Proceso de producción de alimento balanceado

In oil gas pipelines when CO2 is present in produced fluids, carbon steel

corrodes [5]. CO2 corrosion, also referred to as sweet corrosion, is one of the

most significant mechanisms of failure of oil and gas pipelines, as shown previously [5]. Other mechanisms of corrosion exist, such as hydrogen sulphide corrosion, preferential weld corrosion and galvanic corrosion that were beyond the scope of this thesis that contribute to corrosion failures of oil and gas pipelines [5]. The fundamental mechanisms of aqueous corrosion of a metal are explained in this section.

2.4.1 Fundamentals of Aqueous Corrosion

Corrosion is the destructive attack of a metal by electrochemical reaction with its environment [22]. Corrosion degradation can be broadly classified into two categories, uniform (or general) corrosion and localised corrosion. Uniform corrosion occurs when the corrosion mechanisms are consistent across an entire surface area, whereas localised corrosion occurs over much smaller regions on a surface [5, 23, 24]. An electrochemical corrosion reaction requires an anode, a cathode, a metallic conductor and an electrolytic conductor, with corrosion reactions separated into anodic and cathodic

reactions [23]. A typical anodic reaction of a corroding metal is defined by Equation (2.2):

𝑀 → 𝑀𝑛+_{+ 𝑛𝑒}− _(2.2)

where 𝑛 is the number of electrons (𝑒−) released by the metal, 𝑀. The metal ions are transported from the anode to the cathode through the electrolyte, typically a liquid containing a concentration of electrochemically active species [23]. Anodic and cathodic sites can exist on the same metal sample, which also acts as the metallic conductor, providing three of the four components required for electrochemical corrosion [23]. The final component, the electrolyte, consists of anions, negatively charged ions that move towards the anode, and cations, positively charged ions that move towards the cathode, where they are reduced [23]. The electrons produced in the anodic reaction remain on the corroding metal and migrate to the cathode [25]. The cathodic reaction consumes the electrons, achieved through reducible species in the electrolyte adsorbing onto the metal surface, removing the electrons [25]. Metals corrode as they have a tendency to return to a low energy state after being extracted from their ore [23]. The tendency of a metal to corrode is defined by the Gibbs free energy change (𝛥𝐺) [22]. If the free energy change as a result of a reaction is negative, then the reaction is in a lower energy state and is more stable, so the reaction occurs spontaneously. If the free energy change is positive, then the reaction does not occur spontaneously [23]. The more negative the Gibbs free energy, the greater the tendency of the reaction to occur [22]. The Gibbs free energy change is defined by Equation (2.3):

𝛥_𝑟𝐺̅ = −𝑛𝐹𝐸 (2.3)

where Δ_𝑟𝐺̅ is the electrochemical Gibbs energy change, 𝐸 is the potential, 𝑛 is the number of electrons and 𝐹 is the Faraday constant. The Gibbs free energy at standard conditions of ambient temperature and atmospheric pressure is defined by Equation (2.4):

𝛥_𝑟𝐺̅̅̅̅ = −𝑛𝐹𝐸0 0 _(2.4)

where 𝛥_𝑟𝐺̅̅̅̅0 is the standard electrochemical Gibbs energy change, 𝐸0 is the standard redox potential, defined as the potential of a metal in contact with its own ions at a concentration equal to unit activity at 25°C [23]. The Nernst equation is used to determine the potential of a metal in a solution in which the ions are not at unit activity or in a solution of ions other than its own, Equation (2.5) [22]:

E = 𝐸0₋𝑅𝑇

𝑛𝐹ln

[𝑐_{𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠}] [𝑐𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠]

(2.5) where 𝑅 is the universal gas constant, 𝑐𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 is the concentration of the

product species and 𝑐𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠 is the concentration of the reactant species. The

Nernst equation does not contain a current term, which is typically of interest from a corrosion perspective as it can be used to calculate the corrosion rate of a metal [24]. The Buttler-Volmer equation is used to relate electrical current to changes in metal potential from an external source, Equation (2.6) [24]:

𝑖 = 𝑖_{𝑐𝑜𝑟𝑟}[𝑒((1−𝛼)𝑛𝐹(𝐸−𝐸𝑅𝑇 𝑐𝑜𝑟𝑟))_{− 𝑒}(−𝛼𝑛𝐹(𝐸−𝐸𝑅𝑇 𝑐𝑜𝑟𝑟))_] (2.6)

where i is the external current density flowing to or from an electrode as a result of an applied potential, icorr is the corrosion current density, α is a

coefficient ranging from 0 to 1 and Ecorr is the free corrosion potential.

2.4.2 Electrical Double Layer

When a metal corrodes, metal atoms are removed from their lattice sites to ionise as cations, leaving a negatively charged metal surface due to excess electrons [26]. Water molecules surround the metal ions as they leave the lattice, hydrating them. Some of the positively charged metal ions remain at the surface due to the negative charge [24]. The water layer around the ions prevents them from making contact with excess electrons and reducing to form metal atoms. Positively charged ions present in the bulk solution, are also attracted to the negatively charged metal surface [26]. The electrolyte layer adjacent to an electrode surface contains water molecules and ions from the metal and the bulk electrolyte. The metal surface and the adjacent electrolyte is referred to as the electrical double layer (EDL), shown in Figure 2.2, with the typical components of a representative electrical circuit of the EDL also shown [24].

The EDL is typically represented as a capacitor due to the physical separation of two oppositely charged planes. The EDL also acts as a resistor, as the metal resists transferring excess electrons to the electrochemically active species. Electrochemically active species diffuse from the bulk solution to the metal surface and discharge the EDL [24].

Figure 2.2 EDL and equivalent circuit consisting of a solution resistance (Rs),

charge transfer resistance (Rct) and capacitance (Cedl) [24, 26]