D E P A R T M E N T O F M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G M A S S A C H U S E T T S I N S T I T U T E O F T E C H N O L O G Y 3.014 Materials Laboratory Fall 2006 LABORATORY 3: Electrochemical Corrosion Gibbs Free Energy, Anodic Corrosion & the EMF Series Instructor: Professor Linn W. Hobbs Objectives • Learn about galvanic (anodic) corrosion and the driving force that causes galvanic corrosion of a metal. • Understand the connection between a change in Gibbs free energy and the cell potential (emf) in an electrochemical or corrosion reaction. • Appreciate the active (anodic) and noble (cathodic) behavior of a selection of common metals. • Explore the effectiveness of sacrificial anodes in corrosion control and prevention. Materials and Equipment Ag, Al, Au, Co, Cu, Fe, Mg, Ni, Pb, Pt, Zn and brass foils 1018-steel cylinder Piece of scrap iron Darkened (sulfidized) silver object 220-grit silicon carbide abrasive paper NaCl (rocksalt) powder NaHCO3 (sodium bicarbonate) powder De-ionized water Digital voltmeters, ammeters Digital power supply Digital thermometer Stopwatch Two electrochemical cells Magnetic stirrer and heater Background Corrosion is a trillion-dollar problem worldwide, most commonly associated with metals and alloys, but also affecting ceramic and polymeric materials. Corrosion is generally divided into -1-three categories: high-temperature gaseous (dry) corrosion (oxidation, sulfidation); aqueous (electrochemical) corrosion; and biological corrosion (attack by microbes). In this laboratory we will investigate the second of these, aqueous electrochemical corrosion of metals and alloys. Corrosion involves chemical reactions, in which the original material reacts with a chemical agent, to form a new compound or to dissolve into the chemical agent, in both cases involving charge transfer. Corrosion in aqueous solutions is an electrochemical reaction that involves charge transfer, either oxidation, involving loss of electrons; or reduction involving gain of electrons. As shown in fig. 1, corrosion is a metal dissolution process, with metal atoms being converted to metal ions and going into solution. M2+H+H+H2H+H+H+H+H2OH2OH2OH2OH+H+H2OH2OH2ONeutralAcidicSolutionH2OOH-H2e-e-M2+M2+H+H+H+H+H2H2H+H+H+H+H+H+H+H+H2OH2OH2OH2OH2OH2OH2OH2OH+H+H+H+H2OH2OH2OH2OH2OH2ONeutralAcidicSolutionH2OH2OOH-OH-H2H2e-e-Fig. 1 Electrochemical attack of a metal surface immersed in an aqueous medium (after Fontana and Greene1). The products of the reaction depend on the pH (= –log[H+]) of the aqueous solution, which is a measure of the H+ hydrogen ion concentration. Notice the dissolution involves the transfer of charge, in this case electrons. At the metal surface, neutral metal atoms are oxidized to positive metal ions in what is known as an oxidation or anodic reaction: Oxidation or Anodic Reaction M → Mn+ + ne– (1) (e.g. Fe → Fe2+ + 2e–, Zn → Zn2+ + 2e–) The corresponding reduction reactions depend on the nature of the environment, particularly the pH. Typical reduction, or cathodic, reactions are: Reduction or Cathodic Reactions Hydrogen 2H+ + 2e– → H2 (acidic solutions) (2a) -2-Oxygen O2 + 4H+ + 4e– → 2H2O (acidic solutions) (2b) O2 (dissolved) +2H2O + 4e– → OH– (neutral/basic solutions) (2c) Water 2H2O + 2e– → H2 +OH– (neutral/basic solutions) (2d) Corrosion measurements are quantified by constructing a corrosion cell (an electrochemical cell, Fig. 2) in which the oxidation and reduction reactions generally take place at separate electrodes in the cell, which develop an electrical potential difference between them. The (open circuit) cell potential is a measure of the tendency of a metal to corrode. When the two electrodes are in electrical contact, an electrical circuit is formed in which electron current flows through the electrical connection between the electrodes and a corresponding ion current flows through the electrolyte between the electrodes. The current flow is a measure of the corrosion taking place at the anode. ELECTROCHEMICAL CELL Anode: Oxidation reaction site (negative) +++Cathode: Reduction reaction site (positive) E: Cell potential or cell emf (volts) Fig. 2. An electrochemical cell. Relationship between Gibbs Free Energy and the Cell Potential1,2 One can assert the tendency of a metal to corrode from thermodynamic considerations. The change ΔG in Gibbs free energy for the corrosion reaction predicts whether or not a corrosion reaction occurs spontaneously; it does so spontaneously if ΔG < 0. The change in free energy can be calculated from a measurement of the cell potential E. The maximum amount of electrical energy (or work done) that can be delivered by an electrochemical cell in a given state is nF E, which is equivalent to the change in Gibbs free energy ΔG = –nF E, (1) Cell E - + e -anion flow cation flow -Anode (oxidation)Cathode (reduction) - +e-anion flowcation flow--Anode (oxidation)Cathode (reduction)Electrolyte -3-where n is the number of moles of electrons exchanged in an electrochemical reaction, F is Faraday’s constant (96,485 C /mol), and E is the cell potential (in volts) for the cell in a given state. For cell conditions in a standard state, ΔG0 = –nF E0 (2) where E0 represents the standard-state electrochemical cell potential, and ΔG0 represents the Gibbs free energy change for constituents in their standard states. The standard Electromotive Force (emf) series for metals shown in most text books1,2 registers values for E0. How the cell potential varies with cell conditions is established by the Nernst equation. The Nernst equation codifies the fundamental relationship in electrochemical reactions that expresses the electrochemical cell potential in terms of reactants and products of the reaction. It can be derived based on Gibbs free energy criterion for chemical reactions. 2 For a general chemical reaction, the change in Gibbs free energy is related to the activities of the reactants and products of reaction, as follows: ΔG – ΔG0= RT ln (aproducts / areactants ), or (3) ΔG – ΔG0 = 2.303 RT log10 (aproducts / areactants ), (4) where ΔG and ΔG0 represent changes in the free energy of products and reactants in non-standard and standard states, respectively; R