Electrochemistry Notes PDF

Electrochemistry It is the study of chemical reactions produced by passing electric current through an electrolyte or the production of electric current through chemical reactions. The cells which convert chemical energy to electric energy is called electro chemical cells and cells which convert el...

Electrochemistry It is the study of chemical reactions produced by passing electric current through an electrolyte or the production of electric current through chemical reactions. The cells which convert chemical energy to electric energy is called electro chemical cells and cells which convert electric energy to chemical energy is called electrolytic cells. Electrochemical cell- Daniel cell Daniel cell consists of two half cells in which zinc and copper electrodes are immersed in zinc salt and copper salt solutions respectively. The two half cells are internally connected by a salt bridge and externally by a metallic wire. At anode, the zinc electrode undergoes oxidation and loses electrons. The electrons liberated migrate to the copper half-cell. At cathode, the cupric ions accept these electrons, undergo reduction, and get deposited on copper electrode as copper atoms. The electrons generated at the Zn-electrode are consumed at the copper electrode. The flow of electrons (negative electricity) in the cell is from the Zn-electrode to the Cu electrode. Consequently, the flow of positive electricity i.e. the conventional current in the cell is from Cu- electrode to the Zn-electrode. The cell reaction can be obtained by adding equations of Anode and Cathode half-cells.  Zn(s) + Cu2+aq) → Zn2+aq) + Cu(s) or simply as Zn + Cu2+ Zn2+ + Cu Also note that the flow of positive electricity, ie., the conventional current is from Cu to Zn electrode. When the electrons flow through the external wire, for the completion of circuit, there should be the flow of ions between the compartments. For the transference of ions between the compartments a salt bridge or a porous partition is used. The Daniel cell may be represented as Zn(s) / Zn2+(aq) // Cu2+(aq) / Cu(s) or simply as Zn / Zn2+// Cu2+ / Cu NB: Anode is written on the left and cathode is written on the right, “//” represents the salt bridge & the ionic species are written on either side of the salt bridge. Differences between electrochemical cells and electrolytic cells Table 1: Differences between electrochemical cells (Eg. Daniel cell) and electrolytic cells (Eg. Electroplating). Standard electrode Potential, EO is defined as the potential developed at electrode- electrolyte interface, when a metal is in contact with its own ions at 1M concentration and 298K. If the electrodes involve gases, the gas should be passed through the electrolyte at 1 atm pressure. EMF is defined as the potential difference between the two electrodes of a galvanic cell which causes the flow of current from the electrode of higher reduction potential to the electrode of lower reduction potential. It is denoted as Ecell. Ecell = Ecathode – Eanode Under standard conditions, EOcell = E0cathode – EOanode Salt Bridge and its Significance (a) It connects the solutions of two half - cells and completes the cell circuit. (b) It prevents transference or diffusion of the solutions from one half cell to the other. Eg: in Danel cell, the diffusion of Cu2+ to the Zn compartment must be prevented as it gets deposited over the Zn metal. (c) Salts like KCl, NH4NO3 are used for making salt bridge. Ions like K+ and Cl- has equal mobility and get drifted to the cathode and anode compartments respectively under the potential difference. Types of Electrodes 1. Metal/Metal ion electrode (M/Mn+ When a metal is immersed in the respective metal salt solution, a metal/metal ion electrode is formed. Eg. Zinc electrode: a zinc rod dipped in a solution of zinc sulphate Zn/Zn2+, silver electrode: a silver rod dipped in a solution of silver nitrate, Ag/Ag+, or copper electrode: copper rod dipped in copper sulphate solution, Cu/Cu2+ etc. 2. Gas Electrode. In this type, the gas is passed through an inert electrode like platinum which is immersed in a solution containing its own ion. Eg: Hydrogen electrode: Hydrogen gas is bubbled through a platinum electrode immersed in HCl solution, (Pt/H2/H+) or Chlorine electrode: chlorine gas is bubbled through a platinum electrode immersed in NaCl solution (Pt/Cl2/Cl− ). 3. Metal/insoluble metal salt/common ion electrode. In this type the metal is covered with a paste of its insoluble metal salt which is in contact with a solution containing a common ion with the insoluble salt. Eg. Calomel electrode: Mercury in contact with an insoluble precipitate of Hg2Cl2 which is in contact with KCl solution. (Hg/Hg2Cl2/Cl−). Another example is Silver- silver chloride electrode: Silver in contact with a precipitate of AgCl which in contact with KCl solution (Ag/AgCl/Cl−). 4. Redox electrode. In this type an electrode like platinum is in contact with a redox system. Example a platinum electrode immersed in a mixture of ferric chloride and ferrous chloride solution, Pt/F e2+/F e3+. Another example a platinum electrode immersed in a mixture of cerric sulphate and cerrous sulphate solution, Pt/Ce4+/Ce3+. 5. Ion selective electrode. In this type, the sensing part of the electrode is usually made of an ion-specific membrane. The membrane can be glass membrane, crystalline membrane and ion- exchange resin membrane. Eg. glass electrode. Here the glass membrane is made of an ion- exchange type of glass which are sensitive to specific ions like H+; Na+; Ag+ etc. Single Electrode Potential and Helmholtz electrical double layer The tendency of a metal to get Oxidized or Reduced when it is placed in a solution of its own salt is called Electrode Potential. M Mn++ ne (Oxidation) Mn+ + ne M (Reduction) In the first case (oxidation) some metal ions enter the electrolyte solution leaving behind the electrons on the electrode. Thus, electrode gets negative charge and solution side get positive charge due to excess positive metal ions. In the second case (reduction) the positive metallic ions from the solution take electrons from the electrode and get deposited as metal on the electrode surface. Now the electrode gets a positive charge and the solution side gets a negative charge (due to deficiency of positive metal ions). Equilibrium is reached in the vicinity of the electrode, due to the electrostatic force; any further transference of metal ions does not take place. Thus, the positive charge and negative charge remain close to the metal surface forming a double layer. This is called Helmholtz electrical double layer. As a result, a difference in potential is set up between the metal and the solution this potential is called single electrode potential. The tendency of the electrode to lose electrons is called oxidation potential and the tendency of the electrode to gain electrons is called reduction potential. This potential difference becomes a constant at equilibrium. This constant potential difference developed when a metal is in contact with its own salt solution of concentration 1 M at 25oC is called the standard electrode potential of the metal. Reference Electrodes Primary Reference Electrodes Standard Hydrogen Electrode (SHE) or Normal Hydrogen Electrode (NHE) The electrode potential is found out by coupling the electrode with a primary reference electrode, the potential of which is arbitrarily fixed as zero. The important primary reference electrode used is a standard hydrogen electrode. Standard hydrogen electrode (SHE) consists of, ▪ A Platinum wire in an inverted glass tube. ▪ Hydrogen gas passed through the tube at 1 atm. ▪ A platinum foil attached at the end of the wire. ▪ An electrode immersed in 1M H+ ion solution at 25°C. The electrode potential of SHE is zero at all temperatures. It is represented as: Pt/H2(g)1 atm/H+ (1M) It is a reversible electrode, which can act as a cathode or anode, depending on the potential of the electrode to which it is coupled. If the potential of the coupled electrode is less than zero eg (Zn, Mg etc), reduction takes place in SHE and it acts as cathode. H+ + e- ½ H2(g) (reduction) If the potential of the coupled electrode is greater than zero eg (Cu, Ag etc),oxidation takes place in SHE and it acts as anode ½ H2(g) H+ + e- (oxidation) Determination of single electrode potential of Zn-electrode using SHE To measure the standard electrode potential of Zn (ie., Zn electrode dipped in 1M ZnSO 4 solution). The Zn electrode is coupled with SHE through a salt bridge. Zn has greater tendency for oxidation than SHE, hence Zn acts as anode and SHE acts as cathode. The digital voltmeter reading gives the cell emf = 0.76V.  E0cell = E0 cathode - E0 anode ie., E0cell = E0 SHE - E0 Zn/Zn2+ Since E0 of SHE is zero 0.76V = 0 - E0Zn/Zn2+ EOZn/Zn2+ = -0.76 V Limitations of SHE ✓ It is very cumbersome to set up a SHE ✓ It requires considerable volume of solution ✓ The hydrogen electrode will be poisoned by compounds like Hg, As, S, Fe3+ ✓ It cannot be used in solutions containing redox systems ✓ Difficult to maintain pressure of H2 gas at 1 atm and difficult to maintain the concentration of 1M HCl. Electrochemical series By measuring the potentials of various electrodes versus standard hydrogen electrode (SHE), a series of standard electrode potentials has been established. It is the series in which the standard reduction potentials are arranged in the descending order. According to IUPAC only reduction reactions or reduction potentials are taken into consideration. Reduction potential is a measure of the tendency of a chemical species to acquire electrons and thereby be reduced. Reduction potential is measured in volts (V), or millivolts (mV). Applications of electrochemical series 1) Reactivity of metals: Electrodes having negative values of standard electrode potentials with respect to hydrogen have tendency to undergo oxidation, whereas electrodes having positive values have tendency to undergo reduction in aqueous solutions. 2) Displacement reactions: To predict whether a given metal will displace other. Case 1: A metal having lower EO will displace the metal having higher EO values from its salt solution. Eg: Zn metal can displace Cu from CuSO4 solution. EOZn2+/ Zn= -0.76 V and EOCu2+/ Cu= +0.34 V. Zn + CuSO4 ZnSO4 + Cu Case 2: Metals having negative reduction potentials can displace hydrogen from acids. Eg: Fe, Zn, Al etc can displace hydrogen from acids. Zn + H2SO4 ZnSO4 + H2 Case 3:A non metal having higher reduction potential can displace a non-metal having lower reduction potential. Eg: If we pass fluorine gas through NaCl solution, displacement reaction takes place generating chlorine gas and sodium fluoride. 2NaCl + F2 2NaF + Cl2 F2 + 2e- 2F- E0 = +2.87 Cl2 + 2e- 2Cl- EO = +1.36 3) Determination of standard emf of a cell: Standard emf of a cell can be calculated from the standard electrode potentials of respective electrodes. E= ER-EL= Ecathode- Eanode 4) Spontaneity of redox reaction can be judged Eg: if we add K2Cr2O7 to acidified mixture of Ferrous chloride and Potassium iodide. Oxidation of I- to I2 takes place first than oxidation of Fe2+ to Fe3+. The reaction having more cell emf is more spontaneous which is evident form following reduction potentials. Cr2O7−2 + 14H+ + 6e 2Cr3+ + 7H2O; E0 = +1.33V F e3+ + e − F e2+; E0 = +0.77V I2 + 2e − 2I−; EO = +0.54V Similarly acidified dichromate solution is not capable of liberating Cl2 gas form NaCl, but acidified permanganate can liberate Cl2 gas from NaCl, which can be explained form following reduction potentials. Cr2O72− + 14H+ + 6e 2Cr3+ + 7H2O; E0 = +1.33V MnO4− + 8H+ + 5e − Mn2+ + 4H2O; E0 = +1.51V Cl2 + 2e − 2Cl−; EO = +1.36V The cell emf of permanganate- Cl2/Cl− cell is positive and the reaction is spontaneous. whereas cell emf of dichromate- Cl2/Cl− cell is -ve and the reaction is non spontaneous. 5) Analysis of products of electrolysis: For example, if we electrolyze a mixture aqueous solution of KCl and KI using platinum electrodes. H2 is liberated at cathode (-ve terminal) and I2 is liberated at +ve terminal. The ions approaching -ve terminal are H+ and Na+, The specious having high reduction potential (less -ve) undergo reduction (EOH+/H2 = -0.41V at pH = 7 and EONa+/Na = -2.71V). Thus, H2 is liberated rather than Na metal. Similarly, the ions approaching +ve terminal are Cl−; I− and OH−. The specious having low reduction potential get oxidized (EO Cl2/Cl− = +1.36V , EOI2/I− = +0.54V and EO O2/OH− = +0.82V at pH = 7). Thus, I2 is liberated rather than O2 or Cl2 gas. 6)Decomposition potential: Decomposition potential of water can be calculated from electrochemical series. Water can be split into hydrogen and oxygen just by applying 1.23V which is evident from the following reduction potentials of oxygen and hydrogen electrode. O2 + 4H++ 4e- 2H2O (1 M H+ acidic medium) EO= +1.23 V 2H+ + 2e- H2 (1 M H+ acidic medium) EO= 0 Relation between EMF of the Cell and Free energy Free Energy (∆G) is defined as energy available (per mole) that can completely convert into mechanical work. Electrical energy is a form of energy that can be completely converted into mechanical work. So, decrease of free energy is equal to electric energy produced in a reversible cell. Electric energy produced per mole = quantity of charge in one mole × EMF For the reaction Mn+ (aq) + ne- M(s) Quantity of charge in one mole = nF Therefore, electric energy per mole = nFE Since decrease of free energy is converted to electrical energy ∆G=-nFE Where F is called Faraday which is the charge of one mole of electrons i.e., F = 6.023 × 1023 × 1.6 × 10−19 = 96500C Nernst equation for the electrode potential It gives us the effect of electrolyte concentration on electrode potential. For the reaction Mn+ (aq) + ne- M(s) The decrease in free energy, ∆G accompanying this process is given by the well-known thermodynamic equation. (Vant Hoff’s equation) ∆G = ∆G0 + RT lnQ (1) where Q is called reaction quotient, which is the ratio of activities of products to the reactants in a non-equilibrium situation. (At Equilibrium ∆G = 0 and Q = K, equilibrium constant) In a reversible cell, electrical energy is produced at the expense of decrease in free energy, and free energy is related with EMF of the cell as follows ∆G = -nFE (2) Substituting equation (2) in (1) -nFE = -nFE0 + RTlnQ Dividing by -nF (3) Substituting the values of constants at the temperature T = 298K; R = 8.314 JK-1mol-1; F = 96500C (Faraday)  (4) Equations (3) and (4) are different forms of Nernst’s equations. Case 1: Nernst Equation for single electrode: For a single electrode reaction Mn+ (aq) + ne- M(s) reaction quotient Q can be written as In dilute solution, activities are replaced by molar concentration. Also the concentration terms [e] is taken as one and for pure solid [M] is also taken as one, therefore, Thus equation (4) can be written as, Case 2: Nernst Equation for a complete cell : For a cell reaction aA + bB cC + dD Then, And corresponding equation (4) will be For example, for Daniel cell, Here concentration of pure solids are taken as one, hence So, the corresponding Nernst equation will be, Applications of Nernst Equation (Related problems were discussed in the class) 1. It can be used to study the effect of electrolyte concentration on electrode potential. 2. It can be used for the calculation of the potential of a cell under non-standard conditions. 3. The unknown concentration of one of the ionic species in a cell can be determined with the help of Nernst Equation, provided that E0cell and concentration of other ionic species are known. 4. The pH of a solution can be calculated using this equation. 5. It can also be used for finding the valency of an ion or the number of electrons involved in the electrode reaction. 6. Solubility of sparingly soluble salt at any temperature can be determined. Secondary reference electrodes Standard hydrogen electrode is the primary reference electrode. Calomel electrode and silver/silver chloride/KCl electrodes are commonly used secondary reference electrodes. We often use secondary reference electrodes due to limitations of primary reference electrode SHE. Calomel Electrode Calomel electrode is a secondary reference electrode. This is a commonly used reference electrode since it is difficult to set up and maintain SHE. This is an example of metal-metal insoluble metal salt electrode and a solution of its common ion (mercury-mercurous chloride electrode, KCl). It consists of a glass tube at the bottom of which a small amount of Hg is placed. This is covered with a paste of solid mercurous chloride (Hg2Cl2 or calomel) which is further in contact with a solution of KCl. A Pt wire dipped into the Hg layer is used for making electrical contact. The side tube is used for making electrical contact with a salt bridge.  Calomel electrode can be represented as Hg/Hg2Cl2(s)/KCl Or Hg/ Hg2Cl2(s)/Cl- Net Reaction Hg2Cl2 +2e- 2Hg + 2Cl- Calomel electrode is a reversible electrode. So, it can act as either anode or cathode depending on the coupled electrode. If Anode 2Hg + 2Cl- Hg2Cl2 +2e- If Cathode Hg2Cl2 +2e- 2Hg + 2Cl- The Nernst equation of this electrode at 25oC is given by From the above it is seen that the potential of the calomel electrode varies with the concentration of the KCl solution. Saturated calomel electrode is generally used, since it is easy to setup and its emf will not change with chemical reaction. Measurement of electrode potential using calomel electrode It is impossible to know the absolute value of a single electrode potential because neither oxidation nor reduction takes place independently. Hence, we find the relative electrode potential by coupling it with a reference electrode. The primary reference electrode is the normal or standard hydrogen electrode (NHE or SHE), but it is very difficult to setup. So we often go for secondary reference electrode like saturated calomel electrode. Zn-electrode: To measure the standard electrode potential of Zn electrode, (Zn- electrode dipped in 1M ZnSO4 solution), the Zn electrode is coupled with saturated calomel electrode through a salt bridge and the emf of the cell is measured. The electrode having higher reduction potential undergoes reduction and is the cathode and the electrode having lower reduction potential undergoes oxidation and is the anode. The reduction potential of saturated calomel electrode (Esce = 0.2422) is higher than standard reduction potential of Zn2+/Zn = -0.76V. Therefore, reduction reaction takes place at SCE and the Zn electrode undergo oxidation. Cathode: Hg2Cl2 + 2e- 2Hg + 2Cl-; E0 = 0.2422V Anode: Zn Zn2+ + 2e- The cell may be represented as Zn/Zn2+//Cl-/Hg2Cl2/Hg Ecell = ER - EL Ecell = Esce - EZn2+/Zn when concentration of [Zn2+] = 1M and tempetrature = 25oC Ecell = 0.2422 - Eo Zn2+/Zn By measuring cell emf E Zn2+/Zn can be calculated. In the case of Copper Eo Cu2+/Cu it will act as cathode and SCE will be anode. The cell may be represented as Hg /Hg2Cl2/ Cl- // Cu2+/Cu Advantages of Calomel Electrode: It is simple to construct. The electrode potential is reproducible and stable. It can be conveniently handled in lab as a reference electrode Silver-Silver chloride electrode The electrode can be easily constructed by placing a precipitate of AgCl on a loop of silver wire immersed in KCl solution The electrode can be represented as Ag/AgCl/KCl Net Reaction AgCl(s) + e - Ag(s) + Cl- E0 = 0.22V  Electrode Potential of Silver- Silver Chloride Electrode Ecal = Eocal – 0.0591 log [Cl-] Ecal = 0.22 – 0.0591 log [Cl-] Glass electrode It is a pH sensitive electrode widely used for pH determination. ▪ For pH measurement H+ ion selective glass electrode is used. ▪ It consists of low melting glass having high electrical conductivity. ▪ The glass electrode assembly consists of a long glass tube with a thin glass bulb filled with 0.1 M HCl. ▪ The inner surface of the glass is in contact with a AgCl coated silver electrode or simply a platinum contact electrode. ▪ HCl in the bulb furnishes a constant H+ ion concentration. ▪ For measuring the pH of a test solution, the glass bulb is immersed in the test solution such that the outer surface of the glass bulb comes in contact with the test solution The glass membrane functions as an ion-selective resin, and an equilibrium is set up between the Na+ ions of glass and H+ ions in solution H+ +Na+ (glass) Na+ + H+(glass) Electrode can be represented as Ag/AgCl/ 0.1M H+ /glass/H+ (c=?) The Nernst equation for the electrode EG = EoG – 0.0591 log 1 [H+] EG = EoG – 0.0591 pH Determination of pH using glass electrode To determine pH of an unknown solution, the glass electrode is combined with a secondary reference electrode such as calomel electrode and the glass-calomel electrode assembly is dipped in the solution whose pH is to be determined. The two electrodes are connected to a digital potentiometer where EMF of the cell is measured. Unknown pH can be determined using the following equations. E cell = Ecalomel – Eglass Ecell = 0.24 – (E0G -0.0591pH) Ecell = (0.24 – E0G) + 0.0591 pH Or glass electrode assembly is dipped in the unknown pH solution and connected to a digital pH meter, which directly displays the pH of the unknown electrolyte. Conductivity The substances which allow the passage of electric current through them are called conductors while those which do not allow flow of electric current are called insulators. Conductors are of two types 1. Electronic conductors: example: metals and semiconductors. 2. Electrolytic conductors: example: salt solutions, acids, molten salts Conductivity of electrolytic conductors (10-3Scm-1) are very low when compared with metals (107Scm-1). This is due to lesser number of charge carriers per unit volume and lower mobility of ions when compared with electrons. Electrical resistance (R) Electrical resistance measures the obstruction to the flow of current. Resistance of a conductor is proportional to length(l) and inversely promotional to area of crossection (a) of the conductor The constant of proportionality ‘ρ’ is called specific resistance or resistivity of the material. Thus specific resistance is the resistance of a conductor of unit length and unit area of cross section. The unit of resistivity is ohm cm. Electrical conductance (C) It is a measure of the ease with which the current flows through a conductor or a solution. The reciprocal of resistance is called conductance. The unit of conductance is ohm-1 or mho or Siemen, S.  C = 1/ R Specific conductance or conductivity (κ) The reciprocal of specific resistance (ρ) is called specific conductance. It is defined as the conductance of one cm3 of a conductor or solution held between two electrodes of 1 cm2 area placed on the opposite pairs of faces of cube of edge length 1 cm. Where l/a is called cell constant. conductivity(κ) = conductance(C) × cell constant. Cell constant: In a conductivity cell, the distance between two electrodes (l) and the area of electrodes are fixed. This constant is called cell constant. cell constant= l/a. Unit of cell constant is cm-1 or m-1. Conductivity cell A conductivity cell consists of two electrode plates coated with finely divided platinum black (to minimize polarization effects). These two are welded with the ends of platinum wires that are sealed through two glass tubes. The tubes are strongly fixed such that distance between electrodes remains unaltered. Determination of cell constant Cell constant may be obtained by measuring ‘l’ and ‘a’. But it is very difficult to measure the actual area precisely in this type of small cells. In actual practice cell constant is measured by filling the conductivity cell with a solution of known conductivity like 0.1 M KCl solution. The conductivity (specific conductance) of 0.1M KCl solution is 0.01288Scm-1 at 298K. Conductivity of 0.1 M KCl solution(κ) = conductance(C) × cell constant By measuring conductance of 0.1 M KCl solution in the given cell, the cell constant can be determined. Problem.Calculate the conductivity of given sample of water at 298 K, which shows a conductance of 560 µS in the given cell at 298 K. A standard solution of 0.1 M KCl shows a conductance of 12.34 mS in that cell. (Given that conductivity of 0.1M KCl at 298 K is 0.01288Scm-1) Prepared by Dr. Sreedevi P., Assistant Professor (Ad-Hoc), NSS Engg. College, Palakkad