Introduction

Ligands containing more than one donor sites are extensively used for the preparation of metal complexes with interesting properties. Transition metal complexes of these ligands containing pyridine, furan rings and their derivatives are also of great interest due to their ability to absorb visible light to act as electron reservoirs are promising factors in their applications as photosensitizers [1–6]. Many pyridines find application in areas where bioactivity is important, as in medicinal drugs and in agricultural products such as herbicides, insecticides, fungicides and plant growth regulators [7]. Chelating agents containing O and N donor atoms show broad biological activity [8–13] and are sensitive to the molecular environment [14, 15]. Also, Schiff base complexes show high catalytic activities in various chemical reactions such as oxidation [16, 17], epoxidation of olefins [18], homogeneous catalyst [19, 20] and polymerization of ethylene to obtain narrow molecular weight distribution [21]. Palladium complexes play an important role in bioinorganic studies in addition to their catalysis applications in many reactions [22, 23]. Schiff base complexes also have catalytic activity toward the oxidation of olefins and alcohols [24]. A lot of attention has been paid on the mono-condensed Schiff bases derived from 1,2-diamines and different aldehydes/ketones. Therefore, the present study deals with the synthesis and characterization of Mn(II) and Pd(II) complexes with mono-condensed heterocyclic Schiff base ligands, 2-[(pyridin-2-ylmethylidene)amino]-6-aminopyridine (L1) and 2-[(2-furylmethylene)]phenylenediamine (L2) (Fig. 1). All compounds were characterized by elemental analyses, UV–Vis, IR, 1H-NMR spectroscopy, conductance measurements, magnetic susceptibility measurements and thermogravimetric analyses (TG). Density functional theory (DFT) calculations were done in order to correlate between the theoretical and experimental results.

Fig. 1
figure 1

Structure of ligands and their complexes

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Experimental

Reagents

All chemicals used were of analytical reagent grade (AR) and of the highest purity available. 2-Furancarboxaldehyde (Sigma), 2-pyridine carboxaldehyde, 2-phenylenediamine, palladium(II) chloride, manganese(II) chloride tetrahydrate and 2,6-diaminopyridine were supplied by Aldrich. The organic solvents are reagent grade chemicals and used as received.

Instruments

Elemental analysis was performed using PerkinElmer 2400 apparatus. Infrared spectra were recorded on a Shimadzu FTIR spectrometer. 1H-NMR spectra were run on a Bruker Avance 300 MHz instrument. UV–Vis spectra were measured on a Shimadzu UV 1800 Spectrophotometer. The photoluminescent properties were studied using a Jenway 6270 Fluorimeter. Magnetic susceptibilities were measured at room temperature using Cambridge, England Sherwood Scientific magnetic susceptibility balance. Mass spectrometry measurements of the solid complexes (70 eV, EI) were taken on a Finnigan MAT SSQ 7000 spectrometer. Molar conductances were obtained using conductometer ORION model 150 of 0.6 Pascal’s constant. Thermogravimetric analyses (TG and DTG) were carried out under N2 atmosphere with a heating rate of 10 °C min-1 using a Shimadzu DT-50 thermal analyzer.

Preparation of the ligands and their complexes

The unsymmetrical Schiff bases were synthesized according to our previously published method by mono-condensation of the appropriate diamines with aldehydes [25, 26].

2-[(pyridin-2-ylmethylidene)amino]-6-aminopyridine (L1). A solution of 2,6-diaminopyridine (0.545 g, 5 mmol) in 20 mL of ethanol was added dropwise to 2-pyridinecarboxaldehyde (0.535 g, 5 mmol) in 20 mL ethanol slowly with constant stirring. The mixture was stirred for 1 h and then refluxed for 3 h. The resulting solution was evaporated under vacuum. The precipitate was washed with ethanol and then dried in vacuum.

2-[(2-furylmethylene)]phenylenediamine (L2). A solution of O-phenylenediamine (0.541 g, 5 mmol) in 20 mL methanol was added to a methanolic solution (30 mL) of 2-furancarboxaldehyde (0.480 g, 5 mmol). A few drops of acetic acid were added to the reaction mixture and refluxed for 4 h. The mixture was concentrated by evaporation of the solvent. The product was filtered, washed with methanol and then dried in a vacuum over anhydrous calcium chloride.

Preparation of complex (1) [MnL1Cl2(OH2)2]·2H2O was prepared by the addition of a solution of L1 (0.198 g, 1.00 mmol) in 10 mL THF to a solution of MnCl2·4H2O (0.198 g, 1.0 mmol) in methanol (10 mL). A mixture was left under reflux for 3 h. The product was filtered off, washed with hot petroleum ether and finally dried under vacuum.

Preparation of complex (2) [PdL1Cl2] was prepared by dissolving PdCl2 (0.177 g, 1.00 mmol) in ethanol (10 mL) in the presence of a few drops of concentrated HCl. An ethanolic solution of L1 (0.198 g, 1.00 mmol) was added dropwise to the solution of metals. The resulting mixture was allowed to reflux on a water bath for 2 h. The solid products were filtered, washing with ethanol and dried in vacuum desiccators.

Preparation of complex (3) [MnL2Cl2(OH2)2]·2H2O was prepared following the same procedure as described for [MnL1Cl2(OH2)2]·2H2O, starting with L2 (0.186 g, 1.00 mmol) and MnCl2·4H2O (0.198 g, 1.0 mmol).

Preparation of complex (4) [PdL2Cl2] was prepared following the same procedure as described for [PdL1Cl2], starting with L2 (0.186 g, 1.00 mmol) and PdCl2 (0.177 g, 1.00 mmol).

Catalytic activity studies

The metal complex (0.01 mmol) was mixed with 5 mL of H2O2 (0.02 M) in a flask and thermostated to 35 and 55 °C with constant stirring for 1 h. Then, the extent of undecomposed hydrogen peroxide titrating with 0.01 M KMnO4 in the presence of 0.01 M H2SO4. The difference in titer values of the KMnO4 solution before and after the catalyzed decomposition was recorded.

Antibacterial assay

The assay was carried out on the ligands and their metal complexes against two gram-positive bacteria, B. Subtilis and S. aureus, and also against two gram-negative bacteria, E.coli and P. aeruginosa microorganisms. The antibacterial susceptibility test was carried out using the agar well diffusion technique [27]. The stock solution (10 mg mL-1) of the test chemical was prepared in DMSO solution. The test chemicals were added to sterile blank antimicrobial susceptibility disks. The bacteria were subcultured in agar medium. The Petri dishes were incubated for 24 h at 37 °C. The standard antibacterial drug (tetracycline) was also screened under similar conditions for comparison. Activity was determined by measuring the zones of growth inhibition surrounding the disks. Growth inhibition was compared with the standard drug. The activity results were calculated as a mean of triplicates.

DFT study

All calculations were performed with GAUSSIAN09 program package 25. The gas-phase optimized structures of the studied compounds was obtained from density functional B3LYP levels of theory using 6-31G(d) basis set for ligands and LanL2DZ basis set for complexes. The vibrational frequencies were also calculated at the same level [28].

Results and discussion

The present study is based upon the synthesized complexes of metal salts MnCl2·4H2O and PdCl2 with the two Schiff bases L1 and L2. The spectroscopic and magnetic data together with conductance measurements enable us to predict the possible structure of the metal complexes, which are supported also by the thermal decomposition studies. These complexes are stable in air and soluble in DMF and DMSO. The elemental analysis and physical data of the ligands and their complexes are presented in Table 1. The molar conductance values for the complexes in DMSO at a concentration of 10−3 M at room temperature are in the range of 2.3–4.1 Ω−1 mol−1 cm2. The very low values indicate that these complexes behave as non-electrolytes and are neutral in nature. Mass spectrometry is an effective method for investigation of molar weight and structure of the ligands and complexes. The mass spectra of compounds exhibited a peak at a m/z value corresponding to their molecular weight (Fig. 1S).

Table 1 Analytical and physical data of the Schiff bases and their metal complexes
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IR spectra and mode of bonding

The IR spectra of the ligands L1 and L2 showed two bands at 3333, 3280 and 3315, 3282 cm−1 due to νas(NH2) and νs(NH2), respectively [29]. The IR spectra of the Mn(II) complexes (1) and (3) that exhibited a very strong broadband around 3336–3282 cm−1 may be assigned to the stretching vibrations of ν(NH2) group or stretching vibrations of ν(OH) of water molecules associated with the complex. In addition, the IR spectrum of L1 displayed two bands at 1640 and 1588 cm−1 due to ν(C=N) of azomethine and pyridyl moieties, respectively. The IR spectra of [MnL1Cl2(OH2)2]·2H2O (1) and [PdL1Cl2] (2) showed shifts in stretching vibrations of ν(C=N) for azomethine group to 1657–1627 cm−1 and for pyridyl group to 1607–1591 cm−1, suggesting the coordination of metal ion to azomethine and pyridyl nitrogen atoms [30]. The in-plane ring deformation band of py in free ligand L1 (619 cm−1) was shifted to higher frequencies (636–661 cm−1) confirming the coordination of pyridyl nitrogen to metal [31]. The ν(C=N) stretching vibration at 1603 cm−1 in the free ligand L2 was shifted in the (3) and (4) complexes to 1618 and 1609 cm−1, respectively, indicating the participation of azomethine nitrogen in the coordination to metal ion [32]. A medium intensity band due to ν(C–O–C) stretching vibration of furan appeared at 1240 cm−1 in the ligand L2 [5]. This band shifted in [MnL2Cl2(OH2)2]·2H2O (3) and [PdL2Cl2] (4) complexes to 1246 and 1290 cm−1, respectively, suggesting a coordination through oxygen of furan moiety. Appearance of new bands in the spectra of complexes in the regions of 557–454 cm−1 and 450–415 cm−1 was assigned to ν(M–O) and ν(M–N) stretching vibration [8].

1H NMR spectra

The 1H NMR spectra of the free ligands L1 and L2 showed signals at 8.05 and 8.32 ppm, attributed to the azomethine (–CH=N–) protons, respectively. On complexation, the position of azomethine signal was shifted to up-field region 7.69–8.22 ppm in comparison with that of the free ligands, inferring coordination through the azomethine nitrogen atom of the ligand [33]. The 1H NMR spectra of palladium complexes displayed a broad signal due to NH proton without shift with respect to that of the ligands confirming the non-bonding of NH2 nitrogen to metal. In addition, the multiple signals appeared in the 1H NMR spectra of ligands (L1 and L2), which assigned to aromatic, pyridyl and furan protons. These multiplets were displaced upper field in the 1H NMR spectra of Pd(II) complexes indicating the involvement of nitrogen of pyridine ring and oxygen of furan ring in the coordination. The chemical shifts of the different types of protons in the ligands and their Pd(II) complexes are listed in Table 2.

Table 2 1H NMR spectral data of the Schiff bases and their Pd(II) complexes
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Electronic spectra and magnetic moment studies

Electronic spectra of the Schiff bases and their complexes were measured in DMSO (~1 × 10−4 M) in the range of 200–800 nm (Fig. 2). The UV–Vis absorption spectrum of the L1 ligand in DMSO showed a strong band at 386 nm with a shoulder at 466 nm due to n–π* transitions of the non-bonding electrons present in the nitrogen of the azomethine group in the Schiff base. The electronic spectrum of the Mn(II) complex (1) consisted in four bands with the corresponding maxima at 272, 346, 448 and 590 nm due to π–π*, n–π*, 6A1g → 4T2g (G) and 6A1g → 4T1g (G), respectively [34]. The electronic spectrum of Pd(II) complex (2) showed a band at 340 nm with a shoulder at 450 nm. These transitions may be assigned to 1A1g → 1E1g and 1A1g → 1B1g transitions, respectively [35]. UV–Vis spectrum of L2 shows two bands at 380 and 458 nm due to n–π* transition. The electronic spectrum of the Mn(II) complex (3) showed four bands at 278, 308, 392 and 464 nm. The first two bands are due to π–π* and n–π* transitions, respectively. The other two bands are assigned to the transitions: 6A1g → 4T2g (G) and 6A1g → 4T1g (G), respectively, for a Mn(II) ion in a distorted octahedral field [36, 37]. For palladium complex 4, two absorption bands at 364 and 448 nm are observed in the spectrum due to d–d transitions [38].

Fig. 2
figure 2

Electronic absorption spectra of ligands and their complexes in 10−4 M DMSO

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The diamagnetic nature of the palladium complexes is in accordance with square planar geometry for the Pd(II) complexes [39]. The magnetic moment values for Mn(II) complexes are in the range of 5.85–5.90 B.M. The magnetic moment lies within the region expected for octahedral complexes [40].

Fluorescence spectra studies

Fluorescence spectra of the Schiff base ligands and their complexes were measured in DMSO solution at room temperature (Fig. 3). The emission spectrum of L1 shows band at 435 nm upon excitation at 386 nm. A broad fluorescence band centered at 417 nm observed in the emission spectrum of Mn(II) complex (1) when excited at 346 nm attributed to π–π* transitions [41]. The emission spectrum of Pd(II) complex (2) shows maxima at 394 nm on excitation at 386 nm. The fluorescence spectrum of the free L2 when excited at 380 nm showed a maxima band at 395 nm. The emission intensity of the Mn(II) complex (3) is stronger than that of the ligand L2 and showed a maximum at 346 nm when excited at 308 nm. Complex (4) shows very weak fluorescence at 432 nm when excited at 364 nm. Enhancement of fluorescence through complexation is, however, of much interest as it opens up the opportunity for photochemical applications of these complexes [42].

Fig. 3
figure 3

Emission spectra of ligands and their complexes in 10−4 M DMSO

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Thermal analysis (TG and DTG)

The decomposition stages, temperature ranges, decomposition product as well as mass loss percentages of the complexes are listed in Table 3.

Table 3 Thermogravimetric data of complexes
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[MnL1Cl2(OH2)2]·2H2O

The TG curve of Mn(II) complex (1) shows the four steps of decomposition within the temperature range from 318 to 994 K (Fig. 2S). The first two decomposition stages occur in the temperature range of 318–519 K which is due to the loss of four molecules of water and NH3 gas with an estimated mass loss of 22.50% (calcd. 22.48%). In the second decomposition step, complex (1) shows a mass loss of 17.42% (calcd. 17.89%) in the temperature range of 519–730 K, which corresponds to the removal of Cl2 as gas. The final step of decomposition corresponds to the loss of HCN, 2CN, 3HC≡CH and 2C with the mass loss of 45.90% (calcd. 45.74%) leaving metal as a residue.

$$\begin{array}{*{20}l} {[{\text{Mn}}({\text{C}}_{11} {\text{H}}_{10} {\text{N}}_{4} )({\text{OH}}_{2} )_{2} {\text{Cl}}_{2} ] \cdot 2{\text{H}}_{2} {\text{O}}} \hfill & {\mathop{\longrightarrow}\limits_{{{-}4{\text{H}}_{2} {\text{O}}\;{\text{and}}\;{\text{NH}}_{ 3} }}^{{{318{-}519\;{\text{K}}}}}} \hfill & { [ {\text{Mn}}({\text{C}}_{11} {\text{H}}_{7} {\text{N}}_{3} ){\text{Cl}}_{2} ]} \hfill & {(1)} \hfill \\ { [ {\text{Mn}}({\text{C}}_{11} {\text{H}}_{7} {\text{N}}_{3} ){\text{Cl}}_{2} ]} \hfill & {\mathop{\longrightarrow}\limits_{{{-}{\text{Cl}}_{2} }}^{{{519{-}730\,{\text{K}}}}}} \hfill & {[{\text{Mn(C}}_{ 1 1} {\text{H}}_{ 7} {\text{N}}_{ 3} )]} \hfill & {(2)} \hfill \\ { [ {\text{Mn}}({\text{C}}_{11} {\text{H}}_{7} {\text{N}}_{3} )]} \hfill & {\mathop{\longrightarrow}\limits_{{{-}{\text{HCN,2CN,3HC}} {\equiv} {\text{CH}}}}^{{{730{-}994\,{\text{K}}}}}} \hfill & {\text{Mn}} \hfill & {(3)} \hfill \\ \end{array}$$

[PdL1Cl2]

The thermal decomposition of the Pd(II) complex (2) undergoes in two steps. The first step between 320 and 617 K corresponds to the loss of chlorine molecule. The observed mass loss is 18.41%, which is consistent with the theoretical value of 18.88%. The second stage, which occurs in the temperature range of 617–923 K, corresponds to the decomposition of the organic part of ligand. The observed mass loss for this stage is 52.96% (calcd. 52.79%). The final residue was estimated as palladium metal.

$$\begin{array}{*{20}l} {[{\text{Pd(C}}_{ 1 1} {\text{H}}_{ 1 0} {\text{N}}_{ 4} ) {\text{Cl}}_{ 2} ]} \hfill & {\mathop{\longrightarrow}\limits_{{{-}{\text{Cl}}_{2} }}^{{{320{-}617\,{\text{K}}}}}} \hfill & {[{\text{Pd(C}}_{ 1 1} {\text{H}}_{ 1 0} {\text{N}}_{ 4} )]} \hfill & {(1)} \hfill \\ {[{\text{Pd(C}}_{ 1 1} {\text{H}}_{ 1 0} {\text{N}}_{ 4} )]} \hfill & {\mathop{\longrightarrow}\limits_{{{-}{\text{C}}_{11} {\text{H}}_{10} {\text{N}}_{4} }}^{{{617{-}923\,{\text{K}}}}}} \hfill & {\text{Pd}} \hfill & {(2)} \hfill \\ \end{array}$$

[MnL2Cl2(OH2)2]·2H2O

The curve of Mn(II) complex (3) was decomposed into four steps. The first two decomposition stages begin at 318 K and ends at 643 K. The observed mass loss (29.91%) is attributed to the loss of 4H2O, NH3 and HCN (calcd. 30.21%). The third step relates to the elimination of half chlorine in the temperature range of 643–756 K by a mass loss of 8.96% (calcd. 9.20%). The final step of the decomposition in the temperature range of 756–1272 K is due to removal of ½Cl2, 3HC≡CH and 4C (41.71%) leaving MnO as a final residue (calcd. 42.07%).

$$\begin{array}{*{20}l} {[{\text{Mn(C}}_{ 1 1} {\text{H}}_{ 1 0} {\text{N}}_{ 2} {\text{O)(OH}}_{ 2} )_{ 2} {\text{Cl}}_{2} ] \cdot 2{\text{H}}_{2} {\text{O}}} \hfill & {\mathop{\longrightarrow}\limits_{{{-}4{\text{H}}_{2} {\text{O,NH}}_{ 3} \;{\text{and}}\;{\text{HCN}}}}^{{{318{-}643\,{\text{K}}}}}} \hfill & {[{\text{Mn(C}}_{ 1 0} {\text{H}}_{ 6} {\text{O)Cl}}_{ 2} ]} \hfill & {(1)} \hfill \\ { [ {\text{Mn(C}}_{ 1 0} {\text{H}}_{ 6} {\text{O)Cl}}_{ 2} ]} \hfill & {\mathop{\longrightarrow}\limits_{{{-}1/2{\text{Cl}}_{ 2} }}^{{{643{-}756\,{\text{K}}}}}} \hfill & {[{\text{Mn(C}}_{ 1 0} {\text{H}}_{ 6} {\text{O)Cl}}]} \hfill & {(2)} \hfill \\ { [ {\text{Mn(C}}_{ 1 0} {\text{H}}_{ 6} {\text{O)Cl]}}} \hfill & {\mathop{\longrightarrow}\limits_{{{-}1/2{\text{Cl}}_{2} ,3{\text{HC}} {\equiv} {\text{CH}}\;{\text{and}}\; 4 {\text{C}}}}^{{{756{-}1272\;{\text{K}}}}}} \hfill & {\text{MnO}} \hfill & {(3)} \hfill \\ \end{array}$$

[PdL2Cl2]

The TG curve of Pd(II) complex (4) shows three stages of decomposition within the temperature range of 320–1123 K. The first stage of decomposition occurred in the temperature range of 320–653 K with a net mass loss of 14.10% (calcd. 14.43%), which is consistent with the elimination of ½Cl2 and NH3 gases (Fig. 3S). The second step involved removal of ½Cl2, HCN and HC≡CH with a mass loss of 24.73% (calcd. 24.34%). The third step (923–1123 K) corresponds to loss of 2HC≡CH and 4C to give final PdO as a residue.

$$\begin{array}{*{20}l} {[{\text{Pd(C}}_{ 1 1} {\text{H}}_{ 1 0} {\text{N}}_{ 2} {\text{O)Cl}}_{2} ]} \hfill & {\mathop{\longrightarrow}\limits_{{{-}1/2{\text{Cl}}_{2} \;{\text{and}}\;{\text{NH}}_{ 3} }}^{{{320{-}653\,{\text{K}}}}}} \hfill & {[{\text{Pd(C}}_{ 1 1} {\text{H}}_{7} {\text{NO)Cl}}]} \hfill & {(1)} \hfill \\ { [ {\text{Pd(C}}_{ 1 1} {\text{H}}_{ 7} {\text{NO)Cl}}]} \hfill & {\mathop{\longrightarrow}\limits_{{{-}1/2{\text{Cl}}_{ 2} ,{\text{HCN}}\;{\text{and}}\;{\text{HC}} {\equiv} {\text{CH}}}}^{{{653{-}923\,{\text{K}}}}}} \hfill & {[{\text{Pd(C}}_{8} {\text{H}}_{4} {\text{O)}}]} \hfill & {(2)} \hfill \\ { [ {\text{Pd(C}}_{ 8} {\text{H}}_{ 4} {\text{O)]}}} \hfill & {\mathop{\longrightarrow}\limits_{{{-}2{\text{HC}} {\equiv} {\text{CH}}\;{\text{and}}\; 4 {\text{C}}}}^{{{923{-}1123\;{\text{K}}}}}} \hfill & {\text{PdO}} \hfill & {(3)} \hfill \\ \end{array}$$

Catalytic activity studies

The decomposition of H2O2 catalyzed by complexes can be monitored by titrating the undecomposed H2O2 with a standard KMnO4 solution. Variables such as the temperature were found to have important roles in the decomposition of H2O2 [43]. The Pd(II) and Mn(II) complexes were used for these investigations. The results showed that the percentage of hydrogen peroxide decomposition increased with the temperature (Fig. 4S). Pd(II) complexes are more active than Mn(II) complexes in the decomposition of hydrogen peroxide. This is due to the Pd(II) complexes have square planar structure. This provides vacant coordination sites and facilitates the formation of peroxo intermediates which are highly reactive species, leading to the rapid decomposition of hydrogen peroxide [44].

Antibacterial studies

The antibacterial studies were carried out on the ligands and their metal complexes (10 mg mL-1) in DMSO solution against two gram-positive bacteria, B. Subtilis and S. aureus, and also against two gram-negative bacteria; E.coli and P. aeruginosa microorganisms. The experimental data that indicate the inhibition zone diameter (mm) of the assayed compounds are listed in Table 1S. The results of the antibacterial activity of the compounds compared with the standard drug (Tetracycline) have indicated that compounds were active, but activity was lesser than the standard drug (Fig. 4). Comparative studies of the ligands and their metal complexes indicated that metal complexes exhibit higher antibacterial activities than the free ligands.

Fig. 4
figure 4

Antibacterial activity of ligands and their complexes

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DFT calculations

The optimized structures of the two ligands and the corresponding four complexes are shown in Fig. 5 with the labeling atoms. Their bond lengths and bond angles are presented in Table 4. The bond distances in (1) and (3) between manganese and both of chloro and aqua ligands are in the range (2.28–2.33 Å) and (2.01–1.96 Å), respectively. These values are in reported range [45]. The Pd–Cl, Pd–O and Pd–N distances in complexes (2) and (4) are found to be in the reported range [46]. From the geometric parameters, Mn atoms have distorted octahedral structures, while the complexes with Pd atoms have distorted square planar structures.

Fig. 5
figure 5

Optimized geometry of the ligands and their complexes

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Table 4 Some selected bond lengths (Å) and bond angles (°) of the optimized geometry of ligands and their complexes by using DFT/B3LYP levels of theory
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Some selected calculated parameters, namely total energy and dipole moments, are available from the calculations (Table 5). The dipole moment increases spectacularly from ligands to complexes, where the electropositive metal ions induce strong polarization in the σ- and π-frameworks of the ligand molecule. In addition, the energies of the complexes (1–4) become more negative indicating the high stability of these complexes with respect to the free ligands [46].

Table 5 Total energies (a.u.) and dipole moments (Debye) of ligands and their complexes by using DFT/B3LYP levels of theory
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The calculated vibrational frequency values of the N–H, C=N, C–O–C, M–O and M–N stretching modes for the optimized complexes are tabulated in Table 6. The calculated C=N stretching frequencies for L1, L2 and their Mn and Pd complexes were 1646, 1603, 1621, 1620,1644, 1615 cm−1, respectively. Corresponding experimental frequency values are 1640, 1603, 1627, 1657, 1618 and 1609 cm−1, respectively. Also, the calculated stretching frequency values of the N–H, C–O–C, M–O and M–N modes are close to corresponding experimental IR bands. The observed differences are most probably because the calculations were performed in the gas phase, whereas the experimental results were obtained in the solid phase.

Table 6 Selected calculated IR vibration frequencies (cm−1) of the ligands and their complexes
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Frontier molecular orbital energies

Frontier molecular orbitals, i.e., the highest occupied molecular orbital (HOMO) and lowest unoccupied molecular orbital (LUMO) surfaces and energies for ligands and their complexes, are given in Figs. 6 and 7. Frontier molecular orbitals play an important role in the determination of chemical reactivity. HOMOs of ligands are delocalized over the whole ligand, while LUMOs electron density is not located on amino group. All LUMO of complexes is more located on azomethine, furan and pyridyl rings than HOMO of complexes. The energy of HOMO of L1 (−7.75 eV) and L2 (−5.43 eV) is the lowest for all compounds. The energy separation (ΔE) between the HOMO and LUMO for L1 (1) and (2) is 3.93, 3.46 and 1.03 eV, respectively. Also, the energy gap for L2, (3) and (4) is 6.12, 3.90 and 2.60 eV, respectively. An electronic system with a larger HOMO–LUMO gap should be less reactive than having a smaller gap [47, 48]. The energy gaps between HOMO and LUMO (ΔE) of ligands are greater than those for manganese and palladium complexes. Hence, Mn(II) and Pd(II) complexes have more chemical activity than ligands.

Fig. 6
figure 6

Energy-level diagram of the frontier molecular orbitals for a L1 and its complexes: b (1) and c (2)

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Fig. 7
figure 7

Energy-level diagram of the frontier molecular orbitals for a L2 and its complexes: b (3) and c (4)

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Conclusions

The novel complexes of manganese(II) and palladium(II) with the bidentate ligands 2-[(pyridin-2-ylmethylidene)amino]-6-aminopyridine (L1) and 2-[(2-furylmethylene)]phenylenediamine (L2) were synthesized and characterized using spectroscopic methods. The results of the elemental analysis, mass spectra and thermogravimetric analysis deduced the formation of 1:1 Schiff base/metal ions complexes. The optimized geometries and IR vibrational frequencies of the ligands and their complexes were calculated by the DFT method. Experimental and calculated wavenumbers were found to be in a good agreement. The complexes are biologically active. Fluorescence studies showed that the reported compounds can potentially serve as photoactive materials.