Electrode System PDF

Electrode System An electrochemical cell is a device that converts chemical energy into electrical energy or electrical energy into chemical energy. Thus, there can be two types of electrochemical cells. (i) Electrolytic cells are devices that convert electrical energy into chemical energy. (ii) Gal...

Electrode System An electrochemical cell is a device that converts chemical energy into electrical energy or electrical energy into chemical energy. Thus, there can be two types of electrochemical cells. (i) Electrolytic cells are devices that convert electrical energy into chemical energy. (ii) Galvanic cells (or voltaic cells) are devices that convert chemical energy into electrical energy. E.g. Daniel cell (Fig.1): The Daniel cell consists of two containers, one of us with a zinc rod dipped in zinc sulfate solution and the other is with a copper rod dipped in a copper sulphate solution. The solutions of two containers are connected by a salt bridge. The salt bridge is a U- tube, filled with either a jelly containing KCl or NH4NO3 or a saturated solution Of KCl or NH4NO3, and both the ends of the U-tube are plugged with a porous material. The zinc and copper rods (electrodes) are connected externally through an ammeter by using a wire, the following changes are observed: Zinc starts to dissolve. Copper gets deposited on a copper rod. The ammeter indicates the flow of electrons from zinc to the copper rod. The zinc sulfate solution becomes richer with Zn2+ ions. The copper sulfate solution becomes more dilute with respect to Cu2+ ions. Anode (oxidation): Zn → Zn2+ + 2e- Cathode (reduction): Cu2+ + 2e- → Cu Cell reaction: Zn + Cu2+ → Zn2+ + Cu Fig.1. Daniel cell. 1  The flow of electrons from zinc electrode to copper electrode. Hence the flow of conventional current is from copper to zinc electrode. Cell notation and convention: By convention, the electrode at which oxidation occurs is the anode and the electrode where reduction occurs is the cathode. Thus, in the above cell zinc electrode acts as an anode. And copper electrode acts as a cathode. In a galvanic cell, anode is negative and the cathode is positive. The reason, for this, is that oxidation is accompanied by the liberation of electrons which are given up to that electrode, which thereby acquires a negative charge. Reduction, on the other hand, is accompanied by the absorption of electrons by the reactant in the solution from the electrode, which thereby acquires a positive charge. According to present conventions, a galvanic cell is represented by keeping view the following points: (1) Anode is written on the left-hand side, while the cathode is written on the right-hand side. (2) The electrode on the left (i.e., anode) is written by writing the metal (or solid phase) first and then the electrolyte. Two are separated by a vertical line or a semicolon. The electrolyte may be represented by the formula of the whole compound or by ionic species. Additional information regarding the concentration may also be mentioned in the bracket. Eg., Zn | Zn2+ or Zn ; Zn2+ or Zn | ZnSO4(1M) Pt, H2 (1 atm.) | H+(1M). (3) The cathode of the cell (at which reduction takes place) is written on the right-hand side. In this case, the electrolyte is represented first and the metal (or solid phase) thereafter. The two are separated by a vertical line or a semicolon. E.g., Cu2+| Cu or Cu2+; Cu or Cu2+ (1M) | Cu or Cu2+ (1M); Cu (4) A salt bridge is indicated by two vertical lines, separating the two half cells. Thus applying the above considerations to Daniel cell, we may represent as: Zn | Zn2+ (1M) || Cu2+(1M) | Cu 2 The electromotive force of the cell (EMF): The difference in potential which causes a current to flow from an electrode of higher potential to that of lower potential is called the electromotive force (emf). The cell potential is represented by the Ecell. Ecell = ERHS - ELHS Ecell = Ecathode – Eanode Standard emf of the cell is defined as the emf of the galvanic cell when the reactants and products of the cell reaction are at unit concentration or unit activity, at 298K and 1-atmosphere pressure. E0cell = E0RHS – E0LHS The function of salt bridge: In a galvanic cell, if the salt bridge is not used, then the two solutions should be either in contact with each other or separated by a porous membrane. In either case, positive and negative ions of the electrolytes will migrate across the junction in opposite directions as there is an increase of positive and negative charges at the anode and cathode respectively. Inside the cell, the current is carried by ions, negative ions from cathode to anode and positive ions from anode to cathode. Because of the difference in the migration velocities of these ions, the accumulation of charges across the junction occurs. This gives rise to the formation of an electrical double layer developing a potential called liquid junction potential across the junction. In such situations, the emf of the cell includes liquid junction potential also. The use of the salt bridge, however, avoids the liquid junction potential and provides contact between the two solutions. It consists of a saturated solution of a salt such as KCl or NH4NO3, whose ions have almost the same migration velocities. The positive and negative ions of the salt in the salt bridge migrate with equal speed into cathode and anode compartments respectively, thereby avoiding liquid junction potential. During the cell reaction, either Cl- diffuse from the salt bridge into the zinc half-cell or Zn2+ ions diffuse into the salt bridge to keep the zinc-half cell electrically neutral. At the same time, the copper half-cell is kept electrically neutral by the diffusion of either the positive (K+) ions from the salt bridge to the copper cell or the diffusion of NO3- ions from the copper cell into the salt bridge. Without the salt bridge, no electrical current would be produced by the galvanic cell since electrolytic contact must be maintained for the cell to function. 3 Electrode is a part of an electrochemical setup where an element is in contact with its own ions at which oxidation or reduction takes place. An electrode is a point where current enters and leaves the electrolyte. When the current leaves the electrodes it is known as the cathode and when the current enters it is known as the anode. Electrodes are vital components of electrochemical cells. Types of Electrodes 1. Metal/Metal ion electrode (M/Mn+) When a metal is immersed in its own salt solution such type of electrode is called metal/metal ion electrode is formed. Eg: Zn/Zn2+, Cu/ Cu2+ In general M/Mn+ The electrode reaction is Mn+ + ne−→ M Electrode potential depends on the logarithmic concentration of metal ion. 2. Gas Electrode The gas is bubbled around an inert metal like platinum which is immersed in a solution containing ions of the same gas. Eg: Hydrogen electrode or SHE, chlorine electrode. Electrode potential depends on the logarithmic concentration of the ion and the pressure at which the gas is bubbled. 3. Metal insoluble metal salt/ common ion electrode The metal (M) is covered with an insoluble metal salt (MX) is in contact with a solution containing an anion of the insoluble metal salt(X). Eg: Calomel electrode (Hg/Hg2Cl2/Cl−), silver-silver chloride electrode (Ag/AgCl2/Cl−) Electrode potential is the logarithmic concentration of the common ion (Cl−) 4. Redox electrode In this electrode, an inert metal is immersed in a solution containing ions of the same metal having two different oxidation states. Eg: Pt/Fe2+/Fe3+ Reaction: Fe2+→Fe3+ + e- Electrode potential is the rof logarithmic concentration of all ionic species involved in the redox reaction. 5. Ion selective electrode 4 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. Single electrode potential (Reduction potential): Electrode potential is defined as the potential developed on the electrode at the interface when it is in contact with a solution of its own ions. It is denoted by E. All single electrode potentials represent reduction potentials which is the measure of the tendency of an electrode to undergo reduction. Origin of electrode potential: When a metal is in contact with a solution of its own ions, two types of reactions are possible. The metal shows the tendency to go into the solution as metal ion by losing electrons. M → Mn+ + ne- (oxidation) At the same time, metal ions in the solution show the tendency to get deposited as metal atoms. Mn+ + ne-→ M (reduction) These two opposite tendencies will result in an equilibrium represented by the equation. Mn+ + ne- M If the equilibrium lies over to the left, i.e the dissolution reaction is faster than the deposition the net change when equilibrium is attained, is a few metal atoms have dissolved in the solution liberating electrons. These electrons accumulate on the electrode surface, making it negatively charged. The negatively charged electrode surface attracts a layer of positively charged ions at the interface, developing an electrical double layer at the metal solution interface (Fig.2a). Similarly, if the deposition reaction is faster than the dissolution, at equilibrium the net change is deposition of few metal ions as metal ions by consuming electrons. As a result, the electrode surface develops a layer of positive charges which attracts a layer of negatively charged ions at the interface. (Fig.2b), again establishing an electrical double layer. The double layer is called Helmholtz Electrical Double Layer (HDL). The potential difference across the HDL is cause for the electrode potential. 5 Fig.2. Single electrode potential (a) De-electronation (b) Electronation Reference electrodes Any electrode of constant and reproducible potential used to determine the potential of other electrode. E.g.: Standard hydrogen electrode. Measurement of single electrode potential: The potential of a given electrode is measured using standard hydrogen electrode (SHE) whose potential is arbitrarily taken as zero volt at all temperatures and is the reference point for all potential measurements. Note: It is impossible to measure the absolute value of single electrode potential. It is because a single electrode constitutes only a half cell. A half-cell, not be loss or gain electrons by itself. The loss or gain of electrons can take place only in a complete circuit containing two half cells connected to each other. So, the electrode potential can only be measured by using some electrode as reference electrode. Secondary reference electrodes: Because of the difficulties involved in the use of standard hydrogen electrode as reference electrode, some other electrodes constant electrode potential, referred to as secondary reference electrodes. Two such electrodes are (i) calomel electrode (ii) silver-silver electrode. (i) Calomel electrode (mercury-mercurous electrode): Construction: Mercury is placed at the bottom of glass tube. A paste of calomel (Hg2Cl2) and mercury is placed over the pool of mercury. The remaining part of the tube is filled either a saturated or standard (normal or decinormal) solution of potassium chloride. A platinum wire is fused in a glass tube is dipped into mercury is used for electrical contact. A salt bridge is used to 6 couple with other half-cells (Fig. 3). The calomel electrode is represented as Hg | Hg2Cl2 (s) | KCl (saturated or standard) or Hg, Hg2Cl2 (s) , KCl (saturated or standard). Fig. 3. Calomel electrode Working: Calomel electrode can act as anode or cathode depending on the nature of the other electrode of the cell. The electrode reactions are represented as follows. As anode: As cathode: 2Hg → Hg22+ + 2e- Hg2Cl2 → Hg22+ + 2Cl- Hg22+ + 2Cl- → Hg2Cl2 Hg22+ + 2e- → 2Hg 2Hg + 2Cl- → Hg2Cl2 + 2e- Hg2Cl2 + 2e- → 2Hg + 2Cl- The net reversible electrode reaction is 1 𝐻𝑔2 𝐶𝑙2 + 2𝑒 − ⇋ 2𝐻𝑔 + 2𝐶𝑙 − or 2 𝐻𝑔2 𝐶𝑙2 + 𝑒 − ⇋ 𝐻𝑔 + 𝐶𝑙 − 𝟐.𝟑𝟎𝟑 𝑹𝑻 Electrode potential E = 𝑬𝟎 − 𝐥𝐨𝐠[𝑪𝒍− ] 𝑭 7  At 298K, Ecal = E0cal – 0.0591 log [Cl-] Since the calomel electrode is reversible with respect to chloride ion its electrode potential depends on the concentration of KCl solution. At 298 K, the electrode potentials as follows: KCl Concentration 0.1 N 1N Saturated Electrode potential (V) 0.334 0.281 0.2422 Advantages of calomel electrode (a secondary reference electrode): 1. It is very simple in construction. 2. The potential is reproducible and stable over a long period. 3. Its electrode potential will not vary with temperature. Hence, it is commonly used as secondary reference electrode for potential measurements. Determination of single electrode potential using calomel electrode The test electrode (e.g., Zn2+/Zn) is coupled with a saturated calomel electrode (Fig.4). The cell, so formed, may be represented as Zn(s) | ZnSO4 (1M) || KCl (Saturated solution) | Hg2Cl2 (s) | Hg(l) Fig. 4. Determination of electrode potential of zinc electrode using calomel electrode as reference electrode. The calomel electrode is written on the right as reduction takes place at this electrode. The emf of this cell measured potentiometrically. Then Ecell = Ecathode – Eanode = Ecal - EZn2+/ Zn 8  = 0.2422 - EZn2+/ Zn i.e., EZn2+/ Zn = 0.2422 - Ecell Concentration cells: A concentration cell is an electrochemical cell made of two-half cells having identical electrodes, identical electrolyte, except that the concentrations of electrolyte solutions are different (Figure 5). Two half cells may be joined by a salt bridge. Eg: Fig. 5. Concentration cell The cell is represented by - + Ag | AgNO3 (C1M) || AgNO3 (C2M) | Ag ; (C2 > C1). The Ag electrode which is in contact with lower concentration (C1) of the electrolyte acts as anode where oxidation occurs and enters the solution. The Ag electrode which is in contact with higher concentration (C2) of the electrolyte acts as cathode where reduction occurs, and deposition of Ag occurs. Electrode reactions: Anode: Ag (s) → Ag+(C1M) + e- Cathode: Ag+(C2M) + e- → Ag (s) Cell reaction: Ag+(C2M) → Ag+(C1M) There is no net chemical reaction but only concentration change takes place. Evidently, the emf so developed is due to the mere transference of metal ions from the solution of higher concentration (C2) to the solution of lower concentration (C1). Emf of the cell (Ecell) = Ecathode - Eanode 𝑅𝑇 According to Nernst Eq, E = Eº + [𝑛𝐹] ln [Mn+] 𝑅𝑇 𝑅𝑇 Ecell = [Eº + [𝑛𝐹] ln C2] - [Eº + [𝑛𝐹] ln C1] 9  𝑅𝑇 𝐶2 Ecell = [𝑛𝐹] ln 𝐶1 …………(1) At 298 K, 0.0591 𝐶2 Ecell = [ ] log 𝐶1 …………(2) where, C2> C1 𝑛 Note: (i) When C1 = C2 (concentrations are equal), log [C2 /C1] = 0 and hence no electricity flows (since emf = 0). (ii) When [C2 /C1] > 1 (i.e., C2 > C1), log [C2 /C1] is positive. Thus, the direction of spontaneous reaction is from the more concentrated solution (C2) to less concentrated solution (C1). (iii) Higher the ratio C2/C1, higher is the value of cell potential. For Eg. if the ratio C2/C1 increases from 0.001 to 0.1 (i.e., 100 times) the voltage increases three-fold. Ion-selective electrodes: Ion-selective electrode is the one which selectively responds to a specific ion in a mixture and potential developed at the electrode is a function of concentration of that ion in the solution. The electrode generally consists of a membrane which can exchange the specific ions with the solution with which it is in contact. Therefore, these electrodes are also referred to as membrane electrodes. Applications: Ion selective electrodes are used in determining The concentration of several cations such as H+, Li+, Na+, K+, Ag+, NH4+, Cd2+, Pb2+, Cu2+, Ca2+ and hardness [Ca2+ + Mg2+]; The concentration of anions such as F-, NO3-, CN-, S2- and halide ions; The pH of a solution using glass electrode (H+ ion-selective electrode) and The concentration of gas using gas sensing electrodes. For example, an electrode which measures the level of CO2 in blood samples makes use of glass electrode in contact with a very thin CO2 permeable silicone rubber membrane soaked in a dilute solution of sodium bicarbonate. Glass electrode 10 Construction: A glass electrode (Fig.6.) consists of a long glass tube with a thin-walled glass bulb at one end. Special glass (Corning glass containing 22% Na2O, 6% CaO and 72%SiO2) of low melting point and high electrical conductance is used for the purpose. This glass can specifically sense hydrogen ions up to a pH of about 9. The bulb contains 0.1 M HCl and a Ag/AgCl electrode (as internal reference electrode) is immersed into the solution and connected by a platinum wire for electrical contact. The electrode is represented as, Ag | AgCl (s) | HCl (0.1M) | glass. Fig. 6. Glass electrode Electrode potential of glass electrode (Eg): Principle: If a thin-walled glass bulb containing an acid is immersed in another solution, the membrane undergoes an ion exchange reaction; Na+ ions on the glass are exchanged for H+ ions. The potential is developed across the glass membrane (Fig.7). The potential difference, Eb at the interface also referred to as the boundary potential is the result of difference in potential (E2-E1) developed across the gel layer of glass membrane between the two liquids. 11  Fig. 7. Glass membrane Eb can be related to the difference in the hydrogen ion concentration of the two solutions by the 0.0591 𝐶2 relation, Eb = E2 – E1 = 𝑛 𝑙𝑜𝑔 𝐶1 … … … … … … … … … … … …. (1) Where C1 is the concentration of H+ ions of acid solution inside the glass bulb and C2 is the concentration of the acid solution into which the glass bulb is dipped. 𝑅𝑇 𝑅𝑇 𝐸𝑏 = − ln 𝐶1 + ln 𝐶2 … … … … … … (2) 𝑛𝐹 𝑛𝐹 If the concentration C1 of the solution inside the glass bulb is constant, then the first term on the R.H.S. of equation (2) is constant. Therefore 𝑅𝑇 𝐸𝑏 = 𝐶𝑜𝑛𝑠𝑡𝑎𝑛𝑡 + ln 𝐶2 𝑛𝐹 Substituting the value of R and F at 298K 𝐸𝑏 = 𝐿 + 0.0591 log 𝐶2 … … … … ….. (3) Where L is a constant which depends primarily on the pH of the solution taken in the bulb and glass electrode assembly. Since C2 = [H+] of the solution, Equation (3) written as, 𝐸𝑏 = 𝐿 − 0.0591 𝑝𝐻 … … … … ….. (4) The boundary potential established is mainly responsible for the glass electrode potential Eg and is given by 𝐸𝑔 = 𝐸𝑏 + 𝐸𝐴𝑔/𝐴𝑔𝐶𝑙 … … … … ….. (5) From equation (1), Eb = 0 when C1 = C2. But in practice, it has been observed that even when C1 = C2, a small potential is developed. This is called asymmetric potential (Easy). Hence, Equation (5) can be written as, 𝐸𝑔 = 𝐸𝑏 + 𝐸𝐴𝑔/𝐴𝑔𝐶𝑙 + 𝐸𝑎𝑠𝑦 𝐸𝑔 = 𝐿 − 0.0591 𝑝𝐻 + 𝐸𝐴𝑔/𝐴𝑔𝐶𝑙 + 𝐸𝑎𝑠𝑦 𝐸𝑔 = 𝐸𝑔0 − 0.0591 𝑝𝐻 … … … … … ….. (6) 12  𝑊ℎ𝑒𝑟𝑒 𝐸𝑔0 𝑖𝑠 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 𝑒𝑞𝑢𝑎𝑙 𝑡𝑜 𝐿 + 𝐸𝐴𝑔/𝐴𝑔𝐶𝑙 + 𝐸𝑎𝑠𝑦 Determination of pH using glass electrode: The glass electrode is immersed in the solution, the pH of which is to be determined. It is combined with a reference electrode such as a calomel electrode through a salt bridge (Fig.8). The cell assembly is represented as, Hg(l) | Hg2Cl2 (s) | KCl (Saturated solution) || Solution of unknown pH | glass | HCl (0.1M) | AgCl (s) | Ag (s) Fig. 8. Determination of pH using glass electrode The emf of the above cell, Ecell is measured using an electronic voltmeter with a null type potentiometer circuit (or a pH meter) (An ordinary potentiometer with a low resistance of the galvanometer is not used because of high internal resistance of the glass electrode). Ecell measured is the difference between the calomel electrode (Ecal) and Eg. Ecell = Eg – Ecal Ecell = [Eg° - 0.0591 pH] – Ecal 𝐸𝑔0 − 𝐸𝑐𝑎𝑙 − 𝐸𝑐𝑒𝑙𝑙 Therefore, pH = 0.0591 Advantages of glass electrode: (i) Glass electrode can be employed in the presence of strong oxidizing or reducing substances and metal ions. (ii) It is not poisoned easily. 13 (iii) Accurate results are obtained between pH range 1-9. However, by using special glass electrodes, pH 1-13 can be measured. (iv) It is simple to operated, can be used in portable instruments and therefore extensively used in chemical, industrial, agricultural and biological laboratories. Limitations of glass electrode: (i) In strongly acidic solutions of pH