Corrosion reactions anodic

In a battery, the anode and cathode reactions occur ia different compartments, kept apart by a separator that allows only ionic, not electronic conduction. The only way for the cell reactions to occur is to mn the electrons through an external circuit so that electrons travel from the anode to the cathode. But ia the corrosion reaction the anode and cathode reactions, equations 8 and 12 respectively, occur at different locations within the anode. Because the anode is a single, electrically conductive mass, the electrons produced ia the anode reaction travel easily to the site of the cathode reaction and the 2iac acts like a battery where the positive and negative terminals are shorted together. 

The corrosion rate is controlled mainly hy cathodic reaction rates. Cathodic Reactions 5.2 and 5.3 are usually much slower than anodic Reaction 5.1. The slower reaction controls the corrosion rate. If water pH is depressed. Reaction 5.3 is favored, speeding attack. If oxygen concentration is high. Reaction 5.2 is aided, also increasing wastage hy a process called depolarization. Depolarization is simply hydrogen-ion removal from solution near the cathode. 

The current I is called the total current. In free corrosion, i.e., without the contribution of external currents (see Fig. 2-1), it is always zero, as given by Eq. (2-8). and are known as the anodic and cathodic partial currents. According to Eq. (2-10), generally in electrolytic corrosion anodic total currents and/or cathodic redox reactions are responsible. 

Passivation—a reduction of the anodic reaction rate of an electrode involved in an electrochemical reaction, such as corrosion. 

The ion S " reacts with ferrous Fe ion to form black iron sulfide FeS corrosion product. The hydrogen ions are reduced by electrons produced by anodic reaction in step 1 and form hydrogen atom H ... 

All these types of polarisation will be present to a greater or lesser extent in most corrosion reactions, but if one is more significant than the others it will control the rate of the reaction. This leads to a classification of corrosion reactions according to whether the cathodic or anodic reaction is rate... 

In Section 1.4 it was pointed out that a fundamental principle of corrosion is that the sum of the rates of the cathodic reactions must equal the sum of the rates of the anodic reactions, irrespective of whether the attack is uniform or localised ... 

From these two examples, which as will be seen subsequently, present a very oversimplified picture of the actual situation, it is evident that macroheterogeneities can lead to localised attack by forming a large cathode/small anode corrosion cell. For localised attack to proceed, an ample and continuous supply of the electron acceptor (dissolved oxygen in the example, but other species such as the ion and Cu can act in a similar manner) must be present at the cathode surface, and the anodic reaction must not be stifled by the formation of protective films of corrosion products. In general, localised attack is more prevalent in near-neutral solutions in which dissolved oxygen is the cathode reactant thus in a strongly acid solution the millscale would be removed by reductive dissolution see Section 11.2) and attack would become uniform. 

Metals immersed or partly immersed in water tend to corrode because of their thermodynamic instability. Natural waters contain dissolved solids and gases and sometimes colloidal or suspended matter all these may affect the corrosive projjerties of the water in relation to the metals with which it is in contact. The effect may be either one of stimulation or one of suppression, and it may affect either the cathodic or the anodic reaction more rarely there may be a general blanketing effect. Some metals form a natural protective film in water and the corrosiveness of the water to these metals depends on whether or not the dissolved materials it contains assist in the maintenance of a self-healing film. 

Apart from this effect, increased velocity usually increases corrosion rates by removing corrosion products which otherwise might stifle the anodic reaction and, by providing more oxygen, may stimulate the cathodic reaction... 

This is a simplified treatment but it serves to illustrate the electrochemical nature of rusting and the essential parts played by moisture and oxygen. The kinetics of the process are influenced by a number of factors, which will be discussed later. Although the presence of oxygen is usually essential, severe corrosion may occur under anaerobic conditions in the presence of sulphate-reducing bacteria Desulphovibrio desulphuricans) which are present in soils and water. The anodic reaction is the same, i.e. the formation of ferrous ions. The cathodic reaction is complex but it results in the reduction of inorganic sulphates to sulphides and the eventual formation of rust and ferrous sulphide (FeS). 

Sulphur dioxide in the air originates from the combustion of fuel and influences rusting in a number of ways. For example, Russian workers consider that it acts as a cathodic depolariser , which is far more effective than dissolved oxygen in stimulating the corrosion rate. However, it is the series of anodic reactions culminating in the formation of ferrous sulphate that are generally considered to be of particular importance. Sulphur dioxide in the air is oxidised to sulphur trioxide, which reacts with moisture to form sulphuric acid, and this in turn reacts with the steel to form ferrous sulphate. Examination of rust Aims formed in industrial atmospheres have shown that 5% or more of the rust is present in the form of iron sulphates and FeS04 4H2 0 has been identified in shallow pits . 

Equation 10.2, which involves consumption of the metal and release of electrons, is termed an anodic reaction. Equation 10.3, which represents consumption of electrons and dissolved species in the environment, is termed a cathodic reaction. Whenever spontaneous corrosion reactions occur, all the electrons released in the anodic reaction are consumed in the cathodic reaction no excess or deficiency is found. Moreover, the metal normally takes up a more or less uniform electrode potential, often called the corrosion or mixed potential (Ecotr)-... 

Fig. 10.1 Schematic illustration of the corrosion of steel in an aerated environment. Note that the electrons released in the anodic reaction are consumed quantitatively in the cathodic reaction, and that the anodic and cathodic products may react to produce Fe(OH)2...

The corrosion reaction may also be represented on a polarisation diagram (Fig. 10.4). The diagram shows how the rates of the anodic and cathodic reactions (both expressed in terms of current flow, I) vary with electrode potential, E. Thus at , the net rate of the anodic reaction is zero and it increases as the potential becomes more positive. At the net rate of the cathodic reaction is zero and it increases as the potential becomes more negative. (To be able to represent the anodic and cathodic reaction rates on the same axis, the modulus of the current has been drawn.) The two reaction rates are electrically equivalent at E , the corrosion potential, and the... 

The thermodynamic and electrode-kinetic principles of cathodic protection have been discussed at some length in Section 10.1. It has been shown that, if electrons are supplied to the metal/electrolyte solution interface, the rate of the cathodic reaction is increased whilst the rate of the anodic reaction is decreased. Thus, corrosion is reduced. Concomitantly, the electrode potential of the metal becomes more negative. Corrosion may be prevented entirely if the rate of electron supply is such that the potential of the metal is lowered to the value where it is found that anodic dissolution does not occur. This may not necessarily be the potential at which dissolution is thermodynamically impossible. 

It should be noted that when metals like zinc and aluminium are used as sacrificial anodes the anode reaction will be predominantly 10.18a and 10.186, although self-corrosion may also occur to a greater or lesser extent. Whereas the e.m.f. between magnesium, the most negative sacrificial anode, and iron is =0-7 V, the e.m.f. of power-impressed systems can range from 6 V to 50 V or more, depending on the power source employed. Thus, whereas sacrificial anodes are normally restricted to environments having a resistivity of <6 000 0 cm there is no similar limitation in the use of power-impressed systems. 

The overall process is metal transfer from anode to cathode via the solution. The form of anode corrosion is important, and materials may be added both to the anode metal and to the electrolyte, to influence it. There are important instances where an insoluble anode is used, and the anode reaction becomes the oxidation of water or hydroxyl ions ... 

In order to inhibit corrosion, it is necessary to stop the flow of current. This can be achieved by suppressing either the cathodic or the anodic reaction, or by inserting a high resistance in the electrolytic path of the corrosion current. These three methods of suppression are called cathodic, anodic znA resistance inhibition respectively (Section 1.4). 

It has been shown that paint films are so permeable to water and oxygen that they cannot affect the cathodic reaction, and that the anodic reaction may be modified by certain pigments. There are, however, many types of protective paint which do not contain inhibitive pigments. It is concluded that this class of paint prevents corrosion by virtue of its high ionic resistance, which impedes the movement of ions and thereby reduces the corrosion current to a very small value. 

Evans Diagram diagram in which the E vs. I relationships for the cathodic and anodic reactions of a corrosion reaction are drawn as straight lines intersecting at the corrosion potential, thus indicating the corrosion current associated with the reaction. 

Occluded Cell a corrosion cell of a geometry that prevents intermingling of the anodic reaction products (anolyte) with the bulk solution, resulting in a decrease in pH of the anolyte shielded areas or pits, crevices or cracks in the surface of the metal are examples. 

Partial Reactions anodic reaction (reactions) and cathodic reaction (reactions) constituting a single exchange process or a corrosion reaction. 

In acid electrolytes, carbon is a poor electrocatalyst for oxygen evolution at potentials where carbon corrosion occurs. However, in alkaline electrolytes carbon is sufficiently electrocatalytically active for oxygen evolution to occur simultaneously with carbon corrosion at potentials corresponding to charge conditions for a bifunctional air electrode in metal/air batteries. In this situation, oxygen evolution is the dominant anodic reaction, thus complicating the measurement of carbon corrosion. Ross and co-workers [30] developed experimental techniques to overcome this difficulty. Their results with acetylene black in 30 wt% KOH showed that substantial amounts of CO in addition to C02 (carbonate species) and 02, are... 

The cathodic reactions normally are slower than the anodic reactions and are therefore the corrosion rate-determining steps. 

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