Electrochemistry Part-2 PDF

Electrolytic Cell An electrolytic cell can be defined as an electrochemical device that uses electrical energy to facilitate a non-spontaneous redox reaction. Electrolytic cells are electrochemical cells that can be used for the electrolysis of certain compounds. For example, water can be subje...

Electrolytic Cell An electrolytic cell can be defined as an electrochemical device that uses electrical energy to facilitate a non-spontaneous redox reaction. Electrolytic cells are electrochemical cells that can be used for the electrolysis of certain compounds. For example, water can be subjected to electrolysis (with the help of an electrolytic cell) to form gaseous oxygen and gaseous hydrogen. This is done by using the flow of electrons (into the reaction environment) to overcome the activation energy barrier of the non-spontaneous redox reaction. The three primary components of electrolytic cells are: 1. Cathode (which is negatively charged for electrolytic cells) 2. Anode (which is positively charged for electrolytic cells) 3. Electrolyte: The electrolyte provides the medium for the exchange of electrons between the cathode and the anode. Commonly used electrolytes in electrolytic cells include water (containing dissolved ions) and molten sodium chloride. Diagram and Working of an Electrolytic Cell Molten sodium chloride (NaCl) can be subjected to electrolysis with the help of an electrolytic cell, as illustrated below. Here, two inert electrodes are dipped into molten sodium chloride (which contains dissociated Na+ cations and Cl– anions). When an electric current is passed into the circuit, the cathode becomes rich in electrons and develops a negative charge. The positively charged sodium cations are now attracted towards the negatively charged cathode. This results in the formation of metallic sodium at the cathode. Simultaneously, the chlorine atoms are attracted to the positively charged cathode. This results in the formation of chlorine gas (Cl2) at the anode (which is accompanied by the liberation of 2 electrons, finishing the circuit). The associated chemical equations and the overall cell reaction are provided below. Reaction at Cathode: [Na+ + e- → Na] x 2 Reaction at Anode: 2Cl– → Cl2 + 2e- Cell Reaction: 2NaCl → 2Na + Cl2 Thus, molten sodium chloride can be subjected to electrolysis in an electrolytic cell to generate metallic sodium and chlorine gas as the products. Applications of Electrolytic Cells The primary application of electrolytic cells is for the production of oxygen gas and hydrogen gas from water. They are also used for the extraction of aluminium from bauxite. Another notable application of electrolytic cells is in electroplating, which is the process of forming a thin protective layer of a specific metal on the surface of another metal. The electrorefining of many non-ferrous metals is done with the help of electrolytic cells. Such electrochemical cells are also used in electrowinning processes. It can be noted that the industrial production of high-purity copper, high-purity zinc, and high-purity aluminium is almost always done through electrolytic cells. Process of electrolysis 1. The process of electrolysis is carried out by taking the solution of an electrolyte in a suitable vessel. The vessel is called electrolytic tank. 2. It is made up of either glass or of a material which is a bad conductor of electricity. 3. Two metallic rods or plates are suspended in the electrolytic solution. These are connected to the terminal of a battery with the help of metallic wires. 4. These metallic rods or plates allow the passage of current and are called electrodes. The electrode connected to the positive terminal of the battery is called anode while the electrode connected to the negative terminal of the battery is called cathode. When an electrolyte is dissolved in water, it splits up into negative and positive ions. The positively charged ions are called cations and negatively charged ions are called anions. On passing electric current through the solution, the ions are attracted by the oppositely charged electrodes. As a result, cations move towards cathode while anions move towards anode. This movement of ions in solution is known as electrolytic or ionic conduction and constitutes flow of current through the solution. The anions on reaching the anode give up their electrons. On the other hand, cations take up the electrons from the cathode Therefore, cations and anions get discharged at the respective electrodes and are converted to neutral particles. This is known as primary change. The primary products may be collected as such or they undergo further changes to form molecules or compounds. These are called secondary products and the change is known as secondary change. According to ionic theory, the electrolytes are present as ions in solution and the function of electricity is only to direct these ions to their respective electrodes. The electrolytes can be electrolysed only in the dissolved or molten state. Faraday’s Laws of Electrolysis (1)Faraday’s first law of electrolysis The amount of any substance deposited or liberated at any electrode is directly proportion to the quantity of electricity passed through the electrolytic solution. The amount of any sub stance obtained gives the amount of chemical reaction which occurs at any electrode during electrolysis. Thus, if w gram of the substance is deposited on passing Q coulombs of electricity, then w∝Q w = ZQ where Z in a constant of proportionality and is called electrochemical equivalent. If a current of I amperes is passed for t seconds, then Q=I×t W=Z×Q=Z×I×t If, Q = 1 coulomb I = 1 ampere and t = 1 second, then, w= Z × 1 × 1 w= Z Electrochemical equivalent of a substance may be defined as the mass of the substance deposited when a current of one ampere is passed for one second, i.e. a quantity of electricity equal to one coulomb is passed. 2) Faraday’s second law of electrolysis  It states that when same quantity of electricity is passed through different electrolytic solutions connected in series, the weights of the substances produced at the electrodes are directly proportional to their chemical equivalent weights. For example: When same current is passed through two electrolytic solutions, containing copper sulphate (CuSO4) and silver nitrate (AgNO3) connected in series, the weights of copper and silver deposited are : According to Faraday’s law, the amount of chemical change occurred i.e. the moles of substances deposited or liberated is proportional to the number of moles of electrons exchanged during the oxidation-reduction reactions that occur. By knowing the amount of electricity passed, we can easily calculate the number of moles of products formed from the appropriate electrode reaction. From the moles of the products formed, we can calculate the masses their volumes if they are gases. During the passage of electric current through molten NaCl, sodium gets deposited at cathode and chlorine is liberated at anode. NaCl ———-> Na + ½Cl2 During electrolysis, sodium ions move towards cathode, accept electrons and get deposited as: Na+ (aq) +e¯ ——–> Na(s) The passage of one electron produces one sodium atom. The passage of 1 mol of electrons produce 1 mol of sodium (or 23 g). Similarly, at anode, chloride ions give up electrons and produce Cl atoms as: Cl‾ – e¯ ———> ½ Cl2(g) 2 Cl‾ – 2e¯ ———> Cl2(g) 2 mol of electrons produce 1 mol of Cl2 or 1 mol of electrons produce 1/2 mol of Cl (35.5 g) Charge on an electron = 1.602 x 10-19 C Now, 1 mole of electrons= 6.022 x 1023 electrons  Charge on 1 mole of electrons = 6.022x1023×1.602 x10-19 C The charge on one mole of electrons is called 1 Faraday, F. Thus, 1F = 96485 C or approximately 96500 C. Thus, charge on n mol of electrons will be equal to Q= nF Now, the production of 1 mol of sodium or 23.0 g by reduction of sodium ions require 1 mol of electrons. Therefore, amount of charge required, Q= nF = 1 × 96500 C = 96500 C Similarly, 1 mol of Cl2 is obtained by 2 mol of electrons or 2 × 96500 C of charge during electrolysis of NaCl. Similarly, in the reaction Ag+ + e‾ ——–> Ag(s) One mole of electrons is required for the reduction of 1 mol of silver ions. Therefore, the quantity of electricity required for reduction of 1 mol of Ag+ ions is 96500 C or 1 Faraday. Now 1 mole of copper will be produced by 2 mol of electrons or 2 × 96500 C of charge: Cu2+ + 2e¯ —–> Cu The amount of substance deposited or evolved can be calculated. For example: aluminium gets deposited as: Al3+ +3 e¯ —–> Al Thus, 1 mol of Al will be deposited by 3 mol of electrons or 3 Faraday of electricity. When the same quantity of electricity is passed through different electrolyte solutions, connected in series, the weights of different substances produced at the electrodes can be calculated from the mole ratios of their electrode reactions. BATTERY A Battery is a device consisting of one or more electrical cells that convert chemical energy into electrical energy. Every battery is basically a galvanic cell where redox reactions take place between two electrodes which act as the source of the chemical energy. Battery types Based on the application of the battery, they can be classified again. They are: Household Batteries These are the types of batteries which are more likely to be known to the common man. They find uses in a wide range of household appliances (such as torches, clocks, and cameras). These batteries can be further classified into two subcategories: Rechargeable batteries Nickel. Examples: Cadmium batteries, Lithium-Ion Non-rechargeable batteries Examples: Silver oxide, Alkaline & carbon zinc Industrial Batteries These batteries are built to serve heavy-duty requirements. Some of their applications include railroad, backup power and more for big companies. Some examples are: Nickel Iron, Wet Nickel Cadmium (NiCd) Vehicle Batteries These are more user-friendly and a less complicated version of the industrial batteries. They are specifically designed to power cars, motorcycles, boats & other vehicles. An important example of a vehicle battery is the Lead-acid battery. Batteries can be broadly divided into two major types: Primary Cell / Primary battery Secondary Cell / Secondary battery Primary Cell These are batteries where the redox reactions proceed in only one direction. The reactants in these batteries are consumed after a certain period of time, rendering them dead. A primary battery cannot be used once the chemicals inside it are exhausted. An example of a primary battery is the dry cell – the household battery that commonly used to power TV remotes, clocks, and other devices. In such cells, a zinc container acts as the anode and a carbon rod acts as the cathode. A powdered mixture of manganese dioxide and carbon is placed around the cathode. The space left in between the container and the rod are filled with a moist paste of ammonium chloride and zinc chloride. The redox reaction that takes place in these cells is: At Anode Zn(s) → Zn2+ (aq) + 2e- At Cathode  2e- + 2 NH4+ (aq) → 2 NH3 (g) + H2 (g) 2 NH3 (g) +Zn2+ (aq) → [Zn (NH3)2] 2+ (aq) H2 (g) + 2 MnO2 (S) → Mn2O3 (S) + H2O (l) Thus, the overall cell equation is: Zn(s) + 2 NH4+ (aq) + 2 MnO2 (S) → [Zn(NH3)2] 2+ (aq) + Mn2O3 (S) + H2O (l) Another example of the primary cell is the mercury cell, where a zinc-mercury amalgam is used as an anode and carbon is used as a cathode. A paste of HgO is used as an electrolyte. These cells are used only in devices that require a relatively low supply of electric current (such as hearing aids and watches). Secondary Cell These are batteries that can be recharged after use by passing current through the electrodes in the opposite direction, i.e. from the negative terminal to the positive terminal. For example, a lead storage battery that is used in automobiles and inverters can be recharged a limited number of times. The lead storage battery consists of a lead anode and the cathode is a lead grid packed with lead dioxide. Sulphuric acid with a concentration of 38% is used as an electrolyte. The oxidation and reduction reactions involved in this process are listed below. At Anode  Pb → Pb2++ 2 e– Pb+ SO42– → PbSO4(electrode) + 2 e– At Cathode 2 e–+ PbO2+ 4 H+ → Pb2++ 2 H2O 2 e–+ PbO2+ 4 H++ SO42- → PbSO4(electrode) + 2 H2O In order to recharge these batteries, the charge is transferred in the opposite direction and the reaction is reversed, thus converting PbSO4 back to Pb and PbO2. Another example of the secondary cell is the nickel-cadmium cell. These cells have high storage capacities and their lifespan is relatively long (compared to other secondary cells). However, they are difficult to manufacture and maintain. FUEL CELL What is a Fuel cell? A fuel cell is a device that produces electric energy, through a chemical reaction. Characteristics of Fuel Cell: Fuel cells use a positively charged ion (Hydrogen) and an oxidizing agent (oxygen). There are many types of fuel cells, but they all consist of a cathode, an anode, and an electrolyte that allows positively charged (hydrogen) ions to move between the two sides of the fuel cell. They differ from batteries as they (fuel cells) require a continuous supply of fuel. Both batteries and fuel cells produce direct current (D.C). Working of a Fuel Cell The working of this fuel cell involved the passing of hydrogen and oxygen into a concentrated solution of sodium hydroxide via carbon electrodes. The cell reaction can be written as follows: Cathode Reaction: O2 + 2H2O + 4e– → 4OH– Anode Reaction: 2H2 + 4OH– → 4H2O + 4e– Net Cell Reaction: 2H2 + O2 → 2H2O However, the reaction rate of this electrochemical reaction is quite low. This issue is overcome with the help of a catalyst such as platinum or palladium. To increase the effective surface area, the catalyst is finely divided before being incorporated into the electrodes. CORROSION It is basically defined as a natural process that causes the transformation of pure metals into undesirable substances when they react with substances like water or air. This reaction causes damage and disintegration of the metal, starting from the portion of the metal exposed to the environment and spreading to the entire bulk of the metal. Corrosion is usually an undesirable phenomenon since it negatively affects the desirable properties of the metal. For example, iron is known to have good tensile strength and rigidity (especially alloyed with a few other elements). However, when subjected to rusting, iron objects become brittle, flaky, and structurally unsound. Factors Affecting Corrosion 1. Exposure of the metals to air containing gases like CO2, SO2, SO3 etc. 2. Exposure of metals to moisture, especially salt water (which increases the rate of corrosion). 3. Presence of impurities like salt (For example, NaCl). 4. Temperature: An increase in temperature increases corrosion. 5. Nature of the first layer of oxide formed: Some oxides like Al2O3 form an insoluble protecting layer that can prevent further corrosion. Others, like rust, easily crumble and expose the rest of the metal. 6. Presence of acid in the atmosphere: Acids can easily accelerate the process of corrosion. Corrosion Examples, Reactions and Effects 1. Copper Corrosion When copper metal is exposed to the environment, it reacts with the oxygen in the atmosphere to form copper (I) oxide, which is red in colour. 2Cu(s) + ½ O2(g) → Cu2O(s) Cu2O further gets oxidised to form CuO, which is black in colour. Cu2O(s) + ½ O2(g) → 2CuO(s) This CuO reacts with CO2, SO3 and H2O (present in the atmosphere to form Cu2(OH)2(s) (Malachite), which is blue in colour and Cu4SO4(OH)6(s) (Brochantite), which is green in colour. This is why we observe copper turning bluish-green in colour. A typical example of this is the colour of the Statue of Liberty, which has the copper coating on it turning blue-green in colour. 2. Silver Tarnishing Silver reacts with sulphur and sulphur compounds in the air, giving silver sulphide (Ag2S), which is black in colour. Exposed silver forms Ag2S as it reacts with the H2S(g) in the atmosphere, which is present due to certain industrial processes. 2Ag(s) + H2S(g) → Ag2S(s) + 2H+(g) 3. Corrosion of Iron (Rusting) Rusting of iron, which is the most commonly seen example, happens when iron comes in contact with air or water. The reaction could be seen as a typical electrochemical cell reaction. Consider the diagram given below. Here, metal iron loses electrons and gets converted to Fe2+ (this could be considered as the anode position). The electrons lost will move to the other side, where they combine with H+ ions. H+ ions are released either by H2O or by H2CO3 present in the atmosphere (this could be considered as the cathode position). The Hydrogen, thus formed by the reaction of H+ and electrons, react with oxygen to form H2O. Anode reaction 2Fe(s) → 2Fe2+ + 4e– ; E°Fe2+/Fe= - 0.44 V Cathode reaction O2 + 4H+ (aq) + 4e- H2O(l), E°H+/O2/H2= 1.23 V Overall reaction 2Fe(s) + O2(g) + 4H+(aq) → 2Fe2+(aq) + 2H2O(l) E°cell = 1.67V The Fe2+ ions formed at the anode react with oxygen in the atmosphere, thereby getting oxidised to Fe3+ and forming Fe2O3, which comes out in the hydrated form as Fe2O3.xH2O Fe2+ + 3O2 → 2Fe2O3 Fe2O3 + xH2O → Fe2O3. xH2O (rust) Prevention of Corrosion Preventing corrosion is of utmost importance in order to avoid huge losses. The majority of the structures that we see and use are made out of metals. This includes bridges, automobiles, machinery, household goods like window grills, doors, railway lines, etc. While this is a concerning issue, several treatments are used to slow or prevent corrosion damage to metallic objects. This is especially done to those materials that are frequently exposed to the weather, saltwater, acids, or other hostile environments. Some of the popular methods to prevent corrosion include,