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The corrosion process in Figure 2-1 is an example of a spontaneous electrochemical cell reaction. The driving force of the reaction is a reversible cell voltage, expressed as the difference between the potentials of anode and cathode [16]. Reversible cell voltage and corresponding reversible cell potentials of the electrodes are determined by thermodynamic properties.

2.3.1

The standard hydrogen potential

In a corrosion cell, the potential difference between the anode and cathode can be measured using a voltage measuring device, but the absolute potential of the anode and cathode cannot be measured directly. In order to measure the potential of an electrode, it is necessary to compare this to a reference electrode (the standard hydrogen electrode, SHE). The hydrogen potential has been taken as a standard potential for convenient reference points for measuring and comparing the relative affinities of chemical substances for electrons under specified conditions. Therefore, the potential of hydrogen is taken to be zero at all temperatures. However, this study will consider the Ag/AgCl electrode as the reference electrode and platinum as the auxiliary electrode to measure and compare the potential of the working electrode (carbon steel) in a solution of 3.5% wt NaCl and HAc.

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Figure 2-3: Standard hydrogen electrode, SHE.

The standard hydrogen electrode consists of a platinum electrode in a solution containing H+ ions as shown in Figure 2-3. The solution (H2SO4) has a concentration of 1 mol.dm3.

As the hydrogen gas bubbles over the platinum electrode, the reaction becomes: 2-7

2.3.2

The electrode potential at equilibrium

The electrochemical cell provides information on the equilibrium electrode potential, εeq .

At each metal-electrolyte interface, there is a tendency of metal ions from the solution to deposit on the metal electrode to make it positively charged. Alternatively, metal atoms of the electrode have a tendency to migrate into the solution as ions and leave behind the electrons at the electrode trying to make it negatively charged. Eventually, a point is reached when equilibrium is established where the rate of positively charge builds up is

platinum wire hydrogen at 1 bar salt bridge Magnesium Dilute sulphuric acid [H+]= 1 mol dm-3 [H Magnesium sulphate solution [Mg2+] = 1 mol dm-3 high resistance voltmeter Platinum toil covered in porous platinum

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equal to the negative charge. At this point, a potential difference develops between the electrode and the solution.

2.3.3

The electrode potential during corrosion

In measuring corrosion kinetics, ε – εeq is the difference between the electrode potential

for the metal and a solution where the rates of forward and backward reactions are not the same and the potential at equilibrium. In order to measure the electrode potential ε, an alternative experimental set up in Figure 2-4 is required which includes the voltmeter. In Figure 2-4, the voltmeter measures the electrode potential, ε while corrosion is taking place in the test metal M. A current I is applied that measured by the ammeter represents the corrosion rate which is a function of current and the concentration M2+.



z

M

c

Figure 2-4: Electrochemical cell for electrode potential measurements for a non- equilibrium half-cell

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2.3.4

The electrical double layer (EDL)

When a solid emerges in a polar solvent or an electrolyte solution, a redistribution of charges at the metal-electrolyte interface will develop as shown in Figure 2-5. It is well noticed that this interface will alter the solutions as the interactions between the electrode and the solution is quite different to those in the solution. In addition, there is an influence of the charge transfer for the electrode under potentiostatic measurements. These however lead to strong interactions occurring between the ions/molecules in electrolyte and the electrode surface. This region is known as the electrical double layer The EDL is referred to as a special region of an electrode-electrolyte surface that contains a negative charge of electron and positive charge ions separated the electrolyte in a metal electrochemical reaction [17]. Different models have been developed to explain the process observed when electrochemical measurements are performed in electrolyte solutions.

The Barlow and Erdey-Gruz [18, 19] models are used to explain the interfacial structure of the electrical double layer known as the Helmholtz layer behaving like a charged capacitor as shown in Figure 2-5. The Figure 2-5 shows that some negatively charge ions are absorbed on the metal-electrode surface and polar water covers the rest of the surface, forming a protective layer. The following can also be deducted from Figure 2-5.

 The Inner Helmholtz Plane (IHP) is an ionic layer that consists of adsorbed dipole H2O molecules. The majority of the anions do not penetrate this layer but some

does. The inner potential on the boundary of this ionic plane is Ø1.

 The Outer Helmholtz Plane (OHP) consists of a plane of adsorbed ions due to electrostatic forces in contact with a diffuse ionic layer at an inner potential Ø2

 The diffuse layer (DL) is a thick layer located in a region of diffusely ions in contact with the OHP and the bulk of solution at a potential range of Ø1< Ø1diffuse

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In view of this, the present study will relate the principle of electrical double layer to explain the behaviour that will be encounter when linear polarization resistance, LPR, electrochemical impedance spectroscopy, EIS and potentiodynamic scan, PDP measurements are performed on carbon steel electrode in 3.5% wt. NaCl solution containing different concentrations of acetic acid and mono-ethylene glycol.

Figure 2-5: Electrical double layer at a metal-solution interface [17].

2.3.5 The relationship between EDL chemistry, voltage and current

However, unlike a chemical reaction which depends on the temperature (Arrhenius equation), an electrochemical reaction depends directly on the applied potential at the electrodes. The Butler-volmer equation is used to express the relationship between applied potential to the current as follows:

At the electrode surface, the redox reaction occurs as:

2-8

The reaction rate of the forward direction Kf and backward direction Kb depends on the

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2-9

2-10 In equation 2-9 and equation 2-10, Kf is the forward reaction (metal deposition, cathodic),

Kb is the backward reaction (metal dissolution, anodic), E0 is the standard potential of the

redox reaction, E is the applied potential, R is the gas constant, F is the Faraday constant, T is the temperature, n is the number of electrons transferred (valency), K0 is the standard

rate constant and α is the transfer coefficient.

Combining equation 2-9 and equation 2-10 together gives:

2-11

Where, A is the surface area of the electrode, [Ox]and [Red] are the concentrations of the oxidation and reduction.