Mn2+-complexes of N,O-dihydrazone: Structural studies, indirect band gap energy and corrosion inhibition on aluminum in acidic medium

Article abstract of DOI:10.4067/S0717-97072018000404180

Mononuclear manganese complexes of dihydrazone derived from the condensation of oxaloyldihydrazide with 2-hydroxynaphthaldehyde and 2-methoxybenzaldehyde have been synthesized. The hydrazone Schiff base ligands and their chelates were characterized on the

Full text of DOI:10.4067/S0717-97072018000404180

J. Chil. Chem. Soc., 63, Nº 4 (2018)
Mn2+-COMPLEXES OF N,O-DIHYDRAZONE: STRUCTURAL STUDIES, INDIRECT BAND GAP ENERGY AND
CORROSION INHIBITION ON ALUMINUM IN ACIDIC MEDIUM
AYMAN H. AHMED1,2*, A.M. HASSAN2, HOSNI A. GUMAA2, BASSEM H. MOHAMED2, AHMED M. ERAKY2
1Chemistry Department, College of Science and Arts, Jouf University, Gurayat, Saudi Arabia.
2Department of Chemistry, Faculty of science, Al-Azhar University, Nasr City, Cairo, Egypt.
ABSTRACT
Mononuclear manganese complexes of dihydrazone derived from the condensation of oxaloyldihydrazide with 2-hydroxynaphthaldehyde and
2-methoxybenzaldehyde have been synthesized. The hydrazone Schiff base ligands and their chelates were characterized on the basis of their elemental analyses,
spectral (UV-Vis., IR, mass, 1H/13C NMR), magnetism, ESR and thermal (TGA) measurements. The dihydrazone has been suggested to coordinate to the metal
center in bi-dentate manner forming 1:1[M:L] complex. The complexes are suggested to have tetrahedral/octahedral stereochemistry. Optical transmission spectra
were recorded in the range 190–2100 nm and optical band gap energy was determined. The band gap energy (Eg) for all separated compounds lies in the range of
semiconductors. On the other hand, the inhibition and adsorptive properties of the ligands for the corrosion of aluminum in 1 M HCl solutions were studied using
traditional weight loss measurements. The results revealed that bis(2-methoxy-benzaldehyde)oxaloyldihydrazone has a greater inhibition than bis(2-hydroxy-1-
naphthaldehyde)oxaloyldihydrazone. The adsorption of the inhibitors on metal surface was found to be spontaneous first order reaction and consistent well with
the mechanism of physical adsorption. The adsorption data fitted well to Freundlich, Langmuir and Frumkin adsorption isotherms.
Keyword: Metal complexes; Hydrazones; Structure elucidation; Corrosion.
and electronic properties. The work was also undertaken to explore the use
INTRODUCTION
of selected oxaloyldihydrazones as corrosion inhibitors for aluminum in
hydrochloric acid solution where some kinetic and thermodynamic parameters
were calculated and discussed.
Hydrazones are characterized by the presence of the triatomic
grouping C=N-N-. They can be considered as Schiff bases derived from
acid hydrazides. The most important property of hydrazones is their
high physiological activity [1]. They are important for drug design,
organocatalysis and also for the syntheses of heterocyclic compounds [2].
Furthmore, hydrazones and their transition metal complexes have shown
their applications in biological systems, analytical chemistry and as catalysts
[3]. In the context of complexation, hydrazones (derived from o-hydroxy
aldehydes) containing amide, azomethine and phenol functions are recently
recognized as polyfunctional ligands which can react with metal ions either
in keto form and enol form [4]. Extensive studies have revealed that the lone
pair on the nitrogen of the azomethine group (-CH=N-) is responsible for their
chemical and biological activities [5-7]. Although a series of transition metal
complexes of hydrazones [M= Cu(II), Ni(II), Mn(II), Pd(II), Co(II), V(IV)
and ruthenium(II)] were synthesized and characterized [8-15], a survey of
literature revealed that only a few metal complexes of oxaloyldihydrazone
have been investigated [16-18]. The reason of that arises from their poor
solubility. Oxaloyldihydrazones are only soluble in high polar solvents such as
dimethylformamide (DMF) and dimethylsulfoxide (DMSO) and this requires
much effort to isolate them or their complexes in pure form. Interestingly, some
hydrazone complexes have been accommodated in zeolite-Y by alternative
methods and the structures of the solid complexes inside the zeolite were
sustained by variety of physico-chemical techniques and spectroscopic studies
[19-24].
Inhibitor applications often play an important role in oil extraction and
processing industries, heavy industrial manufacturing, water treatment
facility, water-containing hydraulic fluids, water treatment chemicals, engine
coolants, ferrous metal cleaners, automatic transmission fluids, automotive
component manufacture, cutting fluids etc. to minimize localized corrosions
and unexpected sudden failures. Acid solutions are widely used in industry,
the most important fields of application being acid pickling and industrial acid
cleaning due to the general aggressively of acid solutions; thus enhancing
corrosive attack on metallic materials. The most important aspect of inhibition
normally considered by corrosion scientists is the relation between molecular
structure and corrosion inhibition efficiency. Most of the well-known acid
inhibitors are organic compounds containing p bonds, phosphorus, sulfur,
oxygen and nitrogen as well as aromatic rings in their structure which are the
major adsorption centers [25]. Heteroatoms such as nitrogen, oxygen and sulfur
are capable of forming coordinate covalent bond with metal owing to their
free electron pairs and thus acting as inhibitors. Unfortunately, many common
corrosion inhibitors employed in aqueous media are health hazards [28].
In view of the significant role played by the hydrazones and their metal
complexes in various fields, we motivated to synthesize and characterize of
two oxaloyldihydrazones and their Mn2+ complexes. The band gap energy
of the investigated compounds has been estimated to describe their optical
EXPERIMENTAL
Chemicals and Equipments
The aldehydes were obtained from Merck Company. Manganese (2+)
acetate, diethyl oxalate, hydrazine monohydrate were purchased from Sigma-
Aldrich. Oxalicdihydrazide was prepared by reacting diethyloxalate (1 mol)
with hydrazine hydrate (2 mol) [10], (Exp/Lit. m.p = 240/240 °C). Other
employed chemicals and solvents were of highest purity and used without
further purification. Elemental analyses (CHNM), spectral (UV-Vis., FT-IR,
Mass) and thermal (TG) measurements were executed as mentioned elsewhere
[29,10]. The NMR spectra were recorded on a Varian mercury VX-300 NMR
spectrometer. H spectra were run at 300 MHz and 13C spectra were run at
1
75.46 MHz in dimethylsulphoxide (DMSO-d ). Tetramethylsilane (TMS)
was used as an internal reference and chemical6shifts are quoted in δ (ppm).
Magnetic moments (μeff) of the samples were measured at room temperature
(RT) using Faraday’s method. The ESR spectra of the complexes were recorded
in powder form at room temperature using Bruker-EMX-(X-bands-9.7 GHZ)
spectrometer with 100 KHZ frequency, microwave power 1.008 mW and
modulation/amplitude of 4 GAUSS at national center for radiation research
and technology, Egypt. The band gap energy (E ) of product compounds was
calculated according to Tauc’s equations [30,31].gCorrosion test was performed
on freshly prepared sheet of aluminum of chemical compositions (wt %) Ni
(0.006), Pb (0.01), Zn (0.003), Cu (0.002), Fe (0.27) and Al (99.71). The metal
specimens were prepared, degreased and cleaned as previously described [10].
Synthesis of oxaloyldihydrazone ligands
The
dihydrazone
ligands;
bis(2-hydroxynaphthaldehyde)
oxaloyldihydrazone (L1) and bis(2-methoxybenzaldehyde)oxaloyldihydrazone
(L2) were prepared using the same procedure described in Ref. 10 as follows.
Aqueous methanolic solution of oxaloyldihydrazine was mixed with absolute
methanolic solution of the selected aldehydes (maintaining the molar ratio
at 1:2) in hot condition. The reaction mixture was heated for about 3 hours
accompanied by stirring. The precipitate was separated out on concentrating
the solution to half of its volume and cooling. Afterwards, the isolated
compound was filtered, recrystallized from DMF-MeOHaq mixed solvent,
°
washed thoroughly by acetone and then dried in an electric oven at 80 C for
2 h. The molecular structure of the resulting hydrazones was confirmed by
1
elemental analyses, (IR, mass and H/13C NMR) spectroscopy (Table 1, 2).
The mass spectral data of the ligands (L1, L2) are displayed in Fig. 1 and the
obtained M+ values are set out in Table 1. Mass spectrum of the bis(2-hydroxy-
1-naphthaldhyde)oxaloyldihydrazone (L1) revealed ion peaks at m/e = 427
(20%, molecular peak, (C16H14N4O4)), 410 (25%, M+(C24H18N4O3)), 270 (21%,
M+(C13H10N4O3)), 157 (21.2%, M+(C11H9O)), 257 (50%, M+(C13H10N3O3)), 242
(40%, M+(C13H9N2O3)), 175 (71%, M+(C11H9NO)), 141 (4%, M+(C11H9)), 128
e-mail: ayman_haf532@yahoo.com
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J. Chil. Chem. Soc., 63, Nº 4 (2018)
(37%, M+(C10H8)), 79 (40%, M+(C6H8)), 64.95 (71%, M+(C5H5)) and 50 (100%,
base peak, M+(C4H ). Mass spectrum of the bis(2-methoxybenzaldehyde)
oxaloyldihydrazone 4 (L2) exhibited ion peaks at m/e = 356.45 (33%,
molecular peak, (C18H18N4O4)), 247.35 (26%, M+(C11H11N4O3)), 217.2 (62%,
M+(C10H9N4O2)), m/e = 50 (100%, base peak, M+(C4H4)), 234.2 (42%,
M+(C10H10N4O3)), 202 (25%, M+(C10H9N2O3)), 176.5 (60%, M+(C9H9N2O2)),
149 (92%, M+(C8H9N2O)), 190.6 (46%, M+(C9H8N3O2)), 175.55 (78%,
M+(C9H7N2O2)), 119.4 (53%, M+(C7H7N2)), 104.2 (51%, M+(C7H6N)), 134.6
(82%, M+(C8H8NO)), 107.15 (52%, M+(C7H8O)), m/e = 90.6 (75%, M+(C7H6)),
77.05 (86%, M+(C H )) as well as 66.65 (38%) corresponds to M+(C5H6). The
spectral data assur6ed6the chemical formula of the ligands as well the correct
molecular weights.
Fig.1. Mass spectra of L1 and L2.
Synthesis of solid complexes
in DMF. All test solutions contained 10 ml (20 vol.%) of DMF to ensure
solubility. The concentrations of the used inhibitors (L1 and L2) were 5×10-4, 1
×10-4, 5 × 10-5 and1×10-5 M.
Manganese (2+) complexes were prepared by literature method [10].
Dihydrazone (1 mmol) were dissolved first in 20 ml DMF and then 50 ml
methanol was added. The obtained solution was added slowly to a 20 ml
methanolic solution of Mn (2+) acetate in 1:2 molar ratio. The resulting
mixture was refluxed for 4 hours and then reduced to 30 cm3 by gentle
heating. The desired product obtained on cooling down to room temperature
was filtered off, washed vigorously in beaker with successive portions of hot
solvents (dimethylformamide, methanol and acetone, respectively) and dried in
an electric oven at 80 °C for 8 hrs.
Gravimetric measurements
A simple test for measuring corrosion is the gravimetric method (weight
loss). The method involves exposing a clean weighed piece of the metal or
alloy to the corrosive environment for a specified time followed by cleaning
to remove corrosion products and weighing the piece to determine the loss of
weight.
A previously weighed metal (aluminum) was completely immersed
in 50 ml of the test solution (1M HCl) in an open beaker. The beaker was
maintained at a temperature of 301 K. The aluminum coupons were suspended
in the beaker with the help of glass hooks. The coupons were withdrawn from
the test solution at 1 h interval progressively for 4 h, dipped into saturated
Solutions
The aggressive solutions, 1 M HCl, were prepared by dilution of analytical
grade HCl with distilled water. Stock solutions of the inhibitors were prepared
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J. Chil. Chem. Soc., 63, Nº 4 (2018)
ammonium acetate solution to terminate the corrosion reaction, washed by
concentrations of the acid was determined using the following equation [33]:
Where: W = weight loss to the nearest 0.0001gm, A = area of specimen
to the nearest 0.01 cm2 (2 × area of one face) and t = exposure time (minutes).
The inhibition efficiency (%h) for each inhibitor and the degree of surface
coverage (θ) were calculated using the equations [33]:
scrubbing with a light brush, rinsed in acetone and finally dried in an oven at
°
80 C before reweighing. The difference in weight in grams was taken as the
total weight loss (W) [32].
Where: Wi and Wf are the initial and final weights for aluminum,
respectively.
The corrosion rate, ρ (in gm.cm-2.min-1) for aluminum at different
Where ρ1 and ρ2 are the corrosion rates in the presence and absence of
inhibitor, respectively.
Table 1: Analytical, physical and spectroscopic data of the oxaloyldihydrazones and their related Mn(II) complexes.
1H-NMR/ 13C-NMR
Chemical shift
μeff.
4A/
6A
M+
Found/
calcd.
Found(calcd.)%
M.p(oC)
Color
Eg
(eV)
Compound
Symbol
g┴
g//
gav
C
H
N
M
(δppm
)
12.76(NH, s),
12.57(OH, s),
9.74(CH=N, s),
6.5-8.5 (aromatic
protons, m)/
>300
Yellow
68.1
(67.7)
5.3
(4.3)
12.9
(13.2)
-
-
427.1/
426.4
C24H18N4O4
L1
-
3.06
-
-
-
158.4,158.3,
157.1,155.7,
150.4,149.7, 133.2,
132.5, 128.9, 127.8,
123.6, 120.5,118.8
12.3(NH, s),
11.9(OH, s),
8.95(CH=N, s),
6.8-8.5(aromatic
protons, m),
>300
White
59.2
(61)
5.8
(5.1)
15.7
(15.8)
-
-
355.0/
354.4
C18H18N4O4
L2
-
3.42
-
-
-
3.93(OCH3,
s)/156.2, 148.7,
158.1, 132.1, 125.8,
122.0, 120.9, 55.7.
[Mn(L1-H)
(H O) ].0.5
2H22O
>300
Brown
54.2
(54.9)
4.0
(4.2)
10.2
(10.5)
10.8
(10.7)
4.4/
4.7
I
-
-
-
-
2.38
3.39
1.88
1.85
1.930
1998
1.9
1.9
[Mn(L2-2H)
(OAc)2(DMF)2]
>300
Yellowish
49.1
(49.9)
5.2
(5.7)
12.1
(12.5)
9.2
(8.2)
5.1/
4.7
II
cm-1 indicating the interaction of dihydrazide with the aldehyde. (2) The bands
RESULTS AND DISCUSSION
at 3476, 3166 and (1705(m) +1660(v.s) cm-1 are attributed to ν(OH)
,
ν(NH) and ν(C=O) for L1, respectively, meanwhile, L2 which does not cnoanphtthaoiicn
o-hydroxy group revealed ν(NH) at 3202 cm-1 and ν(C=O) at 1653 cm-1. (3)
Both ν(C=O) and ν(NH) are observed concurrently in the IR spectra of the
two ligands giving an important reflection to the existence of keto forms. (4)
Appearance of ν(C=O) as two bands in case of L1 (1705, 1660 cm-1) and one
band for L2 (1653 cm-1) refers to the mixture of [cis(syn/anti-cis-structure) +
trans(staggered)-structure)] isomers for L1 and trans (staggered) structure for
L2, respectively, (Fig. 2). This suggestion is assumed in the light of the field
effect criteria where the cis isomer is absorbed at higher frequency than trans
one [10]. The band observed at 1705 cm-1 characterized cis form while the other
two bands located at 1660 and 1653 cm-1 indicated the presence of trans-form.
Otherwise, NMR spectroscopy provides satisfactory agreement with this view
as shown later. (5) The band observed at 3227 cm-1 (L2) may concern ν(OH)
The analytical and spectroscopic data as well as physical properties of the
ligands and coordination compounds are listed in Tables 1,2. The results show
that the dihydrazone ligands coordinate to the manganese ion in a 1:1 molar
ratio. The prepared complexes are non-hygroscopic, soluble in strong polar
solvents such as in DMF and DMSO but insoluble in H2O, CHCl , CCl ,alcohol,
acetone, ether as well as benzene. The calculated and found3perce4ntages of
elements (CHNM%) point out that the compositions of the isolated complexes
harmonize with the proposed formulae.
Stereo chemical structure of Mn2+-complexes
IR Spectra
The comparative IR spectral study of the dihydrazones and their manganese
(2+) complexes reveals the coordination mode of the ligands during complex
formation. The important frequencies with their possible assignments are listed
in Table 1.
Free ligand: Prepared oxloyldihydrazones can be represented either in the
trans- (staggered) or cis- (anti-cis- or syn-cis-) configuration as revealing in
Fig. 2 [34]. The structures of the ligands were authenticated on the basis: (1)
the ν(C=N) was observed in the IR spectra of the two ligands at 1600-1618
in H-bond bonding with C=N. (6) The bands corresponding to ν(C-O) in
enolic
the spectra of L and L were locatedat 1287 and 964 cm-1, respectively [35,36].
(7) Also, each1ligand2contains bands in the regions 1400-1600 and one band
at 700-800 cm-1 assignable to ν(C=C) and δ(CH)out of plane of the aromatic ring,
respectively.
Mn2+-complexes: The IR spectra of the complexes (I, II) are depicted
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J. Chil. Chem. Soc., 63, Nº 4 (2018)
in Table 2. Virtually, it is difficult to unambiguously assign the bands due to
foundation of various groups e.g., C=O, C=N, N-C=O, NH, OH, ph ring, C-N
and C-H which absorb in the overlapping regions. On examining the IR spectra
of the complexes in comparison with those of the free ligands, the complexes
structures are suggested based on the following evidences:
of ν(NH) with noticeable of ν(C=O) at 1655 cm-1 upon chelation give a
conclusive evidence to the coordination of metal ion with deprotonated NH
group. Persistence of ν(C=O) as one band, 1655 cm-1 instead of two bands
in L1 with slightly change in its location agrees with the presence of the two
oxygen atoms related to (O=C-C-OH) part, in trans direction. Doubtless the
positive shift observed for ν(C-O) vibration (1287→1299 cm-1) provides a
clue to deprotonation of both phenolic (OH) groups during their coordination
[6,36,39] but remarking ν(OH)naphthoic at 3186 cm-1 points to the existence of
naphthoic (OH) in H-bonding, Fig. 4. This can be explained on the basis of the
deprotonation of one naphthoic (OH) group while the other is in H-bonding
with C=N group.
For I, the splitting of the azomethine band (1600, 1618 cm-1) justifies the
imparting of this group in coordination [37] where two dissimilar azomethine
groups are presented, Fig. 3 [38]. One of them is coordinated with manganese
while the other is free. This occurred due to the enolization of left or right
half of the dihydrazone molecule assuming keto-enol skeleton. Existence of
ν(OH)naphthoic at 3342 cm-1 endorses this suggestion. Interestingly, the obscure
Table 2: Significant IR and electronic absorption data of oxaloyldihydrazones and their manganese complexes.
ν
naphMth-Ooic /
enolic
δC-H
aromatic,
out of
ν(OH)
naphthoic
(enolic)
νC-H
aromatic
(aliphatic)
ν
Symbol
νH2O
νNH
νC=O
νC=N
naphCt-hOoic/
methoxy
Structure
νC-
λmax (nm)/
assignments
OMe
(carbonyl/
methoxy)
plane
392(n-π*, C=N),
375(n-π*, C=O),
330(π -π*, C=N),
315(π -π*,
C=O), 303(π -π*,
Phenyl)
3476
(-)
3043
(2926)
1705
1660
-/-
(-/-)
L1
-
3166
1621
1287
-
741
760
-
350(n-π*, C=N),
338(n-π*, C=O),
306(π -π*,
C=N), 292(π -π*,
C=O), 280(π -π*,
Phenyl)
-
3202
3041
(2940)
-/-
(-/-)
L2
-
1653
1655
1600
-
964
-
(3227)
3186
(3342)
3051
(2925)
1618
1600
567/-
(-/-)
460 (6S→4G)
425 (6S→4P)
Tetra-
hedral
I
3379
-
-
1299
-
-
745
761
-
(-)
3040
(2936)
1652
strong
-/-
(-/640)
430/6A1g→4T1g(G)
380/6A1g→2T2g(G)
Octa-
hedral
II*
3203
1600
964
* νC=O of DMF was observed at 1729 cm-1, νOAc was observed at 1351and 1437 cm-1
In case of II and comparing with the IR spectrum of L2 ligand, The mode
of chelation is supported by following facts: (1) The remaining of the ν(NH),
ν(C=N) and ν(C-OMe) at nearly the same position with disappearance of ν(M-
spectra (L1) implies the ligand is present in the keto form. In the spectrum of the
L2, noticeable of δ(NH) at 12.29 ppm accompanied with very weak intensity of
δ(OH)enolic at 11.92 ppm infers domination of diketo form compared with dienol
or keto-enol form. Surely, NMR spectroscopy can determine the skeleton of
the oxaloydihydrazone. As reported in literature [41], if the dihydrazone exists
in the syn-cis-configuration or staggered configuration, the δOH, δNH and
δCH=N, resonances, each should appear as a singlet. However, the appearance
of these signals in the form of six signals (doublet of doublet) indicates anti-
cis (chair) configuration. Based on this supposition together with the obtained
data it can be concluded that the two ligands are in conformity with syn-cis-
or staggered configuration. But the IR information clearly indicates that the
staggered configuration is well defined.
O
) excludes involvement of the above groups in bonding. The feature of
thmeetihnoxtyeresting band assigned to ν(C=O) is varied from strong to very strong
relative to other neighbor bands indicating the coordination of the above group
with manganese ion. Undoubtedly, the observation of C=O groups in the IR
spectrum of II with obscure of ν(NH) band implies that the ligand prefers
its keto form to coordinate with Mn2+ ion. DMF showed band at 1729 cm-1
assigned to the ν(C=O) stretch and this band is not observed in the free ligand
[40].
Actually, the appearance of new M-O bands in both complexes (Table
3) asserts the formation of metal-ligand bond [41]. The broad bands within
the range 3300–3450 cm-1 in the complex I are attributed to the presence
of crystalline and coordinated water molecules. Besides, the two bands
corresponding to νas and νs carboxylic modes of acetate group are shown at 1351
cm-1 and 1437 cm-1, respectively. The larger difference between the above two
frequencies [ν (OAc) and νs(OAc)] is an indicative of monodentate nature of
acetato group a[s42]. Also, the slight changes in the positions of ν(C=C)ph upon
chelation are due to metal-ligand interaction.
The 13C NMR measurement was performed for the two ligands and the
1
resulting data (Fig. 3, Table 1) correlate well with H NMR spectra without
discrepancies. The 13C NMR spectroscopy data for L1 gave resonances at 155.7
ppm for C=O, 149.7 ppm for C=N and 158.3, 132.5, 128.9, 127.8, 123.6 for
Ar-C2, Ar-C4, Ar-C6, Ar-C5, Ar-C1, respectively, confirming a staggered
symmetry for L1. The greater number of lines than the carbon types in the
proposed configuration may be recognized from the existence of some ligand
molecules in diketo and/or dienol forms. 13C NMR spectrum of L2 was also
1H/13C NMR Spectra
1
in good conformance with H NMR assuming a keto-keto form (staggered
The spectra of the two investigated ligands (Table 1) contain signals
at 6.5–8.5, 9.73, 8.95 12.8, 12.6, 12.3 and 3.9 ppm which are assigned to
aromatic(L1, L2), azomethine (L1), azomethine (L2), -NH (L ), naphthoic –
OH (L1), secondary -NH (L2) and -OCH3(L2) protons [11], re1spectively. The
disappearance of signals due to –OH and –NH protons, on D2O exchange
confirms to the assignment of signals in the ligand spectrum. In the spectrum of
the L1. the signal due to the naphthoic –OH protons appeared as a broad singlet
suggesting its capability to establish intermolecular H-bonding with DMSO
solvent. A probable existence of interamolecular hydrogen bonding between
naphthoic -OH and CH=N group was ruled out owing to its appearance at normal
position. The observation of δ(NH) with nonexistence of enolic OH in the same
structure). The spectrum showed three distinct peaks assigned to C=O (156.2
ppm), C=N (148,7 ppm) and OCH3 (55.7 ppm). The peaks centered at 158.1,
132.1, 125.8, 122.0, 120.9 are due to Ar-C2, Ar-C4, Ar-C6, Ar-C5, Ar-C1,
respectively. The signals observed in-between 25-50 ppm are attributed to the
exploited solvent (DMSO).
Optical absorption and magnetic investigation
The electronic spectral bands for the oxaloylhydrazones and their
manganese complexes in DMF solvent have been given in Table 2. The
electronic spectra of the investigated ligands shows peaks at 350-392, 338-
375, 306-340, 292-315 and 280-303 nm due to (n-π*, C=N), (n-π*, C=O), (π-
π*, C=N), (π-π*, C=O) and (π-π*, aromatic ring), respectively. The proposed
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J. Chil. Chem. Soc., 63, Nº 4 (2018)
geometry of the isolated complexes was deduced from the electronic spectra
overlapped with strong ligand bands [44].
in DMF and magnetic measurements. The optical absorption spectrum of
The room temperature magnetic moments (μeff., Table 1) of manganese
complexes under study supported the proposed stereochemistry derived from
the electronic data [44]. The μeff values of compounds I and II, measured at 4
and 6 amperes (A), are close to spin only value and refer to high spin complexes.
In other words, the results point to the presence of 5 unpaired electrons in both
complexes confirming a paramagnetic nature and illustrating that Mn2+ ions
were not oxidized by dihydrazone ligands.
6
I revealed two absorptions at 460 and 425 nm assigned to S→4G and 4P
transitions, respectively. The observation of these bands suggests a tetrahedral
environment around the Mn2+ ion [43]. The electronic spectra of II showed
6
4
two weak bands at 430 and 380 nm assignable to A1g→4T1g(G) and T2g(G)
transitions, respectively, considering an octahedral configuration around the
4
Mn2+ ions [44]. Other bands attributed to octahedral (6A1g→4Eg, A1g, 4T2g(D)
and 4Eg(D)) and tetrahedral (6S→4D and 4F) structures of Mn2+-complexes are
Fig. 2. Proposed structures of oxaloyldihydrazone ligands:
Where L1 (R=ph, X=H), and L2 (R=H, X=CH3).
Table 3: Corrosion parameters for aluminum coupons in (1M HCl) solution in presence and absence of different concentrations of hydrazone ligands obtained
from weight loss measurement at room temperature after 4 hrs immersion time.
Conc.
(M)
Weight Loss
Rate of corrosion
mg cm-2 h-1
Inhibition efficiency
Inhibitor
θ
(g)
h %
Blank
5 × 10-4
1 × 10-4
5 × 10-5
1 × 10-5
0.0070
0.0046
0.0051
0.0054
0.0059
-
-
0.0210
0.0139 0.0153
0.0163
0.34
0.27
0.22
0.15
34
27
22
15
L1
0.0178
Blank
5 × 10-4
1 × 10-4
5 × 10-5
1 × 10-5
0.0188
0.0067
0.0084
0.0099
0.0114
0.0063
0.0022
0.0029
0.0035
0.0038
-
-
0.64
0.54
0.44
0.39
64
54
44
39
L2
4184

J. Chil. Chem. Soc., 63, Nº 4 (2018)
Fig. 3. 13C NMR spectra of L1 and L2 in DMSO-d6 at RT.
Thermogravimetric analysis
of MnO at the final stage (found/calculated % =11.1/10.5).
The thermogram (TGA) data of complex I exhibited five steps of
decomposition. The first step (50-150 °C) corresponds to the loss of 1/2 H2O-
crystalline (found/calculated % = 2.0/1.7), while the other four stages are related
to the gradual decomposition of the complex with formation of speculated
metal oxides at the final step. TG of II gave four stages of decomposition.
On the basis of the various physico-chemical and spectral
studies presented above and their discussion, the complexes may be
suggested to have structures shown in Fig. 4
Electron paramagnetic resonance (EPR)
The complexes I and II have been analyzed by EPR spectroscopy at room
temperature and their spectra are shown in (Fig. 5). The ground state of both
°
The first one (140-280 C) is attributed to the loss of 2OAc- groups (found/
6
calculated% = 18.0/17.5) while the other three stages (300-380, 400-550, 600-
octahedral and tetrahedral d5 ion (spin free) is A. The energy levels are at
800 °C) concerned to the gradual decomposition of the complex with formation
±5/2, ±3/2, ±½ and a single isotropic resonance at g = 2 is easily attained.
4185

J. Chil. Chem. Soc., 63, Nº 4 (2018)
Accordingly, the EPR spectra of complexes Iand II exhibited one broad signals with two “g” values (g//, g┴), Table 1. The absence of hyperfine splitting for these
complexes ruled out any axial distortion. The average g values were calculated using the formula:
Where, the calculated values are inserted also in Table 2.
Fig. 4. Suggested structures of MnII-oxaloyldihydrazone complexes.
Fig. 5. The plots of (αhʋ)2 vs. hʋ of Mn(II)-oxaloyldihydrazone complexes.
The average g values (gav) of I and II complexes are nearly 1.9 close to
free electron g value indicating the absence of spin-orbit coupling. The g// is a
moderately function for covalent character. The complexes I and II exhibited
g at 1.930 and 1.998, respectively which are less than 2.3, implying covalent
n/a/ ture of manganese-ligand bond in the present complexes, where, g//˃2.3
is character of ionic metal-ligand bond [45,46]. Fairly high values of g are
in conformity with the oxygen and nitrogen coordination in these compounds
[18].
coefficient, Eg is the optical energy gap, A is the constant, m is equal to 1/2 and
2 for direct and indirect transitions, respectively. The absorption coefficient (α)
was calculated from to the relation:
Where d is the width of the cell and T is the measured transmittance.
Optical band gape energy
The energy band gap of the ligands and their manganese complexes has
been calculated with the help of absorption spectra. To calculate the optical
band gap energy from absorption spectra, Tauc’s relation is used [30,31]:.
We plot a graph between (αhʋ)2 versus hν, the extrapolation of the straight
line to (αhʋ)2 = 0 axis gives optical band gap. Figure 6 shows Tauc’s plot for all
the samples. Band gap for ligands and Mn2+-complexes in DMF solvent lies in
the ranges 3.06-3.42 eV and 2.38-3.39 eV, respectively and are mentioned in
Table 1. Careful inspection of Table 1 shows that the optical band gap values of
the two ligands (Eg) are minimized upon complexation. This reduction in band
gap is attributed to the mobilization of the electrons toward the metal [47]. It
is suggested that the manganese raises mobilization of the ligand electrons by
Where hν is the photon energy, h is Planck’s constant, α is the absorption
4186

J. Chil. Chem. Soc., 63, Nº 4 (2018)
accepting them in its shell thus the width of the localized levels in the resulting
complex is expanded and in turn, the band gap is smaller. This result has many
applications in optics, electronics and energy conversion devices [48]. In fact
small band gap facilitates electronic transitions between the HOMO-LUMO
energy levels and makes the molecule more electro-conductive [49]. From the
observed values of optical properties, it is seen that the synthesized compounds
lie in the same range of highly efficient photovoltaic materials [47,50] and can
be used as semiconductors.
Table 4: Values of ΔGads (kJ/mol), adsorption constant (K), Freundlich heterogeneity factor (n), lateral interaction term describing the interaction in adsorbed
layer (α), and correlation coefficient (R2) of inhibitors for aluminum coupons.
Freundlich isotherm
Langmuir isotherm
Frumkin isotherm
Ligands
Kads
∆Gads
Kads
∆Gads
Kads
∆Gads
1/n
n
R2
R2
α
R2
(molL-1) (kJ mol-1)
(molL-1)
1.6994
1.7006
(kJ mol-1)
(molL-1)
(kJ mol-1)
L1
L2
0.207
0.129
4.83
7.75
2.5420
1.2680
-12.39
-10.65
0.9767
0.9624
-11.38
-11.38
0.9984 6.471
0.9991 3.539
0.00013
0.00044
-12.32
-9.29
0.8287
0.9588
Fig. 6. The efficiency for aluminum coupons in 1M HCl solution in the presence and absence of different concentrations of oxaloyldihydrazone
ligands at room temperature.
Weight loss studies
Kinetic treatment
The inhibitive actions of ligands on the dissolution of aluminum in
1 M HCl are investigated (Table 3). Careful inspection of these results
showed that, by increasing the inhibitor concentration, both (θ) and (h%) are
increased while the weight loss is reduced. So, the dissolution of aluminum
in the presence of the investigated inhibitor can be interpreted on the basis
of interface inhibition mode where the adsorbed species mechanically screen
the coated part of the metal surface from the action of the corrosive medium.
This adsorption is related to some discriminated features of the investigated
oxaloyldihydrazones like the functional groups (OH, NH, C=O and C=N),
quite large molecular size, high molecular weight, aromaticity, electron
density at the donor atoms (N:, :O:) and also the p-orbital character of
the azomethine, carbonyl and phenyl groups. Adsorption and desorption
of inhibitor molecules continuously occur at the metal surface and an
equilibrium exists between the two processes at a particular temperature.
The inhibition achieved by dihydrazone compounds decreases in the
sequence L2 < L and this behavior can be discussed on the basis that the methoxy
group (in L2) is1more electron donating group than naphthoic hydroxyl (in L2).
Positive mesomeric effect (+M effect) and negative inductive effect (–I effect)
of -OCH lead to push the electron density toward the donor atoms of O=C-NH-
N=CH p3art via the phenyl ring as a result of delocalization and subsequently
enhance the bond strength between the molecule and the corroding aluminum
surface. Under these circumstances, the inhibition efficiency of (L2) should be
higher than that of L1. Furthers, existence of two naphthyl rings in the structure
of L1 may cause crowding that encourages incomplete coating besides they can
withdraw the electron pair of ortho-OH group toward themselves under the
influence of resonance phenomenon. The iinhibition efficiency values of the
dihydrazone are set out in Table 3 whereas the maximum protection efficiency
percentages of L1 is 34% while that for L is 64%. The obtained data boosts
the application of L2 as plausible corrosion2pickling inhibitor. It is evident that
the inhibition efficiency values is not extremely high owing to the lower ligand
concentrations used.
Spontaneous corrosion reaction is a heterogeneous one, composed of
anodic and cathodic reactions with the same rate. On this basis the kinetic
analysis of the data is considered necessary. In the present work, the initial
weight of aluminum coupon at time t is designated as Wi, the weight loss is ΔW
and the weight change at time t is (Wi - ΔW) or Wf. When logWf was plotted
versus time, a linear variation was remarked (Fig. 7) considering a first order
reaction [51]:
Where: Wi = the initial weight before immersion, K = slope (the rate
constant), and t is time.
From the rate constant values obtained from the graphs, regarding the
adsorption reaction on manganese surface proceeds as first order, the half-life
values (t ) of the metal in the test solutions were calculated using the equation
½
[51]:
Adsorption isotherms and Gibb’s free energy
The adsorption behavior of oxaloydihydrazones on the surface of aluminum
was studied by fitting data calculated for the degree of surface coverage to
different adsorption isotherms such as Freundlich, Temkin, Langmuir,
Frumkin, and Flory-Huggins. The obtained correlation coefficients (R2) were
used to determine the best fits [52]. The Gibb’s free energy of adsorption ΔGads
at room temperature, which can characterize the interaction of adsorbed
molecules and aluminum surface, for all choice isotherms was calculated
using the formula [53]:
4187

J. Chil. Chem. Soc., 63, Nº 4 (2018)
Freundlich isotherm
Freundlich isotherm can be expressed according to the equation [54],
Where C is the inhibitor bulk concentration in mol/L, Kads.(mol-1L) is the
equilibrium constant of adsorption and n is Freundlich heterogeneity factor.
Plots of lnθ versus lnC of all ligands (Fig. 8) were linear with
acceptable correlation coefficient which illustrate the applicability
of Freundlich isotherm. The thermodynamic parameters obtained for
the Freundlich adsorption isotherm were recorded in Table 4. This
demonstrates that values of the adsorption parameter n are greater than unity
implying favored adsorption process [55].
Langmuir adsorption
The assumption establishing Langmuir adsorption isotherm is given by
the equation [56],
Where: C is the concentration of the inhibitor in the bulk electrolyte, θ is the
degree of surface coverage of the inhibitor and K is the equilibrium constant of
adsorptionoftheinhibitor.PlottingoflogC/θagainstlogCat301K(Figure8)forthe
two inhibitors gave straight lines with high correlation coefficients (very close to
unity) indicating that Langmuir adsorption isotherm is applicable to the
adsorption of oxaloylhydrazones on the surface of aluminum. Values
of Langmuir parameters deduced from the isotherms are recorded in
Table 4. The application of Langmuir isotherm to the adsorption of
oxaloylhydrazones on surface of aluminum indicated that there is no
interaction between the adsorbate and adsorbent. The equilibrium constant
(K) of the adsorption process can be calculated from the reciprocal of anti
logarithm of the intercept.
Fig. 7. Variation of log W with time (hour) for aluminum coupons in HCl
solution containing oxaloyldihfydrazone ligands (L1 and L2).
Frumkin adsorption
Frumkin adsorption isotherm can be represented by the equation [57],
Where: K is the adsorption-desorption constant and α is the lateral
interaction term describing the interaction in adsorbed layer.
Plots of logC[θ/(1-θ)] versus θ shown in Fig. 8 were linear with acceptable
correlation coefficients suggesting that the adsorption of oxaloylhydrazones
on the aluminum obeys Frumkin isotherm. The values for Frumkin adsorption
parameters are listed in Table 4. It is of interest to note that, the positive α values
reflect the attractive behavior of the inhibitor on the surface of aluminum [58].
Adsorption consideration
The inhibition effect of inhibitor compound is ascribed to adsorption of
the molecule on metal surface. This adsorption may be physical or chemical
depending on the adsorption strength. For a physical adsorption mechanism,
the value of the free energy (ΔGads,) of adsorption should be less negative than
-40 kJ/mol required for chemical adsorption. The negative values of ΔG set
out in Table 4 indicate that the adsorption of oxaloylhydrazones on alumaidnsum
surface is spontaneous and in the same time it characterizes their strong
interaction with the metal surface [59]. Moreover, the magnitudes of ΔG are
in good agreement with physical adsorption mechanism [59]. Accordinglayd,s the
prohibition mechanism is due to the formation of a protective film on the metal
surface where the additive covers both the anodic and cathodic sites.
Fig. 8. Plots of Freundlich lnθ vs. lnC (A), Langmuir Log(C/θ) vs. LogC
(B), and Frumkin LogC[θ/(1-θ)] vs. θ (C) adsorption isotherms for aluminum
coupons in HCl solution containing oxaloyldihydrazone ligands (L1 and L2).
4188

J. Chil. Chem. Soc., 63, Nº 4 (2018)
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CONCLUSION
Inthepresentstudy,wehavepreparedtwomonomericmanganesecomplexes
and characterized them on the basis of data obtained from physicochemical
and spectroscopic studies. The oxaloyldihydrazone ligands (L and L2) can be
synthesized by usual condensation of oxaloyldihydrazide with1the appropriate
aldehyde (2-hydroxynaphthaldehyde and 2-methoxybenzaldehyde) in a 1:2
molar ratio. The stoichiometry of the isolated solid complexes (I, II) is proposed
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on the hydrazone complexes, it is evident that, the dihydrazone in complex I
functions as a keto-enol bidentate donor through deprotonated secondary NH
and naphthoate oxygen giving mononuclear tetrahedral complex. However, in
the case of the complex II, it behaves as a diketo bidentate donor within the two
carbonyl (C=O) groups providing mononuclear octahedral complex. It can be
observed that the band gap of all the samples lies in the range of wide band gap
semiconductors. The band gap values of ligands have been minimized upon
chelation. Corrosion inhibition results reflect the capability of the two ligands
to mitigate the corrosion of aluminum in HCl solution especially L which
contains an electron donating methoxy group in its structure. The in2hibition
efficiency of the ligands was found to increase with increment of the ligand
concentration following the trend: L2 > L1. The obtained negative values of
Gibb’s free energy indicated that the adsorption of the inhibitor is spontaneous
and follows the mechanism of physical adsorption. The experimental data fit
well with Freundlich, Langmuir and Frumkin adsorption isotherms.
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ACKNOWLEDGEMENT
Corresponding author expresses his gratitude to the scientific research
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