SCH 4U: UNIT 7 LESSONS ELECTROCHEMISTRY (Chap 5-pg & Chap 19-pg )

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1 SCH 4U: UNIT 7 LESSONS ELECTROCHEMISTRY (Chap 5-pg & Chap 19-pg ) 1. Rationale: Oxidation & Reduction reactions occur in many chemical systems. Examples include the rusting of iron, the displacement of metals and photosynthesis in plants. All of these reactions involve the transfer of electrons from one chemical species to another. Because electrons are involved in all these reactions, these changes are described as electrochemical changes and the study of these changes is called electrochemistry. 2. REVIEW a) Oxidation Reduction Reactions (from SCH3U) (5.1-pg 193) Done via lab activity 1 i) Mg + O2 ( burn a pc of Mg) pg 194 ii) Mg + HCl ( drop pc of Mg into HCl sol n test gas) pg 206 iii) Mg + Pb(NO3)2 ( drop pc Mg into sol n - observe) pg Write rx eq ns on board- go over oxidation of Mg in each case b) Activity Series (from SCH3U) what it is? / what it means? - pg Oxidation Numbers (5.1 pg ) 1. What are they? What are they used for? 2. Rules for assigning Oxidation Numbers i) The Oxidation # of all elements is 0 ii) The Oxidation # of any simple ion = charge iii) The Oxidation # of any compound = 0 iv) The Oxidation # of any compound ion = overall charge v) The Oxidation # of Hydrogen = +I, oxygen = -II 3. Do several examples ASSIGN Using Oxidation Numbers to balance REDOX rx equations (5.2 pg ) a) Do a few examples assign the rest to be done over next 2 weeks ASSIGN bonus Quiz: Balancing REDOX equations (2 weeks later)

2 5. Electrochemical Cells ( Galvanic Cells ) (19.4 pg ) DEMO: Zinc/Copper EC cell - Go over the principles of the 2 half rxs being separate - Hook up the volt meter ( no current/voltage!!) - Principles of salt bridge - see overhead Zn/Cu EC Cell 6. Cell Potentials of Electrochemical Cells ( E ) (19.5 pg 866) Define: Standard Oxidation Potential / Std Reduction Potential ( E ) * standard 1/2 cell H2 //H + cell Zinc 1/2 cell connected to std 1/2 cell Copper 1/2 cell connected to std 1/2 cell see overhead Zn & Std Hydrogen cell see overhead Zn & Std H + cell Formation of the Std Oxidation Potential/reduction Potential List see overhead Standard Reduction Potentials Reference the Senior Chemistry Data Sheet Std Oxidation Potentials Zinc 1/2 cell connected to the copper 1/2 cell Copper/Silver Cell & Nickel/Copper Cell **ASSIGN 7.4 see overhead Zn / Cu galvanic cell see overhead galvanic cells 7. Practical applications of EC Cells Batteries (19.9 pg 885) Lead storage battery Demo using car battery (see pg ) see overhead lead storage battery Zinc-carbon dry cell demo with cut up batteries (see pg ) see overhead zinc-carbon battery Nickel-cadmium batteries rechargeables (see pg ) Specialized batteries (see pg ) **ASSIGN 7.5 see overhead mercury battery

3 8. Using Cell Potentials of EC Cells to predict spontaneity (19.6 pg 871) Looking for a positive E electrons will spontaneously flow from a substance that will oxidize in the presence of the appropriate oxidizing agent ( ie the oxidizing agent must be below the substance being oxidized on the Std Oxidation Potential list). Do examples: a) Can you stir chlorine water with a nickel plated spoon? b) Can you stir a KMnO4 sol n with a chromium spoon? c) Will copper react with hydrochloric acid? DEMO d) Will copper react with nitric acid? DEMO **ASSIGN Will copper react with sodium nitrate? 9. Effect of Concentration on Cell Potentials Nernst Eq n (19.8 pg 881) illustrate how LeChatelier predicts that as time proceeds & the [reactants] drops + [products] rises that E should also drop ie it changes it is not a constant. To calculate Ecell at any concentration we use the NERNST equation: Ecell = E /n * log10 [products] p / [reactants] r ** explain n EX: Use the Nernst & Zn/Cu cell with [Cu 2+ ] = 1x10 4 M & [Zn 2+ ]=1x10-5 M ** Calculating concentrations using the Nernst Eq n see examples on overhead see overhead Using the Nernst **ASSIGN Using Electrochemical Principles REDOX Titrations (5.6 pg 218) During any REDOX rx o # moles electrons lost must = # moles electrons gained [ sub oxidized] * vol * ne- lost = [ sub reduced] * vol * ne- gained

4 *10b. LAB ACTIVITY Redox Titrations (2) i) find molar mass and then # moles water on ferrous diammonium disulfate ii) find molar mass and then # moles water on ferrous sulfate i) Find the % sodium hypochlorite in laundry bleach ii) Find the % sodium hypochlorite in pool shock & do cost analysis Requires: *Fe(NH4)2(SO4)2.6 H2O 41.2 g/l *FeSO4.6H2O. g/l *KMnO4 *conc H2SO4/H3PO4 80%/20% mixture Laundry bleach Pool shock Na2S2O3.5H2O 0.6 M KI 9.95 g /100 ml 40% acetic acid 40 ml glacial/60 ml water **ASSIGN: lab write-up + lab questions 11. Cell Potentials & Thermodynamics / equilibrium (19.7 pg 879) a) E and Ke As a reaction proceeds the concentration of all reactants drops and concentration of all products increase until the system reaches a state of Equilibrium where the system can not deliver any more useful energy ( work) so E = 0 So 0 = E /n * log10 [products] p / [reactants] r (Nearnst) or log10ke = ne / (since [prod] & [react] are now constant) Ke = 10 ne / Do example find Ke value for Zn/Cu cell ( 1.4 x )

5 b) E and G Recall from Equilibrium/Thermodynamics ( pg 826 ) G = RT ln Ke or G = RT (2.3)log10Ke so log10ke = G /2.3 RT but log 10 Ke = ne / as well so: G /ne = 2.3RT/ (grouping all constants) G /ne = coul / mol e ( volt=j/coul) or G /ne = 1 Faraday or G = nfe ( see pg 879) Do example find G value for Zn/Cu cell ( kj) Do example find Ke & G value for: NiO2 + 2Cl + 4 H + Cl2 + Ni H2O E = 0.32 v ( 6.5 x ) ( 62 kj) **ASSIGN Electrolysis (19.1 pg ) Passing electricity through a molten ionic compound or a solution of any electrolyte causes a REDOX reaction to occur called electrolysis. Ex 1: Electrolysis of molten sodium chloride see diagram pg 846 * go over anode & cathode and cell reaction Ex 2: Electrolysis of potassium nitrate(aq) see diagram pg 848 * do cell reaction DEMO Ex 3: Electrolysis of copper(ii) bromide (aq) see diagram pg 851 * do cell reaction DEMO ** DO example 19.1 pg 851 assign practice execise 1 pg 852

6 13. Stoichiometric relationships in Electrolysis (19.2 pg 852) *Quantitative Electrolysis 1 mol Ag mol e (1 Faraday) 1 mol Ag 1 mol Cu mol e (2 Faraday) 1 mol Cu 1 mol Al mol e (3 Faraday) 1 mol Al 1 mole of electrons will reduce 1 mol of Ag but only 0.5 mole of Cu and 0.33 mole of Al. * use Q = I * t to find the # coulombs then # Faradays 3 moles reduced # grams see EX 19.3 pg 854 ( calculating time) see EX 19.4 pg 855 ( calculating current) DO practice exercise 2, 3 & 4 pg 855 **ASSIGN Industrial Applications of Electrolysis (19.3 pg ) discuss a) electroplating ( pg 856 ) DEMO b) production of Aluminum ( pg 857 ) see overhead Aluminum production c) production of Mg/ Na / Cu ( pg ) **ASSIGN : summary question: 7.18 MAJOR TEST #7: Electrochemistry

Chapter 17. Electrochemistry

Chapter 17. Electrochemistry Chapter 17 Electrochemistry Contents Galvanic cells Standard reduction potentials Cell potential, electrical work, and free energy Dependence of cell potential on concentration Batteries Corrosion Electrolysis