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Mechanism underlying the effectiveness of deferiprone in alleviating Parkinson's disease symptoms Yingying Sun, An Ninh Pham, and T. David Waite ACS Chem. Neurosci., Just Accepted Manuscript • DOI: 10.1021/acschemneuro.7b00478 • Publication Date (Web): 30 Jan 2018 Downloaded from http://pubs.acs.org on February 3, 2018

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ACS Chemical Neuroscience

Mechanism underlying the effectiveness of deferiprone in alleviating Parkinson's disease symptoms Yingying Sun, An Ninh Pham and T. David Waite*

School of Civil and Environmental Engineering, The University of New South Wales, Sydney, NSW 2052, Australia

Re-Submitted ACS Chemical Neuroscience January, 2018

* Corresponding author. Professor T. David Waite, Tel.: +61 2 9385 5059, Email: [email protected]

ACS Paragon Plus Environment

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Abstract Elevation in iron content as well as severe depletion of dopamine (DA) as a result of iron-induced loss of dopaminergic neurons has been recognized to accompany the progression of Parkinson’s disease (PD). To better understand the mechanism of the mitigating effect of the iron chelator deferiprone (DFP) on PD, the interplay between iron and DFP was investigated both in the absence and presence of DA. The results show that DFP was extremely efficient in scavenging both aqueous iron and iron that was loosely bound to DA with the entrapment of iron in Fe-DFP complexed form critical to halting the iron catalyzed degradation of DA and associated generation of toxic metabolites. The DFP related scavenging of dopamine semiquinone (

DA•− ) and superoxide ( O •− 2 ) may also contribute to its positive effects in the treatment of PD.

Keywords: Deferiprone, Iron overload, Dopamine, Parkinson’s disease, Mechanism and Oxidative stress


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Introduction Parkinson’s disease (PD) is the second most common neurodegenerative disorder and affects more than 2% of the population over 65 years of age.1, iron content

3, 4

2

The progression of PD is often accompanied by an elevation in

and severe depletion of dopamine (DA) as a result of preferential loss of dopaminergic

neurons from the substantial nigra pars compacta (SNpc).5,

6

Even though DA is an indispensable

neurotransmitter,7 the DA-mediated generation of toxic quinones and reactive oxygen species (ROS) have long been associated with the occurrence and the progression of the pathological hallmarks of PD.8, 9 It has been reported that the presence of iron may induce the iron catalyzed oxidation of DA,10, 11 which may aggravate the progression of PD by enhancing the accumulation of DA-derived quinones or even the production of strongly oxidizing hydroxyl radicals ( • OH ).12 As such, given the increasing iron content of cells on aging, the redox state of iron present and the transformation of iron between the +II and +III oxidation states are considered to be of particular significance in the progression of PD. Treatments resulting in reduction of the iron content in the diseased brain are generally considered to be a promising approach to disease mitigation in view of the possible effect of such treatment in reducing the iron aggravated generation of DA-derived toxic metabolites and ROS.

Even though a chelation strategy has long been successfully used in treating iron overload diseases such as thalassemia,13 interest in the use of chelators as a therapeutic strategy in neurodegenerative disorders only commenced in 2003.14, 15 Of all the chelators investigated, the small compound, deferiprone (also referred to as CP20, L1 and DMHP shown in Scheme 1 and denoted as DFP hereafter) has been suggested to be a promising candidate. Despite the absence of a successful Phase III trial,16 DFP has been investigated extensively in pilot trials for the mitigation of PD15, 17 in view of its oral ability, high iron affinity, blood brain barrier (BBB) permeability and its apparently manageable influence on systemic iron loss.18-21 Clinical trials have shown that DFP can be used to efficiently scavenge the plasma non-transferrin-bound-iron (NTBI) in patients, a source of iron which is generally thought to be largely responsible for oxidative stress.22 2



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In order to effectively prevent oxidative stress arising from the presence of loosely bound iron and to reduce the iron content in patients experiencing iron overload disease, it is critical for DFP to effectively outcompete other ligands to bind NTBI. However, even though DFP has already achieved full marketing authorization in Europe,23 limited information regarding the kinetics of formation of the key tris-complex (i.e., FeIIIDFP3) has been obtained, especially in the presence of competing ligands. Instead, the effectiveness of DFP in excess iron depletion is mainly assessed by iron excretion and quantitative changes in serum iron (present as either ferritin or NTBI).23 While useful, these assessments may result in mis-interpretation of the effectiveness of DFP, in part as a result of the complicated homeostasis of iron with processes such as erythropoiesis and interactions between ferritin and NTBI influencing the speciation of iron.24,

25

Even

though the DFP-related alleviation in iron-induced oxidative stress is widely recognized, the direct relationship between iron transformations and the concomitant generation of ROS has rarely been reported with most previous studies simply noting the reduced extent of lipid peroxidation in cell cultures or animal models.26-28 While several fundamental studies have been conducted with particular attention given to the reducibility of DFP bound Fe(III) and interaction between DFP chelated Fe(II) and H2O2,21, 29-32 details of iron transformation in the presence of DFP are still far from well understood. To improve our understanding of the therapeutic mechanism of DFP used in the treatment of iron overload diseases, we focus in this study on the formation of the stable non-ionic tris-complex, FeIIIDFP3 and the impact of this formation on the iron induced generation of toxic metabolites. Given our particular interest in the rectifying effect of DFP on PD, instead of choosing the typical NTBI, citrate or albumin bound iron,33 FeIIIDA2 was herein chosen as the target compound in order to investigate the ability and effectiveness of DFP in scavenging DA-bound iron and thereby alleviating the production of toxicants induced by the ironcatalyzed transformation of DA. FeIIIDA2 is the dominant iron-DA complex in PD brains, particularly in the later stage when significantly elevated iron concentrations and the presence of toxic intermediate-induced DA leakage from synaptic vesicles is likely.34 As both Fe(II) and DA semiquinone ( DA•− ) may be slowly released from the FeIIIDA2 complex with resultant production of toxic metabolites,11, 35, 36 the efficiency of the scavenging of DA bound iron by DFP is of great importance in understanding the effectiveness of DFP in mitigating the progression of PD. As such, a kinetic model is developed based on the experimental data 3


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collected in this study with this model enabling a more quantitive description of the iron transformations occurring in the presence of DFP. Even though the model developed does not perfectly describe the experimental data under all conditions examined, it does provide important insights into the chelation mechanism of DFP and should assist in the selection and design of more effective chelators in the future.

Results and Discussion Effectiveness of DFP on the chelation of ferric iron The affinity of DFP for Fe(III) was examined by monitoring the rate and extent of formation of the stable complex, FeIIIDFP3, as a result of its ability to compete with Fe(III) precipitation and, possibly, mobilization of iron from the precipitated amorphous ferric oxyhydroxide (structurally similar to ferrihydrite and ferritin, denoted as AFO hereafter). As shown in Figure 1, addition of 5 µM Fe(III) to a solution containing 10, 20 and 50 µM DFP at pH 7.4 gave rise to the formation of the stable, fully coordinated FeIIIDFP3 complex with the initial concentration of this complex increasing on increase in DFP concentration. A gradual increase in the concentration of FeIIIDFP3 was also observed following the rapid initial formation, especially in the first 10 – 20 minutes at lower DFP to Fe(III) concentration ratios. At high DFP to Fe(III) concentration ratios (e.g., [DFP]/[Fe(III)] = 50:5), almost all the Fe(III) was rapidly incorporated into the stable FeIIIDFP3 complex.

4



Figure 1 Formation of FeIIIDFP3 followed by the addition of 5 µM Fe(III) to air-saturated 0.1 M NaCl solutions containing () 0 µM, (●) 10 µM, () 20 µM and (◆) 50 µM DFP at pH 7.4. Error bars are standard errors from triplicate measurements and solid lines represent the model fit. As shown in Scheme 1, under physiological pH 7.4, DFP mainly exists in the form of HDFP0 (1b) in view of the low pKa1 value of 3.62.37 As a result of resonance, the 1c form renders HDFP0 capable of chelating iron via the two phenol functional groups.

Scheme 1 Acid-base equilibrium of deferiprone Note: a) H2DFP+ represents the protonated deferiprone, HDFP0 represents deferiprone and DFP- represents the deprotonated deferiprone; b) pKa values take from Motekaitis and Martell 38

In a manner similar to other bidentate chelators, DFP could form three different kinds of complexes with aqueous Fe(III) via the hard O donor atom20 with the zero net charged tris-complex being the most stable and dominant species at physiological pH 7.4. Once the rate-limiting mono-complex, FeIIIDFP, is formed, the remaining coordinated H2O molecules can subsequently be replaced by another two DFP molecules (reactions 2 and 3 in Table 1). Not surprisingly, FeIIIDFP is most readily formed via interaction of Fe(III) with the deprotonated DFP species, DFP- (1d) while formation of FeIIIDFP is least favored with the protonated form, H2DFP+. The particular efficacy of DFP- can be mainly attributed to the presence of the fully deprotonated phenol groups with deprotonation of the functional group generally favouring the formation rather than the dissociation of iron-DFP complexes (see details in SI 1). Similar behavior is observed for transferrin for which the tendency for iron chelation decreases with pH.21 Additionally, it is not surprising that the concentration of FeIIIDFP3 increases significantly on increase in DFP concentration since 5


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the proportional increase in the concentration of mono- and fully deprotonated phenol groups of DFP would be expected to be even more effective in preventing the hydrolysis of Fe(III).39 The gradual increase in the FeIIIDFP3 concentration in the first 20 minutes shown in Figure 1 may arise from the formation and detachment of a surface complex between DFP and any iron present as AFO precipitate; i.e.

AFO + DFP → FeIII DFP

(1)

While the strength of the Fe(III)-DFP complex will be the key determinant of the rate and extent of AFO dissolution, the initial association between DFP and AFO may be facilitated by electrostatic attraction between the positively charged AFO surface (provided pH < pHpzc for AFO) and the DFP anion. Table 1 Modelled reactions and rate constants for interaction between DFP and ferric iron Rate constants No.

Reactions

Reference (M-1s-1 or s-1)

1

k1 Fe(III) + Fe(III)I  → AFO + nH +

k1 = 5.0 × 106

1

k2 = 1.49 × 106

This study

k−2 = 2.81

This study

k3 = 1.16 × 108

This study

k−3 = 0.65

This study

k4 = 1.16 × 108

This study

k−4 = 3.6 × 10-3

This study

k5 = 51.7

This study

2

3

4

5

k5 AFO + DFP  → FeIII DFP

Fe(III), inorganic ferric ion; Fe(III)I, total inorganic Fe(III); AFO, amorphous ferric oxyhydroxide (ferrihydrite) and DFP, deferiprone. (1) Pham et al. 39

6



The DFP induced dissolution process is consistent with the behaviour observed in previous studies using both the siderophore desferrioxamine (DFO)40 and the bidentate ligand, DA.11 While the rate constant for this reaction is not available, it has been reported that the iron mobilization rate of DFP is 10 times greater than that of DFO.41 In comparison, based on the results obtained here at physiological concentrations, the apparent rate constant for DFP-mediated iron mobilization (i.e. kAFO+DFP = 51.7 M-1s-1 in Table 1) is over twenty times greater than that of DA (kAFO+DA = 2.43 M-1s-1 shown in SI 2).11 In addition to the effect of steric hindrance, the enhanced iron mobilization in the presence of DFP may be attributed, at least partially, to the increased proportion of mono- or even fully deprotonated phenol groups as a result of the much lower pKa1 value of DFP (pKa1 = 3.62) compared to either DA (pKa1 = 10.58; pKa2 = 12.07)42 or DFO (pKa1 = 8.32; pKa2 = 8.96 and pKa3 = 9.55).38 From this perspective, the low pKa1 value and the associated higher proportion of deprotonated functional groups present may be one of the advantages of DFP with regard to its use in mitigating iron overload with this feature potentially of importance in the selection and design of chelators in the future.

Scavenging of Fe(II) by DFP and the resultant FeIIIDFP3 and H2O2 formation

The fate of Fe(II) and its potential for oxidant production in the presence of DFP was examined by measuring the decay of Fe(II) and the formation of FeIIIDFP3 and H2O2. In general, in the absence of DFP, the oxidation of Fe(II) was relatively slow at pH 7.4 with approximately 50% of the Fe(II) remaining in solution after 25 minutes (Figure 2a). In the presence of DFP, however, the rate of oxidation of 5 µM Fe(II) increased dramatically on increasing DFP concentration (Figure 2a) with this process accompanied by the formation of both the FeIIIDFP3 complex and H2O2 (Figures 2b and c). A biphasic trend with rapid initial decay (within the first 1 – 2 minutes) followed by much slower decay was evident when the concentration ratio of [DFP]/[Fe(II)] was less than three. Consistently, the initial concentration of FeIIIDFP3 increased stoichiometrically on increasing DFP concentration. In comparison, an instantaneous loss of almost all of the Fe(II) coupled with the formation of steady-state concentrations of FeIIIDFP3 and H2O2 was evident once the ratio of [DFP]/[Fe(II)] was greater than three (Figures 2a and b).

7


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Figure 2. Oxidation of Fe(II) (panel a), formation of FeIIIDFP3 (panel b) and generation of H2O2 (panel c) followed by the addition of DFP into 5 µM Fe(II) and 0.1 M NaCl containing solutions at pH 7.4. Error bars are standard errors from triplicate measurements and solid lines represent the model fit. Note that the symbols for 20 µM and 50 µM DFP are overlapping.

It is recognized that the use of a high Fe(III) affinity chelator results in the removal of Fe(II) from solution43 with this removal facilitated by a decrease in the reduction potential of the dominant iron species present as a result of the formation of weak ferrous complexes. For DFP, even though it has been proposed that three different Fe(II) complexes can be formed in a manner similar to that for Fe(III), the instantaneous dissociation of the bis- and tris-complexes coupled with the likely extremely reducing nature of these species renders the mono-complex, FeIIDFP, the only effective species in binding free Fe(II).29, 30 The reduction potential EFe0 III DFP → FeII DFP can be roughly estimated by applying the Nernst equation as follows: 0 Fe 3+ + e − → Fe 2 + , EFe , ∆G 0 = − FEFe0 3+ → Fe 2+ 3+ → Fe 2+

(2)

K

Fe 2+ + DFP → Fe II DFP, K FeII DFP , ∆G 0 = − 2.303RT log10FeIIDFP K

Fe III DFP → Fe3+ + DFP, 1/K FeIII DFP , ∆G 0 = + 2.303RT log10Fe

(3)

III DFP

Fe III DFP + e − → Fe II DFP, EFe0 III DFP → FeII DFP , ∆G 0 = − FEFe0 III DFP → FeII DFP

(4)

(5)

Thus,

0 EFe0 III DFP → FeII DFP = E(Fe 3+

0 where E(Fe 3+

→ Fe 2+ )

→ Fe2+ )

 K III − 0.059log10  Fe DFP  K II  Fe DFP

  

(6)

= 0.77 V 44 is the standard reduction potential of the Fe3+/Fe2+ couple and K FeIII DFP and

K FeII DFP are the stability constants of the complexes FeIIIDFP and FeIIDFP shown in Table S4 in section SI 3.

Therefore, EFe0 III DFP → FeII DFP is calculated to be 0.19 V. The predicted increase in the reduction potential as a 9


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result of the change in speciation is in accordance with results of a previous study in which a positive shift in reduction potential was observed on decrease in pH.45 From the calculations above, it is reasonable to deduce that the initial rapid decay of Fe(II) in the presence of low DFP concentrations may arise from the formation of the weak FeIIDFP complex followed by its oxidation and formation of the FeIIIDFP3 complex. This phenomenon is consistent with the previously reported extremely active nature of the weak FeIIDFP complex.29 Once the majority of DFP is retained in the FeIIIDFP3 complex, the dominant ferrous species present in solution converts from FeIIDFP to the unbound Fe(II). As such, even though the dominant ferric iron species would be DFP bound Fe(III) as a result of the instantaneous complexation of the oxidized inorganic Fe(III), the reduction potential controlling the oxidation process would increases dramatically from 0.19 V to 0.77 V, which renders the Fe(II) species less active toward oxygenation after the initial rapid oxidation step. While not particularly significant, a slight decrease in the concentration of FeIIIDFP3 over time was evident in the presence of low DFP concentrations (2, 5 and 10 µM) (Figure 2b). This observation indicates that iron may be gradually released as a result of either the dynamic dissociation of the complex or internal ligand-to-metal charge transfer (LMCT) (possibilities are discussed further in the section below). Theoretically, hydrolysis and precipitation of Fe(III) would be expected to accompany the dissociation process as a result of ineffective competition for the vacant coordination sites of Fe(III) at low concentrations of DFP.

The oxidation of Fe(II) is accompanied by the generation of H2O2 either through two consecutive oneelectron transfers between iron and O2 or as a result of the disproportionation of the one electron transfer 46, 47 product, O •− 2 .

Fe(II) + O 2 → Fe(III) + O •− 2

(7)

Fe(II) + O •− → Fe(III) + H 2 O 2 2

(8)

Fe II DFP + O 2 → Fe III DFP + O •− 2

(9)

10



O •− + O •− → H 2O2 + O2 2 2

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(10)

As shown in Figure 2c, significant steady state concentrations of H2O2 formed almost instantaneously when the concentration ratio of [DFP]/[Fe(II)] was greater than 3 while a low but gradual formation of H2O2 was evident at lower [DFP]/[Fe(II)] ratios. The gradual production of H2O2 in the presence of low DFP concentrations suggests that a balance between the oxygenation and peroxidation of ferrous iron may exist. According to the rate law for the production of • OH shown in Eq (11), d [ • OH] = k[Fe(II)]tot [H 2 O 2 ] dt

(11)

• the rate of generation of strongly oxidizing OH is dependent on the concentration of Fe(II) (in both the free

and DFP bound forms). Even though it has been reported that the rate constant k of the peroxidation process is positively related to DFP concentration,30 the biphasic depletion of iron coupled with the extremely low H2O2 concentrations at low DFP concentrations indicate that peroxidation of ferrous iron may be important under these conditions. Conversely, the instantaneous depletion of all the Fe(II) coupled with the concomitant formation of the stable FeIIIDFP3 complex and the extensive accumulation of H2O2 shown in Figure 2 suggests that the generation of • OH is very unlikely at high concentration ratios of [DFP]/[Fe(II)]. As such, taking into account that in vivo peak concentrations of DFP of 10 – 50 µM are expected as a result of typical DFP dosages (25 – 100 mg/kg/day)23, 48 and the extremely low free Fe(II) concentrations expected in a mature iron homeostatic system, the application of DFP should result in the effective scavenging of • Fe(II) with the likelihood of formation of toxic OH through this process being negligible. Reactions

associated with the effect of DFP on the transformation of Fe(II) are provided in Table 2 below.

11


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Table 2 Reactions involved in the interaction between Fe(II) and DFP Rate constants No.

Reactions

Reference (M-1s-1 or s-1) k6 = 0.77

2

k−6 = 1.50 × 108

3

k7 = 1 × 107

3

k8 = 1.33 × 104

4

6

7

k7 Fe(II) + O •−  → Fe(III) + H 2 O 2 2

8

k8 Fe(II) + H 2 O 2  → Fe(III) +

9

k9 O •− + O •−  → H2O2 + O2 2 2

k9 = 1.9 × 105

5

10

k10 AFO + O •−  → AFO + Fe(II) + O 2 2

k10 = 3.70 × 105

2

k11 = 6.76 × 106

This study

k−11 = 8.0 × 103

6

k12 = 6.7 × 104

This study

k−12 = 1.5 × 108

3

k13 = 6.8 × 104

7

•

OH + OH -

11

12

13

k13 Fe II DFP + H 2 O 2  → Fe III DFP +

•

OH + OH -

Fe(III), inorganic ferric ion; AFO, amorphous ferric oxyhydroxide (ferrihydrite); Fe(II), inorganic ferrous • ion; DFP, deferiprone; O•− 2 , superoxide; H2O2, peroxide and OH , hydroxyl radicals

(2) Sun et al. 11 (3) Rush and Bielski 49 (4) González-Davila et al. 50 (5) Zafiriou 51 (6) Merkofer et al. 29 and (7) Merkofer et al. 30

Removal of DA bound Fe(III) by DFP

To mimic in vivo conditions, we have investigated the effectiveness of DFP in the scavenging of DA bound iron within the peak concentration range of DFP mentioned above. In order to avoid the complexity arising from the oxidative transformation of DA in the interpretation of the results, the chelation of DA bound iron by DFP was examined in the absence of O2. Even though it has been reported that the presence of metals such as Mn3+ and Fe3+ can result in the anaerobic oxidation of DA,52, 12


53

the effect of Fe(III) on the


transformation of DA, especially in the presence of DFP, is expected to be negligible. Compared with Mn3+, the 3d orbitals of Fe(III) are half filled with this relatively stable state rendering Fe(III) a much more inactive oxidant than its electron scarce counterpart, Mn(III). Additionally, the second order rate constant of LMCT within DA bound Fe(III) decreases dramatically from 0.23 M-1s-1 for FeIIIDA, the dominant Fe-DA complex under acidic conditions, to 7.26 × 10-5 M-1s-1 for FeIIIDA2, the dominant DA bound iron species at the physiological pH of 7.4 used in this work.11, 36 As such, compared with the instantaneous iron transformation in the presence of DFP shown below, the Fe(III) catalyzed degradation of DA in the absence of O2 is expected to be negligible. While the ligand exchange experiments were conducted in the absence of O2, the results obtained herein can be reasonably extrapolated to physiological conditions as a result of i) the negligible role played by O2 in the ligand exchange process between DA and DFP, ii) the non-participation of DFP in the generation of O•− 2 shown in SI 4 and iii) the substantially lower O2 concentration in brain tissues than that of air-saturated solutions.54, 55 As shown in Figure 3, despite the presence of considerable excess concentrations of DA (390 µM), addition of DFP into solutions containing 5 µM FeIIIDA2 resulted in instantaneous formation of the FeIIIDFP3 complex (Figure 3a) and, accordingly, a concomitant decrease in the concentrations of FeIIIDA2 (Figure 3b). Once formed, the concentrations of FeIIIDFP3 were stable over time, which indicates that the previously mentioned LMCT is very unlikely to occur within this complex.

13


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Figure 3. Formation of FeIIIDFP3 (panel a) and concomitant loss of FeIIIDA2 (panel b) followed by the addition of (◆) 0 µM, (●) 5 µM, (▲) 10 µM and (■) 50 µM DFP to deoxygenated 0.1 M NaCl solutions containing 5 µM FeIIIDA2 and 390 µM DA at pH 7.4. Error bars are standard errors from triplicate measurements and solid lines represent the model fit.

The rapid ligand exchange between DA and DFP for Fe(III) observed in Figure 3 may be attributed to one or more of the following aspects: i) the presence of vacant coordination sites on the FeIIIDA2 complex, ii) the more extensive deprotonation of phenol groups within DFP compared to that of DA as a result of the lower pKa1 values of DFP, and iii) a possible enhancement in the crystal field stabilization energy (CFSE) obtained as a result of the ligand induced rearrangement of d orbitals of the central iron atom. Compared with the reported ligand exchange rate constant for the most common NTBI in blood plasma, Fe(III)-citrate (43 M-1s1 56

), the much larger rate constant proposed herein for the DA bound iron (4.05 × 106 M-1s-1 for FeIIIDA2)

may be attributed to the accessibility of the central iron atom. As Fe(III) has a coordination number of six, the vacant coordination sites of FeIIIDA2 are readily accessed by DFP, while the polymerization of Fe(III)citrate complexes may dramatically hinder the accessibility and reduce the exchange rate between DFP and citrate. Indeed, similar effects arising from the polymerization of Fe(III)-citrate have been reported for the ineffective chelation of iron by transferrin.57, 58 In addition, the instantaneous transfer of Fe(III) to DFP 14



shown in Figure 3 indicates that, despite possessing a smaller logK value, DFP has a much greater capacity to bind Fe(III) than does DA as is evident from the pFeIII values for the Fe(III) complexes with these ligands (pFeIII of DA is 14.1 and pFeIII of DFP is 20.5).4, 59 Note that the pFeIII values presented here are –log[Fe3+] values when [Fe(III)]total = 10-6 M and [ligand]total = 10-5 M at pH 7.4.4 While the lower pKa resulting in the enhanced deprotonation of phenol groups within DFP is a clear benefit with regard to iron chelation, another advantage of DFP over DA may arise from the possible enhancement in the CFSE obtained (compared to the DA bound Fe(III)). As a result of complexation, the 3d orbitals of the central iron atom are split into two levels with one possessing an elevated energy, eg, and the other one with a lowered energy, t2g. According to ligand field theory,60, 61 in the presence of a strong field ligand, the electrons in the 3d orbitals of iron tend to occupy the t2g level initially as the splitting energy ∆ is greater than the electron pairing energy P. As such, a more stable complex would be derived from the one that has obtained the greatest CFSE via the complexation process. In general, the splitting energy ∆ can be roughly estimated from the absorbance of the complex since the energy needed for the d-d transition of the electrons is related to the peak absorbance of the complex. As evident from the shorter absorbance wavelength of the Fe(III)-DFP complexes compared to that of the Fe(III)-DA complexes, DFP should be a stronger field ligand than DA and result in the formation of a more stable complex with iron.

Alleviation in the toxicity induced by iron and DA on addition of DFP

To better understand the rectifying mechanism of DFP on iron and DA induced toxicity, the formation of H2O2 was investigated in the presence of ferric iron, DA and DFP. As shown in Figure 4, the addition of DFP into Fe(III) and DA containing solutions resulted in a significant decrease in the generation of H2O2. In the presence of 10 µM DFP, only around 100 nM H2O2 was generated at the conclusion of the two-hour experiments, while almost 600 nM H2O2 was produced in the absence of DFP.

15


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Figure 4. Formation of H2O2 in air-saturated 0.1 M NaCl solutions containing 20 µM DA and 5 µM Fe(III) at pH 7.4 in the presence of (■) No DFP, (▲) 5 µM and (◆) 10 µM. Error bars are standard errors from triplicate measurements and solid lines represent the kinetic model fit.

Obviously, the formation of the more stable FeIIIDFP3 complex via the instantaneous transformation of Fe(III) from the FeIIIDA2 species to a species bound by the strong ligand DFP would be expected to halt the previously reported iron catalyzed oxidation of DA.11 However, if the scavenging of iron loosely bound to DA is the only reason that DFP is effective in alleviating the progression of PD, the suppression of the formation of H2O2 would not be positively related to the concentrations of DFP. Under the conditions investigated herein, the addition of aqueous Fe(III) into the DA and DFP containing solutions would result in initial competition between the chelation and precipitation processes. The comparable intrinsic rate constant for

these

two

process

(

kFe(III)+DFP = 1.49 × 106 M -1s -1

, kFe(III)+DA = 4.15 × 105 M -1s -1

and

kFe(III)+Fe(III)I = 5.0 × 106 M -1s-1 shown in SI 2) coupled with the low concentrations of DFP and DA used

herein would be expected to result in most of the added iron forming AFO. Theoretically, increase in DFP concentration would dramatically reduce the concentration of precipitated iron rather than that of H2O2 since the chelation of aqueous iron is not involved in the production of H2O2 (as shown in SI 4). Once the aqueous and DA bound Fe(III) is transformed into DFP after the initial interaction, the H2O2 generated from the DA16



Page 18 of 35

mediated iron mobilization process should be identical as a result of the presence of the same amounts of DA. Thus, the obvious discrepancy in H2O2 production as a consequence of the variation in DFP concentrations suggests that the presence of DFP may result in the quenching of radicals such as DA•− and since the oxidative transformation of these radicals is closely related to H2O2 production.11 O•− 2 Thermodynamically, the feasibility of this process can be estimated by examination of the reduction potentials of different redox couples. If the reduction potential of the DFP•− / DFP couple is assumed to be similar to that of its parent compound, 4-hydroxypyridine (E0 = 1.24 V),62 theoretically, DA•− ( 63 0 = 1.35 V ) is capable of being reduced back to DA by DFP. The improved agreement between EDA •− /DA

model and experimental data on addition of this reaction coupled with the clinically observed improvement in the availability of DA for patients provided with DFP 27 indicates that this hypothesis might be reasonable. As a result of the interaction between DA•− and DFP, DFP•− would be generated. However, the formation of DFP•− as well as its dimerization product have only been detected in the absence of O2 after irradiation in the presence of an electron acceptor.32 The almost 100% in vivo recovery of DFP when the DFP glucuronide conjugate, unchanged DFP and DFP bounded metals are summed 64, 65 indicates that, despite the presence of low in vivo O2 concentrations, the dimerization product of DFP is unlikely to be formed. As such, it is reasonable to deduce that there may be an O2 mediated transformation of any DFP•− that is formed, which is consistent with observations by Timoshnikov et al. 32. Here, we hypothesise that DFP•− might be reduced by with the resultant reformation of DFP and O2. By taking into account the reduction potentials O•− 2 0 0 0 = 0.77 V ,62, 66 EDFP = 1.24 V and EDFP EDFP − /DFP q /DFP

q

/DFP

−

can be roughly estimated using Eq (12):

E = ( E1 + E 2 ) / 2

(12)

where E1 and E2 are the half-cell reduction potentials for the first and second reduction steps.67 Thus, the 0 calculated EDFP

q

/DFP

−

should be around 0.3 V. As such, theoretically, the reduction of DFP•− rather than

oxidation by O•− 2 is more thermodynamically favourable. In addition, the presence of the positive charge on the -N group within the DFP•− molecule may further favour the reduction as a result of its electron scarcity and affinity. Reactions involved in the interactions between iron, DA and DFP are shown in Table 3. 17


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Table 3 Modelled reactions and rate constants for interaction between iron, DA and DFP Rate constants No.

Reactions

Reference (M-1s-1 or s-1)

14

k14 FeIII DA2 + DFP  → FeIII DFP + 2DA

k14 = 4.05 × 106

This study

15

k15 DFP + DA•−  → DFP •− + DA

k15 = 2.94 × 106

This study

16

k16 DFP •− + O •− DFP + O 2 → 2

k16 = 8 × 106

This study

•− DA, dopamine; DA•− , semiquinone radical; O•− , deferiprone 2 , superoxide; DFP, deferiprone; DFP

semiquinone radical; FeIIIDA2, bis-complex formed during the interaction between iron and DA (details of the interaction between iron and DA can be found in SI 2).

It is well recognized that iron plays a critical role in the non-enzymatic degradation of DA35, 68 with the toxicity induced by iron closely related to the iron-catalyzed generation of DA-derived toxic quinones as well as the generation of ROS.1, 8 The essentially instantaneous transformation of Fe(III) from the commonly existing DA-bound iron in PD brains to DFP-bound iron shown in this study is presumably critical to the success of the application of DFP in attenuating the progression of PD. Note that the concentration of DAderived o-quinone (DAQ) was too low to be measured in the studies described here in view of its high reactivity. However, to visualize the effect of DFP on the production of DAQ, results of the kinetic model developed here can be used. In order to better mimic in vivo conditions, we assume that the chelation of 5 µM FeIIIDA2 by DFP occurred in the presence of a constant concentration of O2 (60 µM, the typical concentration in brain tissues)54, 55 and excess DA (390 µM). As shown in Figure S5, DAQ achieves a steady state concentration in the time period of in vivo DFP clearance (the half-life is around 47 – 134 min)64, 69 with the concentration of DAQ reduced significantly on addition of DFP (from 8 × 10-11 M in the absence of DFP to 0.25 × 10-11 M in the presence of 50 µM DFP). Despite the presence of a mature in vivo H2O2 removal system,70 the DFP induced reduction in H2O2 concentration as a result of the formation of inactive DFP bound Fe(III) or even the DFP related quenching of radicals would be expected to directly attenuate iron and

18



DA induced oxidative stress. As such, the use of DFP should result in the attenuation of both ROS, including •

OH , and DA metabolites associated with the progression of PD.71, 72 O•− 2 or even

19


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Materials and methods A complete description of the materials and methods used in this paper is presented in SI 6. All analytical grade chemicals were purchased from Sigma-Aldrich (or as otherwise stated) and were used without further refinement. All solutions were prepared using 18 MΩ.cm ultrapure Milli-Q water (MQ). All glassware was acid washed in 5% (v/v) HCl for at least one week before use. Stock solutions were kept in dark bottles and were refrigerated at 4 oC when not in use. All experiments were conducted under dark conditions and performed at a controlled room temperature of 22 ± 0.6 oC.

Experimental measurements

The concentration of Fe(II) was quantified spectrophotometrically using the modified FZ method73 in a 10 cm cuvette. A Cary 60 spectrophotometer was used for measurement of absorbance of FeIIFZ3 at 562 nm with baseline correction undertaken at 690 nm.

The H2O2 formed during the course of DA oxidation was quantified using the modified DPD method.11, 74, 75 As shown in Figure S6, the presence of DFP and iron-DFP complexes does not influence the H2O2 measurement. The influence of Fe(II) in the presence of low concentrations of DFP could be eliminated by the addition of 2,2′-bipyridyl (as shown in Figure S7).

The concentration of FeIIIDFP3 formed in air-saturated solutions was determined spectrophotometrically in a 10 cm cuvette by measuring the absorbance at 460 nm.29, 45 The measurement was performed by using a Cary 60 spectrophotometer with baseline correction undertaken at 800 nm.

20



Page 22 of 35

The concentrations of FeIIIDA2 present in the ligand exchange experiments were determined spectrophotometrically in a 10 cm cuvette by using a Cary 60 spectrophotometer under deoxygenated conditions. Briefly, the solutions were sparged for 1 h using a special gas mixture of 297 ± 6 ppm CO2 in argon (BOC) prior to the addition of DA. The solution was then bubbled for another 10 min before the addition of Fe(III). To guarantee the complete formation of the FeIIIDA2 complex, the DA and Fe(III) containing solution was bubbled for at least one hour before the addition of DFP. Sparging was continued during the course of the experiment in order to maintain deoxygenated conditions. Theoretically, the absorbance at a particular wavelength is the sum of the contributions from different species as a result of the overlap of the spectrum. In view of the coexistence of FeIIIDA2 and FeIIIDFP3, according to the method published on previous studies,76 the concentration of each species can be determined by solving a system of two linear equations: III

III

Fe DA 2 Fe DFP3 A700 = ε 700 CFeIII DA l + ε 700 CFeIII DFP l 2

III

(13)

3

III

Fe DA 2 Fe DFP3 A460 = ε 460 CFeIII DA l + ε 460 CFeIII DFP l 2

(14)

3

where A is the overall absorbance at a specific wavelength, ε i j is the molar absorptivity of component j at wavelength i, Cj is the concentration of component j and l is length of the cuvette, which is 10 cm in this study. Instead of using a peak absorbance at 580 nm,77, 78 the concentration of FeIIIDA2 was calculated by using linear regression curves at 700 nm with baseline correction at 850 nm where the interference of III

Fe DFP3 FeIIIDFP3 on FeIIIDA2 is negligible. As such, the ε 700 CFeIII DFP l term in Eq (13) can be omitted since the 3

molar absorptivity of FeIIIDFP3 at 700 nm is around 0. Accordingly, the concentration of FeIIIDFP3 in the presence of DA was calculated by using Eq (14). Speciation modeling

The concentrations of the various DFP and Fe(III)-DFP species present over a range of DFP concentrations were determined using the program Visual Minteq79 with the equilibrium reactions and stability constants

21


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provided in Table S4. Details of the distribution of DFP and Fe(III)-DFP species can be found in Figures S2 and S3. Kinetic modeling

A kinetic model was developed and used to describe the experimental data at pH 7.4 over a range of conditions using the program Kintek Explorer.80 The model developed here to describe the chelation of DA bound iron by DFP was based on a model describing interactions between iron and DA published previously (and shown in SI 2).11 Details of the model and justification for the various reactions used can be found in Tables 1 – 3 and SI 2, respectively. The sensitivity of the model to changes in individual rate constants was determined using the program Kintecus81 combined with a Visual Basic for Applications (VBA) program.

Conclusions and implications The results of this study show that, as a result of the low pKa1 value and resultant high proportion of deprotonated phenol groups present, DFP can efficiently scavenge both aqueous and loosely DA bound Fe(III) with resultant formation of the non-ionic and inactive FeIIIDFP3 species. The observation of a gradual increase in the concentration of FeIIIDFP3 species suggests that DFP may be capable of slow removal of iron from iron oxide precipitates and possibly even the iron storage protein, ferritin. The instantaneous depletion of Fe(II) at concentrations typical of in vivo conditions indicates that the use of DFP could result in the prevention of generation of oxidants (such as • OH ) via Fenton processes. In addition to the reduction in iron content and the cessation of iron catalyzed oxidative transformation of DA, the results presented here (and summarized in the schematic shown in Figure 5) suggest that DFP may possibly play a role in the quenching of radicals, such as DA•− , which may be critical in the prevention of the generation of DA-derived toxic quinones and resultant attenuation of the progression of PD.

22



Figure 5. Schematic showing range of processes involving interaction of dopamine (DA), deferiprone (DFP) and iron (Fe). It is suggested here that the particular efficacy of DFP as an iron chelator may result from the presence of the -N group within the ring. As a consequence of the strong affinity for electrons, the presence of -N may significantly reduce the electron density on the functional phenol groups, with this property accounting for the stability of the reduced form of the ligand and resultant aversion to LMCT. As such, selection and design of chelators with strong electron withdrawing groups which act to reduce the electron density on the functional groups of the ligand may be critical in the design of chelators that will be effective in retarding the progression of iron-induced neurodegenerative diseases.

23


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Associated content Supplementary data associated with this article can be found in the online version.

Author Information Corresponding Author: *Tel.: +61 2 9385 5059, Email: [email protected] Author Contributions: All the authors were involved in the experimental design. YYS conducted the

experiments and prepared the manuscript. ANP and TDW revised the manuscript. All authors reviewed the results and approved the final version of the manuscript. Conflict of interest: The authors declare that they have no conflicts of interest with the contents of this

article.

Acknowledgements We gratefully acknowledge the China Scholarship Council and the University of New South Wales for scholarship support to Yingying Sun.

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References [1]

Spillantini, M. G., Crowther, R. A., Jakes, R., Hasegawa, M., and Goedert, M. (1998) α-Synuclein in filamentous inclusions of Lewy bodies from Parkinson’s disease and dementia with Lewy bodies, Proceedings of the National Academy of Sciences 95, 6469-6473.

[2]

Blum, D., Torch, S., Lambeng, N., Nissou, M.-F., Benabid, A.-L., Sadoul, R., and Verna, J.-M. (2001) Molecular pathways involved in the neurotoxicity of 6-OHDA, dopamine and MPTP: contribution to the apoptotic theory in Parkinson's disease, Prog Neurobiol 65, 135-172.

[3]

Baudrexel, S., Nürnberger, L., Rüb, U., Seifried, C., Klein, J. C., Deller, T., Steinmetz, H., Deichmann, R., and Hilker, R. (2010) Quantitative mapping of T1 and T2* discloses nigral and brainstem pathology in early Parkinson's disease, Neuroimage 51, 512-520.

[4]

Hider, R. C., Roy, S., Ma, Y. M., Le Kong, X., and Preston, J. (2011) The potential application of iron chelators for the treatment of neurodegenerative diseases, Metallomics 3, 239-249.

[5]

Burns, R. S., Chiueh, C. C., Markey, S. P., Ebert, M. H., Jacobowitz, D. M., and Kopin, I. J. (1983) A primate model of parkinsonism: selective destruction of dopaminergic neurons in the pars compacta of the substantia nigra by N-methyl-4-phenyl-1, 2, 3, 6-tetrahydropyridine, Proceedings of the National Academy of Sciences 80, 4546-4550.

[6]

Harnois, C., and Di Paolo, T. (1990) Decreased dopamine in the retinas of patients with Parkinson's disease, Invest Ophth Vis Sci 31, 2473-2475.

[7]

Hornykiewicz, O. (1966) Dopamine (3-hydroxytyramine) and brain function, Pharmacol Rev 18, 925964.

[8]

Segura‐Aguilar, J., Paris, I., Muñoz, P., Ferrari, E., Zecca, L., and Zucca, F. A. (2014) Protective and toxic roles of dopamine in Parkinson's disease, J Neurochem 129, 898-915.

[9]

Berman, S. B., and Hastings, T. G. (1999) Dopamine oxidation alters mitochondrial respiration and induces permeability transition in brain mitochondria, J Neurochem 73, 1127-1137. 25


Page 26 of 35

Page 27 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


[10] Jiang, D., Shi, S., Zhang, L., Liu, L., Ding, B., Zhao, B., Yagnik, G., and Zhou, F. (2013) Inhibition of the Fe (III)-catalyzed dopamine oxidation by ATP and its relevance to oxidative stress in Parkinson’s disease, ACS Chem Neurosci 4, 1305-1313. [11] Sun, Y., Pham, A. N., and Waite, T. D. (2016) Elucidation of the interplay between Fe (II), Fe (III), and dopamine with relevance to iron solubilization and reactive oxygen species generation by catecholamines, J Neurochem 137, 955-968. [12] Fahn, S., and Cohen, G. (1992) The oxidant stress hypothesis in Parkinson's disease: evidence supporting it, Ann Neurol 32, 804-812. [13] Olivieri, N. F., and Brittenham, G. M. (1997) Iron-chelating therapy and the treatment of thalassemia, Blood 89, 739-761. [14] Gaeta, A., and Hider, R. C. (2005) The crucial role of metal ions in neurodegeneration: the basis for a promising therapeutic strategy, Brit J Pharmacol 146, 1041-1059. [15] Kaur, D., Yantiri, F., Rajagopalan, S., Kumar, J., Mo, J. Q., Boonplueang, R., Viswanath, V., Jacobs, R., Yang, L., and Beal, M. F. (2003) Genetic or pharmacological iron chelation prevents MPTPinduced neurotoxicity in vivo: a novel therapy for Parkinson's disease, Neuron 37, 899-909. [16] Bernhardt, P. V. (2007) Coordination chemistry and biology of chelators for the treatment of iron overload disorders, Dalton Transactions, 3214-3220. [17] Cohen, A. R., Galanello, R., Piga, A., Dipalma, A., Vullo, C., and Tricta, F. (2000) Safety profile of the oral iron chelator deferiprone: a multicentre study, Brit J Haematol 108, 305-312. [18] Mounsey, R. B., and Teismann, P. (2012) Chelators in the treatment of iron accumulation in Parkinson's disease, International journal of cell biology 2012, 1-12. [19] Sohn, Y.-S., Breuer, W., Munnich, A., and Cabantchik, Z. I. (2008) Redistribution of accumulated cell iron: a modality of chelation with therapeutic implications, Blood 111, 1690-1699.

26



[20] Liu, Z. D., and Hider, R. C. (2002) Design of iron chelators with therapeutic application, Coordin Chem Rev 232, 151-171. [21] Evans, R. W., Kong, X., and Hider, R. C. (2012) Iron mobilization from transferrin by therapeutic iron chelating agents, Biochimica et Biophysica Acta (BBA)-General Subjects 1820, 282-290. [22] Esposito, B. P., Breuer, W., Sirankapracha, P., Pootrakul, P., Hershko, C., and Cabantchik, Z. I. (2003) Labile plasma iron in iron overload: redox activity and susceptibility to chelation, Blood 102, 26702677. [23] Hoffbrand, A. V., Cohen, A., and Hershko, C. (2003) Role of deferiprone in chelation therapy for transfusional iron overload, Blood 102, 17-24. [24] Brittenham, G. M., Cohen, A. R., McLaren, C. E., Martin, M. B., Griffith, P. M., Nienhuis, A. W., Young, N. S., Allen, C. J., Farrell, D. E., and Harris, J. W. (1993) Hepatic iron stores and plasma ferritin concentration in patients with sickle cell anemia and thalassemia major, Am J Hematol 42, 8185. [25]

Balfour, J. A. B., and Foster, R. H. (1999) Deferiprone, Drugs 58, 553-578.

[26] Dexter, D. T., Statton, S. A., Whitmore, C., Freinbichler, W., Weinberger, P., Tipton, K. F., Della Corte, L., Ward, R. J., and Crichton, R. R. (2011) Clinically available iron chelators induce neuroprotection in the 6-OHDA model of Parkinson’s disease after peripheral administration, J Neural Transm 118, 223-231. [27] Devos, D., Moreau, C., Kluza, J., Laloux, C., Petrault, M., Devedjian, J.-C., Ryckewaert, G., Garcon, G., Rouaix, N., and Jissendi, P. (2012) Disease Modifying Strategy Based upon Iron Chelation in Parkinson's Disease: A Translational Study (P02. 240), Neurology 78, P02. 240-P202. 240. [28] Aguirre, P., Mena, N. P., Carrasco, C. M., Muñoz, Y., Pérez-Henríquez, P., Morales, R. A., Cassels, B. K., Méndez-Gálvez, C., García-Beltrán, O., and González-Billault, C. (2015) Iron chelators and

27


Page 28 of 35

Page 29 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


antioxidants regenerate neuritic tree and nigrostriatal fibers of MPP+/MPTP-lesioned dopaminergic neurons, PloS one 10, e0144848. [29] Merkofer, M., Domazou, A., Nauser, T., and Koppenol, W. H. (2006) Dissociation of CP20 from iron (II)(CP20) 3: A pulse radiolysis study, Eur J Inorg Chem 2006, 671-675. [30] Merkofer, M., Kissner, R., Hider, R. C., Brunk, U. T., and Koppenol, W. H. (2006) Fenton chemistry and iron chelation under physiologically relevant conditions: electrochemistry and kinetics, Chem Res Toxicol 19, 1263-1269. [31] Devanur, L. D., Neubert, H., and Hider, R. C. (2008) The Fenton activity of iron (III) in the presence of deferiprone, J Pharm Sci 97, 1454-1467. [32] Timoshnikov, V., Klimentiev, V., Polyakov, N., and Kontoghiorghes, G. (2014) Photoinduced transformation of iron chelator deferiprone: Possible implications in drug metabolism and toxicity, Journal of Photochemistry and Photobiology A: Chemistry 289, 14-21. [33] Hider, R., Silva, A., Podinovskaia, M., and Ma, Y. (2010) Monitoring the efficiency of iron chelation therapy: the potential of nontransferrin‐bound iron, Ann Ny Acad Sci 1202, 94-99. [34] Volpicelli-Daley, L. A., Luk, K. C., Patel, T. P., Tanik, S. A., Riddle, D. M., Stieber, A., Meaney, D. F., Trojanowski, J. Q., and Lee, V. M.-Y. (2011) Exogenous α-synuclein fibrils induce Lewy body pathology leading to synaptic dysfunction and neuron death, Neuron 72, 57-71. [35] Hare, D. J., and Double, K. L. (2016) Iron and dopamine: a toxic couple, Brain, aww022. [36] El-Avaan, U., Herlinger, E., Jameson, R., and Linert, W. (1997) Anaerobic oxidation of dopamine by iron (III), J Chem Soc, Dalton Trans 16, 2813-2818. [37] Dobbin, P. S., Hider, R. C., Hall, A. D., Taylor, P. D., Sarpong, P., Porter, J. B., Xiao, G., and van der Helm, D. (1993) Synthesis, physicochemical properties, and biological evaluation of N-substituted 2alkyl-3-hydroxy-4 (1H)-pyridinones: orally active iron chelators with clinical potential, J Med Chem 36, 2448-2458. 28



[38] Motekaitis, R. J., and Martell, A. E. (1991) Stabilities of the iron (III) chelates of 1, 2-dimethyl-3hydroxy-4-pyridinone and related ligands, Inorg Chim Acta 183, 71-80. [39] Pham, A. N., Rose, A. L., Feitz, A. J., and Waite, T. D. (2006) Kinetics of Fe(III) precipitation in aqueous solutions at pH 6.0–9.5 and 25°C, Geochim Cosmochim Acta 70, 640-650. [40] Kraemer, S. M. (2004) Iron oxide dissolution and solubility in the presence of siderophores, Aquat Sci 66, 3-18. [41] Piga, A., Gaglioti, C., Fogliacco, E., and Tricta, F. (2003) Comparative effects of deferiprone and deferoxamine on survival and cardiac disease in patients with thalassemia major: a retrospective analysis, Haematologica 88, 489-496. [42] Sánchez-Rivera, A., Corona-Avendano, S., Alarcón-Angeles, G., Rojas-Hernández, A., Ramırez-Silva, M., and Romero-Romo, M. (2003) Spectrophotometric study on the stability of dopamine and the determination of its acidity constants, Spectrochimica Acta Part A: Molecular and Biomolecular Spectroscopy 59, 3193-3203. [43] Harris, D. C., and Aisen, P. (1973) Facilitation of Fe (II) antoxidation by Fe (III) complexing agents, Biochimica et Biophysica Acta (BBA)-General Subjects 329, 156-158. [44] Schumb, W. C., Sherrill, M. S., and Sweetser, S. B. (1937) The measurement of the molal ferricferrous electrode potential, J Am Chem Soc 59, 2360-2365. [45] Merkofer, M., Kissner, R., Hider, R. C., and Koppenol, W. H. (2004) Redox properties of the iron complexes of orally active iron chelators CP20, CP502, CP509, and ICL670, Helv Chim Acta 87, 3021-3034. [46] Pham, A. N., and Waite, T. D. (2008) Oxygenation of Fe(II) in natural waters revisited: Kinetic modeling approaches, rate constant estimation and the importance of various reaction pathways, Geochim Cosmochim Acta 72, 3616-3630.

29


Page 30 of 35

Page 31 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


[47] Sun, Y., Pham, A. N., and Waite, T. D. (2015) Effect of chloride driven copper redox cycling on the kinetics of Fe (II) oxidation in aqueous solutions at pH 6.5–8.0, Geochim Cosmochim Acta 161, 118127. [48] Porter, J., Huehns, E., and Hider, R. C. (1989) The development of iron chelating drugs, Bailliere's clinical haematology 2, 257-292. [49] Rush, J. D., and Bielski, B. (1985) Pulse radiolytic studies of the reactions of HO2/O2-with Fe(II)/Fe(III) ions. The reactivity of HO2/O2- with ferric ions and its implication on the occurrence of the Haber-Weiss reaction, J Phys Chem 89, 5062-5066. [50] González-Davila, M., Santana-Casiano, J. M., and Millero, F. J. (2005) Oxidation of iron (II) nanomolar with H2O2 in seawater, Geochim Cosmochim Acta 69, 83-93. [51] Zafiriou, O. C. (1990) Chemistry of superoxide ion-radical (O2-) in seawater. I. pK*aswb (HOO) and uncatalyzed dismutation kinetics studied by pulse radiolysis, Mar Chem 30, 31-43. [52] Segura-Aguilar, J., and Lind, C. (1989) On the mechanism of the Mn3+-induced neurotoxicity of dopamine: Prevention of quinone-derived oxygen toxicity by DT diaphorase and superoxide dismutase, Chem-biol Interact 72, 309-324. [53] Paris, I., Martinez‐Alvarado, P., Perez‐Pastene, C., Vieira, M. N., Olea‐Azar, C., Raisman‐Vozari, R., Cardenas, S., Graumann, R., Caviedes, P., and Segura‐Aguilar, J. (2005) Monoamine transporter inhibitors and norepinephrine reduce dopamine‐dependent iron toxicity in cells derived from the substantia nigra, Journal of neurochemistry 92, 1021-1032. [54] Koppenol, W., and Butler, J. (1985) Energetics of interconversion reactions of oxyradicals, Advances in Free Radical Biology & Medicine 1, 91-131. [55] Ndubuizu, O., and LaManna, J. C. (2007) Brain tissue oxygen concentration measurements, Antioxidants & redox signaling 9, 1207-1220.

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[56] Faller, B., and Nick, H. (1994) Kinetics and mechanism of iron (III) removal from citrate by desferrioxamine B and 3-hydroxy-1, 2-dimethyl-4-pyridone, J Am Chem Soc 116, 3860-3865. [57] Bates, G. W., Billups, C., and Saltman, P. (1967) The kinetics and mechanism of iron (III) exchange between chelates and transferrin I. The complexes of citrate and nitrilotriacetic acid, J Biol Chem 242, 2810-2815. [58] Bates, G. W., Billups, C., and Saltman, P. (1967) The Kinetics and Mechanism of Iron (III) Exchange between Chelates and Transferrin II. The Presentation and Removal with Ethylenediamineteraacetate, J Biol Chem 242, 2816-2821. [59] Charkoudian, L. K., and Franz, K. J. (2006) Fe (III)-coordination properties of neuromelanin components: 5, 6-dihydroxyindole and 5, 6-dihydroxyindole-2-carboxylic acid, Inorg Chem 45, 36573664. [60] Hauser, A. (2004) Ligand field theoretical considerations, Spin Crossover in Transition Metal Compounds I, 49-58. [61] Figgis, B. N. (1987) Ligand field theory, Comprehensive Coordination Chemistry 1, 213-279. [62] Das, T. N., and Neta, P. (1998) Reduction potentials of naphthoxyl and pyridoxyl radicals in aqueous solutions, The Journal of Physical Chemistry A 102, 7081-7085. [63] Pham, A. N., and Waite, T. D. (2014) Cu(II)-catalyzed oxidation of dopamine in aqueous solutions: Mechanism and kinetics, J Inorg Biochem 137, 74-84. [64] Kontoghiorghes, G. J., Goddard, J. G., Bartlett, A. N., and Sheppard, L. (1990) Pharmacokinetic studies in humans with the oral iron chelator 1, 2‐dimethyl‐3‐hydroxypyrid‐4‐one, Clin Pharmacol Ther 48, 255-261. [65] Al‐Refaie, F., Sheppard, L., Nortey, P., Wonke, B., and Hoffbrand, A. (1995) Pharmacokinetics of the oral iron chelator deferiprone (L1) in patients with iron overload, Brit J Haematol 89, 403-408.

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Page 32 of 35

Page 33 of 35 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


[66] Yadegari, H., Jabbari, A., Heli, H., Moosavi-Movahedi, A., and Majdi, S. (2008) Electrochemistry of deferiprone as an orally active iron chelator and HIV-1 replication inhibitor and its determination, J Brazil Chem Soc 19, 1017-1022. [67] Meisel, D., and Czapski, G. (1975) One-electron transfer equilibriums and redox potentials of radicals studied by pulse radiolysis, The Journal of Physical Chemistry 79, 1503-1509. [68] Lotharius, J., and Brundin, P. (2002) Pathogenesis of Parkinson's disease: dopamine, vesicles and αsynuclein, Nat Rev Neurosci 3, 932-942. [69] Kontoghiorghes, G., Bartlett, A., Hoffbrand, A., Goddard, J., Sheppard, L., Barr, J., and Nortey, P. (1990) Long‐term trial with the oral iron chelator 1, 2‐dimethyl‐3‐hydroxypyrid‐4‐one (L1) I. Iron chelation and metabolic studies, Brit J Haematol 76, 295-300. [70] Drechsel, D. A., and Patel, M. (2010) Respiration-dependent H2O2 removal in brain mitochondria via the thioredoxin/peroxiredoxin system, J Biol Chem 285, 27850-27858. [71] Spillantini, M. G., Schmidt, M. L., Lee, V. M.-Y., Trojanowski, J. Q., Jakes, R., and Goedert, M. (1997) α-Synuclein in Lewy bodies, Nature 388, 839-840. [72] Bianco, C. L., Ridet, J., Schneider, B., Deglon, N., and Aebischer, P. (2002) α-Synucleinopathy and selective dopaminergic neuron loss in a rat lentiviral-based model of Parkinson's disease, Proceedings of the National Academy of Sciences 99, 10813-10818. [73] Garg, S., Ito, H., Rose, A. L., and Waite, T. D. (2013) Mechanism and kinetics of dark iron redox transformations in previously photolyzed acidic natural organic matter solutions, Environ Sci Technol 47, 1861-1869. [74] Bader, H., Sturzenegger, V., and Hoigne, J. (1988) Photometric method for the determination of low concentrations of hydrogen peroxide by the peroxidase catalyzed oxidation of N, N-diethyl-pphenylenediamine (DPD), Water Res 22, 1109-1115.

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[75] Voelker, B. M., and Sulzberger, B. (1996) Effects of fulvic acid on Fe (II) oxidation by hydrogen peroxide, Environ Sci Technol 30, 1106-1114. [76] Sun, Y., Pham, A. N., and Waite, T. D. (2018) Effect of release of dopamine on iron transformations and reactive oxygen species (ROS) generation under conditions typical of coastal waters, Environmental Science: Processes & Impacts, DOI: 10.1039/C7EM00497D [77] Avdeef, A., Sofen, S. R., Bregante, T. L., and Raymond, K. N. (1978) Coordination chemistry of microbial iron transport compounds. 9. Stability constants for catechol models of enterobactin, J Am Chem Soc 100, 5362-5370. [78] Sever, M. J., and Wilker, J. J. (2004) Visible absorption spectra of metal–catecholate and metal– tironate complexes, J Chem Soc, Dalton Trans, 1061-1072. [79] Gustafsson, J. P. (2005) Visual MINTEQ 3.0 user guide, Dep. of Land and Water Resour. Eng., KTH Royal Inst. of Technol., Stockholm, Sweden. [80] Johnson, K. A., Simpson, Z. B., and Blom, T. (2009) Global kinetic explorer: a new computer program for dynamic simulation and fitting of kinetic data, Anal Biochem 387, 20-29. [81] Ianni, J. C. (2003) A comparison of the Bader-Deuflhard and the Cash-Karp Runge-Kutta integrators for the GRI-MECH 3.0 model based on the chemical kinetics code Kintecus, Comput Fluid Solid Mech, 1368-1372.

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