Benzoic hydroxamate-based iron complexes as model compounds for humic substances: Synthesis, characterization and algal growth experiments

Article abstract of DOI:10.1039/c5ra25256c

A series of monomeric and dimeric Fe^III complexes bearing benzoic hydroxamates as O,O-chelates has been prepared and characterized by elemental analysis, IR spectroscopy, UV-Vis spectroscopy, electrospray ionization mass spectrometry (ESI-MS), cyclic voltammetry, EPR spectroscopy and for some examples by X-ray diffraction analysis. The stability of the synthesized complexes in pure water and seawater was monitored over 24 h by means of UV-Vis spectrometry. The ability to release iron from the synthesized model complexes has been investigated with algae growth experiments.

Full text of DOI:10.1039/c5ra25256c

RSC Advances
View Article Online
View Journal | View Issue
PAPER
Benzoic hydroxamate-based iron complexes as
model compounds for humic substances: synthesis,
characterization and algal growth experiments†
Cite this: RSC Adv., 2016, 6, 40238
Ewelina Orlowska,a Alexander Roller,a Hubert Wiesinger,a Marc Pignitter,b
a
a
Franz Jirsa,ac Regina Krachler, Wolfgang Kandioller and Bernhard K. Kepplera
*
A series of monomeric and dimeric FeIII complexes bearing benzoic hydroxamates as O,O-chelates has
been prepared and characterized by elemental analysis, IR spectroscopy, UV-Vis spectroscopy,
electrospray ionization mass spectrometry (ESI-MS), cyclic voltammetry, EPR spectroscopy and for some
examples by X-ray diffraction analysis. The stability of the synthesized complexes in pure water and
seawater was monitored over 24 h by means of UV-Vis spectrometry. The ability to release iron from the
synthesized model complexes has been investigated with algae growth experiments.
Received 27th November 2015
Accepted 29th March 2016
DOI: 10.1039/c5ra25256c
www.rsc.org/advances
ferric hydroxide aggregates.21 In fact, 99.97% of dissolved iron
in seawater exists as complexes with organic chelators.22–25
Introduction
However, the bioavailability of iron does not solely depend on
its solubility moreover other parameters such as complex
hydrolysis and photo- or bioreductions are essential.26 The
accumulated complexed Fe(III) can be reduced by bioreductants
or more importantly photochemical reactions. The formed Fe(II)
dissociates and is prone to oxidation reactions with O2 or H2O2
(Fenton reaction). This reaction cycle increases the steady-state
concentration of dissolved bioavailable iron.27
Iron has been recognized as an essential trace element for all
living organisms, ranging from archaea to mammals and is
required for the function of many proteins and enzymes. Also
photosynthesis as well as nitrogen assimilation is dependent on
iron-based enzyme cofactors.1–3 More than 50% of the photo-
synthetic activity on earth is attributed to phytoplankton, which
consists of photoautotrophic microscopic organisms.3–5 They
are responsible for a major part of the global biosynthesis of
organic compounds from inorganic carbon dioxide dissolved in
water and therefore play a key role in the global CO2 cycle.
Although iron is a common element on earth, large areas of the
ocean show very low iron concentrations.6–8 In so called high-
nutrient, low-chlorophyll (HNLC) areas nutrients like nitrate
and phosphate are present in high concentrations; however,
chlorophyll concentrations, as indicator for phytoplankton
density are low.9,10 Evidences have shown that insufficient
concentrations of iron in these areas are responsible for the
inhibition of growth, development and productivity of marine
phytoplankton.11–13 The presence of bioavailable iron in
seawater is crucial for a high productivity of phytoplankton.14–20
The observed low iron concentrations can be explained by the
chemical speciation of iron in seawater. Thermodynamically
stable iron(III) precipitates in oxygenated seawater at pH 8 as
Until now it was believed that the main sources of iron in the
ocean are atmospheric deposition of volcanic dust, terrestrial
sand and extraterrestrial dust, the upwelling of iron rich waters
as well as hydrothermal vents and ice melting;26,28 however, the
role of river water has been mainly neglected. Upon mixing of
river water with seawater in the brackish water zone of estuaries
the majority of the transported iron gets sequestered into the
sediment due to the major change in ionic strength.29–34 Recent
investigations in the Mississippi Delta and North Scotland have
shown that a signicant proportion of iron(III) complexes of
aquatic humic substances (AHS) in the river remain in colloidal
solution in sea water and can contribute to iron-fertilization of
coastal waters.35–37 AHS, resistant against salt-induced occu-
lation, have been found in the entire Atlantic, Arctic and Pacic
oceans. The identication of AHS as important carriers for iron
explains the increased bioavailability of iron in natural seawater
compared to articial one.38,39 AHS have been under investiga-
tion for a very long time, and their composition has not been
claried in detail. The complex structure of this substance class
origins from the incomplete decomposition of organic material,
such as lignins which undergo chemical, physical
and microbiological transformations.40,41 Furthermore other
organic molecules like amino acids, sugars and peptides can be
incorporated leading to a broad range of different structural
aInstitute of Inorganic Chemistry, Faculty of Chemistry, University of Vienna,
Waehringer Str. 42, A-1090 Vienna, Austria. E-mail: wolfgang.kandioller@univie.ac.at
bDepartment of Nutritional and Physiological Chemistry, Faculty of Chemistry
University of Vienna, Althanstr. 14/UZA II, A-1090 Vienna, Austria
cDepartment of Zoology, University of Johannesburg, Auckland Park, 2006 South Africa
† Electronic supplementary information (ESI) available. CCDC 1436729–1436734.
For ESI and crystallographic data in CIF or other electronic format see DOI:
10.1039/c5ra25256c
40238 | RSC Adv., 2016, 6, 40238–40249
This journal is © The Royal Society of Chemistry 2016

View Article Online
Paper
RSC Advances
features.42,43 Unfortunately, the very low concentration of those spectrophotometer and Perkin Elmer lambda 35 with PTP
natural iron chelators in seawater makes the detection and 6 (Peltier Temperature Programmer) and Julabo AWC 100
characterization of them challenging.44 Being aware of the recirculating cooler in the range of 200–800 nm in both distilled
importance of those natural macromolecules for marine water and seawater at 25 ꢀC. Electrospray ionization mass
microorganisms it is of large interest to nd synthetic spectra were measured with a Bruker maXis ESI-QqTOF spec-
compounds that are comparable to AHS, regarding their iron trometer in the positive and negative mode using ACN/MeOH
binding and transport properties in sea water, as well as their 1% H2O as solvent. Elemental analyses were performed at the
capability of providing bioavailable iron to algae. Comparing Microanalytical Laboratory of the University of Vienna with
the latest EXAFS studies45 and the current information about a Perkin-Elmer 2400 CHN Series II elemental analyzer or
AHS we put further requirements for the model compounds a Eurovector EA3000 elemental analyzer and are within 0.4% of
together. Overall they include mononuclear Fe(III) complexes, the calculated values. ATR-IR spectra were measured using
mainly O,O coordination, functional groups such as phenols a Bruker Vertex 70 Fourier transform IR spectrometer. Cyclic
and aromatic carboxylic groups and good water solubility. Iron voltammograms were measured in a three-electrode cell using
shows very strong interaction with AHS even at low pH values, a 2.0 mm- and 3.0 mm-diameter glassy carbon working elec-
which is similar to iron complexes bearing hydroxamic acid trode, a platinum auxiliary electrode, and an Ag|Ag+ reference
derivatives.46 Thus ligands with hydroxamate functional groups electrode containing 0.1 M AgNO3. Measurements were per-
seemed to be suitable models for humic substances as proposed formed at room temperature using an EG & G PARC 273A
by Frimmel et al.47,48 The strong chelating properties of potentiostat/galvanostat. Deaeration of solutions was accom-
hydroxamates are exploited by many microorganisms like plished by passing a stream of argon through the solution for 5
bacteria which produce and release low molecular siderophores min prior to the measurement and then maintaining a blanket
to bind and transport iron. Signicant nitrogen content49–51 and atmosphere of argon over the solution during the measure-
possible presence of nitrogen in the rst and second coordi- ment. The potentials were measured in DMF containing 0.10 M
nation sphere of iron in AHS supports our decision to use them [n-Bu4N][BF4], using [Fe(h5-C5H5)2] (E1/2 ¼ +0.6–0.68 V vs. NHE)
as models.
as internal standard and are quoted relative to the normal
Herein we report the synthesis of benzhydroxamic acids and hydrogen electrode NHE.52 Electron paramagnetic resonance
their respective iron(III) coordination compounds. The synthe- (EPR) spectra were acquired on a Bruker Elexsys-II E500 CW-
sized complexes were investigated and characterized by stan- EPR spectrometer operating at X-band with 100 kHz modula-
dard analytical methods such as elemental analysis, UV-Vis, IR, tion frequency. All samples were dissolved in DMF and trans-
EPR, cyclic voltammetry, mass spectrometry and X-ray diffrac- ferred to a 3 mm inner diameter quartz EPR tube, which was
tion. Moreover, the stability of the iron complexes in (sea)water placed into a high-sensitivity cavity (SHQE1119). The spectra of
was investigated. Algal growth experiments were performed all samples and the solvent alone were recorded at 90 ꢁ 1 K with
with the marine single-cell chlorophyte Chlorella salina and a variable nitrogen temperature control system using the
haptophyte Diacronema lutheri to elucidate the potential of following settings: center eld: 6000 G; sweep width: 12 000 G;
these iron complexes to act as model compounds for humic sweep time: 335.5 s; modulation amplitude: 20.37 G; microwave
substances in regards of their biological function in sea water power: 15 mW; conversion time: 81.92 ms; resolution: 4096
environment.
points; averaged scans: 3. The rhombic symmetry of the S ¼ 5/2
spin system was determined by the experimentally observed
effective g-values using visual RHOMBO.53 X-ray diffraction
measurements were performed on a D8 Venture at 100 K
experimental parameters of the X-ray diffraction measurements
Experimental
Materials and methods
All chemicals were purchased from Sigma Aldrich (salicylic are listed in Table S1–S13.† Distortion parameters for all
acid, 4-hydroxybenzoic acid, 3,4,5-trimethoxy benzoic acid, 3,4- diffraction measurements are listed in Table S14.† SAINT54 was
dihydroxyhydrocaffeic acid, o-coumaric acid, iron(III) chloride used to process the data. The structures were solved by direct
hexahydrate, potassium hydroxide, hydroxylamine hydrochlo- methods and rened by full-matrix least-squares techniques.
ride, palladium on carbon 10 wt% loading), Alfa (2,4-dihy- Non-hydrogen atoms were rened with anisotropic displace-
droxybenzoic
acid,
2,3-dihydroxybenzoic
acid,
3,5- ment parameters. Hydrogen atoms were inserted at calculated
dihydroxybenzoic acid), Fluka (N-methylmorpholine, ethyl positions and rened with a riding model or as free rotating
chloroformate, 3-hydroxybenzoic acid), Riedl de Haen (iron(III) groups. The following computer programs were used: structure
chloride anhydrous) or Acros (vanilic acid, cinnamic acid, solution, SHELXS-97;55 renement, SHELXL-2013 (ref. 55)
syringic acid, benzyl bromide, sodium bicarbonate) and used OLEX2;56 SHELXLE;57 molecular diagrams. The crystallographic
without further purication. All solvents were of analytical data les of the complexes have been deposited at the Cam-
grade and used without further purication. If not otherwise bridge Crystallographic Database as CCDC 1436734 (6a),
stated the substances were synthesized and puried according 1436733 (6b), 1436732 (6d), 1436731 (6e), 1436730 (6g), 1436729
to the general procedure. 1H NMR spectra were recorded on (6j), respectively.†
a Bruker Avance III™ 500 MHz spectrometer in DMSO-d6 at 298
K using standard pulse programs at 500.10 MHz for H exper- cultures of the unicellular chlorophyte species C. salina, strain
iments. UV-vis experiments were performed on an Agilent 8453 SAG 8.86 and haptophyte D. lutheri, strain SAG 926-1 obtained
Algal growth experiments were carried out with batch
1
This journal is © The Royal Society of Chemistry 2016
RSC Adv., 2016, 6, 40238–40249 | 40239

View Article Online
RSC Advances
Paper
from the Culture Collection of Algae at Goettingen University.
C
30H36FeN3O15: C, 49.06; H, 4.94; N, 5.72; O, 32.67; found: C,
These algae species were chosen because of their widespread 48.78; H, 4.60; N, 5.60; O, 32.56. IR (ATR, selected bands, nmax):
occurrence and abundance in the Northern Atlantic Ocean. 3220, 1567, 1518, 1489, 1414, 1362, 1298, 1242, 1124, 1077,
Experiments were performed in modied sterile f/2 medium,58 1030, 978, 858, 828, 771, 719 cmꢂ1
.
containing EDTA as complexing agent, prepared with 35&
[Fe(cinammic hydroxamate)3] (6b). FeCl3 (87 mg, 0.53 mmol)
salinity articial seawater as described by Kester59 at pH 8.2. dissolved in methanol (2 mL) was added to the suspension of 4b
Cultures were grown in 200 mL Schott asks kept at 21 ꢁ 0.5 ꢀC (272 mg, 1.64 mmol) and KOH (89.4 mg, 1.6 mmol) in methanol
by means of a water bath, stirred with 300 rpm and supplied (10 mL). The dark red solution was stirred for 2 h, ltered and
with ltered air. Plant grow-uorescent lamps with a 16 : 8 h the solvent was removed in vacuo. The residue was extracted
light : dark cycle were used to provide algae with light at mean with acetone (10 mL) and le in refrigerator overnight. The
intensities, directly measured at the ask surface, of 165 mmol suspension was ltered and diethyl ether was slowly added
mꢂ2 sꢂ1. All cultures were carried out in triplicates; for each yielding to the formation of a red precipitate. The product was
approach three different control samples were prepared: full f/2 ltered off and dried in vacuo. Yield: 181 mg, 0.32 mmol, 61%.
medium (+Fe, +EDTA), f/2 medium without iron (ꢂFe, +EDTA) X-ray diffraction quality single crystals were grown in ethanol
and f/2 medium without EDTA (+Fe, ꢂEDTA) (see S15†). As (slow evaporation). ESI-MS: m/z 565 [M + Na]+, 380 [M ꢂ L + H]+;
a negative control for our studies, we utilized iron-free samples anal. calcd for C27H24FeN3O6$H2O: C, 57.87; H, 4.68; N, 7.50;
where we used extra pure sodium chloride for the seawater found: C, 57.85; H, 4.66; N, 7.33. IR (ATR, selected bands, nmax):
preparation to avoid any iron contamination. To test benzoic 2983, 1639, 1514, 1449, 1360, 1249, 1205, 1068, 1012, 968, 852,
hydroxamates-based iron complexes, the respective iron 752, 698, 670, 580 cmꢂ1
concentration was added in form of the respective complex into [Fe(syringic hydroxamate)3] (6c). FeCl3 (96 mg, 0.59 mmol)
.
the f/2 medium and no additional EDTA was used. All the dissolved in methanol (5 mL) was added to the suspension of 5c
nutrient stock solutions were sterilized by passing through a 0.2 (380 mg, 1.78 mmol) and KOH (100 mg, 1.78 mmol) in meth-
mm capsule lter (Sartorius Sartobran 300). Algae were precul- anol (12 mL). The dark red solution was stirred for 2 h, ltered
tured in full medium, at the beginning of the experiment an and the solvent was removed in vacuo. The residue was extracted
inoculum of 2–5 mL was used to obtain an initial concentration with MeOH/CHCl3 1 : 1 (100 mL) and le in refrigerator over-
of app. 9 ꢃ 104 cells per mL. In the rst and second experiment night. The suspension was ltered and diethyl ether was slowly
on C. salina, ve and three model compounds were tested added yielding the formation of a red precipitate. The product
respectively. The algal growth experiment on D. lutheri was was ltered off and then dried in vacuo. Yield: 170 mg, 0.23
carried out with six model compounds. The concentrations of mmol, 37%. ESI-MS: m/z 731 [M + K]+, 715 [M + Na]+, 693 [M +
tested substances were 11.7 mmol Lꢂ1. Moreover in each H]+; anal. calcd for C27H30FeN3O15$0.75CHCl3: C, 42.63; H, 3.96;
experiment one sample with isolated humic acid fraction and N, 5.37; found: C, 42.61; H, 4.22; N, 5.40. IR (ATR, selected
three control samples were (described above) tested. The bands, nmax): 3270, 1576, 1491, 1424, 1353, 1289, 1251, 1223,
experiments were carried out over a period of 24 to 30 days 1117, 1073, 979, 832, 776, 724, 571 cmꢂ1
(depending on algal growth) and the algae concentration was [Fe(salicylic hydroxamate)3] (6d). FeCl3 (127 mg, 0.78 mmol)
.
monitored aer 7 days from beginning of experiment and then dissolved in methanol (3 mL) was added to the suspension of 5d
every other day. The number of cells in the culture was esti- (352 mg, 2.34 mmol) and KOH (128 mg, 2.34 mmol) in meth-
mated with a Neubauer improved cell counting chamber with anol (20 mL). The dark red solution was stirred for 2 h, ltered
a 0.1 mm depth and microscope. Because of the mobility of D. and the solvent was removed in vacuo. The residue was extracted
lutheri, in order to count the cells, 1 mL of each sample was with methanol (20 mL) and stored at 4 ꢀC overnight. The
collected and algae were xed adding 10 mL of 10% formic acid suspension was ltered and the solvent was evaporated to small
solution. C. salina cells were counted without any treatment.
volume. Slow diffusion of diethyl ether into the methanolic
phase resulted in the formation of dark red precipitate. The
product was ltered off and then dried in vacuo. Yield: 199 mg,
0.34 mmol, 43%. X-ray diffraction quality single crystals were
Synthesis of iron complexes
[Fe(eudesmic hydroxamate)3] (6a). FeCl3 (18.5 mg, 0.11 grown in acetone (slow evaporation). ESI-MS: m/z 535 [M + Na]+;
mmol) dissolved in methanol (2 mL) was added to the anal. calcd for C21H18FeN3O9: C, 49.24; H, 3.54; N, 8.20; found:
suspension of eudesmic hydroxamic acid 4a (79 mg, 0.35 mmol) C, 49.40; H, 3.85; N, 8.11. IR (ATR, selected bands, nmax): 2970,
and KOH (20 mg, 0.35 mmol) in methanol (5 mL). The dark red 1603, 1566, 1480, 1385, 1312, 1245, 1154, 1098, 1057, 1034, 919,
solution was stirred for 2 h, ltered and the solvent was 858, 748, 666, 578 cmꢂ1
removed in vacuo. The residue was extracted with dichloro-
.
[Fe(3-hydroxybenzoic hydroxamate)3] (6e). FeCl3 (68 mg, 0.42
methane (20 mL), the solution was ltered and the solvent was mmol) dissolved in methanol (3 mL) was added to the
removed. The obtained residue was dissolved in a small amount suspension of 5e (193 mg, 1.3 mmol) and KOH (71 mg, 1.26
of methanol (2 mL) and stored overnight at 4 ꢀC. The formed red mmol) in methanol (20 mL). The dark red solution was stirred
precipitate was ltered off and then dried in vacuo. Yield: 66 mg, for 2 h, ltered and the solvent was removed in vacuo. The
0.09 mmol, 82%. X-ray diffraction quality single crystals were residue was extracted with acetone (20 mL) and stored in
grown in acetone/water (slow evaporation). ESI-MS: m/z 773 [M + refrigerator overnight. The suspension was ltered, the solvent
K]+, 757 [M + Na]+, 547 [M ꢂ L + K]+, anal. calcd for was evaporated and the residue was dissolved in a small
40240 | RSC Adv., 2016, 6, 40238–40249
This journal is © The Royal Society of Chemistry 2016

View Article Online
Paper
RSC Advances
amount of acetone. Slow diffusion of hexane into the acetone The precipitate was washed with water and dried in vacuo at 50
solution resulted in the formation of dark red precipitate. The ꢀC. Mother liquor was le overnight and second fraction of the
product was ltered off and then dried in vacuo. Yield: 178 mg, product crystallized. Yield: 175 mg, 0.29 mmol, 48%. ESI-MS: m/z
0.29 mmol, 69%. X-ray diffraction quality single crystals were 583 [M + Na]+; elemental analysis: calcd for C21H18FeN3O12$2H2O:
grown in acetone (slow evaporation). ESI-MS: m/z 535 [M + Na]+; C, 42.30; H, 3.72; N, 7.04; O, 37.57; found: C, 41.97; H, 3.56; N, 6.78;
m/z 360 [M ꢂ L]+; anal. calcd for C21H18FeN3O9$H2O$1.5C3H6O: O, 37.19. IR (ATR, selected bands, nmax): 3196, 1565, 1487, 1358,
C, 49.61; H, 4.73; N, 6.81; found: C, 49.34; H, 4.46; N, 6.59. IR 1306, 1210, 1158, 1125, 1063, 1001, 983, 847, 815, 675 cmꢂ1
.
(ATR, selected bands, nmax): 3201, 1699, 1566, 1520, 1481, 1357,
1236, 1129, 1058, 1000, 967, 845, 791, 680 cmꢂ1
[Fe(p-hydrocoumaric hydroxamate)3] (6j). FeCl3$6H2O (168
mg, 0.62 mmol) dissolved in water (2 mL) was added to the
.
[Fe(4-hydroxybenzoic hydroxamate)3] (6f). FeCl3$6H2O (133 solution of 5j (337 mg, 1.87 mmol) and KOH (105 mg, 1.87
mg, 0.49 mmol) dissolved in water (2 mL) was added to the mmol) in water (20 mL) at 50 ꢀC. Dark red precipitate was
solution of 5f (225 mg, 1.48 mmol) and KOH (83 mg, 1.48 mmol) formed and the suspension was stirred at 50 ꢀC for 2 h and
in water (20 mL) at 50 ꢀC. Dark red precipitate was formed and ltered while hot. The precipitate was washed with water and
the suspension was stirred at 50 ꢀC for 2 h and ltered hot. The dried in vacuo at 50 ꢀC. Mother liquor was le overnight and
ꢀ
precipitate was washed with water and dried in vacuo at 50 C. second fraction of the product crystallized. Crystals were
ꢀ
Mother liquor was le overnight at 4 ꢀC and a second fraction of washed with water and dried in vacuo at 50 C. Yield: 130 mg,
the product crystallized. The obtained crystals were washed 0.21 mmol, 43%. ESI-MS: m/z 619 [M + Na]+; anal. calcd for
with water and dried in vacuo at 50 ꢀC. Yield: 150 mg, 0.29
C27H30FeN3O9$0.5H2O: C, 53.57; H, 5.16; N, 6.94; found: C,
mmol, 59%. ESI-MS: m/z 535 [M + Na]+; anal. calcd for C21H18- 53.53; H, 5.44; N, 6.76. IR (ATR, selected bands, nmax): 3220,
FeN3O9$0.5H2O: C, 48.39; H, 3.67; N, 8.06; O, 29.16; found: C, 1596, 1512, 1455, 1382, 1238, 1105, 1069, 997, 907, 825,
48.31; H, 3.30; N, 7.90; O, 29.27. IR (ATR, selected bands, nmax): 720 cmꢂ1
.
3197, 1597, 1487, 1348, 1241, 1176, 1146, 1048, 914, 840, 739,
628 cmꢂ1
[Fe(dihydrocaffeic hydroxamate)3] (6k). FeCl3 (195.8 mg, 1.2
mmol) dissolved in methanol (5 mL) was added to the
.
[Fe2(hypogallic hydroxamate)4Cl2] (6g). FeCl3 (73.4 mg, 0.45 suspension of 5k (710 mg, 3.6 mmol) and KOH (201 mg, 3.6
mmol) dissolved in methanol (3 mL) was added to the mmol) in methanol (25 mL). The dark brown solution was
suspension of 5g (230 mg, 1.36 mmol) and KOH (76 mg, 1.36 stirred for 2 h, ltered and the solvent was removed in vacuo.
mmol) in methanol (10 mL). The dark red suspension was The dry residue was extracted with MeOH/CHCl3 1 : 1 (60 mL)
stirred for 2 h, ltered and the solvent was evaporated. The and le in refrigerator overnight. The suspension was ltered and
residue was extracted with methanol (20 mL) and le in diethyl ether was slowly added yielding the formation of a black
refrigerator overnight. The solution was ltered and the solvent precipitate. The product was ltered off and then dried in vacuo at
was evaporated to small volume (3 mL). Slow diffusion of diethyl r.t. Yield: 168 mg, 0.21 mmol, 18%. ESI-MS: m/z 667 [M + Na]+; m/z
ether into the methanolic solution resulted in the formation of 449 [ML2 + H]+; anal. calcd for C27H30FeN3O12$0.75CHCl3$
dark red crystals. The product was ltered off and then dried in 0.75MeOH: C, 45.16; H, 4.49; N, 5.54; found: C, 45.23; H, 4.78; N,
vacuo at r.t. Yield: 81 mg, 0.1 mmol, 21%. X-ray diffraction 5.21. IR (ATR, selected bands, nmax): 3321, 2970, 1595, 1518,
quality single crystals were grown in methanol (diffusion of 1449, 1359, 1282, 1238, 1195, 1114, 1068, 1032, 1000, 862, 805,
diethyl ether into the methanolic solution). ESI-MS: m/z 390 [M 717 cmꢂ1
.
ꢂ 2Cl ꢂ 4H]2ꢂ; anal. calcd for C28H24Cl2Fe2N4O16$H2O: C,
[Fe(vanillic hydroxamate)3] (6l). FeCl3$6H2O (177 mg, 0.65
38.52; H, 3.00; N, 6.42; O, 31.15; found: C, 38.68; H, 3.05; N, 6.30; mmol) dissolved in water (2 mL) was added to the solution of 5l
O, 31.20. IR (ATR, selected bands, nmax): 2978, 1565, 1520, 1463, (362 mg, 1.ꢀ99 mmol) and KOH (110 mg, 1.96 mmol) in water (20
1318, 1277, 1240, 1158, 1062, 987, 863, 827, 792, 732, 633, 601 mL) at 50 C. Red precipitate was formed and the suspension
cmꢂ1
.
was stirred at 50 ꢀC for 2 h and ltered hot. The precipitate was
ꢀ
[Fe(b-resorcylic hydroxamate)3] (6h). FeCl3$6H2O (155 mg, washed with water and dried in vacuo at 50 C. Yield: 200 mg,
0.57 mmol) dissolved in water (2 mL) was added to the solution 0.31 mmol, 47%. ESI-MS: m/z 625 [M + Na]+; elemental analysis:
of 5h (293 mg, 1.73 mmol) and KOH (97 mg, 1.73 mmol) in calcd for C24H24FeN3O12$2H2O: C, 45.16; H, 4.42; N, 6.58;
ꢀ
water (20 mL) at 40 C. A dark red precipitate was formed and found: C, 45.13; H, 4.41; N, 6.11. IR (ATR, selected bands, nmax):
ꢀ
the suspension was stirred at 40 C for 2 h and ltered while 3202, 1597, 1485, 1339, 1288, 1233, 1121, 1054, 1029, 955, 852,
hot. The precipitate was washed with water and dried in vacuo at 788, 708, 649 cmꢂ1
.
50 ꢀC. Yield: 151 mg, 0.26 mmol, 45%. ESI-MS: m/z 583 [M +
Na]+; anal. calcd for C21H18FeN3O12$1.5H2O: C, 42.95; H, 3.60;
Isolation of the seawater-soluble humic acid fraction
N, 7.16; found: C, 42.89; H, 3.34; N, 7.04. IR (ATR, selected
bands, nmax): 3242, 1603, 1464, 1380, 1256, 1191, 1151, 1098, Seawater-soluble humic acid fraction was isolated according to
1065, 981, 908, 836, 638, 587 cmꢂ1
the previously published procedure.60 1 L ltered creek water
.
[Fe(a-resorcylic hydroxamate)3] (6i). FeCl3 (96 mg, 0.59 (Craggie Burn, Scotland UK) was concentrated under reduced
mmol) dissolved in water (2 mL) was added to the solution of 5i pressure to approximately 10 mL. 20 mL of articial seawater
acid (300 mg, 1.78 mmol) and KOH (99 mg, 1.77 mmol) in water were added and the solution was kept in darkness at 4 ꢀC for 24
(20 mL) at 50 ꢀC. A dark redꢀ precipitate was formed and the h. The sample was lled up to 50 mL with articial seawater and
suspension was stirred at 50 C for 2 h and ltered while hot. ltrated through 0.45 mm cellulose-acetate membrane. 3 mL of
This journal is © The Royal Society of Chemistry 2016
RSC Adv., 2016, 6, 40238–40249 | 40241

View Article Online
RSC Advances
Paper
the ltrate were treated with 1 mL of concentrated nitric acid to 82%) depending on the utilized purication protocol. In order
estimate the iron content with GF-AAS (6.131 mg Lꢂ1). For algal to conrm the formation of the desired products, 1H-NMR
growth experiments, the applied volume of the concentrated spectroscopy, electrospray ionization mass spectrometry (ESI-
humic acid sample contained the same iron amount as the MS) and elemental analyses were performed for the synthe-
modied sterile f/2 medium58 (1.17 ꢃ 10ꢂ5 mol Lꢂ1).
sized ligands (see ESI†). Due to the paramagnetic properties of
Fe(III) complexes, 1H-NMR spectroscopy could not be utilized for
characterization and therefore other methods were used
instead (elemental analysis, ESI-MS, X-ray diffraction analysis).
The ESI-MS spectra of the complexes (ESI-MS, positive and
negative mode) were measured in methanol or in acetonitrile.
Overall the positively charged H+, Na+ or K+ adducts were
detected. For example complex 6a showed peaks with m/z values
Results and discussion
Synthesis – general overview
The hydroxamate ligands were synthesized starting from the
respective commercially available benzoic acids, which were
converted into the corresponding methyl esters 1c–1l. Methyl
esters bearing free phenolic groups were protected with benzyl
groups (2c–2l) to avoid undesired side reactions. In the next step
the methyl esters were cleaved under alkaline conditions and
converted to the corresponding hydroxamic acids in a one-step
approach using ethyl chloroformate to generate intermediately
the more reactive anhydrides which were treated with hydrox-
ylamine yielding the desired hydroxamic acids (4b–4l).61
Cleavage of the benzyl groups was performed by palladium-
catalyzed hydrogenation yielding the desired ligands (5c–5l) in
good yields (Scheme 1). 3,4,5-Trimethoxybenzhydroxamic acid
4a was synthesized according to the previously published
procedure by reaction of 3,4,5-trimethoxybenzoic acid with
thionyl chloride and subsequent conversion with hydroxyl-
amine hydrochloride under alkaline conditions.62 The Fe(III)
complexes were synthesized by deprotonation of the ligand with
potassium hydroxide and addition of iron chloride in aqueous
or methanolic solution (Scheme 2).
of 773, 757 and 547, which were assigned to the [Fe(C10H12
-
NO5)3K]+, [Fe(C10H12NO5)3Na]+ and [Fe(C10H12NO5)2]+ species,
respectively. In the case of the dinuclear Fe(III) complex 6g
a peak at m/z 390, which can be assigned to [Fe(C7H6NO4)2 ꢂ
2H+]ꢂ. All detected molecular peaks of the complexes showed
the typical isotopic iron pattern. Also the electrochemical
behaviour of the complexes 6a–l and their corresponding
ligands was studied by means of cyclic voltammetry. Due to the
redox activity of the ligands and subsequent reactions aer
irreversible reduction of iron(III) to iron(II), interpretation of the
cyclic voltammograms and clear assignment of the iron reduc-
tion peak was not clearly achievable. What can be derived is that
all complexes (see S20–S31†) showed only irreversible reduction
waves. Depending on the ligand scaffold, the possible reduction
peaks of iron(III) were found within a broad potential range but
in most cases between ꢂ1.6 and ꢂ0.5 V vs. NHE. However, the
measured values are distinctly lower compared to the reported
potentials of Fe(III)-EDTA or Fe(III)-SOD (+0.12 and +0.27 V,
respectively)27 and therefore more in the region of ferrienter-
obactin and ferrioxamine (ꢂ0.75 and ꢂ0.45 V, respectively).63 In
this case Fe(III) cannot be reduced by biological reduction
agents and the Fenton reaction is not feasible.
The purication of the complexes were found to be the
crucial step of the synthesis, because the complexes were good
to highly soluble in protic solvents as the formed by-product
KCl. Separation via Sephadex G-10 or LH 20 failed and there-
fore a broad range of different purication strategies were
tested. In some cases (6f, 6h, 6i, 6j and 6l) work up could be
facilitated by using water as reaction solvents. However,
Single-crystal X-ray analysis
multiple extraction steps were necessary for complexes 6c and Single crystals suitable for X-ray diffraction analysis were
6k to isolate the salt-free compound explaining the poor yields. obtained for complexes 6a, 6b, 6d, 6e, 6g, 6j by either precipi-
Overall, the complexes were isolated in poor to good yields (18– tation from acetone, water, methanol, ethanol or slow diffusion
Scheme 1 Procedure for the synthesis of benzoic hydroxamic acids.
40242 | RSC Adv., 2016, 6, 40238–40249
This journal is © The Royal Society of Chemistry 2016

View Article Online
Paper
RSC Advances
Scheme 2 Synthesis of monomeric and dimeric Fe(III) complexes bearing benzoic hydroxamates.
of diethyl ether into acetone or methanol solution (see Experi- respectively. The complex 6a crystallized in the triclinic
ꢀ
mental section, see Fig. 1 and S14–S19†). The complexes 6j and centrosymmetric space group P1 and 6e in the orthorombic
6g crystallized in the monoclinic space group P21/n and C2/c, space group Pbca. In case of monomeric iron complexes, the
octahedral Fe(III) center is surrounded by three chelate ligands
bound via oxygen atoms of the hydroxamate groups (Table
S14†). Compounds 6g and 6b crystallized in the monoclinic
space group P21/c (Fig. 1). X-ray diffraction studies of 6g
conrmed the formation of a dinuclear Fe(III) complex. Catechol
groups are well known for their chelating properties and strong
interaction with iron; however, the hydroxamic groups seem to
have higher affinity towards iron explaining the observed
structure. Important bond lengths between iron and the donor
atoms are shown in Table 1.
Comparison with the EXAFS studies45 of natural humic acids
showed the similarity to the model compounds. In all crystal-
˚
lized complexes the Fe–O bond distances were around 2 A and
are in the same range as the distances in humic acids. The bond
length indicates the octahedral iron(III) conguration. Complex
˚
6b, with the Fe–O bond length of 1.98 A showed best agreement
with experimental values from EXAFS studies. The Fe/C
˚
˚
distance in model compounds was between 2.78 A and 2.81 A
and signicantly shorter as in natural complexes where this
˚
bond is about 2.95 A. Due to the shorter distance to carbon, the
Fe–O–C bond angle was around 112ꢀ and also smaller as
described in the literature namely 125ꢀ.
The EXAFS studies of natural humic acids indicate the
presence of monomeric iron complexes in the solution which
ts to all models with exception of 6g. In 6g the Fe/Fe bond
˚
length of 3.28 A is in agreement with the values measured for
the precipitated fulvic acids where the dimeric/polymeric form
is predominant. Overall the synthesized hydroxamate-based
Fe(III) complexes seem to be appropriate models for humic acid.
EPR spectroscopy
Fig. 1 ORTEP view of one independent molecule of 6b (top) and 6g
(bottom). Thermal displacement ellipsoids are drawn at the 50%
probability level.
EPR measurements revealed the presence of rhombic high spin
Fe(III) in all samples (Fig. 2). The EPR line at geff ¼ 4.25 is
This journal is © The Royal Society of Chemistry 2016
RSC Adv., 2016, 6, 40238–40249 | 40243

View Article Online
RSC Advances
Paper
Table 1 Selected bond distances (A˚) and angles (degree) for 6a, 6b, 6d, 6e, 6g, and 6j
EXAFS45
6a
6b
6d
6e
6g
6j
Fe–O(C)
Fe–O(N)
Fe–C
1.98–1.99
—
2.95
2.045
1.987
2.808
1.986(3)
2.038(3)
2.80
2.021
2.002
2.792(6)
2.026(3)
2.001
2.785(6)
2.000
2.037(3)
1.991(3)
2.794(6)
2.038(6)
2.7805
3.647
Fe–Fe
Fe–C–C
Fe–N
3.3
—
—
—
—
—
—
3.282
—
4.239(6)
2.800(3)
113
4.217(3)
2.787(6)
113
4.219(6)
2.81
114
4.218
2.798
113
4.1915
2.839(6)
114.5
4.247(6)
2.791(6)
113
Fe–O–C
125
similar to our models have been already estimated.66 The
determined values were in the range between 28.8 and 29.7 for
the formation of trihydroxamates. Logarithmic stability
constants of the complexes with benzoic hydroxamic acid and
propanohydroxamic acid, which show structural similarities
with our models, were 28.8 and 28.44 respectively. Comparing
those values with the stability constants of Fe(III) EDTA which is
25.1,67 iron trihydroxamates are supposed to be more stable.
However it should be considered that EDTA possesses six
binding sides and the formation of its complex with iron(III) is
thermodynamically favoured in comparison to bidentate
hydroxamates. In order to prove this presumption, we investi-
gated the stability of 6e (25 nM) treated with EDTA in distilled
water and seawater. Aer addition of EDTA (25 nM) to the
solution of 6e, signicant changes in UV-Vis spectra were
observed over 24 h (see S12–S13†) indicating ligand exchange
reactions. Therefore EDTA was not added to the complex solu-
tions used for algal growth experiments. The aqueous stability
of the synthesized complexes in seawater at pH z 8.2–8.4 (see
Fig. S1–S4†) and distilled water (see Fig. S5–S9†) was monitored
over 24 h by means of UV-Vis spectrometry. Complex 6k bearing
hydrocaffeic hydroxamate was found to be the most instable
over this time range, whereas most of the investigated
complexes showed no changes in their UV-Vis spectra under
these conditions (Fig. 3). Furthermore, 6k showed an increased
stability in distilled water over 24 h. The possible explanation
for the changes in the recorded UV-Vis spectra of 6k can be
attributed to the presence of a uncoordinated catecholic moiety,
which is prone to photo-oxidation in seawater; however, stabi-
lization can be achieved by coordination of iron(III) ions.
Fig.
2 Representative EPR spectra for all mononuclear Fe(III)
complexes (6e) and the dinuclear complex 6g recorded at 9.43 GHz
and 90 K. Arrows indicate the most prominent effective g values
relevant for identification of the high spin Fe(III).
a characteristic signal of mononuclear ferric complexes.64 This
high-intensity signal results from the three coinciding g values
for the I ꢁ 3/2 > doublet at the 5/2 system. The appearance of
a weak feature at approximately geff ¼ 9 is also commonly seen
in spectra of rhombic Fe(III) complexes and can be attributed to
the 1 4 2 transition. For the dinuclear Fe(III) complex 6g, the
EPR signals almost disappeared most likely due to spin
coupling effects.65 The rhombicity (h) as determined by the zero-
eld splitting parameters, was calculated to be 0.327 for all the S
¼ 5/2 systems. As a completely rhombic system is reached at h ¼
0.33, the crystal eld at the Fe(III) center shows a strong ortho-
rhombic character.
The obtained results conrm that the Fe(III) is tightly bound
to the hydroxamate and the free catechol group is able to react
with other metal ions or undergo oxidation. In seawater, the
intensity of the signal at 335 nm increases over 24 h and also
a new signal arose at 380 nm over time. In distilled water only
the signal at 335 nm increased but no differences in absorbance
at 380 nm was observed. Measurements at pH 11 showed the
increase in absorbance at 335 nm and also the signal at 380 nm,
which appeared directly aer pH-adjustment from 7 to 11. On
the basis of these results, the signal at 335 nm can be assigned
to the oxidation of catecholic moiety, which proceeds more
slowly in distilled water. The signal at 380 nm seems to be
associated with the deprotonation of 6k which occurs very fast
at pH 11 in contrast to pH 8.2–8.4 (seawater). Free catecholic
Stability studies
The stability of these model compounds in aqueous systems is
important for further biological investigations, because the
complexes have to be stable enough to enable cellular iron
accumulation and if the complex decomposes too fast, the
released iron will precipitate in water. According to literature
iron shows high affinity to ligands bearing hydroxamic groups
and iron monohydroxamates are, under the condition that no
other instable groups are present, in most cases stable in
aqueous solutions. At pH regions of our interest, namely 8.2–8.4
(seawater) the formation of trihydroxamates is highly preferred.
Stability constants of the complexes with ligand scaffolds
40244 | RSC Adv., 2016, 6, 40238–40249
This journal is © The Royal Society of Chemistry 2016

View Article Online
Paper
RSC Advances
Fig. 4 Time dependent UV-Vis spectra of 6f (A, B) and 6l (C, D) in
seawater (A, C) and in distilled water (B, D). (The spectra were
measured over 24 h in an 1 h interval, start and end point of the
measurement are indicated by an arrow in (A) and (C).)
components because no inuence of pH change on stability has
been observed.
Algal growth experiments
C. salina and D. lutheri showed signicantly different growth
response depending on the tested model compounds and
control samples. Fig. 5 shows the respective growth curves of C.
salina batch cultures using different iron sources and chelating
agents. Without iron supply the culture can't grow effectively
aer its own intracellular reserve is consumed. If our
compounds do not supply algae with iron, the culture growth
will be comparable or worse. As second negative control we
supplied the culture with iron but we waived the chelator EDTA.
Fig. 3 Time dependent UV-Vis spectra of 6k (above) and 6e (below) in
seawater over 24 hours (spectra were measured in an 1 h interval, start
and end point of the measurement are indicated by an arrow).
moieties are also present in the structure of 6g, but in contrast
to 6k, this model compound remained stable in solution. This
contrary behaviour might be explained by the formation of
stable hydrogen bonds between the nitrogen and the oxygen
atoms of the ligand scaffold. To investigate the impact of the
ligand scaffold on the observed different behaviour of 6g and
6k, the stability of the respective free ligands, namely dihy-
drocaffeic hydroxamic acid and hypogallic hydroxamic acid,
were determined in both seawater and distilled water. Both
were found to be instable in seawater and probably oxidation
occurred over time. Complex 6k was not the only instable
compound in seawater. We also saw signicant changes in UV-
Vis spectra of the complexes 6f, 6i and 6l (Fig. 4). In contrast to
the decay in seawater, these compounds were found to be
remarkably stable in distilled water. The observed results for 6f,
6i and 6l might be explained by the reaction or interaction with
metal ions from seawater. Also the pH value of approximately
8.2–8.4 can contribute to the observed changes. To verify this
assertion the stability of those complexes at pH 11 was inves-
tigated. In case of the complexes 6f and 6l it is most likely that
both of them are deprotonated at higher pH value (see Fig. S10
and S11†). The increasing signal in seawater over 24 h is present Fig. 5 Growth curves of C. salina (error bars: ꢁSD) treated with model
compounds 6c, 6d, 6e, 6h, 6j and isolated fraction of humic acid
in the spectra obtained at pH 11 immediately aer base addi-
compared to control samples (+Fe, +EDTA; +Fe, ꢂEDTA; ꢂFe,
tion. The complex 6i seems to interact with seawater
+EDTA).
This journal is © The Royal Society of Chemistry 2016
RSC Adv., 2016, 6, 40238–40249 | 40245

View Article Online
RSC Advances
Paper
Table 2 Relation between the end algae concentration of control
samples to samples treated with 6c, 6d, 6e, 6h, 6j, and isolated humic
acid
In this case the bioavailability of iron and other metal ions is
decreased. The EDTA complex with iron has good aqueous
solubility; however, the metal is bound very strongly to
chelating ligand, limiting its bioavailability. Nevertheless the
complex is light sensitive, which results in reduction of Fe(III) to
Fe(II) and release of iron (lower affinity of EDTA towards Fe(II)).68
Thus the growth of C. salina without EDTA is expected to be
worse than in the ideal medium. This negative control evaluates
the impact of the chelator and other metal ions on the culture. It
also shows if the algal species can uptake iron and other metal
ions at very low concentrations and how far the growth will be
inhibited. Samples with ideal f/2 medium without iron showed,
as expected, poor growth.
6c
6d
6e
6h
6j
Humic acid
+Fe, +EDTA in %
ꢂFe, +EDTA
13.1
2
0.3
63.4
10
1.6
31
5
0.8
7.7
1.2
0.2
10.2
1.6
0.3
115
18
3
+Fe, ꢂEDTA
compound 6e was promoting growth of the C. salina and the
samples algae concentration reached 31% of the concentration in
samples under ideal conditions. Thus our experiments have
shown that iron can be released from the hydroxamate
complexes and absorbed by algae. Furthermore the solubility of
the complexes plays also an important role for the bioavailability
of iron. Complexes 6e and 6d with one free hydroxy group were
more efficient promoting growth of C. salina than the complexes
with methoxy groups or aliphatic chains with relative poor water
solubility like 6j and 6c. Another aspect of the rst experiment
was the different growth patterns of the cultures. Samples with
HS and iron without EDTA started to grow very soon but we
noticed the short delay in growth of the sample with the ideal
medium.
This delay was much more signicant in case of model
compounds. We observed the distinct changes in algae
concentration aer approximately two weeks. That lead us to
the conclusion that the stability of the complexes in water
seems to play an important role for the release and bioavail-
ability of iron. Unfortunately it was not possible to monitor the
growth of the culture longer than 30 days due to the limited
stability of f/2 medium over this time period. Further observa-
tion would not provide reliable values due to the changing
conditions between several cultures. In the second experiment,
three model compounds namely 6b, 6i and 6k were tested. Fig. 6
shows growth curves of C. salina batch cultures under different
conditions including standard deviation.
Controls with medium with iron but without the chelator
EDTA showed moderate growth rates whereas samples with full
medium with iron and EDTA have grown to algae concentration
of approximately 7.7 ꢃ 106 ꢁ 6.6 ꢃ 105 cells per mL. It is
important to mention, that the samples without EDTA showed no
further growth aer 12 days. At the beginning of experiment the
removal of inorganic iron was not necessarily complete. The low
concentration of dissolved iron could be sufficient to support the
growth of culture. Aer achieving the cell number of 3 ꢃ 106 ꢁ
2.2 ꢃ 105 the requirements for iron were probably too high to
support further growth. Surprisingly the sample containing iso-
lated humic acid fraction had the biggest impact on the culture,
reaching the concentration near 9 ꢃ 106 ꢁ 3.6 ꢃ 106 cells per mL
at the end of the experiment, resulting in slightly better growth
than in the ideal medium. Moreover this sample showed not only
the highest cell concentration at the end of experiment but also
the fastest response and growth. We assume that, similar to the
Fe-EDTA complex, the light sensitivity of the Fe(III)–AHS
complexes and photoreduction of the Fe(III) to Fe(II) may
contribute to the observed effects in C. salina. Moreover Fe(III)–
AHS complexes can bind to the cell surface of the algae and
undergo kinetically favoured biochemical reduction.69 The
control sample without chelator EDTA also showed quick
response and grew fast. The explanation for this could be the
precipitation of iron(III) and formation of amorphous Fe(OH)3,
which can supply algae with iron. The solubility of amorphous
Fe(OH)3 is much better than the solubility of crystalline Fe(OH)3.
Over days amorphous Fe(OH)3 can convert to the crystalline
form. Therefore the bioavailability of iron is too low aer a few
days to cover the requirements.70 The model compounds showed
mostly moderate to good impact on the culture growth with
exception of 6c and 6j, which didn't signicantly enhance the
growth rate of C. salina comparing to the controls without iron. It
is important to mention that the impact of model compounds on
the growth of C. salina was noticeably weaker than for the control
sample with chelator and iron but it was stronger compared to
the control without iron. The relation between the concentration
of algae in samples treated with model compounds and control
samples at the end of the experiment is shown in Table 2. The
most efficient growth promoting model compound was 6d with
salicylic hydroxamic acid, where the treated samples reached the
end concentration of 63.4% compared to ideal conditions.
Moreover the algal growth was also better than in samples con-
Fig. 6 Growth curves of C. salina (error bars: ꢁSD) treated with model
compounds 6b, 6i and 6k compared to control samples (+Fe, +EDTA;
taining iron chloride without chelator. The similar model +Fe, ꢂEDTA; ꢂFe, +EDTA).
40246 | RSC Adv., 2016, 6, 40238–40249
This journal is © The Royal Society of Chemistry 2016

View Article Online
Paper
RSC Advances
Table 3 Relation between the end algae concentration of control
samples to samples treated with 6b, 6i and 6k
6b
6i
6k
Humic acid
+Fe, +EDTA in %
ꢂFe, +EDTA
0.7
0.05
0.1
17.3
1.3
2.1
29.7
2.2
3.6
166
17.6
29.2
+Fe, ꢂEDTA
The samples kept at ideal conditions reached aer 24 days
the algae concentration of approximately 4.5 ꢃ 106 ꢁ 8.2 ꢃ 104
cells per mL whereas the growth of cultures lacking EDTA or
iron was poor. In contrast to the rst experiment, control
samples with chelator and without iron reached slightly higher
concentration as control samples with iron without chelator.
Nevertheless both values were low. Similar to the rst experi-
ment, samples containing isolated humic acid fraction had the
biggest positive impact on the culture, reaching at the end of
the experiment the concentration near 10 ꢃ 106 ꢁ 6.7 ꢃ 105
cells per mL which was better as the end concentration of the
ideal medium. Model compounds 6b and 6i did not signi-
cantly support the growth of C. salina. In fact, in samples
treated with 6b the lowest nal algal concentration was
observed. This can be explained with hydrophobic character
and low solubility of the cinnamic hydroxamic acid in the
complex 6b. UV-Vis studies of this compound have shown
signicant decrease in absorbance aer 24 h due to the
precipitation. Together with the results from the algal growth
experiments, we assume that the iron bound in 6b is not
sufficiently bioavailable for the C. salina. Samples containing
6k showed best response to the treatment and reached the
concentration of 1.3 ꢃ 106 ꢁ 5.2 ꢃ 105 cells per mL which is
30% of the algal concentration in ideal medium (Table 3). This
effect of 6k on algae growth can be explained by the moderate
stability of the complex in seawater which was already
mentioned in UV-Vis studies. Moreover the ligand in 6k has
two possible metal-binding sides. Dihydrocaffeic hydroxamic
acid is bound via hydroxamate to iron and the cathecholic
group is free for coordination which might additionally
enhance the bioavailability of other metal ions as for example
Zn and Cu.
This is an important factor because it is unlikely to achieve
the same growth rate of the culture as under the ideal condi-
tions using model complexes instead of EDTA and iron chlo-
ride. EDTA is not only chelator for iron but also for other metals
included in micronutrient solution like for example Cu, Zn and
Mn. In order to compare the results from the rst two experi-
ments with C. salina, we performed the same tests with the
photoautotrophic algal species D. lutheri. In this experiment, six
model compounds namely 6c, 6d, 6e, 6g, 6j and 6k were
investigated. Fig. 7 shows the determined growth curves of D.
lutheri batch cultures under different conditions including
standard deviation. The culture of D. lutheri reached at all tested
conditions signicantly lower cell numbers in the growth
experiments compared to C. salina. Control samples with ideal
medium with iron and EDTA have grown to algae concentration
of approximately 2.1 ꢃ 106 ꢁ 7.1 ꢃ 105 cells per mL.
Fig. 7 Growth curves of D. lutheri treated with model compounds 6c,
6d, 6e, 6g, 6j and 6k compared to control samples (+Fe, +EDTA; +Fe,
ꢂEDTA; ꢂFe, +EDTA).
Table 4 Relation between the end algae concentration of control
samples to samples treated with 6c, 6d, 6e, 6g, 6j, 6k and isolated
humic acid
6c 6d
6e
6g
6j 6k
Humic acid
+Fe, +EDTA in %
ꢂFe, +EDTA
0
0
0
38.6 36.6 27.2
0
0
0
28.1 155.7
0.8
0.4
0.8
0.4
0.6
0.3
0.6
0.3
3.4
1.6
+Fe, ꢂEDTA
Surprisingly the samples containing medium with iron but
without the chelator EDTA had almost the same impact on the
culture as the positive control whereas the samples lacking iron
reached the end concentration of 1 ꢃ 106 ꢁ 6 ꢃ 104 cells per
mL.
These ndings might be explained by slower growth due to
the lower cell number, higher intracellular iron pool or better
iron uptake at low concentrations. As already observed in C.
salina, samples treated with soluble humic acid fraction showed
the best growth response. None of the investigated model
compounds showed a big impact on the culture and the
concentrations at the end of the experiment remained under
the concentration of samples treated with ideal medium
without iron. The complexes 6c and 6j did not support the
growth of D. lutheri at all. Those two complexes had also been
found to be the less effective in the culture of C. salina. The
relation between the concentration of algae in samples treated
with model compounds and control samples at the end of the
experiment is shown in Table 4.
In all performed algal experiments, both C. salina and D.
lutheri showed the typical growth pattern in full medium.71,72
In iron-free samples (ꢂFe, +EDTA) the growth was strongly
inhibited, which means, iron deciency conditions were given.
The relative long time between the beginning of each experi-
ment and start of the exponential growth is typical and known
This journal is © The Royal Society of Chemistry 2016
RSC Adv., 2016, 6, 40238–40249 | 40247

View Article Online
RSC Advances
Paper
from large-scale iron fertilization experiments in HNLC
regions.9,73
9 J. W. Pitchford and J. Brindley, J. Plankton Res., 1999, 21, 525–
547.
10 A. M. Edwards, T. Platt and S. Sathyendranath, Ecol. Modell.,
2004, 171, 103–125.
11 A. Watson, P. Liss and R. Duce, Limnol. Oceanogr., 1991, 36,
1960–1965.
Conclusions
Herein, we present the synthesis of 11 novel iron(III) complexes
as promising model compounds for humic acids. All complexes
were characterized by elemental analysis, ESI-MS, IR, EPR,
cyclic voltammetry and some of them by X-ray diffraction
analysis. Comparison of the bond lengths, angles and distances
in evaluated X-ray structures with the EXAFS studies45 of natural
humic acids showed that the selected model structures are
suitable as AHS model compounds. The ligand scaffold and the
presence of different functional groups seem to have a big
inuence on the stability and solubility of the compounds,
which was monitored in water and seawater over 24 h. We
determined the impact of the synthesized model compounds on
growth of the algae species C. salina and D. lutheri. The study
revealed that several of our models support the growth of the
algae culture conrming that the iron bound in the complexes
is bioavailable. The results obtained from algal growth experi-
ments support our proposal that iron complexes with benzoic
hydroxamates are suitable models for the research on iron
binding ability of humic substances; however, the induced
growth upon treatment with the model compounds is still lower
compared to ideal conditions and might be explained by the
observed low redox potentials of the complexes. Therefore
optimization with regard on the algal growth and redox poten-
tials is required and currently ongoing. Furthermore, the
synthesized model compounds will be used as reference
substances in EXAFS experiments, where the similarity in the
nearest surrounding of iron in the natural humic acids will be
investigated.
12 J. H. Martin and S. E. Fitzwater, Nature, 1988, 331, 341–343.
13 S. E. Fitzwater, K. H. Coale, R. M. Gordon, K. S. Johnson and
M. E. Ondrusek, Deep Sea Res., Part II, 1996, 43, 995–1015.
14 D. A. Hutchins and K. W. Bruland, Nature, 1998, 393, 561–
564.
15 J. Lewandowska and A. Kosakowska, Oceanologia, 2004, 46,
269–287.
16 B. W. Frost, Mar. Ecol.: Prog. Ser., 1987, 39, 49–68.
17 C. L. Moloney and J. G. Field, Limnol. Oceanogr., 1989, 34,
1290–1299.
18 C. L. Moloney and J. G. Field, J. Plankton Res., 1991, 13, 1003–
1038.
19 N. M. Price, L. F. Andersen and F. M. M. Morel, Deep–Sea
Res., 1991, 38, 1361–1378.
20 W. G. Sunda and S. A. Huntsman, Mar. Chem., 1995, 50, 189–
206.
21 J. F. Wu, E. Boyle, W. Sunda and L. S. Wen, Science, 2001,
293, 847–849.
22 J. F. Wu and G. W. Luther, Mar. Chem., 1995, 50, 159–177.
23 E. L. Rue and K. W. Bruland, Mar. Chem., 1995, 50, 117–138.
24 A. E. Witter and G. W. Luther, Mar. Chem., 1998, 62, 241–258.
25 C. M. G. Van Den Berg, Mar. Chem., 1995, 50, 139–157.
26 K. Hunter and D. Turner, The Biogeochemistry of Iron in
Seawater, Wiley, West Sussex, England, 2001.
27 J. L. Pierre and M. Fontecave, BioMetals, 1999, 12, 195–199.
28 J. K. Moore and O. Braucher, Biogeosciences, 2008, 5, 631–
656.
29 E. R. Sholkovitz, Geochim. Cosmochim. Acta, 1976, 40, 831–
845.
Acknowledgements
30 J. M. Eckert and E. R. Sholkovitz, Geochim. Cosmochim. Acta,
1976, 40, 847–848.
31 E. A. Boyle, J. M. Edmond and E. R. Sholkovitz, Geochim.
Cosmochim. Acta, 1977, 41, 1313–1324.
32 E. R. Sholkovitz and D. Copland, Geochim. Cosmochim. Acta,
1981, 45, 181–189.
We cordially thank the FWF Austrian Science Fund (Project
number P 25849-N19) for nancial support and Markus Gruber
for his valuable advices regarding algae cultivation.
33 K. A. Hunter, M. R. Leonard, P. D. Carpenter and J. D. Smith,
Colloids Surf., A, 1997, 120, 111–121.
Notes and references
1 R. Osterber, Nature, 1974, 249, 382–383.
34 U. Nowostawska, J. P. Kim and K. A. Hunter, Mar. Chem.,
2008, 110, 205–210.
2 F. Egami, J. Biochem., 1975, 77, 1165–1169.
3 J. A. Raven, New Phytol., 1988, 109, 279–287.
4 W. W. Gregg, M. E. Conkright, J. E. O'Reilly, F. S. Patt,
M. H. H. Wang, J. A. Yoder and N. W. Casey, Appl. Opt.,
2002, 41, 1615–1628.
35 R. Krachler, F. von der Kammer, F. Jirsa, A. Suphandag,
R. F. Krachler, C. Plessl, M. Vogt, B. K. Keppler and
T. Hofmann, Global Biogeochem. Cycles, 2012, 26, GB3024.
36 R. Krachler, R. F. Krachler, F. von der Kammer,
A. Suephandag, F. Jirsa, S. Ayromlou, T. Hofmann and
B. K. Keppler, Sci. Total Environ., 2010, 408, 2402–2408.
5 C. B. Field, M. J. Behrenfeld, J. T. Randerson and
P. Falkowski, Science, 1998, 281, 237–240.
6 P. W. Boyd and M. J. Ellwood, Nat. Geosci., 2010, 3, 675–682. 37 R. Krachler, F. Jirsa and S. Ayromlou, Biogeosciences, 2005, 2,
7 R. F. Nolting, L. J. A. Gerringa, M. J. W. Swagerman,
311–315.
K. R. Timmermans and H. J. W. de Baar, Mar. Chem., 1998, 38 X. W. Liu and F. J. Millero, Mar. Chem., 2002, 77, 43–54.
62, 335–352.
39 T. Hiemstra and W. H. van Riemsdijk, Mar. Chem., 2006, 102,
181–197.
8 J. H. Martin and R. M. Gordon, Deep–Sea Res., 1988, 35, 177–
196.
40248 | RSC Adv., 2016, 6, 40238–40249
This journal is © The Royal Society of Chemistry 2016

View Article Online
Paper
RSC Advances
40 S. A. Waksman and T. C. Cordon, J. Am. Chem. Soc., 1936, 58, 58 R. R. Guillard and J. H. Ryther, Can. J. Microbiol., 1962, 8,
969–972. 229–239.
41 S. A. Waksman and I. J. Hutchings, Soil Sci., 1936, 42, 119– 59 D. R. Kester, I. W. Duedall, D. N. Connors and
130. R. m. Pytkowic, Limnol. Oceanogr., 1967, 12, 176–179.
42 G. W. Ding, J. D. Mao and B. S. Xing, Commun. Soil Sci. Plant 60 R. Krachler, R. F. Krachler, G. Wallner, S. Hann, M. Laux,
Anal., 2001, 32, 1991–2005.
43 J. B. Jahnel and F. H. Frimmel, Acta Hydrochim. Hydrobiol.,
1995, 23, 31–35.
44 E. Ibisanmi, S. G. Sander, P. W. Boyd, A. R. Bowie and
K. A. Hunter, Deep Sea Res., Part II, 2011, 58, 2113–2125.
45 J. W. J. van Schaik, I. Persson, D. B. Kleja and
J. P. Gustafsson, Environ. Sci. Technol., 2008, 42, 2367–2373.
M. F. C. Recalde, F. Jirsa, E. Neubauer, F. von der Kammer,
T. Hofmann and B. K. Keppler, Mar. Chem., 2015, 174, 85–93.
61 A. S. Reddy, M. S. Kumar and G. R. Reddy, Tetrahedron Lett.,
2000, 41, 6285–6288.
62 R. Takahashi, T. Nunokawa, T. Shibuya, R. Tomita,
Y. Tatewaki, S. Okada, T. Kimura, S. Shimada and
H. Matsuda, Bull. Agric. Chem. Soc. Jpn., 2012, 85, 236–244.
46 F. H. Frimmel and J. Geywitz, Fresenius' Z. Anal. Chem., 1983, 63 S. R. Cooper, J. V. Mcardle and K. N. Raymond, Proc. Natl.
316, 582–588.
Acad. Sci. U. S. A., 1978, 75, 3551–3554.
47 F. H. Frimmel, Z. Wasser- Abwasser-Forsch., 1981, 14, 7–10.
48 L. Bauer and O. Exner, Angew. Chem., Int. Ed. Engl., 1974, 13,
376–384.
64 W. R. Hagen, Biomolecular EPR Spectroscopy, CRC Press, Boca
Raton, Fl, USA, 2009.
65 A. M. N. Silva, X. Kong, M. C. Parkin, R. Cammack and
R. C. Hider, Dalton Trans., 2009, 8616–8625.
49 D. C. Glick and A. Davis, Org. Geochem., 1991, 17, 421–430.
50 H. Knicker, P. G. Hatcher and A. W. Scaroni, Int. J. Coal Geol., 66 E. Farkas, E. Kozma, M. Petho, K. M. Herlihy and C. Micera,
1996, 32, 255–278. Polyhedron, 1998, 17, 3331–3342.
51 A. Vairavamurthy and S. Wang, Environ. Sci. Technol., 2002, 67 H. Gomathi, Bioelectrochemistry, 2000, 16, 459–465.
36, 3050–3056.
68 F. M. M. Morel, Abstr. Pap., Jt. Conf.–Chem. Inst. Can. Am.
Chem. Soc., 2005, 230, U1712.
69 H. Lis, Y. Shaked, C. Kranzler, N. Keren and F. M. M. Morel,
ISME J., 2015, 9, 1003–1013.
52 W. C. Barrette, H. W. Johnson and D. T. Sawyer, Anal. Chem.,
1984, 56, 1890–1898.
53 W. R. Hagen, Mol. Phys., 2007, 105, 2031–2039.
54 M. R. Pressprich and J. Chambers, SAINT + Integration 70 S. M. Kraemer, Aquat. Sci., 2004, 66, 3–18.
Engine, Program for Crystal Structure Integration, Bruker 71 A. S. Miron, M. C. C. Garcia, F. G. Camacho, E. M. Grima and
Analytical X-ray systems, Madison, 2004.
Y. Chisti, Enzyme Microb. Technol., 2002, 31, 1015–1023.
55 G. M. Sheldrick, Acta Crystallogr., Sect. A: Found. Crystallogr., 72 G. E. Fogg and B. Thake, Algal Cultures and Phytoplanktion
2007, 64, 112.
Ecology, The University of Wisconsin Press, Wisconsin,
United States of America, 1987.
73 M. Fujii and F. Chai, Deep Sea Res., Part II, 2009, 56, 2936–
2947.
56 O. V. Dolomanov, L. J. Bourhis, R. J. Gildea, J. A. K. Howard
and H. Puschmann, J. Appl. Crystallogr., 2009, 42, 339–341.
57 C. B. Hubschle, G. M. Sheldrick and B. Dittrich, J. Appl.
Crystallogr., 2011, 44, 1281–1284.
This journal is © The Royal Society of Chemistry 2016
RSC Adv., 2016, 6, 40238–40249 | 40249