Synthesis, characterization, corrosion inhibition of mild steel in HCl (0.5?N) solution and solid-state electrical conductivity of new Co(II), Ni(II), Cu(II) and Zn(II) complexes

Article abstract of DOI:10.1002/aoc.3572

This work consists of a study of the corrosion-inhibiting and semiconducting properties of new binuclear Co(II), Ni(II), Cu(II) and Zn(II) complexes with a Schiff base, H₂L, obtained from the 2:1 M condensation of salicylaldehyde and o-dianisidine, respectively. Elemental, spectral and thermal analyses were used to characterize these complexes. The magnetic susceptibilities of these complexes were also determined. Weight loss, potentiodynamic polarization and scanning optical microscopy were the techniques used to investigate the efficiency of these new compounds as corrosion inhibitors. The antibacterial activity of the compounds was measured against sulfate-reducing bacteria. It was found that inhibition occurs via the chemisorption of metal complexes on the steel surface. This absorption obeys the Langmuir adsorption isotherm model.

Full text of DOI:10.1002/aoc.3572

Received: 15 April 2016
DOI 10.1002/aoc.3572
Revised: 23 June 2016
Accepted: 1 July 2016
F U L L PA P E R
Synthesis, characterization, corrosion inhibition of mild steel in
HCl (0.5 N) solution and solid‐state electrical conductivity of new
Co(II), Ni(II), Cu(II) and Zn(II) complexes
*
A. M. Nassar | A. M. Hassan | A. N. Elkmash
Chemistry Department, Faculty of Science, Al‐
Azhar University, Nasr City, Cairo, Egypt
This work consists of a study of the corrosion‐inhibiting and semiconducting prop-
erties of new binuclear Co(II), Ni(II), Cu(II) and Zn(II) complexes with a Schiff
base, H2L, obtained from the 2:1 M condensation of salicylaldehyde and o‐
dianisidine, respectively. Elemental, spectral and thermal analyses were used to
characterize these complexes. The magnetic susceptibilities of these complexes were
also determined. Weight loss, potentiodynamic polarization and scanning optical
microscopy were the techniques used to investigate the efficiency of these new com-
pounds as corrosion inhibitors. The antibacterial activity of the compounds was
measured against sulfate‐reducing bacteria. It was found that inhibition occurs via
the chemisorption of metal complexes on the steel surface. This absorption obeys
the Langmuir adsorption isotherm model.
Correspondence
Amr M. Nassar, Chemistry Department, Faculty of
Science, Al‐Azhar University, Nasr City, Cairo,
Egypt.
E‐mail: nassar_tanta@yahoo.com
KEYWORDS
anticorrosion, bimetallic, Complexes, DC electrical conductivity, SRB
1
|
INTRODUCTION
complexes with mild steel is greatly affected by their standard
electrode potentials, their reactivity and the nature of the
Mild steel is one of the major industrial materials. It is, for
instance, used as a main raw material in fabrication of
equipment, a chief material of construction and in fabrica-
tion of weapons due to its good mechanical properties
and low cost.[1] Acid solutions are widely used for acid
cleaning and acid descaling etching of mild steel and heat
exchangers to remove deposits and other corrosion prod-
ucts.[2] Unfortunately, acids are very corrosive to most
common metals and alloys particularly mild steel. The
use of synthetic inhibitors is one of the practical methods
for the protection of mild steel against corrosion in acidic
media.[3] Schiff base complexes are part of these synthe-
sized inhibitors. In fact, these complexes have several appli-
cations in catalysis,[4] antimicrobial[5] and anticancer[6]
photochemistry,[7] nanomaterials synthesis,[8] polymer
synthesis[9] and semiconductors.[10]
ligand used to stabilize the metallic complexes. The presence
of aromatic rings, electronegative nitrogen and oxygen atoms
and azomethine group is likely to enhance the adsorption of
the inhibitors on a steel surface promoting effective corrosion
inhibition.[17]
Many microorganisms have effects on metals, known as
biochemical corrosion, ascribed to microbiologically
influenced corrosion (MIC).[18] Sulfate‐reducing bacteria
(SRB) are reported to be the main reason for MIC by acceler-
ating the corrosion rate and inducing stress corrosion and
pitting corrosion.[19] In previous works,[14,15] we have studied
the effect of Schiff base metal complexes as SRB growth
inhibitors in the protection of metals from microbial corrosion.
In view of the obtained results showing the significant
role of anticorrosion and antimicrobial corrosion properties
of Schiff base complexes, we had the idea to prepare Co
(II), Ni(II), Cu(II) and Zn(II) complexes with the Schiff base
(H2L), and to study their structures, DC electrical conductiv-
ity and anticorrosion and anti‐SRB behaviour.
It appears from a literature survey that transition metal
complexes have anticorrosion properties.[11–15] According
to Mahdavian and Attar,[16] the interaction of transition metal
Appl. Organometal. Chem. 2016; 1–9
wileyonlinelibrary.com/journal/aoc
Copyright © 2016 John Wiley & Sons, Ltd.
1

2
A. M. NASSAR ET AL.
2
| EXPERIMENTAL
2.1 | Materials and methods
All the chemicals and reagents were of analytical grade and
were used as such without any further purification. Dis-
tilled water was used throughout all experiments. The
organic solvents were HPLC grade or extra‐pure grades
and were used without further purification. The electrolyte
solution (HCl; 0.5 N) was prepared from analytical‐grade
HCl and distilled water. All corrosion tests were performed
at room temperature according to the methods mentioned in
previous work.[11–13] The biological activity was assayed
using the most probable number technique.[18] The electri-
cal conductivity (σ) of samples was measured in air in
the temperature range 30–100°C. Good contact of elec-
trodes with the surfaces of the disc was obtained by paint-
ing the surfaces carefully with a silver paste. The
temperature was measured via a digital thermometer placed
close to the sample. The currents (I) at different applied
voltage (V) were measured at room temperature to establish
the I–V characteristics of the samples.
FIGURE 1 Structure of Schiff base ligand H2L
to 25 ml of an ethanolic solution of o‐dianisidine (1.22 g,
0.005 mol) with stirring. The mixture was then heated under
reflux for 2 h. An orange‐coloured product was precipitated
on concentrating the solution to half of its initial volume
and on cooling to room temperature. The product was fil-
tered, washed with diethyl ether and dried in a desiccator over
anhydrous calcium chloride. The pure compound was
recrystallized from hot ethyl acetate. The purity of the ligand
was checked using the TLC technique.
Yield 87.3%; m.p. 180°C, molecular weight 452.3 g mol
−1. Anal. Calcd for C28H24N2O4 (%): C, 74.3; H, 5.3; N, 6.3.
Found (%): C, 74.5; H, 5.3; N, 6.2. Main IR peaks (KBr, cm
1
−1): ν(OH) 3445, ν(C═N) 1616. H NMR (DMSO‐d6, δ,
ppm): 13.9 (s, OH, 2H), 8.9 (s,─CH═N, 2H), 6.7–7.6 (m,
Ar─H, 14H), 3.9 (s,─OCH3, 6H) (Figures S1 and S2). MS,
m/z = 453 (Figure S3).
2.2 | Instrumentation
2.4 | Synthesis of metal complexes
Elemental analysis for C, H and N was conducted using a
CHNS‐932 (LECO) Vario elemental analyser. Magnetic sus-
ceptibility measurements were carried out at room tempera-
Cobalt(II), nickel(II), copper(II) and zinc(II) complexes of
H2L were synthesized by mixing 30 ml of a methanolic solu-
tion containing 0.010 mol of each metal(II) chloride with H2L
(2.26 g, 0.005 mol) solution in hot methanol (25 ml) in metal‐
to‐ligand molar ratio of 2:1. The Ni(II) and Zn(II) complexes
were formed as insoluble precipitates after refluxing the reac-
tion mixture for 2–3 h, whereas the Co(II) and Cu(II) com-
plexes were precipitated immediately on stirring at room
temperature. The metal complexes were filtered, washed sev-
eral times with diethyl ether and finally dried in a desiccator
at room temperature. Below are mentioned the most impor-
tant characteristics of the obtained complexes.
ture with
a Sherwood Scientific magnetic balance.
Electronic spectra were measured in the UV–visible range
(195–1100 nm) using a PerkinElmer Lambda 35 UV–visible
spectrometer. Infrared (IR) spectra were recorded using KBr
pellets with a PerkinElmer 1430 spectrometer in the region
200–4000 cm−1 1H NMR and 13C NMR spectra were
.
recorded with a GEMINI‐300BB NMR 500 MHz spectrom-
eter. Electron spin resonance (ESR) spectra were recorded
with a Bruker EMX spectrometer working in expand
9.7 GHz with 100 kHz frequency and microwave power of
1 MW and modulation amplitude of 4 G. Mass spectra were
2.4.1
| Complex Co2L(H2O)2(Cl)2·3H2O (1)
recorded with
a GC‐MSA‐QP 5050 A (Shimadzu).
Brown solid; m.p. > 300°C; molecular weight 728.98 g mol
−1. Anal. Calcd for C28H32Co2N2O9Cl2 (%): C, 46.9; H, 4.4;
N, 3.8; Co, 16.2. Found (%): C, 47.0; H, 4.4; N, 3.9; Co,
15.9. Main IR peaks (KBr, cm−1): 1598 (C═N), 499
(Co─N), 1261 (C─O), 3402 (ν H2O), 580 (Co─O), 304
(Co─Cl). Λm (Ω−1 mol−1 cm2): 5.05. UV–visible (DMF,
nm): 608, 668 (d–d transition). μeff (BM): 4.9.
Potentiodynamic polarization measurements were carried
out using an Autolab‐PGSTAT302N‐HP, with a glassy car-
bon working electrode, platinum wire auxiliary electrode
and Ag/AgCl reference electrode. Scanning optical micros-
copy was conducted using a Zeiss Axiovert 40 mat micro-
scope at Central Metallurgical Research and Development
Institute (CMRDI), Helwan, Egypt. The electrical conductiv-
ity (σ) of samples was measured in air in the temperature
range 30–100°C using a Super Magohmmeter electrometer.
2.4.2
| Complex Ni2L(H2O)2(Cl)2·2H2O (2)
Brownish yellow solid; m.p. > 300°C; molecular weight
710.38 g mol−1. Anal. Calcd for C28H30Ni2N2O8Cl2 (%): C,
47.3; H, 4.2; N, 3.9; Ni, 16.5. Found (%): C, 47.4; H, 4.2;
N, 3.9; Ni, 16.0. Main IR peaks (KBr, cm−1): 1600 (C═N),
498 (Ni─N), 1267 (C─O), 3384 (ν H2O), 579 (Ni─O), 320
(Ni─Cl). Λm (Ω−1 mol−1 cm2): 8.02. UV–visible (DMF,
nm): 465 (d–d transition). μeff (BM): 3.5.
2.3 | Preparation of Schiff base ligand
The Schiff base ligand (H2L), the structure of which
presented in Figure 1, was prepared and characterized using
a published method.[20,21] An ethanolic solution (25 ml) of
salicylaldehyde (1.065 ml, 0.010 mol) was added dropwise

A. M. NASSAR ET AL.
3
2.4.3
|
Complex Cu2L(H2O)2(Cl)2·2H2O (3)
3.1 | Characterization of complexes
Dark brown solid; m.p. 280°C; molecular weight
720.00 g mol−1. Anal. Calcd for C28H30Cu2N2O8Cl2 (%):
C, 46.6; H, 4.2; N, 3.9; Cu, 17.6. Found (%): C, 46.7; H,
4.1; N, 83.9; Cu, 17.1. Main IR peaks (KBr, cm−1): 1602
(C═N), 499 (Cu─N), 1263 (C─O), 3454‐(ν H2O), 579
(Cu─O), 315 (Cu─Cl). Λm (Ω−1 mol−1 cm2): 9.50. UV–visi-
ble (DMF, nm): 568 (d–d transition). μeff (BM): 2.8.
3.1.1
| IR spectra
On comparing the IR spectra of the complexes with that of
the free ligand (Table 1), we note a shift to lower frequencies
between 1597 and 1602 cm−1 of the νC═N band observed in
the spectrum of the ligand at 1616 cm−1. This shift could be
explained by coordination via the azomethine nitrogen with
the metal ions.[23] This explanation could be confirmed by
the appearance of new bands assigned to νM─N observed
in the far‐IR region of the spectra of the complexes in the
range 478–499 cm−1. It also appears that the phenolic‐OH
of the ligand is deprotonated during complexation with
metal(II). This could be confirmed by the shift of the pheno-
lic νC─O to lower wavenumber in the range 1261–1267 cm
2.4.4
| Complex Zn2L(H2O)2(Cl)2·3H2O (4)
Deep orange solid; m.p. 270°C; molecular weight
741.76 g mol−1. Anal. Calcd for C28H32Zn2N2O9Cl2 (%):
C, 45.3; H, 4.3; N, 3.8; Zn, 17.6. Found (%): C, 45.4; H,
4.3; N, 3.9; Zn, 17.2. Main IR peaks (KBr, cm−1): 1597
(C═N), 478 (Zn─N), 1267 (C─O), 3390‐(ν H2O), 571
(Zn─O), 305 (Zn─Cl). Λm (Ω−1 mol−1 cm2): 7.32. 1H
NMR (DMSO‐d6, δ, ppm): 4.1 (s, lattice H2O, 6H), 4.2 (s,
coordinated H2O, 4H), 9.04(s, ─CH═N, 2H), 6.8–8.2 (m,
Ar─H, 14H), 3.9 (s,─OCH3, 6H). 13C NMR (DMSO‐d6, δ,
ppm): 162 (C─O), 160.8 (─C═N), 56.5 (─OCH3), 115–130
(Ar─C). UV–visible (DMF, nm): 425 (charge transfer),
diamagnetic.
−1
which appears in the spectrum of the free ligand at
1285 cm−1. We can say that the participation of the ligand
with the metal ions occurs through a phenolate oxygen.[24]
On the other hand, a very broad band is observed in the spec-
tra of the complexes in the range 3384–3453 cm−1 which
supports the presence of lattice and coordinated water mole-
cules in the complexes.[25] Weak bands appear in the far‐IR
spectra between 571 and 580 cm−1 and can be attributed to
νM─O. Furthermore, other weak bands in the range 304–
320 cm−1 are observed in the spectra of the complexes and
are ascribable to νM─Cl.
3
| RESULTS AND DISCUSSION
The metal complexes which were synthesized in this work are
homobimetallic and are obtained according to the following
reaction:
3.1.2
measurements
| Electronic spectra and magnetic susceptibility
UV–visible spectra of the complexes were recorded in DMF
solution. Complex 1 shows a magnetic moment μeff of 4.9
BM. The low‐intensity shoulders at 608 and 668 nm are
probably due to 4A2(F) → 4 T1(P) and 4A2(F) → 4 T1(F) indi-
cating a tetrahedral geometry of the ligands around Co (II)
ions.[26] The magnetic moment of complex 2 is 3.5 BM.
The electronic spectrum of this complex shows a very broad
2MCl2⋅nH2O þ H2L→½M2LCl2ꢀ⋅nH2O þ 2HCl
where M = Co(II), Ni(II), Cu(II) or Zn(II).
The metal complexes are insoluble in water and in com-
mon organic solvents such as ethanol, methanol, acetone,
chloroform, benzene, diethyl ether and methylene chloride,
but are soluble in dimethylformamide (DMF) and
dimethylsulfoxide (DMSO). The structures of the complexes
3
3
band at 465 nm corresponding to T1(F) → T1(P) indicat-
ing the tetrahedral configuration of the Ni(II) complex.[27]
The spectrum of complex 3 shows a broad band at 568 nm.
1
2
2
were established using elemental analysis and IR, H NMR,
Hence, the copper complex appears to be B1g → A1g and
the μeff value (2.8 BM) corresponds to a tetrahedral geometry
arrangement of the ligands around the Cu(II)ions.[28] The
electronic absorption spectrum of complex 4 shows a charge
transfer band at 425 nm. The diamagnetic behaviour of this
complex is due to the d10 configuration of Zn(II) ions.[29]
13C NMR, ESR and UV–visible spectra, thermogravimetric
analysis (TGA) as well as their magnetic susceptibilities.
The conductance values of the complexes in DMSO solution
at room temperature are in the range 5.05–9.50 Ω−1 mol
−1 cm2 which indicates that these complexes are neutral.[22]
TABLE 1 IR spectral data of ligand (H2L) and its complexes
Compound
ν(H2O)
—
ν(OH)
3445
—
ν(CH)ar
3061
ν(CH)aliph
2961
ν(C═N)
1616
ν(C─O) phenolic, methoxy
1285, 1244
ν(M─O)
—
ν(M─N)
—
ν(M─Cl)
—
H2L
1
3408
3384
3419
3453
3045
2960
1598
1261, 1243
580
499
304
2
—
3034
2930
1600
1267, 1244
579
498
320
3
—
2945
2847
1602
1263, 1243
579
499
315
4
—
3048
2942
1597
1267, 1239
571
478
305

4
A. M. NASSAR ET AL.
3.1.3
|
1H NMR and 13C NMR spectra
occurs through azomethine nitrogen and phenolic oxygen.
1
In H NMR and 13C NMR spectra of the zinc complex,
The 1H NMR spectrum of the Zn(II) complex 4 was recorded
in DMSO‐d6 at room temperature. The broad resonance sig-
nal relating to Ar─OH appears in the spectrum of H2L at
13.9 ppm (s, 2H) and completely disappears in that of its
Zn(II) complex. This indicates the deprotonation of the phe-
nolic‐OH group during complexation confirming the subse-
quent involvement of deprotonated hydroxyls in chelating
with metal ions. Two new signals are observed in the spec-
trum of the zinc complex at 4.1 and 4.2 ppm, which are not
found in the spectrum of the free ligand, and can be attributed
to the resonance of protons of lattice and coordinated water
molecules, respectively.[25] The signal at 8.9 ppm (s, 2H) in
1H NMR spectrum of H2L and assigned to the azomethine
protons is slightly shifted downfield to 9.04 ppm (s, 2H) in
the spectrum of the zinc complex, indicating coordination
through the azomethine nitrogen to Zn(II) ions. The reso-
nance of phenyl protons in the Zn(II) complex show down-
field shift at 6.8–8.2 ppm (m, 14H) as compared to the free
ligand at 6.7–7.6 ppm (m, 14H), due to coordination in the
complex. The spectrum of the zinc complex shows a wide
chemical shift range of aromatic resonance (1.4 ppm) which
may be ascribable to the electron deshielding effect of zinc
metal atom. The 13C NMR spectrum of the free ligand shows
signals at 162.6 ppm (C─OH,ipso) and 161.3 ppm
(C═N).[21] The 13C NMR spectrum of the zinc complex in
DMSO‐d6 shows chemical shifts to higher values corre-
sponding to 162.3 ppm for C─OH and to 160.8 ppm for
C═N in comparison with the signals of the free ligand.
These shifts suggest that the coordination with metal ions
the resonances of methoxy group protons at 3.9 ppm (s,
6H) and carbons at 56.5 ppm remain unchanged compared
to those of the free ligand, indicating that this group does
not participated in coordination with Zn(II) ions.
3.1.4
| ESR spectra
The ESR spectrum of the copper complex (Figure S4) in the
solid state gives axial signal shape and has symmetric bands
with two g values: g|| = 2.2, g⊥ = 2.03. These values suggest
a tetrahedral stereochemistry for the complex. The average g
value equal to 2.08 was equation:
À
Á
1
3
gav
¼
g∥ þ 2g⊥
The value of g|| < 2.3 suggests a covalent character of the
copper–ligand bonding in this complex.[27] From the
observed trend g|| > g⊥ > 2.0023, it is clear that the unpaired
electron lies predominantly in dx2−y2 orbital, giving 2B1g as
the ground state.[28] The exchange interaction (axial) parame-
ter (G), calculated as G= (g∥ − 2.0023)/(g⊥ − 2.0023), and
found for the copper complex to be equal to 7.13, suggests
that the local tetragonal axes are aligned parallel and the
exchange interactions between copper centres in the solid
state are negligible.[29] The ESR spectral parameters of the
Cu(II) complex suggest that the complex has tetrahedral
geometry. These data are in good accordance with other
reported values.[27–29]
TABLE 2 TGA data for complexes
Compound
1
Stage
First
Temp. range (°C)
40–90
Calculated (%)
12.10
5.50
Found (%)
11.70
6.50
Assignment
3 Crystalline H2O + N2
Second
Third
90–163
2 Coordinated H2O
163–285
285–478
478–628
40–629
10.50
26.00
27.00
21.50
9.50
10.10
23.50
27.60
20.50
9.80
Cl2
Fourth
Fifth
Decomposition of organic part of ligand (2C7H4)
Decomposition of organic part of ligand (2C7H7)
1.3 CoO
Residue
First
2
39–130
2 Crystalline H2O + N2
Second
Third
130–179
179–454
454–650
39–650
5.40
6.60
2 Coordinated H2O
31.10
37.50
18.80
5.00
30.40
34.30
18.70
6.30
Decomposition of organic part of ligand (2C7H5) + O2
Decomposition of organic part of ligand(C7H7) + Cl2
2 NiO
Fourth
Residue
First
3
4
38–133
2 Crystalline H2O
Second
Third
133–379
379–612
39–613
33.60
43.80
14.90
7.30
34.10
45.20
14.50
7.70
2 Coordinated H2O + Decomposition of organic part of ligand (2C7H5) + N2
Decomposition of organic part of ligand (2C7H6) + Cl2 + 2O2
1.35 CuO
Residue
First
40–150
3 Crystalline H2O
Second
Third
150–295
295–546
546–649
40–650
8.10
8.90
2 Coordinated H2O + N2
36.80
28.50
18.60
36.60
28.40
18.30
Decomposition of organic part of ligand(2C7H5) + Cl2 + 1.5O
Decomposition of organic part of ligand(2C7H6) + O2
2 ZnO
Fourth
Residue

A. M. NASSAR ET AL.
5
3.1.5
|
TGA
energy of the conduction process, k is the Boltzmann con-
stant and T is the absolute temperature.
The TGA curves for the complexes (Figures S5‐S8) were
obtained at a heating rate of 10°C min−1 and under
flowing nitrogen atmosphere over a temperature range
from ambient to 800°C. The proposed decomposition data
are recorded in Table 2. The decomposition temperatures
and the weight losses of the complexes were calculated
from TGA curves.
Plots of log σ against 1/T are found to have a linear rela-
tionship with temperature in the range 30–100°C, indicating
the semiconducting behaviour of the complexes.[32] The acti-
vation energy values are seen to decrease while the conduc-
tivity increases. This may be due to the overlap between the
electrons in the antibonding π‐orbitals of the ligand and the
empty d‐orbitals of the metals. This leads to a delocalization
of the π electronic charges of the molecules.[33]
Based on the results discussed above, the proposed struc-
ture for the metal complexes is shown in Figure 2.
3.3 | Corrosion inhibition efficiency
3.2 | Solid‐State electrical conductivity
3.3.1
| Weight loss technique
The solid‐state DC electrical conductivity of the synthe-
sized complexes in compressed pellet form (5 ton cm−2
)
The corrosion rate (CR) in millimetres per year was calcu-
lated using the following equation:[34]
with diameter of 10 mm and thickness of 1 to 2 mm
coated with silver paste was measured in the range 30–
100°C. The pellets were dried at 100°C for 24 h in air
before being cooled slowly to room temperature. The con-
ductivity is expected to increase in the presence of halo-
gen atoms due to their ability to accept electrons.[30] It
is observed from Figure 3 that, as the temperature
increases, there is a gradual decrease in DC resistivity.
For the studied complexes, the conductivity increases with
temperature. A linear dependence of log σ = f(103/T) is
observed from the Arrhenius plots of electrical conductiv-
ity, indicating the semiconducting behaviour of these com-
pounds[31] according to the relation:
W×534
CR ¼
(1)
A×d×t
where W is the weight loss (g), d the density of mild steel
(7.85 g cm−3), A the specimen area (cm2) and t the exposure
time (h). The percentage corrosion inhibition efficiency
(IE% = μw%) was calculated using the following equation:[34]
C0R−Ci
μw% ¼
R ×100
(2)
CR0
where C0R and CiR are the corrosion rate in the absence and
presence of inhibitor, respectively. From the evaluated weight
ꢀ
ꢁ
−Ea
kT
TABLE 3 Corrosion rate, surface coverage and inhibition efficiency for the
corrosion of mild steel in aqueous solution of 0.5 M HCl in the absence and
presence of various concentrations of inhibitors from weight loss
measurements
σ ¼ σ0 exp
where σ0 is the pre‐exponential constant value correspond-
ing to 1/T = 0, Ea represents the ordering activation
Compound
Blank
Concentration (M)
0.5 M HCl
1 × 10−4
3 × 10−4
5 × 10−4
7 × 10−4
1 × 10−4
3 × 10−4
5 × 10−4
7 × 10−4
1 × 10−4
3 × 10−4
5 × 10−4
7 × 10−4
1 × 10−4
3 × 10−4
5 × 10−4
7 × 10−4
1 × 10−4
3 × 10−4
5 × 10−4
7 × 10−4
CR (mm yr−1
0.923
)
θ
IE%
—
—
H2L
0.0012
0.0006
0.0004
0.0002
0.0009
0.0004
0.0003
0.0001
0.0006
0.0006
0.0004
0.0002
0.0014
0.0011
0.0006
0.0002
0.0014
0.0008
0.0004
0.0002
0.511
0.732
0.807
0.909
0.605
0.817
0.841
0.940
0.721
0.756
0.835
0.883
0.414
0.543
0.725
0.877
0.562
0.748
0.874
0.936
51.1
73.2
80.7
90.9
60.5
81.7
84.1
94.0
72.1
75.6
83.5
88.3
41.4
54.3
72.5
87.7
56.2
74.8
87.4
93.6
1
2
3
4
FIGURE 2 Proposed structure of metal complexes
FIGURE 3 Arrhenius plots of conductivity for the four complexes

6
A. M. NASSAR ET AL.
loss, surface coverage (θ) was calculated using the following
maximum inhibition efficiency corresponds to 7 × 10
−4 mol dm−3. Complex 1 exhibits the greatest impact on cor-
rosion inhibition among the compounds studied. The order-
ing of the efficiency of the compounds is as follows:
1 > 2 > 4 > H2L > 3.
equation:[34]
C0R−CiR
θ ¼
(3)
CR0
The values of μw%, CR and θ obtained from the weight
loss method at different concentrations of all compounds at
293 K are summarized in Table 3. From the data, the Schiff
base and its complexes show appreciable corrosion inhibition
for mild steel in 0.5 N HCl medium. The effect of the inhib-
itor concentration on the inhibition efficiency shows that the
3.3.2
| Potentiodynamic polarization
The linear Tafel segments[35] of anodic and cathodic curves
were extrapolated to the corrosion potential to obtain corro-
sion current densities (icorr). From the obtained polarization
curves, icorr was calculated by curve fitting using the But-
ler–Volmer equation:[36]
ꢃ
ꢀ
ꢀ
ꢂ
ΔE
ba
ΔE
bc
I ¼ icorr exp 2:3
Þ− exp −2:3
Þ
(4)
where icorr is the corrosion current density, ΔE = E
(Applied potential)‐Ecorr(corrosion potential) and ba and
bc are tafel constants can be measured from experimental
data. The inhibition efficiency was evaluated from the
measured icorr values using the relationship[36]
i0corr−ii
μp% ¼
corr ×100
(5)
i0corr
where i0
and iicorr are the corrosion current densities in
corr
FIGURE 4 Polarization curves for mild steel in presence of various concen-
trations of complex 2 as corrosion inhibitor
the absence and presence of inhibitor, respectively. The
TABLE 4 Polarization parameters of Schiff base and its complexes
Compound
Concentration (M)
Blank
ba
BC
Ecorr
Icorr
119.3
31.1
20.2
14.8
11.3
119.3
27.4
26.7
18.2
10.9
119.3
52.1
36.4
22.2
12.5
119.3
39.6
28.1
16.0
14.5
119.3
36.5
21.3
16.8
11.6
CR
IE%
—
H2L
47.7
15.6
18.0
78.4
26.8
47.7
21.3
42.8
32.0
23.75
47.7
38.6
38.1
24.2
22.4
47.7
33.4
49.35
30.9
25.8
47.7
23.1
37.1
18.9
22.3
71.3
17.2
17.7
74.0
22.1
71.3
18.6
65.4
32.9
21.45
71.3
41.7
45.5
34.7
20.0
71.3
30.3
54.25
31.1
23.0
71.3
28.0
35.7
20.1
19.6
−478.3
−464.9
−466.9
−467.0
−455.7
−478.3
−451.0
−464.0
−464.3
−463.9
−478.3
−475.5
−467.0
−462.9
−463.6
−478.3
−472.4
−460.9
−452.5
−456.1
−478.3
−472.5
−464.1
−466.8
−465.5
0.923
0.240
0.156
0.114
0.087
0.923
0.212
0.206
0.140
0.084
0.923
0.403
0.281
0.172
0.096
0.923
0.306
0.217
0.123
0.112
0.923
0.282
0.165
0.130
0.090
1 × 10−4
3 × 10−4
5 × 10−4
7 × 10−4
Blank
1 × 10−4
3 × 10−4
5 × 10−4
7 × 10−4
Blank
1 × 10−4
3 × 10−4
5 × 10−4
7 × 10−4
Blank
1 × 10−4
3 × 10−4
5 × 10−4
7 × 10−4
Blank
73.9
83.0
87.6
90.5
—
1
2
3
4
77.0
77.6
84.7
90.8
—
56.3
69.5
81.4
89.5
—
66.8
76.4
86.6
87.8
—
1 × 10−4
3 × 10−4
5 × 10−4
7 × 10−4
69.4
82.1
85.9
90.3

A. M. NASSAR ET AL.
7
electrochemical polarization parameters of steel corrosion
in 0.5 N HCl were determined in the absence and pres-
ence of the inhibitors at various concentrations. Figure 4
shows the polarization curves in the presence of complex
2 as inhibitor. The corresponding corrosion efficiencies
determined using the Nova 1.10 program are given in
Table 4. From the obtained results, we note an increase
in the inhibition efficiency for all the compounds used
as the inhibitor concentration increases. According to the
polarization measurements, the ordering of inhibition effi-
ciency is as follows: 1 > 2 > 4 > H2L > 3. These results
are in good agreement with those obtained using the
weight loss method. We also remark that the addition of
the complexes affects both the anodic and the cathodic
reactions. Therefore, these compounds can be classified
as mixed types (anodic/cathodic) inhibitors (Figure S9).
3.3.4 | Scanning optical microscopy
The obtained optical images of mild steel samples show a
degradation of the mild steel in the absence of inhibitors.
The degradation is more accentuated at grain boundaries,
since these regions are most susceptible to corrosion, and
may be responsible for the high rate of corrosion. Optical
microscopy images were obtained of mild steel after corro-
sion in acid medium without and with inhibitors. Figure 6
shows adsorbed layer of complex 1 on the metal surface.
The inhibitor may form a protective film adsorbed on the
steel surface, resulting in a decrease in the contact between
steel and the aggressive solution.
3.3.5
| Mechanism of corrosion inhibition
Apparently, most of the metal complexes have higher corro-
sion inhibition efficiency than the free organic ligand. This
inhibition may be due to the adsorption mechanism.[39] The
adsorption process can occur either by physisorption via
van der Waals forces or by chemisorption via chemical
3.3.3
| Adsorption isotherm
The increase in the inhibition efficiency for mild steel in
0.5 N HCl solution with increasing inhibitor concentration
can be explained on the basis of additive adsorption. In fact,
for this medium, the additives show linear plots of C/θ versus
C. This behaviour suggests that the adsorption obeys the
Langmuir adsorption isotherm (as shown in Figure 5)
according to the following equation:[37]
θ
¼ KC
(6)
1−θ
where θ is the surface coverage, C the concentration of inhib-
itor and K the adsorption equilibrium constant representing
the degree of adsorption.
The constant K is obtained from the reciprocal of the
intercept of Langmuir plot lines. K and the slope of these
lines are near unity which means that each inhibitor molecule
occupies one active site on the metal surface. The high value
of K indicates that the inhibitor is strongly adsorbed on the
metal surface. The negative free energy values during adsorp-
tion can be justified by the adsorption of the inhibitor mole-
cules on the metal surface.[38]
FIGURE 5 Langmuir isotherms of H2L and its complexes at 20°C on mild
steel surface
FIGURE 6 Scanning optical microscopy images of mild steel samples in
acidic medium: (A) without inhibitors; (B) with complex 1 as inhibitor

8
A. M. NASSAR ET AL.
solid‐state DC electrical conductivity data indicate semicon-
ducting properties of the complexes.
ACKNOWLEDGMENTS
We are grateful to Dr Mustafa Abo El‐Soud from the
National Research Center, Dokki, Giza for recording the
antibacterial activity.
FIGURE 7 Chemisorption of metal complex with mild steel surface via
coordination bonds
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4
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A. M. NASSAR ET AL.
9
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