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Electrochemical Cells, Types of Electrodes, Nernst equation & its application

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Conducto ppt

  1. 1. CONDUCTOMETRY LAB PRESENTATION  Electrochemical cells  Types of electrodes  Nernst Equation & its Application
  2. 2. Prepared & Presented by: Sidra Safdar Durrani M.Sc. Final year Presented to: Miss Atiya Firdous For the Course of: Conductometry lab
  3. 3. ELECTROCHEMICAL CELLS MITSUBISHI IMIEV - PURE ELECTRIC CAR Powered by a 330 V Li-Ion Rechargeable battery Top Speed km/h 130 $ 36,000 US $ 50,000 Can Plugs into your house and takes 14 hours to charge -100 km for $ 0.60
  4. 4. ELECTROCHEMICAL CELLS Electrochemical cells are Batteries Alkaline Batteries Car Batteries Cell Phone batteries
  5. 5. ELECTROCHEMICAL CELLS Electrochemical cells are Batteries Lithium Coin Cell Space Ship Batteries Ni-Metal Hydride
  6. 6. ELECTROCHEMICAL CELLS An electrochemical cell – a system of salt electrodes , bridge that electrolytes , allow and oxidation and reduction reactions to occur and electrons to flow through an external circuit .
  7. 7. FOR ANY CELL  Oxidation a l w a ys occurs at the anode and reduction at the cathode.  Electrons flow through the wire and go from anode to cathode.  Anions (- ions) migrate to the anode and cations (+ions) migrate towards the cathode usually through the salt bridge.
  8. 8. ANALYZING THE ELECTROCHEMICAL CELLS  The reaction that is higher on the reduction chart is the reduction and the lower is oxidation and is written in reverse.  The salt bridge allows ions to migrate from one half-cell to the other without allowing the solutions to mix.
  11. 11. A H 2/ A g E L E C T R O C H E M I C A L C E L L WITH A KNO3 SALT BRIDGE
  12. 12. The term ―electrochemical cell‖ is often used to refer to a: Voltaic Cell – one with a spontaneous reaction SOA over SRA on the activity series Eocell greater than zero = spontaneous Electrolytic cell – one with a nonspontaneous reaction SOA below SRA – i.e. zinc sulfate and lead solid cell Eocell less than zero= nonspontaneous
  13. 13. VOLTAIC CELLS (AKA GALVANIC CELL) A device that spontaneously produces electricity by redox Composed of two half-cells; which each consist of a metal rod or strip immersed in a solution of its own ions or an inert electrolyte. The electrons flow from the anode to the cathode (―a before c‖) through an electrical circuit rather than passing directly from one substance to another A porous boundary to maintain cell neutrality (mostly is a salt bridge containing an inert aqueous electrolyte ; such as Na2SO4(aq) or KNO3(aq)).
  14. 14. ELECTROLYTIC CELLS Electrolytic Cell – a cell in which a nonspontaneous redox reaction is forced to occur; a combination of two electrodes, an electrolyte and an external power source. Electrolysis – the process of supplying electrical energy to force a nonspontaneous redox reaction to occur 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
  15. 15. TYPES OF ELECTRODES TYPES OF ELECTRODES ON THE BASIS OF USES:- Electrodes are classified into two types on the basis of their uses:•Reference Electrode •Indicator Electrode TYPES OF ELECTRODES ON THE BASIS OF THEIR WORKING:- Whether the electrode is indicator or reference, it is classified into the following types: •Electrode of First kind •Electrode of Second kind •Electrode of Third kind
  16. 16. REFERENCE ELECTRODE:A reference electrode is an electrode which has a stable and well-known electrode potential. A half-cell with an accurately known electrode potential, Eref, that is independent of the concentration of the analyte or any other ions in the solution and always treated as the lefthand electrode. Common examples are: Normal hydrogen electrode (NHE Saturated calomel electrode (SCE) Copper-copper (II) sulfate electrode (CSE) Silver chloride electrode pH-electrode
  17. 17. INDICATOR ELECTRODE:- It can serve either as an anode or a cathode, depending on the applied polarity. One of the electrodes in some ""classical two-electrode"" cells can also be considered a ""working"" (""measuring,"" ""indicator,"" or ""sensing"") electrode). Indicator electrode, responds rapidly and reproducibly to changes in the concentration of an analyte ion (or groups of analyte ions).
  18. 18. ELECTRODE OF FIRST KIND:This type of electrode always contains a single chemical element in contact with an electrolyte solution containing its own ions, or charged species originating from this element (e.g., silver immersed in a silver nitrate solution,or hydrogen gas and oxonium ions, etc.). The equilibrium electrode potential of this electrode is a function of the activity of the ions in the solution and the activity of the chemical element. Subtypes: metal/metal ion electrode, gas electrode, amalgam electrode.
  19. 19. Silver chloride electrode:Silver/ chloride electrode is a type of reference electrode, commonly used in electrochemical measurements. For example, it is usually the internal reference electrode in pH meters. As another example, the silver chloride electrode is the most commonly used reference electrode for testing cathodic protection corrosion control systems in sea water environments. The electrode functions as a redox electrode and the reaction is between the silver metal (Ag) and its salt — silver chloride (AgCl, also called silver (I) chloride). .
  20. 20. ELECTRODE OF SECOND KIND:A metal electrode assembly with the equilibrium potential being a function of the concentration of an anion in the solution.Typical examples are the silver/silver-chloride electrode and the calomel electrode. Contrast with electrode of the first kind and electrode of the third kind, the assembly consists of a metal, in contact with a slightly soluble salt of this metal (or metal - oxide), immersed in a solution containing the same anion as that of the metal salt (e.g., silver/silver chloride/potassium chloride solution). The potential of the metal is controlled by the concentration of its cation in the solution, but this, in turn, is controlled by the anion concentration in the solution through the solubility product of the slightly soluble metal salt.
  21. 21. ELECTRODE OF THIRD KIND:A metal electrode assembly with the equilibrium potential being a function of the concentration of a cation, other than the cation of the electrode metal, in the solution. These have been used, with limited success, in sensors for metal ions for metals that are not stable in aqueous solutions, e.g., calcium and magnesium. Contrast with electrode of the first kind and electrode of the second kind, the assembly consists of a metal in contact with two slightly soluble salts (one containing the cation of the solid metal, the other the cation to be determined, with both salts having a common anion) immersed in a solution containing a salt of the second metal, (e.g. zinc metal/zinc oxalate/calcium oxalate/calcium salt solution).
  22. 22. NERNST EQUATION In electrochemistry, the Nernst equation is an equation that relates the equilibrium reduction potential of a halfcell in an electrochemical cell (or the total voltage (electromotive force) for a full cell) to the standard electrode potential, temperature, activity, and reaction quotient of the underlying reactions and species used. It is named after the German physical chemist who first formulated it, Walther Nernst.
  23. 23. APPLICATIONS OF EQUATION 1. Oxygen and the aquatic environment The presence of oxygen in the atmosphere has a profound effect on the redox properties of the aquatic environment— that is, on natural waters exposed directly or indirectly to the atmosphere, and by extension, on organisms that live in an aerobic environment. This is due, of course, to its being an exceptionally strong oxidizing agent and thus a lowlying sink for electrons from most of the elements and all organic compounds. Those parts of the environment that are protected from atmospheric oxygen are equally important because it is only here that electrons are sufficiently available to produce the "reducing" conditions that are essential for processes varying from photosynthesis to nitrogen fixation.
  24. 24. APPLICATIONS OF EQUATION 2. Analytical chemistry application A very large part of Chemistry is concerned, either directly or indirectly, with determining the concentrations of ions in solution. Any method that can accomplish such measurements using relatively simple physical techniques is bound to be widely exploited. Cell potentials are fairly easy to measure, and although the Nernst equation relates them to ionic activities rather than to concentrations, the difference between them becomes negligible in solutions where the total ionic concentration is less than about 10–3 M.
  25. 25. APPLICATIONS OF EQUATION 3. Determination of solubility products The concentrations of ions in equilibrium with a sparingly soluble salt are sufficiently low that their direct determination can be quite difficult. A far simpler and common procedure is to set up a cell in which one of the electrode reactions involves the insoluble salt, and whose net cell reaction corresponds to the dissolution of the Salt. Nernst equation can be modified in the given form in order to determine Ksp Ecell = E - (0.0591 / n) log Ksp