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Introduction to electrochemistry - Basics of all techniques -

Introduction to electrochemistry - Basics of all techniques -. Electrochemistry is the study of phenomena at electrode-solution interfaces. Two quite different aspects of the field of electrochemistry. An introduction to redox equilibria and electrode potentials.

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Introduction to electrochemistry - Basics of all techniques -

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  1. Introduction to electrochemistry- Basics of all techniques -

  2. Electrochemistryis the study of phenomena at electrode-solution interfaces

  3. Two quite different aspects of the field of electrochemistry

  4. An introduction to redox equilibria and electrode potentials The more negative the value, the stronger reducing agent the metal is. The more positive the value, the stronger oxidising agent the metal ion is.

  5. Redox Potentials for non-metal and other systems • Chlorine gas is the strongest oxidising agent (E° = +1.36 V). • A solution containing dichromate(VI) ions in acid is almost as strong an oxidising agent (E° = +1.33 V). • Iron(III) ions are the weakest of the three new ones (E° = +0.77 V). • None of these three are as strong an oxidising agent as Au3+ ions (E° = +1.50 V).

  6. Looking at this from an equilibrium point of view Suppose you have a piece of magnesium in a beaker of water. There will be some tendency for the magnesium atoms to shed electrons and go into solution as magnesium ions. The electrons will be left behind on the magnesium.

  7. A dynamic equilibrium will be established when the rate at which ions are leaving the surface is exactly equal to the rate at which they are joining it again.

  8. At that point there will be a constant negative charge on the magnesium, and a constant number of magnesium ions present in the solution around it.

  9. Copper is less reactive and so forms its ions less readily. Any ions which do break away are more likely to reclaim their electrons and stick back on to the metal again. You will still reach an equilibrium position, but there will be less charge on the metal, and fewer ions in solution.

  10. standard hydrogen electrode As the hydrogen gas flows over the porous platinum, an equilibrium is set up between hydrogen molecules and hydrogen ions in solution. The reaction is catalysed by the platinum.

  11. The standard hydrogen electrode is attached to the electrode system you are investigating - for example, a piece of magnesium in a solution containing magnesium ions.

  12. Magnesium has a much greater tendency to form its ions than hydrogen does. The position of the magnesium equilibrium will be well to the left of that of the hydrogen equilibrium. That means that there will be a much greater build-up of electrons on the piece of magnesium than on the platinum.

  13. What if you replace the magnesium half cell by a copper one?

  14. standard electrode potentials • The standard electrode potential of a metal / metal ion combination is the electro-motive force (emf) measured when that metal / metal ion electrode is coupled to a hydrogen electrode under standard conditions.

  15. In the copper case:

  16. The two equilibria which are set up in the half cells are:

  17. Obviously, the voltmeter will show that the zinc is the negative electrode, and copper is the (relatively) positive one.

  18. (a) Galvanic and (b) electrolytic cells

  19. Scope of electrochemistry Introduction Introduction • Investigation of chemical phenomena associated with a charge transfer reaction • To assure electroneutrality two (or more) half-reactions take place in opposite directions (oxidation/reduction) • If the sum of free energy changes at both electrodes is negative electrical energy is released battery • If it is positive, external electrical energy has to be supplied to oblige electrode reactions  electrolysis

  20. Reactions and electrodes • The overall chemical reaction taking place in a cell is made up of two independent half-reactions, which describe the real chemical changes at the two electrodes. • Most of the time one is interested in only one of these reactions, and the electrode at which it occurs is called the working (or indicator) electrode, coupled with an electrode that approaches an ideal nonpolarizable electrode of known potential, called the reference electrode. In experiments, the current is passed between the working electrode and an auxiliary(or counter) electrode. • Three electrodes are frequently placed in three compartments separated by a sintered-glass disk.

  21. Reference electrode: • A reference electrode is used in measuring the working electrode potential of an electrochemical cell. • The reference electrode acts as a reference point for the redox couple.  A Luggin capillary is often used to position the sensing point of a reference electrode to a desired point in a cell.

  22. Reference electrode ● The potential of the working electrode is monitored relative to a separate reference electrode, positioned with its tip near the working electrode. ● The internationally accepted primary reference is the standard hydrogen electrode (SHE) or normal hydrogen electrode (NHE), which is Pt/H2(a=1)/H+(a=1,aqueous) ●By far the most common reference is the saturated calomel electrode (SCE) and the Silver/Silver Chloride (Ag/AgCl) electrodes.SCEis Hg/Hg2Cl2/KCl (sat’d in water). Its potential is 0.242 V vs. NHE.

  23. The device minimizes any iR drop in the electrolyte associated with the passage of current in an electrochemical cell.

  24. working electrode A fixed potential difference is applied between the working electrode and the reference electrode. This potential drives the electrochemical reaction at the working electrode's surface.  The current produced from the electrochemical reaction at the working electrode is balanced by a current flowing in the opposite direction at the counter electrode.

  25. Materials of working electrode • A wide variety of working electrodes are now available. Originally the carbon paste electrode was developed but this was soon replaced by more "convenient" and stable carbon-based working electrodes including those made from glassy carbon, pyrolytic carbon and porous graphite.   Metals such as platinum, gold, silver, nickel, mercury, gold-amalgam and a variety of alloys are now also commonly used as working electrode materials.

  26. Choice of working electrode • Carbon paste electrodes cannot be used with mobile phases containing high amounts of organic modifier because the electrode will dissolve unless a polymeric binder is used. • The optimal working electrode choice is dependent upon many factors, including the usable applied potential range, involvement of the electrode in the redox reaction, and kinetics of the electron transfer reaction. 

  27. Three-electrode cell and notation for the different electrodes

  28. Potential window • A working electrode will only function within a defined potential window. For example, electrolysis of many compounds will readily occur on a glassy carbon working electrode up to approximately 1300mV vs. a silver/silver chloride reference electrode. • The applied potential to the working electrode is dependent upon both the working electrode material and the pH of the mobile phase.

  29. The potential window of various working electrodes under acidic and basic conditions.

  30. Kinetics of the Electron Transfer Reaction • Electron-transfer reactions can be either kinetically fast or slow. For a fast reaction most of an analyte will react at the working electrode's surface. For slow reactions not all of the analyte reaching the working electrode's surface will have time to react.  To drive the electrolytic reaction at a faster rate a much higher potential or "overpotential" must be used.  It is usually found that organic species react more favorably on one particular working electrode material than another.

  31. Factors affecting electrode reaction rate In general, the electrode reaction rate is governed by rates of processes such as: • Mass transfer (e.g., from the bulk solution to the electrode surface). (2) Electron transfer at the electrode surface. (3)Chemical reactions preceding or following the electron transfer. (4)Other surface reactions. ◆ The magnitude of this current is often limited by the inherent sluggishness of one or more reactions called rate-determining steps.

  32. Conditions for electrochemical experiments Reproducible experimental conditions must be given Interfering side effects must be avoided as Migration effects High solution resistance -these effects can be minimised by adding an inert supporting electrolyte (around 1 mol/L) Undefined or large diffusion layer • A complete study of the electrode process requires the measurement of kinetic as well as thermodynamic parameters.

  33. Faradaic and nonfaradaic processes • Charges (e.g., electrons) are transferred across the electrode-solution interface and causes oxidation or reduction to occur. Since these reactions are governed by Faraday’s law, they are called faradaic processes. • Under some conditions, processes such as adsorption and desorption can occur, and the structure of the electrode-solution interface can change with changing potential or solution composition, these processes are called nonfaradaic processes.

  34. Capacitance and charge of an electrode • The behavior of the electrode-solution interface is analogous to that of a capacitor. When a potential is applied across a capacitor, charge will accumulate on its electrode plates. • At a given potential there will exist a charge on the metal electrode, qM, and a charge in the solution, qs. At all times, qM=-qs. • At a given potential the electrode-solution interface is characterized by a double-layer capacitance, Cd, typically in the range of 10 to 40μF/cm2.

  35. The nature of electrode reactions • Electrode reactions are heterogeneous and take place in the interfacial region between electrode and solutiondiffusion layer • The charge separation at each electrode is represented by acapacitance • the difficulty of charge transfer by a resistance • The electrode can act as (1) a sourceof electrons (cathode) reduction ,(2) a sink of electrons transferred from species in solution (anode) oxidation • The amount of electrons transferred is related to the current flowing between the two electrodes

  36. Thermodynamics and kinetics • Thermodynamics and kinetics • The potential at which a reduction or oxidation takes place (measured relative to the normal hydrogen electrode) is given by theNernst equation i : stoichiometric numbers: positive for reduced species, negative for oxidised species E0 : standard electrode potential ci : concentration(ai has to be applied if activity coefficient is not 1) E = E0 – (RT/nF) i ln ci

  37. Thermodynamics and kinetics ●The concentration of species at the electrode interface depends on its mass transport coefficientkd ●The rate of the electrode reaction is expressed by the standard rate constantkf0which is the rate when E = E0 • reversible reaction  kf0>> kd • irreversible reversible reaction  kf0<< kd, an overpotential  has to be applied additionally to overcome this kinetic barrier ●A behaviour in between these extremes is called quasireversiblereaction

  38. Reference • A.J. Bard, L.R. Faulkner, Electrochemical Methods—Fundamentals and Applications. New York: John Wiley & Sons,1980.

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