Engg. Chemistry Chapter I-IV AU Notes 2024-25 PDF

ENGINEERING CHEMISTRY NAME: ROLL NUMBER: BRANCH & SECTION: 1  CHAPTER-I MOLECULAR STRUCTURE AND SPECTROSCOPY Syllabus: Molecular structure and Spectroscopy (9L): Introduction, Concept of atomic and molecular orbita...

ENGINEERING CHEMISTRY NAME: ROLL NUMBER: BRANCH & SECTION: 1  CHAPTER-I MOLECULAR STRUCTURE AND SPECTROSCOPY Syllabus: Molecular structure and Spectroscopy (9L): Introduction, Concept of atomic and molecular orbitals, LCAO, Molecular orbitals of diatomic molecules, Molecular orbital energy level diagrams of diatomic molecules (N2 and O2). Pi-molecular orbitals of butadiene. Spectroscopic techniques: Principles of spectroscopy, selection rules and applications of electronic spectroscopy (UV & Visible). Vibrational and rotational spectroscopy (IR spectroscopy) - Applications. Atom: A smallest indivisible particle of an element/matter/substance is called atom. In an atom, subatomic particles like protons and neutrons are present inside the nucleus and negatively charged electrons are present outside the nucleus. Orbit: A fixed path in which electrons revolve around the nucleus of an atom is called orbit. Orbital: A region having the highest probability of finding electrons (outside the nucleus) in an orbit called orbital. Atomic orbital: A mathematical function describing the location and wave-like behavior of an electron in an atom called atomic orbital. Atomic orbitals are labelled as s, p, d and f orbitals. All four orbitals have different shapes and different energies. The energy of atomic orbitals is in the order of s 35 g/L of dissolved salts The following techniques are carried out for desalination of Brackish water: 1) Reverse Osmosis, 2) Electrodialysis & 3) Distillation. 34 ***1) REVERSE OSMOSIS**** Reverse osmosis (RO) is a commonly used in water purification, where it removes impurities like salts, bacteria, and other contaminants from water, producing clean drinking water. Osmosis: When two solutions of different concentrations are separated by a semi permeable membrane (SPM), solvent flows from a low concentration region to a high concentration region is called as Osmosis. This process is a natural process. This natural process aims to equalize the concentrations on both sides of the membrane. Reverse Osmosis: The osmosis can be reversed by applying a pressure higher than the osmatic pressure on the high concentration side. Then, the solvent flows from higher concentration solution to lower concentration solution, which is called reverse osmosis. Osmatic pressure: The external pressure which is required to stop the flow of solvent from low concentration region to high concentration region is called osmatic pressure. Semi permeable membranes having the pores in the range of (0.0001 – 0.001 mm in diameter) and these are made by cellulose triacetate or cellulose butyrate or polyamide. Osmosis Reverse Osmosis Applied Pressure Pure Water Semi Permeable Semi Permeable Membrane (SPM) Membrane (SPM) Contam Contam inants inants Direction of Water Flow Direction of Water Flow Advantages of RO technique are as follows 1. It can remove ionic and non-ionic, inorganic, colloidal and high molecular weight organic impurities. 2. By removing chlorine, chemicals, and organic compounds, RO water typically tastes and smells better than untreated tap water. 3. RO can remove heavy metals like lead, arsenic, and mercury, as well as nitrates, fluoride, and other potentially harmful substances, contributing to better health. 35 4. The lifetime of membrane is quite high (>1 year). 5. The maintenance cost is only the replacement of membrane. 6. It removes particles with size up to 0.1 nm. Disadvantages a) Membranes have limited pH tolerance. They degrade at temperatures greater than 35 oC. b) They are vulnerable to bacteria. c) They have poor tolerance for free chlorine. TREATMENT OF BOILER FEED WATER OR SOFTENING OF WATER: There are two types of treatments are available for boiler feed water. 1) Internal treatment. 2) External treatment. 1) Internal treatment: The principle involved in internal treatment is the conversion of scale forming salts into sludges. Internal treatment methods are generally followed by blow-down operation. So that an accumulated sludge is removed. Important internal condition methods are: a) Colloidal conditioning: In low pressure boilers scale formation can be avoided by adding organic substances like kerosene, tannin, agar-agar (a gel) etc, which get coated over the forming precipitates, thereby yielding non-sticky and loose deposits which can be easily removed by blow- down operators. b) Phosphate conditioning: In high pressure boilers, scale formation can be avoided by adding sodium phosphate, which reacts with calcium and magnesium salts forming non-adherent and easily removable soft sludge of calcium and magnesium phosphates, which can be removed by blow-down operation. 3CaCl2 + 2Na3PO4 → Ca3(PO4)2 + 6NaCl The main phosphates employed are (a) NaH2PO4, sodium dihydrogen phosphate (acidic); (b) Na2HPO4, disodium hydrogen phosphate (weakly basic); (c) Na3PO4, trisodium phosphate (basic). ***c) Calgon conditioning: The word Calgon means Calcium gone, i.e., the removal of Ca+2. Sodium hexameta phosphate {Na2[Na4(PO3)6]} is Calgon. When we add Calgon to boiler water, it prevents the scale and sludge formation by forming soluble complex compound with CaSO4. Na2[Na4(PO3)6] + 2CaSO4 → Na2[Ca2(PO3)6] + 2Na2SO4 (Calgon) (Soluble complex ion) 36 2. EXTERNAL TREATMENT OF BOILER FEED WATER Treating the water before feeding it into boiler is called external treatment. ****Q: Explain the external treatment of water by ion exchange method and what are its merits and demerits? i) ION EXCHANGE PROCESS (DEIONIZATION OR DE-MINERALIZATION) This process removes almost all the ions present in water. Ion exchangers are resins (Polymers). Based on the functional groups, resins are two types: (These functional groups are responsible for exchange of ions. Ions with a positive charge are called cations. Ions with a negative charge are called anions. Diagram of Ion Exchange Method (i) Cation exchange resins: The resins which are capable of exchanging cations present in hard water with H+ ions are called as cation exchange resins. Cation exchange resins contain acidic functional groups like sulphonic acids (-SO3H) or carboxylic acids (-COOH). Example: Sulphonated polystyrene. In this step, cations like Na+, K+, Ca+2, Mg+2 present in hard water are exchanged with H+ ions of resin. R-H + M+ → R-M + H+ (Resin) (Cations) (Resin-Metal) (ii) Anion exchange resin: The resins which are capable of exchanging anions present in hard water with OH- ions is called as anion exchange resin. In this step, anions like Cl-, SO4-2, HCO3- present in hard water are exchanged with OH- of the resins. Anion exchange resins contain basic functional groups like quaternary ammonium salts etc. Example: Polystyrene with quaternary ammonium salts. 37  R-OH + Aˉ → R-A + OHˉ (Resin) (Anions) (Resin-Metal) Step-I Step-II Water passed through cation chamber Na+, K+, Step-I water passed through anion chamber, Ca+2, Mg+2, etc. are exchanged with H+ ions of then Cl-, SO4-2, HCO3- etc. ions are exchanged resin. with OH- of the resin. R-H + Na+ → R-Na + H+ R-OH + Cl- → R-Cl + OH- 2R-H + Mg+2 → R2-Mg + 2H+ 2R-OH + SO4-2 → R2SO4 + 2OH- 2R-H + Ca+2 → R2-Ca + 2H+ Free from almost all the ions and it is neutral Water is acidic in nature. (deionized or demineralized water) REGENERATION OF RESINS: (i) The exhausted cation exchange resin is regenerated by passing dilute HCl solution through it. R-Na + HCl → R-H + NaCl (Cation exchange Resin) R2-Ca + 2HCl → 2R-H + CaCl2 (Cation exchange Resin) (ii) The exhausted anion exchange resin is regenerated by passing dilute NaOH solution through it. R-Cl + NaOH → R-OH + NaCl (Anion exchange Resin) Advantages 1) The water hardness can be reduced up to 2 ppm, hence it is suitable for use in high pressure boilers. 2) Highly acidic or highly basic water can be softened by using this process. Limitations 1) The resins used in the process are quite expensive. 2) If water contains turbidity, the efficiency of the process will be reduced. 3) Water containing Fe and Mn cannot be treated because they form stable product with the resins. ---------------------------------------------------------------------------- Yours: Dr. Mallesham Godumala (Anurag University) 38  CHAPTER-III ELECTROCHEMISTY and BATTERIES SYLLABUS: Electrode, electrode potential, galvanic cell, cell reactions and cell notation, cell EMF, Nernst equation, Numerical problems, types of electrodes (Normal Hydrogen Electrode and Saturated Calomel Electrode), Determination of pH. Electrochemistry is an important branch of chemistry. It deals with the relation between electrical energy and chemical energy. It is the study of “conversion of electrical energy into chemical energy and chemical energy into electrical energy”. Electrode: A metal rod dipped in its own salt solution is called an electrode. If the metal is losing electrons (oxidation), it's referred to as the anode. On the other hand, if the metal ions in the solution that are gaining electrons (reduction), it's known as the cathode. Both anode and cathode play a crucial role in various electrochemical processes. Examples for metal electrode: Zinc electrode: Zn/ZnSO4 , Copper electrode: Cu/CuSO4 Magnesium electrode: Mg/MgSO4 , Silver electrode: Ag/AgNO3 n Electrode Cu Electrode Ag Electrode n - - - Cu - - - - - - - Ag - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - Zn/ZnSO4 Cu/CuSO4 Ag/AgNO3 Electrode potential (Ecell): The tendency of metals to either lose electrons (undergo oxidation) or gain electrons (undergo reduction) when placed in their own salt solution called electrode potential. Means, when a metal [M] placed into its own salt solution containing its own ions [M n+], then the metal may undergo either oxidation or reduction. If metal undergoes oxidation, then the positive metal ions will pass into the solution M → Mn+ + ne- (Oxidation) If the metal undergoes reduction, then the negative ions may get deposited over the metal. Mn+ + ne- → M (Reduction) 39 Standard electrode potential (𝑬𝟎𝒄𝒆𝒍𝒍 ): The tendency of a metal to either lose electrons (oxidation) or gain electrons (reduction) when placed in its own salt solution with a concentration of 1.0 M (unimolar) at 25°C temperature is called as standard electrode potential. Standard electrode potential is a fundamental concept in electrochemistry and is crucial for understanding the behavior of metals in electrochemical reactions. Cell (Galvanic Cell or Electrochemical cell): Cell is made up of two electrodes, i.e., anode half-cell and cathode half-cell. Oxidation reaction (lose of electrons) takes place at one electrode called as anode half-cell and reduction reaction (gain of electrons) takes place at another electrode called as cathode half-cell. Cell = Anode half-cell + Cathode half-cell. Cell reactions: In an electrochemical cell, flow of electrons takes place from one electrode to another electrode due to the oxidation reaction at anode and reduction reaction at cathode. So, the net chemical change can be obtained by adding the two half-cell reactions called as cell reaction. Cell reaction = Anode half-cell reaction + Cathode half-cell reaction Example, Galvanic cell [Zn/ZnSO4//CuSO4/Cu] Anode half-cell reaction: Zn → Zn+2 + 2e- (Oxidation) Cathode half-cell reaction: Cu+2 + 2e- → Cu (Reduction) ……………………………………………………………………………………… Cell reaction or net reaction: Zn + Cu+2 → Zn+2 + Cu ……………………………………………………………………………………… Electromotive Force (EMF) or Cell Potential: In an electrochemical cell, a potential difference which is required to flow of electricity from anode to cathode is called as electromotive force (emf) or cell potential. Representation and notation of electrochemical cells The Galvanic cell, or electrochemical cells, is represented on paper with the help of symbolic representation of single electrodes. Example: A Galvanic cell is represented as follows: Zn/ZnSO4//CuSO4/Cu Steps for construction the Cell: (i) Single vertical line (/): It represents a phase boundary between the metal and its salt solution. Zn/ZnSO4 (Anode half-cell), and CuSO4/Cu (Cathode half-cell). (ii) The Anode half-cell is written on the left side, and cathode half-cell is written on right side. 40 (iii) Anode half-cell and cathode half cells are separated by two vertical lines (//) called salt bridge. Example: Zn/ZnSO4//CuSO4/Cu (iv) Symbolic formation of the electrode, like the Platinum electrode, is often enclosed in brackets. Mg/Mg+2//H+/H2 (Pt) (v) The emf value of a cell is written on the right side of the cell diagram. Zn/ZnSO4//CuSO4/Cu, E = 1.12 eV NERNST EQUATION: This equation gives the relationship between electrode potential, concentration of the products and concentration of the reactants, and absolute temperature. Let us consider a cell reaction aA + bB → cC + dD Where, a, b, c, d represent the number of moles of A, B, C and D components, respectively. The Nernst equation for the cell is C c D d 2.303 RT Ecell = Eo cell log nF A a B b 0.0591 C c D d log Ecell = Eocell a b n A B Where, R = gas constant = 8.314 J K-1mol-1, T = 298 K, F = 96500 coulombs (Faraday), n = no of electrons involved in the cell reaction, 𝑜 𝐸𝑐𝑒𝑙𝑙 = Standard electrode potential & Ecell = Cell potential (EMF) Nernst equation is used to determine single electrode potential and also cell potential. Electrochemical cell: Cell which converts electrical energy into chemical energy or chemical energy into electrical energy called electrochemical cell. Galvanic cell/Voltaic Cell/Daniel Cell: A device which converts chemical energy into electrical energy due to spontaneous redox reaction. 41  Zn/ZnSO4//CuSO4/Cu n n+2 + 2e Cu+2 + 2e- Cu Salt bridge: It is an inverted U-shaped glass tube, which contains a paste of KCl or KNO3 or NH4NO3 with Agar-agar. Use of salt bridge: Salt bridge completes the circuit and allows the passage of electric current. It also maintains the electric neutrality in both half cells. Types of electrodes: 1) Working electrode, 2) Reference electrode & 3) Counter electrode Reference electrode: The electrode which is used to determine the electrode potential of another electrode is called a reference electrode. The most commonly used reference electrodes are: 1) Standard Hydrogen Electrode (SHE) or Normal Hydrogen electrode (NHE): This is one of the best reference electrodes because its electrode potential at all temperatures is considered to be 'zero'. 2) Standard Calomel Electrode (SCE) 1) Standard Hydrogen Electrode (SHE) or Normal Hydrogen Electrode (NHE): (a) Introduction: The Standard Hydrogen Electrode (SHE) is a primary reference electrode that is employed to determine both the electrode potential and the pH of an unknown sample solution. (b) Construction: The Standard Hydrogen Electrode (SHE) comprises platinum electrode coated with platinum black, which is immersed in a 1.0 M acidic solution containing H+ ions, and is maintained at 25°C. A stream of pure hydrogen is bubbled around the platinum foil at a constant pressure of one atmosphere (1 atm). This process causes the gas at the platinum electrode to form H + ions and electrons. 42  Pt wire H2 gas (1atm) Glass envelop Pt electrode coated with Pt black Acidic solution (H+, 1.0M) (c) Cell Representaion: (Pt), H2 (g)/H+ solution (1.0 M) SHE acts as both anode and cathode. If it acts as anode H2 ⇌ 2H+ + 2e- If it acts as cathode 2H+ + 2e- ⇌ H2 (d) Determination of pH of any unknown sample solution using standard hydrogen electrode The hydrogen electrode can be employed to determine the pH value of an unknown solution. In this process, the solution of unknown pH is taken in a vessel, and an electrode is half-dipped in it. The half-cell formed is connected to a standard hydrogen electrode (which contains a solution of 1N HCl) through a salt bridge. The electromotive force (emf) of the cell is determined using a potentiometer. Since the emf of the reference electrode is zero, the observed emf directly represents the emf of the half-cell containing the solution under test. Potentiometer Glass envelop H2 gas (1atm) H2 gas (1atm) Salt bridge ---------------------------- -- ------------------------- Pt electrode coated ---------------------------- with Pt black ---------------------------- ---------------------------- ---------------------------- Hydrogen electrode (HE) Acidic solution Acidic solution Standard Hydrogen electrode (Anode) (H+, pH to be determined) (H+, 1.0M) (SHE) (Cathode) 43 The cell represented as; Anode (Pt), H2 (g)/H+ (pH to be determined)//Acidic solution (H+ = 1M)/H2 (g), (Pt) (Cathode) Ecell = Eanode + Ecathode H2 ⇌ 2H+ + 2e- Ecell = EHE + ESHE 0.0591 [P] Ecell = E°cell - log = (0.0591 pH) + 0 n [R] = 0.0591pH 0.0591 [H+] 2 𝐸𝑐𝑒𝑙𝑙 Ecell E°cell log pH = = - 0.0591 2 [H2] 0.0591 [H+] 2 Ecell = 0 - log 2 1 0.0591 Ecell = - 2 x log [H+] 2 Ecell = 0.0591pH (B) STANDARD CALOMEL ELECTRODE (SCE) (a) Introduction: Standard calomel electrode (SCE) is a secondary reference electrode used to determine the electrode potential and pH of any unknown solution. (b) Construction: Pt wire Salt bridge Saturated KCl solution Hg2Cl2 and Hg paste Mercury Calomel Electrode (c) Cell Representaion: Standard calomel electrode representation: (Pt), Hg/Hg2Cl2, Saturated KCl Predominantly it acts as cathode. Hg2Cl2 + 2e− → 2Hg + 2Cl− Hg2Cl2 → 2Hg+ + 2Cl− 2Hg+ + 2e− → 2Hg Hg2Cl2 → 2Hg + 2Cl− 44 Electrode potential of calomel electrode depends on the concentration of electrolyte. If saturated KCl is used as an electrolyte, the electrode potential is 0.241 V. If 1.0N KCl is used as an electrolyte, the electrode potential is 0.280 V. If 0.1N KCl is used as an electrolyte, the electrode potential is 0.333 V. (d) Determination of pH of unknown solution using standard calomel electrode: To determine the pH of an unknown solution, a Hydrogen Electrode (HE) containing solution of unknown pH is coupled with standard calomel electrode. Potentiometer Glass envelop H2 gas (1atm) Salt bridge Pt electrode coated Saturated with Pt black Cl solution Hg+Hg2Cl2 paste Pure Hg Hydrogen electrode (HE) Acidic solution Standard Calomel electrode (Anode) (H+, pH to be determined) (SCE) (Cathode) The above cell can be represented as (Pt), H2(g)/H+ (pH to be determined) // Sat.KCl/ Hg2Cl2, Hg (Anode) (Cathode) 45  Ecell = Ecathode + EAnode H2 ⇌ 2H+ + 2e- 0.0591 [P] Ecell = ESCE + EHE Ecell = E°cell - log n [R] Ecell = 0.241 + 0.0591pH 0.0591pH = Ecell – 0.241 0.0591 [H+] 2 Ecell = E°cell - log pH = Ecell – 0.241 2 [H2] 0.0591 0.0591 [H+] 2 Ecell = 0 - log 2 1 0.0591 Ecell = - 2 x log [H+] 2 Ecell = 0.0591pH Problems based on Nernst equation: 1) Calculate the single electrode potential of the following electrochemical cell reaction. Cu+2/Cu where E0Cu+2/Cu = 0.36V (0.01M) Ans: Net reaction: Cu+2 + 2e- → Cu (Reduction) 0 0.0591 [𝑃] Nernt Equation: Ecell = Ecell - log [𝑅] 𝑛 ECu+2/Cu = 0.0591 Cu log n Cu+2 = 0.0591 0.36 - log [0.01] 2 = 0.0591 0.36 - log 2 = 0.0591 0.36 - x 2 log 2 = 0.36 - 0.0591 ECu+2/Cu = 0.3009V 2) Calculate the single electrode potential of the following electrochemical cell reaction. Ag+/Ag where E0Ag+/Ag = 0.8V (0.01M) Ans: Net reaction: Ag+ + e- → Ag (Reduction) 0 0.0591 [𝑃] Nernt Equation: Ecell = Ecell - log [𝑅] 𝑛 EAg+/Ag = 0.0591 [Ag] log n Ag+ 46  = 0.0591 0.8 - log [0.01] 1 = 0.8 - 0.0591 log = 0.8 - 0.0591 x 2 log = 0.8 - 0.1182 EAg+/Ag = 0.6818 V 3) Calculate the single electrode potential of the following electrochemical cell reaction. (Pt) Cl2/Cl- where E0Cl2/Cl- = 1.36V (0.01M) Ans: Net reaction: Cl2 + 2e- → 2Cl- (Reduction) 0 0.0591 [𝑃] Nernt Equation: Ecell = Ecell - log [𝑅] 𝑛 ECl2/Cl- = 0.0591 [Cl-]2 - log n Cl2 = 0.0591 [0.01]2 1.36 log 2 1 = 0.0591 1.36 - log [10-2]2 2 = 0.0591 1.36 - x (-4) log 2 = 1.36 + 0.1182 ECl2/Cl- = 1.4782V 4) Calculate the single electrode potential of the following electrochemical cell reaction. Zn/Zn+2 where E0Zn+2/Zn = -0.76V (0.1M) Ans: Net reaction: Zn → Zn+2 + 2e- (Oxidation) 0 0.0591 [𝑃] Nernt Equation: Ecell = 𝐸𝑐𝑒𝑙𝑙 - log [𝑅] 𝑛 Ecell = 0 0.0591 [𝑃] 𝐸𝑐𝑒𝑙𝑙 - log 𝑛 [𝑅] EZn/Zn+2 = 0.0591 n+2 log n n = 0.0591 [0.1] 0.76 - log 2 = 0.76 - 0.0295 log [10-1] = 0.76 + 0.0295 47  EZn/Zn+2 = 0.7895 V 5) Calculate the EMF of cell potential of the following electrochemical cell reaction. Zn/Zn+2//Ag+/Ag where E0Zn/Zn+2 = 0.76V and E0Ag+/Ag = 0.8V (1.0M) (1.0M) 0 0 0 Ans: 𝐸𝑐𝑒𝑙𝑙 = 𝐸𝐴𝑛𝑜𝑑𝑒 + 𝐸𝐶𝑎𝑡ℎ𝑜𝑑𝑒 = E0Zn/Zn+2 + E0Ag+/Ag = 0.76 + 0.80 0 𝐸𝑐𝑒𝑙𝑙 = 1.56V Reaction at anode : Zn → Zn+2 + 2e- Reaction at cathode : 2Ag+ + 2e- → 2Ag Net cell reaction : Zn + 2Ag+ → Zn+2 + 2Ag 0 0.0591 [𝑃] Nernt Equation: Ecell = Ecell - log [𝑅] 𝑛 Ecell = 0.0591 [Zn+2] Ag 2 log n [Ag+]2 n = 0.0591 1 2 1.56 log 2 2 1 = 0.0591 1.56 - log 2 = 1.56 + 0 Ecell = 1.56V 6) Calculate the EMF of cell potential of the following electrochemical cell reaction. Fe/Fe+2//Cu+2/Cu where E0Fe+2/Fe = -0.44V and E0Cu+2/Cu = 0.34V (0.1M) (0.01M) 0 0 0 Ans: 𝐸𝑐𝑒𝑙𝑙 = 𝐸𝐴𝑛𝑜𝑑𝑒 + 𝐸𝐶𝑎𝑡ℎ𝑜𝑑𝑒 = E0Fe+2/Fe + E0Cu+2/Cu = 0.44 + 0.34 0 𝐸𝑐𝑒𝑙𝑙 = 0.78V Reaction at anode : Fe → Fe+2 + 2e- Reaction at cathode : Cu+2 + 2e- → Cu Net cell reaction : Fe + Cu+2 → Fe+2 + Cu 0 0.0591 [𝑃] Nernt Equation: Ecell = Ecell - log [𝑅] 𝑛 48  Ecell = 0.0591 Fe+2 Cu log n Fe Cu+2 = 0.0591 0.1 1 0.78 log 2 1 0.01 = 0.0591 0.78 - log 2 = 0.78 - 0.0295 Ecell = 0.7505V 7) Calculate the EMF of cell potential of the following electrochemical cell reaction. Zn/Zn+2//Ag+/Ag where E0Zn/Zn+2 = 0.76V and E0Ag+/Ag = 0.8V (0.1M) (0.01M) 0 0 0 Ans: 𝐸𝑐𝑒𝑙𝑙 = 𝐸𝐴𝑛𝑜𝑑𝑒 + 𝐸𝐶𝑎𝑡ℎ𝑜𝑑𝑒 = E0Zn/Zn+2 + E0Ag+/Ag = 0.76 + 0.80 0 𝐸𝑐𝑒𝑙𝑙 = 1.56V Reaction at anode : Zn → Zn+2 + 2e- Reaction at cathode : 2Ag+ + 2e- → 2Ag Net cell reaction : Zn + 2Ag+ → Zn+2 + 2Ag 0 0.0591 [𝑃] Nernt Equation: Ecell = Ecell - log [𝑅] 𝑛 Ecell = 0.0591 [Zn+2] Ag 2 log n [Ag+]2 n = 0.0591 [0.1] 2 1.56 log 2 0.01 2 0.0591 10 1 1.56 log 2 10 2 2 = 0.0591 1.56 - log 2 = 1.56 - 0.0295 x 3 = 1.56 - 0.0885 Ecell = 1.47V 49 ---------------------------------------------------------------------------- MID-II Syllabus ---------------------------------------------------------------------------- BATTERIES Introduction to cell and battery, Primary (lithium cell) and secondary cells, (Lithium ion cells). Fuel cells – Hydrogen – Oxygen fuel cell, advantages and engineering applications of fuel cells. Cell: A cell is a single unit of device that converts chemical energy into electrical energy or electrical energy into chemical energy. Cells are two types: Galvanic cell/Voltaic Cell/Daniel Cell: Electrochemical cells which convert chemical energy to electrical energy. Here, Anode (-Ve charge) and Cathode (+Ve charge). Electrolytic cell: Electrochemical cells which convert electrical energy to chemical energy. Here, Anode (+Ve charge) and Cathode (-Ve charge) Battery: Battery is the collection of cells that converts chemical energy into electrical energy and vice versa. Cells and batteries can store chemical energy and release it as electrical energy when required. Battery has 3 main components: Anode (-Ve charge), Cathode (+Ve charge) & Electrolyte. Characteristics of Battery a) A battery should have high capacity (very small variation of voltage during discharge). b) Battery should have high energy efficiency. c) A battery should have high cycle life means it should have more number of charging and discharging cycles before failure. d) Long shelf-life is required. e) High tolerance to different service conditions like variation in temperature, vibrations, shocks etc. f) Cost should be affordable. Types of batteries: Batteries are two types. Primary battery/cell and Secondary battery/cell. Primary cells/battery Secondary cells/battery 1. The cell in which reactions are not reversible 1. Cells in which reactions are reversible. It can called as primary cell. Once reactants are be used for many cycles by charging and converted to products, the cell becomes dead. discharging. 2. These are not rechargeable 2. These are rechargeable 50  3. These can be used only once 3. These can be used for multiple times 4. These are cheap 4. These are expensive 5. Ex: Lithium cell, Dry cells 5. Ex: Lead acid storage cells, Lithium-ion cell 1) LITHIUM CELLS (A Primary cell): In Lithium cell, Lithium (Li) acts as anode Depending on the cathode and electrolyte Lithium cells are classified into two types. a) Lithium Cell with Solid Cathode (LSC): Ex. Li-MnO2 b) Lithium cell with Liquid Cathode (LLC): Ex. Li-SO2 and Li-SOCl2 a) Lithium Cell with Solid Cathode (LSC): In these cells the cathode is solid. Li-MnO2 is the example for LSC Anode: Li Cathode: MnO2 Electrolyte: Propylene carbonate and 1,2-dihydroxyehtane. At Anode: Li Li+ + e- At Cathode: Li+ + e- + MnO2 LiMnO2 Net Cell reaction: Li + MnO2 LiMnO2 Li-MnO2 battery delivers voltage about 3V. Applications: (i) Cylindrical type of batteries is used in cameras and remotes etc. (ii) Coin type batteries are used in wrist watches and calculators. b) Lithium Cell with Liquid Cathode (LLC): In these cells the cathode is liquid. Li-SO2 and Li-SOCl2 are the examples for LLC Li-SO2: Anode: Li Cathode: SO2 Electrolyte: Propylene carbonate or acrylonitrile At Anode: Li Li+ + e- At Cathode: Li+ + e- + 2SO2 LiS2O4 Net Cell reaction: Li + 2SO2 LiS2O4 51 Li-SOCl2: Anode: Li Cathode: SOCl2 Electrolyte: SOCl2 At Anode: 4Li 4Li+ + 4e- At Cathode: 4Li+ + 4e- + 2SOCl2 4LiCl + SO2 + S Net Cell reaction: 4Li + 2SOCl2 4LiCl + SO2 + S These batteries deliver voltage up to about 3.3 – 3.5 V Applications: They are used in electronic circuit boards. They have many military and space applications. They can be used in medical devices like neuro stimulations and drug delivery system. SECONDARY CELLS Secondary cells act as both galvanic/voltaic and electrolytic cells. When it converts chemical energy into electrical energy functions as voltaic cell and while converting electrical energy into chemical energy functions as electrolytic cells. 2) LEAD-ACID STORAGE CELL: Lead storage battery is also known as the lead acid battery. It is a kind of rechargeable battery. It is a secondary storage battery. This secondary storage battery can undergo reversible chemical reaction. Lead storage batteries are the energy storage device. Anode: Lead (Pb) plates Cathode: Pb plates coated with PbO2 Electrolyte: Diluted H2SO4 (38%) Construction: Pb (Anode) Pb plate coated with PbO2 (Cathode) 2H+ SO4 2 H2SO4 (38 ) When Lead-Acid storage battery converting chemical energy into electrical energy it is said to be presenting discharging phase. 52 DISCHARGING PHASE At Anode: Pb Pb+2 + 2e- Pb+2 + SO4-2 PbSO4 Net reaction at anode: Pb + SO4-2 PbSO4 + 2e- - - - - - - - - - - - - (1) At Cathode: PbO2 + 2e- + 4H+ Pb+2 + 2H2O Pb+2 + SO4-2 PbSO4 Net reaction at Cathode: PbO2 + 2e- + 4H+ + SO4-2 PbSO4 + 2H2O - - - - - - - - - - - - (2) Overall reaction during discharging (1) + (2) Reaction at anode: Pb + SO4-2 PbSO4 + 2e- Reaction at Cathode: PbO2 + 2e- + 4H+ + SO4-2 PbSO4 + 2H2O Net Cell reaction: Pb + PbO2 + 4H+ + 2SO4-2 2PbSO4 + 2H2O + Energy Lead-acid battery delivers voltage 2V. In this case, PbSO4 completely formed at anode and cathode the battery stops functioning as voltaic cell, at this condition cell needs to be recharged. Charging can be done by sending electricity with voltage more than 2V. Charging Phase: When Lead-acid storage cell converts the electrical energy into chemical energy it is said to be in the charging phase, during charging, the chemical reactions taking place in the discharging phase will be reversed (Reverse to reactions 1 &2). Net reaction at anode: PbSO4 + 2e- Pb + SO4-2 Net reaction at Cathode: PbSO4 + 2H2O PbO2 + 2e- + 4H+ + SO4-2 Net Cell reaction: 2PbSO4 + 2H2O + Energy Pb + PbO2 + 4H+ + 2SO4-2 A car battery contains six lead-acid storage cells, resulting in a voltage delivery above 12V. Typically, for every 1°C decrease in temperature, there is a corresponding decrease in voltage of approximately 1.5 x 10-4 V. Consequently, starting a car becomes challenging in cold climates, necessitating the heating of the car battery under such conditions. 53 Applications: Lead-acid batteries are used for supplying electricity for automobiles (all kind of vehicles), electronic devices, laboratories, hospitals, railways, broad-casting stations, etc. 3) LITHIUM ION CELL Anode: LiCoO2 (LiNiO2 or LiMnO4) mixed with conductor and binder Cathode: Carbon mixed with binder. Electrolyte: LiPF6 in organic solution. Voltameter ------------------------------------ ------------------------------------ 1st Metal foil coated ------------------------------------ 2nd Metal foil with LiCoO2 + ------------------------------------ coated with binder + conductor ------------------------------------ Carbon + binder (Anode) ------------------------------------ (Cathode) ------------------------------------ Electrolyte (LiPF6) DISCHARGING PHASE: When cell converts chemical energy into electrical energy is said to be present in discharging phase. Reaction at anode: LimCoO2 Lim-nCoO2 + nLi+ + ne- Reaction at Cathode: C + nLi+ + ne- CLin Net Cell reaction: LimCoO2 + C Lim-nCoO2 + CLin Lithium-ion cell delivers voltage about 4.1V RECHARGING PHASE: Charging can be done by sending current with voltage more than 4.1V, then chemical reaction taking place in discharging phase will be reversed. Reaction at anode: Lim-nCoO2 + nLi+ + ne- LimCoO2 Reaction at Cathode: CLin C + nLi+ + ne- Net Cell reaction: Lim-nCoO2 + CLin LimCoO2 + C Applications: Lithium-ion batteries used in laptops and mobile phones. 54 Q: What is a fuel cell? Explain the working of H2-O2 fuel cell and its applications. Fuel Cell Fuel cell is an electrochemical cell, where chemical energy is converted into electrical energy by sending reactants into the battery. The basic principle of fuel cell is identical to that of an electrochemical cell. But only the difference is that chemical energy provided by oxidant and fuel are stored outside of battery/cell. Example: Hydrogen – oxygen fuel cells (H2 - O2). Characteristics of Fuel cells: Energy conversion is very high & thermal pollution is low. Hydrogen-Oxygen Fuel Cell (H2 - O2) Similar to a galvanic cell, fuel cell also has two half cells. Both half cells have porous carbon/graphite electrode with a catalyst (platinum or lead or silver). The electrodes are placed in the aqueous solution of NaOH or KOH which acts as an electrolyte. Hydrogen is supplied at anode and oxygen is supplied at cathode. H2 gas O2 gas Unreacted H2 Unreacted O2 e e Porous carbon with Porous carbon with Pt or Pd (Anode) Pt or Ag (Cathode) OH OH H2O H2O Aqueous OH solution (electrolyte) The two half-cell reactions are as follows: 2 At Anode: 2H2 + 4OH 4H2O + 4e At Cathode: O2 + 4e + 2H2O 4OH Net reaction: 2H2 + O2 2H2O (Fresh water) 55 The EMF of this cell is measured to be 1.23V. A number of such fuel cells are stacked together in series to make a battery. Applications of Hydrogen-Oxygen Fuel Cell (H2 - O2): The most important application of a fuel cell is its use in Space Vehicles, Submarine or Military Vehicles. Advantages of Hydrogen-Oxygen Fuel Cell (H2 - O2): Energy conversion is very high in Hydrogen-Oxygen fuel cell. It has low thermal pollution. Low maintenance cost. Product H2O is a drink water source for Astronauts. It is an energy storage system for space applications. Disadvantages of Hydrogen-Oxygen Fuel Cell (H2 - O2): ❖ The actual lifetime of Hydrogen-Oxygen (H2 - O2) fuel cell is not known. ❖ Initial cost of Hydrogen-Oxygen fuel cell is very high. ❖ Storage of large amount of Hydrogen is difficult and the problems of durability. Yours: Dr. Mallesham Godumala (Anurag University) 56  CHAPTER-IV CORROSION AND ITS CONTROL: (8 L) Corrosion: Introduction, types of corrosion: chemical (oxidation corrosion) and electrochemical corrosion, factors affecting the rate of corrosion: nature of the metal, galvanic series, purity of metal, nature of corrosion product, nature of environment: effect of temperature, effect of pH, humidity. Corrosion control methods: Cathodic protection: sacrificial anode method and impressed current cathode method. Protective coatings: metallic coatings (anodic and cathodic), methods of applications of metallic coatings: electroplating and electroless plating. INTRODUCTION Metals and alloys (combination of two or more metals) are used for fabrication or construction of engineering materials. If the metals or alloy-based components are not properly maintained, they slowly decay by the action of atmospheric gases, moisture and other chemicals. “The process of decay of metal by environmental attack is called as corrosion”. When metals undergo corrosion they convert into their oxides, hydroxides, carbonates, sulphides, etc. ***Examples for corrosion: (i) Rusting of iron (Fe) – when Iron (Fe) is exposed to the atmospheric conditions, a layer of Reddish-brown scale formed is called as rusting of Iron [Fe2O3.xH2O]. Reddish-brown film Environment Fe Fe (ii) Formation of a Green film on the surface of Copper (Cu) when exposed to open atmosphere containing CO2 due to the formation of basic carbonate [CuCO3 + Cu(OH)2]. Green film Environment Cu Cu Units for corrosion: mm/year, cm/year or inches/year. Reasons for corrosion 1. The metals exist in nature in the form of their “Raw materials or Ores’’ in their stable forms like oxides, hydroxides, carbonates, sulphides, chlorides and silicates. 57 2. Metals are extracted by giving some amount of energy to the Ore of the metal. So, the extracted metal has high energy and it is thermodynamically unstable. Hence, metals have natural tendency to go back to their stable Ore form (oxides, hydroxides, carbonates, sulphides, chlorides and silicates). Corrosion (oxidation) Metallic Metal compound (Or) Ore of metal (High energy, Metal extraction thermodynamically unstable) (Reduction) (Low energy, thermodynamically stable) [Corrosion is the reverse process of metal extraction] Disadvantages of corrosion 1. Life span and efficiency of metallic parts of machinery is reduced. 2. Metals loss metallic properties like electrical conductivity, flexibility and elasticity. 3. Wastage of metal in the form of its compounds. TYPES OF CORROSION (1) Chemical corrosion/Dry corrosion & (2) Electrochemical corrosion/Wet corrosion (1) CHEMICAL CORROSION (OR) DRY CORROSION Attack of atmosphere on the surface of a metal in the absence of moisture is called dry corrosion or chemical corrosion. This type of corrosion occurs mainly through the direct chemical action of atmospheric gases like O2, N2, halogens, H2S, SO2. Dry corrosion is classified into three types. (A) Oxidation corrosion, (B) Corrosion by other gases & (C) liquid metal corrosion. (a) Oxidation corrosion: Attack of oxygen (O2) on the surface of any metal in the absence of moisture called oxidation corrosion. Generally, at low-temperature alkali and alkali earth metals undergo oxidation corrosion. At high temperature the remaining all metals undergo oxidation corrosion [except silver (Ag), Gold (Au) and Platinum (Pt)]. 58  Mechanism: Formation of Metal oxide [MO] at the point of meeting Mn+ & O-2 O-2 Mn+ O-2 Mn+ O-2 Mn+ O2 Mn+ O-2 Mn+ O-2 Mn+ M M+2 + 2e- (Oxidation) 1/2O2 + 2e- O-2 (Reduction) Net Reaction: M + 1/2O2 M+2 + O-2 MO (Metal Oxide) The extent of oxidation corrosion depends on the nature of the metal oxide layer formed on the surface of the metal. Following are the types of metal oxide [MO] layers formed on metals. (a) Stable metal oxide layer: The metal oxide layer formed is stable and it prevents further corrosion. Stable MO layer Metal + O2 Metal No further corrosion Eg.: Al, Cu, Sn, Pb, etc. form stable oxide layers on surface, thus preventing further oxidation. (b) Unstable metal oxide layer: If the formed MO layer is unstable, then it decomposes back to metal and oxygen. Therefore, oxidation corrosion is not possible in such cases. Ex: Ag, Au and Pt do not undergo oxidation corrosion. Unstable MO layer Long time MO decompose Metal + O2 Metal Metal + O2 Rapidly 59 (c) Volatile metal oxide layer: The metal oxide layer formed is volatile in nature and evaporates as soon as it is formed. Example: Mo- molybdenum forms volatile MoO3 layer. Fresh surface for Volatile Metal Oxide further attack Evaporation Metal + O2 Metal Metal of MO (d) Porous metal oxide layer: If the metal oxide layer is porous, that layer has pores or cracks. In this case the atmospheric oxygen penetrates through the pores or cracks and corrode the underlying metal surface. This cause continuous corrosion till the complete conversion of metal into its oxide. Example: Alkali and alkaline earth metals (Li, Na, K, Mg etc). Porous metal oxide Further attack through Metal + O2 Metal pores and cracks continues **Pilling Bedworth rule** “The extent of protection given by the metal oxide layer to the parent metal” is Pilling Bedworth rule. It is expressed in terms of Specific Volume Ratio. Volume of Metal Oxide Layer Specific Volume Ratio = Volume of Parent Metal Specific volume ratio increases the further oxidation corrosion decreases. Specific volume ratio decreases the further oxidation corrosion increases. Example. The specific volume ratio of W, Cr, and Ni are 3.6, 2.0 and 1.6, respectively. Ni has lower specific volume ratio, so it undergoes rapid oxidation corrosion. (B) Corrosion by other gases Gases like CO2, SO2, Cl2, F2 and H2S also cause corrosion of metals in absence of moisture. The rate of corrosion depends on the reactivity of gases on metal surface. Examples: (i) Attack of Cl2 on Ag forms stable AgCl layer. Stable ‘AgCl’ that Ag + Cl2 Ag prevents corrosion (ii) Attack of Cl2 on Sn forms volatile SnCl4 layer. 60  Volatile ‘SnCl4’ Sn + Cl2 Sn [Fresh metal for further attack] (iii) Attack of H2S on steel forms porous ‘FeS’ layer. Porous ‘FeS’ layer Steel + H2S Steel [Further corrosion takes place] (c) Liquid Metal Corrosion “Attack of liquid metal on the surface of any solid metal in absence of moisture”. It is due to inter penetration of liquid metal through the solid metal (or) dissolution of solid metal in liquid metal. Example: Attack of Mercury (Hg) on any metal. 2) ELECTROCHEMICAL CORROSION OR WET CORROSION***** Attack of atmosphere on the surface of a metal in presence of moisture is called as Electrochemical Corrosion or Wet Corrosion. In this, formation of electrochemical cells takes place. Electrochemical cell involves, i) Anode and cathodes, which are separated by conducting medium (electrolyte) ii) At anode, metal undergoes oxidation to form metallic ions M → Mn+ + ne- iii) At cathode, H+ ions/H2O/O2 takes electrons and forms non-metallic ions like OH- or O-2 iv) Metallic ions formed at anode and non-metallic ions formed at cathode combines each other and forms corrosion product (like Rust) between anode and cathode. Wet corrosion takes place in two ways a) Evolution of Hydrogen (H2) & b) Absorption of Oxygen (O2) a) Evolution of Hydrogen (H2): This type of wet corrosion mainly occurs in “acidic medium” and also possible in neutral medium (H2O). In hydrogen evolution type wet-corrosion, anodic areas are large and cathodic areas are small. In acidic medium: 61  ------------------------------- Acidic medium - - - - -H+ - - - - - - - - - - - -- - - - -H+ - - - - - - (Electrolyte) - - - - - - -H+ - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - -+ - - - -- - - - - - - - - - -H+ - - Small cathodic area 2 H + 2e H2 Fe Fe 2 2e 2 Fe Fe 2e Large Anode Area Large Anode Area e- e- Fe Metal At Anode: Fe Fe+2 + 2e- At Cathode: 2H+ + 2e- H2 Net reaction: Fe + 2H+ Fe+2 + H2 In neutral medium (H2O): At Anode: Fe Fe+2 + 2e- At Cathode: 2H2O + 2e- 2OH- + H2 Net reaction: Fe + 2H2O Fe+2 + 2OH- + H2 Fe(OH)2 4Fe(OH)2 + O2 + 2 H2O → 4Fe(OH)3 → 2[Fe2O3.3H2O] [Rust] (b) Absorption of Oxygen (O2): This type of corrosion takes place in basic or neutral medium in presence of oxygen. Usually, in presence of oxygen the surface of iron is coated with a thin film of iron oxide. If the film develops cracks, anodic areas are created on the surface and the rest of the metal surface acts as cathodes. In absorption of oxygen type corrosion, anodic areas are small and cathodic areas are large. 62  Neutral aqueous solution of electrolyte (NaCl) O2 + H2O + 2 e 2OH O2 + H2O + 2 e 2OH Rust Metal oxide Large Cathodic Area Large Cathodic Area e Small Anodic Area Fe Fe+2 + 2e Fe Metal At Anode: Fe Fe+2 + 2e At Cathode: O2 + H2O + 2e 2OH Net reaction: Fe + O2 + H2O Fe+2 + 2OH Fe(OH)2 If there is enough oxygen is present, Ferrous hydroxides [Fe(OH)2] will be converted to Ferric hydroxides [Fe(OH)3] followed by rust. (𝑜𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛) (𝑜𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛) 4Fe(OH)2 + O2 + 2 H2O → 4Fe(OH)3 → 2[Fe2O3.3H2O] [Rust] Differences between Dry corrosion and wet corrosion Dry corrosion Wet corrosion 1. The process of attacking of atmosphere 1. The process of attacking of atmosphere on metal surface in absence of moisture on metal surface in presence of moisture. 2. It is a slow process. 2. It is a rapid process 3. It is a uniform corrosion. 3. It is non-uniform corrosion. Metal Metal 4. It is explained by absorption mechanism 4. Explained by electrochemical mechanism 5. In this corrosion, product is formed at the 5. In this process corrosion product takes place where corrosion is taking place. place in between anode and cathode. 63 Q: Discuss the various factors effecting the rate of corrosion. ***FACTORS EFFECTING THE RATE OF CORROSION*** Always corrosion is the relation between METAL and ENVIRONMENT. Therefore, the rate of corrosion depends on the nature of metal and nature of environment. 1. Nature of Metal 2. Nature of Environment (a) Position of metal in the Galvanic Series (a) Temperature (b) Purity of Metal (b) Humidity (c) Nature of surface film (c) pH (d) Nature of corrosion product 1. NATURE OF METAL (a) Position of metal in the Galvanic Series: Metals present at the top of the Galvanic Series undergo rapid corrosion due to high oxidation potential. Metals present at bottom of the Galvanic Series undergo less corrosion due to low oxidation potentials. Mg, Mg-alloy, Zinc, Al, Fe, Steel, Pb-Sn alloy, Pb, Sn, Brass, Cu, Bronze, Ag, Graphite, Ti, Au, Pt Decreasing rate of corrision → (b) Purity of Metal: Rate of corrosion is inversely proportional to purity of the metal. Means as purity of metal decreases the rate of corrosion increases, and as purity of metal increases the rate of corrosion decreases. 𝟏 Rate of corrosion ∝ % 𝐎𝐟 𝐩𝐮𝐫𝐢𝐭𝐲 𝐨𝐟 𝐌𝐞𝐭𝐚𝐥 Impurities present in metal can lead to heterogeneity, resulting in the formation of an electrochemical cell. Consequently, this accelerates the rate of corrosion. Example: Zinc metal. % Of purity 99.999 99.99 99.9 99.0 Rate of Corrosion 1 2650 5000 7200 (c) Nature of surface film: The relation between nature of surface film and the rate of corrosion can be explained by Pilling-Bedworth rule. 𝐕𝐨𝐥𝐮𝐦𝐞 𝐨𝐟 𝐌𝐞𝐭𝐚𝐥 𝐎𝐱𝐢𝐝𝐞 𝐋𝐚𝐲𝐞𝐫 Specific Volume Ratio (SVR) = 𝐕𝐨𝐥𝐮𝐦𝐞 𝐨𝐟 𝐏𝐚𝐫𝐞𝐧𝐭 𝐌𝐞𝐭𝐚𝐥 64  If SVR increases, the rate of corrosion decreases. (Non porous, protective surface film Metal If SVR decreases, the rate of corrosion increases. (Porous, non- Metal protective surface film) (d) Nature of corrosion product (i) Volatility of corrosion product: If the corrosion product is volatile, it will evaporate as soon as it is forms, exposing a fresh metal surface to further corrosion. Fresh surface for Volatile Metal Oxide further attack Evaporation Molybdenum O2 Mo Molybdenum of MO (ii) Solubility of corrosion product: If corrosion product is soluble in the corroding medium, then the rate of corrosion is High. If corrosion product is insoluble in the corroding medium, then the rate of corrosion is Low. H2SO4 corroding medium ------------ ------------ ------------ ------------ H2SO4 ------------ ------------ ------------ PbSO4 Pb Pb Corrosion product PbSO4 is insoluble in H2SO4, so rate of corrosion of Pb is low. 2. NATURE OF ENVIRONMENT (a) Temperature: The rate of corrosion is directly proportional to the temperature. Means, the rate of corrosion increases with increase in temperature and vice versa. Rate of corrosion ∝ Temperature Reasons: As temperature increases, the rate of reaction increases and diffusion of gases on metal surface increases. (b) Humidity in air: The rate of corrosion is directly proportional to the humidity. The rate of corrosion increases with increase in humidity/moisture present in air and vice versa. Rate of corrosion ∝ Humidity of air. 65 Reasons: (i) Humidity present in atmosphere provides water to the electrolyte which is essential for setting up of an electrochemical cell. (ii) The metal oxide film formed on metal surface has the tendency to absorb moisture which creates another electrochemical cell. Example: (i) Fe undergoes rapid corrosion in wet conditions than in dry conditions. ***Critical Humidity: The humidity above which the rate of corrosion increases rapidly is called as critical humidity. (c) pH: Rate of corrosion is inversely proportional to pH. Means, in acidic medium metals undergo rapid corrosion compared to neutral/basic medium. 𝟏 Rate of corrosion ∝ 𝐩𝐇 Acidic medium: PH < 7 - Corrosion is More Basic medium: PH > 7 - Corrosion is Less Fe undergoes rapid corrosion in H2CO3 (acidic) than H2O and NaOH (basic medium). ---------------------------------------------------------------------------- CORROSION CONTROLLING METHODS Corrosion controlling methods are two types: (i) Cathodic protection & (ii) Metallic coatings. (i) CATHODIC PROTECTION The method of protecting the base metal by forcibly making it to behave like a cathode there by preventing corrosion is called as Cathodic Protection. It is classified into two types. (a) Sacrificial Anodic Protection & (b) Impressed Current Cathodic Protection. (A) SACRIFICIAL ANODIC PROTECTION In this method, the metal to be protected (base metal) is connected to a more anodic metal (having high oxidation potential metal) through a wire, then the newly introduced metal acts as anode and undergoes corrosion, whereas the base metal acts as cathode i.e., that will be protected from corrosion. The attached more anodic metal is called Sacrificial Anode. Generally, sacrificial anodes are Mg, Al and Zn. 66  Base Metal Mg/Zn/Al (Sacrificial Metal Anode) APPLICATIONS 1) Marine structures and Ship hulls can be protected from corrosion. ‘’‘’ ‘’‘’ ‘’‘’ ‘’‘’ ‘’‘’ ‘’‘’ ‘’‘’ ‘’‘’ ‘’‘’ ‘’‘’ ‘’‘’ ‘’‘’ Base Metal Sacrificial Anode ( n) Water 2) Underground pipelines can be protected Ground level Mg Underground pipeline Sacrificial Metal Anode 3) Industrial Hot Water tank can be protected. ------------ ------------ Sacrificial Anode Base Metal ------------ (Mg/Zn/Al) ------------ ------------ (B) IMPRESSED CURRENT CATHODIC PROTECTION 67 In this method, an impressed current is allowed through the circuit from an external DC source in the opposite direction to nullify the corrosion current. As a result, the anode metal is converted into cathode and protected from corrosion. DC line Ground level Graphite -------- Underground pipeline Base metal. Back fill Initially acts as Anode but after Composed of Gypsum passing current acts as cathode Applications 1. Underground pipelines can be prevented from corrosion 2. Marine structure can be prevented from corrosion 3. Hot water tanks can be prevented from corrosion 4. Condensers can be prevented from corrosion 5. Transmission line towers can be prevented from corrosion Limitations i) Capital investment cost is very high. ii) Maintenance cost is high iii) Regular inspection is needed (II) METALLIC COATINGS The process of coating of a pure metal (coat metal) on the surface of the base metal which is to be protected (base metal) is called metallic coating. Metal II (Coating Metal) Base Metal Based on coating, metallic coatings are classified into two types. i) Anodic coating and ii) Cathodic coating. A) ANODIC COATING: The process of coating more anodic metal (high oxidation potential metal) on the surface of the base metal is called anodic coating. Ex: Coating of Mg/Al/Zn on any base metal like Iron (Fe) or Steel. 68  More Anodic Metal ( n) Base Metal Anodic coating prevents corrosion of base metal due to its high oxidation potential, i.e., anodic coating undergoes corrosion while base metal protected from corrosion. Corrosion Environment Anodic Coating Exposed part of base Anodic Coating (Anode) Metal (Cathode) (Anode) e e Base Metal If anodic coating contains any imperfections like pores, breaks or discontinues, those areas can act as cathodic parts (cathode). Hence, no corrosion takes place to the base metal until entire anodic coating gets corrosion. B) CATHODIC COATING: The process of coating of more cathodic metal (low oxidation potential metal) on the surface of the base metal is called cathodic coating. Coating of Tin (Sn) on Iron (Fe) is known as Tinning. More Cathodic Metal (Sn) Base Metal Cathodic coating prevents corrosion of base metal due to its noble behavior towards corrosion, i.e., less tendency towards corrosion due to high reduction potential. Corrosion Environment Cathodic Coating Exposed part of Cathodic Coating (Cathode) Base Metal (Anode) (Cathode) e e Base Metal 69 If cathodic coating contains any imperfections like pores, breaks or discontinues, then those areas of base metal acts as anode and undergo corrosion. Therefore, a properly applied and intact cathodic coating is crucial for effective corrosion protection. Example: Tinning (coating of Sn on base metal) Difference between anodic coating and cathodic coating Anodic coating Cathodic coating i) Coating of more anodic metal on the i) Coating of more cathodic metal on the surface of the base metal. surface of the base metal. ii) The coated metal should have high ii) The coated metal should have low oxidation potential than the base metal. oxidation potential than the base metal. iii) It prevents corrosion of the base metals iii) It prevents corrosion of the base metals artificially. by its noble behavior iv) If it contains any pores, breaks and iv) If it contains any pores, breaks and discontinues, those areas of the base metal discontinues, those areas of the base metal acts as cathode and prevents corrosion acts as anode and gets corrosion. v) No corrosion is formed until the entire v) No corrosion is formed unless the coating anodic coating gets corrosion. metal contains pores, breaks and Ex: Galvanization discontinues. Ex: Tinning 70 ****METHODS AND APPLICATIONS OF METALLIC COATING 1) Electroplating, 2) Electroless plating & 3) Galvanisation. 1) ELECTROPLATING/ELECTRODEPOSITION (Coating of Copper on any metal) Coating of a coat metal on the surface of any base metal by passing a direct current called Electroplating. Electrolyte to refill the loss in the case of inert anode Cu DC Line + Wooden Tank Cathode (Article to be Mn+ electroplated) - - - Mn+- - - - - Electrolyte (Solution of soluble salt of coating metal) Anode (made of pure metal to be deposited on other metal) ▪ In this method article to be electroplated is subjected to solvent cleaning using trichloroethylene to remove oils and grease present on the surface. Then the article is subjected to acid picking using dilute HCl to remove oxides. ▪ The cleaned article is made as the cathode of electrolytic cell. ▪ The pure metal to be deposited is made as the anode of the cell. Anode and cathodes are dipped into the electrolytic solution of soluble salt of coating metal. ▪ When electricity is passed through electrodes, then the anode metal undergoes oxidation and forms metal ions and the liberated electrons accumulate at the cathode. The formed metal ions migrate towards cathode and deposited on the surface of article (cathode) which results the formation of smooth, bright and thin layer of coating metal on the surface of base metal (article). ▪ Example: Electroplating of Copper (Cu) on the surface of article (base metal) with smooth and bright layer. It should maintain the following conditions. Anode : Cu pellets Cathode : Article to be electroplated (Base metal) 71  Electrolyte : CuSO4 or AgCuCl2. Temperature : 40 – 70 ⁰C Current density : 20 – 30 mA/cm2 Applications Metals and non-metals can be electroplated. In this case of metals, electroplating increases corrosion resistance and chemical resistance, lifetime, strength, hardness, and surface properties. In the case of non-metals like plastic, glass and wood, electroplating increases strength and improves decorative values. d) ELECTROLESS PLATING/AUTOCATALYTIC PLATING (COATING OF NICKEL ON ABS) Coating of a coat metal on the surface of any base metal by using reducing agent instead of electricity is called electroless plating. In this process, the reducing agent supplies electrons for the reduction of Metal Ions into Metal. M+ + R M + R Metal ion Reducing agent Metal Following are the requirements for electroless plating a) The coating metal should be in the salt soluble form. b) Reducing agents like Formaldehyde or Sodium Hypophosphite (NaH2PO2) are used. c) Accelerators like fluorides and Succinates are used to accelerate the reduction rate. d) To improve the brightness of the coat, brighteners like Pb +2 and Cd+2 are added. e) Buffer solution used to maintain PH and temperature. Example: Electroless plating of Nickel (Ni) on the surface of ABS ABS = Acrylonitrile Butadiene Styrene Ni --------------------- NiSO4 + Sodium hypophosphite --------------------- ABS (NaH2PO2) + Sodium succinate + Cd+2 + Sodium acetate --------------------- 72 In this process, reducing agent (NaH2PO2) converts Ni+2 ions into Ni Ni+2 + H2PO2ˉ+ H2O → Ni + H2PO3ˉ + 2H+ Thus formed Ni-metal will be deposited on the surface of catalytically active ABS (non-metal)/(metal). Applications In the case of metals, electroless plating increases corrosion resistance, strength, hardness and surface properties. In case of non-metals, electroless plating increases strength. Dr. Mallesham GODUMALA (Anurag University) 73