Digital content preparation-Batteri.pptx

Reading materials about Batteries

1. ENGINEERING CHEMISTRY (25CTXXX) MODULE 2- UNIT - 4 Topic: ELECTROCHEMICAL ENERGY SYSTEMS By Dr. Y S L V Narayana Assistant Professor School of Applied Sciences & Humanities Department of Chemistry VFSTR, Vadlamudi

2. Contents Introduction Classification of Batteries Primary and Secondary Batteries Construction, working and applications of Lead-acid battery Construction, working and applications of Li-ion battery

3. Contents to Introduction Redox reactions Electrode potential EMF of an electrochemical cell Electrochemical series Electrochemistry Anode and Cathode Motivation

4. Electrochemistry is crucial for engineering students due to its applications in energy conversion and storage, materials science, and industrial processes.  Understanding electrochemistry is essential for developing sustainable technologies like batteries, fuel cells, and electrolysis, as well as for advancements in areas like corrosion control and electroplating Motivation

5. ELECTROCHEMISTRY Electrochemistry - a branch of chemistry that deals with the interrelation ship between electrical energy and chemical energy based on spontaneous oxidation-reduction reactions .

6. Redox or Oxidation-Reduction reactions In oxidation: a species looses one or more electrons resulting in an increase in its oxidation number . ( Oxidation is the loss of one or more electrons ) In reduction : the species gains one or more electrons; there is a decrease in oxidation number. ( Reduction is the gain of one or more electrons ). A reduction process necessarily accompanies an oxidation process since electrons can be lost by a species only when there is a counterpart to receive them. 6

7. Anode (Negative, Oxidation): Anode and Cathode The electrode where oxidation (loss of electrons) takes place. Electrons flow from the anode into the external circuit.  Cathode (Positive, Reduction): The electrode where reduction (gain of electrons) takes place.  Electrons flow from the external circuit into the cathode. An electrode is a component that facilitates the flow of electrical current between a circuit's metallic portion and non-metallic parts, such as electrolytes, semiconductors, or vacuums . 

8. Electrode potential Natural tendency of a metal to undergo Oxidation (or) Reduction reaction when it is in contact with an aqueous solution of its own ions is expressed in terms of electrode potential. Definition : It is the measure of the tendency of a metallic electrode to lose (Oxidation Potential) or gain (Reduction potential) electrons, when it is in contact with its own salt solution of 1 M Concentration. (OR) Electrode potential is the potential difference between the metal and its salt solution . Ex: Zinc rod in contact with zinc sulphate is called zinc electrode . Similarly copper rod in contact with copper sulphate solution is called copper electrode

9. The tendency of an electrode to lose electrons is direct measure of its tendency to get oxidized and this tendency is called “ Oxidation Potential ”. Similarly, the tendency of an electrode to gain electrons is a direct measure of its tendency to get reduced and this tendency is known as “ Reduction Potential ”. Units for electrode potential is “Volts”. Electrode potential Thus “ Standard Electrode Potential ” of a metal is the measure of tendency of a metallic electrode to lose or gain electrons, when it is in contact with its own salt solution of unit molar concentration at 25 o C

10. Half cell - Since a cell is a combination of two electrodes, each electrode is referred to as a single electrode or half- cell. Half cell- consists of a metal in contact with a solution of its own ion .” The electrode where oxidation half reaction occurs is called anode half cell. The where reduction half reaction occurs is called cathode half cell . ELECTROCHEMCAL CELL

11. There are two types of Electrochemical cells. 1.Galvanic cell or voltaic cell: The one, which converts chemical energy into electrical energy . 2.Electrolytic cell: The one, which converts electrical energy into chemical energy Electrochemical cells

12. Cell representation Anode half cell Cathode half cell

13. EMF of an electro chemical cell The cell electromotive force, or cell EMF, is the net voltage between the oxidation and reduction half-reactions taking place between two half cells (algebraic sum of the single electrode potentials ) . E.M.F of galvanic cell can be calculated by using,  E cell = E Right - E Left where E cell = EMF of the cell E right = reduction potential of right-hand side electrode E left = reduction potential of left-hand side electrode Standard Electrode potential: of a re When all the substances involved in the reaction are in their standard states that is solutions are at 1M concentrations, gases at 1 atm pressure and solids and liquids are in pure form with all at 25 o C.

14. Cell EMF The larger the difference between E  red values, the larger E  cell . In a voltaic (galvanic) cell (spontaneous) E  red (cathode) is more positive than E  red (anode). E  cell = E  red (cathode) - E  red (anode) Note Cell Diagram: Anode || Cathode

15. Electrochemical Series What Is Electrochemical Series? Importance of Electrochemical Series? The electrochemical series, also known as the activity series, is  a list that describes the arrangement of elements in the order of their increasing electrode potential values. The series has been established by measuring the potential of various electrodes versus standard hydrogen electrodes (SHE). The arrangement of different electrode potentials of different electrodes from highest -ve to highest +ve are called electrochemical series. It's a crucial tool in chemistry for predicting the outcome of redox reactions and understanding the behavior of elements in electrochemical cells. 

16. Electrochemical Series

17. Applications of Electrochemical series 1 . Oxidizing and Reducing Strengths The electrochemical series helps us to identify a good  oxidizing agent  or reducing agent . All the substances appearing on the top of the electrochemical series are good oxidizing agents 2 . Calculation of Standard emf (E ) of Electrochemical Cell The standard emf of the cell is the sum of the standard reduction potential of the two half cells: reduction half cell and oxidation half cell Example: For a reaction, 2Ag + ( aq ) + Cd → 2Ag + Cd +2 ( aq ) The standard reduction potential given are: Ag + / Ag = 0.80 volt, Cd +2 / Cd = -0.40 volt

18. 3. Predicting the Feasibility of Redox Reaction Any redox reaction would occur spontaneously if the free energy change (ΔG) is negative. The free energy is related to cell emf in the following manner: ΔG o = nFE o Where n is the number of electrons involved, F is the Faraday constant, and E o is the cell emf. ΔG o can be negative if E o is positive. When E o is positive, the cell reaction is spontaneous and serves as a source of electrical energy. If it comes out to be negative, then the spontaneous reaction cannot take place. The resultant value of E o for redox reaction is important in predicting the stability of a metal salt solution when stored in another metal container .

19. Galvanic Cells and Batteries The electrochemical series is at the heart of galvanic cells and batteries. The potential difference between two half-reactions in a galvanic cell generates electrical energy. The electrochemical series helps select the right combination of electrode materials to create the desired voltage and current. In batteries, the series is used to design and optimize the electrochemical reactions that produce electrical energy, leading to the development of batteries with varying capacities, sizes, and applications in electronics, transportation, and renewable energy storage.

20. Batteries (Electrochemical Cells)

21. A container consisting of one or more cells, in which chemical energy is converted into electricity and used as a source of power Batteries The cell is a single unit, so it is light and compact, whereas the battery is a collection of cells. 

22. Cell vs Battery

23. Batteries An electrochemical cell ( Battery) is a device capable of either deriving electrical energy from chemical reactions, or facilitating chemical reactions through the introduction of electrical energy . A battery depending on the need can be made up of with two (or) more Electrochemical or Galvanic cell either in series (or) parallel (or) both, to convert chemical energy into electrical energy through a redox reactions.

24. Classification of Batteries Batteries are classified as follows 1. Primary batteries (non-rechargeable batteries) 2. Secondary batteries (rechargeable batteries) 3. Flow battery or fuel cell

25. Batteries History with Timeline

26. Classification of Batteries : Cells can be divided into two major classes: Primary batteries and Secondary batteries . Primary batteries are not rechargeable and must be replaced once the reactants are depleted. Examples of primary cells include: Carbon- Zinc (Leclanche or Dry cell), Alkaline- Manganese, Mercury- Zinc, Silver- Zinc, and Lithium cells (Lithium- Manganese dioxide, Lithium- Sulfur dioxide, and Lithium- Thionyl chloride).

27. Secondary Batteries Secondary batteries are rechargeable and require a DC charging source to restore reactants to their fully charged state. Examples of secondary cells include: Lead- Lead dioxide (Lead- acid), Nickel- Cadmium, Nickel- Iron, Nickel- Hydrogen, Nickel- Metal hydride, Silver- Zinc, Silver- Cadmium, and Lithium- Ion batteries.

28. Lead Acid Battery It is invented in 1859 by French physicist Gaston Plante , are the oldest type of rechargeable battery. The electrolyte in this cell is an acid and electrodes made up of with lead, that’s why it is lead acid storage cell. Components of lead acid batteries: Anode or Negative Electrode : Pb grid with spongy Pb Cathode or Positive Electrode : Pb -Sb coated with PbO 2 Electrolyte : 28-30% H 2 SO 4 Separator : Nylon cloth Container : Polypropylene Open-circuit voltage : 2.10 V Cell representation: Pb /PbSO 4 (s), H 2 SO 4 ||PbSO 4 (s), PbO 2 (s)/ Pb

29. (Polyethylene) Rectangular Ebonite or Polymeric case Lead Acid Battery

30. Lead Acid Battery

31. Lead Acid Battery

32. Advantages: It is made easily. It produces very high current. The self- discharging rate is low. Low maintenance, long life cycle. 97% of the lead is recycled and reused. Disadvantages: Lead causes environmental hazards. Cannot be stored in a discharged condition - sulfation process Mechanical strain and normal bumping reduces battery life.

33. French scientist, Dry Battery (Leclanché Cell) : First described by the Georges Leclanché, in 1868 . In which Zn(0) is oxidized to Zn(II) at the anode, Mn(IV) is reduced to Mn(III) at the cathode, and the resulting net cell potential is roughly 1.4 V. Primary Cell (Dry Cell)

34. Primary Cell (Dry Cell) Starting around the last decade of the 19th called “wet century, the so- form” of the transformed into the Leclanché cell was gradually modern “dry cell”, which is the version most commonly used today . The dry cell or dry battery does not contain any liquid electrolyte in it.

35. Primary Cell (Dry Cell) Polymer Case (leak proof) Brass Cap

36. Primary Cell (Dry Cell) Construction of Dry Cell: The Zinc vessel serves as anode. The cathode is a graphite rod in the center of the cell. Graphite rod is surrounded by the electrolyte which consists of a paste of NH 4 Cl, ZnCl 2 , MnO 2 and traces of acetylene black, graphite powder . Starch is added to make the mixture like a thick paste.

37. Primary Cell (Dry Cell) Cell representation of Dry Cell: Cell reactions involved: +

38. Primary Cell (Dry Cell) Cell reactions involved: The liberated NH 3 , reacts with Zn 2+ ions to form a complex [Zn(NH 3 ) 2 Cl 2 ] The cell is a primary cell and gives a voltage of 1.5 V. Gradually, the voltage drops due to anodic and cathodic reactions. Advantages: These cells are used in transistors, tape recorders, toys, portable electronic devices.

39. Primary Cell (Dry Cell) Disadvantages: These cells do not have long life because the acidic NH 4 Cl corrodes the container even when the cell is not in use. When the current is rapidly drawn from the cell, voltage drop takes place due to the building up of products at the electrodes.