Electrochemical Cells & The Electrochemical Series 1.A simple cell from an exothermic redox reaction Zn (s) + Cu 2+ (aq)A simple cell from an exothermic.

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Electrochemical Cells & The Electrochemical Series 1.A simple cell from an exothermic redox reaction Zn (s) + Cu 2+ (aq)A simple cell from an exothermic redox reaction Zn (s) + Cu 2+ (aq) 2.Redox Equilibrium – zinc in waterRedox Equilibrium – zinc in water 3.Standard Hydrogen Reference Half CellStandard Hydrogen Reference Half Cell 4.Standard electrode potential – Zn/Zn 2+Standard electrode potential – Zn/Zn 2+ 5.Standard electrode potential – Cu/Cu 2+Standard electrode potential – Cu/Cu 2+ 6.Standard electrode potential – Fe 2+ /Fe 3+Standard electrode potential – Fe 2+ /Fe 3+ 7.The Electrochemical SeriesThe Electrochemical Series 8.Calculation of Cell e.m.f.Calculation of Cell e.m.f. 9.QuestionsQuestions 10.Useful Internet LinkUseful Internet Link

Cu 2+ (aq) H 2 O (l) Electrochemical Cell Zn V SALT BRIDGE Zn (s)  Zn 2+ (aq) + 2e - Cu 2+ (aq) + 2e -  Cu (s) HALF EQUATION

Salt Bridge Maintains electrical neutrality. Solution of ionic compound, usually potassium bromide KBr. Positive (K + ) ions move into negative half cell Negative (Br - )ions move into positive half cell Br - K+K+

Very High Resistance Voltmeter No current drawn. Measures the electromotive force [e.m.f.] of the cell. The potential for the cell to provide energy

OXIDATION ANODE OXID A TION An anode is a place of oxidation Zn (s)  Zn 2+ (aq) + 2e -

REDUCTION CATHODE REDU C TION A cathode is a place of reduction Cu 2+ (aq) + 2e -  Cu (s)

Redox Equilibrium [PRESS ZINC] Zinc Zn Zn 2+ e - Zne - Zn 2+ Zn (s) Ý Zn 2+ (aq) + 2e - e - Zn 2+ Add stronger oxidising agentAdd stronger reducing agent ADDING REDOX AGENTS

Redox Equilibrium 2 MAGNESIUM e - Mg Mg 2+ e - Mg Mg 2+ Mg (s) Ý Mg 2+ (aq) + 2e - e - Mg 2+ Equilibrium for magnesium lies further to right than for zinc Magnesium

magnesium sheds electrons and forms ions more readily than zinc does. can't measure the absolute voltage between the metal and the solution don't need to be able to measure the absolute voltage between the metal and the solution. enough to compare the voltage with a standardised system [standard hydrogen reference electrode].standard hydrogen reference electrode

Oxidising agent added, e.g. Cu 2+ : Zinc Zn 2+ Zn (s) + Cu 2+ (aq)  Zn 2+ (aq) + Cu (s) e - Cu 2+ Zn 2+ Cu EQUATION

Reducing agent added, e.g. Mg: Zinc Zn 2+ Mg (s) + Zn 2+ (aq)  Mg 2+ (aq) + Zn (s) e - Mg 2+ Zn 2+ e - Mg EQUATION Zn

1 mol dm -3 H + (aq) Standard Hydrogen ReferenceElectrode 2H + (aq) + 2e -  H 2 (g) HALF EQUATION HYDROGEN GAS 1 atmosphere 298K Pt Platinum Electrode

1 mol dm -3 Zn 2+ (aq) E θ Zn 2+ / Zn Zn V SALT BRIDGE Zn 2+ (aq) + 2e -  Zn (s) HALF EQUATION 1 mol dm -3 H + (aq) Pt H 2 (g) Pt (s)  H 2 (g), 2H + (aq)  Zn 2+ (aq)  Zn(s) CELL NOTATION E θ = -0.76V EθEθ

1 mol dm -3 Cu 2+ (aq) E θ Cu 2+ / Cu Cu V SALT BRIDGE Cu 2+ (aq) + 2e -  Cu (s) HALF EQUATION 1 mol dm -3 H + (aq) Pt H 2 (g) Pt (s)  H 2 (g), 2H + (aq)  Cu 2+ (aq)  Cu(s) CELL NOTATION E θ = +0.34V EθEθ

1 mol dm -3 Fe 2+ (aq) E θ Fe 3+ / Fe 2+ Pt V SALT BRIDGE Fe 3+ (aq) + e -  Fe 2+ (aq) HALF EQUATION 1 mol dm -3 H + (aq) Pt H 2 (g) Pt (s)  H 2 (g), 2H + (aq)  Fe 3+ (aq), Fe 2+ (aq)  Pt(s) CELL NOTATION E θ = +0.77V EθEθ 1 mol dm -3 Fe 3+ (aq)

Electrochemical Series Li + (aq) + e − Ý Li(s)−3.05V K + (aq) + e − Ý K(s)−2.93V Ca 2+ (aq) + 2e − Ý Ca(s)−2.76V Na + (aq) + e − Ý Na(s)−2.71V Mg 2+ (aq) + 2e − Ý Mg(s)−2.38V Al 3+ (aq) + 3e − Ý Al(s)−1.68V Zn 2+ (aq) + 2e − Ý Zn(s)−0.76V Fe 2+ (aq) + 2e − Ý Fe(s)−0.44V 2H + (aq) + 2e − Ý H 2 (g) 0.00V Cu 2+ (aq) + e − Ý Cu+(aq)+0.16V Cu 2+ (aq) + 2e − Ý Cu(s)+0.34V O 2 (g) + 2H 2 O(l) + 4e - Ý 4OH - (aq)+0.40V Cu + (aq) + e − Ý Cu(s)+0.52V I 2 (s) + 2e − Ý 2I − (aq)+0.54V Fe 3+ (aq) + e − Ý Fe 2+ (aq)+0.77V Ag + (aq) + e − Ý Ag(s)+0.80V Br 2 (aq) + 2e − Ý 2Br − (aq)+1.09V Cl 2 (g) + 2e− Ý 2Cl − (aq)+1.36V Cr 2 O 7 2− (aq) + 14H + + 6e − Ý 2Cr 3+ (aq) + 7H 2 O(l)+1.38V MnO 4 − (aq) + 8H + + 5e − Ý Mn 2+ (aq) + 4H 2 O(l)+1.51V F 2 (g) + 2e − Ý 2F − (aq)+2.87V

1 mol dm -3 Cu 2+ (aq) 1 mol dm -3 Zn 2+ (aq) Cell e.m.f. Zn V SALT BRIDGE Cu Zn (s)  Zn 2+ (aq)  Cu 2+ (aq)  Cu(s) CELL NOTATION E θ CELL = – (-0.76) = +1.10V E θ CELL opposite Questions

1 mol dm -3 Zn 2+ (aq) 1 mol dm -3 Cu 2+ (aq) Cell e.m.f. Cu V SALT BRIDGE Zn Cu (s)  Cu 2+ (aq)  Zn 2+ (aq)  Zn(s) CELL NOTATION E θ CELL = – (+0.34) = -1.10V E θ CELL Questions

QUESTIONS Feasibility Of Reactions Electrochemical Cells

QUESTIONS Feasibility of Reactions For each of the following mixtures predict whether a reaction is or isn’t feasible at s.t.p. If a reaction is feasible then write the balanced ionic equation. 1.Li(s) and Cu 2+ (aq) 2.Zn(s) and Al 3+ (aq) 3.O 2 (g), H 2 O(l) and Mg(s) 4.Br 2 (aq) and Fe(s) 5.Acidified Cr 2 O 7 2- and I - (aq)

QUESTIONS electrochemical cells For each of the following redox reactions write the cell notation and calculate the e.m.f. of the cell at s.t.p. 1.Li(s) + Ag + (aq)  Li + (aq) + Ag(s) 2.3Na(s) + Al 3+ (aq)  3Na + (aq) + Al(s) 3.Cr 2+ (aq) + Fe 3+ (aq)  Fe 2+ (aq) + Cr 3+ (aq) 4.O 2 (g) + 2H 2 O(l) + 2Cu(s)  4OH - (aq) + 2Cu 2+ (aq) 5.Ca(s) + 2K + (aq)  Ca 2+ (aq) + 2K(s)