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Evans Diagrams. Where we left off Tafel Equation The Tafel slope is an intensive parameter and does not depend on the electrode surface area. i 0 is.

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Presentation on theme: "Evans Diagrams. Where we left off Tafel Equation The Tafel slope is an intensive parameter and does not depend on the electrode surface area. i 0 is."— Presentation transcript:

1 Evans Diagrams

2 Where we left off

3 Tafel Equation The Tafel slope is an intensive parameter and does not depend on the electrode surface area. i 0 is and extensive parameter and is influenced by the electrode surface area and the kinetics or speed of the reaction. Notice that the Tafel slope is restricted to the number of electrons, n, involved in the charge transfer controlled reaction and the so called symmetry factor, . n is often = 1 and although the symmetry factor can vary between 0 and 1 it is normally close to 0.5. This means that the Tafel slope should be close to 120 mV if n = 1 and 60 mV if n = 2. (The latter is normally not the case)

4 Butler-Volmer Equation where: I = electrode current, Amps I o = exchange current density, Amp/m 2 E = electrode potential, V E eq = equilibrium potential, V A = electrode active surface area, m 2 T = absolute temperature, K n = number of electrons involved in the electrode reaction F = Faraday constant R = universal gas constant α = so-called symmetry factor or charge transfer coefficient dimensionless The equation is named after chemists John Alfred Valentine Butler and Max Volmer

5 Butler-Volmer Equation – High Field Strength i a and i c are the exhange current densities for the anodic and cathodic reactions These equations can be rearranged to give the Tafel equation which was obtained experimentally

6 Butler Volmer Equation - Tafel Equation

7 Current Voltage Curves for Electrode Reactions Without concentration and therefore mass transport effects to complicate the electrolysis it is possible to establish the effects of voltage on the current flowing. In this situation the quantity E - E e reflects the activation energy required to force current i to flow. Plotted below are three curves for differing values of i o with α = 0.5.

8 Current Voltage Curves for Single Electrode Reactions Current Voltage Electrochemical reactions of different i 0 or degrees of reversibility The iE curves from the previous slide have been rotated.

9 Single Chemical Reaction Only at appreciable over- potentials does the reverse reaction become negligible At E e the forward and reverse currents are equal

10 Electrochemical reaction which has a large exchange current density, i 0, This means that a small applied voltage results in an appreciable increase in current. Electrode reactions which have a high exchange current density are not easily polarised. Examples are the hydrogen evolution reaction on Pt and AgCl + e ↔ Ag + Cl - The H + /H 2 (Pt) and Ag/AgCl make good reference electrodes because they are not easily polarised

11 11 Electrochemical reaction in which the i 0 value is very low. This means that it takes an appreciable over- potential to produce a significant current. This electrode is easily polarisable since a small current would result in a significant change in voltage

12 12 At low overpotential the Butler Volmer equation is linear (Stern Geary equation)

13 So far we have looked mainly at single electrochemical reactions

14 14 KINETICS OF AQUEOUS CORROSION Anodic and cathodic reactions are coupled at a corroding metal surface Schematics of two distinct corrosion processes. (a) The corrosion process M + O  M n+ + R showing the separation of anodic and cathodic sites. (b) The corrosion process involving two cathodic reactions.

15 Wagner Traud Method The cathodic and anodic reactions are drawn together on the same graph to show how the currents are equal at the corrosion potential Butler Volmer graphs for two electrochemical reactions

16 Note in the previous diagram that: i a = i c = i corr at the corrosion potential E corr E corr is a mixed potential which lies between (E e ) c and (E e ) a. In this case it is closer to (E e ) a because the i 0 and the kinetics of the anodic reaction is faster. The metal dissolution is driven by the anodic activation overpotential η a = E corr - (E e ) a The cathodic reaction is driven by the cathodic activation overpotential η c = E corr - (E e ) c The thermodynamic driving force ΔE = (E e ) c - (E e ) a ΔE is usually large enough to put E corr in the Tafel region for both reactions, i.e. the reverse reaction is negligible.

17 Evans Diagrams It is convenient to represent the linear plots of i and E as log i/E plots with the negative cathodic current plotted positively, i.e. both the anodic and cathodic current appear in the positive quadrant. The linear region gives us the Tafel slopes The i 0 for the individual reactions can be obtained by extrapolating back to (E e ) a and (E e ) c if these values are known.

18 Evans Diagrams The intersection of the two curves at E corr gives us i corr Of course you do not see the portion of the E/logi c and E/logi a at potentials more positive and more negative of E corr respectively. However, it is important to realise that they exist. I believe it is worthwhile to look at your Tafel type measurements as a linear representation of current and voltage. The logarithmic plots involve a mathematical manipulation of data and errors can be introduced. Nevertheless Evans Diagrams are a convenient way of viewing electrochemical reactions

19 Evans Diagrams In this case the cathodic reaction with the higher oxidation potential is controlling the reaction

20 Evans Diagrams In this example because of the faster kinetics. the cathodic reaction taking place at the lower oxidation (+ve) Potential is influencing the corrosion rate more,

21 Evans Diagrams The situation in the previous example often occurs for a metal corroding in acid, compared with the metal corroding in dissolved oxygen. Despite the thermodynamic driving force, E e, being greater for oxygen than H 2 /H +, the acid corrosion is faster. In some cases the oxygen and acid have a synergistic effect. For example in the case of Ni corrosion. The reaction is quite slow in sulphuric acid (0.5 M) and it is also slow in water saturated with air at pH 7. In the latter case a passive protective oxide film is formed. However, in the presence of sulphuric acid and air. The corrosion rate is relatively rapid. The acid dissolves the protective oxide film allowing oxygen to corrode the metal.

22 Evans Diagrams The relative corrosion rates of metals depends on the i 0 and mass transfer. With acid corrosion: 2H + + e → H 2 i 0 can vary from – A cm -2 The Tafel slope  120 mV/decade For oxygen corrosion O 2 + H 2 O + 4e → 4OH - I 0 is difficult to difficult to determine because it is very low, but it is of the order of < A cm -2 The Tafel slope >120 mV/decade

23 Exchange Current Densities in 1 Molal H 2 SO 4 Electrode Material-log 10 (A/cm 2 Palladium3.0 Platinum3.1 Rhodium3.6 Nickel5.2 Gold5.4 Tungsten5.9 Niobium6.8 Titantium8.2 Cadmium10.8 Manganese10.9 Lead12 Mercury12.3

24 α Values for Some Reactions MetalSystemα PtFe 3+ + e ↔ Fe PtCe 4+ + e ↔ Ce HgTi 4+ + e ↔ Ti Hg2H + + 2e ↔ H Ni2H + + 2e ↔ H AgAg + + e ↔ Ag0.55

25 Evans Diagrams The slowest reaction controls the rate of corrosion. Normally this is the cathodic reaction. In this example: A small changes in kinetics of cathode have a large effect on corrosion rate. A small changes in kinetics of anode have small effect on corrosion

26 Mass Transfer Control If the cathodic reagent at the corrosion site (e.g., dissolved O 2 in the O 2 reduction) is in short supply, mass transfer of the reagent can become rate limiting. The cathodic charge-transfer reaction at the metal/solution interface is fast enough to reduce the concentration of the reagent at interface (cathodic sites) to a value less than that in the bulk solution. This sets up a concentration gradient and the reaction becomes diffusion controlled.

27 Mass Transfer Control When the corrosion rate is limited by mass transfer it can be increased by: By altering the bulk concentration By stirring and reducing the thickness of the Nernst diffusion layer

28 Mass Transfer Control Diffusion or Mass Transfer Controlled Activation Controlled

29 Mass Transfer Control Increase in corrosion potential, E corr, and the corrosion current, i corr, due to an increase in mass transfer caused by stirring.

30 Mixed Transfer Control Polarization curve for the cathodic process showing: 1.Activation polarization 2.Joint activation- concentration polarization 3.Mass transport- limited corrosion control The cathodic Tafel plot often shows deviation from ideal Tafel behavior

31 Evans Diagrams Anodic Control Mixed Control Cathodic Control

32 Galvanic Corrosion – Influence of i 0

33 Cyclic Voltammetry at a Pt Electrode in Sulphuric Acid Solution Oxygen Adsorption Pt-O Oxygen Evolution O 2 ↑ Reduction of adsorbed oxide film (Pt-O) Hydrogen Evolution H 2 ↑ Reduction of adsorbed H (Pt-H) Formation of adsorbed H (Pt-H) The peak height of the adsorption/desorption processes is directly proportional to scan, i.e., the charge iE or area under the curve. This contrasts with a diffusion process where the peak height is proportional to the square root of the scan rate.

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