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Molar absorption coef cients and stability constants of metal complexesof 4-(2-pyridylazo)resorcinol (PAR): Revisiting common chelating probefor the study of metalloproteins

Anna Kocył a, Adam Pomorski, Artur Krężel ⁎Laboratory of Chemical Biology, University of Wroc  ł aw, ul. Joliot-Curie 14a, 50-383 Wroc  ł aw, Poland

a b s t r a c ta r t i c l e i n f o

 Article history:Received 19 May 2015Received in revised form 25 August 2015Accepted 26 August 2015Available online 5 September 2015

Keywords:

4-(2-Pyridylazo)resorcinol (PAR)Molar absorption coef cientDissociation constantStability constantZinc protein

4-(2-Pyridylazo)resorcinol (PAR) is one of the most popular chromogenic chelator used in the determination of the concentrations of various metal ions from the  d, p and f  blocks and their af nities for metal ion-bindingbiomolecules. The most important characteristics of such a sensor are the molar absorption coef cient and themetal–ligand complex dissociation constant. However, it must be remembered that these values are dependenton the specic experimental conditions (e.g. pH, solvent components, and reactant ratios). If one uses thesevalues to process data obtained in different conditions, the  nal result can be under- or overestimated. Weaimed to establish the spectral properties and the stability of PAR and its complexes accurately with Zn2+,Cd2+, Hg2+, Co2+, Ni2+, Cu2+, Mn2+ andPb2+ at a multiplepH values. Theobtained results account forthe pres-enceof different speciesof metal–PARcomplexesin thephysiological pH range of 5 to 8 andhave beenfrequentlyneglected in previous studies. The effective molar absorption coef cient at 492 nm for the ZnH x(PAR)2 complexat pH 7.4 in buffered water solution is 71,500 M−1 cm−1, and the dissociation constant of the complex in theseconditions is 7.08 × 10−13 M2. To conrm these values and estimate the range of the dissociation constants of zinc-binding biomolecules that can be measured using PAR, we performed several titrations of zinc  nger pep-tides and zinc chelators. Taken together, our results provide the updated parameters that are applicable to anyexperiment conducted using inexpensive and commercially available PAR.

© 2015 Elsevier Inc. All rights reserved.

1. Introduction

Proteins utilize a large number of cofactors to achieve a variety of functions and structures. Critical among these cofactors are metal ions,that differ substantially from organic cofactors. For example, bioinfor-matic analyses of the human genome suggest that up to 3000 proteinsmay participate in Zn2+ binding [1]. This number corresponds to ~10%of all encoded proteins, which may additionally contain variousnumbers of zinc domains and motifs with different metal af nities [2].One of the critical steps in the metalloprotein studies is monitoring of the metallic cofactors association or dissociation with proteins. More-over, metalloproteins, particularly those that bind metal ions throughsulfur donors of cysteine residues, might be oxidized or chemicallymodied depending on the number of biological reactive species thattypically decrease metal ion-to-protein af nity [3,4]. The best examplesof such cysteine-containing proteins are the metallothioneins and zincnger domains [5–11].

The release of metal ions from proteins and signicant decreases inmetal ion-to-protein af nities are dif cult to observe in in vitro condi-tions due to the typically low concentrations of the metalloproteins

used in such experiments. Weak spectroscopic properties of some metalions, especially Zn2+, do not aid the spectrophotometric determinationof metal ions concentrations. To measure the micro- or submicromolarmetal ion concentrations precisely, specialized chelating chromophores(also termed metallochromic indicators), that change their spectral prop-erties upon metal ion binding and have appropriate af nities towardmetal ions, are strictly required [12]. In analytical chemistry, one of thebest known (from the historical perspective) metallochromic indicatorsfor the Hg2+, Pb2+, Cu

2+, Zn2+, Co2+ and Ni2+ detection is dithizone(Fig. 1). It requires hydrophobic solvents and thus is not suitable formetalloprotein-based applications [13–14]. One of the most popularchelating chromophores for the bioinorganic analysis of Zn2+ is water-soluble 4-(2-pyridylazo)resorcinol (PAR) [15]; however, a numberof other compounds have been applied to measure Zn2+ , such asZincon, Eriochrome Blue SE, Eriochrome red B, Naphtylazoxine 6Sand SNAZOXS [16,17] (Fig. 1, Fig. S1, ESI). Among these compounds,PAR and Zincon are the most popular in bioanalytics. The former onehas a higher Zn2+ af nity, and its complex has a 3-fold higher molarabsorption coef cient than the latter one [18,19].

The last two decades have produced a number of useful  uorogenicchelating chromophores thatare frequently termed zinc probes such asTSQ, Fura, FluoZin, ZnAF, Zinpyr families that are used in similar applica-tions [20–24] (Fig. 1, Fig. S1, ESI). Fluorogenic probes with uorescence

 Journal of Inorganic Biochemistry 152 (2015) 82–92

⁎   Corresponding author.E-mail address: artur.krezel@uwr.edu.pl (A. Krężel).

http://dx.doi.org/10.1016/j.jinorgbio.2015.08.024

0162-0134/© 2015 Elsevier Inc. All rights reserved.

Contents lists available at  ScienceDirect

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modulated by metal ion binding are used in most applications atmuch lower concentrations than classical metallochromic indicatorsdue to the greater sensitivities of  uorescent techniques and theunique chemical properties of  uorogenic probes. The importantdrawbacks of the majority of  uorescent chromophores are limitedavailabilities from commercial suppliers and high costs. Additionally,many uorescent chromophores are soluble only in water at the lowmicromolar concentrations and precipitate when used above thatrange. Despite the many advantages of  uorescent chelating probes,particularly in molecular biology studies, classical chromophores,including PAR, are the most popular for in vitro studies due to theirprotein stability and Zn2+ release characteristics. Classical chromo-phores can be easily used at submillimolar levels and are inexpen-sive, which makes them the  rst choice probes for measuring metalion binding and release from metalloproteins or other metal-binding molecules [25–35].

The most important physicochemical parameters of a metal chelat-

ing chromophore are its molar absorption coef cient and the stabilityconstant of the formed complex. In general, thehigher molar absorptioncoef cients are associated with the lower metal ion detection limits.Molar absorption coef cients are used in metalloprotein studies dually:to establish the metal ion concentration and to quantitatively monitorcompetition with macromolecules to determine the metal bindingaf nities of the macromolecules. In both cases, anyone using themolar absorption coef cient and dissociation constant values avail-able in the literature should know the conditions for which thosevalues are valid and employ precisely the same setup in their exper-iments. The data from the literature, listed in  Table 1, indicate thatthese values have been determined in various conditions  [36–49].Application of the molar absorption coef cients determined in sig-nicantly different conditions and the use of inappropriately deter-

mined reference values obviously results in major experimentalerrors.

Here, we aimed to re-determine the molar extinction coef cients of thecommonly used chelatingchromophore PARandits complexes witheight biogenic and toxic metal ions (i.e., Zn2+, Cd2+, Hg2+, Co2+, Ni2+,Cu2+, Mn2+, and Pb2+) due to the signicant differences in the litera-ture data or lack of it (Table 1) [36–49]. Because the application of PAR for determination of the metalloproteins af nity constants requireswell-dened stability constants, we performed potentiometric andspectroscopic studies across a wide range of pH to re-determine thedissociation constants. These values were then applied to determinethe zinc af nities of several previously characterized zinc ngers andzinc chelators. In this study, we provide re-determined values of themolar absorption coef cients and stability constants that are valid for

various experimental conditions.

2. Experimental

 2.1. Materials

ZnSO4·7H2O, 4-(2-pyridylazo)resorcinol (PAR), CdCl2, NiCl2·6H2O,HgCl2, MnSO4·H2O, Pb(NO3)2, CuCl2·2H2O, Na2HPO4, Na3PO4·6H2O,sodium acetate, potassium hydrogen phthalate (KHP), 4-(2-hydroxyethyl)-1-piperazineethanesulfonic acid (HEPES), 2-(N -morpholino)ethanesulfonic acid (MES), 1,2-ethanedithiol (EDT),special quality HNO3, 1,4,8,11-tetraazacyclotetradecane (cyclam),ethylene-bis(oxyethylenenitrilo)tetraacetic acid (EGTA), tris(2-carboxyethyl)phosphine hydrochloride (TCEP), perchloric acid, aceticanhydride, thioanisole, anisole, and the standard solution of 0.1 M

Fig. 1. Examples of chelating chromophores and  uorophore: Zincon (a), Eriochrome Blue SE (b), Dithizone (c), Naphtylazoxine 6S (d), SNAZOXS (e), and TSQ (f).

 Table 1

Molar absorption coef cients of PAR –metal ion complexes reported in the literature.

Metalion

Wavelength(nm)

ε (M−1 cm−1) Experimental conditions References

Zn2+ 492 84 100 50 mM ammonia solution   [36]500 80 000 50 mM MOPS pH 7.3,

100 NaCl[37]

495 63 400 Ammonia solution, pH 8   [38]495 77 400 Borate, pH 9   [39]495 81 000 Borate, pH 8   [40]493 83 000 Carbonate, pH 9.7   [41]485 38 750 5 mM Tris–HCl, pH 8.6   [42]485 43 300 5 mM Tris–HCl, pH 7.4,

100 mM NaClO4

[43]

497 46 700 50 mM HEPES, pH 7.4, 4 MGdnHCl

[44]

500 66 000 40 mM HEPES, pH 7.0   [45]Cd2+ 494 84 400 50 mM ammonia solution   [36]

485 21 666 5 mM Tris–

HCl, pH 8.6   [45]485 32 000 5 mM Tris–HCl,pH 7.4, 0.1 M NaClO4

[43]

Hg2+ 540 20 059 pH 3.1–3.3   [46]Co2+ 508 58 600 50 mM ammonia solution   [36]

513 33 300 50 mM HEPES, pH 7.4, 4 MGdnHCl

[44]

Ni2+ 494 72 200 50 mM ammonia solution   [36]Cu2+ 496 67 000 50 mM ammonia solution   [36]

515 68 700 Borate, pH 7.96–9.40,0.1 M NaClO4

[47]

Mn2+ 496 78 000 50 mM ammonia solution   [36]490 38 300 NaOH solution, pH 10.0   [48]496 86 500 100 mM phosphate buffer,

pH 11.2[49]

500 78 000 Ammonia solution,pH 9.7–10.7

[40]

Pb2+ 518 35 900 50 mM ammonia solution   [36]

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NaOH were purchased from Sigma-Aldrich. The metal-chelating resinChelex 100® was from BioRad. Acetonitrile (ACN) and Co(NO3)2·6H2Owere fromMerck Millipore.NaCl, NaOH, HCl, aceticacid, and ethyl etherwere from Avantor Performance Materials Poland (Gliwice, Poland).N ,N -dimethylformamide (DMF), dichloromethane (DCM), 1-methyl-2-pyrrolidinone (NMP), N ,N ,N ′,N ′-tetramethyl-O-(1H -benzotriazol-1-yl)uronium hexauorophosphate (HBTU), triuoroacetic acid (TFA),N ,N -diidopropylethylamine (DIEA), piperidine, TentaGel R Ram and all

Fmoc-protected amino acids were obtained from Iris Biotech GmbH(Marktredwitz, Germany). The concentration of the stock solutions of the metal ion salts was 0.05 M, and the exact concentrations wereconrmed by representative series of ICP-MS measurements [50]. Allof the pH buffers used in this study were treated with Chelex-100resin to eliminate trace metal ion contamination   [51–53]. A PAR solution in DMSO was prepared freshly before each experiment. Itshould be noted that DMSO solutions should not be stored at a roomtemperature longer than one week due to degradation. PAR watersolutions are stable for hours and were prepared immediately beforethe measurements from DMSO stock solutions.

 2.2. Peptide synthesis

The tetracysteine motif (TC), MTF1-1 zinc nger and C@E mutant of the ZF133-11 zinc nger peptides (Table S1, ESI) were synthesized viasolid phase synthesis using a Fmoc microwave-assisted synthesizer(CEM)   [54]. The reagent excess, cleavage and purication wereperformed essentially as previously described [55–57]. The peptideswere acetylated on the N-terminus using acetic anhydride in thepresence of DIEA and cleaved from the resin with a mixture of TFA/thioanisole/EDT/anisol (90/5/3/2 v/v/v/v) over a period of 2 h followedby precipitation in cold (−20 °C) diethyl ether. The crude peptidepellets were collected by centrifugation, dried and puried via HPLC(Dionex Ultimate 3000) on Phenomenex Aeris Peptide 3.6  μ m 100 ÅC18 columns using a gradient of ACN in 0.1% TFA/water from 0 to 50%over 30 min. The puried peptides were identied by ESI massspectrometry with an API 2000 Applied Biosystems instrument. Theidentied/calculated monoisotopic molecular masses were 3137.3/

3137.5 (ZF133-11 C@E), 3600.4/3600.7 (MTF1-1), and 1356.4/1356.5(TC motif).

 2.3. Potentiometry

The protonation constants of PAR and the stability constants weremeasured at 0.1 M ionic strength (from 96 mM KNO3 in 4 mM HNO3)at 25 °C using pH-metric titration over the pH range of 2.5 to 11.0(Molspin automatic titrator, Molspin) with 0.1 M NaOH as the titrant.The exact concentration of the titrant wasestablishedvia previous titra-tions of 3 mM KHP. Due to the limited solubility of PAR in water in acidicconditions, 1% DMSO was added to all media used in the experiments.Changes in pH were monitored using a combined glass-Ag/AgClelectrode (Biotrode, Methrom) that was calibrated daily in hydrogen

concentrations achieved with HNO3 titrations [58]. Sample volumes of 2.0 ml and PAR concentrations of ca. 600  μ M were used. The datawere analyzed using the Superquad program [59]. Value of 13.83 wasused in the data analysis to account for the ionic product of water. Itrepresents a 0.1 M ionic strength in 1% of DMSO  [58,60].

 2.4. Spectroscopic titrations

The absorbance spectra were recorded on a Jasco V-650 spectropho-tometer at 25°C overthe spectral range of 300–650nm in 1.0-cm quartzcuvettes. The spectroscopic pH titrations of 20 μ M PAR were performedin0.1 M NaClO4 over the pH range of 2 to 13 using Mettler ToledoInLabSemi Micro glass electrodes to capture all deprotonation events [61].The spectroscopic titrations of 20  μ M PAR with metal ions to their

nal 10  μ M (Zn2+

, Co2+

, Cd2+

, Ni2+

, Hg2+

, and Mn2+

) or 20  μ M

(Pb2+ and Cu2+) concentrations were performed at pH of 7.0, 7.4, 8.0,9.0 and 9.9 using 50 mM HEPES or 50 mM Tris buffers with  I  = 0.1 Mfrom NaCl. The volume of each sample was 2.0 ml, and the increase involume during the titrations was included in the data processing. Thesamples were equilibratedfor 1–2 min afterthe addition of each portionof 1.0 mM metal ion solution; however, the equilibrations with Ni 2+

and Co2+ required 4–5 h of incubation prior to the measurements.

 2.5. Molar absorption coef  cients

The molar absorption coef cients of the PAR species that werepresent in highly basic pH were calculated based on the spectroscopicpH-titration data of 20 μ M PAR and the molar fraction plot from thepotentiometric data using multiple linear regression. The effectivemolar absorption coef cients were determined as the slopes of in-creases of thelinearabsorption observedafter theaddition of theappro-priate metal ion in the range of 0–10 μ M to 100 μ M PAR at pH 7.4 and25 °C (50 mM HEPES buffer,  I  = 0.1 M) and 11.0 (50 mM phosphatebuffer, I  = 0.1 M). Additional measurements of the effective molar ab-sorption coef cients were performed for Zn2+ across a wide pH rangeusing 50 mM acetate (pH 4.0, 4.5, and 5.0), 50 mM MES (pH 5.5, 6.0,and 6.5), 50 mM HEPES (pH 7.0, 7.4, and 8.0), 50 mM Tris (pH 9.0,and 9.9) with I  = 0.1 M from NaCl. The molar absorption coef cientsof the individual ZnH x(PAR) y species were calculated by combiningthe potentiometric and spectroscopic data using multiple linearregression.

 2.6. Competition studies

The competition studies of 100 μ M PAR with the zinc binding pep-tides (zinc  ngers and the TC motif) and the zinc chelators (cyclamand EGTA) were conducted in 50 mM HEPES buffer at pH 7.4 and 25°C, I  = 0.1 M from NaCl. TCEP, which is non-metal-binding protein di-sulde reducing agent, was added to the peptides containing cysteinesto protect them from oxidation to its  nal concentration of 100 μ M[62]. PAR was partially saturated by the addition of Zn2+ toits nal con-centration of 5 μ M. The 2.0 mM stock solutions of the zinc nger pep-

tides were prepared in 5 mM HCl to avoid oxidation and were thenused for the competition assays in the range of 0–12 μ M of the peptide[51,53]. The zinc chelators were prepared similarly; however the stocksolutions were prepared in 50 mM HEPES buffer. Eachsample was incu-bated for 10 h after the addition of the specied amounts of the Zn2+

binding ligand. The samples with EGTA and cyclam were incubated for2 h without the addition of TCEP. All incubations were performed inglass vials to prevent the adhesion of the metal ion PAR complexes toplastic tubes. The exact concentrations of the ZnH x(PAR)2 complexespresent in each sample after equilibration were calculated based onthe absorbances at 492 nm using the appropriate effective molar coef -cients determined in this study. All calculations of the dissociationconstants of the zinc  ngers and zinc complexes with chelators wereperformed using the effective dissociation constant of ZnH x(PAR)2

at pH 7.4 that was determined in this study. The results of all of thecalculations are presented in Tables S2-S6 (ESI).

3. Results and discussion

 3.1. Acid–base properties of PAR

The PAR molecule contains three protonating groups, i.e., twochemically inequivalent hydroxyl groups of resorcinol moieties andone pyridyl functional group (Fig. 2a). The dissociation constants of these groups were measured independently with the two techniques,i.e., spectrophotometry and potentiometry. These methods differin terms of numerical accuracy. Potentiometry is superior in this re-spect; however, due to the limitations of its application at basic pH,

potentiometry cannot be used to determine p K a  values above

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10.5–11.0 with high accuracy. Moreover, PAR has limited solubility atacidic pH required in potentiometry. To avoid the limitation of watersolubility 1% DMSO was added to all media used in the potentiometricexperiments. Spectrophotometric titration of 20   μ M PAR over thewide range of pH from 1.5 to 13.4 allowed us to characterize all of thedeprotonation events due to the signicant changes in the absorptionspectra that are associated with all acid–base equilibria (Fig. S2a, ESI).The pyridyl group exhibited the highest acidity with a pK a1 = 2.92(Table 2), which is comparable to other pyridyl derivatives with p K avalues that vary between 2 and 4 [63,64].

The most basic deprotonation occurred above pH 10.5 with apK a3 = 12.1, and this deprotonation was associated with the formationof a new characteristic band at 489 nm (Fig. S2b, ESI). Due to the inabil-ity to determine this pK a value with potentiometry, it was obtainedspectrophotometrically and then xed as a constant value in the dataprocessing using Superquad program [59]. All pK a values obtained byboth techniques are provided in Table 2. These data are highly conver-gent, which demonstrates the lack of concentration-dependence of the deprotonation processes. Comparisons of these datawith the values

that were previously determined in 50%–50% dioxane-water revealedsignicant differences, especially regarding the second deprotonation,which differed by ~ 1.5 log units (Table 2) [65,66]. Similar observationswere made regarding methanol–water mixtures in which the pK a2values varied by 1.3 log units as the concentration of methanol wasincreased from 0 to 90% [67]. This difference was due to the expositionof the 1-hydroxyl group to the solvent and the lack of additional inter-actions of this group. In contrast, the second 3-hydroxyl group partici-pates in the formation of a hydrogen bond with a nitrogen atom aspresented in Fig. 2a. This interaction increases the basicity of thehydroxyl group compared with the pK a values of the chemically equiv-alent groups of resorcinol, which are 9.2 and 10.9, respectively  [68].

The dissociationconstants of PAR, presentedin Table2, allowed us toplot the distributions of the molar fractions of all protonated and

deprotonated species acrossa wide range of pH (Fig. 3). The absorbance

changes measured at 342, 413, 458and 489nm (whichcorresponded tothe absorbance maxima of the given species) were well correlated withthe molar fraction distributions. The known molar fractions of the indi-vidual species in combination with the absorbance values at 489 nmallowed us to calculate the molar absorption coef cient (ε) of 32300 ± 300 M−1 cm−1 for the fully deprotonated L 2− species (Fig. S2b,ESI) using multiple linear regression.

 3.2. Complex formation

 3.2.1. Metal binding properties of PAR

The relatively exible PAR molecule containing two nitrogen do-nors and one oxygen donor serves as an attractive ligand for manymetal ions with various radii, oxidation states and geometries.Early literature indicates that PAR is able to associate with largenumbers of metal ions from the  d,  p  and   f  blocks of the periodictable [69–74]. The binding of metal ions to the PAR molecule is asso-ciated with the deprotonation of the phenolic group (p K a3), which

has major implication for the electronic spectra of PAR (Fig. 2a). Allmetal complexes exhibit the formation of a new band between 480to 540 nm at a neutral or basic pH (Figs. S3–S12, ESI). The basicity of the phenolic group decreases depending on themetal ion andobviouslyaffects the stabilities of the formed complexes. Because the PAR mole-cule provides only three donors, ML and ML 2 complexes are commonlyformed. Moreover, both complexes might exist in water solutions as

Fig. 2. Structuresof the fully protonatedPAR molecule: H3L + (a), MHL + (b), andMH2L 2 (c)metal complexspecies. M refers to any bivalentmetal ionsexamined in thisstudy.Greencolor

represents protons that dissociate at basic pH transforming complexes to ML, MHL 2− and ML 2

2− species with the same coordination mode.

 Table 2

Dissociation constants of PAR determined by potentiometric and spectrophotometric(UV –vis) titration at 25 °C and I  = 0.1 M and the values reported in the literature.

UV –vis Potentiometry Potentiometryc UV-visd

pK a1   2.92 ± 0.03 2.86 ± 0.03a 2.3 3.07–2.45

pK a2   5.43 ± 0.02 5.45 ± 0.03a 6.9 5.50–6.28

pK a3   12.10 ± 0.02 12.10b 12.4 12.04–12.77

Ref. This study This study   [65,66] [67]

a log β H2L = 17.55 ± 0.01, log β H3L + = 20.41 ± 0.02 with xedlog β HL= 12.1.The

values were determined in 1% DMSO water solution (see Experimental section).b The value taken from UV –vis titration.c Values determined in 50% dioxane–water mixture.d

Values determined in methanol–

water mixtures (from 0 to 90% methanol).

Fig. 3. Comparisonof thespectrophotometric pH titration of20 μ MPARin0.1NaClO4 at25°C with the distribution of the molar fractions of PAR based on potentiometry. Blue, or-ange, red and green correspond to the absorptions measured at 342, 413, 458 and

489 nm, respectively, which were the maximal absorbances of each species.

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differently protonated forms, such as MHL +, ML, MH2L 2, MHL 2− and

ML 22−, which additionally complicates the equilibria. MHL + and

MH2L 2 complexes have hydroxyl groups at position 1 (Fig. 2b, c) thatare deprotonated at basic pH (ML and ML 2   species, respectively).MHL 2

− is a chemical form with one fully deprotonated PAR molecule.It should be noted that hydroxyl group deprotonation in those com-plexes does not change the coordination mode of themetal ion. Becauseall of thespecies present in water solution exhibitspectroscopic proper-

ties, it is vitally important to know the stabilities and molar absorptioncoef cients of the species that are formed upon metal binding. The PAR molecule is frequently treated as a simple ligand that forms only theM(PAR)2 complex when it is used in excess over a metal ion. Due tothe acid–base properties of PAR, several species with various spectralproperties and protonation states are simultaneously present in awide range of pH. The use of molar absorption coef cients determinedin different conditions results in the over- or underestimation of metalion concentrations.

 3.2.2. Potentiometric and spectroscopic studies

The complexation of metal ions to chromophoric chelating com-pounds can be studied either by potentiometry or with pH-dependentand pH-constant spectroscopic studies. Potentiometric techniques re-quire relatively high concentrations of the compound. The addition of 1%DMSOtoa4mMHNO3 water solution is required for potentiometricstudies and provides a solubility of PAR in the range of ~230 μ M, whichis acceptable for this technique. It must be remembered that PAR in a

DMSO solution is not stable for more than a week and thus should befreshly prepared. The protonation constants agree well with the spec-troscopic measurements (Table 2). Unfortunately, the addition of all of themetal ions studied here at weaklyacidicpH,results in themajor pre-cipitation of the complexes, most likely as neutral MH2L 2 complexes.The usage of a high 7:1 excess of PAR over the metal, which is thelimit for this technique, also results in some precipitation in weaklyacidic conditions. This eliminates the possibility of the potentiometryuse to study metal ion complexation by PAR. In a very early study,

Fig. 4. Spectrophotometric titrations of 20  μ M PAR with Zn2+, Co2+, Cd2+, Ni2+, Hg2+, Mn2+, Pb2+, and Cu2+ at pH 7.0 (green), 7.4 (red), 8.0 (yellow), 9.0 (blue), and 9.9 (black). The

titrations were performed in acetate, MES, HEPES, and Tris buffers at 25 °C,  I  = 0.1 M from NaCl. The dashed lines denote the stoichiometric points.

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Corsini et al. applied a 50% water solution of dioxane for potentiometricstudies [65,66]. In these conditions, metal complexes are easily solublein a wide range of pH. However, such conditions are not attractive forother researchers due to the high organic solvent content, which affectboth, the stability constants and the spectroscopic properties of PAR and its complexes.

Decreasing the PAR concentration to 20  μ M is ideal for the spectro-scopic studies and prevents the precipitation of complexes in a wide

range of pH. Fig. 4 presents titrations of 20 μ M PAR with Zn

2+

, Cd

2+

,Hg2+, Co2+, Ni2+, and Mn2+ to  nal metal:ligand molar ratio of 1:2and 1:3 with Pb2+ and Cu2+ at pH 7, 7.4, 8, 8.5, and 9.0. The absorptionvalues were measured at the characteristic wavelength of each metalion as listed in Table 3. All of the metal ions exhibited a tendency tobind to PAR; however, the complexations were strictly dependent onthe pH and metal ion. In the case of Co2+ and Ni2+ the titration curveshapes were almost the same between the pH 7 and 9, which indicatesformation of highly stable MHx(PAR)2 complexes. Here, the long-termincubation with PAR was required to reach equilibrium. The Zn2+ titra-tion plot also reveals the formation of ZnH x(PAR)2 complexes; withcharacteristic inection point above pH 8.0. In contrast, the Cd2+,Hg2+ and Mn2+ titration plots were parabolic, and full saturationswere not observed even at pH 9, which might be attributable to thesignicantly lower metal-to-ligand af nities. However, Pb2+ exhibiteda tendency to primarily form 1:1 complexes, which was easily observ-able in both, the pH-constant titrations (Fig. 4) and the pH-dependenttitrations (Figs. S9 and S10, ESI). Absorbance increased linearly up tothe molar ratio of 1:1, which might be explained by the size of themetal ion and its strong tendency to from complexes with coordinationnumber of 4. Cu2+ exhibits the dual properties of forming both stoichi-ometries, i.e., CuHx(PAR)2   complexes at a neutral pH and mixedCuH x(PAR)2 and CuH xPAR complexes at a basic pH.

Figs. S3–S12 (see ESI) demonstrate the electronic spectra of 20 μ MPAR with 10 μ M and 20 μ M metal ions (for Pb2+ and Cu2+) recordedfrom pH 2.5 to 11. Similar to the experiment above, the full complexa-tions of Co2+, Ni2+, Zn2+, and Cd2+ were reached above the pH of 6,7, 9 and 11, respectively. In the case of Mn2+, no absorbance plateauwas observed. Above pH 10, the absorption at 498 nm decreased

dramatically, which most likely demonstrates the dissociation of Mn2+ due to the formation of hydroxide-complexes [75]. The titrationsof Pb2+ with PAR at reactant ratios of 1:2 and 1:1 demonstrated thatthe absorbances at the saturation points were ~0.3 and ~ 0.6, respec-tively, which conrms the formation of PbH xPAR complexes at pHabove 7–7.5. Similar to the Mn2+ titrations, slight decreases in absor-bance wereobserved abovepH 9 forthe 2:1 ratios dueto the partial dis-sociation of Pb2+ and the formation of hydroxide-complexes [76]. Thestoichiometries of the Cu2+ complexes strictly depend on the reagentmolar ratios. At the ratio of 1:1, the formation of CuH xPAR complexeswas observed above pH 7. At a metal-to-ligand ratio of 1:2, two

equilibria events were observed over the wide range of pH. Similar tothe 1:1 ratio, the   rst saturation was observed at pH ~ 7, and thesecond oneoccurred at a basic pH. Thesendings conrm the equilibriabetween CuH xPAR and CuH x(PAR)2 complexes in these conditions.

The spectroscopic results presented above allow us to state thatMH x(PAR)2 complexes of Zn

2+, Cd2+, Hg2+, Co2+, Mn2+ and Ni2+

are formed, depending on the metal ion, at a neutral and basic pH.However, 1:1 complexes might also be present in neutral or slightly

acidic conditions. Cu

2+

clearly forms both 1:1 and 1:2 complexes in areagent-ratio dependent manner. Pb2+ formed only 1:1 complexesindependent of the reagent ratio. The use of PAR at a higher metal-ionratios increases the concentration of MH x(PAR)2 complexes (with theexception of Pb2+) and is frequently used to prevent the formation of 1:1 complexes.

 3.2.3. Molar absorption coef  cients

The absorbance of a metal-chromophoric chelator complex at itscharacteristic wavelength depends on both, the molar fraction of thecomplex andits molar absorption coef cient. An accurate molar absorp-tion coef cient (ε) value can be determined at specic conditions onlywhen the fraction of the metal complex is known. Frequently, multiplemetal complex species with various spectroscopic properties are simul-taneously present or incomplete complexation of the metal ions occursin certain experimental conditions. In such cases, the molar absorptioncoef cient should refer to the total amount of metal (variouslycomplexed and free) present in the solution, and this quantity is termedas the effective molar absorption coef cient (εeff ) in this article. Usingthis value, one is able to precisely calculate the total concentration of the metal ion for any condition required by an experiment. If incom-plete complexation occurs, e.g., in acidic pH, the application of the effec-tive molar absorption coef cient does not provide information aboutthe absorption species concentration but rather about the total amountof metal. However, in most cases, chelating chromophores are used inhigh excess over metal ions, ensuring that 100% of the metal ions arecomplexed. Notably, a difference in the probe concentration or in thepH changes the εeff  value. To avoid major errors, the effective molarcoef cient should be used for the same experimental conditions.

Taking these facts into consideration, we measured the absorbancesat the maximum wavelengths in buffered solutionsat pH of 7.4 and 11.0(Fig. 5). To avoid the formation of MH xPAR complexes in the cases of Zn2+, Cd2+, Hg2+, Co2+, Ni2+, and Mn2+, high molar excesses of PAR over the metal ions were used. PAR at a concentration of 100  μ M wastitrated at a constant pH with the specied metal ion from 0 to 10 μ M.Linear   ts representing the effective molar absorption coef cientswere obtained in all cases and are presented in Table 3 along with thestandard deviations. With the exceptions of Ni2+ and Co2+, the absorp-tion coef cients were lower at pH 7.4 than at pH 11.0 due to eitherfractional saturation of PAR or the presence of differentially protonatedcomplexes with various absorptivities. As discussed above, the com-plexes of Co2+ and Ni2+ with PAR are extremely stable, and thosemetalionswerefully complexed at pH 7.4 and 11.0 to form ML 2 species

(Figs. S4 and S6, ESI). Therefore, the effective molar coef cients wereidentical for these metal ions in both conditions.

Table 1 illustrates the molar absorption coef cients obtained in dif-ferent experimental conditions, methods and determined at variouswavelengths. A prime example is the Zn2+-PAR system; the effectivemolar coef cients that have been determined for this system at variouspH are in the absorption range of 485 to 500 nm. In the literature, themost widely cited molar extinction coef cient of the Zn2+-PAR systemis 66 000 M−1 cm−1, which was determined at pH 7.0 at 500 nm byHunt et al.   [45]. Because the fractions of particular species of theZn2+-PAR system differ signicantly at neutral pH, the effective coef -cients determined at experimental pH values should be used to avoiderrors. To highlight this issue, we determined the effective molarabsorption coef cients of Zn2+ complexes across the very wide range

of pH value from 4.0 to 11.0 (Table 4,  Fig. 6). The effective molar

 Table 3

Effective molar absorption coef cients of PAR complexes with metal ions determined atthe indicated wavelengths, which correspond to the maximum absorbances. The valueswere obtained from the slopes of the titrationsof 100 μ M PAR withthe appropriate metalions over the of 0–10 μ M at pH 7.4 (50 mM HEPES buffer) or 11.0 (50 mM phosphatebuffer) at 25 °C, I  = 0.1 M from NaCl. The slope ± the SD corresponds to the εeff .

Metal ion   λmax (nm)   εeff 7.4

(×103 M−1 cm−1)εeff 11.0

(×103 M−1 cm−1)

Zn2+ 492 71.5 ± 0.4 81.5 ± 0.2Cd2+ 495 49.3 ± 0.3 74.6 ± 0.4Hg2+ 505 15.19 ± 0.09 48.9 ± 0.3Co2+ 508 51.3 ± 0.3 51.4 ± 0.2Ni2+ 494 61.1 ± 0.4 61.2 ± 0.5Cu2+ 508 44.6 ± 0.2 59.3 ± 0.1Mn2+ 498 7.71 ± 0.06 67.2 ± 0.3Pb2+ 520 30.5 ± 0.3 36.0 ± 0.1a

a The value was determined in non buffered conditions due to the precipitation of 

Pb3(PO4)2 in phosphate buffer. The pH value was adjusted manually using 2 M NaOH.

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coef cients (at 492 nm) for the pH of 7.0, 7.4, and 8.0 were 60 800, 71500, and 76 800 M−1 cm−1, respectively. These values differ fromthosedeterminedat 500 nm(Table 4). Plotting the determined effectivemolar coef cients as a function of pH resulted in a Zn2+ complexation

isotherm curve similar to that obtained from spectrophotometricpH-titration (Fig. 7 and the inset of  Fig. 6). This  nding conrms thestatement that the changes in the molar fractions of ZnH2L 2, ZnHL 2

−

and ZnL 22− (when PAR is used in excess over the metal ion) from pH 5

to 9 critically affect the average absorbance at 492nm andconsequentlyalter the effective molar coef cients. Therefore, it is imperative to usethe appropriate coef cient values that are determined in the sameconditions as the experiment of interest.

 3.2.4. Stability constants of Zn 2+ complexes with PAR

PARis widelyused for determinationsof Zn2+ concentrations, its re-lease from metalloproteins or determinations of the stability constantsof proteins or other colorless zinc complexes. For the  rst two applica-

tions, the molar absorption coef cient is required; additionally, for thelatter application, the stability constants of the PAR complexes are alsocritical. There are two reports in the literature from Tanaka et al. andPollak et al. on the stability constants of Zn2+–PAR complexes [18,77].The protonation and stability constant values were determined spectro-photometrically and are presented in Table 5 as cumulative constants

Fig. 5. Determinationof effectivemolar absorption coef cientsoftheZn2+, Cd2+, Hg2+, Co2+, Ni2+, Cu2+, Mn2+ andPb2+ complexeswithPARat pH7.4 (red circles,50 mMHEPES, I =0.1M)or pH 11.0 (black circles, 50 mM phosphate buffer,  I  = 0.1 M).

 Table 4

Effective molar absorption coef cients of the Zn2+–PAR system determined at 492and 500 nm across a wide range of pH. The values were obtained from the slopes of thetitrations of 100 μ M PAR with 0–10  μ M Zn2+ in 50 mM acetate, MES, HEPES, Tris or

phosphate buffersat 25 °C, I = 0.1M fromNaCl. Theslope± theSD corresponds tothe εeff (×103 M−1 cm−1).

pH   εeff  (×103 M−1 cm−1)

492 nm 500 nm

4.0 2.80 ± 0.03 3.40 + 0.034.5 8.05 ± 0.04 9.41 ± 0.045.0 16.1 ± 0.1 17.7 ± 0.15.5 24.7 ± 0.1 26.3 ± 0.16.0 37. 7 ± 0.1 37.9 ± 0.26.5 49.9 ± 0.4 49.8 ± 0.47.0 60.8 ± 0.5 59.8 ± 0.37.4 71.5 ± 0.4 70.1 ± 0.48.0 76.8 ± 0.4 75.0 ± 0.39.0 81.1 ± 0.4 78.5 ± 0.29.9 81.4 ± 0.3 78.9 ± 0.311.0 81.5 ± 0.2 79.0 ± 0.2

Fig. 6. Determination of effective molar absorption coef cients of the Zn2+–PAR system.PAR (100 μ M) was titrated with Zn2+ over the range of 0–10 μ M at various pH values.

The inset demonstrates the pH-dependence of the effective molar absorption coef cients.

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( β ijk). Using our protonation constants that were determined potentio-metrically and spectrophotometrically (Fig. 7), we were able to recalcu-

late these values. The results are comparable with those reported byTanaka et al. and Pollak et al.; however, the values differ signicantlyfor the ML 2 complexes (Table 5).

Because the ZnH x(PAR) y   system includes ZnHL +, ZnL, ZnH2L 2,

ZnHL 2− and ZnL 2

2− complexes that are present across a wide range of pH, it was not possible to create experimental conditions that couldmaintain only one species with the exception of conditions that in-volved high pH andhigh molar excesses over Zn2+. Themost frequentlyused  K d value for the   “Zn(PAR)2” complex at pH 7.0 presented in thereport of Hunt et al. was calculated by an unknown method based onthe numerical data from the reports of Tanaka et al. and Pollak et al.[18,45,77]. Fig. 7a demonstrates the molar fraction distributions of theZnH x(PAR) y  complexes (10  μ M Zn

2+ and 20  μ M PAR) over a widerange of pH. This gure shows that at neutral pH, there are several spe-

cies that differ not only in stoichiometry but also in spectral propertiesand stability. The overlap of the absorbance increases at 492 nm,which demonstratesthat ZnL, ZnHL 2

− andZnL 22− signicantlycontribute

to the average absorbance at this wavelength. For the 100  μ M:10 μ M

ratio of PAR over Zn2+ (Fig. 7b) at pH 7.4, the percentages of the ZnL,ZnH2L 2, ZnHL 2−, and ZnL 22− complexes are 5.2, 5.8, 52.1, and 36.9%, re-spectively. For this ratio or higher ones, the amount of ZnL complexcan be neglected, which signicantly simplies all calculations withouta loss in accuracy (data not shown). Providing the simple effective K d of a particular complex is notuseful because all of the differently protonat-ed ZnH x(PAR)2 complexes are simultaneously present across a widerange of pH. Therefore, the most convenient method for describing thestability of the system is to provide the effective dissociation constantof ZnH x(PAR)2, which is the sum of all of the 1:2 complexes at a givenpH. Table 6 includes the effective dissociationconstants that were calcu-lated forseveral pH values from 7.0 to 9.0. These values are based on thecorrected stability values and molar fraction speciations with high PAR excesses over Zn2+ presented in this study. The effective dissociationconstants can be used directly for any competition experiment at a

specied pH value when PAR is used in high excess (at least 10) overthe total Zn2+.

 3.3. Competitive studies: determinations of metalloprotein af  nities

 3.3.1. General considerations

The determination of metal ion-to-protein af nities is broadly inter-esting to many researchers who specialize in chemistry, biochemistry,biophysics and molecular biology. Many methods and techniquescan be applied for direct determination of the stability constants of metalloprotein complexes [2,25–35,54]. If the dissociation constantis below theconcentration of the metal-binding protein used in a partic-ular technique, indirect methods, such as competition with other metal

ions or other ligands, should be applied[51,54,78,79]. For indirect meth-od, the exact stabilityof the metal ion relativeto the competitor must beknown along with the spectral properties at a given experimentalconditions.

Fig. 7. Speciation plots of the Zn2+–PAR complexes calculated using thecorrectedproton-ation andZn2+ stability constants presented in thisstudy.The molar fraction distributionsare plottedfor themolar ratios of(a) 10μ M Zn2+:20μ MPARand(b)10μ M Zn2+:100 μ M.The circles included in the top plot demonstrate the absorptions of the Zn2+-PAR com-plexes at 407(green) and492 nm (cyan) recordedwith the samereactantconcentrationsas a function of pH. The dotted line indicates the pH of 7.4.

 Table 5

Stability constants and molar absorption coef cients of the Zn-PAR complexes.

Species   β ijka

Tanaka et al. β ijk

a

Pollak et al. β ijkThis study

ε  (λ) (×103 M−1 cm−1)Tanaka et al.

ε  (λ) (×103 M−1 cm−1)Pollak et al.

ε  (λ) (×103 M−1 cm−1)This study

ZnHL + 17.8 16.7 17.6 10.7(490) 10.7(490) 14.0 ± 0.4(492)

ZnL 11.9 11.5 11.7 31.8(490) 28.5(495) 48.0 ± 0.6(492)

ZnH2L 2   36.2 31.99 35.15 17.5(495) 16.8(495) 19.0 ± 0.3(492)

ZnHL 2− 29.75 26.29 28.7 67.4(495) 70.9(490) 75.5 ± 0.5(492)

ZnL 22− 22.2 20.5 21.15 95.8(495) 92.9(490) 81.3 ± 0.2(492)

a

 β  ZnH xL  y = [ZnH xL  y] / ([Zn2+

][H] x

[L] y

), where [L] is the concentration of fully deprotonated PAR (L 2−

).

 Table 6

Effective dissociation constants (K deff )a oftheZnH x(PAR)2 speciesat variouspH values from

7 to 9. The values were calculated based on the protonation and corrected Zn2+ stabilityconstants reported in this study.b

pH   −log K deff 

7.0 11.677.2 11.897.3 12.027.4 12.15

7.5 12.297.6 12.437.8 12.758.0 13.088.2 13.448.4 13.818.6 14.198.8 14.579.0 14.96

a K deff =[Zn2+][L]2 / [Zn(L)2], where [L]is theconcentrationof unboundPAR atdifferent

protonation states (i.e., the of the H3L +, H2L, HL 

− and L 2− species); [Zn(L)2] refers to thesum of all protonated and deprotonated Zn(PAR)2 complexes (ZnH2L 2, ZnHL 2 and ZnL 2).

b To avoid the presence of ZnH xPAR complexes, a twenty-fold excess of PAR (100μ M) over Zn2+ (5 μ M) was used in the calculations of  K d

eff .

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PAR is relatively often used for indirect measurements of Zn2+-to-protein af nities [51,53]. These measurements can be performed eitherby metalloprotein equilibration with PAR (Zn2+ transfer from protein toPAR) or by the titration of partially Zn2+-saturated PAR with apoprotein(Zn2+ transfer from PARto protein).In both approaches,the knowledgeabout the Zn2+-to-PAR af nity and the spectral properties is essential.However, one needs to remember that application of PAR for studyingZn2+ transfer from zinc protein can be a problematic issue, mostly

due to time required for the equilibration. In many cases PAR causesdissociation of Zn2+ from metalloproteins but the amount of metal re-leased is lower than one expected based on thermodynamic stabilityof protein and PAR. This is because of the slow dissociation of Zn2+

(low  koff ) from many zinc proteins. Long equilibration of cysteine-containing proteins with PAR may result in protein oxidation and itssubsequent denaturation, when non-reducing conditions are used.The application of disulde reductants, e.g. DTT (dithiothreitol) addi-tionally complicates equilibria due to their high af nity for metal ions[80]. Use of PAR at a high concentration, usually above 300  μ M, mayresult in non-specic interactions with protein, precipitation of PAR orits neutral metal complexes during experiments.

Alternatively, the application of partially saturated PAR with Zn2+

for competition with apoproteins avoids problematic issue of kinetics.Relatively high koff  of ZnH x(PAR)2 complexes and high kon of most zincproteins result in a reasonable timeframe equilibration. To avoidapoprotein oxidation during its incubation with PAR, TCEP, known asa “non-metal binding” reducing agent, shouldbe used [62]. Its negligibleaf nity toward transition metal ions makes it a great reducing agentthat does not interfere with PAR-Zn2+-protein equilibria.

It is known that stability constants of metalcomplexes with low andhigh molecularweightligands are dependenton many factorsincludingsolvent, pH value, ionic strength, temperature, etc. It is obvious, butfrequently overlooked, that af nity constant determined in one set of conditions cannot be used for the interpretation of experiments thatare run in other conditions. Besides solvent composition, pH and ionicstrength are the most neglected factors in choosing the right competitorand experimental conditions. It has been shown that a 0.5 difference inpH value mayresult in oneor even more orders of magnitude difference

in stability constants [78,81]. Similarly, differences in ionic strengthresult in major difference in stability constants. Dissociation constantof theoretical zinc complex determined at   “zero” ionic strength and0.2 M differs by ~0.5 logarithmic scale [82]. Composition of buffer isanother factor. It is well known that some buffer components, such asphosphates, have signicant af nities toward metal ions that cannotbe neglected, especially when reference complex is not very stable.Some frequently used buffers, e.g. Tris buffer also forms complexeswith metal ions, such as Cu2+ [83].

 3.3.2. Competition between ZnH  x(PAR)2 complexes with apoproteins

Our main goal was to provide spectral and stability data acrossa wide range of pH for accurate use in any experiment of interest.Table 4 shows the effective molar absorption coef cients from pH 4 to

11 as determined for conditions involving a high excess of PAR overZn2+ to simplify the system to include only the ZnH x(PAR)2 species.Table 6 shows the effective dissociation constants of ZnH x(PAR)2 frompH 7 to 9.  Table 5  presents pH-independent cumulative constantsthat may be used to calculate speciation in any condition of interest(i.e., any pH and reactant ratio condition).

To examine numerical data obtained in this study we aimed to usePAR to determine the stability constants of several chemically differentmetal ion binding molecules, including zinc nger peptides (MTF1-1,C@E ZF133-1, Table S1), zinc binding motifs used foruorescent proteinlabeling (TC) and the known zinc ion chelators, including cyclamand EGTA. These compounds were chosen because of their well-established Zn2+-to-ligand af nities  [53,54,84–86]. The previouslyuncharacterized C@E mutant of ZF133-1 was chosen to close the stabil-

ity and af nity gap between EGTA and the MTF1-1 zinc   nger. All

targets were equilibrated with partially Zn2+-saturated PAR, and theabsorbances at 492 nm were measured after equilibration in glass

vials (Fig. 8a). The changes in absorbance were used to calculate thetotal and equilibrated concentrations of the reagents at all of the exper-imental points. Because all of the experiments were performed atpH 7.4, the concentrations of the ZnH x(PAR)2 complex were calculatedusing 71 500 M−1 cm−1 as the effective molar absorption coef cient.The exchange constants (K ex) of the PAR and ligand competitions forZn2+ (Eq. (1)) were calculated according to the Eq.(2) and are present-ed in Tables S2–S6 (ESI).

ZnH x  PAR ð Þ2 þ L → 2H xPAR  þ ZnL    ð1Þ

K ex ¼   H xPAR ½ 2   ZnL ½ =   ZnH x  PAR ð Þ2

    L ½

  ð2Þ

The dissociation constants of the zinc complexes with the examinedpeptides and other ligands (K d

L ) were calculated using the effectivedissociation constant (K d

eff  = 7.08 × 10−13 M2) of ZnHx(PAR)2 fromTable 6 based on Eq. (3).

K dL  ¼ K d

eff  1=K ex   ð3Þ

All of the determined dissociation constants are presented in Table 7and were convergent well with the previously published values. Rela-tive to the dissociation constants and molar absorption coef cientsfromHunt etal., we obtained either different data or were unable to de-termine the values due to negative concentration values obtained in thecalculations. Moreover, the standard deviations obtained using our setof data were much lower than those obtained by examining the

previously published values [45].

Fig. 8. Competition between ZnH x(PAR)2 complexes with ligands that form ZnL complexstoichiometries with variousZn2+ stabilities. a) Titrationof 100μ M PAR partially saturatedbyZn2+ (5 μ M) with ligandsover a concentration range of 0–11.9 μ M,25°C,and I =0.1M.The gray, blue, dark green, green and red circles correspond to cyclam, ZF133-11 C@E,EGTA, TC motif, and MTF1-1 ZF, respectively b) Simulated titration curves of ZnH x(PAR)2complexes with ligands with different af nities for Zn2+.

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The zinc-binding models used in the PAR competitions exhibitedaf nities toward Zn2+ that ranged widely from high micromolar(cyclam, pK d

L  = 8.03) to subpicomolar (MTF1-1, pK dL  = 11.40) values

as also indicated in the theoretical curves presented in  Fig. 7b, whichwere generated using the data obtained in this study. These resultsconrm that the PAR probe is a good choice for determinations of Zn2+ af nities for metalloproteins and other Zn2+-binding compoundsin competition studies. Depending on the reactant concentrations, therange of the dissociation constants that can be determined using PAR 

varies between 10−8 and 10−12 M assuming a 1:1 stoichiometry of the zinc–protein complex.

4. Conclusions

We demonstrated that at physiological pH values, complexes of 4-(2-pyridylazo)resorcinol (PAR) with various metal ions exhibitdifferent protonation states and stoichiometries (i.e., MH x(PAR) y).Particular species have different molar absorption coef cients andcontribute to the average absorbance value. Therefore, measurementsof metal ionconcentrations require theuse of an effective molar absorp-tion coef cient that is determined in the same conditions as those of planned experiment. Because PAR is commonly used in Zn2+-relatedbiochemical studies, we have provided ready-to-use effective molar

absorption coef cients forpH rangingfrom 4 to 11. Based on thesend-ings, we have also re-determined the values of the dissociationconstants of ZnH x(PAR)2 complexes across a wide range of pH. Wehave also established the range of zinc-to-proteins af nities withinwhich PAR can easily be used in competition studies. Using six differentmodels of zinc-binding molecules, we conrmed that the dissociationconstants of the resultant Zn2+ complexes convergent well with thepreviously publish stability data. Taken together, the values of  εeff  andK d presented here allow for the precise determination of metal concen-trations and can be easily applied to metallome studies.

 Acknowledgments

This research was supported by the National Science Centre (NCN)

under OPUS grant No. 2011/01/B/ST5/00830 and in part by aPRELUDIUM grant (No. 2012/07/N/NZ1/03079). The instrumentationwas supported by the Polish Foundation for Science under Focusprojects F1/2010/P/2013, and FG1/2010. The authors thank WojciechŚmigiel, Ł ukasz Winkler, and Tomasz Kochańczyk for help with thepeptide synthesis.

 Appendix A. Supplementary data

Supplementary data to this article can be found online at  http://dx.doi.org/10.1016/j.jinorgbio.2015.08.024.

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 Table 7

Dissociation constants (K dL )ofZn2+-binding ligands used in this work determined at pH7.4 (50 mM HEPES, I = 0.1 M from NaCl) based on the competition withZnH x(PAR)2 complexes.

Thevalues werecalculatedusing the effective molar absorptioncoef cient of71 500 M−1 cm−1 and the effective dissociation constant of ZnH x(PAR)2, −logK deff = 12.15 reportedhere or

in the data reported by Hunt et al. [45].

Ligand   −logK dL  of ZnL complex calculated based on

stability constants presented in this work−logK d

L  of ZnL complex calculated basedon stability constants from refs [1–3]

−logK dL  of ZnL complex

reported in the literatureReference

Cyclam 8.03 ± 0.08 No dataa 7.97   [84,85]ZF133-11 C@E 8.44 ± 0.04 8.06 ± 0.34 No datab –EGTA 9.15 ± 0.06 9.06 ± 0.27 9.20   [86]

TC motif 10.17 ± 0.13 9.98 ± 0.35 10.11   [53]MTF1-1 ZF 11.40 ± 0.25 11.14 ± 0.61 11.44, 11.62   [54]

a It was not possible to determine the constant due to the negative concentrations obtained during the calculation.b The stability constant of the Zn2+ complex with ZF133-11 C@E is reported here for the rst time.

91 A. Kocy ł a et al. / Journal of Inorganic Biochemistry 152 (2015) 82 –92

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