in volts)  q (in coulombs) (15:3) Since the unit of charge is the coulomb, V equals J/C The relationship between current and potential difference in a circuit is expressed by Ohm’s Law: i¼ E R (15:4) where i is the current; E, the potential difference, and R, the resistance in the circuit The units of resistance are V/A or ohms, V 15.2 ELECTROCHEMICAL CELLS At the heart of electrochemistry is the electrochemical cell We will consider the creation of an electrochemical cell from the joining of two half-cells When an electrical conductor such as a metal strip is immersed in a suitable ionic solution, such as a solution of its own ions, a potential difference (voltage) is created between the conductor and the solution This system constitutes a half-cell or electrode (Fig 15.1) The metal strip in the solution is called an electrode and the ionic solution is called an electrolyte We use the term electrode to mean both the solid electrical conductor in a half-cell (e.g., the metal strip) and the complete half-cell in many cases, for example, the standard hydrogen electrode, the calomel electrode Each half-cell has its own characteristic potential difference or electrode potential The electrode potential measures the ability of the half-cell to work, or the driving force for the half-cell reaction The reaction between the metal strip and the ionic solution can be represented as M0 ! Mnỵ ỵ ne (15:5) 922 Chapter 15 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com Figure 15.1 A half-cell composed of a metal electrode M0 in contact with its ions, Mỵ, in solution The salt bridge or porous membrane is shown on the lower right side where M0 is an uncharged metal atom, Mnỵ is a positive ion, and e2 is an electron The number of electrons lost by each metal atom is equal to n, where n is a whole number This is an oxidation reaction, because the metal has lost electrons It has been oxidized from an uncharged atom to a positively charged ion In the reaction, the metal ions enter the solution (dissolve) By definition, the electrode at which oxidation occurs is called the anode We say that at the anode, oxidation of the metal occurs according to the reaction shown in Eq (15.5) Some examples of this type of half-cell are: Cd(s) ! Cd2ỵ ỵ 2e Ag(s) ! Agỵ ỵ e Cr(s) ! Cr3ỵ ỵ 3e Note that in normal usage, the zero oxidation state of the solid metal is understood, not shown with a zero superscript It has been found that with some metals the spontaneous reaction is in the opposite direction and the metal ions tend to become metal atoms, taking up electrons in the process This reaction can be represented as Mnỵ ỵ ne ! M0 (15:6) This is a reduction reaction because the positively charged metal ions have gained electrons, lost their charge, and become neutral atoms The neutral atoms deposit on the electrode, a process called electrodeposition This electrode is termed a cathode At the cathode, reduction of an electroactive species takes place An electroactive species is one that is oxidized or reduced during reaction Electrochemical cells also contain nonelectroactive (or inert) species such as counterions to balance the charge, or electrically conductive electrodes that not take part in the reaction Often these inert electrodes are made of Pt or graphite, and serve only to conduct electrons into or out of the half-cell It is not possible to measure directly the potential difference of a single half-cell However, we can join two half-cells to form a complete cell as shown in Fig 15.2 In this example, one half-cell consists of a solid copper electrode immersed in an aqueous solution of CuSO4 ; the other has a solid zinc electrode immersed in an aqueous solution Electroanalytical Chemistry 923 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com Figure 15.2 A complete Zn/Cu galvanic cell with a salt bridge separating the half-cells of ZnSO4 The two half-cell reactions and the net spontaneous reaction are shown: Anode (oxidation) reaction: Zn(s) ! Zn2ỵ ỵ 2e Cathode (reduction) reaction: Cu2ỵ ỵ 2e ! Cu(s) Net reaction: Zn(s) ỵ Cu2ỵ ! Zn2ỵ ỵ Cu(s) No reaction will take place, and no current will flow, unless the electrical circuit is complete As shown in Fig 15.2, a conductive wire connects the electrodes externally through a voltmeter (potentiometer) A salt bridge, a glass tube filled with saturated KCl in agar gel, physically separates the two electrolyte solutions The salt bridge permits ionic motion to complete the circuit while not permitting the electrolytes to mix The reason we need to prevent the mixing of the electrolytes is that we want to obtain information about the electrochemical system by measuring the current flow through the external wire If we had both electrodes and both ionic solutions in the same beaker, the copper ions would react directly at the Zn electrode, giving the same net reaction but no current flow in the external circuit In the electrolyte solution and the salt bridge the current flow is ionic (ion motion), and in the external circuit the current flow is electronic (electron motion) This cell and the one in Fig 15.3 show the components needed for an electrochemical cell: two electrical conductors (electrodes), suitable electrolyte solutions and a means of allowing the movement of ions between the solutions (salt bridge in Fig 15.2, a semipermeable glass frit or membrane in Fig 15.3), external connection of the electrodes by a conductive wire and the ability for an oxidation reaction to occur at one electrode, and a reduction reaction at the other Of course there are counterions present in each solution (e.g., sulfate ions) to balance the charge; these ions are not electroactive and not take part in the redox reaction They flow (ionic motion) to keep charge balanced in the cell A cell that uses a spontaneous redox reaction to generate electricity is called a galvanic cell Batteries are examples of galvanic cells A cell set up to cause a nonspontaneous reaction to occur by putting electricity into the cell is called an electrolytic cell The complete cell has a potential difference, a cell potential, which can be measured by the voltmeter The potential difference for the cell, Ecell , can be considered to be equal 924 Chapter 15 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com Figure 15.3 A schematic complete cell with a porous frit separating the half-cells to the difference between the two electrode potentials, when both half-cell reactions are written as reductions That is, Ecell ¼ Ecathode À Eanode (15:7) when the potentials are reduction potentials This convention is necessary to calculate the sign of Ecell correctly, even though the anode reaction is an oxidation The cell potential is also called the electromotive force or emf While we cannot measure a given single electrode potential directly, we can easily measure the cell potential for two half-cells joined as described So now we have a means of measuring the relative electrode potential for any half-cell by joining it to a designated reference electrode (reference half-cell) In real cells, there is another potential difference that contributes to Ecell , called a junction potential If there is a difference in the concentration or types of ions of the two half-cells, a small potential is created at the junction of the membrane or salt bridge and the solution Junction potentials can be sources of error When a KCl salt bridge is used the junction potential is very small because the rates of diffusion of Kỵ and Cl2ions are similar, so the error in measuring a given electrode potential is small 15.2.1 Line Notation for Cells and Half-Cells Writing all the equations and arrows for half-cells and cells is time consuming and takes up space, so a shorthand or line notation is often used For example, the half-cell composed of a silver electrode and aqueous 0.0001 M Agỵ ion (from dissolution of silver nitrate in water) is written in line notation as Ag(s) j Agỵ (0:0001 M) The vertical stroke or line between Ag and Agỵ indicates a phase boundary, that is, a difference in phases (e.g., solid j liquid, solid j solid 2, or liquid j gas) in the constituents of the half-cells that are in contact with each other The complete cell in Fig 15.2 can be represented as Zn(s) j Zn2ỵ (0:01 M) k Cu2ỵ (0:01 M) j Cu(s) Electroanalytical Chemistry 925 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com The double vertical stroke after Zn2ỵ indicates a membrane junction or salt bridge The double stroke shows the termination of one half-cell and the beginning of the second This cell could also be written to show the salts used, as shown: Zn(s) j ZnSO4 (aq) k CuSO4 (aq) j Cu(s) It is conventional to write the electrode that serves as the anode on the left in an electrochemical cell The other components in the cell are listed as they would be encountered moving from the anode to the cathode 15.2.2 Standard Reduction Potentials 15.2.2.1 The Standard Hydrogen Electrode In order to compile a table of relative electrode potentials, chemists must agree upon the half-cell that will serve as the reference electrode The composition and construction of the half-cell must be carefully defined The value of the electrode potential for this reference half-cell could be set equal to any value, but zero is a convenient reference point In practice, it has been arbitrarily decided and agreed upon that the standard hydrogen electrode (SHE) has an assigned electrode potential of exactly zero volts at all temperatures The SHE, shown in Fig 15.4, consists of a platinum electrode with a surface coating of finely divided platinum (called a platinized Pt electrode) immersed in a solution of M hydrochloric acid, which dissociates to give Hỵ Hydrogen gas, H2 , is bubbled into the acid solution over the Pt electrode The finely divided platinum on the electrode surface provides a large surface area for the reaction 2Hỵ ỵ 2e ! H2 (g) E0 ẳ 0:000 V (15:8) Figure 15.4 The standard hydrogen electrode (SHE) This design is shown with a presaturator containing the same M HCl solution as in the electrode to prevent concentration changes by evaporation (Aikens et al., by permission, Waveland Press Inc., Long Grove, IL, 1984 All rights reserved.) 926 Chapter 15 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com In addition, the Pt serves as the electrical conductor to the external circuit Under standard state conditions, that is, when the H2 pressure equals atm and the ideal concentration of the HCl is M, and the system is at 258C, the reduction potential for the reaction given in Eq (15.8) is exactly V (The potential actually depends on the chemical activity of the HCl, not on its concentration The relationship between activity and concentration is discussed subsequently For an ideal solution, concentration and activity are equal.) The potential is symbolized by E 0, where the superscript zero means standard state conditions The term standard reduction potential means that the ideal concentrations of all solutes are M and all gases are at atm; other solids or liquids present are pure (e.g., pure Pt solid) By connecting the SHE half-cell with any other standard half-cell and measuring the voltage difference developed, we can determine the standard reduction potential developed by the second half-cell Consider, for example, a cell at 258C made up of the two half-cells: Zn(s) j Zn2ỵ (aq, M) k Hỵ (aq, M) j H2 (g, atm) j Pt(s) This cell has a Zn half-cell as the anode and the SHE as the cathode All solutes are present at ideal M concentrations, gases at atm and the other species are pure solids, and so both half-cells are at standard conditions The measured cell emf is ỵ0.76 V and this is the standard cell potential, E 0, because both half-cells are in their standard states From Eq (15.7), we can write: 0 ¼ Ecathode À Eanode Ecell (15:9) The total voltage developed under standard conditions is ỵ0.76 V But the voltage of the SHE is by definition; therefore the standard reduction potential of the Zn half-cell is: þ0:76 V ¼ 0:000 À EZn EZn ¼ À0:76 V Therefore we can write: Zn2ỵ ỵ 2e ! Zn(s) E0Zn ¼ À0:76 V We have determined the Zn standard reduction potential even though the galvanic cell we set up has Zn being oxidized By substituting other half-cells, we can determine their electrode potentials (actually, their relative potentials) and build a table of standard reduction potentials If we set up a galvanic cell with the SHE and Cu, we have to make the SHE the anode in order for a spontaneous reaction to occur This cell, Pt(s) j H2 (g, atm) j Hỵ (aq, M) k Cu2ỵ (aq, M) j Cu(s) has a measured cell emf ¼ þ0.34 V Therefore the standard reduction potential for Cu is þ0:34 V ¼ E0Cu À 0:000 V ¼ þ0:34 V ECu The quantity E is the emf of a half-cell under standard conditions A half-cell is said to be under standard conditions when the following conditions exist at a temperature of 258C: All solids and liquids are pure (e.g., a metal electrode in the standard state) All gases at a pressure of atm (760 mmHg) All solutes are at M concentration (more accurately, at unit activity) Electroanalytical Chemistry 927 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com The true electrode potential is related to the activity of the species in solution, not the concentration For a pure substance, its mole fraction and its activity ¼ For a pure substance that is not present in the system, its mole fraction and its activity ¼ From general chemistry, you should remember that Raoult’s Law predicts that for an ideal solution, the mole fraction of the solute and its activity are equal However, most solutions deviate from linearity because of interactions between the solute and solvent molecules These deviations can be positive or negative, because the species may attract or repel each other The amount of attraction or repulsion affects the activity of the solute For dilute solutions, the activity is proportional to concentration (Henry’s Law) Activity is equal to the concentration times the activity coefficient for the species in solution That is aion ẳ ẵMion gion (15:10) where aion is the activity of given ion in solution; [Mion], the molar concentration of the ion; and gion , the activity coefficient of the ion Activity depends on the ionic strength of the solution If we compare a M solution and a 0.01 M solution, the more concentrated solution may act as though it is less than 100 more concentrated than the dilute solution It is then said that the activity of the M solution is less than unity or the activity coefficient is less than For solutions with positive deviations from Henry’s Law, the activity coefficient will be greater than For very dilute solutions (low ionic strength) the activity coefficient g approaches 1, so concentration is approximately equal to activity for very dilute solutions We will use concentrations in the calculations in this text instead of activities, but the approximation is only accurate for dilute solutions (,0.005 M) and for ions with single charges Details on activity corrections can be found in most analytical chemistry texts, such as the ones by Harris or Enke listed in the bibliography Looking back at our cell in Fig 15.2, with the Zn half-cell as the anode and the Cu half-cell as the cathode, we can calculate the standard cell potential for this galvanic cell In the spontaneous reaction, Zn is oxidized and Cu2ỵ is reduced, therefore Zn is the anode and Cu is the cathode Zn2ỵ ỵ 2e ! Zn0 E0 ẳ 0:76 V Cu2ỵ ỵ 2e ! Cu0 E0 ẳ ỵ0:34 V The standard cell potential developed is calculated from Eq (15.9): ỵ 0:34 (0:76 V) ẳ ỵ1:10 V In tables of standard potentials, all of the half-cell reactions are expressed as reductions The sign is reversed if the reaction is reversed to become an oxidation In a spontaneous reaction, when both half-cells are written as reductions, the half-cell with the more negative potential will be the one that oxidizes The negative sign in Eq (15.9) reverses one of the reduction processes to an oxidation Some standard reduction potentials for common half-cells are given in Appendix 15.1 These can be used to calculate E for other electrochemical cell combinations as we have done for the Zn/Cu cell More complete lists of half-cell potentials can be found in references such as Bard et al., listed in the bibliography In the reaction of our example, zinc metal dissolves, forming zinc ions and liberating electrons Meanwhile, an equal number of electrons are consumed by copper ions, which plate out as copper metal The net reaction is summarized as Zn0 ỵ Cu2ỵ ! Cu0 ỵ Zn2ỵ ỵ1:10 V 988 Chapter 15 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com The addition of surfactants is not always recommended in normal and differential pulse polarography, because their presence may reduce the sensitivity of these methods in certain circumstances Many other forms of polarography have been suggested and tested Prominent among these are AC polarographic methods, which use sinusoidal and other periodic waveforms Most modern instruments offer a range of techniques to the electroanalytical chemist The solvent used plays a very important role in polarography It must be able to dissolve a supporting electrolyte and, if necessary, be buffered In many cases the negative potential limit involves the reduction of a hydrogen ion or a hydrogen on a molecule It is therefore vital to control the pH of the solution with a buffer The buffer normally serves as the supporting electrolyte Furthermore, it must conduct electricity These requirements eliminate the use of many organic liquids such as benzene A polar solvent is necessary, the most popular being water To dissolve organic compounds in water, a second solvent, such as ethanol, acetone, or dioxane, may first be added to the water The mixture of solvents dissolves many organic compounds and can be conditioned for polarography As an alternative to the aqueous/nonaqueous systems, a pure polar solvent may be used, especially if water is to be avoided in the electrolysis Commonly used solvents are acetonitrile, dimethylformamide, dimethylsulfoxide, and propylene carbonate Tetralkylammonium perchlorates and tetrafluoroborates are useful supporting electrolytes because their cations are not readily reduced A list of functional groups that can be determined by polarography is shown in Table 15.6 In general, simple saturated hydrocarbons, alcohols, and amines are not readily analyzed at the DME However, aldehydes and quinones are reducible, as well as ketones Similarly, olefins (in aqueous solutions) may be reduced according to the equilibrium R22C55C22R0 ỵ 2e ỵ 2H2 O ! R22CH2 CH222R0 þ 2OHÀ Table 15.6 Typical Functional Groups that can be Determined by Polarography Functional group RCHO RCOOH RR0 C5 5O R2 2O2 2N5 5O R2 2N5 5O R2 2NH2 R2 2SH Name E1=2 (V) Aldehyde Carboxylic acid Ketone Nitrite Nitroso Amine Mercaptan 21.6 21.8 22.5 20.9 20.2 20.5 20.5 Note: Electrochemical data on a large number of organic compounds are compiled in the CRC Handbook Series in Organic Electrochemistry Electroanalytical Chemistry 989 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com In short, polarography can be used for the analysis of C22N, C22O, N22O, O22O, S22S, and C22S groups and for the analysis of heterocyclic compounds Also, many important biochemical species are electroactive, such as vitamin C (ascorbic acid), fumaric acid, vitamin B factors (riboflavin, thiamine, niacin), antioxidants such as tocopherols (vitamin E), N-nitrosamines, ketose sugars (fructose and sorbose), and the steroid aldosterone 15.3.5 Voltammetry We mentioned that polarography is a form of voltammetry in which the electrode area does not remain constant during electrolysis It is possible, however, to use electrode materials other than mercury for electroanalyses, provided that the potential “window” available is suitable for the analyte in question Table 15.7 summarizes the accessible potential ranges for liquid mercury and for solid platinum electrodes The precious metals and various forms of carbon are the most common electrodes in use, although a great many materials, both metallic and semiconducting, find use as analytical electrode substrates Voltammetry is conducted using a microelectrode as the working electrode under conditions where polarization at the working electrode is enhanced This is in sharp contrast to both potentiometry and coulometry where polarization is absent or minimized by experimental conditions In voltammetry, very little analyte is used up in the measurement process unlike coulometry where complete consumption of the analyte is desired In polarography the DME is renewed regularly during the voltage sweep, with the consequence that the bulk concentration is restored at the electrode surface at the start of each new drop at some slightly higher (or lower) potential At a solid electrode, the electrolysis process initiated by a voltage sweep proceeds to deplete the bulk concentration of analyte at the surface without interruption This constitutes a major difference between classical polarography and sweep voltammetry at a solid electrode Normal pulse polarography is a useful technique at solid electrodes, because the bulk concentration is restored by convection at the surface during the off-pulse, provided that this is at least 10 longer than the electrolysis period It is possible to use a mercury electrode in a so-called “static” or “hanging” mode Such a stationary Hg droplet may be suspended from a micrometer syringe capillary Alternatively, commercially available electrodes have been developed Table 15.7 Potential Windows for Commonly Used Electrodes and Solvents Range (V vs SCEa) Mercury Aqueous M HClO4 M NaOH Nonaqueous 0.1 M TEAP/CH3CNb Platinum Aqueous M H2SO4 M NaOH Nonaqueous 0.1 M TBABF4/CH3CNc a ỵ0.05 to 21.0 20.02 to 22.5 ỵ0.6 to 22.8 ỵ1.2 to 20.2 ỵ0.6 to 20.8 ỵ2.5 to 22.5 The SCE may only be used in nonaqueous systems if traces of water are acceptable b TEAP, tetraethylammonium perchlorate c TBABF4 , tetrabutylammonium tetrafluoroborate 990 Chapter 15 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com that inject a Hg drop of varying size to the opening of a capillary These electrodes find application for stripping voltammetry (see subsequently) A powerful group of electrochemical methods uses reversal techniques Foremost among these is cyclic voltammetry, which is invaluable for diagnosing reaction mechanisms and for studying reactive species in unusual oxidation states In cyclic voltammetry, a microelectrode is used as the WE The potential is increased linearly and the current is measured The current increases as the potential of the electroactive material is reached The area of the working electrode and the rate at which the analyte can diffuse to the electrode surface limits the current A single voltage ramp is reversed at some time after the electroactive species reacts and the reverse sweep is able to detect any electroactive products generated by the forward sweep This is cyclic voltammetry A cyclic voltammogram of a typical reversible oxidation – reduction reaction is shown diagrammatically in Fig 15.34 Usually, an XY recorder or computer data system is used to track the voltage on the time axis so that the reverse current appears below the peak obtained in the forward sweep, but with opposite polarity The shapes of the waves and their responses at different scan rates are used for diagnostic purposes Substances are generally examined at concentrations around the millimolar level, and the electrode potentials at which the species undergo reduction and oxidation may be rapidly determined The height of the current peak of the first voltage sweep can be calculated from the Randles– Sevcik equation: ip ¼ 2:69  105 n2=3 AD1=2 Cv1=2 (15:41) where n is the number of electrons transferred; A, the electrode area (cm2); D, the diffusion coefficient of the electroactive species (cm2/s); C, the concentration of the electroactive species (mol/cm3); and v, the potential scan rate (V/s) The standard potential E is related to the anodic and cathodic peak potentials, Epa and Epc : E0 ẳ (Epc ỵ Epa ) (15:42) The peak separation is related to the number of electrons involved in the reaction: Epa À Epc ¼ 0:059 n (15:43) assuming that the IR drop (resistance) in the cell is not too large, which will increase the peak separation Cyclic voltammetry is not primarily a quantitative analytical technique The references at the end of this chapter provide additional guidance to its applications and interpretation Its real value lies in the ability to establish the nature of the electron transfer reactions—for example, fast and reversible at one extreme, slow and irreversible at the other—and to explore the subsequent reactivity of unstable products formed by the forward sweep Suffice it to say that such studies are valuable for learning the fate and degradation of such compounds as drugs, insecticides, herbicides, foodstuff contaminants or additives, and pollutants 15.3.5.1 Instrumentation for Voltammetry A cyclic voltammetry mode of operation is featured on many modern polarographs, or a suitable voltage ramp generator may be used in combination with a potentiostat The threeelectrode configuration is required Pretreatment of the working electrode is necessary for Electroanalytical Chemistry 991 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com Figure 15.34 (a) Excitation waveform and (b) current response for a reversible couple obtained in cyclic voltammetry reproducibility (unless a hanging MDE is used), and normally this entails polishing the electrode mechanically with successively finer grades of abrasives Some confusion results in choosing an initial voltage at which to start the voltammogram It is best to measure the open-circuit voltage between the working and reference electrodes with a high-impedance digital voltmeter This emf is known as the rest potential Erp Scans may then be made in the negative and/or positive directions starting at the rest potential set on the potentiostat By this means the potentials for reductions and oxidations, respectively, and the electrolyte limits (the “windows”) are established In other words, it is generally necessary to commence at a potential at which no faradaic reaction is possible, or else voltammograms will be distorted and irreproducible Scan rates typically vary from 10 mV/s to V/s if an XY recorder is used for current output The lower limit is due to thermal convection, which is always present in an 992 Chapter 15 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com electrolyte To record faster responses an oscilloscope or fast transient recorder is required The currents in successive scans differ from those recorded in the first scan For quantitation purposes the first scan should always be used 15.3.5.2 Stripping Voltammetry The technique of stripping voltammetry may be used in many areas requiring trace analyses to the ppb level It is especially useful for determining heavy metal contaminants in natural water samples or biochemical studies Either anodic or cathodic stripping is possible in principle, but analyses by anodic stripping voltammetry are more often used Stripping analysis is a two-step technique involving (1) the preconcentration of one or several analytes by reduction (in anodic stripping) or oxidation (in cathodic stripping) followed by (2) a rapid oxidation or reduction, respectively, to strip the products back into the electrolyte Analysis time is on the order of a few minutes The overall determination involves three phases: preconcentration quiescent (or rest) period stripping process, for example, by sweep voltammetry or differential pulse It is the preconcentration period that enhances the sensitivity of this technique In the preconcentration phase precise potential control permits the selection of species whose decomposition potentials are exceeded The products should form an insoluble solid deposit or an alloy with the substrate At Hg electrodes the electroreduced metal ions form an amalgam Usually the potential is set 100–200 mV in excess of the decomposition potential of the analyte of interest Moreover, electrolysis may be carried out at a sufficiently negative potential to reduce all of the metal ions possible below hydrogen ion reduction at Hg, for example Concurrent Hỵ ion reduction is not a problem, because the objective is to separate the reactants from the bulk electrolyte In fact, methods have been devised to determine the group I metals and NHỵ ion at Hg in neutral or alkaline solutions of the tetraalkylammonium salts Exhaustive electrolysis is not mandatory and 2–3% removal suffices Additionally, the processes of interest need not be 100% faradaically efficient, provided that the preconcentration stage is reproducible for calibration purposes, which is usually ensured by standard addition Typical solid substrate electrodes are wax-impregnated graphite, glassy (vitreous) carbon, platinum, and gold; however, mercury electrodes are more prevalent in the form of either a hanging mercury drop electrode (HMDE) or thin-film mercury electrode (TFME) TFMEs may be electro-deposited on glassy carbon electrodes from freshly prepared Hg(NO3)2 dissolved in an acetate buffer (pH 7) There is an art to obtaining good thin films, and usually some practice is necessary to get uniform, reproducible coverage on the carbon substrates It is recommended that literature procedures be adhered to carefully Some procedures recommend simultaneously pre-electrolyzing the analyte and a dilute mercury ion solution (1025 M) so that the amalgam is formed in a single step It is important to stir the solution or rotate the electrode during the preconcentration stage The purpose of this is to increase the analyte mass transport to the electrode by convective means, thereby enhancing preconcentration In general, in electroanalysis one seeks to obtain proper conditions for diffusion alone to permit mathematical expression of the process rate (the current) In this and controlled flow or rotation cases it is advantageous to purposely increase the quantity of material reaching the electrode surface Pre-electrolysis times are typically or longer In the rest period or quiescent stage, the stirrer is switched off for perhaps 30 s but the electrolysis potential is held This permits the concentration gradient of material within Electroanalytical Chemistry 993 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com the Hg to become more uniform The rest period is not obligatory for films produced on a solid substrate A variety of techniques have been proposed for the stripping stage Two important methods are discussed here In the first, in anodic stripping, for example, the potential is scanned at a constant rate to more positive values With this single sweep voltammetry the resolution of a TFME is better than that of an HMDE, because stripping of the former leads to a more complete depletion of the thin film As illustrated in Fig 15.35, it is possible to analyze for many metal species simultaneously The height of the stripping peak is taken to be directly proportional to concentration Linearity should be established in the working range with a calibration curve The second method of stripping involves the application of a differential pulse scan to the electrode As in polarographic methods, an increase in sensitivity is obtained when the differential pulse waveform is used Extremely high sensitivities in trace analyses require good analytical practice, especially in the preparation and choice of reagents, solvents, and labware Glass cells and volumetric glassware must be soaked 24 h in trace-metal purity M HNO3 Plastic electrochemical cells are recommended if loss of sample by adsorption to the vessel walls is a likely problem The inert gas (N2 or Ar) used to remove dissolved oxygen should be purified so as not to introduce additional contaminants Solid catalysts and drying agents are recommended for oxygen and water removal A presaturator is recommended to reduce electrolyte losses by volatilization, especially when low vapor pressure organic solvents are used A major source of contamination is the supporting electrode itself This may be purified in part by recrystallization, but the use of sustained preelectrolysis at a mercury pool electrode may ultimately be required Dilute standards and samples ought to be prepared daily because of the risk of chemical (e.g., hydrolysis) or physical (e.g., adsorption) losses 15.3.5.3 Applications of Anodic Stripping Voltammetry Anodic stripping voltammetry is readily applicable for those metals that form an amalgam with mercury, for example, Ag, As, Au, Bi, Cd, Cu, Ga, In, Mn, Pb, Sb, Sn, Tl, and Zn One important cause of interferences is intermetallic compound formation of insoluble alloys Figure 15.35 Stripping voltammogram of metal-ion species at a TFME 994 Chapter 15 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com between the metals within the amalgam For example, In/Au and Cu/Ni can form in the Hg drop and then not respond to the stripping stage, so their measured concentrations will be too low It is imperative to carefully select the electrolyte such that possibly interfering compounds are complexed and electroinactive This is also a means of improving resolution when there are two overlapping peaks Some elements can be analyzed in aqueous electrolytes only with difficulty (groups I and II), but, fortunately, their analysis by flame or atomic absorption methods are sensitive and easy to carry out Anions of carbon compounds can also be stripped by either (1) anodic preconcentration as sparingly soluble Hg salts, (2) adsorption and decomposition, or (3) indirect methods, such as displacement of a metal complex Examples of complexes used include thiourea, succinate, and dithizone Anions may be determined as mercurous or silver(I) salts if these are sparingly soluble Based on solubility determinations, it is possible to estimate some theoretical values for the minimum determinable molar concentrations of anions at mercury: Cl2,  29 22 1026; Br2,  1026; I2,  1028; S22,  1028; CrO22 ,  10 ; WO4 ,  27 26 26 22 22 10 ; MoO4 ,  10 ; and C2O4 ,  10 With the exception of Cl , the mercurous salts are less soluble than the silver salts In conclusion, stripping voltammetry is an inexpensive, highly sensitive analytical tool applicable to multicomponent systems; in fact, it is not recommended for metal-ion samples whose concentrations are greater than ppm Careful selection of operating conditions and especially the electrolyte buffer is necessary The sensitivity of stripping voltammetry is less for nonmetallic and anionic species than for metals More recently, flow-through systems have been devised for continuous monitoring purposes Stripping voltammetry has been applied to numerous trace metal analyses and environmental studies, for example, to determine impurities in oceans, rivers, lakes, and effluents; to analyze body fluids, foodstuffs, and soil samples; and to characterize airborne particulates and industrial chemicals 15.4 LC DETECTORS Electrochemical detectors for the determination of trace amounts of ionic and molecular components in liquid chromatographic effluents are sensitive, selective, and inexpensive For many separations they provide the best means of detection, outperforming, for example, UV/VIS and fluorescence spectroscopy and refractive index methods Two major categories of electrochemical detectors will be described subsequently: voltammetric and conductometric An efficient redox center in the ion or molecule to be detected is necessary for voltammetric methods, but this is not mandatory in the species actually undergoing separation, because many organic compounds can be derivatized with an electroactive constituent before or after column separation Alternatively, certain classes of organic compounds can be photolytically decomposed by a post-column online UV lamp into electroactive products, which can be detected For example, organic nitrocompounds can be photolyzed to nitrite and nitrate anions, which undergo electrolysis at suitable working electrode potentials In the case of conductivity detectors the method does not require species that are redox active, or chromophoric; rather, changes in the resistivity of the electrolyte eluent are monitored Examples of species that may be detected are simple ions such 32 32 ỵ as halides, SO22 , PO4 , NO , metal cations, NH4 , as well as organic ionic moieties 15.4.1 Voltammetric Detection Various waveforms and signal presentation techniques are available, such as steady-state voltammetry with amperometric or coulometric (integrated current) outputs and sweep Electroanalytical Chemistry 995 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com Figure 15.36 Schematic of an LC system with electrochemical detection procedures such as differential pulse In the simplest mode a sufficient steady-state potential (constant E) is applied to a working electrode to reduce or oxidize the component(s) of interest The output signal is a transient current peak collected on a Y-time recorder (Fig 15.36) or for modern chromatographs, collected and processed by a computer data system To obtain reasonable chromatographic resolution, the electrochemical cell, positioned at the outlet of the column, must be carefully designed The volume of the detector cell must be small to permit maximum band resolution As the reduction or oxidation is performed under hydrodynamic flow conditions, the kinetic response of the electrode/redox species should be as rapid as possible Glassy (vitreous) carbon is a popular electrode material, and it is possible to chemically alter the electrode surface to improve the response performance However, other substrates are available, such as Pt, Hg, and Au; these prove useful in cases where they are not passivated by the adsorption of organic reaction products Platinum and gold electrodes are particularly prone to blockage from the oxidation products of amino acids, carbohydrates, and polyalcohols, for example Typically, a three-electrode thin-layer cell is used, with the reference and auxiliary electrodes downstream from the working electrode Dissolved oxygen may interfere in some assays and may therefore have to be removed by prior degassing 15.4.2 Conductometric Detection Many important inorganic and organic ions at the trace level are not easily detected by reduction–oxidation or spectroscopic methods This has led to the development of sophisticated separation and conductometric detectors capable of ultrasensitive analyses It is difficult to measure small changes in conductivity due to a trace analyte ion when the background level is orders of magnitude larger; consequently procedures have been developed that are capable of suppressing high ionic strength backgrounds, allowing the net signal to be measured more easily Combined with miniature conductance detectors, these procedures have led to a technique known as ion chromatography with eluent suppression (cf 13.2.1; Fig 13.24) Figure 15.37 shows the setup for a typical application: the analysis of the constituents of acid rain samples 996 Chapter 15 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com Figure 15.37 Schematic of the eluent suppression approach for ion chromatography that permits conductometric detection of analytes In the first ion exchange column, whose resin exists in its carbonate form and SO22 (R2ỵCO22 ), the sample species, such as Cl , are selectively separated and eluted as sodium salts in a background of sodium carbonate Detection of the chloride and sulfate by conductivity at this stage would not be possible because of the large excess of conducting ions Naỵ and CO22 In the second column, the cations undergo ion exchange to produce highly conducting acids (such as the strong acids HCl and H2SO4 , which are fully ionized) with a background of slightly ionized carbonic acid: H2 CO3 ! 2Hỵ ỵ CO2 Now the highly ionized chloride and sulfate can be measured because the carbonic acid has a low conductance A major problem is that the capacity of the second column needs frequent regeneration if large amounts of Na2CO3 eluent are used Various approaches have been tested to improve the practical operation of this form of separation The first objective is to shorten the time needed to resolve the dilute sample constituents This is done with a low-capacity separator, which decreases the resolution time and thereby the exchange consumption of the suppressor column The next objective is to provide continuous regeneration of the suppressor column by automation or by a novel ion exchange membrane methodology that continuously replenishes the column For cation analyses, lightly sulfonated beads of Electroanalytical Chemistry 997 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com cross-linked polystyrene have been employed; for anion analysis composite beads are formed by electrostatically coating these beads with anion-exchanging spheres Liquid chromatographic electrochemical detection has been widely used for metabolite studies in complex matrices and has general applicability in many fields, for example, the pharmaceutical industry, forensic science, medicine, the explosives industry, and agriculture BIBLIOGRAPHY Adams, R Electrochemistry at Solid Electrodes; Marcel Dekker, Inc.: New York, 1969 Aikens, D.; Bailey, R.; Moore, J.; Giachino, G.; Torukins, R Principles and Techniques for an Integrated Chemistry Laboratory; Waveland Press, Inc.: Prospect Heights, IL, 1985 ASTM Annual Book of ASTM Standards, Water and Environmental Technology; ASTM: West Conshohocken, PA, 2000; Vols 11.01 and 11.02 Bard, A.J., Ed Electroanalytical Chemistry; Marcel Dekker, Inc.: New York, 1993 Bard, A.J.; Faulkner, L.R Electrochemical Methods—Fundamentals and Applications; John Wiley and Sons, Inc.: New York, 1980 Bard, A.J.; Parsons, R.; Jordan, J., Eds Standard Potentials in Aqueous Solution; Marcel Dekker, Inc.: New York, 1985 Baizer, M.M.; Lund, H Organic Electrochemistry; Marcel Dekker, Inc.: New York, 1983 Bates, R.G Determination of pH; John Wiley and Sons, Inc.: New York, 1973 Berezanski, P In Handbook of Instrumental Techniques for Analytical Chemistry; Settle, F.A., Ed.; Prentice-Hall PTR: NJ, 1997 Bockris, J.O Modern Electrochemistry; Plenum Press: New York, 1970; Vols and Bond, A.M Modern Polarographic Methods in Analytical Chemistry; Marcel Dekker, Inc.: New York, 1980 Dean, J.A Analytical Chemistry Handbook; McGraw-Hill, Inc.: New York, 1995 Diefenderfer, A.J.; Holton, B.E Principles of Electronic Instrumentation, 3rd ed.; Saunders College Publishing: Philadelphia, PA, 1994 Dryhurst, G Electrochemistry of Biological Molecules; Academic Press: New York, 1977 Enke, C.G The Art and Science of Chemical Analysis; John Wiley and Sons, Inc.: New York, 2000 Evans, A Potentiometry and Ion Selective Electrodes; John Wiley and Sons, Inc.: New York, 1987 Ewing, G.W., Ed Analytical Instrumentation Handbook, 2nd ed.; Marcel Dekker, Inc.: New York, 1997 Fritz, J.S Acid–Base Titrations in Non-Aqueous Solvents; G Frederick Smith Chemical Company: Columbus, OH, 1952 (This publication is still available from GFS Chemicals upon request.) Fry, A.J Synthetic Organic Electrochemistry; Harper and Row: New York, 1972 Greenberg, A.; Clesceri, L.; Eaton, A., Eds Standard Methods for the Examination of Water and Wastewater, 18th ed.; American Public Health Association: Washington, DC, 1992 Harris, D.C Quantitative Chemical Analysis, 5th ed.; W.H Freeman and Co.: New York, 1999 Hart, J.P Electroanalysis of Biologically Important Compounds; Ellis Horwood: New York, 1990 Huber, W Titrations in Nonaqueous Solvents; Academic Press: New York, 1967 Ives, D.J.G.; Janz, G.J., Eds Reference Electrodes—Theory and Practice; Academic Press: New York, 1961 Kalvoda, R Operational Amplifiers in Chemical Instrumentation; Ellis Horwood: Chichester, 1975 Kissinger, P.T.; Heineman, W.R Laboratory Techniques in Electroanalytical Chemistry, 2nd Ed.; Marcel Dekker, Inc.: New York, 1996 Koryta, J.; Stulik, K Ion-Selective Electrodes, 2nd Ed.; Cambridge University Press: Cambridge, 1983 Kucharsky, J.; Safarik, L Titrations in Nonaqueous Solvents; Elsevier: Amsterdam, 1965 Malmstadt, H.V.; Enke, C.G.; Crouch, S.R Microcomputers and Electronic Instrumentation: Making the Right Connections; American Chemical Society: Washington, DC, 1994 Meites, L.; Zuman, P.; Rupp, E.B., Eds CRC Handbook Series in Organic Electrochemistry; CRC Press, Inc.: Boca Raton, FL, 1982; Vols – Michael, A.; Wightman, R.M In Laboratory Techniques in Electroanalytical Chemistry, 2nd Ed.; Kissinger, P., Heineman, W., Eds.; Marcel Dekker, Inc.: New York, 1996 998 Chapter 15 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com Settle, F.A., Ed Handbook of Instrumental Techniques for Analytical Chemistry; Prentice-Hall PTR: NJ, 1997 Streuli, C.A Titrimetry: acid–base titrations in non-aqueous solvents In Treatise on Analytical Chemistry, Part I; Kolthoff, I.M., Elving, P.J., Eds.; Wiley-Interscience: New York, 1975; Vol II Vydra, F.; Stulik, K.; Julakova, E Electrochemical Stripping Analysis; Ellis Horwood, 1976 Wang, J Stripping Analysis; VCH Publishers, Inc.: Deerfield Beach, FL, 1985 Warner, M Anal Chem 1994, 66, 601A SUGGESTED EXPERIMENTS 15.1 Run a DC polarogram with a known volume of M KNO3 electrolyte that has not been deaerated Observe and identify the two reduction waves of dissolved oxygen 15.2 Purge the electrolyte in problem 15.1 with purified nitrogen gas for 20 or so Add sufficient lead nitrate solution to make the electrolyte 1.0 1024 M in Pb2ỵ ion Record the DC polarogram and measure the limiting current and the half-wave potential (E1=2 ) for the lead ion reduction wave 15.3 Measure the mercury flow rate with the column height set as for problem 15.2 Collect about 20 droplets under the electrolyte (why?) and determine the drop lifetime with a stopwatch Use the Ilkovic equation to calculate the diffusion equation for Pb2ỵ ion and compare your derived value with a literature value 15.4 Record a normal pulse polarogram of (a) 1.00  1025 M Zn2ỵ, (b) Cd ion and, (c) Cu ion in a degassed M KNO3 electrolyte Construct a graph of potential (E ) vs log[(iL i)/i] Determine the value of n from the slope and comment on the reversibility of this reaction How does the electrolyte temperature affect the slope of this curve? 15.5 Record a differential pulse polarogram of a background electrolyte without and with added tapwater Which metal species are present? Use a standard additions calibration to quantify one of the metal impurities Compare Epeak with E1/2i 15.6 Use a F2 ion selective electrode per manufacturer’s instructions to find the concentration of F2 ion in your drinking water and in your brand of (a) toothpaste and (b) mouthwash Use a standard additions calibration procedure Some sample preparation may be needed for the toothpaste Several extraction procedures are available in the literature 15.7 Collect some samples for pH analyses, for example, rainfall, snow, lake water, pool water, soft drinks, fruit juice Standardize the pH meter with a standard buffer of pH close to the particular sample of interest Design some experiments to change the temperature of these samples What is the influence of temperature on pH measurements? 15.8 (a) Use a glass pH electrode combination and titrate unknown strength weak acid(s) with 0.2000 N NaOH Stir the electrolyte continuously (b) Calculate the strength(s) of the acid(s) (i.e., the pKa) and explain the shape of the curve 15.9 Copper sulfate may be analyzed by gravimetry by exhaustive electrolysis at a weighed platinum electrode Add mL H2SO4 , mL HNO3 , and g urea to 25 mL of the copper solution Suggest a stoichiometry for your copper sulfate based on a variable water of hydration, for example, CuSO4 2H2O 15.10 Tap waters, as well as a variety of biological fluids and natural waters, can be analyzed for heavy metals by anodic stripping voltammetry at a hanging mercury Electroanalytical Chemistry 999 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com drop electrode Clean glassware and sample bottles by leaching for about h in N HNO3 First, analyze for Cd2ỵ and Pb2ỵ ions at acidic pH (2–4) Deposition is made with gentle stirring for exactly After a 15 s quiescent period, scan from 20.7 V SCE in the positive direction Zn2ỵ and Cu2ỵ ions can be analyzed by raising the pH to 8–9 with M NH4Cl/1 M NH4OH buffer Deposit at 21.2 V SCE and strip as before It will be necessary to adjust the deposition time and/or current output sensitivity to achieve best results 15.11 Cyclic voltammetry can be used to study the reversible reduction – oxidation couple of [Fe(CN)6]32 þ e2 $ [Fe(CN)6]42 You will need solutions of M KNO3 ,  1023 M K3[Fe(CN)6] in M KNO3 , and  1023 M VOSO4 in M KNO3 A CV equipped with a three-electrode cell (Pt microelectrode as the working electrode, Pt foil electrode as the auxiliary electrode and a calomel reference electrode) and an X– Y recorder is required The potassium nitrate solution serves as the baseline All solutions must be purged with nitrogen to eliminate oxygen before scanning the potential A sweep from 0.80 to 20.12 V and back vs SCE is suitable for the Fe system The range may need to be changed for the vanadium system Can you evaluate n for each system? Is the value what you expect for each reaction? Evaluate E for each system Do these values agree with those in the literature? (Experiment courtesy of Professor R.A Bailey, Department of Chemistry, Rensselaer Polytechnic Institute.) PROBLEMS 15.1 15.2 Give an example of a half-cell The absolute potential of a half-cell cannot be measured directly How can the potential be measured? Using the Nernst equation, complete the following table: pH of solution Concentration of Hỵ E of hydrogen half-cell 1.2 1025 10210 12 14 15.3 15.4 15.5 Describe and illustrate an SCE Write the half-cell reaction How is pH measured with a glass electrode? Why does a glass electrode give pH readings lower than the actual pH in strongly basic solutions? What other errors can occur in pH measurement with a glass electrode? Calculate the theoretical potential of the following cell Is the cell as written galvanic or electrolytic? PtjCr3ỵ (3:00102 M),Cr2ỵ (4:00105 M)kSn2ỵ (2:00102 M), Sn4ỵ (2:00104 M)jPt 1000 Chapter 15 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com 15.6 15.7 15.8 15.9 15.10 15.11 15.12 15.13 15.14 15.15 15.16 15.17 15.18 15.19 15.20 15.21 15.22 A salt of two monovalent ions M and A is sparingly soluble The E for the metal is ỵ0.621 V vs SHE The observed emf of a saturated solution of the salt is 0.149 V vs SHE What is the solubility product Ksp of the salt? (Neglect the junction potential.) Describe the process of electrodeposition and how it is used for electrogravimetric analysis Can two metals such as iron and nickel be separated completely by electrogravimetry? Explain your answer and state any assumptions you make State Faradays Law A solution of Fe3ỵ has a volume of 1.00 L, and 48,246 C are required to reduce the Fe3ỵ to Fe2ỵ What was the original molar concentration of iron? What is the molar concentration of Fe2ỵ after the passage of 12,061 C? What are the three major forms of polarography? State the reasons why pulse polarographic methods are more sensitive than classical DC polarography What are the advantages of mercury electrodes for electrochemical measurements? What are the advantages of the dropping mercury electrode vs a Pt microelectrode for polarography? What are the disadvantages of the DME? Describe the method of anodic stripping voltammetry What analytes can it be used to determine? Why is stripping voltammetry more sensitive than other voltammetric methods? Describe the principle of an ISE Why is the term ion-specific electrode not used? How can conductometric measurements be used in analytical chemistry? Give two examples Briefly outline two types of electrochemical detectors used for chromatography The fluoride ISE is used routinely for measuring fluoridated water and fluoride ion in dental products such as mouthwash A 50 mL aliquot of water containing sodium fluoride is analyzed using a fluoride ion electrode and the method of standard additions The pH and ionic strength are adjusted so that all fluoride ion is present as free F2 ion The potential of the ISE/reference electrode combination in a 50 mL aliquot of the water was 20.1805 V Addition of 0.5 mL of a 100 mg/L F2 ion standard solution to the beaker changed the potential to 20.3490 V Calculate the concentration of (1) fluoride ion and (2) sodium fluoride in the water sample Copper is deposited as the element on a weighed Pt cathode from a solution of copper sulfate in an electrolytic cell If a constant current of 0.600 A is used, how much Cu can be deposited in 10.0 min? (Assume no other reductions occur and that the reaction at the anode is the electrolysis of water to produce oxygen.) Why does coulometry not require external calibration standard solutions? Explain why a silver electrode can be an indicator electrode for chloride ion What are the three processes by which an analyte in solution is transported to an electrode surface? What single transport process is desired in polarography? Explain how the other transport processes are minimized in polarography Would you expect the half-wave potential for the reduction of copper ion to copper metal to be different at a Hg electrode from that at a platinum electrode? Explain your answer Sketch a schematic cyclic voltammogram for a nonreversible reduction reaction (See Fig 15.34 for a CV of a reversible reaction.) Electroanalytical Chemistry 1001 Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com APPENDIX 15.1 SELECTED STANDARD REDUCTION POTENTIALS AT 2588 C Half-reaction E (V) F2(g) ỵ 2e2 ! 2F2 Co3ỵ ỵ e2 ! Co2ỵ H2O2 ỵ 2Hỵ ỵ 2e2 ! 2H2O Ce4ỵ ỵ e2 ! Ce3ỵ PbO2(s) ỵ 4Hỵ ỵ SO22 ! PbSO4(s) ỵ 2H2O ỵ 2e ỵ MnO4 ỵ 8H ỵ 5e ! Mn2ỵ þ 4H2O Au3þ þ 3e2 À! Au(s) Cl2(g) þ 2e2 ! 2Cl2 ỵ ! 2Cr3ỵ ỵ 7H2O Cr2O22 þ 14H þ 6e þ MnO2(s) þ 4H þ 2e ! Mn2ỵ ỵ 2H2O O2(g) ỵ 4Hỵ ỵ 4e2 ! 2H2O Br2(l) ỵ 2e2 ! 2Br2 ỵ NO2 ! NO(g) ỵ 2H2O ỵ 4H ỵ 3e 2ỵ 2Hg ỵ 2e ! Hg2ỵ 2 ! 2Hg(l) Hg2ỵ ỵ 2e Agỵ ỵ e2 ! Ag(s) Fe3ỵ þ e2 À! Fe2þ O2(g) þ 2Hþ þ 2e2 À! H2O2 MnO2 ! MnO2(s) ỵ 4OH2 ỵ 2H2O þ 3e 2 I2(s) þ 2e À! 2I O2(g) þ 2H2O þ 4e2 À! 4OH2 Cu2þ þ 2e2 À! Cu(s) AgCl(s) ỵ e2 ! Ag(s) ỵ Cl2 Cu2ỵ ỵ e2 ! Cuỵ Sn4ỵ ỵ 2e2 ! Sn2ỵ 2Hỵ ỵ 2e2 ! H2(g) Pb2ỵ ỵ 2e2 ! Pb(s) Sn2ỵ ỵ 2e2 ! Sn(s) Ni2ỵ ỵ 2e2 ! Ni(s) Co2ỵ ỵ 2e2 ! Co(s) Cd2ỵ ỵ 2e2 ! Cd(s) Fe2ỵ ỵ 2e2 ! Fe(s) Cr3ỵ ỵ 3e2 ! Cr(s) Zn2ỵ ỵ 2e2 ! Zn(s) 2H2O ỵ 2e2 ! H2(g) ỵ 2OH2 Mn2ỵ ỵ 2e2 ! Mn(s) Al3ỵ ỵ 3e2 ! Al(s) Naỵ ỵ e2 ! Na(s) Ca2ỵ ỵ 2e2 ! Ca(s) Ba2ỵ ỵ 2e2 ! Ba(s) Kỵ ỵ e2 ! K(s) Liỵ ỵ e2 ! Li(s) ỵ2.87 ỵ1.82 ỵ1.77 ỵ1.74 ỵ1.70 þ1.51 þ1.50 þ1.36 þ1.33 þ1.23 þ1.23 þ1.07 þ0.96 þ0.92 þ0.85 þ0.80 þ0.77 þ0.68 þ0.59 þ0.53 þ0.40 þ0.34 þ0.22 þ0.15 þ0.13 0.00 20.13 20.14 20.25 20.28 20.40 20.44 20.74 20.76 20.83 21.18 21.66 22.71 22.87 22.90 22.93 23.05 Note: All species in aqueous solution (aq) unless otherwise indicated All dissolved species ¼ M; all gases ¼ atm Simpo PDF Merge and Split Unregistered Version - http://www.simpopdf.com ... are now used not only for trace metal ion analyses, but also for the analysis of organic compounds, for continuous process analysis, and for studying the chemical reactions within a single living... orders of magnitude, thus providing a means of analysis at the ppb level In practice, electrochemistry not only provides a means of elemental and molecular analysis, but also can be used to acquire... species One disadvantage of exhaustive electrolysis is the time required for analysis, and faster methods of electrochemical analysis are described In summary, important practical considerations in