possesses negative potential & undergoes oxidation preferably act as ANODE

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1 ELECTROCHEMISTRY Introduction: Electrochemistry is the area of Chemistry dealing with the interconversion of electrical energy and chemical energy. There are many applications of this in every day life. Batteries, control of corrosion, metallurgy and electrolysis are just a few examples of the applications of electrochemistry. In general, metals tend to make good reducing/oxidizing agents because they can donate or accept electrons. The reducing or oxidizing ability of the metal is given by the electrochemical series. The more active metal is able to reduce the less active metal cation. This electrochemical series is: Li K Ca Na Mg Al Zn Cr Fe Ni Sn Pb H 2 Cu Hg Ag Pt Au possesses negative potential & undergoes oxidation preferably act as ANODE possesses positive potential & undergoes reduction - preferably act as CATHODE So for example, magnesium metal is able to reduce copper (II) ions in solution to form magnesium ions and copper metal: Mg(s) + Cu 2+ (aq) Mg 2+ (aq) + Cu(s) Note: Removal of electron is known as oxidation. Acceptance of electron is known as reduction.

2 Definition of Emf of a cell, single electrode potential and standard electrode potential: Emf of a cell: The potential difference between two electrodes of the cells which causes the flow of current is called electromotive force. It is denoted by E cell. Emf of cell = Reduction potential of cathode - Reduction potential of anode E cell = E cathode - E anode Single electrode potential: Electrode potential is the electrical potential developed at the interface between metal and its solution when it is in contact with a solution of its own ions and the two solutions are in equilibrium with each other. It is denoted by E. Standard electrode potential: Standard electrode potential is the potential developed when the pure metal electrode is in contact with a solution of its ions of unit concentration at 298K. It is denoted by E 0. Explain the origin of single electrode potential. When a metal electrode is in contact with a solution of its own ions, one of the following two chemical reactions take place. i. Metal passes into the solution as metal ions leaving behind the electrons on the metal as follows M M n+ + ne - (Dissolution) Thus, the metal acquires the ve charge. ii. Positive metal ions from the solution deposits on the metal electrode as metal atoms as follows M n+ + ne - M (Deposition) Thus, the metal acquires the +ve charge. If the dissolution reaction is faster than the deposition reaction, the liberated e - s accumulates on the surface of metal making it vely charged. This vely charged electrode surface attracts a layer of +vely charged ions at the interface. Thus a double layer called Helmholtz electrical double layer is formed at the metal solution interface. Similarly, if the deposition reaction is faster than the dissolution reaction, a layer of +ve charges formed on the surface attracts a layer of ve charges, develops an electrical double layer. Across this double layer, develops a potential called single electrode potential.

3 Nernst Equations Walther H. Nernst ( ) received the Nobel Prize in 1920 "in recognition of his work in thermochemistry". His contribution to chemical thermodynamics led to the well-known equation correlating chemical energy and the electric potential of a galvanic cell or battery. Consider an general equation M n+ + ne M For above reaction, decrease in the free energy (G) is given by, G = - nfe.( 1 ) Where, n = number of moles of electrons involved in the reaction. F= Faradays constant & E = Electrode potential As per the Gibbs-Helmholtz equation, the free energy change can be related to equilibrium constant K, G= G 0 + RT ln K..( 2 ) Equilibrium constant for reduction reaction of the type, M n+ + ne M is given by K = M M On substituting (1) & (3) in (2), we get n - nfe = - nfe o + RT ln.(3) M M n.. (4) Divide equation (4) by - nf, on both sides, and convert natural log into base 10, We get

4 E = E o RT n F log 10 M n. (5) M Under standard conditions [M] = 1, hence the above equation (5) becomes E = E o - E = E o RT n F 2.303RT n F On substituting the values for R = J/K/mol T = 298 K log 10 1 M n log 10 [M n+ ] F = C/mol, the above equation reduces to E = E o log10 [M n+ ].(6) n The equation (6) is known as the Nernst equation for single electrode potential. Note: The General Nernst Equation for a cell is given by E = E o log10 CathodeSpecies n AnodeSpecies What is an electrochemical cell? How are electrochemical cells classified? Electrochemical cell is a device which converts chemical energy into electrical energy or electrical energy into chemical energy using an electrochemical redox reaction. These can be classified into two types, 1. Galvanic cell or Voltaic cell: This is the one which converts chemical energy into electrical energy. Eg: Daniel cell Galvanic cell is further classified into three types. a. Primary Cell (Irreversible): In this, the cell reaction is not reversible. They are not rechargeable. Eg: Dry cell. b. Secondary Cell (reversible): In this, cell reaction is completely reversible and rechargeable. Therefore we use over and over again. These are called as storage cells. Eg: Ni - Cd cell, Lead - acid cell. c. Concentration Cell: These types of cell are made up of the same metal immersed in same ions solution with different concentrations. Eg: Zn Zn 2+ C1 Zn 2+ C2 Zn.

5 Potential difference arises due to transfer of substance from a solution of higher concentration to a solution of lower concentration. 2. Electrolytic cell: This is the one which converts electrical energy into chemical energy. Eg: Electroplating, electrolysis. Differences between Galvanic and Electrolytic cells. Galvanic cell Electrolytic cell 1 It requires a source of external energy It requires a source of electrical energy 2 It converts chemical energy to It converts electrical energy into electrical energy. chemical energy. 3 The redox reaction is spontaneous The redox reaction is non-spontaneous 4 The electrodes used are dissimilar metals. The electrodes used may be dissimilar or same metals. 5 Each metal electrode is dipped in its Both the electrodes are immersed in the own ionic solution and both have same electrolyte solution (single separate compartments. compartment). 6 Salt bridge is required. No salt bridge is required. 7 Cathode is positive electrode and anode is negative electrode. Cathode is negative electrode and anode is positive electrode. Describe the construction and working of a Galvanic cell? Explain the function of salt bridge. Galvanic cell is a device which converts chemical energy into electrical energy. Typical example of galvanic cell is Daniel cell. Construction and working of a Daniel cell: It consists of a Zn rod dipped in 1M solution of ZnSO 4 forms a half cell and Cu rod dipped in 1M solution of CuSO 4 forms another half-cell. The two half cells are internally connected by a salt bridge, filled with KCl solution and the two ends of the salt bridge are plugged with a porous material. The two electrodes are connected externally by a wire through voltmeter. The two half-cell reactions are represented as follows:

6 At Anode: Zn oxidizes to Zn 2+ liberating two electrons. Zn Zn 2+ +2e - At Cathode: Cu 2+ reduces to Cu by accepting two electrons. Cu e - Cu Thus the net cell reaction: Zn + Cu 2+ Zn 2+ + Cu And cell can be represented as: Zn(s) Zn 2+ (1M) Cu 2+ (1M) Cu(s) The electron liberated at the zinc half-cell builds up an electrical potential difference. i.e., electrons move from the zinc electrode to the copper electrode producing a current in the circuit which is indicated by voltmeter. In the galvanic cell, the electrode where oxidation takes place is referred as anode which is assigned with a negative sign and the electrode where reduction takes place referred as cathode which is assigned with a positive sign. Function of salt bridge: 1. During the cell reaction either Cl - ions diffuse into the zinc half-cell or Zn 2+ ions diffuse into the salt bridge to keep the zinc half-cell electrically neutral. Similarly K + diffuses into the copper half-cell or SO 2-4 ions diffuse into the salt bridge to keep the copper half-cell electrically neutral. Thus it maintains electric neutrality. 2. The salt bridge allows the current to flow through the cell without allowing the contents to mix. Without the salt bridge no electrical current would be produced and electrolytic contact must be maintained for the cell to function. List of conventions used in representing the electrodes and cells. Galvanic cell is represented as Zn ZnSO 4 (1M) CuSO 4 (1M) Cu

7 1. The electrode at which oxidation occurs (anode) is written on the left hand side whereas the electrode at which reduction occurs (cathode) is always written on the right hand side. 2. Single vertical line indicates the phase boundary between the metal and the solution. 3. Double vertical line indicates the salt bridge. 4. The concentration of corresponding solutions is represented in the parenthesis. 5. The arrow mark indicates the direction of flow of electrons. 6. The term electrode potential always refers to reduction potential and is represented as E M n+ /M What is a concentration cell? Derive an expression for emf of a concentration cell. It is an electrochemical cell in which two identical electrodes are immersed in the same ionic solution but of different concentration. As a result the potential difference arises due to transfer of substance from a solution of higher concentration to a solution of lower concentration. Derivation of expression for the cell potential of a concentration cell: Concentration cell can be represented as Zn Zn 2+ (C 1 ) Zn 2+ (C 2 ) Zn The half-cell reactions are At anode: Zn Zn 2+ (C 1 ) +2e - At cathode: Zn 2+ (C 2 ) +2e - Zn Net cell reaction: Zn 2+ (C 2 ) Zn 2+ (C 1 ) C 1 and C 2 are concentrations of Zn 2+ in the two half cells respectively and C 2 > C 1. The left hand electrode is the anode and the right hand electrode is the cathode. ZnSO 4 (C 1 ) ZnSO 4 (C 2 ) Emf of a concentration cell can be calculated using Nernst equation as follows:

8 E anode = E RT log C 1 nf E cathode = E RT log C 2 nf E cell = E cathode E anode = (E RT log C 2 ) _ (E RT log C 1 ) nf nf At 298 K, E cell = 2.303RT log C 2 nf C 1 E cell = log C 2 n C 1 E cell = log[ions at cathode] n [ions at anode] Problems: 1. Represent a cadmium-copper cell. Give the electrode and net cell reactions. If the standard reduction potentials of Cd and Cu are 0.40 and 0.34 V respectively, calculate the standard emf of the cell. In this problem, E o Cu 2+ /Cu > Eo Cd 2+ /Cd, Cu behaves as cathode and Cd acts as anode (i) Cell representation: Cd Cd 2+ Cu 2+ Cu (ii) Electrode reactions: At anode: Cd Cd e At cathode: Cu e Cu Net cell reaction:cd + Cu 2+ Cd 2+ + Cu (iii) Calculation of standard emf of the cell: E o cell= E o cathode E o anode = 0.34 ( 0.40) E o cell = 0.74V

9 2. Write the electrode reaction and calculate the emf of the following cell at 298K given Cu (s) /Cu 2+ ( 1x10-2 M ) // Ag + (1x10-1 M)/Ag (s). E o cell is 1.30V. Electrode reactions: At anode: Cu(s) Cu e At cathode: [Ag + + e Ag(s)] x 2 Net cell reaction: Cu(s) + 2Ag + Cu Ag (s) The emf of the cell is E cell = E o log [Ag + ] 2 E cell = log [1x10-1 ] 2 2 [1x10-2 ] = [ x log1] = 1.30V What are the types of electrodes? What do you mean by a reference electrode? Mention the types of reference electrodes. The following are the important types of electrode: 1) Metal-Metal ion electrodes: These are the electrodes where a metal is dipped in its metal salt solution. Eg: Cu CuSO 4, Zn ZnSO 4 2) Gas electrodes: These are the electrodes where a gas is in contact with an inert metal dipped in an ionic solution of gas molecules. Eg: Hydrogen electrode. 3) Metal insoluble salt electrodes: These are the electrodes where a metal will be in contact with its insoluble salt. Eg: Calomel electrode (Hg Hg 2 Cl 2 ), Silver electrode (Ag AgCl). 4) Ion selective electrodes: These are the electrodes which are sensitive to particular ionic species and will develop a potential when a membrane is in contact with an ionic solution. Eg: Glass electrode Reference electrodes: These are the electrodes which are used to determine single electrode potential of other electrodes and are also called as standard reference electrodes. There are two types of reference electrodes. 1. Primary reference electrode. Eg: Standard hydrogen electrode. 2. Secondary reference electrode. Eg: Calomel electrode, Silver electrode, ion selective electrode.

10 Explain the construction and working of Calomel electrode Calomel electrode consisting of a glass tube with two side tubes. At the bottom of which mercury is placed above which a layer of mercury and mercurous chloride (called calomel) and remaining portion of the tube is filled with saturated KCl solution through one of the side tube. The other side tube is connected to a salt bridge. A platinum wire is dipped into the mercury for external electrical contact. Electrode potential of the cell depends on the concentration of KCl used. KCl concentration 0.1N 1N Saturated Electrode potential (V) The calomel electrode acts as both anode and cathode depending upon the other electrode used. When it acts as anode the electrode reactions is, 2Hg + 2Cl - Hg 2 Cl 2 + 2e - When it acts as cathode the electrode reaction is Hg 2 Cl 2 + 2e - 2Hg + 2Cl - The overall reaction: Hg 2 Cl 2 + 2e - 2Hg + 2Cl - Representation of calomel electrode: Hg Hg 2 Cl 2 KCl (sat) The electrode potential (reduction potential) is given by E cal = E log [Cl - ] The electrode potential is decided by the concentration of chloride ions and the electrode is reversible with respect to chloride ions. Advantages: 1) Calomel electrode is simple to construct.

11 2) Cell potential is reproducible 3) Potential stable over a long period. 3) Cell potential does not vary with temperature. Hence it is used as secondary reference electrode. Applications: 1. It is used as secondary reference electrode in measurement of single electrode potential of other electrodes instead of SHE. 2. It is used as reference electrode in all potentiometric titrations. 3. It is used as reference electrode with glass electrode in ph determination. BATTERY TECHNOLOGY What is a Battery? Explain the classification with examples. Battery is a collection of cells connected either in series or in parallel to get required amount of energy. Classification of Batteries; Batteries are classified into three types as follows. a) Primary Battery b) Secondary Battery a) Primary Batteries: These are the batteries which serve as a source of energy only as long as the active chemical species are present in the battery or in the cell. The cell reactions are irreversible. These are designed for only single discharge and cannot be charged again. Ex: Dry Cell, b) Secondary Batteries: These batteries are chargeable and can be used again and again. The cell reactions are reversible and are often called reversible batteries. During discharging the cells acts like voltaic cell converting chemical energy into electrical energy. During charging the cell acts like electrolytic cell by converting electric energy into chemical energy, hence these batteries are called as storage battery. Ex: Lead- acid Battery, Ni- Cd battery etc. Construction & Working of dry cell: Construction: It consists of Zinc anode in the shape of container containing paste of ammonium chloride in water. Carbon at the center acts as cathode and mixed with manganese dioxide to improve the conductivity. In general dry cell can be represented as follows:

Zn(s) NH 4 Cl(aq) MnO 2(s) C Working: As cell switched on zinc is oxidized and manganese dioxide is reduced: Anode: Cathode: Zn Zn 2+ + 2e - 2NH 4 + + 2 MnO 2 + 2e - Mn 2 O 3 + 2NH 3 + 2 H 2 O

12 Zn(s) NH 4 Cl(aq) MnO 2(s) C Working: As cell switched on zinc is oxidized and manganese dioxide is reduced: Anode: Cathode: Zn Zn e - 2NH MnO 2 + 2e - Mn 2 O 3 + 2NH H 2 O Overall: Zn + 2NH MnO 2 Zn 2+ + Mn 2 O 3 + 2NH H 2 O The voltage of the cell is about 1.5 V Advantages: Low cost and non - toxic materials. The alkaline electrolyte does not readily react with zinc (compare Zn- C cell above) giving a much longer shelf- life (5 years) and

13 the current and voltage are steady due to the strong base/alkali electrolyte having a smaller resistance the ammonium chloride - carbon paste. Disadvantages: Cannot be recycled, more expensive due to extra sealing and low power. Fuel Cells Fuel cells are the galvanic cells in which chemical energy of fuel is directly converted into electrical energy. Classification of Fuel cells Depending on temperature, these are classified into three types as follows. 1] Low temp fuel cells: Which operates at the temp range below 100 o C 2] Moderate temp fuel cells: Which operates at the temp range about 100 o C to 250 o C 3) High temp fuel cells: Which operates at the temp range about 500 o C Based on type of electrolyte used, these are classified into three types as follows 1. Alkaline fuel cells (AFC): 1. These fuel cells containing alkali such as KOH or NaOH as electrolyte. Hydrogen is used as fuel and oxygen gas is used as an oxidant. The cell operates at a temp of 80 0 C. Uses: These are used in emergency lights and portable power, generations, space applications, military applications etc 2. Phosphoric-acid fuel cell (PAFC): These fuels cells consisting of 98% phosphoric acid, 2% water as electrolyte, O 2 is used as oxidant. Hydrogen LPG, NPG etc, are used as fuels. These operate at a temp 190 to C. Platinum alloys such as platinum- cobalt- chromium, are used as electro catalyst. These cells are used to provide light and heat in large buildings. 3. Solid oxide fuel cell (SOFC): These contains ZrO 2, Y 2 O 3 are solid electrolytes. Cathode is made up of porous strontium doped with LaMnO 3 or In 2 O 3 and SnO 2. Anode is made up of cobalt, nickel, or ZrO 2. Operating temperature is C. These cells are used in KW power plants, water heating etc, 4. Molten carbonate fuel cell (MCFC): These fuel cells consisting of molten carbonates such as lithium carbonate 26.2% and potassium carbonate (K 2 CO 3 ) 23% and lithium. Aluminium carbonate as electrolyte. The anode is made up of nickel and cathode made up of nickel oxide. Operating temperature is C. These are used in chemical industries such as aluminum Cloro alkali industries. 5. Proton-Exchange-Membrane fuel cell (PEMFC): These contain ion exchange membrane as solid electrolyte for ionic conduction nafion - R membranes which

14 are chemically and electrochemically stable at C are used.operating temperature of the cell is 80 0 C. The electrodes are made up of platinum and noble metals are used as electro catalysts. Uses: Used in the manned Gemini terrestrial orbital missions. 6. Biochemical fuel cell (BCFC): An electrochemical power generator in which the fuel source is bioorganic matter; air is the oxidant at the cathode, and microorganisms catalyze the oxidation of the bioorganic matter at the anode. Explain the construction and working of H 2 - O 2 fuel cell. Hydrogen - oxygen fuel cells consisting of two porous graphite electrodes, which is impregnated with an electro catalyst such as finely, divided Pt or Ni. Concentrated KOH is used as an electrolyte. Hydrogen gas and oxygen gas are continuously supplied to the anode and cathode respectively. The hydrogen undergoes combustion generating electric current. The cell delivers an emf of 1.23V. The cell reactions are as follows. Anode Reaction: H 2 + 2OH - 2 H 2 O + 2 e - Cathode Reaction: 1/2 O 2 + H 2 O + 2 e - 2OH - Overall Reaction: H 2 + 1/2 O 2 H 2 O Uses: The H2-O2 cells are used in Space vehicles, military and mobile power systems. Difference between battery and fuel cell

15 Sl.No. Battery Fuel cell 1. It makes electricity from the chemical reactant it has stored inside the battery makes its electricity from fuel supplied externally 2. Comparatively lifetime of the battery is lesser lifetime of the fuel cells is couple of years longer than batteries 3. Used chemicals to produce electricity are hazardous to environment Eco friendly fuels are used to produce electricity 4. It need re-charging to supply energy continuously It won t need re-charging; it can supply energy continuously 5. The electrodes are consumed The electrodes are not consumed Question bank: 1. Define single electrode potential. Derive Nernst equation for single electrode potential. 2. Explain the construction and working of Calomel electrode. 3. Explain the construction and working of Daniel cell (Galvanic cell). 4. Discuss construction and working of dry cell 5. With neat sketch explain working of hydrogen oxygen fuel cell. 6. What are concentrations cells? Explain with an example. 7. What is a Battery? Explain the classification of Batteries with examples Objective type questions: 1. Converts chemical energy into electrical a) Galvanic cell b) Daniel cell c) Dry cell d) all 2. In the anodic chamber reaction takes place a) Oxidation b) Reduction c) Addition d) Substitution 3. In the cathodic chamber reaction takes place a) Oxidation b) Reduction c) Addition d) Substitution 4. Origin of electrode potential is explained in a) Nernst theory b) Helmholtz double layer theory c) galvanic theory d) Electrochemical theory 5. In two half cells, the one which is having high negative value acts as a) anode b) Cathode c) Dry cell d) None 6. EMF of a concentration cell depends on a) [M1] & [M2] b) No. of charges c) temperature d) All of the above 7. Calomel electrode potential is dependent of a) Cl- concentration b) Hg2Cl2 c) Temperature d) None 8. EMF of the cell comprising of Zn (E 0 = ) and Cu (E 0 = ) is 1.1V. 9. Standard electrode potential is nothing but single electrode potential at a) Unit concentration b) 1 atm. pressure c) 298 K d) All of the above 10. Double vertical line ( ) represents a) Solid liquid interphase b) Salt bridge c) solid - solid or Liquid - liquid interphase d) None 11. Daniel cell is a combination of standard electrodes of a) Cu & Ag b) Zn & Cd c) Zn & Cu d) Cu & Cd 12. The electrolyte used in alkaline fuel cells a) KOH b) H3PO4 c) K2CO3 d) ZrO2

Oxidation-Reduction (Redox)

Oxidation-Reduction (Redox) Electrochemistry involves the study of the conversions between chemical and electrical energy. Voltaic (galvanic) cells use chemical reactions to produce an electric current.