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BASICS OF ELECTROCHEMISTRY J.Daniell and M.Faraday – English scientists who gave basics for electrochemistry (18-19 century)

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Presentation on theme: "BASICS OF ELECTROCHEMISTRY J.Daniell and M.Faraday – English scientists who gave basics for electrochemistry (18-19 century)"— Presentation transcript:

1 BASICS OF ELECTROCHEMISTRY J.Daniell and M.Faraday – English scientists who gave basics for electrochemistry (18-19 century)

2 Electrochemistry is the study of interchange of chemical and electrical energy Electrochemical processes – redox processes 1. GALVANIC processes 2. ELECTROLYTIC processes COROSSIVE processes COROSSIVE processes

3 1. GALVANIC processes Chemical reaction is spontaneous and produces electricity It is also called a Galvanic cell or a Voltaic cell, named after Luigi Galvani, or Alessandro Volta respectively in Galvanic cell: In an electrochemical cell – an electric potential is created between two dissimilar metals Chemical energy is converted to Electrical energy Electrochemical cell – derives electrical energy from spontaneous redox reactions taking place within the cell

4 Rod of metal immersed in a salt solution is called an electrode METAL ELECTRODE HALF-CELL Metal electrode potential will be either positive or negative Metal electrode potential is measured in comparison with Hydrogen electrode potential

5 HYDROGEN ELECTRODE POTENTIAL Pt, H 2 /H + - such special electrode is assigned a potential of ZERO volts Standard hydrogen electrode allows H 2 gas molecules to interact directly with H + dissolved in water and with the electrons from the external circuit simultaneously

6 METAL ACTIVITY ROW Standard metal electrode potential is determined at solute concentrations of 1 Molar, gas pressure of 1 atm, and a standard temperature which is usually 25°C Li-K-Ca-Mg-Al-Zn-Cr-Fe-Cd-Pb-H-Cu-Hg-Ag-Pt-Au -3.0 -2.4 -0.76 -0.4 0.0V +0.8 +1.7 Standard half-cell potential is denoted by a degree sign as a superscript   o Me standardhydrogenelectrode 1.Measured against – standard hydrogen electrode 1 Molar 2.Concentration – 1 Molar 1 atmosphere 3.Pressure – 1 atmosphere 25°C 4.Temperature – 25°C METAL ELECTRODE POTENTIAL

7 GALVANIC CELLS Anode (negative electrode, active metal) is written on the left side, and Cathode (positive electrode, less active metal) – on the right side Such galvanic cell can be formulated: (-)Zn│ZnSO 4 ││ CuSO 4 │Cu(+) Galvanic cell – two half-cells of different metals connected with a conductor

8 C (+) Less active Me n+ + ne → Me 0 (reduction) GALVANIC PROCESSES A (-) More active Me 0 – ne → Me n+ (oxidation) It is customary to visualize the cell reaction in terms of two half- reactions: oxidation an oxidation half-reaction and reduction a reduction half-reaction By definition: Anode – the electrode where oxidation (loss of electrons) takes place; the anode attracts anions (Zinc electrode is the anode)Anode – the electrode where oxidation (loss of electrons) takes place; the anode attracts anions (Zinc electrode is the anode) Cathode – the electrode where reduction (gain of electrons) takes place; the cathode attracts cations (Copper electrode is the cathode)Cathode – the electrode where reduction (gain of electrons) takes place; the cathode attracts cations (Copper electrode is the cathode)

9 Cathode (+) Cu 2+ + 2e → Cu 0 (reduction) GALVANIC PROCESSES Anode (-) Zn 0 – 2e → Zn 2+ (oxidation) Total reaction: Zn 0 + CuSO 4 → Cu 0 + ZnSO 4 (redox)

10 ELECTROMOTIVE FORCE Cell potential can then expressed: Cell potential – E cell (also called electromotive force or EMF) has a contribution: E cell = reduction potential – oxidation potential from the anode which is a measure of its ability to lose electrons (called as oxidation potential), and from the anode which is a measure of its ability to lose electrons (called as oxidation potential), and from the cathode which has a contribution based on its ability to gain electrons (called as reduction potential) from the cathode which has a contribution based on its ability to gain electrons (called as reduction potential)

11 In our case: E cell =  o (Cu 2+ /Cu) –  o (Zn 2+ /Zn) E cell =  o (Cu 2+ /Cu) –  o (Zn 2+ /Zn) ELECTROMOTIVE FORCE E cell =  o reduction –  o oxidation E cell = 0.34 – (- 0.76) ═ 1.1 V Standard potential for the cell is equal: positive  valuenegative  value to more positive  o value minus more negative  o value  0 reduction – cathode standard potential, V  0 oxidation – anode standard potential, V

12 TYPES of GALVANIC CELLS Wetcelldry cell Wet cell battery has a liquid electrolyte, while dry cell has an electrolyte that is immobilized as a paste, which is only liquid enough to allow the flow of electrons There are different number of battery cell types which include: Batteries can also be made up for different types of materials : wet cell batterywet cell battery dry cell batterydry cell battery molten-salt cell batterymolten-salt cell battery reserve cell batteryreserve cell battery concentration cell batteryconcentration cell battery fuel cell batteryfuel cell battery zinc-carbon batteryzinc-carbon battery alkaline batteryalkaline battery nickel-zinc batterynickel-zinc battery lithium-ion baterrylithium-ion baterry lead-acid batterylead-acid battery Moltensalt Molten salt batteries use molten salt as an electrolyte – used in electric vehicles Reserve Reserve battery – a battery that is often stored in an unassembled form and only activates when the inner parts are assembled. Can be stored for long periods of time and are often a part of emergency kits. Provide power for a few minutes

13 CONCENTRATION CELL Is a kind of a wet galvanic cell that has two equivalent half-cells of the same material differing only in concentrations Rods of Cu immersed in salt solution (CuSO 4 ) of different concentration Potential of half-cell:  (Me) =  o (Me) – (0.0592/n)·log 10 [Me n+ ] (A) Cu | CuSO 4 (0.05 M) || CuSO 4 (2.0 M) | Cu (C) (A) Cu 0 – 2e → Cu 2+ (0.05 M) (C) Cu 2+ (2.0 M) + 2e → Cu 0 One can calculate the potential developed by such a cell using the Nernst Equation

14 Zinc-carbon Zinc-carbon cell battery Electrolyte – immobilized as a paste (mixture of NH 4 Cl and starch or C powder) (–) Anode – zinc box (cell is encased in Zinc casing) (+) Cathode – carbon (graphite) rod, surrounded by a layer of MnO 2 (-A) Zn  NH 4 Cl  MnO 2  C (+C) DRY HERMETIC CELL Cells are often encased in a plastic or metal casing, with two points of a negative (-) and a positive (+) sides

15 REDOX REACTIONS: (-A) (-A) Zn + 2OH − - 2e → ZnO + H 2 O (oxidation) (+C) (+C) 2MnO 2 + H 2 O + 2e− →Mn 2 O 3 + 2OH − (reduction) ALKALINE CELL Anode Anode – zinc powder Cathode Cathode – MnO 2 Electrolyte Electrolyte – KOH The usual voltage for an alkaline battery is 1.5 V

16 FUEL CELL convertshydrogenoxygen waterproduceselectricity Hydrogen fuel cell converts the chemicals hydrogen and oxygen into water, and in the process it produces electricity Fuelcellelectrochemicalenergyconversion hydrogenoxygen Fuel cell is an electrochemical energy conversion device – most fuel cells in use today use hydrogen and oxygen as the chemicals Battery storedinside “goes dead” Battery has all of its chemicals stored inside, and converts those chemicals into electricity. This means that a battery eventually “goes dead” Withafuelcellconstantlyflownever “goes dead” With a fuel cell, chemicals constantly flow into the cell so it never “goes dead” As long as there is a flow of chemicals into the cell, the electricity flows out…

17 Hydrogenfuelcell anodecathodeOxygen hydrogen Hydrogen fuel cell operates similar to a battery – it has two electrodes, an anode and a cathode, separated by a membrane. Oxygen passes over one electrode and hydrogen over the other FUEL CELL Hydrogen negativelyelectronspositivelyions Electronselectricalenergy Hydrogen reacts to a catalyst on the anode that converts the hydrogen gas into negatively charged electrons (e-) and positively charged ions (H+). Electrons flow out of the cell to be used as electrical energy Hydrogenions combinewithoxygenproducewater Hydrogen ions move through the electrolyte membrane to the cathode where they combine with oxygen and the electrons to produce water neverrunout Unlike batteries, fuel cells never run out Reactions involved in a fuel cell are as follows: (-) Anode side: 2H 2 - 4e - => 4H + (+)Cathode side: O 2 + 4H + + 4e - => 2H 2 O Net "redox" reaction: 2H 2 + O 2 => 2H 2 O

18 ALKALINE FUEL CELL Alkalinefuelcells electricalenergy water Alkaline fuel cells (AFCs) were one of the first fuel cell technologies developed, and they were the first type widely used in the U.S. space program to produce electrical energy and water on-board spacecrafts Anode Reaction (Oxidation): 2H 2 + 2O 2− → 2H 2 O + 4e − Cathode Reaction (Reduction): O 2 + 4e – → 2O 2− Overall Cell Reaction (Redox): 2H 2 + O 2 → 2H 2 O The reactions involved in a fuel cell are as follows: potassium hydroxideelectrolyte These fuel cells use a solution of potassium hydroxide in water as the electrolyte and can use a variety of metals as a catalyst at the anode and cathode

19 FUEL CELLS – PROBLEMS TO BE SOLVED Regenerative Fuel Cells: Regenerative Fuel Cells: produce electricity from H 2 and O 2 and generate heat and water as byproducts, just like other fuel cells. However, regenerative fuel cell systems can also use electricity from solar power or some other source to divide the excess water into oxygen and hydrogen fuel by "electrolysis" PRODUCTION: Hydrogen can be produced PRODUCTION: Hydrogen can be produced using diverse, domestic resources including fossil fuels, such as natural gas and coal (with carbon sequestration); nuclear; biomass; and other renewable energy technologies, such as wind, solar, geothermal, and hydro-electric power. The overall challenge to hydrogen production is cost reduction STORAGE: On-board hydrogen storage STORAGE: On-board hydrogen storage for transportation applications continues to be one of the most technically challenging barriers to the widespread commercialization of hydrogen-fueled vehicles. Specific hydrogen storage material classes: on-board reversible metal hydrides, hydrogen adsorbents, and chemical hydrogen storage materials

20 Difference between Cell and Battery battery stackpilegalvaniccells – A battery is basically nothing but a stack or pile of galvanic cells – battery consists of multiple cells voltaic pileelectricalbattery Volta was the inventor of the voltaic pile, the first electrical battery Usualvoltagealkaline cell1.5 V addingmorecells Usual voltage for a single alkaline cell is 1.5 V and the voltage can be increased by adding on more cells 9-Voltbattery 6 alkaline cells Depending on application of the battery, the cells are combined to provide a higher voltage, for example a 9-Volt battery would have 6 alkaline cells with a 1.5 V charge

21 2. ELECTROLYTIC processes decompositionwater hydrogenoxygenbauxitealuminium Important examples of electrolysis are the decomposition of water into hydrogen and oxygen, and bauxite into aluminium and other chemicals in Electrolytic cell: Electrical energy is converted to Chemical energy Electrolytic cell redox reactionelectricalenergyapplied Electrolytic cell is an electrochemical cell that undergoes a redox reaction when electrical energy is applied decompose electrolysis lysisbreakup It is most often used to decompose chemical compounds, in a process called electrolysis (the Greek word lysis means to break up)

22 AnodeA Anode is positive electrode +(A) CathodeC Cathode – negative electrode - (C) ELECTROLYSIS Electrolyticcellelectrolyte cathodeanode Electrolytic cell has three component parts: an electrolyte and two electrodes (a cathode and an anode) electrolytesolutionwater The electrolyte is usually a solution of water in which ions are dissolved Moltensalts Molten salts such as sodium chloride are also used as electrolytes

23 (+A) Cl - - e → Cl o (oxidation) (Cl + Cl → Cl 2 ) flow electrodes charge-transferring reactions Electrolyte provides ions that flow to the electrodes, where charge-transferring reactions take place Molten salt: NaCl → Na + + Cl - (-C) Na + + e → Na o (reduction)

24 ELECTROLYSIS – application 1. Rechargeablebattery 1. Rechargeable battery – cell battery, which acts as a galvanic cell when discharging (converting chemical energy to electrical energy), and an electrolytic cell when being charged (converting electrical energy to chemical energy 2. Electrorefining 2. Electrorefining – process of electrolytic refining of metals is used to extract impurities from crude metals. Here in this process a block of crude metal is used as anode, a diluted salt of that metal is used as electrolyte and plates of that pure metal is used as cathode 3. Electrolysis 3. Electrolysis – used to produce acids, alkalis, non-metals (H 2, Cl 2, O 2, As)

25 ELECTROLYSIS – application 4. Electroforming 4. Electroforming – reproduction (making copies) of objects by electro- deposition. The surface of the wax mould which bears exact impression of the object, is coated with graphite powder in order to make it conducting. Then the mould is dipped into the electrolyte solution as cathode and metal will be deposited on the graphite coated surface of the mould 5. Electroplating 5. Electroplating – cathode is an object on which the electroplating to be done (e.g. iron being copper-platted by using copper electroplating)

26 RECHARGEABLE BATTERY rechargeablebattery A rechargeable battery (e.g. lead-acid battery) acts as: Galvanic cell dischargingGalvanic cell – when discharging (converting chemical energy to electrical energy), and Electrolytic cell beingchargedElectrolytic cell – when being charged (converting electrical energy to chemical energy)

27 Discharging In the discharged state both the positive and negative plates become lead(II) sulfate (PbSO 4 ) and the electrolyte loses much of its dissolved sulfuric acid and becomes primarily water. The discharge process is driven by the conduction of electrons from the negative plate back into the cell at the positive plate in the external circuit Negative plate Negative plate reaction (Anode Reaction): Pb(s) + HSO 4 − (aq) → PbSO 4 (s) + H + (aq) + 2e - Positive plate Positive plate reaction (Cathode Reaction): PbO 2 (s) + HSO 4 − (aq) + 3H + (aq) + 2e - → PbSO 4 (s) + 2H 2 O(l) Total reaction Total reaction can be written: Pb(s) + PbO 2 (s) + 2H 2 SO 4 (aq) → 2PbSO 4 (s) + 2H 2 O(l)

28 Charging In the charged state, each cell contains negative plates of elemental lead (Pb) and positive plates of lead(IV) oxide (PbO 2 ) in an electrolyte of approximately 33.5% v/v sulfuric acid (H 2 SO 4 ). The charging process is driven by the forcible removal of electrons from the positive plate and the forcible introduction of them to the negative plate by the charging source Negative plate Negative plate reaction: PbSO 4 (s) + H + (aq) + 2e - → Pb(s) + HSO 4 − (aq) Positive plate Positive plate reaction: PbSO 4 (s) + 2H 2 O(l) → PbO 2 (s) + HSO 4 − (aq) + 3H + (aq) + 2e - Overcharging with high charging voltages generates oxygen and hydrogen gas by electrolysis of water, which is lost to the cell. Periodic maintenance of lead acid batteries requires inspection of the electrolyte level and replacement of any water that has been lost

29 Ion motion During discharge, H+ produced at the negative plates and from the electrolyte solution moves to the positive plates where it is consumed, while HSO 4 − is consumed at both plates. The reverse occurs during charge Due to the freezing-point depression of the electrolyte, as the battery discharges and the concentration of sulfuric acid decreases, the electrolyte is more likely to freeze during winter weather when discharged

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