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Electrochemistry Thermodynamics at the electrode.

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Presentation on theme: "Electrochemistry Thermodynamics at the electrode."— Presentation transcript:

1 Electrochemistry Thermodynamics at the electrode

2 Learning objectives  You will be able to:  Identify main components of an electrochemical cell  Write shorthand description of electrochemical cell  Calculate cell voltage using standard reduction potentials  Apply Nernst equation to determine free energy change  Apply Nernst equation to determine pH  Calculate K from electrode potentials  Calculate amount of material deposited in electrolysis

3 Energy in or energy out  Galvanic (or voltaic) cell relies on spontaneous process to generate a potential capable of performing work – energy out  Electrolytic cell performs chemical reactions through application of a potential – energy in

4 Redox Review  Oxidation is...  Loss of electrons  Reduction is...  Gain of electrons  Oxidizing agents oxidize and are reduced  Reducing agents reduce and are oxidized

5 Redox at the heart of the matter  Zn displaces Cu from CuSO 4 (aq)  In direct contact the enthalpy of reaction is dispersed as heat, and no useful work is done  Redox process:  Zn is the reducing agent  Cu 2+ is the oxidizing agent

6 Separating the combatants  Each metal in touch with a solution of its own ions  External circuit carries electrons transferred during the redox process  A “salt bridge” containing neutral ions completes the internal circuit.  With no current flowing, a potential develops – the potential for work  Unlike the reaction in the beaker, the energy released by the reaction in the cell can perform useful work – like lighting a bulb

7 Labelling the parts

8 Odes to a galvanic cell  Cathode  Where reduction occurs  Where electrons are consumed  Where positive ions migrate to  Has positive sign  Anode  Where oxidation occurs  Where electrons are generated  Where negative ions migrate to  Has negative sign

9 The role of inert electrodes  Not all cells start with elements as the redox agents  Consider the cell  Fe can be the anode but Fe 3+ cannot be the cathode.  Use the Fe 3+ ions in solution as the “cathode” with an inert metal such as Pt

10 Anode Catho de Oxidati on Reduct ion

11 Cell notation  Anode on left, cathode on right  Electrons flow from left to right  Oxidation on left, reduction on right  Single vertical = electrode/electrolyte boundary  Double vertical = salt bridge Anode: Zn →Zn e Cathode: Cu e →Cu

12 Vertical │denotes different phase  Fe(s)│Fe 2+ (aq)║Fe 3+ (aq),Fe 2+ (aq)│Pt(s)  Cu(s)│Cu 2+ (aq)║Cl 2 (g)│Cl - (aq)│C(s)

13 Connections: cell potential and free energy  The cell in open circuit generates an electromotive force (emf) or potential or voltage. This is the potential to perform work  Energy is charge moving under applied voltage

14 Relating free energy and cell potential  The Faraday: F = C/mol e Standard conditions (1 M, 1 atm, 25°C)

15 Standard Reduction Potentials  The total cell potential is the sum of the potentials for the two half reactions at each electrode  E cell = E cath + E an  From the cell voltage we cannot determine the values of either – we must know one to get the other  Enter the standard hydrogen electrode (SHE)  All potentials are referenced to the SHE (=0 V)

16 Unpacking the SHE  The SHE consists of a Pt electrode in contact with H 2 (g) at 1 atm in a solution of 1 M H + (aq).  The voltage of this half-cell is defined to be 0 V  An experimental cell containing the SHE half-cell with other half-cell gives voltages which are the standard potentials for those half-cells E cell = 0 + E half-cell

17 Zinc half-cell with SHE  Cell measures 0.76 V  Standard potential for Zn(s) = Zn 2+ (aq) + 2e = 0.76 V

18 Where there is no SHE  In this cell there is no SHE and the measured voltage is 1.10 V

19 Standard reduction potentials  Any half reaction can be written in two ways:  Oxidation: M = M + + e (+V)  Reduction: M + + e = M (-V)  Listed potentials are standard reduction potentials


21 Applying standard reduction potentials  Consider the reaction  What is the cell potential?  The half reactions are:  E° = 0.80 V – (-0.76 V) = 1.56 V  NOTE: Although there are 2 moles of Ag reduced for each mole of Zn oxidized, we do not multiply the potential by 2.

22 Extensive v intensive  Free energy is extensive property so need to multiply by no of moles involved  But to convert to E we need to divide by no of electrons involved  E is an intensive property

23 The Nernst equation  Working in nonstandard conditions

24 Electrode potentials and pH  For the cell reaction  The Nernst equation  Half-cell potential is proportional to pH

25 The pH meter is an electrochemical cell  Overall cell potential is proportional to pH  In practice, a hydrogen electrode is impractical

26 Calomel reference electrodes  The potential of the calomel electrode is known vs the SHE. This is used as the reference electrode in the measurement of pH  The other electrode in a pH probe is a glass electrode which has a Ag wire coated with AgCl dipped in HCl(aq). A thin membrane separates the HCl from the test solution

27 Cell potentials and equilibrium  Lest we forget…  So then  and

28 Cell potential a convenient way to measure K

29 Many pathways to one ending  Measurement of K from different experiments  Concentration data  Thermochemical data  Electrochemical data

30 Batteries  The most important application of galvanic cells  Several factors influence the choice of materials  Voltage  Weight  Capacity  Current density  Rechargeability

31 Running in reverse  Recharging a battery requires to run the process in reverse by applying a voltage  In principle any reaction can be reversed  In practice it will depend upon many factors  Reversibility depends on kinetics and not thermodynamics  Cell reactions that involve minimal structural rearrangement will be the easiest to reverse

32 Lithium batteries  Lightweight (Molar mass Li = 6.94 g)  High voltage  Reversible process

33 Fuel cells – a battery with a difference  Reactants are not contained within a sealed container but are supplied from outside sources

34 Store up not treasures on earth where moth and rust…  An electrochemical mechanism for corrosion of iron. The metal and a surface water droplet constitute a tiny galvanic cell in which iron is oxidized to Fe 2+ in a region of the surface (anode region) remote from atmospheric O 2, and O 2 is reduced near the edge of the droplet at another region of the surface (cathode region). Electrons flow from anode to cathode through the metal, while ions flow through the water droplet. Dissolved O 2 oxidizes Fe 2+ further to Fe 3+ before it is deposited as rust (Fe 2 O 3 ·H2O).

35 Mechanisms  Why does salt enhance rusting?  Improves conductivity of electrolyte  Standard reduction potentials indicate which metals will “rust”  Aluminium should corrode readily. It doesn’t. Is thermodynamics wrong?  No, the Al 2 O 3 provides an impenetrable barrier

36 No greater gift than to give up your life for your friend  A layer of zinc protects iron from oxidation, even when the zinc layer becomes scratched. The zinc (anode), iron (cathode), and water droplet (electrolyte) constitute a tiny galvanic cell. Oxygen is reduced at the cathode, and zinc is oxidized at the anode, thus protecting the iron from oxidation.

37 Electrolysis  Electrolysis of a molten salt using inert electrodes  Signs of electrodes:  In electrolysis, anode is positive because electrons are removed from it by the battery  In a galvanic cell, the anode is negative because is supplies electrons to the external circuit

38 Electrolysis in aqueous solutions – a choice of process  There are (potentially) competing processes in the electrolysis of an aqueous solution  Cathode  Anode

39 Thermodynamics or kinetics?  On the basis of thermodynamics we choose the processes which are favoured energetically  But…chlorine is evolved at the anode

40 The role of overpotentials  Thermodynamic quantities prevail only at equilibrium – no current flowing  When current flows, kinetic considerations come into play  Overpotential represents the additional voltage that must be applied to drive the process

41  In the NaCl(aq) solution the overpotential for evolution of oxygen is greater than that for chlorine, and so chlorine is evolved preferentially  Overpotential will depend on the electrolyte and electrode. By suitable choices, overpotentials can be minimized but are never eliminated  The limiting process in electrolysis is usually diffusion of the ions in the electrolyte (but not always)  Driving the cell at the least current will give rise to the smallest overpotential

42 Electrolysis of water  In aqueous solutions of most salts or acids or bases the products will be O 2 and H 2

43 Quantitative aspects of electrolysis  Quantitative analysis uses the current flowing as a measure of the amount of material  Charge = current x time  Moles = charge/Faraday

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