Electrochemistry

Basic

1. Basic of electrochemistry BY: Dr.G.C.Wadhawa Karmaveer Bhaurao Patil College, Vashi, Navimumbai India

2. An electrochemical system is comprised of a vessel containing an electrolyte into which two electrodes are dipped. The electrodes are connected by first-class conductors either with a source of electric current. During electrolysis, chemical reactions of ions occur at the electrodes due to the passage of electric current. In an electrolytic cell, a flow of current produces a chemical reaction. In a galvanic cell, electric current is generated in consequence of chemical reactions proceeding at the electrodes. 1 Electrochemistry

3. Electrochemical cells An electrochemical cell is a device capable of either generating electrical energy from chemical reactions or facilitating chemical reactions through the introduction of electrical energy. An electrochemical cell consists of two half-cells. Each half-cell consists of an electrode and an electrolyte. The chemical reactions in the cell may involve the electrolyte, the electrodes, or an external substance. In a full electrochemical cell, species from one half- cell lose electrons (oxidation) to their electrode while species from the other half-cell gain electrons (reduction) from their electrode.

4. A salt bridge is often employed to provide ionic contact be tween two half-cells with different electrolytes and prevent the solutions from mixing and causing unwanted side reactions. e.g., filter paper soaked in KNO 3 or some other electrolyte . An alternative to a salt bridge is to allow direct contact (and mixing) between the two half-cells. Also it allows the flow of negative or positive ions to maintain a steady-state charge distribution between the oxidation and reduction vessels.

5. Electrochemical cells are classified into two types: Electrolytic cells are those in which electrical energy from an external source causes non spontaneous chemical reactions to occur. (  G > 0) Voltaic cells are those in which spontaneous chemical reactions produce electricity and supply it to an external circuit. (  G < 0) Both types of cells contain two electrodes connected to an external circuit that provides an electrical connection between systems. When circuit is closed, electrons flow from the anode to the cathode; electrodes are connected by an electrolyte, which is an ionic substance or solution that allows ions to transfer between the electrodes, thereby maintaining the system’s electrical

6. Electrolytic cells In electrolytic cell, nonspontaneous chemical reactions are forced to occur by the input of electrical energy . This process is called electrolysis . ele c t r o d e s , an curren t . Anode 3. The Mig r ations o f i o ns in the s o l u tion is a n electr i c a l 2. A n ele c t r o l y t ic c e l l con s i sts o f t w o electrolyte and external power source . Cat h ode

7. … Electrodes are surfaces on which oxidation or reduction half-reactions occur. They may or may not participate in the reactions. Those that do not react are called inert electrodes . Inert electrodes are often used so that they do not affect product of electrolysis. The external power source acts as an “ electron pump ”; the electric energy is used to do work on the electrons to cause an electron transfer Electrons are pulled from the anode and pushed to the cathode by the battery or power supply.

8. … Regardless of the kind of cell, electrolytic or voltaic, the electrodes are identified as follows. The cathode is defined as the electrode at which reduction occurs as electrons are gained by some species. The anode is the electrode at which oxidation occurs as electrons are lost by some species. Cu(s) + Zn +2 ↔ Cu +2 + Zn(s) Cu(s) ↔ Cu +2 + 2e - (oxidation) Zn +2 + 2e - ↔ Zn(s) (reduction) Each of these can be either the positive or the negative electrode. In electrolytic cell, anode is positive while cathode is negetive

9. + - battery Na (l) Ano d e Na + Cl - Na + Cathode Na + + e -  Na 2Cl -  Cl 2 + 2 e - Cl 2 (g) escapes Electrolysis of molten NaCl Observe the reactions at the electrodes NaCl (l) ( - ) Cl - (+) Cl - Na + Overall cell reaction: 2Na + + 2Cl -  2Na + Cl 2 Non-spontaneous reaction!

10. + - battery e - e - NaCl (l) (-) anode Na + Cl - Cl - (+) Cl - Na + Na + cathode Na + + e -  Na Cations migrate toward (-) electrode Anions migrate toward (+) electrode At the microscopic level 2Cl -  Cl 2 + 2 e -

11. Voltaic cells Voltaic or galvanic cells are electrochemical cells in which spontaneous redox reaction produce electrical energy. In all voltaic cells , electrons flow spontaneously from the negative electrode (anode) to the positive electrode (cathode). Composed of two separated half-cells ; which each consist of a metal rod or strip immersed in a solution of its own ions or an inert electrolyte. The electrodes are connected by wire and the circuit the two cells is completed by salt bridge . This can be any medium through which ions can slowly pass. The electrons flow from the anode to the cathode through an electrical circuit rather than passing directly from on substance to another .

12. Cu 1.0 M CuSO 4 Zn 1.0 M ZnSO 4 Zn ele c tr o de erodes or d i s s o l ves cathode half-cell Cu +2 + 2e -  Cu Cu deposits on electrode anode half-cell Zn  Zn +2 + 2e - + What about the sign of the electrodes? What happened at each ele c tr o de? e - e - K + Salt bridge Cl -

13. Co n t ’ d… Salt bridge serves the following • It allows electrical contact between the two solutions. • It prevents mixing of the electrode solution. • It maintains electrical neutrality in each half cell as ions flow into and out of the salt bridge.

14. ELECTROLYTIC CELLS GALVANIC CELLS It requires EMF It produces EMF Electric energy is converted into chemical energy Chemical energy is converted into electrical energy Anode is +ve and cathode ‘-‘ ve Anode is –ve cathode is ‘+’ ve Oxidation takes place at anode and reduction at cathode Oxidation takes place at anode and reduction at cathode. Discharges of ion occur at both electrodes. Discharges of ions occur only at cathode Non – spontaneous reaction occurs. Spontaneous reaction occurs These are irreversible. These may be reversible Flow of electrons is from anode to cathode Flow of electrons is from anode to cathode Electrons leave the cell at anode and enter the cell at cathode. Electrons leave the cell at anode and enter the cell at cathode

15. Electrolysis During electrolysis, the anode carries a positive charge and the cathode a negative charge due to the influence of an external source. Anions travel through the electrolyte toward the anode, cations travel toward the cathode. Electrochemical reactions depend primarily on the nature of the electrolyte , on temperature , the current density, and on the material of the electrodes

16. In terms of the material used, we classify electrodes as inert and active. Inert elect r od e s do not parti c ipate in electrochemical reactions proceeding at them. Example : platinum electrode Active electrodes take part in electrochemical reactions. may either dissolve or their material may react with the corresponding ions. electrode   Examples of the  2 produ  cts discharging at an inert 2 Zn  e  Zn 2Cl  Cl  2e

17. Reversible electrodes Reversible electrode is the electrode at which the small change in potential can reverse the direction of the net current flow at this electrode. A metallic electrode that will dissolve when a current is passed from it into a solution and that will have plated on it metal from the solution when the current is passed in the reverse direction When current is led through a Daniell cell(Zn/ZnSO 4 (aq)//CuSO 4 (aq)// Cu) from the zinc to the copper: more ions of zinc are formed from the solid electrode. if another current be sent the reverse way through the cell, it will carry some of the zinc ions with it, and deposit them in the metallic state on the electrode. This, then, is a case of a reversible electrode, and in particular it is reversible with respect to a cation, that is Zn +2

18. Reversible electrodes may be divided into three groups. Electrodes of the first kind: A) These include cationic electrodes (metal , amalgam, gas electrodes, the hydrogen electrode ) equilibrium is established between atoms or molecules of the substance and the corresponding cations in solution and anionic electrodes. B) Electrodes of the second kind: These electrodes consist of three phases. The metal is covered by a layer self metal of its sparingly s o lubl e sal t and i s immers e d i n solutions a containing the anions of this salt.

19. Because of the two interfaces, equilibrium is established between: the metal atoms and the anions in solution. the metal and its cation in the sparingly soluble salt the anion in the solid phase of the sparingly soluble salt and the anion in solution can make eqlibrium with each other c) Oxidation-reduction electrodes : An inert metal (usually Pt, Au, or Hg) is immersed in a solution of two soluble oxidation forms of a substance. Equilibrium is established through electrons. This type of electrode differs from electrodes of the first kind only in that both oxidation states can b

20. Cell Diagrams and IUPAC Conventions In a cell diagram: The anode is written on the left, the cathode on the right. Their charges may, but do not have to, be marked. The phase interface is marked by a single vertical bar” | ”. the salt bridge, which has two phase boundaries, is shown by a double vertical line “ || ” W ith t h e ele c tr o lyte w e n o te its m o l al i t y o r a c tivity in brackets. For gas electrodes, the gas pressure or fugacity is given in brackets. A dash ( — ) is wr i tten betwe e n a gas ele c tr o de and its carrier.

21. 1 Here’s a cell diagram for Zn/Cu cell Zn( s )| Zn 2+ ( aq , 1 M ) || Cu 2+ ( aq , 1 M ) | Cu( s ) Anode Salt Bridge Cathode 2 Cell components for the Zn-Cu cell are: A metallic Cu strip immersed in 1.0 M copper (II) sulfate. A metallic Zn strip immersed in 1.0 M zinc (II) sulfate. A w ire and a salt bridge to complet e circuit

22. Liquid junction potential ( E j ) Liq u id j u ncti o n p o ten t i a l occurs when two s o l u t i o ns of different concentrations are in contact with each other. The m o r e conc e nt r a ted soluti o n will ha v e a tende n cy to diffuse into the comparatively less concentrated one. The rate o f d i f f u si o n o f eac h i o n will b e r o u ghly proportional to its speed in an electric field. If the anions diffuse more rapidly than the cations, they will diffuse ahead into the dilute solution, leaving the latter negatively charged and the concentrated solution positively charged . This will result in an electrical double layer of positive and negative charges at the junction of the two solutions .

23. Commercial Galvanic Cells Galvanic cells can be self-contained and portable and can be used as batteries and fuel cells. A battery (storage cell) is a galvanic cell (or a series of galvanic cells) that contains all the reactants needed to produce electricity. A fuel cell is a galvanic cell that requires a constant external supply of one or more reactants in order to generate electricity. Batteries Two basic kinds of batteries (i) Disposable, or primary , batteries in which the electrode reactions are effectively irreversible and which cannot be recharged.

24. ( Rechargeable , or secondary , batteries, which form an insoluble product that adheres to the electrodes; can be recharged by applying an electrical potential in the reverse direction, which temporarily converts a rechargeable battery from a galvanic cell to an electrolytic cell. Major difference between batteries and galvanic cells is that commercial batteries use solids or pastes rather than solutions as reactants to maximize the electrical output per unit mass. The Dry Cell One example of a dry cell is flashlight and radio batteries. The cell’s container is made of zinc which acts as an electrode .

25. The space between the electrodes is filled with a mixture of: ammonium chloride, NH 4 Cl manganese (IV) oxide , MnO 2 The electrodes in a secondary cell can be regenerated by the addition of electricity. The s e c e l l s ca n b e swi t ch e d fro m vo l tai c to electrolytic cells .

26. he electrolyte is dilute sulfuric acid. One example of a secondary voltaic cell is the lead storage or car battery. The Lead Storage Battery the lead storage battery the electrodes are two sets of lead alloy grids (plates). Hol e s i n on e of th e grid s are f illed wit h lea d (I V ) oxide, PbO 2 . The other holes are filled with spongy lead.

27. Lithium – iodine ba ttery Water-free battery Consists of two cells separated by a metallic nickel mesh that collects charge from the anodes The anode is lithium metal, and the cathode is a solid complex of  2 Electrolyte is a layer of solid Li  that allows Li + ions to diffuse from the cathode to the anode Highly reliable and long-lived Used in cardiac pacemakers, medical implants, smoke alarms, and in computers Disposable

28. The Nickel-Cadmium (Nicad) Cell Nicad batteries are the rechargeable cells used in calculators, cameras, watches, etc. a water-based cell with a cadmium anode and a highly oxidized nickel cathode. This design maximizes the surface area of the electrodes and minimizes the distance between them, which gives the battery both a high discharge current and a high capacity. Lightweight, rechargeable, and high capacity but and do not store well ; tend to lose capacity quickly also p re s e n t s di s po sal p r o blem s becaus e of thetoxicity of cadmium .

29. Fuel Cells A galvanic cell that requires an external supply of reactants because the products of the reaction are continuously removed. Does not store electrical energy but allows electrical energy to be extracted directly from a chemical reaction . Have reliability problems and are costly 58

30. Used in space vehicles. Hydrogen is oxidized at the anode. Oxygen is reduced at the cathode Types of fuel cells PEMFC, Proton Ex c hange Membrane Fuel Cell. DMFC , D ire c t Methano l Fuel Cell. PAFC, Ph o sph o ric Acid F u el Cell. AFC, Alkaline Fuel Cell. MCFC, Molten Carbonate Fuel Cell . SOFC, Solid Oxide Fuel Cell . 59

31. What is electrochemistry? Electrochemistry is defined as the branch of chemistry that examines the phenomena resulting from combined chemical and electrical effec ts Types of processes This field covers: Electrolytic processes : Reactions in which chemical changes occur on the passage of an electrical current Galvanic or Voltaic processes : Chemical reactions that result in the production of electrical energy

32. What is Electrode Potential? When a metal is placed in a solution of its ions, the metal acquires either a positive or negative charge with respect to the solution. On account of this, a definite potential difference is developed between the metal and the solution. This potential difference is called electrode potential . For example, when a plate of zinc is placed in a solution having Zn 2+  ions, it becomes negatively charged with respect to solution and thus a potential difference is set up between zinc plate and the solution. This potential difference is termed the electrode potential of zinc.

33. Example 1) similarly, when copper is placed in a solution having Cu 2+  ions, it becomes positively charged with respect to solution. A potential difference is set up between the copper plate and the solution. The potential difference thus developed is termed as electrode potential of copper. The potential difference is estab­lished due to the formation of electrical double layer at the Following two changes occur when a metal rod is dipped in its salt solution, 

34. (a) Oxidation:  Metal ions pass from the electrode into solution leaving an excess of electrons and thus a negative charge on the electrode. The conversion of metal atoms into metal ions by the attractive force of polar water molecules. M →   M n  + ne - The metal ions go into the solution and the electrons remain on the metal making it negatively charged. The tendency of the metal to change into ions is known as electrolytic solution pressure. (b) Reduction:  Metal ions in solution gain electrons from the electrode leaving a positive charge on the electrode. Metal ions start depositing on the metal surface leading to a positive charge on the metal. M n +  +  ne -  →  M In the beginning, both these changes occur with different speeds but soon an equilibrium is established.

35. In practice, one effect is greater than the other, If first effect is greater than the second, the metal acquires a negative charge with respect to solution and If the second is greater than the first, it acquires positive charge with respect to solution, thus in both the cases a potential difference is set up. The magnitude of the electrode potential of a metal is a measure of its relative tendency to lose or gain electrons, i.e., it is a measure of the relative tendency to undergo oxidation (loss of electrons) or reduction (gain of electrons )

36. Depending on the nature of the metal electrode to lose or gain electrons, the electrode potential may be of two types: Oxidation potential:  When electrode is negatively charged with respect to solution, i.e., it acts as anode. Oxidation occurs. M →   M n +  + ne - Reduction potential:  When electrode is positively charged with respect to solution, i.e., it acts as cathode. Reduction occurs. M n +  + ne -  →  M It is not possible to measure the absolute value of the single electrode potential directly. Only the difference in potential between two electrodes can be measured experimentally.  It is, therefore, necessary to couple the electrode with another electrode whose potential is known. This electrode is termed as  reference electrode . The EMF of the resulting cell is measured experimentally.  The EMF  of the cell is equal to the sum of potentials on the two electrodes. Emf of  the cell =   E Anode  + E Cathode    = Oxidation potential of anode + Reduction potential of cathode Knowing the value of reference electrode, the value of other electrode can be determined.

37. Standard Electrode Potential In order to compare the electrode potentials of various electrodes, it is necessary to specify the concentration of the ions present in solution in which the electrode is dipped and the temperature of the half-cell. The potential difference developed between metal electrode and the solution of its ions of unit molarity (1M) at 25°C (298 K) is called  standard electrode potential. According to the IUPAC convention, the reduction potential alone be called as the electrode potential (E O ), i.e., the given value of electrode potential be regarded as reduction potential unless it is specifically mentioned that it is oxidation potential. Standard reduction potential of an electrode means that reduc­tion reaction is taking place at the electrode. If the reaction is reversed and written as oxidation reaction, the numerical value of electrode potential will remain same but the sign of standard potential will have to be reversed. Thus Standard reduction potential = - (Standard oxidation potential)   or Standard oxidation potential = - (Standard reduction potential)

38. Reference Electrode ( Standard Hydrogen Electrode, SHE) 1) Hydrogen electrode is the primary standard electrode. 2) It con­sists of a small platinum strip coated with platinum black as to adsorb hydrogen gas. 3) A platinum wire is welded to the platinum strip and sealed in a glass tube as to make contact with the outer circuit through mercury. 4) The platinum strip and glass tube is surrounded by an outer glass tube which has an inlet for hydrogen gas at the top and a number of holes at the base for the escape of excess of hydrogen gas. 5) The platinum strip is placed in an acid solution which has H +  ion concentration 1 M. 6) Pure hydrogen gas is circulated at one atmospheric pressure. 7) A part of the gas is adsorbed and the rest escapes through holes. This gives an equilibrium between the adsorbed hydrogen and hydrogen ions in the solution. H 2   2H +  + 2e - 8) The temperature of the cell is maintained at 25 C. By international agreement the standard hydrogen electrode is arbitrarily assigned a potential of exactly ± 0.000 . .. volt. 9) the hydrogen electrode thus obtained forms one of two half-cells of a voltaic cell. When this half-cell is connected with any other half-cell, a voltaic cell is constituted. The hydrogen electrode can act as cathode or anode with respect to other electrode. HE half reaction        Electrode potential  H 2  →  2H +  + 2e -   0.0 V (Anode)  2H +  + 2e -  →  H 2 0.0 V (Cathode)

39. Measurement of Electrode Potential The measurement of electrode potential of a given electrode is made by constituting a voltaic cell, i.e., by connecting it with a standard hydrogen electrode (SHE) through a salt bridge. 1 M solution is used in hydrogen half-cell and the temperature is maintained at 25 o C.

40. The EMF of the cell is measured either by a calibrated potentiometer or by a high resistance voltmeter, i.e., a valve voltmeter. The reading of the voltmeter gives the electrode potential of the electrode in question with respect to the hydrogen electrode. The standard electrode potential of a metal may be determined as it is the potential difference in volt developed in a cell consisting of two electrodes: the pure metal is contact with a molar solution of one of its ions and the standard hydrogen electrode.

41. Cell Representation Oxidation half reaction Reduction half reaction Zn|Zn 2+  ( aq )/Anode(-) || 2H( aq )| H 2  (g)/Cathode (+)    Zn →  Zn 2+  + 2e - 2H +  + 2e -  →  H 2 ↑ Determination of Standard Electrode Potential of Zn/Zn 2+  Electrode A zinc rod is dipped in 1 M zinc sulphate solution. This half-cell is combined with a standard hydrogen electrode through a salt bridge. Both the electrodes are con­nected with a voltmeter. The deflection of the voltmeter indicates that current is flowing from hydrogen electrode to metal electrode or the electrons are moving from zinc rod to hydrogen electrode. The zinc electrode acts as an anode and the hydrogen electrode as cathode and the cell can be represented as            

42. The EMF of the cell is 0.76 volt E Cell  = E o Anode  + E o Cathode 0.76 =  E o Anode  + 0 or E o Anode  = +0.76 V As the reaction on the anode is oxidation, i.e., Zn →  Zn 2+  + 2e, E o Anode  is the standard oxidation potential of zinc. This potential is given the positive sign.  

43. E o ox  (Zn/Zn 2+ ) = +0.76 volt So standard reduction potential of Zn, i.e., E o  (Zn/Zn 2 + ) = E o ox   = -(+0.76 ) = -0.76 volt The EMF of such a cell gives the positive value of standard oxidation potential of metal M.  The standard reduction potential ( E o ) is obtained by reversing the sign of standard oxidation potentia l.

44. Some Other Reference Electrodes Since a standard hydrogen electrode is difficult to prepare and maintain, it is usually replaced by other reference electrodes, which are known as secondary reference electrodes. These are convenient to handle and are prepared easily. Two important secondary reference electrodes are described here.

45. Calomel Electrode It consists of mercury at the bot­tom over which a paste of mercury- mercurous chloride is placed. A solution of potassium chloride is then placed over the paste. A platinum wire sealed in a glass tube helps in making the electrical contact. The electrode is connected with the help of the side tube on the left through a salt bridge with the other electrode to make a complete cell. 3) The potential of the calomel electrode depends upon the concentration of the potassium chloride solution . 4) If potassium chloride solution is saturated, the electrode is known as saturated calomel electrode (SCE) and if the potassium chloride solution is 1 N, the electrode is known as normal  calomel electrode  (NCE) while for 0.1 N potassium chloride solution, the electrode is referred to as decinormal calomel electrode (DNCE). The electrode reaction when the electrode acts as cathode is: 1/2 Hg 2 Cl 2  + e -  Hg + Cl - The reduction potentials of the calomel electrodes on hydrogen scale at 298K are as follows:

46. Concentration of KCl Reduction Potential Saturated KCl 0.2415 V 1.0NKCl 0.2800 V  0.1NKCl 0.3338 V The electrode potential of any other electrode on hydrogen scale can be measured when it is combined with calomel electrode. The emf of such a cell is measured. From the value of electrode potential of calomel electrode, the electrode poten­tial of the other electrode can be evaluated.

47. Ferrous – Ferric Electrode 1) An example is a Pt wire dipping in a solution containing ferrous and ferric ions. Such a cell is described as: Pt | Fe 2+  (C 1 ), Fe 3+  (C 2 ) The comma is used to separate two chemical species in the same solution. The electrode reaction is 2)  Fe 3+  + e –  →   Fe 2+ 3) The function of a Platinum wire is to pick up the electrons and provide electrical contact to the electrode.

48. Quin – Hydrone Electrode 1) Oxidation reduction electrodes can also be made with organic molecules that can exist in two different oxidation states. 2) A generally used material of this type is related to important biochemical oxidation – reduction reactions, is the system of hydroquinone, which can form the oxidation reduction system, between quinone (Q) and hydroquinone (QH 2 ). The presence of a Pt electrode in a solution containing these two species again clearly provides an electrode that can donate or accept electrons. 3) If hydro – quinone  is represented by QH 2  and quinone by Q, the electrode is abbreviated as Pt .|QH 2 ,Q,H +  (C) 4) This electrode is called a quin-hydrone electrode, because of the charged complex that is formed between QH 2  and Q.

49. Metal – Metal ion Electrode 1) An example of this is a metallic silver electrode in an AgNO 3  solution. The electrode is represented as Ag | Ag +  (C) and the electrode reaction is Ag +  + e –   →   Ag. 2) Note :    Very active metals react directly with water itself and cannot be used for such electrodes. 3) Another example is Silver-silver chloride electrode 4) This is another widely used reference electrode. It is reversible and stable and can be combined with cells containing chlorides without insetting liquid junctions.5) Silver chloride is deposited electrolytically on a silver or platinum wire and it is then immersed in a solution containing chloride ions. Its standard electrode potential with respect to the standard hydrogen electrode is 0.2224 V at 298 K. the electrode is represented as:   Ag|AgCl|Cl -  and the electrode reaction is,  AgCl+e - →  Ag + Cl -  

50. Amalgam Electrode In a variation of  the previous electrode, the metal is in the form of an amalgam, i.e., it is dissolved in mercury, rather than in the pure form. Electrical contact is made by a Pt wire dipping into the amalgam pool. The reaction is the same as in the metal–metal ion electrode, with the Hg playing no role. A sodium amalgam electrode is represented as Na [ in Hg at C 1 ] | Na +  (C 2 ).  

51. The working electrode (WE) represents the most important component of an electrochemical cell. It is at the interface between the WE and the solution that electron transfers of greatest interest occur. The selection of a working electrode material is critical to experimental success. Several important factors should be considered. Firstly, the material should exhibit  favorable redox behavior  with the analyte , ideally fast, reproducible electron transfer without electrode fouling. Secondly, the  potential window  over which the electrode performs in a given electrolyte solution should be as wide as possible to allow for the greatest degree of analyte characterization. Additional considerations include the cost of the material, its ability to be machined or formed into useful geometries, the ease of surface renewa l following a measurement, and toxicity.

52. The most commonly used working electrode materials are  platinum ,  gold ,  carbon , and  mercury . Among these, platinum is likely the favorite, demonstrating good electrochemical inertness and ease of fabrication into many forms. The biggest disadvantage to the use of platinum, other than its high cost , is that the presence of even small amounts of water or acid in the electrolyte leads to the reduction of hydrogen ion to form hydrogen gas (hydrogen evolution) at fairly modest negative potentials (E = -0.059 x pH). This reduction obscures any useful analytical signal.

53. Gold electrodes behave similarly to platinum, but have limited usefulness in the positive potential range due to the oxidation of its surface. It has been very useful, however, for the preparation of  modified electrodes  containing surface structures known as  self-assembled monolayers (SAMs) . Carbon electrodes allow scans to more negative potentials than platinum or gold, as well as good anodic potential windows. Carbon paste The most common form of carbon electrode is  glassy carbon , which is relatively expensive and difficult to machine.  Carbon paste electrodes  are also useful in many applications. These electrodes are made from a paste of finely granulated carbon mixed with an oil substrate like Nujol . The paste is then packed into a cavity in an inert electrode body. Carbon paste electrodes have the disadvantage of being prone to mechanical damage during use.

54. Mercury  has historically been a widely used electrode material, primarily as a spherical drop formed at the end of a glass capillary through which the liquid metal is allowed to flow. It displays an excellent potential window in the cathodic direction, but is severely limited in the anodic direction by its ease of oxidation. A   dropping mercury electrode (DME) , in which drops are formed and fall off repeatedly during a potential scan, being replaced by a “fresh” electrode every second or so, was commonly in past years the first electrode many students encountered in their studies . The toxicity of mercury has lead to a limited use these days, though it still is a very useful surface in methods that involve the preconcentration of a metallic analyte prior to potential scan, such as is done in  anodic stripping voltammetry (ASV) . Many practitioners now make use of  mercury films  formed on the surface of solid electrodes rather than the pure metal. Under these conditions, the small volume of the film allows analyte to concentrate at large values, with rapid diffusion times .

55. Material Advantages Limitations Pt available wire, flat plate & tube large range of sizes Pt-Rh alloy for rigidity low hydrogen overvoltage so cathodic potential range limited expensive Au configurations same as Pt larger cathodic potential range larger cathodic potential range anodic window limited by surface oxidation, expensive Carbon many types and configurations good cathodic potential range quality varies greatly hard to shape C-paste wide potential range low background current inexpensive unstable in flow cells cannot use in organic solvents Hg excellent cathodic window easy to “refresh” forms amalgams limited anodic window due to mercury oxidation toxic The approximate useful potential ranges for platinum, mercury, and carbon electrodes in aqueous electrolyte solutions along with those for platinum in a number of non-aqueous systems . Solid electrodes for voltammetric measurements are most often fabricated by encapsulating the electrode material in a nonconducting sheath of glass or inert polymeric material like Teflon, Kel -F (poly- chlorotrifluoroethylene ) or PEEK (poly- etheretherketone ). Most commonly, the exposed electrode material is in the form of a disk. Common commercially available disk diameters are 1.0, 3.0 and 10.0 mm . Electrodes of this size generally produce measured currents in the μA to low mA range for analytes at concentrations near 1 mM. Two common commercial sources of working electrodes, are ESA,

56. solvent The is a common but avoidable source of contamination, as methods of solvent purification are well known. Deionized water should be used to make the electrolyte for aqueous electrochemistry because tap water contains trace metal ions. For non-aqueous electrochemistry, small amounts of water contamination can usually be accommodated, but one should make an effort to minimize water content, as it is fairly easy to obtain water concentrations below 5 parts per million (ppm). Common solvents for non-aqueous electrochemistry include N,N-dimethylformamide (DMF), acetonitrile ( MeCN ), tetrahydrofuran (THF),  dimethylsulfoxide (DMSO), 1,2-Dimethoxyethane (DME), and dichloromethane (DCM), but electrochemical experiments can also be performed in supercritical gasses or ionic liquids. Sufficient drying is usually possible by passing the solvent through a column of activated alumina, followed by storage of the solvent over activated molecular sieves.

57. Other methods of drying include stirring or refluxing over a drying agent, such as calcium hydride (CaH 2 ), sodium hydride ( NaH ), sodium metal (Na),  phosphorus pentoxide (P 2 O 5 ), or molecular sieves, followed by distillation under an inert atmosphere. Non-aqueous solvents for electrochemistry should be stored under an inert atmosphere and over activated molecular sieves. Drying agents have to be carefully chosen according to the solvent because some are more effective with certain solvents and others violently react with the solvent. Consult the literature or your labmates for your labs accepted methods of drying.

58. in  electrochemistry, according to an IUPAC  definition, is an electrolyte containing chemical species that are not electroactive (within the range of potentials used) and which has an ionic strength and conductivity much larger than those due to the electroactive species added to the electrolyte. Supporting electrolyte is also sometimes referred to as  inert electrolyte  or  inactive electrolyte . Supporting electrolytes are widely used in electrochemical measurements when control of electrode potentials is required. This is done to increase the conductivity of the solution (to practically eliminate the so-called IR drop), to eliminate the transport of electroactive species by ion migration in the electric field, to maintain constant ionic strength, to maintain constant pH, etc supporting electrolyte