Electrolysis Notes PDF

WEEK FIVE; LESSON ONE ELECTROLYSIS  ELECTROLYSIS Electrolysis is the decomposition of a molten or aqueous ionic compound (an electrolyte) by passing an electric current through it. The solution must contain a cathode (negative electrode) and an anode (positive electro...

WEEK FIVE; LESSON ONE ELECTROLYSIS  ELECTROLYSIS Electrolysis is the decomposition of a molten or aqueous ionic compound (an electrolyte) by passing an electric current through it. The solution must contain a cathode (negative electrode) and an anode (positive electrode). Electrolyte is a compound either in aqueous solution or molten form which conducts electricity and is decomposed in the process. Electrolytes conduct electric current by movement of ions.  ELECTROLYSIS Electrolytes vary in their ability to conduct electricity. Strong electrolytes like strong acids and salts conduct electricity readily since they are composed entirely of free moving ions. Weak electrolytes like weak acids ionize only slightly and so do not conduct electricity readily. Non-electrolytes do not conduct electricity because they do not ionize. They are covalent and mainly organic compounds like urea, ethanol, benzene, trichloromethane, ether and tetrachloromethane.  ELECTROLYSIS Electrodes are conductors in the form of wires, rods or plates through which an electric current enters or leaves the electrolyte. Anode is the positive electrode by which conventional current enters the electrolyte or by which electrons leave an electrolyte. It is the electrode which is joined to the positive terminal of the electric supply. Cathode is the negative electrode by which conventional current leaves the electrolyte or by which electrons enter the electrolyte. It is the electrode which is joined to the negative terminal of the electric supply.  Faraday’s First Law of Electrolysis Faraday’s first law of electrolysis states that the amount or mass of substance produced or liberated at each electrode is directly proportional to the quantity of charge flowing through the cell.  Faraday’s Second Law of Electrolysis Faraday’s second law of electrolysis states that when the same quantity of electricity is passed through different electrolytes, the relative number of moles of the elements discharged are inversely proportional to the charges on the ions of the elements.  Electrolytic Reactions Electrolytic reactions are redox reactions since they involve the transfer of electrons. Oxidation occurs at the anode where the anions lose electrons. This reaction at the anode is known as anodic half-reaction. The cathodic half-reaction which takes place simultaneously at the cathode is a reduction reaction since the cations gain electrons. The overall reaction is obtained by the algebraic addition of the two half-reactions.  Preferential Discharge of Ions Certain products are discharged at each electrode during electrolysis. In general, metals or hydrogen are discharged at the cathode while non-metals (except hydrogen) are discharged at the anode. The product formed at the electrode depend on the nature or state of the electrolyte. Where the electrolyte is a solution, the products formed may vary because the solvent, which is usually water, will also ionize. The cations and anions of both the electrolyte and the solvent will migrate to the cathode and the anode respectively where they will compete with one another to be discharged.  Preferential Discharge of Ions The product which is formed at the electrode will depend on which ions are preferentially discharged-the ions from the electrolyte or those from the solvent. The discharge of the ions is governed by three conditions, namely, 1. The position of the ions in the electrochemical series; 2. The concentration of the ions in the electrolyte 3. The nature of the electrode.  Uses of Electrolysis Electrolysis is of great importance in industry. Some of the applications of electrolysis include: 1. Extraction of elements; metals such as Na, K, Mg, Ca, Al, Zn and non-metals such as H2, F2, Cl2, are obtained either by electrolysis of their fused compounds or their aqueous solutions 2. Electroplating of one metal by another; coating the surface of one metal with another metal, usually copper, silver, chromium, nickel or gold, by means of electrolysis, for decoration or protection against corrosion. 3. Purification of metals such as Cu, Hg, Ag, Au 4. Preparation of certain important compounds such as sodium hydroxide and sodium trioxochlorate (v).  ELECTROCHEMICAL CELLS An electrochemical cell is a system consisting of electrodes that dip into an electrolyte and in which a chemical reaction either uses or generates an electric current. An electrochemical cell consists of two half-cells that are electrically connected. Each half-cell is the portion of an electrochemical cell in which a half-reaction takes place. The first half-reaction, in which a species loses electrons, is the oxidation half-reaction. The electrode at which oxidation occurs is called the anode. The second half- reaction, in which a species gains electrons, is the reduction half-reaction. The electrode at which reduction occurs is called the cathode. The sum of the two half-reactions is the net reaction that occurs in the voltaic cell; it is called the cell reaction.  ELECTROCHEMICAL CELLS Electrochemical cells are of two types: voltaic and electrolytic. In a voltaic cell, a spontaneous reaction generates electricity and does work on the surroundings. In an electrolytic cell, the surroundings supply electricity that does work to drive a nonspontaneous reaction. In both types of cell, two electrodes dip into electrolyte solutions; oxidation occurs at the anode, and reduction occurs at the cathode. The fundamental difference between the two types of electrochemical cells is based on whether the overall redox reaction in the cell is spontaneous (free energy is released) or nonspontaneous (free energy is absorbed)  ELECTROCHEMICAL CELLS A voltaic, or galvanic, cell is an electrochemical cell in which a spontaneous reaction generates an electric current. An electrolytic cell is an electrochemical cell in which an electric current drives an otherwise nonspontaneous reaction. Voltaic cells are used commercially as portable energy sources (batteries). In addition, the basic principle of the voltaic cell is employed in the cathodic protection of buried pipelines and tanks.  ELECTROCHEMICAL CELLS Electrolytic cells represent another type of electrochemical cell. They use an external voltage source to push a reaction in a nonspontaneous direction. The electrolysis of an aqueous solution often involves the oxidation or reduction of water at the electrodes. Electrolysis of concentrated sodium chloride solution, for example, gives hydrogen at the cathode. The amounts of substances released at an electrode are related to the amount of charge passed through the cell. This relationship is stoichiometric and follows from the electrode reactions.  ELECTROCHEMICAL CELLS A cell potential is a measure of the driving force of the cell reaction. This reaction occurs in the cell as separate half-reactions: an oxidation half- reaction and a reduction half-reaction. The electrode potential is an intensive property. This means that its value is independent of the amount of species in the reaction. Thus, the electrode potential for the half-reaction  ELECTROCHEMICAL CELLS The standard cell potential, Ecell, is the emf of a voltaic cell operating under standard-state conditions (solute concentrations are each 1 M, gas pressures are each 1 atm, and the temperature has a specified value usually 25˚C. The standard electrode potential, Ecell, is the electrode potential when the concentrations of solutes are 1 M, the gas pressures are 1 atm, and the temperature has a specified value (usually 25˚C. The strongest oxidizing agents in a table of standard electrode potentials are the oxidized species corresponding to half-reactions with the largest (most positive) E˚ values. The strongest reducing agents in a table of standard electrode potentials are the reduced species corresponding to half-reactions with the smallest (most negative) E˚ values.  ELECTROCHEMICAL CELLS The electrode with the higher electrical potential to give up its electrons is designated the anode, and the other electrode is the cathode. The electrical energy the cell produces is proportional to the difference in electrical potential between the two electrodes, which is called the cell potential (Ecell); it is also called the voltage of the cell or the electromotive force (emf). Electrons flow spontaneously from the negative to the positive electrode, that is, toward the electrode with the more positive electrical potential (anode to cathode). Thus, when the cell operates spontaneously, there is a positive cell potential: Ecell > 0 for a spontaneous process.  ELECTROCHEMICAL CELLS A positive Ecell arises from a spontaneous reaction. The more positive it is, the more work the cell can do, and the farther the reaction proceeds to the right as written. A negative Ecell is associated with a nonspontaneous cell reaction. If Ecell = 0, the reaction has reached equilibrium and the cell can do no more work 𝐸𝑐𝑒𝑙𝑙 = 𝐸𝑐𝑎𝑡ℎ𝑜𝑑𝑒 − 𝐸𝑎𝑛𝑜𝑑𝑒  ELECTROCHEMICAL CELLS When calculating cell potential values, chemists employ two approaches, both of which rely on using standard reduction potential values. One approach requires that you combine half-reactions and their standard potential values (reduction and oxidation); the other involves using an equation to calculate the difference between standard reduction potentials.  ELECTROCHEMICAL CELLS The measurement of cell potentials gives us yet another way to obtain equilibrium constants. Combining the previous equation, ˚ ΔG = −nFEcell with the equation ΔG = -RT ln K nFE˚ = RT ln K or ˚ RT 2.303RT Ecell = ln K = log K nF nF Substituting values for R and F at 25˚C gives the equation ˚ 0.0592 Ecell = logK (values in volts at 25˚C) n  ELECTROCHEMICAL CELLS The figure below summarizes the various relationships among K, ΔG˚, and E˚cell.  ELECTROCHEMICAL CELLS The cell potential of a cell depends on the concentrations of ions and on gas pressures. For that reason, cell potentials provide a way to measure ion concentrations. We can relate cell potentials for various concentrations of ions and various gas pressures to standard electrode potentials by means of an equation first derived by the German chemist Walther Nernst (1864–1941). Nernst Equation The free-energy change, ΔG, is related to the standard free-energy change, ΔG˚, by the following equation ΔG = ΔG˚ + RTlnQ where Q is the thermodynamic reaction quotient.  ELECTROCHEMICAL CELLS The reaction quotient has the form of the equilibrium constant, except that the concentrations and gas pressures are those that exist in a reaction mixture at a given instant. We can apply this equation to a voltaic cell. In that case, the concentrations and gas pressures are those that exist in the cell at a particular instant. If we substitute ΔG = -nFE˚cell and ΔG˚= -nFE˚cell into this equation, we obtain -nFEcell = -nFE˚cell +RT ln Q  ELECTROCHEMICAL CELLS This result rearranges to give the Nernt equation, an equation relating the cell potential to its standard potential and the reaction quotient. ˚ RT Ecell = Ecell − lnQ nF ˚ 2.303RT Ecell = Ecell − logQ nF If we substitute 298K (25˚C) for the temperature in the Nernst equation substitute in values for R and F, we get ˚ 0.0592 Ecell = Ecell − logQ n (values in volts at 25˚C)  ELECTROCHEMICAL CELLS For nonstandard conditions, the Nernst equation shows that Ecell depends on E˚cell and a correction term based on Q. Ecell is high when Q is small (high [reactant]), and it decreases as the cell operates. At equilibrium, ΔG and Ecell are zero, which means that Q = K.  ELECTROCHEMICAL CELLS Equilibrium Constants from Cell Potentials Some of the most important results of electrochemistry are the relationships among cell potential, free-energy change, and equilibrium constant. The free energy change ΔG for a reaction equals the maximum useful work of the reaction ΔG = wmax For a voltaic cell, this work is the electrical work, -nFEcell (where n is the number of moles of electrons transferred in a reaction), so when the reactants and products are in their standard states, we have ˚ ΔG = −nFEcell With this equation, cell potential measurements become an important source of thermodynamic information. WEEK FIVE; LESSON THREE FUEL CELLS AND BATTERIES  FUEL CELLS Fuel cells (sometimes called a flow battery) are like batteries; the key difference is that a battery is self- contained, while in a fuel cell the reactants need to be constantly replenished from an external source. With use, normal batteries lose their ability to generate voltage because the reactants become depleted as electrical current is drawn from the battery. In a fuel cell, the reactants-the fuel provided from an external source-constantly flow through the battery, generating electrical current as they undergo a redox reaction A fuel cell is essentially a battery, but it differs in operating with a continuous supply of energetic reactants, or fuel.  FUEL CELLS In a direct reaction between hydrogen and oxygen, oxygen atoms gains directly from hydrogen atoms. In hydrogen-oxygen fuel cells, the same redox reactions, but the hydrogen and oxygen are separated, forcing electrons to travel through an external wire to get from hydrogen to oxygen. These moving electrons constitute an electric current. Fuel cells employ the electron- gaining tendency of oxygen and electron-losing tendency of hydrogen to force electrons to move through a wire to create the electricity that provide power for a home or an electric automobile.  FUEL CELLS The most common fuel cell is the hydrogen-oxygen fuel cell shown below. In this cell, hydrogen gas flows past the anode (a screen coated with platinum catalyst) and undergoes oxidation: Oxidation (Anode): 2H2(g) + 4OH-(aq) → 4H2O (l) + 4e- Oxygen gas flows past the cathode (a similar screen) and undergoes reduction: Reduction (Cathode): O2(g) + 2H2O(l) + 4 e-→4 OH-(aq) The half-reactions sum to the following overall reaction: Overall reaction: 2H2(g) + O2(g) → 2H2O(l)  FUEL CELLS The half-reactions sum to the following overall reaction: Notice that the only product is water. In the space shuttle program, hydrogen-oxygen fuel cells consume hydrogen to provide electricity and astronauts drink the water that is produced by the reaction. In order for hydrogen-powered fuel cells to become more widely used, a more readily available source of hydrogen must be developed.  FUEL CELLS A type of fuel cell being developed for use in cars is the proton exchange membrane (PEM) cell, which uses H2 as the fuel and has an operating temperature of around 80˚C shown below. The cell reactions are: On one side of the cell, the anode, hydrogen passes through a porous material containing a platinum catalyst, allowing the following reactions to occur: + H2(g) → 2H(aq) + 2e− (anode)  FUEL CELLS + The H(aq) ions then migrate through a proton- exchange membrane to the other side of the cell to participate in the cathode reaction with O2(g) : + O2(g) + 4H(aq) + 4e− → 2H2 O(l) (cathode) The sum of the half-reactions (note that the cathode reaction must be multiplied by 2 prior to adding) is H2(g) + O2(g) → 2H2 O(l)  FUEL CELLS which is the net reaction in the fuel cell. The first applications of PEM fuel cells were in space, but more recently, they have provided for lighting, emergency power generators, communications equipment, automobiles and buses. Other types of cells using materials and fuels such as hydrocarbons or methanol are either in commercial or under development.  RECHARGEABLE BATTERIES In general, a battery consists of self-contained voltaic cells arranged in series (plus-to-minus-to- plus, and so on), so that their individual voltages are added. Batteries are voltaic cells arranged in series and are classified as primary (e.g., alkaline, mercury, and silver), secondary (e.g., lead-acid, nickel– metal hydride, and lithium-ion), or fuel cells. Supplying electricity to a rechargeable (secondary) battery reverses the redox reaction, re-forming reactant.  RECHARGEABLE BATTERIES Secondary (Rechargeable) Batteries Some types of cells are rechargeable after use, however. An important example is the lead storage cell. This voltaic cell consists of electrodes of lead alloy grids; one electrode is packed with a spongy lead to form the anode, and the other electrode is packed with lead dioxide to form the cathode. Both are bathed in an aqueous solution of sulfuric acid, H2SO4. The half- cell reactions during discharge are White lead(II) sulfate coats each electrode during discharge, and sulfuric acid is consumed.  RECHARGEABLE BATTERIES Each cell delivers about 2 V, and a battery consisting of six cells in series gives about 12V.  RECHARGEABLE BATTERIES Both the anode and the cathode are immersed in aqueous sulfuric acid (H2S04 (aq)). As electrical current is drawn from the battery, both electrodes become coated with PbSO4. If the battery is run for a long time without recharging, too much PbSO4(s) develops on the surface of the electrodes and the battery goes dead. The lead-acid storage battery can be recharged by an electrical current (which must come from an external source such as an alternator in a car). The current causes the preceding reaction to occur in reverse, converting the PbSO4 back to Pb(s) and PbO2(s).  RECHARGEABLE BATTERIES After the lead storage battery is discharged, it is recharged from an external electric current. The previous half-reactions are reversed. Some water is decomposed into hydrogen and oxygen gas during this recharging, so more water may have to be added at intervals. However, newer batteries use lead electrodes containing some calcium metal; the calcium–lead alloy resists the decomposition of water. These maintenance free batteries are sealed.  RECHARGEABLE BATTERIES Lead-Acid Storage Battery: a lead-acid storage consists of six cells wired in series. Each cell contains a porous lead anode and a lead oxide cathode, both immersed in sulphuric acid.  RECHARGEABLE BATTERIES The ubiquity of power electronic products such as laptops, cell phones, and digital cameras, as well as the growth in popularity of hybrid electric vehicles, has driven the need for efficient, long-lasting, rechargeable batteries. The most common types include the nickel- cadmium (NiCad) battery, the nickel-metal hydride (NiMH) battery, and the lithium-ion battery. The Nickel- Cadmium (NiCad) Battery Nickel-cadmium batteries consist of an anode composed of solid cadmium and a cathode composed of NiO(OH)(s). The electrolyte is usually KOH(aq). During operation, the cadmium is oxidized and the NiO(OH) is reduced according to the equations:  RECHARGEABLE BATTERIES The nickel–cadmium cell (nicad cell) is a common storage battery. It is a voltaic cell consisting of an anode of cadmium and a cathode of hydrated nickel oxide (approximately NiOOH) on nickel; the electrolyte is potassium hydroxide. Nicad batteries are used in calculators, portable power tools, shavers, and toothbrushes.  RECHARGEABLE BATTERIES Recently, nickel metal hydride (NiMH) and lithium batteries have been replacing nicad batteries in many applications that require rechargeable batteries. The NiMH batteries use a metal hydride for the anode instead of the more toxic cadmium. The NiMH battery employs the same cathode reaction as the NiCad battery but a different anode reaction. In the anode of a NiMH battery, hydrogen atoms held in a metal alloy are oxidized. If we let M represent the metal alloy, we can write the half-reactions as follows:  RECHARGEABLE BATTERIES NiHM batteries are currently used to supply electric power for hybrid-electric vehicles. Lithium batteries typically use a carbon anode and lithium cobalt dioxide or a lithium manganese compound as the cathode. Lithium batteries are commonly used to power consumer electronics. In addition to being more environmentally friendly than NiCad batteries, NiMH batteries also have a greater energy density (energy content per unit battery mass), as we can see in Table 18.2.  RECHARGEABLE BATTERIES In some cases, a NiMH battery can carry twice the energy of a NiCad battery of the same mass, making NiMH batteries the most common choice for hybrid electric vehicles. The Lithium-Ion Battery The most common type of rechargeable battery is the lithium-ion battery. Since lithium is the least dense metal (0.53 g cm- 3) , lithium batteries have high energy densities (see Table 18.2). The lithium battery works differently than the other batteries we have examined so far. Operation of the lithium battery is due primarily to the motion of lithium ions from the anode to the cathode.  RECHARGEABLE BATTERIES The anode is composed of graphite into which lithium ions are incorporated between layers of carbon atoms. Upon discharge, the lithium ions spontaneously migrate to the cathode, which consists of a lithium transition-metal oxide such as LiCoO2 or LiMn204. The transition metal is reduced during this process. Upon recharging, the transition metal is oxidized, forcing the lithium to migrate back into the graphite. The flow of lithium ions from the anode to the cathode causes a corresponding flow of electrons in the external circuit.  RECHARGEABLE BATTERIES Lithium-ion batteries are commonly used in applications where light weight and high energy density are important. These include cell phones, laptop computers, and digital cameras. Its key drawbacks are cost and flammability of the organic solvent.