Kuliah 2. Aspek Elektrokimia S2-20.pptx

Basic Theory of aqueous corrosion
Electrochemical aspect
Thermodynamics aspect
Kinetics aspect

1. Aspek Elektrokimia Kuliah 2 Korosi dan Perlindungan Logam

2. Topics : Basic Theory of aqueous corrosion Electrochemical aspect Thermodynamics aspect Kinetics aspect

3. Elektrokimia adalah ilmu yang mempelajari fenomena pada antarmuka elektroda - larutan Interface ( antarmuka ) elektroda larutan

4. Gambar 1, sel elektrokimia pada sel korosi , electron mengalir melalui kawat logam penghantar dari sisi yg terjadi reaksi anodic ke tempat terjadinya reaksi katodik . Ion2 ( partikel muatan mengalir melalui elektrolit untuk supaya aliran elektronnya seimbang . Anion2 ( muatan ion yang negative dari reaksi katodik ) mengalir kearah anoda , dan kation2 ( ion bermuatan positif dari anodanya sendiri ) mengalir kearah katoda . Dengan demikian Anoda terkorosi , katoda tidak terkorosi . Disini juga terjadi perbedaan potensial antara Anoda dan Katoda . Fe  Fe 2+ + 2e 2H + + 2e  H 2

5. Copy Fig 2.4 zaki

6. electrochemistry ionic electrodic Electrolyte thermodynamics Mass Transport & Conductivity Electrode thermodynamics Electrode Kinetics The Division Of Electrochemistry

7. Reaksi Korosi Logam Reaksi Anodik : ciri : 1. oksidasi logam menjadi ionnya dan muatannya Fe –Fe 2+ + e 2. melepas electron 3. valensinya meningkat ( Fe dg valensi teroksidasi jadi valensi +2) Reaksi oksidasi untuk logam dan paduannya menunjukkan terjadinya korosi . Maka , Kalau reaksi oksidasi di stop, maka korosi juga stop

8. Reaksi Katodik : reaksi katodik = reaksi reduksi pada katoda . electron yang lepas dari anoda di konsumsi pada permukaan katoda . maka , valensi pada reaksi katodik akan berkurang . Reaksi katodik yang umum : 2H + + 2 e  H 2 ( dalam larutan asam ) O 2 + 4H + + 4e  2 H 2 O ( dalam larutan asam ) 2H 2 O + O 2 + 4e  4 OH - ( dalam larutan pH netral dan basa ) Fe 3+ + e  Fe 2+ ( reduksi ion logam dalam larutan garam ferric cl)) Pengendapan loga : M 2+ + 2e  M ; Ni 2+ + 2e  Ni Reduksi Bakteri Sulfat : SO 2- 4 + 8 H + + 8e  S - + 4H 2 O

9. Jenis-jenis Sel KOROSI bias beragam : (1) Galvanic Cells (2) Concentration Cell (3) Electrolytic Cells (4) Different temperature Cells

10. Galvanic Cells Sel Galvanik : Anoda dan Katoda beda logam di dalam suatu larutan . Anoda dan Katoda sama logamnya , dalam kondisi elektrolit beda Logam yang lebih mulia sebagai katoda Arus mengalir dari Anoda ke Katoda didalam elektrolit .

11. Concentration Cells : Seperti sel galvanic, Anoda dan Katoda logamnya sama , lingkungannya elektrolit heterogen . Pipa di dalam tanah : Ada kandungan oksigen yang berbeda di dalam tanah Ada perbedaan moisture di dalam tanah Ada komposisi yang berbeda di dalam tanah Fig 2.6 Zaki p 13

12. Electrolytic Cells Electrolytic cell terjadi saat ada arus dari luar ke dalam system. Anoda pada kutub (+) Katoda Kutub (-) KNAP Sel elektrolitik ini biasa untuk aplikasi CP

13. Proses korosi digambarkan sebagai proses elektrokimia dimana terjadi - peristiwa oksidasi di anoda dan - peristiwa reduksi di katoda .

14. Mechanism of Corrosion Iron 1. In corrosion Electrodes are pieces of metal on which an electrochemical reaction is occurring. From definition of corrosion is the degradation of metals resulting from their chemical interaction with the environment. Example : a mild steel immersed in acid solution The metal in this environment is not stable and shows a natural tendency toward spontaneous oxidation. Fe2+ lepas dari anoda karena oksidasi dan OH- direduksi katodik pada permukaan logam

15. Zaki p 14 reaksi

16. Concept of Free Energy Hubungan antara free energy dan konstanta kesetimbangan .

17. Oxidation of mild steel in contact with an acid solution : Fe  Fe++ + 2e Each atom of metal loses two electrons. The reaction occurs at the metal-solution interface. The positevely charge ion ( cations , Fe++) is ejected from the metal and goes into solution, while the two liberated electrons remain in the metal. The result is separation of electrical charge. The cation leaves the metal, then form a positively charged layer on solution side, while the electrons in metal form a negatively charge layer. So this creates an electrical double layer .

18. This mechanism of double layer formation useful approach to amodern understanding of electrodes and their process. Moreover, many corrosion phenomena can be adequately discussed in these terms.

19. Electrochemical Nature of aqueous corrosion Metallic corrosion process involve transfer of electronic charge in aqueous solutions Thus, it is necessary to discuss ‘the electrochemical nature of corrosion’ before discussing of the various form of corrosion.

20. Electrochemical reactions : Example of corrosion between Fe and hydrochloric acid, represented by : Fe + 2 HCl  FeCl 2 + H 2 (1) Fe reacts with the acid solution forming soluble Fe chloride and liberating hydrogen bubbles on the surface.

21. Acids and Bases An acid is a substance that produces excess hydrogen ions (H + ) when dissolved in water examples are HCl, H 2 SO 4 A base is a substance that produces excess hydroxyl ions (OH - ) when dissolved in water examples are NaOH, KOH

22. Corrosion of Zinc in Acid Zinc dissolves with hydrogen evolution Zn + 2HCl  ZnCl 2 + H 2 Zinc known as a base or active metal One atom of zinc metal plus two molecules of hydrogen chloride (hydrochloric acid) reacts to form goes to one molecule of zinc chloride plus one molecule of hydrogen gas

23. Anodic Reactions Examples Zn  Zn 2+ + 2e - zinc corrosion Fe  Fe 2+ + 2e - iron corrosion Al  Al 3+ + 3e - aluminium corrosion Fe 2+  Fe 3+ + e - ferrous ion oxidation H 2  2H + + 2e - hydrogen oxidation 2H 2 O  O 2 + 4H + + 4e - oxygen evolution Oxidation reactions Produce electrons

24. Cathodic reactions : Hydrogen evolution 2H + + 2e  H 2 Oxygen reduction (acid) O 2 + 4H + + 4e  2H 2 O Oxygen reduction(basic): O 2 + 2H 2 O+4e4OH - Metal reduction : M 3+ + e  M 2+ Metal deposition ; M + + e  M read Fontana p 17

25. Electrochemistry of corrosion Permukaan logam yg sama diekspose di aqueous electrolyte Biasanya memiliki sisi untuk suatu oksidasi dan reduksi . Sisi ini membentuk corrosion cell

26. The electrode - solution interface is also hydrated because of the charge effects. ( N.B. the potential difference between metal and solution is of the order of  1V, remember, - cannot be measured - but over a molecular distance amounts to an electric field of  10 8 V/cm). Schematic of a charged interface and the locations of cations at the electrode surface.

27. Because of adsorbed H 2 O on surface and primary hydration shell, ions can only approach to within a fixed distance from the electrode: the OUTER HELMHOLTZ PLANE . The resulting double layer of charge is often treated theoretically as a capacitor . . . . Simplified double layer at a metal aqueous interface.

28. Voltaic Cells In spontaneous oxidation-reduction (redox) reactions, electrons are transferred and energy is released.

29. Voltaic Cells In the cell, then, electrons leave the anode and flow through the wire to the cathode. As the electrons leave the anode, the cations formed dissolve into the solution in the anode compartment.

30. Voltaic Cells As the electrons reach the cathode, cations in the cathode are attracted to the now negative cathode. The electrons are taken by the cation, and the neutral metal is deposited on the cathode.

31. Electromotive Force (emf) Water only spontaneously flows one way in a waterfall. Likewise, electrons only spontaneously flow one way in a redox reaction—from higher to lower potential energy.

32. Electromotive Force ( emf ) dan Ecell The potential difference between the anode and cathode in a cell is called the electromotive force ( emf ) . It is also called the cell potential , and is designated E cell . Any work performed can only be done through a decrease in Free Energy of the cell reaction, hence: ΔG = - nFE ΔG = free energy change (joules) E =standard potential of the reaction (volts)

33. Energi bebas reaksi tidak mungkin untuk mendapat nilai absolutnya . Tapi , perubahan energi bebas dapat diukur . Bila nilai Δ Go adalah negatif , maka reaksi mungkin terjadi . Tanda (-) atau (+) dari Δ Go hanya memprediksi kespontanan reaksi . Emf dan Energi bebas

34. REVERSIBLE ELECTRODE POTENTIAL Adanya reaksi reversibel antara logam dengan elektrolitnya , akan menimbulkan Pembentukan lapis ganda (double layer)

35. REVERSIBLE POTENTIAL AND STANDARD POTENTIAL Erev : Potensial yang timbul saat logam dicelup dalam larutan yang mengandung ion logamnya . Contoh : Zn dicelup dalam ZnSO4 Eo : Potensial logam murni yang diukur dengan referensi SHE pada keadaan standard (P,T)