Stock solutions preparation

NaClO₄ was obtained from Fluka. KNO₃, Cd and CdCl₂H₂O (Fw: atomic absorption standard (98%)) were purchased from Aldrich.

The 5.10^-3 M Cd²⁺ reference solution was prepared by weighing out the required amount of CdCl₂H₂O and dissolving it in 100 mL Milli-Q® water.

The 10^-5 M Cd²⁺ solution used in this study was prepared by diluting 5.10^-3 M Cd²⁺ with 10 mL 0,1 M KNO₃ or 0,1 M NaClO₄ in the polarographic cell.

Na-DDTC (Fw: 225.331) was purchased from Sigma Aldrich. Na-DDTC solutions with Ct from 1.10^-4 to 3.10^-4 M were prepared by weighing out their required amounts and dissolving them in the same solution of 0.1 M KNO₃ or 0.1 M NaClO₄, as those used for preparing the Cd²⁺ solution.

H₂SO₄ (Fw: 98.08, purity of 99.99%) used in WL measurements was purchased from Sigma Aldrich.

Equipments

DPV measurements were performed using a Trace-Lab50 from Radiometer, which included a polarographic analyser (Pol150) and stand (MDE150), monitored by Trace Master 5 software. These measurements were carried out using a conventional three-electrode system. The working electrode was a HMDE, with a capillary diameter of 70 µm. Potential’s values were expressed vs. an Ag/AgCl (3 M KCl) electrode separated from the test solution by a salt bridge containing the solvent/SE. The auxiliary electrode was a Pt wire.

Spectrophotometric measurements were performed on a UV-Visible Thermo Electron Corporation Nicolet Evolution 100 device, using a matched 10 mm quartz cell.

The pH measurements were carried out employing a pH-meter from Denver Instrument Company.

Electrochemical method

In DPV experiments, the pulse duration was programmed to 0.04 s, and the scan rate was 125 mV/s. Initial and final potential’s values were -400 and -700 mV, respectively. A 10 mL SE was placed in the polarographic cell. A magnetic stirrer provided the convective transport. Pure N₂ was bubbled through the experimental solution to remove dissolved O₂. Data were analyzed using Lingane’s equations.

Lingane’s equation

For the reaction denoted as M^m+ + qL( (MLq)^m+ (where M is a metal with a charge of m⁺ and L is a neutral ligand with a stoichiometric coefficient of q), the polarographic half wave or E_p shifts more negatively with higher Ct of Na-DDTC added to Cd ion solution. Lingane’s ²² classical equation is derived in eq. (1), to calculate ( when activity coefficients were not considered.

(1)

where n value of was = 2. This equation has been successfully applied by several researchers for determining ( and the ligand’s stoichiometric coefficient complexation reactions ^22,23.

Extended Lingane's equation

If the current’s effect is included, the classical Lingane’s equation should be extended as follows ^24,25 in eq. (2):

(2)

This equation can only be applied when the system involves a complex highly stable.

Spectrophotometric method

For determining stoichiometry and K_f, A of a series of metal salt and ligand mixtures was initially measured at 298 K. Then, in order to perform ( variation with changes in T, A shifts at different T were studied.

Data were analyzed using Job’s method, also called CVM, which was selected, since it is easy, simple and commonly employed ²⁶ to determine anions and organometallic compounds, and to help understand ligands/metal action modes ^26,27. In analytical chemistry, it is mainly used to determine the stability of a complex, its K_f and composition. Optical measurements were employed to study a system with two or three components.

The procedure included making a volume of 2 mL metal complex solutions with different Ct of Cd ions and ligands (Table 1).

Table 1: Metal complex solutions containing different Ct of Cd ions and ligands. 

Then, the sum of the Ct from Cd²⁺ was calculated by eq. (3):

(3)

The optical densities of the solutions, prepared in the previous step, were determined at the wavelength of a light strongly absorbed by Cd-Na-DDTC. The metal-ion and the ligand did not absorb at this wavelength.

A was plotted against f_L (= C_L /Ct). This plot is referred to as a Job’s plot. For calculating the Ct of Cd-Na-DDTC, eq. 3 was rewritten as eq. (4):

(4)

Since, as in eq. (5):

(5)

Eq. (4) can be reduced to:

(6)

From eqs. (5) and (6) q can be obtained in eq. (7):

(7)

Job’s diagram, as shown in Fig. 6, consists of two straight lines intersecting at A₀, at a given value of f_L, which indicates Cd:Na-DDTC ratio in ML_q.

Once q was determined, K_f was calculated ^29,30 by eqs. (8) and (9):

(8)

(9)

A was plotted at λ_max vs. f_L, in order to obtain Job’s diagram. A and f_L were given by eqs. (10) and (11):

(10)

(11)

The deviation from A_o values was due to Cd-Na-DDTC’s complex dissociation. Thus, ( was determined from theoretical lines deviations (Fig. 6). A_max denoted A value at the highest point on the experimental curve, indicating the maximum amount of the complex formed with a dissociation degree (α). A_o represents absorbance value, where the theoretical straight lines intersected due to the complex, when its maximum Ct was formed, with a value of α = 0. A_α represents the absorbance value of the complex’s dissociated part, which is the difference between A₀ and A_max. In order to quantify (, it was imperative to determine α, which was calculated by employing eq. (12):

(12)

At equilibrium, eq. (13) is as follows:

(13)

So, K _f was given by:

(14)

CI study

Corrosive solution

0.5 M H₂SO₄ was used as an aggressive solution, prepared by diluting a 99.99% analytical reagent grade with ultra-pure water.

For WL measurements, 100 mL of the test solution were used, and the Ct of the studied inhibitors varied from 50 to 500 mg/L.

WL measurements

Rectangular samples of CS XC38 (AFNOR/Euronorm: C35E CS and US specification: SAE1035), of 3 × 1.4 × 0.25 cm, were obtained by cutting plates with the chemical composition (wt.%): C = 0.37%, S = 0.016%, Cr = 0.077%, Mn = 0.68%, Si = 0.23%, Ti = 0.011%, Ni = 0.059%, Co = 0.009%, Cu = 0.160% and remainder iron (Fe), which were used for WL measurements. The samples were gradually polished with 200 to 1200 grade emery paper, with a pitch of 200. They were rinsed with double distilled water, degreased, washed thoroughly with double distilled water, dried in a stream of hot air, and weighed. Then they were immersed in the test aggressive solution without and with inhibitor, at different Ct, for 1 h, at T in the range from 298 to 323 K. At the end of the test period, the samples were withdrawn, cleaned with water, dried and reweighed. All measurements were repeated at least thrice, to ensure the results reproducibility.

The measured CR (mg/cm²/h) was determined using eq. (15):

(15)

where m₀ and m₁ are the weight of samples before and after immersion, respectively, t is immersion time (1 h) and A is the samples’ total exposed surface area.

According to several authors ^31-33, IE(%) may be calculated using eq. (16):

(16)

where v₀ and v denote average WL in a H₂SO₄ solution without and with inhibitor, respectively.

Voltammetric behavior

( of Cd-Na-DEDTC complex was calculated by DPV. KNO₃ and NaClO₄ solutions were tested as SE, with an ionic strength of 0.1 M, at 298 K.

Under these conditions, 10^-5 M Cd²⁺ showed a peak at about -500 mV, without the ligand. With Na-DEDTC added, at different Ct, to the Cd ion, its E_p shifted towards more negative values than those of the simple and free states. DPV values of Cd²⁺ in different Ct of Na-DDTC are shown in Figs. 1 and 2.

Figure 1: DPV of- (A) 10^-5 M Cd²⁺ and (B) different Ct of Na-DDTC in 0.1 M NaClO₄. 

Figure 2: DPV (A) without and (B) with Na-DDTC, at different Ct in 0.1 M KNO₃. 

This significant displacement was used to determine ( and q, using simple and extended Lingane’s equations. The obtained results are listed in Tables 2 and 3.

Table 2: ∆E_p vs. Ag/AgCl and peak current of a 10^-5M Cd²⁺ solution as a function of the Ct from Na-DDTC with 0.1 M NaClO₄. 

Table 3: ∆E_p vs. Ag/AgCl and peak current of a 10^-5 M Cd²⁺ solution as a function of the Ct from Na-DDTC with 0.1 M KNO₃. 

( values of Cd-Na-DDTC complex obtained from Eq. 1 were equal to 17.52 for KNO₃ and 20.2 for NaClO₄. The values of q were 4 in KNO₃ and 5 in NaClO₄. The values obtained from Eq. 2 for logβ were 21.20 and 24.48, in KNO₃ and NaClO₄, respectively. Those of q were 5 in KNO₃ and 6 in NaClO₄. Thus, it is noted that there was a significant difference in the results obtained by eqs. (1) and (2), which seems to be due to the change in peak intensity.

Validity of approximation

In this step, the same approximations that were used by ³⁴ were adopted. Lingane’s equation was applied by considering Na-DDTC’s activity instead of its Ct. Ionic strength was considered 0.1 M, and the ligand’s activity coefficient was found to be 0.714. Thus, obtained results are given in Table 4, which shows that log( on KNO₃ and NaClO₄ originated errors of -0.75 and -3.59%, respectively. Concerning q, on KNO₃ and NaClO_4, 0.064% and 0.00%, respectively, were obtained. That is why it was concluded that Lingane’s equation should be used in its extended form, and that the approximation γ_Mxγ_L ^q/γ_ML = 1 was not favorable. Thus, and still in agreement with ³⁴, q and log( values were obtained by plotting n∆E_p/0.059 + log(I_M/I_ML), as an activity’s function (Fig. 3). In this context, even if ∆E_p remains constant, β and q can be calculated by plotting log(I_ML/I_M) as log C_L’s function ³⁵.

Cd²⁺ reversibility in Na-DDTC presence was verified by DPV. As Fig. 4 shows, according to E_pa - E_pc, values were close to 30 mV, and to I_pa/I_pc, were near to unity. This indicated that the system, including the cation with Na-DDTC, was reversible. A similar result was found by ³⁶, who have studied complexation reactions of Zn²⁺, Pb²⁺, Cd²⁺ and Tl⁺ metal cations by 5,7- diiodo-8-hydroxyquinoline in non-aqueous solvents ³⁶.

Table 4: Results of q and logβ determination for Cd-Na-DDTC calculated from DPV at 298 K. 

Figure 3: Linear dependence of ∆E_p on Na-DDTC activity and C_L in - (A) KNO₃ and (B) NaClO₄. 

Figure 4: DPV of 10^-5 M Cd²⁺ with 10^-4 M Na-DDTC in- (A) 0.1 M KNO₃ and (B) 0.1 M NaClO₄; scan rate of 125 mV/s, at 25 °C. 

Spectrophotometric behavior

The obtained results of ( and q in both SE were checked using spectrophotometric technique. Electronic spectra of Na-DDTC and Cd(II) complex in KNO₃ are shown in Fig. 5. When Cd²⁺ ion was added to Na-DDTC in the KNO₃ solution, two absorption peaks were seen. The first, at λ_max = 256 nm, corresponds to n-π transition of S-C=S, and the second, at λ_max = 280 nm, is assigned to the intraligand π-π* transition of N-C-SS. S absorption was determined ^37,38. In perchlorate salts, the first absorption peak appeared in the same place, while a 20 nm bathochromic shift occurred at 300 nm.

Figure 5: Cd-Na-DDTC absorption spectra in a solution with decreasing Na-DDTC, increasing CS and constant total Ct in 0.1 M NaClO₄, at 298±0.1 K. 

According to the results in Tables 5 and 6, Job’s diagrams were drawn by plotting absorbance as f_L function, at 280 nm and 300 nm, in KNO₃ and NaClO₄ solutions, respectively. ( and q were derived from eqs. (7) and (14).

Table 5: Spectrophotometric data for Cd-Na-DDTC system in the 0.1 M KNO₃ solution, with pH of 7.5, at different T. 

Table 6: Spectrophotometric data for Cd-Na-DDTC system in the 0.1 M NaClO₄ solution, with pH = 7.5, at different T. 

Fig. 6 depicts Job’s diagram in NaClO₄, showing that the straight lines intersected at f_L = 0.85, which indicates that Cd²⁺ coordinated with six Na-DDTC ligands. The same was seen for KNO₃. In this case, q was found to be 5. All measurements were taken at a pH of 7.5, since Na-DDTC was dissolved in acidic media. These results can be compared with those obtained by DPV. In addition, in this area, hydroxyl groups formation was suppressed. When the results obtained in both SE were compared, it was found that ( in KNO₃ was lower than that in NaClO₄. From this, it can be concluded that nitrate ions have a complexing character, and compete with Cd²⁺ for Na-DDTC.

Figure 6: Job’s diagram for Cd-Na-DDTC complex in 0.1 M NaClO₄ (λ_max = 300 nm). 

Effect of T

Thermodynamic parameters provide information about the structure, the kind of complex, and the nature of the interaction between Na-DDTC and Cd²⁺. These parameters were calculated using the eqs. (17-19):

(17)

(18)

These equations can be arranged to give eq. (19):

(19)

The plot of lnK or ln( vs. that was not herein reported, gave a straight line. The slope and intercept of this straight line were (∆H°)/R and (∆𝑆°)/𝑅 (complexation entropy), respectively.

(, q and thermodynamic parameters of the complexation process were determined in KNO₃ or NaClO₄ solutions at different T (Table 7). These results show that ΔGº was negative, which confirms the spontaneous nature of the complexation process. Moreover, ( values decreased with increasing T. At 323 K, q was found to be 4. This shows that the complex’s formation was strongly favored at lower T. On the other hand, ΔHº negative values indicate an exothermic interaction between Na-DDTC and Cd²⁺. ΔSº positive values led to an increase in the solution, and accordingly, the complex’s structure disordering. ΔSº value in the NaClO₄ solution was higher than that from the KNO₃ one, showing an increase in steric factor ³⁹. K_f was more favored in the NaClO₄ solution.

Table 7: ( for Cd-Na-DDTC calculated from Job’s method, at various T, in NaClO₄ and KNO₃ solutions, with different thermodynamic parameters. 

Effect of Ct

DPV showed that, even in acidic media, Cd-Na-DDTC complex was formed. For this reason, IE(%) of the complex on CS corrosion in H₂SO₄ was studied. Corrosion parameters for CS in 0.5 M H₂SO₄ were determined, from WL measurements, at different Ct (50-500 mg/L) of the studied inhibitor, at various T (Fig. 7).

Figure 7: Relationship between CR and Na-DDTC Ct, at different T. 

Fig. 7 shows that, at the initial Ct of Na-DDTC (50 mg/L), CR stayed high. Thus, this Ct was not enough to cover the entire CS’s surface, and IE(%) was low. When it was increased, CR decreased for all T, except 323 K, and the inhibitor formed an effective film that protected the CS’s surface. This was accompanied by a decrease in CR and an increase in IE(%), as a function of the inhibitor’s Ct. Maximum IE(%) (>96) was obtained with a Ct of 500 mg/L.

Concerning general dissolution mechanism for CS, the one proposed by ⁴⁰, which found that CS is positively charged in a H₂SO₄ medium, was herein adopted. SO₄ ^2- anions act as connecting bridges between protonated organic inhibitors and positively charged metal surfaces, as expressed in eqs. (20-22).

(20)

(21)

(22)

Furthermore, CS corrosion in H₂SO₄ solutions was an electrochemical process. The anodic reaction was:

(23)

And cathodic behavior corresponded to Volmer’s reaction:

(24)

followed by Tafel’s reaction:

(25)

or Heyrovsky’s reaction:

(26)

Effect of T

To investigate T effect on CR, WL measurements were performed at different T (298-323 K), without and with inhibitor, at various Ct (Table 8).

Table 8: CR and IE(%) for CS in a 0.5 M H₂SO₄ solution with different Ct of Na-DDTC, at various T. 

Inspection of the results reveals that CR decreased with higher inhibitor’s Ct. So, at 500 mg/L, IE(%) reached maximum values of 96.43, 96.98, and 92.2%, at 293, 303 and 308 K, respectively.

The fact that IE(%) decreased with higher T suggests a physical adsorption mechanism. This may be due to an increase in the solubility of the protective barrier. On the other hand, the decrease in IE(%) with higher T may also be due to a possible shift in K_ads towards the desorption of the adsorbed inhibitor molecules. Inversely, IE(%) increased with T from 293 to 303 K. That suggests a chemical adsorption mechanism ⁴¹. So, physisorption was predominant at higher T, while chemical adsorption predominated at lower ones (≤ 303 K).

Synergistic effect of Cd²⁺

To show the effect of Cd²⁺ addition to Na-DDTC on its IE(%), the metal’s Ct was increased from 4.43 to 22.53 mg/L, while the ligand’s Ct was kept constant. Table 9 shows that the decrease in CR of CS in 0.5 M H₂SO₄ with Cd²⁺ was more pronounced than that without it. This table also shows that CR visibly decreased with increased Ct of Cd²⁺. As a result, IE(%) also substantially increased. This result may be explained by Cd ions and Na-DDTC molecules co-adsorption onto the CS surface, which involved a possible increase in Cd²⁺ synergistic effect. Therefore, it was concluded that Cd-Na-DDTC complex was more hydrophobic in the CS surface, and it had low electronegativity and easy polarization in the electronic layer.

Table 9: CR and IE(%) for XC38 CS in a 0.5 M H₂SO₄ solution, at 303 K, at different Ct of Cd²⁺ with 100 mg/L Na-DDTC. 

It is worthwhile to note that, when H₂SO₄ was added to Na-DDTC without Cd ion, the synergistic effect was not seen. Inversely, when Cd ion was added to Na-DDTC without H₂SO₄, IE(%) increased. This could mean that although H₂SO₄ can degrade Na-DDTC, it cannot dissolve the complex.

Adsorption study

Most often, organic CI act on the metal surface by adsorption. It is known that the determination of adsorption types and thermodynamic parameters serves to understand the interaction between metal surface and inhibitor.

Assuming that CI mechanism was due to Na-DDTC adsorption, according to ⁴², the degree of metal surface coverage SC (θ) was calculated from WL measurements, using the following equation:

(27)

where Vcorr and V´corr are CR for solutions without and with inhibitor, respectively. Therefore, SC was calculated using eq. (28):

(28)

In order to verify the suitable adsorption mode, various isotherms, including Temkin’s, Frumkin’s, Flory Huggins’ and Langmuir’s, were tested. The best fit was obtained with Langmuir’s isotherm. According to it, the inhibitor Ct was related with SC by eq. (29):

(29)

K_ads denotes the interaction strength between adsorbent and adsorbate ⁴³. The plot of C_𝑖𝑛ℎ/𝜃 vs. C_inh, at different T, yielded straight lines (Fig. 8 (A)).

Figure 8: Langmuir’s adsorption isotherm on CS in 0.5 M H₂SO₄ for- (A) Na-DDTC and (B) Cd-Na-DDTC. 

Table 10 shows R² values and the slope of all straight lines, which were close to 1. This confirms that the inhibitor adsorption onto the CS surface obeyed Langmuir’s isotherm. However, a deviation of the slopes from unity can be seen, which was due to the interaction between adsorbates on the metal surface ^44-46, and to the fact that adsorption heat changed with the increasing SC ⁴⁴. So, Langmuir’s isotherm could not be rigorously applied. Thus, Na-DDTC’s behavior was interpreted by a modified Langmuir’s isotherm, which suggested that each molecule inhibitor could occupy n’s adsorbed sites (31). This modified expression is given in eq. (30):

(30)

K_ads value obtained from the intercept enabled to calculate ΔGº_ads, as in eq. (31):

(31)

where 55.5 is the molar Ct of water in the solution in mol/L^-1.

From the intercept in Fig. 8 (A) and from Eq. (31), K_ads and ΔGº_ads (Table 10) for Na-DDTC were determined at different T. K_ads larger value obtained at 308 K suggests that corrosion IE(%) of Na-DDTC improved at this T, leading to the formation of a protective film at the metal/solution interface, which reduced chemical attacks on CS. ΔGº_ads negative values at all tested T revealed that the adsorption process was spontaneous.

According to ⁴⁵, ΔGº_ads values up to -20 kJ/mol cause physisorption, while values lower than -40 kJ/mol lead to chemisorption. In this study, ΔGº_ads values were from -20 to -40 kJ/mol. This suggests that Na-DDTC went through both physical and chemical adsorption ⁴⁷.

The plot of 𝐶_𝑖𝑛𝑏/𝜃 vs. C_inh, in Cd-Na-DDTC complex’s case, was a straight line (Fig. 8 (B)) with the slope equal to unity and R² equal to 0,99. This confirmed that the inhibitor adsorption followed Langmuir’s isotherm. The intercept gave K_ads, as represented in Table 10.

Table 10: Thermodynamic parameters obtained from gravimetric study for Na-DDTC and Cd-Na-DDTC adsorption onto the CS surface in 0.5 M H₂SO₄. 

The obtained results indicate that K_ads values were higher than those obtained by Na-DDTC alone, which might be due to the presence of a strong bond between adsorbate and adsorbent. More importantly, ΔGº_ads value for Cd-Na-DDTC complex around -40 kJ/mol led to chemisorption process. This was due to the transfer of an unshared electronic pair of organic molecules to the metal surface, thus forming a dative bond.

Synergistic effect mechanism

S_p for different Ct of Cd²⁺ in 0.5 M H₂SO₄ with 100 mg/L Na-DDTC was determined from the relationship proposed by ⁴⁸ (eqs. (32) and (33)).

(32)

with

(33)

where P₁, P₂ and P'₁₊₂ are IE(%) values of Na-DDTC, Cd²⁺ and of Cd-Na-DDTC, respectively. S_p values are given in Table 11.

Table 11: S_p for different Ct of Cd²⁺ + 100 mg/L^-1 Na-DDTC. 

It was noted that all S_p values were greater than unity. This result shows that Cd-Na-DDTC high CI on CS was due to the synergy effect. As a matter of fact, ⁴⁹ explained the synergy effect either by competitive or by cooperative adsorption between two compounds. In the first case, the two compounds are adsorbed on different sites onto the electrode surface. In the second case, one compound is chemisorbed onto the metal surface and the other is physisorbed onto it. If S_p < 1, thre is competitive adsorption. If S_p > 1, cooperative adsorption occurs. The values obtained in this study were well above unity, which shows cooperative adsorption between Na-DDTC and Cd²⁺. Thus, the synergistic effect can be explained as follows: Na-DDTC was protonated in the acidic solution. Then, the protonated Na-DDTC and Cd²⁺ could attach onto XC38 (CS) through electrostatic interaction with the negatively charged surface, which was provided with the specifically adsorbed SO₄ ^2- anions onto the FeSO₄ ^2- surface. When Na-DDTC adsorbed onto the CS surface, coordinate bonds were formed by the partial transference of electrons from unprotonated N and S atoms and delocalized π electrons in S atoms groups to the metal surface and to the Cd²⁺ ions, which were able to form stable complexes with the ligand. Therefore, in the adsorption process, both physical and chemical adsorption took place.