1 / 36

Electrochemical Equilibria

Electrochemical Equilibria. Nernst Equation, Standard Reduction Potential and Galvanic Series. Component of an Electrochemical Cell – Cell Potential. A cell can be made from 2 electrodes (metals or electronic conductors) dipping into an electrolyte solution and connected by an external circuit.

huong
Download Presentation

Electrochemical Equilibria

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Electrochemical Equilibria Nernst Equation, Standard Reduction Potential and Galvanic Series

  2. Component of an Electrochemical Cell – Cell Potential • A cell can be made from 2 electrodes (metals or electronic conductors) dipping into an electrolyte solution and connected by an external circuit. • The electrical potential or voltage, E, between the two electrodes in a cell can be measured by connecting a voltmeter. • The measured voltage is known as the cell potential, E. Older terminology refers to the electromotive force, EMF. • The cell potential is related to the Gibbs Free Energy change for the overall reaction carried out by the cell.

  3. Cell Potential • ΔG = -nFE • ΔG = Gibbs Free Energy change, J mol- • n = number of electrons transferred in the cell reaction • F = Faraday’s constant, 96,485 Coulombs mol-1 • E = cell potential, volts, V • Thus the cell potential is equivalent to a measure of how much work can be done by the electrons flowing through the external circuit of the cell. It is strictly a THERMODYNAMIC measurement of the cell.

  4. Current – Rate of Reaction • The current flowing through a cell can be directly related to the rate of the cell reaction. Thus measurement of the current enable the kinetics of the cell reaction to be studied. • Rate = I/nF • Rate, mol s-1 • I = current, Amps, A • n = number of electrons transferred in the cell reaction. • F = Faraday’s constant, 96485 coulombs mol-1

  5. Standard State • The standard state is defined as a temperature of 298 K and a pressure of 1 bar. Pure substances (solids, liquids and gases) should be in their normal state at this temperature and pressure. Solutions should have an activity* equal to one, which we will approximate here as a concentration equal to 1 mol L-1 • The potential of a cell under standard conditions is denoted Eo. • Hence we can find the standard Gibbs Free Energy change, and the equilibrium constant for the cell reaction. • ΔGo = -nFEo = -RTlnK • The concept of activity relates concentration to its effective thermodynamic equivalent.

  6. The Standard Hydrogen Electrode • Since a cell is made from two half cells, it is useful to identify the contribution of each half cell to the cell potential. • It is impossible to measure the potential of a single half cell, this is because every attempt to take a measurement using a voltmeter will introduce a wire (metal/solution interface) or another half cell. • To overcome this problem we need a standard for comparison – the standard hydrogen electrode (SHE). • H+ (aq) (1 mol L-1 ) І H2 (g) (1 bar) І Pt at 298o K • Technically we should use the activity = 1, rather than concentration.

  7. The Standard Hydrogen Electrode The standard hydrogen electrode: Platinized platinum electrode hydrogen gas flow solution of the acid with activity of H+ = 1 mol kg-1 hydroseal for prevention of the oxygen interference reservoir through which the second half-element of the galvanic cell should be attached. This creates an ionically conductive path to the working electrode of interest.

  8. You cannot measure the absolute potential of a half cell. You can only measure a potential difference between half cells or a change in potential with respect to a reference electrode

  9. The Standard Hydrogen Electrode Pt • 2H+(s) + 2e ↔ H2(g)

  10. The Standard Hydrogen Electrode

  11. Exchange Current Densities in 1 Molal H2SO4

  12. The Standard Half Cell Potential • All other half cells can be compared to the SHE. A half cell in its standard state is connected as the cathode to the SHE connected as the anode. • The measured potential is called the standard half cell potential, Eo. • This connection is made regardless of which half cell may actually be anode or cathode, and will result in negative values for some half cells. • The half cell reaction should always be shown as a standard reduction potential. • Eo can be related to the tendency of the half cell to reduce H+ (aq) (1 mol L-1 ) to H2 (g) (1 bar) • The larger (more positive) the value of Eo, the more favorable that process.

  13. Standard Half Cell Potential • The metal at the more negative potential is more easily oxidised or corroded. • The metal at the more positive potential is more easily reduced.

  14. Standard Half Cell Potentials • Consider the following half cell reduction reactions • Cu2+(aq) + 2e → Cu(s) E0red = +0.340 V • Zn2+(aq) + 2e → Zn(s) E0red = -0.763 V • Therefore copper will be reduced and zinc oxidised • Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) • The voltage of the cell E0cell = 0.340 – (0.763) = 1.103V

  15. Standard Half Cell Potentials • Concentration cells • Fe → Fe2+ + 2e • An iron electrode in a solution of 10-2M Fe2+ will have a potential of -0.56 V and a similar electrode in 10-5M Fe2+ will have a potential of -0.74 volts. • If the electrodes were connected the cell with the lower concentration of Fe2+ would corrode to produce more Fe2+ and the cell with a higher concentration of Fe2+ would undergo reduction to decrease the Fe2+ concentration.

  16. The Standard Half Cell Potential The reason the half cell reaction should be shown as a standard reduction potential is that the potential given will be the same as the potential measured with a voltmeter. A list of standard half cell reductions potentials is given to enable you to predict reactions that will occur. Half cell oxidation potentials are also given in the literature. For a particular half cell reaction the potential is the same but opposite in sign to the corresponding reduction potential. This is done for thermodynamic reasons. However, it causes confusion and since they are of no practical benefit they can be ignored.

  17. Non – Standard Half Cell Potentials • This could be measured for any half cell not in its standard state, connected as the cathode to the SHE. • We can calculate the non-standard half cell (equilibrium) potential using the Nernst Equation:

  18. Non – Standard Half Cell Potentials • Where: • E = non-standard half cell potential • E0 = standard half cell potential • n = number of electrons in standard reduction reaction • Q = reaction quotient for standard reduction reaction. • Q, the reaction quotient has the same appearance as the equilibrium constant, but does not necessarily relate to equilibrium conditions.

  19. Non – Standard Half Cell Potentials • Non standard cell potential can then be calculated from the non standard half cell potentials: • Ecell = Ecathode - Eanode

  20. Galvanic Series • The galvanic series relates the potential of metals in seawater and enables the corrosion engineer to predict the corrosivity of metals and whether the coupling of two metals will be compatible or not. • While you can predict whether a metal coupled to another metal may undergo galvanic corrosion it cannot predict the rate of corrosion

  21. Key Thermodynamic Quantity • Another key thermodynamic quantity, the entropy change, ΔS, can be determined from the variation of the cell potential with temperature. • Where: • ΔS = Entropy change, J K-1 mol-1 • n = number of electrons transferred in the cell reaction • F = Faraday’s constant, 96485 C mol-1

  22. Key Thermodynamic Quantity • After the entropy change is determined, the enthalpy change can be determined from the fundamental thermodynamic relationship • ΔG = ΔH – TΔS • Where: • ΔG = Gibbs Free Energy change, J mol-1 • ΔH = Enthalpy change, J mol-1 • ΔS = Entropy change, J K-1 mol-1 • It can also be noted that the equilibrium potential, E, gives the minimum free energy that will need to be supplied to an electrolytic cell to force the cell reaction.

  23. Standard Half Cell Potentials Consider the following half cell reduction reactions Cu2+(aq) + 2e → Cu(s) E0red = +0.340 V Zn2+(aq) + 2e → Zn(s) E0red = -0.763 V Therefore copper will be reduced and zinc oxidised Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) The voltage of the cell E0cell = 0.340 – (0.763) = 1.103V

  24. Galvanic Cell Cu/Cu2+ Zn/Zn2+ Cu Zn CuSO4 ZnCl2

  25. Galvanic Cell Cu/Cu2+ Zn/Zn2+

  26. Galvanic Cell Cu/Cu2+ Zn/Zn2+

  27. Galvanic Cell Cu/Cu2+ Zn/Zn2+

  28. Galvanic Corrosion • As a zinc atom provides the electrons, it becomes a positive ion and goes into aqueous solution, decreasing the mass of the zinc electrode. On the copper side, the two electrons received allow it to convert a copper ion from solution into an uncharged copper atom which deposits on the copper electrode, increasing its mass. The two reactions are typically written • Zn(s) -> Zn2+(aq) + 2e- • Cu2+(aq) + 2e- -> Cu(s) • The letters in parentheses are just reminders that the zinc goes from a solid (s) into a water solution (aq) and vice versa for the copper. It is typical in the language of electrochemistry to refer to these two processes as "half-reactions" which occur at the two electrodes. • Zn(s) -> Zn2+(aq) + 2e-

  29. Galvanic Corrosion • The zinc "half-reaction" is classified as oxidation since it loses electrons. The terminal at which oxidation occurs is called the "anode". For a battery, this is the negative terminal. The copper "half-reaction" is classified as reduction since it gains electrons. The terminal at which reduction occurs is called the "cathode". For a battery, this is the positive terminal. Cu2+(aq) + 2e- -> Cu(s) • In order for the voltaic cell to continue to produce an external electric current, there must be a movement of the sulfate ions in solution from the right to the left to balance the electron flow in the external circuit. The metal ions themselves must be prevented from moving between the electrodes, so some kind of porous membrane or other mechanism must provide for the selective movement of the negative ions in the electrolyte from the right to the left.

  30. Galvanic Corrosion • Energy is required to force the electrons to move from the zinc to the copper electrode, and the amount of energy per unit charge available from the voltaic cell is called the electromotive force (emf) of the cell. Energy per unit charge is expressed in volts (1 volt = 1 joule/coulomb). • Clearly, to get energy from the cell, you must get more energy released from the oxidation of the zinc than it takes to reduce the copper. The cell can yield a finite amount of energy from this process, the process being limited by the amount of material available either in the electrolyte or in the metal electrodes. For example, if there were one mole of the sulfate ions SO42- on the copper side, then the process is limited to transferring two moles of electrons through the external circuit. The amount of electric charge contained in a mole of electrons is called the Faraday constant, and is equal to Avogadro's number times the electron charge:

  31. Galvanic Corrosion • Faraday constant = F = NAe = 6.022 x 1023 x 1.602 x 10-19 = 96,485 Coulombs/mole The energy yield from a voltaic cell is given by the cell voltage times the number of moles of electrons transferred times the Faraday constant. • Electrical energy output = nFEcell • The cell emf Ecell may be predicted from the standard electrode potentials for the two metals. For the zinc/copper cell under the standard conditions, the calculated cell potential is 1.1 volts.

  32. EMF of Cells • http://video.google.com/videoplay?docid=566682988710852020

  33. http://web.mst.edu/~gbert/Electro/Electrochem.html • http://www.funsci.com/fun3_en/electro/electro.htm • http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/electroChem/voltaicCell20.html

More Related