CHEM-E6185 Applied Electrochemistry and Corrosion

Size: px

Start display at page:

Download "CHEM-E6185 Applied Electrochemistry and Corrosion"

Shannon Brown
6 years ago
Views:

1 CHEM-E6185 Applied Electrochemistry and Corrosion Lecture 1, electrochemical reactions and Faradays law Contents 1. Introduction 2. Electrode potential 3. Reaction rates Faraday s law Mixed potential theory and corrosion cell 2 1

2 Introduction Electrochemistry is the background for leaching of raw materials, electrolysis processes and some of the solution purification and product recovery processes. Electrochemical reactions happen on the surface of a solid and liquid phase and involve charge transfer between the phases. Every system needs at least one oxidation reaction and one reduction reaction. 3 Introduction In an electrochemical reaction the chemical energy stored in a material is spontaneosusly converted to energy or by using electrical energy the reactions are driven to produce elements and compounds. Two types of reactions: Anodic reaction is oxidation reaction that releases electrons, like dissolution of a metal. Cathodic reaction is reduction reaction that consumes electrons, like deposition of a metal. 4 2

3 Introduction Thermodynamics will describe what is the tendency for a chemical system to move towards a certain equilibrium state. In electrochemistry thermodynamics is described by potential. Reaction kinetics will describe how fast or how slowly the system moves towards the equilibrium. Reaction kinetics will determine the reaction rate and in electrochemistry reaction rate is measured by the number of electrons in time, that is the electric current. 5 Introduction Any deviation from equilibrium state results in reactions that proceed from higher potential to lower potential. If the potential of a metal electrode is higher than the potential in solution, the metal ions move to the solution and the electrode acts as an anode. If the potential of the metal electrode is lower then the metal ions can move from solution to electrode and it acts as a cathode. By changing the magnitude and direction of the potential difference it is possible to change the rate and direction of an electrochemical reaction. 6 3

4 Electrode potential Electrode potential, E (IUPAC) Electromotive force of a cell in which the electrode on the left is a standard hydrogen electrode and the electrode on the right is the electrode in question. Standard electrode potential, E (IUPAC) The value of the standard emf of a cell in which molecular hydrogen under standard pressure is oxidized to solvated protons at the lefthand electrode. Electrode potential (EN-ISO 8044) Voltage measured in the external circuit between an electrode and a reference electrode in contact with the same electrolyte. 7 Electrode potential Equilibrium state can be calculated from thermodynamic values, the standard electrode potential, E o. In the equilibrium state dynamic equilibrium causes the reaction to proceed in both anodic and cathodic directions at equal rates. No net effect is seen. A lowe o value indicates that the reaction has tendency to proceed into the anodic (oxidation, dissolution) direction. A high potential value indicates tendency to proceed into the cathodic (reduction, deposition) direction. 8 4

5 Electrode potential Metal/metal ion E 0 vs. H 2 /H + [V] equilibrium Au/Au Noble Pt/Pt Pd/Pd Ag/Ag Hg/Hg Cu/Cu H 2 /H Pb/Pb Sn/Sn Ni/Ni Co/Co Fe/Fe Cr/Cr Zn/Zn Al/Al Mg/Mg Na/Na K/K Active Gold is noble and will not dissolve easily. Iron is active and often dissolves easily 9 Electrode potential Graphite Titanium "acid resistant" Monel "stainless" 70/30 CuNi Lead 90/10 CuNi Si-bronze Copper Brass Al-bronze Steel, cast iron Al-alloys Zinc Magnesium POTENTIAL, mv vs. AgCl 10 5

6 Electrode potential 11 Electrode potential Reference electrodes Electrode Filling electrolyte E/mV vs. SHE at 25EC Temperature range,ec Temperature coefficient, mv/ec Use Ag/Ag2S 4M Na2S Kraft liquors HgTl/TlCl 3.5M KCl Hot media Cu/CuSO4 Sat. CuSO Soil, waters Hg/HgO/OH- 1M NaOH Alkaline media Ag/AgCl 3M KCl General, hot media Hg/Hg2Cl2/Cl - Sat. KCl General Ag/AgCl 0.1M KCl General Hg/HgSO4 Sat. K2SO Sulfate containing media, alkaline media 12 6

7 Electrode potential E 0 Fe H + /H 2 CuSO 4-0,44 0 0,32 E / V H + /H 2-0,78 Zn / sea water 0,25 AgCl / sea water For conversion from one potential scale to another subtract or add the reference potential difference. For thermodynamical calculations use always the hydrogen electrode scale. 13 Electrode potential The standard electrode potential of an electrochemical reaction is E o. It is calculated in standard conditions pressure 100 kpa (used to be 1 atm), temperature 25 C and activities of reacting species areone. In the equilibrium state reaction proceeds in both directions at equal rates. The system is in dynamic equilibrium and not net change is noticed. 14 7

8 Dynamic equilibrium Me + Me e - Me(s) Metal surface Me + Solvated metal ions 15 Electrode potential Standard electrode potential E o is calculated from Gibbs energy change E o = - DG z F If not in standard state equilibrium potential is calculated from Nernst equation. E = E o - RT zf [ RED] ln [ OX ] IUPAC definition is to write the equations as reduction reactions. 16 8

9 Faradays laws Quantitative laws used to express magnitudes of electrolytic effects, first described by the English scientist Michael Faraday in 1833: 1. The amount of chemical change produced by current at an electrode-electrolyte boundary is proportional to the quantity of electricity used. 2. The amounts of chemical changes produced by the same quantity of electricity in different substances are proportional to their equivalent weights. In electrolytic reactions, the equivalent weight of a substance is the gram formula weight associated with a unit gain or loss of electron. 17 Faradays laws The quantity of electricity that will cause a chemical change of one equivalent weight unit has been designated a faraday. Faraday is equivalent to coulombs of electricity (Encyclopedia Britannica value). 1 coulomb (C) = 1 Ampere-second (As) 18 9

10 Faradays law n m = = M I t z F M ekv = g / Ah z 26.8 For iron the electrochemical equivalent is 55.9 g/mol / 2*96500 As/mol = 0.29 mg/as or 1.04 g/ah. 19 Faraday s law In electrochemical research it is a rule to use current density instead of current. Current density is the measured current divided by the geometric area of the reacting electrode. When the current density on an electrode is known, calculation of the amount of reacting substance is possible. m[g/m 2 s] = ekv[g/as] i [A/m 2 ] 20 10

11 Faradays law Current density can be converted to mass of reacted substance by using Faradays law. When the reacted mass is related to reacting area and experiment length, the dissolution/corrosion/deposition rate can be expressed as weight loss/area/time. Weight loss/area/time can be converted to change of thickness by dividing by material density. 21 Faradays law Dissolution of iron at current density of 1 ma/cm 2 results in Mass loss = 0.29 mg/as * A/cm 2 * sec/day = 25.0 mg/cm 2 /day Mass loss per area and time can be converted to loss of thickness. The density of iron is 7.87 g/cm 3. When the mass loss of 25.0 mg/cm 2 /day is divided by density, the loss of thickness becomes: Loss of thickness = g/cm 2 /day / 7.87 g/cm days/year = 1.16 cm/year A rule of thumb: For common engineering metals with molar weight about 60 g/mol, density about 8 g/cm 3, and dissolving with two electrons, current density of 1 ma/cm 2 equals g/m 2 /day or mm/year

12 Faradays law What if the material is not a pure metal but an alloy or compound? Use the electrochemical equivalents of pure elements and their mass fractions ekv [g/ah] = 1 f i ekv i For example CuZn40 alloy, ekv = 1/(0.6/ /1.220) = g/ah 23 Faradays law What is the growth rate in mm/h when depositing copper using current density 270 A/m 2? Corrosion rate of Fe18Cr10Ni stainless steel in mm/year, when corrosion current density is 1.5 ma/cm 2? 24 12

13 Electrochemical reactions are heterogeneous reactions. They occur in several phases in subsequent steps: Transfer of reactive species from electrolyte to electrode surface Adsorption Charge transfer step at the surface Desorption Transfer of reaction products from the surface to the electrolyte. The slowest one determines the total reaction rate. 25 AQUEOUS BULK ELECTROLYTE O 2 O 2 O 2 O 2 ANODE CATHODE Adsorption of oxidizer from bulk electrolyte O 2 Fe Fe 4 e - Charge transfer reactions and electron transfer ANODE CATHODE from anode to cathode areas Fe = Fe e - O 2 + H 2 O + 4e- = 4 OH - Fe 2+ OH- OH - OH - Fe 2+ OH - Fe Fe ANODE CATHODE Removal of reaction products 26 13

14 The rate of an electrochemical reaction cannot be calculated theoretically. It must be measured by using electrochemical techniques, weight change or other methods. The plots showing electrode potential (V) and current density (A/m 2 ) are called polarization curves and they can be used to estimate reaction rate at different potentials. The higher is the current density the higher is the reaction rate DIN sea water, T = 10 C CURRENT DENSITY, ma/cm Mild steel Stainless steel POTENTIAL, mv vs. SCE 28 14

15 i, ma/cm Gelatiinin vaikutus kuparin saostumiseen elektrolyysissä 0 ppm 1 ppm 2 ppm 3 ppm 5 ppm 10 ppm 20 ppm The rate of electrochemical reaction depends on surface potential. The important factor on the reaction rate is not the potential itself but how much the reaction has been polarized from its equilibrium state. E, mv vs. Cu/CuSO

16 In the dynamic equilibrium the rate of the exchange reaction is described by exchange current density i 0. Exchange current density depends on the general rate constant and concentrations of reacting species. i 0 = z F k 0 a c OX (1-a ) c RED Depending on the reaction and solution the values of exchange current densities can be A/cm In the dynamic equilibrium no net reaction happens. The system is on its equilibrium potential E o and the reaction rate is described by exchange current density i 0. To get some net effect in the system, it is necessary to change the potential of the system from its equilibrium value. As system or electrode that is not in its equilibrium state is polarized

17 1. At the equilibrium E o and i 0 2. Potential change E > E o 3. Anodic reaction rate increases, cathodic decreases 4. Net current at E is i = i a i c > 0 33 Butler-Volmer equation 500 i = i 0 Ø Œe º Anodic i 0 e a z F h RT Cathodic i 0 a z F h RT - e -(1-a ) z F h RT partial reaction -(1-a ) z F h RT -e partial reaction ø œ ß CURRENT DENSITY, ma/cm Anodic partial reaction Cathodic partial reaction Sum reaction OVERPOTENTIAL, mv 34 17

18 The deviation of electrode potential from equilibrium potential is polarization and the difference is called overpotential or overvoltage h = E - E 0 Overvoltage h is positive for anode (E > E 0 ) and negative for cathode (E < E 0 ). Current is also positive for anode and negative for cathode. 35 Polarisation is caused by finite rate of all the reaction steps. One of the steps is often slower than others limiting the total reaction rate. Types of overpotential activation overpotential charge transfer diffusion overpotential reaction overpotential crystallization overpotential resistance overpotential mass transfer chemical reaction nucleation and growth into crystal structure voltage drop caused by solution resistance, not related to the reaction mechanism 36 18

19 Activation polarization or activation overvoltage describes the rate of the charge transfer step. A large activation polarization means that it is necessary to change the electrode potential far away from the equilibrium potential to get significant reaction rates. Increasing temperature decreases activation polarization VIRRANTIHEYS, ma/cm i 0 = 10-2 ma/cm 2 i 0 = 10-5 ma/cm YLIPOTENTIAALI, mv 38 19

20 = a + b log(i) 39 = a + b log(i) 40 20

21 41 There are three mechanisms for mass transfer between electrode surface and electrolyte: Migration is the movement of charged particles in an electric field Diffusion is caused by concentration differences and the species diffuse from higher to lower concentration. Convection is natural or forced flow of the solution

22 Mass transfer or flux is the amount of moles through an area in certain time, J [mol/cm 2 /s]. In a liquid phase the mass transfer is the sum of diffusion caused by concentration gradient, migration caused by electric field and convection caused by solution flow. Nernst and Planck equation: J= J diff J = D + J migr + J konv c E + z F c x x + c v

23 When diffusion coefficient and concentration gradient are known it is possible to calculate the diffusion current density. i d x z F DA ( ca -ca) = d and when the surface concentration = 0 the theoretical maximum current or limiting current density is reached. i d z F D = ilim = d A c A 45 h d RT i = ln( 1- ) zf i lim 46 23

24 10 The polarisation phenomena will sum up. Charge transfer, mass transfer and solution resistance will all slow down the total reaction rate by some amount. CURRENT DENSITY, ma/cm 2 1 h a h c h W OVERPOTENTIAL, mv 47 Electrochemical cell For an electrochemical reaction to proceed it is required that the electrons in the electrode move somewhere. This can be done only by creating a system with two different electrode reactions. Two electrode reactions + solution + conductor = electrochemical cell, corrosion cell The system can be spontaneous (corrosion, leaching) or forced (electrolysis)

25 Corrosion cell The metals are inhomogeneous materials in inhomogeneous environment. Spontaneous formation of corrosion pairs is evident. Parts of the structure or surface with high potential become cathodic and on them a reduction reaction happens. Parts with low potential become anodic, and oxidation reaction happens. The cathodic reactions that proceed on the surface start the anodic reactions. 49 Corrosion cell Anode - Oxidation reaction - Transfer of oxidized species to solution e - H 2 Cathode - Transfer of oxidant to the electrode surface - Reduction reaction - Transfer of reduced species from surface back to solution H + Me + Bulk electrolyte with an oxidant (O 2, H + etc.) 50 25

26 Corrosion cell The driving force of a spontaneous electrochemical system is the potential difference between anodic and cathodic reactions. The reaction rate depends on the polarization of the electrodes. The system approaches to a state, where the total current delivered by anodic reactions equals the total current consumed by cathodic reactions. This non-equilibrium situation is described by corrosion potential E corr and corrosion current density i corr. 51 Corrosion cell driving force For a spontaneous electrochemical system the equilibrium potential of the anodic oxidation reaction is lower than that of the cathodic reduction reaction. The electromotoric force EMF is E o (cathode) E o (anode) For a forced system using external current the potential of the anodic oxidation reaction is higher than that of the cathodic reduction reaction. Cell voltage is E o (anode) E o (cathode) 52 26

27 Corrosion cell driving force When analysing spontaneous systems the task is usually to find out if the system is thermodynamically possible. Cell voltage or electromotoric force is EMF = E o (cathode) E o (anode) EMF = E o (reduction) E o (oxidation) If EMF > 0 the system can proceed spontaneously. If EMF < 0 the system is not thermodynamically possible and external current is needed. DG = - z F EMF 53 Corrosion cell driving force Example on calculation of cell voltage. Reactions are H + /H 2 and Fe/Fe 2+ Temperature 25 o C. H 2 H Fe + ( p = 1 atm) H ( a + = 1) Fe( a 2+ 1) H Fe 2 = H e + 2 e - - = H = Fe + Fe = H Fe 2+ Fe 54 27

28 Corrosion cell driving force The systems are in standard state so the cell voltage can be taken as difference of tabulated values E(H + /H 2 ) = 0 V ja E(Fe/Fe 2+ ) = V. Potential is calculated by subtracting the potential of oxidation reaction (iron) from reduction reaction (hydrogen). Cell voltage or electromotoric force or driving force is E = = V. Because E > 0 the reaction proceeds to right producing ferrous ions. Hydrogen evolution as cathodic reaction causes iron to dissolve. 55 Corrosion cell driving force If the concentrations of the reacting species change, the equilibrium potentials will also change. Potentials are calculated using Nernst equation. Example, can acid solution dissolve copper? Assume hydrogen ion concentration 20 mol/l, i.e. ph = -1,3 and copper concentration 10-4 mol/l. 8, E(H 2 / H + ) = ln Ł 8, E(Cu / Cu 2+ ) = 0, = 0, 077V ł ln 1 = 0, 219V Ł10 4 ł 56 28

29 Mixed potential theory Mixed potential theory was developed by Wagner and Traud in 1939 to explain reactions on a corroding metal surface. Graphical analysis of corroding systems was developed by Ulick R. Evans and colleagues in 1930 s, so-called Evans diagrams. Same idea was later adopted in hydrometallurgy and now the diagrams were called Ritchie diagrams after Australian professor Ian M. Ritchie. 57 Mixed potential theory Mixed potential theory states that on the surface are randomly distributed anodic and cathodic sites. The location of theses sites changes with time. In a corrosion cell all electrons released in anodic reaction are consumed in cathodic reaction. When the rate of the anodic or cathodic reaction changes the rate of the other reaction will follow

30 Corrosion cell CURRENT DENSITY, ma/cm i corr E 0 (anode) E corr E 0 (cathode) POTENTIAL, mv 59 Corrosion cell i a decreases 10-2 i / ma/cm x10-4 1x i corr decreases E corr shifts to anodic direction i a =10-3 ma/cm 2 i a =10-4 ma/cm 2 i a =10-5 ma/cm E / mv 60 30

31 Corrosion cell i c =10-4 ma/cm 2 i c =10-5 ma/cm 2 i c =10-6 ma/cm 2 i c decreases i / ma/cm x10-4 1x i corr decreases E corr shifts to cathodic direction E / mv 61 Corrosion cell Evans diagrams E c E c Potential Cathodic control Potential E corr Anodic control E corr E a E a I corr Current I corr Current 62 31

32 Corrosion cell Evans diagrams E c E c Potential E corr Mixed control Potential Ohmic control E a E a I corr Current I corr Current 63 Corrosion cell Evans diagrams E c E c Potential E' c Potential E' c E corr E' a E' a E a E a I corr I' corr I corr I' corr Current Current 64 32

33 Corrosion cell Evans diagrams 10 2 CURRENT DENSITY, ma/cm i corr E corr E corr (noble) i corr (noble) E cathode E anode E anode (noble) POTENTIAL, mv 65 Corrosion cell Evans diagrams 10 CURRENT DENSITY, ma/cm E anode i corr i corr i lim decreases E corr decreases E cathode POTENTIAL, mv 66 33

34 Leaching electrochemistry Anode reaction E o, mv Cathode reaction MnO H 2 O = MnO H e H 2 O H e - = 2 H 2 O 1501 O H e - = 3 H 2 O Au = Au e Cl e- = 2 Cl O H e - = 2 H 2 O Pt = Pt e Fe 3+ + e - = Fe 2+ CuS = Cu 2+ + S + 2e Cu 2 S = 2 Cu 2+ + S+ 4e CuFeS 2 = Cu 2+ + Fe S+ 4 e FeS 2 = Fe S+ 2e ZnS = Zn 2+ + S + 2e Cu 2+ + e - = Cu + FeS = Fe 2+ + S + 2e H e - = H 2 67 Sulfide mineral dissolution The extraction of metals from sulfides requires oxidation of sulfur to elemental sulfur or sulfate. Oxidation sulfur to higher oxidatlon states (0, +2, +4 and +6) can occur if the redox-potential of the oxidant in solution, e.g. O 2, Fe 3+, is higher than the equilibrium potential of the sulfide. Sulfide oxidation S 2- = S e - e - Me 2+ Oxygen reduction O H e - = 2 H 2 O H 2 O O 2, H

35 Hydrogen reduction 69 Summary Reactions happen in a corrosion cell, driving force or cell voltage is the potential difference between anodic and cathodic reactions. Reaction rates depend on polarization, low polarization preferred in hydrometallurgy, high in corrosion. Current densities are converted to amount of reacted material using Faradays law. Reacted material either as mass change / area / time or thickness change / time. This conversion is done using material density

CHAPTER 6 Modern Theory Principles LECTURER SAHEB M. MAHDI

CHAPTER 6 Modern Theory Principles LECTURER SAHEB M. MAHDI Modern Theory principles in Corrosion and their applications :- Corrosion studies can be carried-out by two methods 1 Thermodynamics. or 2 By