Chapter 4: Electrolysis Calculations In Electrolysis – Faraday’s Laws (pages 123 to 128) • Faraday’s First law of Electrolysis: * the amount of any substance produced at an electrode during electrolysis is directly proportional to the quantity of electric charge, Q, passed through the cell ie, Q = It Q = electric charge – measured in coulombs (C) I = current – in amperes (A) t = time – in seconds (s) - since the charge on one electron is 1.602 X 10-19 C - the charge on one mole of electrons is (6.023 X 1023 x 1.602 X 10-19 C) = 9.649 X 104 C referred to the faraday (F) or Faraday constant (usually given as 96500 Cmol-1) • Faraday’s Second law of Electrolysis: * in order to produce 1 mole of substance by electrolysis, 1, 2, 3 or another whole number of moles of electrons (faradays) must be consumed, according to the relevant half-equation eg: Na+ + e- Na 1 faraday (1 mol of e-) must be passed to produce 1 mol of Na atoms Cu2+ + 2e- Cu 2 faraday (2 mol of e-) must be passed to produce 1 mol of Cu atoms Al3+ + 3e- Al 3 faraday (3 mol of e-) must be passed to produce 1 mol of Al atoms 2Cl- Cl2 + 2e2 faraday (2 mol of e-) must be passed to produce 1 mol of Cl2 molecules Q ie, Q = n(e-)xF or n(e-) = F • Applying Faraday’s Laws of Electrolysis: Q - use: Q = It and n(e-) = F Sample Problem 4.1 (page 125,126) A solution of copper(II) sulfate is electrolysed for 30.0 minutes using a current of 0.500 A. Calculate: (a) the mass of copper deposited at the cathode (b) the volume of oxygen gas (at SLC) evolved at the anode ……………………………………………………………………………………………………………. ……………………………………………………………………………………………………………. ……………………………………………………………………………………………………………. ……………………………………………………………………………………………………………. ……………………………………………………………………………………………………………. ……………………………………………………………………………………………………………. ……………………………………………………………………………………………………………. ……………………………………………………………………………………………………………. ……………………………………………………………………………………………………………. ……………………………………………………………………………………………………………. ……………………………………………………………………………………………………………. ……………………………………………………………………………………………………………. Q’s: 8,9,10,11(p.126); Review Q’s: 14,15,16,17,18,19,20,21,22,23,24,25,27,28,29(p.140) Sample Prob 4.2 (page 126,127) – Calculating the Charge on an Ion from Electrolysis Experiments Q’s: 12,13(p.127) • Faraday’s Laws In Industry: * Calculations based on Faraday’s Laws are critical to industrial electrolytic processes * To determine efficiency of process, knowledge of the theoretical maximum amount is required Sample Problem 4.3 (page 127,128) Q: 14(p.128); Review Q: 26(p.140) Secondary Cells – Rechargeable Batteries (pages 128 to 134) • often referred to as rechargeable batteries; can be recharged after they become ‘flat’ • are essentially galvanic cells using spontaneous redox reactions to produce electricity during discharge; then they become electrolytic cells converting electrical energy back into chemical energy when they are recharged – for this to happen they are designed so that the discharge products remain in contact at the electrodes • recharging involves connecting the negative terminal of the charger to the negative terminal of the battery, and the positive to the positive; electrons are forced to travel in the reverse direction and, because the products are still in contact with the electrodes, the original reactions are reversed – therefore recharging the cell or battery • are initially more expensive than primary cells, but ability to be recharged many hundreds of times makes them a cheaper long term alternative • The Lead-Acid Accumulator: - is the most common & durable secondary cell - developed in late 1880’s - relatively long lifetime, high current, cheap to produce - 12 V lead-acid storage battery - consists of six 2 V cells connected in series - largely used in transport applications; rely on vehicle’s generator or alternator supplying enough voltage to reverse spontaneous electrochemical reaction - each cell consists of: * lead electrodes or grids; grid structure provides a larger surface area for electrode reactions anode packed with spongy lead cathode packed with lead(IV) oxide, PbO2 * electrodes immersed in 4M H2SO4 * electrodes separated by porous separator * The Discharging Process in a Lead-Acid Accumulator: Pb(s) + SO42-(aq) PbSO4(s) + 2e- PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- PbSO4(s) + 2H2O(l) Overall: ……………………………………………………………………………………. - discharge produces electric power to start car; results from spontaneous redox reaction - at both anode & cathode lead(IV) sulfate, PbSO4, deposits on the grid - pH increases during discharge cycle * The Recharging Process in a Lead-Acid Accumulator: - since the products remain as a deposit on the electrodes, reactions can be reversed by passing a current through the cell in the opposite direction - while car in motion PbSO4(s) + 2e- Pb(s) + SO42-(aq) PbSO4(s) + 2H2O(l) PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- Overall: ……………………………………………………………………………………. - pH decreases during recharge cycle - not a spontaneous reaction, so direct current must be applied to make it proceed – achieved by the alternator (motor driven electrical source of higher voltage than the battery) which supplies 14 V (Note – eg of Electrolysis) - in theory battery can be recharged indefinitely; in practice lasts about 4 years; due to small amounts of lead(IV) sulfate, PbSO4, falling from the electrodes and dropping to the bottom – recharge process eventually becomes ineffective Q’s: 15,16(p.133); Review Q’s: 30,31(p.140) • The Nickel Metal Hydride (NiMH) Rechargeable Cell: - has largely replaced the first widely used rechargeable cell, the nickel cadmium (NiCd) - shares a number of features with the NiCd, but is environmentally safer due to the absence of cadmium - doesn’t have the NiCd problem of the so-called memory effect: if only partly discharged before being recharged, it would not receive a full charge. NiMH cells show less of this effect - have nearly 50% more charge/gram, can recharge faster & can run longer on each charge * similar to NiCd batteries, except a hydrogenabsorbing metal is used as the anode; these metals (M) can be alloys of lanthanum and rare earths * potassium hydroxide is the electrolyte * produce 1.2 volts - have many advantages, but also some disadvantages: they suffer from self-discharge, a problem that is worse at high temperatures; they require more complicate charging devices to prevent overcharging - used in laptops, electric shavers, toothbrushes, cameras, camcorders, mobile phones and medical instruments, and are being developed for car batteries - Electrode half-equations: MH(s) + OH-(aq) M(s) + H2O(l) + eM = metal NiO(OH)(s) + H2O(l) + e Ni(OH)2(s) + OH (aq) Overall: ………………………………………………………… Revision Q’s: 19,20(p.133) • Rechargeable Lithium Ion Cell: - uses lithium compounds rather than lithium metal - rapidly becoming the most popular battery choice for portable devices - consists of: *porous graphite anode; lithium is interspersed between the layers of graphite *porous metal oxide cathode; cobalt IV oxide containing some lithium ions stored in gaps of the lattice structure commonly used; As a result: cobalt is present in both +4 & +3 oxidation states * in an electrolyte of lithium salt held in an organic solvent - Equations: Li1-xCoO2 + xLi+ +xe- LiCoO2 CLix C + xLi+ + xeOverall: ………………………………………………… Q’s: 17,18(p.133); Review Q’s: 32,33(p.141) • Factors Affecting Battery Life: - while there are specialist batteries designed to operate at extreme temps, the majority of batteries used today are designed to operate over an optimum range of temps that include room temp; if used outside this range, breakdowns in both the electrochemistry & physical components may occur rendering it inoperative & even dangerous - battery specific problems: Lead-acid accumulators: - overdischarge thickens the layers of PbSO4 on the electrodes, this can break off & fall, lowering recharge amounts; short circuits … Lithium Ion Battery: - at low temps insertion of lithium ions into anode & cathode lattices spaces can become difficult as contraction reduces space size; high temps cause ‘thermal runaway’ - generating excessive heat & breakdown
