Synthesis, characterization, and x-ray crystal structures of cyclam derivatives. 8. Thermodynamic and kinetic appraisal of lead(II) chelation by octadentate carbamoyl-armed macrocycles

Inorg. Chem. 2005, 44, 7895−7910

Synthesis, Characterization, and X-ray Crystal Structures of Cyclam

Derivatives. 8. Thermodynamic and Kinetic Appraisal of Lead(II)

Chelation by Octadentate Carbamoyl-Armed Macrocycles¹

Franc¸ois Cuenot, Michel Meyer, Enrique Espinosa, and Roger Guilard*

Laboratoire d’Inge´nierie Mole´culaire pour la Se´paration et les Applications des Gaz

(LIMSAG, UMR 5633 du CNRS), UniVersite´ de Bourgogne, Faculte´ des Sciences,

6 bouleVard Gabriel, 21100 Dijon, France

Received May 18, 2005

En route toward the development of hybrid organic−inorganic extracting materials incorporating lead-selective chelators

and their implementation in water purification processes, the lead(II) binding properties of three N-carbamoylmethyl-

substituted 1,4,8,11-tetraazacyclotetradecanes (cyclams) have been fully investigated by spectroscopic (IR, UV

−

1

vis, MALDI-TOF MS, H and ¹³C NMR), X-ray crystallographic, potentiometric, and kinetic methods. Solution NMR

+

studies revealed that the Pb²ion is entrapped in a molecular cage constituted by the four macrocyclic nitrogen

and four amidic oxygen atoms. Protonation and lead binding constants determined in aqueous solution were shown

to be linearly dependent, so that all three derivatives possess a similar affinity at any pH value. Thermodynamic

and kinetic parameters revealed the crucial role played by the intramolecular hydrogen bonds also evidenced in

the crystal structure of the tetraacetamide derivative L¹, which involve the lone pair of each macrocyclic tertiary

amine and one amidic hydrogen atom belonging to the appended arm. In contrast to L¹, the absence of such

intramolecular interactions for N-(dimethyl)carbamoylmethyl- and N-(diethyl)carbamoylmethyl-substituted cyclams

(L²and L³, respectively) accounts for the 2

affinities. Stopped-flow kinetic measurements enabled unraveling the formation process of the three lead(II) complexes

that proceeds in a single rate-limiting step according to the Eigen Winkler mechanism, while the apparent rate

constants were found to increase in the order L³< L²L¹as a consequence of the more acidic character of L¹.

−3 orders of magnitude enhancement of their proton and lead binding

−

,

A common proton-assisted dissociation mechanism has been found for the three lead(II) complexes, which involves

the rapid formation of a protonated, six-coordinate intermediate followed by either a unimolecular decomposition or

a bimolecular attack of a second hydronium ion.

Introduction

threats. While two centuries before our era Nicander already

recognized the toxicity of cerussite (PbCO₃),²documentation

of numerous cases of poisoning, some with fatal conse-

quences, has led to a resurgence in concern over the past

decades. The toxicology of lead is nowadays well docu-

mented,⁵although its behavior with respect to biological

materials remains yet a field of intense activity. Divalent

lead mediates its deleterious effects in part by interfering

with zinc and calcium metabolism. It binds to the active site

of numerous metalloenzymes and thus inhibits their biologi-

cal activity.⁵Replacement of the Zn²⁺center in δ-amino-

levulinate dehydratase, the first cytosolic enzyme in the haem

Historically, lead has been one among a handful of readily

available metals that shaped human evolution.^2-4Because

lead is not biodegradable it is a persistent contaminant in

the environment and remains one of the major public health

* To whom correspondence should be addressed. Phone: (33) 3 80 39

61 11. Fax: (33) 3 80 39 61 17. E-mail: Roger.Guilard@u-bourgogne.fr.

(1) For the previous paper in this series, see: Meyer, M.; Fre´mond, L.;

Espinosa, E.; Brande`s, S.; Vollmer, G. Y.; Guilard, R. New J. Chem.

2005, 29, 1121-1124.

(2) Grandjean, P. EnViron. Qual. Saf. Suppl. 1975, 2, 6-75.

(3) Wittmers, L. E.; Aufderheide, A.; Rapp, G.; Alich, A. Acc. Chem.

Res. 2002, 35, 669-675.

(4) Claudio, E. S.; Godwin, H. A.; Magyar, J. S. In Progress in Inorganic

Chemistry; Karlin, K. D., Ed.; Wiley: New York, 2003; Vol. 51, pp

1-144.

(5) Godwin, H. A. Curr. Opin. Chem. Biol. 2001, 5, 223-227.

Inorganic Chemistry, Vol. 44, No. 22, 2005 7895

10.1021/ic0508019 CCC: $30.25

Published on Web 10/01/2005

Cuenot et al.

biosynthetic pathway, provides a classical example.^6,7Lead

is particularly harmful to children and has uniquely

deleterious effects on the brain and central nervous system,

increasing the risk for neurocognitive decrements, but long-

term exposure is also responsible for kidney injury, anemia,

and infertility.⁸In the United States, lead poisoning is still

considered to be the most frequent environmentally caused

disease as a consequence of an average body burden about

100-1000 times that estimated for prehistoric populations.

Blood lead levels higher than 100 µg L^-1were estimated

for 2.2% of children aged 1-5 years (434 000 children with

a 95% confidence interval from 189 000 to 846 000) ac-

cording to the National Health and Nutrition Examination

Survey conducted between 1999 and 2000.⁹Besides absorp-

tion of airborne particles through the skin and more ef-

ficiently by the lungs due to a larger surface area and a high

blood irrigation, lead uptake in humans occurs by ingestion

from the gastrointestinal tract of contaminated food and

beverages but also dust and old paint chips. The latter source

is an important and well-known cause of poisoning for young

children, which absorb and eventually accumulate lead more

efficiently (50%) than adults (10%). Since most divalent lead

compounds are only weakly soluble in vivo, only a limited

portion is absorbed, whereas the liposoluble tetraalkyl

lead(IV) derivatives easily cross the cellular membranes, and

thus most of the absorbed quantity ends up in the brain.

Because excretion via the urine and feces is extremely slow,

once lead enters the blood the majority (ca. 97%) is rapidly

taken up by the erythrocytes, which distribute lead to all soft

body tissues and vital organs before being irreversibly

deposited in bones (half-life > 20 years), where the Pb²⁺

cations displace the calcium centers in the hydroxyapatite

matrix.

BCE by the roman architect Vitruvius, who recommended

in its treaty entitled De Architectura conveying water

preferentially in terracotta pipes. In the middle of the

nineteenth century several reports highlighted the impact of

lead-contaminated water consumption on human health.¹¹

Several cases of acute lead intoxications were reported as,

for example, in the Saxon city of Dessau in 1886,¹²which

could be jugulated by treatments with limestone and soda

ash. As far as regulation is concerned, the total ban of lead

pipes imposed in 1878 in the German state of Wurtenberg

is probably one of the earliest public-health decisions in that

area.¹³On the basis of these historical considerations, it is

more than puzzling to realize that the first groundbreaking

research programs intended to gain a deeper insight of

plumbosolvency and consequently define corrosion-control

strategies were launched only in the late 1970s and early

1980s and spurred by the 1978 WHO recommendations that

set for the first time a maximal allowed concentration for

lead in drinking water.^11,13-16

Contamination occurs most exclusively through corrosion

of services pipes, joints, and solders, although other sources

might be responsible for lead levels exceeding in some

households 300-400 µg L^-1. Dissolution of the mineral

coatings mostly composed of hydrocerussite, the basic lead

carbonate Pb₃(CO₃)₂(OH)₂, deposited at the inner pipe wall

or release of scales through erosion are the principal sources

for old distribution networks.¹⁶However, “lead-free”

materials and equipment that can contain several percents

of lead, such as brass fixtures, tin solders on copper pipes,

galvanized steel, and even poly(vinyl chloride) pipes, can

significantly contribute to deterioration of the water quality

through lead leaching.¹⁷An estimated 3.3 million lead service

lines are located throughout the United States,¹⁸while a

European Community survey carried out in 1995 revealed a

much similar, if not worse, situation for West-European

countries.¹⁹For France alone the number of service pipes

was estimated to be close to 4 million (38%), corresponding

to an approximate total length of 40 000 km, with a much

higher occurrence in big cities such as Paris (70%). Fur-

thermore, about the same proportion of home properties (ca.

10 million) were still equipped with internal lead plumbing.²⁰

Although replacement of existing lead installations is the

ultimate way to ensure that the new 10 µg L^-1guideline

To avoid all risks of increasing blood lead levels in infants

and protect them against the adverse neurotoxic effects, the

World Health Organization (WHO) since 1994 has recom-

mended a maximum uptake of 25 µg per week and kilogram

of corporal mass in conjunction with a maximum acceptable

concentration in drinking water of 10 µg L^-1.¹⁰Following

these guidelines, in 1998 the Council of the European

Communities approved the 98/83/CE directive related to

the quality of consumption water, requesting the national

governments to promulgate bylaws intended to reduce the

maximal allowed lead concentration at the consumer’s tap

from 50 to 25 µg L^-1since Dec 2003 and eventually 10 µg

L^-1in 2013.

(11) Schock, M. R.; Wagner, I. In Internal Corrosion of Water Distribution

Systems; Crnkovich, S., Ed.; AWWA Research Foundation and

DVGW-Forschungstelle Cooperative Research Report: Denver, 1985;

pp 213-316.

(12) Wolffhu¨gel, G. Arbeiten am Kaiserlichen Gesundheitsamte II 1887,

484.

Health hazards related to water transport through lead

tubing were already recognized as far back as the first century

(13) Kuch, A.; Wagner, I. Water Res. 1983, 17, 1303-1307.

(14) Schock, M. R. J.sAm. Water Works Assoc. 1980, 72, 695-704.

(15) Schock, M. R.; Gardels, M. C. J.sAm. Water Works Assoc. 1983,

75, 87-91.

(6) Warren, M. J.; Cooper, J. B.; Wood, S. P.; Shoolingin-Jordan, P. M.

Trends Biochem. Sci. 1998, 23, 217-221.

(7) Pauza, N. L.; Cotti, M. J. P.; Godar, L.; de Sancovich, A. M. F.;

Sancovich, H. A. J. Inorg. Biochem. 2005, 99, 409-414.

(8) Chisolm, J. J. In Diagnosis and Treatment of Lead Poisoning; Chisolm,

J. J., Mahaffey, K. R., Aronson, A. L., Eds.; MSS Information: New

York, 1976; pp 158-183.

(9) Second National Report on Human Exposure to Environmental

Chemicals; CDC publication 02-0716; U.S. Department of Health and

Human Services, Centers for Disease Control and Prevention: Atlanta,

2003.

(16) Schock, M. R. J.sAm. Water Works Assoc. 1989, 81, 88-100.

(17) Schock, M. R.; Neff, C. H. J.sAm. Water Works Assoc. 1988, 80,

47-56.

(18) Boyd, G. R.; Tarbet, N. K.; Kirmeyer, G.; Murphy, B. M.; Serpente,

R. F.; Zammit, M. J.sAm. Water Works Assoc. 2001, 93, 74-87.

(19) Jackson, P.; van den Hoven, T.; Wagner, I.; Leroy, P. The Financial

and Economic Implications of a Change of the MAC for Lead;

European Commission Report: 1995.

(10) Guidelines for drinking water quality, 2nd ed.; World Health

Organization: Geneva, 1994; Vol. 1 (Recommendations).

(20) Leroy, P. J. Water Supply: Res. Technol., AQUA 1993, 42, 233-

238.

7896 Inorganic Chemistry, Vol. 44, No. 22, 2005

Cyclam DeriWatiWes

Chart 1

simultaneous spectrometer. MALDI-TOF mass spectra were ob-

tained on a Bruker ProFLEX III spectrometer using dithranol as

matrix. The ¹H NMR and proton-decoupled ¹³C NMR spectra were

recorded in DMSO-d₆on a DRX Avance spectrometer (Bruker)

operating at 300 MHz. Chemical shifts are expressed in parts per

million relative to the residual solvent peak.²⁶Infrared absorption

spectra were measured as KBr pellets on a IFS 66v (Bruker) Fourier

transform spectrophotometer from 4000 to 400 cm^-1at 2 cm^-1

resolution. Electronic absorption spectra were collected with a

uniform data point interval of 1 nm on a Cary 50 (Varian)

spectrophotometer equipped with a thermostated cell holder con-

nected to a RE106 (Lauda) water circulator. The optical path length

of the quartz cell (Hellma) was 1 cm.

Preparation of Compounds. Unless otherwise noted, all

chemicals and starting materials were obtained commercially and

used without further purification. Cyclam (1,4,8,11-tetraaza-

cyclotetradecane) was synthesized following the procedure of

Barefield²⁷modified by Guilard et al.,²⁸while L¹, L², and L³were

prepared according to previously patented methods.²⁵

[Pb(L¹)](ClO₄)₂‚3H₂O. Pb(ClO₄)₂‚3H₂O (460 mg, 1 mmol) and

L¹(429 mg, 1 mmol) were dissolved in hot double-deionized water

(18.2 MΩ cm). The resulting clear mixture (15 mL) was magneti-

cally stirred at 75-80 °C for 30 min. Slow evaporation at room

temperature afforded white crystals, which were isolated and dried

in vacuo. Yield: 765 mg, 86%. (+) MALDI-TOF MS: m/z ) 429.5

[L¹+ H]⁺, 635.4 [Pb(L¹) - H]⁺, 735 [Pb(L¹) + ClO₄]⁺. ¹H NMR

(300 MHz, DMSO-d₆, T ) 300 K): δ 1.57 (m, 2H, CH₂CH₂CH₂),

value will not be exceeded, retrofitting remains an expensive

and long-term task. This prompted us to explore the

feasibility of a new concept for a lead-uptake cartridge that

could be implemented directly at the consumer’s faucet. The

principle relies on solid/liquid extraction using a selective

macrocyclic lead-sequestering agent covalently attached onto

the surface of silica gel.²¹To successfully remove lead in

flowing tap water, the ideal chelator should fulfill a number

of criteria including the fast formation of stable lead

complexes in the typical pH range of drinking water with

high selectivity toward the alkaline and earth-alkaline ele-

ments usually present at much higher concentrations. To

assess the proper ligand design and choice, we thoroughly

examined the solution and solid-state behavior of N-

substituted polyazamacrocycles bearing pendant amide groups

with respect to a range of metals commonly present in natural

waters.^22-24As a part of the characterization of tetrafunc-

tionalized cyclam derivatives that might be appropriate

models of the sequestering agents incorporated in the solid-

phase extracting systems,²⁵thermodynamic stabilities of the

lead complexes formed with one primary and two tertiary

tetraamides (Chart 1) together with their formation and

proton-assisted dissociation rates have been evaluated using

a range of physical methods including potentiometry, spec-

troscopy, and stopped-flow absorption spectrophotometry.

1.78 (m, 2H, CH₂CH₂CH₂), 2.66 (m, 4H, CH₂CH₂CH₂), 2.79 (bs,

2

8H, NCH₂CH₂N), 2.95 (m, 4H, CH₂CH₂CH₂), 3.48 (dd, J_AB

)

15.5 Hz, δν ) 23 Hz, 8H, NCH₂CO), 7.88 (s, 4H, CONH₂), 8.26

(s, 4H, CONH₂). ¹³C NMR (75 MHz, DMSO-d₆, T ) 300 K): δ

25.3 (CH₂CH₂CH₂), 55.0 (NCH₂), 59.2 (NCH₂CO + NCH₂), 175.7

²⁰⁷Pb-¹³

(s + d, J

_C) 14 Hz, CH₂CO). IR (KBr): V_max) 3500 (ν_OH),

3363 (ν_{NH asym}), 3171 (ν_{NH sym}), 2912 (ν_CH), 2867 (ν_CH), 2837 (ν_CH),

1658 (ν_CdO), 1451 (δ_CH), 1139 (ν₃ClO₄^-), 1109 (ν₃ClO₄^-), 1091

2

(ν₃ClO₄^-), 626 cm^-1(ν₄ClO₄^-). UV (H₂O): λ_max) 274 nm; ꢀ )

6850 M^-1cm^-1. Anal. Calcd (found) for C₁₈H₃₆N₈Cl₂O₁₂Pb‚

3H₂O: C, 24.33 (24.62); H, 4.76 (5.22); N, 12.61 (12.64); Pb, 23.32

(23.10).

[Pb(L²)](ClO₄)₂‚0.5H₂O. Pb(ClO₄)₂‚3H₂O (322 mg, 0.7 mmol)

and L²(378 mg, 0.7 mmol) were dissolved in a hot methanol/water

(10:1 v/v) mixture. The resulting clear solution (22 mL) was

magnetically stirred with refluxing for 30 min and then concentrated

by removing one-half of the solvent. Slow evaporation at room

temperature afforded white crystals, which were isolated and dried

in vacuo. Yield: 633 mg, 94%. (+) MALDI-TOF MS: m/z ) 540.7

[L²+ H]⁺, 746.9 [Pb(L²) - H]⁺, 846.4 [Pb(L²) + ClO₄]⁺. ¹H NMR

(300 MHz, DMSO-d₆, T ) 300 K): δ 1.52 (m, 2H, CH₂CH₂CH₂),

1.79 (m, 2H, CH₂CH₂CH₂), 2.79 (m, 16H, CH₂CH₂CH₂+ NCH₂-

CH₂N), 2.90 (s, 12H, NCH₃), 3.05 (s, 12H, NCH₃), 3.81 (dd, ²J_AB

) 18 Hz, 8H, NCH₂CO). ¹³C NMR (75 MHz, DMSO-d₆, T ) 300

K): δ 26.1 (CH₂CH₂CH₂), 35.6 (NCH₃), 36.8 (NCH₃), 55.0

Experimental Section

Safety Note. Although no problems were experienced in handling

perchlorate compounds, these salts when combined with organic

ligands are potentially explosiVe and should be manipulated with

care and used only in Very small quantities.

Physical and Spectroscopic Measurements. Microanalyses

were performed on a Fisons EA 1108 CHNS instrument. The lead

content was determined by inductively coupled plasma atomic

emission spectrometry (ICP-AES) using a Varian VISTA AX

²⁰⁷Pb-¹³

C

(ΝCH₂), 57.9 (NCH₂CO), 58.9 (NCH₂), 172.4 (s + d, J

)

18 Hz, CH₂CO). IR (KBr): V_max) 3426 (b, ν_OH), 2940 (b, ν_CH),

2915 (b, ν_CH), 2840 (b, ν_CH), 1609 (ν_CdO), 1144 (ν₃ClO₄^-), 1120

(ν₃ClO₄^-), 1089 (ν₃ClO₄^-), 626 cm^-1(ν₄ClO₄^-). UV (H₂O):

λ_max) 274 nm; ꢀ ) 8070 M^-1cm^-1. Anal. Calcd (found) for

(21) Cuenot, F.; Meyer, M.; Bucaille, A.; Guilard, R. J. Mol. Liq. 2005,

118, 89-99.

(22) Meyer, M.; Dahaoui-Gindrey, V.; Lecomte, C.; Guilard, R. Coord.

Chem. ReV. 1998, 178-180, 1313-1405.

(23) Espinosa, E.; Meyer, M.; Berard, D.; Guilard, R. Acta Crystallogr.,

Sect. C 2002, 58, m119-m121.

(26) Gottlieb, H. E.; Kotlyar, V.; Nudelman, A. J. Org. Chem. 1997, 62,

7512-7515.

(27) Barefield, E. K. Inorg. Chem. 1972, 11, 2273-2274.

(28) Guilard, R.; Meunier, I.; Jean, C.; Boisselier-Cocolios, B. U.S. Patent

5 434 262, 1995.

(24) Cuenot, F. Ph.D. Thesis, Universite´ de Bourgogne: Dijon, France,

2004.

(25) Guilard, R.; Roux-Fouillet, B.; Lagrange, G.; Meyer, M.; Bucaille,

A. T. PCT Application WO 01 46202, 2001.

Inorganic Chemistry, Vol. 44, No. 22, 2005 7897

Cuenot et al.

Table 1. X-ray Crystallographic Data for L¹‚2H₂O and

[L¹H₂](NO₃)₂‚2H₂O

Anisotropic thermal parameters were used for non-hydrogen atoms.

Relevant experimental crystal data and refinement details are

summarized in Table 1. Molecular drawings were generated with

the ORTEP-3 for Windows application.³³

L¹‚2H₂O

[L¹H₂](NO₃)₂‚2H₂O

empirical formula

fw

C

₁₈H₄₀N₈O₆

C₁₈H₄₂N₁₀O₁₂

590.62

L¹‚2H₂O. Colorless prismatic single crystals were obtained from

an aqueous solution of L¹by slow evaporation at room temperature.

A high-quality specimen of prismatic shape (0.20 × 0.09 × 0.08

mm³) was selected for the X-ray diffraction experiment carried out

at T ) 173(2) K. A total of 8845 reflections were collected in full

θ range 3.12-26.36° and merged into 2382 independent reflections

(R_int) 0.050). All hydrogen atoms were found in the Fourier

synthesis and refined with a global isotropic thermal factor. While

the H(-C) and the amidic H(-N) hydrogen atoms were placed at

calculated positions using a riding model (C-H ) 0.970 Å, N-H

) 1.009 Å), those belonging to the cocrystallized water molecules

H(-Ow) were constrained to display Ow-H distances of 0.96 Å.

One-half of the molecular unit constitutes the asymmetric unit.

464.58

space group

a, Å

b, Å

P2₁/n

P1h

10.5939(9)

5.1252(3)

22.100(1)

90

100.365(9)

90

8.5200(3)

8.8830(3)

10.6060(4)

98.904(2)

108.550(2)

111.047(2)

676.60(4)

1

c, Å

R, deg

â, deg

γ, deg

V, Å³

Z

1180.4(1)

2

T, K

173(2)

1.307

0.099

110(2)

1.450

0.121

F

_calcd, g cm^-3

µ(Mo KR), mm^-1

R indices [I > 2σ(I)]^aR₁) 0.0672; wR₂) 0.1809 R₁) 0.0406; wR₂) 0.0939

R indices (all data)^aR₁) 0.0872; wR₂) 0.1925 R₁) 0.0587; wR₂) 0.1048

^aR₁) Σ||F_o| - |F_c||/Σ|F_o|, wR₂) {Σ[w(F_o- F_c)²]/Σ[w(F_o)²]}^1/2

.

2

[L¹H₂](NO₃)₂‚2H₂O. Colorless prismatic single crystals were

obtained at room temperature by slow evaporation from an aqueous

solution of L¹containing HNO₃(pH ≈ 3). A high-quality specimen

of prismatic shape (0.24 × 0.20 × 0.20 mm³) was selected for the

X-ray diffraction experiment carried out at T ) 110(2) K. A total

number of 5365 reflections were collected in full θ range 2.13-

28.75° and merged into 3446 independent reflections (R_int) 0.031).

All hydrogen atoms were located in the Fourier synthesis. Their

positional parameters were refined along with a global isotropic

thermal factor. One-half of the molecular unit constitutes the

asymmetric unit.

C₂₆H₅₂N₈Cl₂O₁₂Pb‚0.5H₂O: C, 32.67 (32.77); H, 5.59 (5.76); N,

11.72 (11.70); Pb, 21.68 (21.62).

[Pb(L³)](ClO₄)₂. Pb(ClO₄)₂‚3H₂O (230 mg, 0.5 mmol) and L³

(326 mg, 0.5 mmol) were dissolved in hot double-deionized water

(18.2 MΩ cm). The resulting clear mixture (50 mL) was magneti-

cally stirred at 75-80 °C for 30 min. Slow evaporation at room

temperature afforded white crystals, which were isolated and dried

in vacuo. Yield: 333 mg, 62%. (+) MALDI-TOF MS: m/z ) 653.0

[L³+ H]⁺, 858.8 [Pb(L³) - H]⁺, 958.7 [Pb(L³) + ClO₄]⁺. ¹H NMR

3

(300 MHz, DMSO-d₆, T ) 300 K): δ 1.05 (t, J ) 6.9 Hz, 12H,

3

NCH₂CH₃), 1.13 (t, J ) 6.9 Hz, 12H, NCH₂CH₃), 1.53 (m, 2H,

Solution Preparations. All solutions were prepared with boiled

and argon-saturated double-deionized high-purity water (18.2 MΩ

cm) obtained from a Maxima (USF Elga) cartridge system designed

for trace analysis. The carbonate-free NaOH solution, prepared from

Merck concentrates (Titrisol), was standardized by titrating against

oven-dried (120 °C for 2 h) potassium hydrogenphthalate (Aldrich,

99.99%) and stored under an atmosphere of purified argon using

Ascarite II (Acros, 20-30 mesh) scrubbers in order to prevent

absorption of carbon dioxide. The 0.1 M HClO₄solution was

obtained by dilution of analytical-grade acid (Fischer, 70%) and

standardized against 0.1 M NaOH using the equipment described

below. Equivalent points were calculated by the second-derivative

method. The concentration of the standardized solutions cor-

responded to the average of at least five replicates and was known

with a relative precision of less than 0.15%. All ligand stock

solutions were prepared by careful weighing using a Precisa

262SMA-FR balance (precision (0.01 mg). The mother solution

(0.0325 M) of Pb(ClO₄)₂‚3H₂O (Across, >99%) was standardized

with a 0.100 M Na₂H₂EDTA solution (Titriplex III, Merck) using

xylenol orange as indicator.³⁴In all experiments the ionic strength

was maintained constant at 0.1 M by addition of the appropriate

amount of NaClO₄‚H₂O (Merck).

Potentiometric Titrations. Acid-base titrations were carried

out in a water-jacketed cell connected to a Lauda RE106 water

circulator ensuring a constant temperature of 298.2(2) K. Magneti-

cally stirred solutions were maintained under a low-pressure argon

stream to exclude CO₂from the headspace. Titrant aliquots were

delivered through a polypropylene line from a calibrated automatic

ABU901 (Radiometer-Tacussel) 10 mL piston buret. Volumes were

corrected according to a linear calibration function obtained by

weighing known quantities of water and taking into account the

CH₂CH₂CH₂), 1.86 (m, 2H, CH₂CH₂CH₂), 2.82 (m, 8H, CH₂CH₂-

CH₂+ NCH₂CH₂N), 2.94 (m, 8H, CH₂CH₂CH₂+ NCH₂CH₂N),

3

2

3.34 (m + s, J ) 6.9 Hz, NCH₂CH₃+ HDO), 3.83 (dd, J_AB

)

17.4 Hz, δν ) 11 Hz, 8H, NCH₂CO). ¹³C NMR (75 MHz, DMSO-

d₆, T ) 300 K): δ 12.5 (ΝCH₂CH₃), 13.9 (ΝCH₂CH₃), 26.0

(CH₂CH₂CH₂), 40.6 (NCH₂CH₃), 41.6 (NCH₂CH₃), 55.1 (ΝCH₂),

²⁰⁷Pb-¹³

C

57.6 (NCH₂CO), 59.3 (NCH₂), 171.5 (s + d, J

) 19 Hz,

CH₂CO). IR (KBr): V_max) 2973 (ν_CH), 2936 (ν_CH), 2870 (b, ν_CH),

1474 (δ_CH), 1608 (ν_CdO), 1089 (b, ν₃ClO₄^-), 624 cm^-1(ν₄ClO₄^-).

2

UV (H₂O): λ_max) 274 nm; ꢀ ) 8400 M^-1cm^-1. Anal. Calcd

(found) for C₃₄H₆₈N₈Cl₂O₁₂Pb: C, 38.56 (39.12); H, 6.47 (6.67);

N, 10.58 (10.78); Pb, 19.56 (19.63).

X-ray Crystallographic Data Collection. Selected crystals were

mounted with silicon grease on the tip of a glass capillary.

Diffraction data were collected on

a Nonius KappaCCD

diffractometer,²⁹equipped with a nitrogen jet stream low-temper-

ature system (Oxford Cryosystems). The X-ray source was graphite-

monochromatized Mo KR radiation (λ ) 0.71073 Å) from a sealed

tube. Lattice parameters were obtained by a least-squares fit to the

optimized setting angles of the entire set of collected reflections.

Intensity data were recorded as æ and ω scans with κ offsets. No

significant intensity decay or temperature drift was observed during

the data collections. Data reductions were done by using the

DENZO software,³⁰without applying absorption correction. The

structures were solved by direct methods using the SIR97 program.³¹

Refinements were carried out by full-matrix least squares on F²

using the SHELXL97 program³²and the complete set of reflections.

(29) COLLECT, Data Collection Software; Nonius BV: Delft, The

Netherlands, 1998.

(30) Otwinowski, Z.; Minor, W. Methods Enzymol. 1997, 276, 307-326.

(31) Altomare, A.; Burla, M. C.; Camalli, M.; Cascarano, G. L.; Giaco-

vazzo, C.; Guagliardi, A.; Moliterni, A. G. G.; Polidori, G.; Spagna,

R. J. Appl. Crystallogr. 1999, 32, 115-119.

(32) Sheldrick, G. M. SHELXL-97, Program for the Refinement of Crystal

Structures; University of Go¨ttingen: Go¨ttingen, Germany, 1997.

(33) Farrugia, L. J. J. Appl. Crystallogr. 1997, 30, 565.

(34) Me´thodes d’analyses complexome´triques aVec les Titriplex; E. Mer-

ck: Darmstadt, 1990.

7898 Inorganic Chemistry, Vol. 44, No. 22, 2005

Cyclam DeriWatiWes

buoyancy effect.³⁵A PHM240 (Radiometer-Tacussel) ionometer

was used to record at 0.1 mV resolution the electromotive force

between a high-alkalinity XG200 (Radiometer-Tacussel) glass-bulb

electrode and a XR300 (Radiometer-Tacussel) Ag/AgCl reference

electrode which was separated from the bulk of the solution by a

sintered glass bridge filled with 0.1 M NaClO₄. To avoid precipita-

tion of KClO₄at the liquid junction, the reference compartment

was filled with a saturated NaCl solution. Both instruments were

controlled by the Windows-based multitasking and interactive high-

resolution titration software HRT Acide-Base Titration developed

under the TestPoint environment by C. Mustin.³⁶Prior to each

experiment, the glass electrode was calibrated to read hydronium

ion concentrations (p[H] ) -log [H₃O⁺]) by titrating 4.000 mL of

standardized 0.1 M HClO₄diluted in 25 mL of 0.1 M NaClO₄(pK_w

) 13.78(1)) with 9.010 mL of 0.1 M standardized NaOH in 0.120

mL increments. Calibration data were processed according to the

four-parameter extended Nernst equation which takes into account

liquid junction potentials.^37,38

ionometer. Enough time was allowed after the addition of each base

increment in order to reach the equilibrium. The potential-drift

criterion was set at dE/dt < 0.1 mV min^-1. The collection of

absorption spectra was repeated with 2 min delays between two

consecutive measurements until superimposable spectra were

obtained with optical densities not exceeding 0.5 units. For each

titration point, the pH-meter readings were stable in less than 2

min, and no more than two spectral recordings were required. The

fitting procedure of the experimental data with the Specfit^41-43and

Hyperquad 2000³⁹programs has been fully described elsewhere.³⁸

Species distribution diagrams were computed with the help of the

Hyss software.⁴⁴

Kinetic Measurements. Formation and proton-assisted dissocia-

tion kinetic studies were performed under pseudo-first-order condi-

tions by using a SF-61 DX2 stopped-flow spectrophotometer from

Hi-Tech Scientific. The drive syringes, mixing chamber, and optical

cell were thermostated at 298.2(2) K by water circulating from a

RE106 (Lauda) constant-temperature bath. Due to the limited

solubility of the diperchlorato complexes in water, the total lead

concentration was kept below 5 × 10^-5M. Hence, only the ligand

concentration could be varied to ensure pseudo-first-order condi-

tions, albeit in a narrow range comprised between 0.2 and

1 × 10^-3M because of the restricted solubility of the macrocycles.

During the course of the formation reaction the p[H] was maintained

within (0.02 units by addition of the nonabsorbing and weakly

coordinating 2-(N-morpholino)ethanesulfonate buffer (MES, pK_a

) 6.27 at 298 K). To avoid ligand precipitation the buffering

substance was only introduced in the lead perchlorate solution at a

concentration of 0.05 M (0.025 M in the reaction mixture), the

final p[H] being adjusted with perchloric acid. The reactants

maintained at this temperature for at least 15 min prior to injection

were mixed in less than 1.5 ms in a flow-through optical quartz

cell of 1 cm path length. The kinetic traces, recorded at the

maximum of the UV absorption band assigned to the lead

complexes (274 nm), were averaged out of at least six replicates

and then processed on-line with the fitting routine implemented in

the Kinet Asyst 2 (Hi-Tech Scientific) software. This program fits

up to three exponential functions to the experimental curves by

the Marquardt nonlinear least-squares algorithm.⁴⁵The goodness

of fit was judged in terms of the statistical parameter ø²(in general,

values less than 0.1 were obtained) and by visual inspection of the

residual plots. The pseudo-first-order rate constants thus derived

are included in the Supporting Information. Linear regression and

nonlinear least-squares calculations were performed with Origin

6.0 from Microcal.⁴⁶Standard deviations are reported throughout

in parentheses as the last significant digit.

Protonation constants were determined in the p[H] range 1.8-

11.5 by titrating about 0.1 mmol of ligand (ca. 4 × 10^-3M)

dissolved in 25 mL of supporting electrolyte (0.1 M NaClO₄)

acidified with 4.000 mL of 0.1 M HClO₄. The electrochemical cell

was allowed to equilibrate for at least 1 min after each addition of

a 0.030 mL increment of 0.1 M NaOH. The collected potential

readings were converted into p[H] values with the help of a

Microsoft Excel spreadsheet by iterative solving of the four-

parameter calibration function.^37,38For each system, the individual

titration curves were analyzed by the weighed nonlinear least-

squares procedure implemented in Hyperquad 2000.³⁹In the final

refinement step the total amounts of titrated ligand and initially

added acid were also allowed to vary. The detailed experimental

and data processing procedures have been fully described else-

where.^37,38The final values are reported as the arithmetic mean of

at least three independent measurements together with their standard

deviations, which were systematically higher compared to those

derived for a single experiment from the full variance/covariance

matrix.⁴⁰

Spectrophotometric Titrations. Potentiometric measurements

coupled with a spectrophotometric detection were carried out using

the same titration cell and electrode calibration procedure as

described above. Visible absorption spectra of a ca. 4 × 10^-5

M

ligand solution containing an equimolar amount of Pb(ClO₄)₂were

recorded in situ as a function of p[H] ranging between 3.9 and 9.4

with a Cary 50 Probe (Varian) spectrophotometer equipped with

an immersion probe of 1 cm path length made of SUPRASIL 300

quartz (Hellma, ref 661.202).³⁸Prior to titration, the reference

spectrum of the 0.1 M NaClO₄supporting electrolyte solution was

acquired in the 240-340 nm range. The standard deviation of the

measured absorbance for the baseline was constant over the entire

wavelength region and did not exceed 0.001 absorbance unit.

Aliquots of a standardized 0.01 M NaOH solution containing 0.09

M NaClO₄were added manually with the help of a Gilmont

micropipet (2 µL resolution). The potential of the calibrated glass

electrode was measured with a PHM240 (Radiometer-Tacussel)

Results

X-ray Crystal Structures of L¹‚2H₂O and [L¹H₂](NO₃)₂‚

2H₂O. Single-crystal X-ray structure analysis of the free-

base ligand L¹and its diprotonated form was performed in

order to assess the conformational changes undergone by

the molecule upon protonation. Figure 1 shows the ORTEP

(35) Battino, R.; Williamson, A. G. J. Chem. Educ. 1984, 61, 51-52.

(36) Naja, G.; Mustin, C.; Volesky, B.; Berthelin, J. Water Res. 2005, 39,

579-588.

(41) Gampp, H.; Maeder, M.; Meyer, C. J.; Zuberbu¨hler, A. D. Talanta

1985, 32, 95-101.

(42) Gampp, H.; Maeder, M.; Meyer, C. J.; Zuberbu¨hler, A. D. Talanta

1985, 32, 251-264.

(37) Fre´mond, L.; Espinosa, E.; Meyer, M.; Denat, F.; Guilard, R.; Huch,

V.; Veith, M. New J. Chem. 2000, 24, 959-966.

(43) Gampp, H.; Maeder, M.; Meyer, C. J.; Zuberbu¨hler, A. D. Talanta

1985, 32, 1133-1139.

(38) Meyer, M.; Fre´mond, L.; Tabard, A.; Espinosa, E.; Vollmer, G. Y.;

Guilard, R.; Dory, Y. New J. Chem. 2005, 29, 99-108.

(39) Gans, P.; Sabatini, A.; Vacca, A. Talanta 1996, 43, 1739-1753.

(40) Raymond, K. N.; McCormick, J. M. J. Coord. Chem. 1998, 46, 51-

57.

(44) Alderighi, L.; Gans, P.; Ienco, A.; Peters, D.; Sabatini, A.; Vacca, A.

Coord. Chem. ReV. 1999, 184, 311-318.

(45) Marquardt, D. W. J. Soc. Ind. Appl. Math. 1963, 11, 431-441.

(46) Origin 6.0; Microcal Software Inc.: Northampton, MA.

Inorganic Chemistry, Vol. 44, No. 22, 2005 7899

Cuenot et al.

ascribed to type III. However, each arm is folded toward

the cyclam unit as highlighted by the torsional angles

C1-N1-C11-C12 (-82°), C5′-N1-C11-C12 (148.4°),

C2-N2-C21-C22 (-83.5°), and C3-N2-C21-C22

(146.1°), allowing the nitrogen lone pairs to interact with

one terminal amide proton via weak hydrogen bonds

(N13‚‚‚N1 ) 2.762 Å, N13-H13b‚‚‚N1 ) 101°; N23‚‚‚N2

) 2.723 Å, N23-H23b‚‚‚N2 ) 102°). In the crystal packing

the cohesion occurs primarily through the intricate inter-

molecular hydrogen-bond network surrounding each amide

oxygen atom as acceptor and terminal NH₂group as

donor (N13‚‚‚O22 ) 2.935 Å, N13-H13a‚‚‚O22 ) 178°;

N13‚‚‚O12 ) 2.889 Å, N13-H13b‚‚‚O12 ) 134°;

N23‚‚‚O22 ) 2.927 Å, N23-H23b‚‚‚O22 ) 126°;

N23‚‚‚O12 ) 2.951 Å, N23-H23a‚‚‚O12 ) 168°). Hence,

the twisted conformation of the appended substituents,

highlighted by the N1-C11-C12-N13 (-30.5°) and N2-

C21-C22-N23 (-28.5°) torsion angles, reflects an equi-

librium situation where the spatial arrangement of the arms

is mainly imposed by the stronger intermolecular NH‚‚‚O

hydrogen bonds and further stabilized by weaker intramo-

lecular NH‚‚‚N interactions.

The good quality of the diffraction data for [L¹H₂](NO₃)₂‚

2H₂O enabled the unambiguous identification of two trans-

located quaternary ammonium sites by visual inspection of

the Fourier difference map. Examination of the C(sp³)-N

distances further corroborated the assignment, since the

C-N5 bond lengths (1.506(8) Å) are on average significantly

longer than those related to the N1 atom (1.468(1) Å), the

latter comparing favorably with the mean value found for

the free base (1.466(6) Å). As shown in Figure 1b, proto-

nation of the N5 and N5′ tertiary amines induces a

pronounced molecular rearrangement, as evidenced by the

torsion-angle sequence τ₁-τ₇) -55.9, -69.1, -174.7,

-68.7, -47.0, 169.8, and -173.0°, with τ₁) N1-C2-C3-

C4. Hence, the cyclam skeleton switches to a quadrangular

[3,4,3,4]-C conformation where C3, C6, and the symmetry

related atoms are located at the vertexes. It turns out that

the directionality of each acetamide group with respect to

the N₄plane is now imposed by the conformational layout

of the macrocyclic chain, giving rise to a type IV configu-

ration of the nitrogen atoms.³⁷Moreover, the structural

Figure 1. ORTEP diagrams of (a) the free base and (b) the diprotonated

form of macrocycle L¹showing the atom-labeling scheme for one-half of

the molecule (symmetry related atoms are labeled according to the same

scheme and primed). Thermal ellipsoids are represented at the 50%

probability level. All H(-C) hydrogen atoms, counterions, and water

molecules have been omitted for clarity.

views of L¹and L¹H₂²⁺together with the numbering scheme.

Crystallographic parameters and data collection and refine-

ment information are summarized in Table 1. In both

structures the asymmetric unit contains one-half of the

molecular unit, the macrocycles being centered on crystal-

lographic inversion centers.

According to the extended Dale’s nomenclature that takes

into account the exact position of the heteroatoms within

the cycle,^22,37the 14-membered ring of L¹adopts a fully

extended [3,4,3,4]-A conformation. The tertiary amines

occupy the four corner positions and thus form a quadrilateral

of 3.783 (N1‚‚‚N2) by 5.116 Å (N1‚‚‚N2′) with an average

C-N-C angle of 114.6(1)°. Starting at τ₁, defined as the

N1-C1-C2-N2 torsion angle, the following sequence along

the unique ring fragment is observed: τ₁-τ₇) 167.3, -71.2,

-60.2, 173.5, -168.8, 57.9, and 66.0°. The four pendant

N-acetamide groups are pointing away from the cavity, one

pair belonging to the same propylenediamine subunit is

directed above and the other pair below the macrocyclic

mean plane. Following the stereochemical description of

tetraazamacrocycles introduced by Bosnich et al.,⁴⁷the

relative orientation of the four nitrogen substituents can be

2+

changes encountered in L¹H₂also affect the spatial

extension of the pendant arms. The amide groups born by

the protonated nitrogen atoms are fully extended (N5-C51-

C52-N52 ) 166.0°) and point away from the cavity center

(C4-N5-C51-C52 ) -56.9° and C6-N5-C51-C52 )

71.1°). In contrast, both functions attached to N1 and N1′

are folded, one above and the other below the ring (N1-

C11-C12-O12 ) 38.6°, C2-N1-C11-C12 ) 109.4°, and

C7′-N1-C11-C12 ) -124.4°). Each oxygen atom points

toward the adjacent ammonium proton, giving rise to a

hydrogen bond (N5‚‚‚O12′ ) 2.784(4) Å, N5-H5‚‚‚O12′

) 154.9(2)°). The difference in the C12-O12 (1.245(2) Å)

and C52-O52 (1.224(2) Å) bond distances translates into a

split carbonyl stretching mode in the FTIR spectrum (ν_CdO

) 1668 and 1691 cm^-1). Furthermore, the H5 proton is also

(47) Bosnich, B.; Poon, C. K.; Tobe, M. L. Inorg. Chem. 1965, 4, 1102-

1108.

7900 Inorganic Chemistry, Vol. 44, No. 22, 2005

Cyclam DeriWatiWes

Table 2. Protonation Constants of L¹, L², and L³and Stability

involved in a second, albeit much weaker, hydrogen bond

with the unprotonated N1′ atom (N5‚‚‚N1′ ) 2.843(2) Å,

N5-H5‚‚‚N1′ ) 111.3(1)°). However, the large departure

from the ideal value (360°) of the sum of the three angles

N5-H5‚‚‚N1′, N5-H5‚‚‚O12′, and N1′‚‚‚H5‚‚‚O12′ (348.7°)

rules out formation of three-center hydrogen bonds,⁴⁸like

those found in the diprotonated form of the propionamide

analogue of L¹.⁴⁹The longer arms and the R positioning of

the four amines with respect to the corners ([3,4,3,4]-B

conformer) account for the slightly different spatial arrange-

ment adopted by the lateral functionalities in both

molecules.

Constants of the Corresponding Lead(II) Complexes^a

M

L¹

L²

L³

H⁺

log K₀₁₁

log K₀₁₂

log â₀₁₂

∆log K^b

log â₁₁₀

7.44(2)

6.38(1)

13.82(2)

1.06

8.77(1)

8.09(1)

16.86(1)

0.68

8.85(2)

8.03(1)

16.88(2)

0.82

Pb²⁺

8.52(1)

11.30(3)

11.83(3)

^aI ) 0.1 M NaClO₄; T ) 298.2(2) K. ^b∆ log K ) log K₀₁₁- log K₀₁₂

.

Synthesis and Characterization of the Lead(II) Com-

plexes. Although perchlorates are potentially hazardous

compounds, spectroscopic arguments prompted us to delib-

erately choose this anion in the present study because it does

not absorb light in the UV range, in contrast to the nitrate

salts.⁵⁰Lead(II) complexes of the octadentate N-methylcar-

bamoyl pendant-armed macrocycles were readily obtained

by reacting in aqueous media equimolar quantities of L¹,

L², or L³with Pb(ClO₄)₂‚3H₂O, followed by crystallization.

The MALDI-TOF MS spectra displayed the molecular ion

and its perchlorate-adduct peaks with the correct isotope

pattern together with the signal assigned to the free ligand.

Laser-induced demetalation during the desorption process

is a typical behavior for this type of compounds. Elemental

analyses were consistent with a [Pb(L)](ClO₄)₂‚xH₂O for-

mulation assuming x ) 3 (L¹), 0.5 (L²), and 0 (L³), in

agreement with the presence of broad infrared absorption

bands in the 3200-3500 cm^-1range for the L¹and L²

complexes. Characteristic vibration bands of the cocrystal-

lized perchlorate anions appear in the expected region as a

triplet centered around 1110 cm^-1(ν₃) and a doublet at 625

cm^-1(ν₄), suggesting C_2Vsymmetry in the solid state.⁵¹

Infrared spectroscopy also provided structural evidence for

the binding of the amide groups to Pb²⁺through the oxygen

atoms. The FTIR spectra of L²and L³exhibit a bathochromic

shift of ca. 40 cm^-1of the carbonyl stretching frequencies

as a result of the bond strength weakening upon complex-

ation. Surprisingly, the CdO vibration mode of L¹is not

significantly affected, although the X-ray crystal data

obtained for the nitrate complex shows the encapsulation of

the lead cation inside the molecular cavity created by both

sets of donor atoms, the tertiary amines, and the carbonyl

oxygen atoms.²¹However, occurrence in the solid-state

structure of L¹‚2H₂O of intermolecular hydrogen-bond

interactions involving the amide groups hampers a direct

comparison of the free-ligand and metal-complex spectra. It

is indeed recognized that hydrogen bonding can shift by

several tenths of cm^-1to lower energy carbonyl absorption

bands.

Figure 2. Spectrophotometric titration of L¹in the presence of Pb²⁺as a

function of p[H]. [Pb]_tot) 3.9 × 10^-5M; [L]_tot) 4 × 10^-5M; I ) 0.1 M

NaClO₄; T ) 298.2(2) K; l ) 1 cm. Spectra 1-13: p[H] ) 3.957, 4.312,

4.487, 4.600, 4.719, 4.845, 4.977, 5.135, 5.321, 5.588, 6.143, 8.272, 9.446.

Due to their limited water solubility the solution structure

of each complex was characterized in DMSO-d₆by NMR

1

spectroscopy. H and ¹³C NMR resonances were assigned

using ¹H-¹H COSY and ¹H-¹³C HETCOR correlation

charts. The spectra are in agreement with the proposed

octacoordinated shell structure found in the crystal state for

[Pb(L¹)](NO₃)₂: they exhibit a single set of signals cor-

responding to one-fourth of the total number of proton or

carbon atoms, as expected for dynamic time-averaged C_2V-

symmetric species. Upon metal binding the four pairs of

geminal aliphatic protons become diastereotopic, in agree-

ment with the unidirectional orientation of the four N-

carbamoylmethyl substituents that differentiates both sides

of the cyclam ring. Ligation of the amidic oxygen atoms is

supported by a downfield shift (0.85 and 1 ppm, respectively)

of both CONH₂resonances as well as by a larger peak

separation (∆δ ) 0.22 ppm for L¹vs 0.38 for Pb(L¹)²⁺).

Interestingly, cross-correlation peaks between both terminal

N-methyl resonances but also between them and the AB spin

system appearing at 3.81 ppm assigned to the methylenic

arm protons were observed in the COSY map of Pb(L²)²⁺

(Figure S1), supporting four- and even five-bond scalar

coupling, respectively, within the rigidified amide chelate

rings. It is also worth noting that the ¹³C NMR carbonyl

signal is flanked by two satellite peaks corresponding to the

coupling between the carbon nucleus and the 22.6% abundant

1

²⁰⁷Pb isotope (I ) /₂) with J

≈ 15 Hz, whereas it

²⁰⁷Pb-¹³

C

appears as a singlet in the free ligand. This observation is a

further indication that the encapsulated Pb²⁺cation is bound

to the carbonyl oxygen atoms.

(48) Jeffrey, G. A.; Mitra, J. Acta Crystallogr., Sect. B 1983, 39, 469-

480.

(49) Dahaoui-Gindrey, V.; Lecomte, C.; Gros, C.; Mishra, A. K.; Guilard,

R. New J. Chem. 1995, 19, 831-838.

(50) Meyer, M.; Cuenot, F.; Espinosa, E.; Mangayarkarasi, N.; Bucaille,

A.; Guilard, R. Unpublished results.

Solution Equilibrium Studies. Overall stability constants

â_mlhreported hereafter relate to equilibrium 1, where M, L,

and H refer to the unhydrolyzed aquametal ion, the ligand,

and the hydronium ion, respectively. The associated stoi-

chiometric coefficients are labeled m, l, and h, whereas

(51) Nakamoto, K. Infrared and Raman Spectra of Inorganic and Coor-

dination Compounds; Wiley: New York, 1970.

Inorganic Chemistry, Vol. 44, No. 22, 2005 7901

Cuenot et al.

Figure 3. Distribution diagrams of the lead(II) hydroxo and carbonato species ([Pb]_tot) 2 × 10^-5M) (a) in anaerobic conditions (p_CO) 0 atm) and (b)

2

under CO₂atmosphere (p_CO) 10^-3.5atm), (c) of the protonated forms of L¹(-) and L³(--) ([L]_tot) 2 × 10^-4M), and (d) of the Pb²⁺/L¹(-) and Pb²⁺/L³

(--) systems in the absence ²of CO₂([Pb]_tot) 2 × 10^-5M; [L]_tot) 2 × 10^-4M). Lead hydrolysis and carbonation equilibrium constants were taken from

our own critical survey of literature data.⁵³

charges have been omitted for clarity. Stepwise protonation

constants K_01iare defined by the relation â_01h) Π K_01i.

The entire set of thermodynamic data reported in Table 2

allowed the computation of the species distribution diagrams

shown in Figure 3 as a function of the free proton concentra-

tion. According to these curves, the mono- and diprotonated

forms of L¹, L², or L³predominate in the p[H] range 4-7,

although a small fraction of the free base is also present at

the upper end. Lead(II) exists as a mixture of Pb²⁺and the

soluble hydroxo complex Pb(OH)⁺if the solution is protected

against CO₂ingress, while all three complexes are formed

at more than 80% above p[H] 5.5.

mM + lL + hH a M_mL_lH_h

(1)

The acid-base properties in 0.1 M NaClO₄solutions of

tetraamines L¹, L², and L³were investigated at 298.2 K by

glass-electrode potentiometry. Analysis of the titration curves

acquired between p[H] 1.8 and 11.5 (p[H] ) -log [H₃O⁺])

indicated only two sequential protonation steps. The cor-

responding values of the equilibrium constants are collected

in Table 2.

Formation Kinetics. Taking into account the predomi-

nance diagrams depicted in Figure 3, the lead(II) complex

formation reaction may be written for each ligand L

according to the general formula given by eq 2.

The lead(II) uptake by L¹, L², and L³was monitored at

298.2 K by UV absorption spectrophotometry as a function

of p[H] in 0.1 M NaClO₄. A typical set of spectra recorded

during the titration of equimolar quantities of macrocycle

L¹and Pb²⁺with sodium hydroxide is depicted in Figure 2.

It shows the progressive appearance of the characteristic

absorption band of the mononuclear lead(II) complex with

a maximum at 274 nm. For each system, singular value

decomposition of the complete multiwavelength data array

with the program Specfit^41-43points to the existence of only

Pb(OH)_x^(2-x)++ LH_h^h+f Pb(L)²⁺+ (h - x) H⁺(2)

The time course of the complex formation was monitored

at 274 nm by single-wavelength stopped-flow spectropho-

tometry and treated as a pseudo-first-order process by

degenerating the reaction order with respect to the ligand.

Concentrations after mixing were at least 10 times larger

than those of lead(II). Figure 4 shows a typical absorption

versus time profile recorded on the millisecond time scale

for the insertion of Pb²⁺(2.015 × 10^-5M) into L¹present

in 30-fold excess (5.98 × 10^-4M). Under such conditions

the reaction proceeds to completion in a single rate-limiting

step with no more initial loss of spectrophotometric amplitude

than that expected from the dead time of the employed

instrument. For each system the process was found to be of

first order with respect to lead(II). Indeed, the recorded signal

could be fitted with excellent statistical confidence to a

single-exponential function by nonlinear least squares as

2+

two spectroscopically distinguishable species (LH₂and

Pb(L)²⁺). Subsequent nonlinear least-squares refinement of

the observed absorbance data confirmed the hypothesized

model and afforded best estimates of the complex-formation

equilibrium constants â₁₁₀, which were in excellent agreement

with those calculated by fitting the potentiometric data alone

with Hyperquad 2000³⁹(Table 2). Extinction coefficients

calculated by Specfit were, within experimental error,

identical to the values obtained after dissolving the pure

complex in water (ꢀ_max) 6850, 8070, and 8400 M^-1cm^-1

at λ_max) 274 nm for Pb(L¹)²⁺, Pb(L²)²⁺, and Pb(L³)²⁺,

respectively).

7902 Inorganic Chemistry, Vol. 44, No. 22, 2005

Cyclam DeriWatiWes

Figure 4. Spectrophotometric signal recorded during the complexation

reaction of lead(II) by L¹. [Pb]_tot) 2.015 × 10^-5M; [L¹]_tot) 5.98 × 10^-4

M; p[H] ) 5.48; [MES] ) 0.025 M; I ) 0.1 M NaClO₄; T ) 298.2(2) K;

λ ) 274 nm; l ) 1 cm.

Figure 6. Variations of the rate constants (a) k_fand (b) k_dfor Pb(L¹)²⁺as

a function of the proton concentration. [MES] ) 0.025 M; I ) 0.1 M

NaClO₄; T ) 298.2(2) K.

Figure 5. Variation of the pseudo-first-order rate constant k_obswith the

analytical concentration of L¹at different p[H] values. [Pb]_tot) 2.015 ×

10^-5M; [MES] ) 0.025 M; I ) 0.1 M NaClO₄; T ) 298.2(2) K.

not significantly different from zero, the experimental data

were reprocessed and adjusted to a line passing through the

origin.

Table 3. Formation (k_f) and Dissociation (k_d) Rate Constants of

Lead(II) Complexes^a

k_obs) k_f[L]_tot+ k_d

(3)

L¹

L²

L³

k_f× 10^-3

k_d

k_f× 10^-3

k_d

k_f× 10^-3k_d

d[Pb(L)²⁺]

p[H] (M^-1s^-1

)

(s^-1

)

p[H] (M^-1s^-1

)

(s^-1

)

p[H] (M^-1s^-1) (s^-1

)

V )

) k_obs[Pb]_tot

(4)

dt

4.29

8.8(1) 4.60(4) 5.08

9.2(1) 0.31(6) 5.09

9.3(1)

5.40 18.1(2)

5.71 37.9(1)

5.95 61.0(2)

6.14 98.5(3)

6.39 167.8(7)

6.64 294.2(6)

6.78 402(3)

0

4.45 13.3(1) 3.40(8) 5.40 19.8(1)

4.69 24.5(3) 1.4(3) 5.71 40.6(1)

4.88 42.0(3) 1.1(2) 5.93 67.5(2)

5.09 74.6(3)

5.48 164(1)

5.85 347(2)

6.05 529(5)

0

For each investigated system the values of the second-order

rate constants k_fgathered in Table 3 are strongly p[H]

dependent. Excluding the data points collected near neutrality

where a slight deviation occurs, the variation of the k_fvalues

as a function of [H⁺] on a bilogarithmic scale affords a

straight line with a negative unitary slope supporting an

apparent reaction order of -1 with respect to the proton

concentration. This general trend, which corresponds to a

slowing of the reaction with increasing acid concentration,

suggests that different protonated states of one or both

reactants, each of them having its own reactivity, are involved

in complex formation (Figure 6a). According to the metal

and ligand speciation diagrams displayed in Figure 3, the

global mechanism is outlined in Scheme 1.

0

6.14 112.1(2)

6.37 188.0(3)

6.63 336.9(4)

6.79 452(1)

6.98 543(1)

6.96 589(1)

^a[Pb]_tot) 2.015 × 10^-5M; [MES] ) 0.025 M; I ) 0.1 M NaClO₄; T

) 298.2(2) K.

shown in Figure 4. The apparent formation rate constants

(k_obs) are tabulated in the Supporting Information.

The linear variation of the pseudo-first-order rate constants

(k_obs) with the analytical concentration of the entering ligand

at different p[H] values is graphically displayed in Figure 5

(the raw data may be found in the Supporting Information).

At each p[H] the data were fitted to eq 3 by linear regression,

supporting the apparent rate law given by eq 4. The values

of the second-order k_fand first-order k_drate constants thus

determined are listed in Table 3. The statistical significance

of the intercept was evaluated at the 95% probability level

by the Student t test.^46,52When k_dvalues were negative or

The corresponding rate law derived from Scheme 1 is

given by eq 5.

V ) (k₁[L] + k₃[LH⁺] + k₅[LH₂²⁺])[Pb²⁺] + (k₂[L] +

k₄[LH⁺] + k₆[LH₂²⁺])[Pb(OH)⁺] - (k_-1+ k_-2[OH^-] +

k_-3[H⁺] + k_-4+ k_-5[H⁺]²+ k_-6[H⁺])[Pb(L)²⁺] (5)

According to the mass balance equations, the total ligand

concentration can be approximated by eq 6 as a consequence

of the pseudo-first-order conditions, while the total lead

(52) Commissariat a` l’Energie Atomique. Statistique applique´e a` l’exploitation

des mesures; Masson: Paris, 1978; Vols. 1-3.

Inorganic Chemistry, Vol. 44, No. 22, 2005 7903

Cuenot et al.

Scheme 1

mental time-dependent variation of the absorbance at 274

nm since the Pb(L)²⁺complex is the only absorbing species

at this wavelength. Hence, k_fand k_d(eq 3) are, respectively,

identified as k and kh.

d[Pb(L)²⁺]

[Pb(L)²⁺] - [Pb(L)²⁺]_∞

) - (k[L]_tot+ kh)dt

(11)

[Pb(L)²⁺] - [Pb(L)²⁺]_∞

[Pb(L)²⁺]_t)0- [Pb(L)²⁺]_∞

) exp{-(k[L]_tot+ kh)t} (12)

The formation rate constants (k_i, i ) 1-6) of eq 9 were

treated as adjustable parameters by a weighted nonlinear

least-squares procedure. The employed weighing scheme (ω_i

) (1/k_i^exp)²), which was intended to approximately equalize

the contribution to the mean square sum of each data point,

intrinsically assumes a constant relative error of k_fvalues

over the entire H⁺concentration range. In our conditions

only two of the six second-order rate constants k_icould be

adjusted; inclusion of an additional parameter led systemati-

cally to divergence. Moreover, we were faced with proton

ambiguity since data refinement by two different models

yields the same goodness-of-fit criterion ø².⁴⁶All the results

summarized in Table 4 are consistent with a reaction

mechanism involving the unhydrolyzed Pb²⁺cation and the

monoprotonated form of each macrocycle (k₃). Considering

this sole pathway in the p[H] region where LH₂²⁺and Pb²⁺

predominate, eq 9 reduces to k ≈ (k₃/K₀₁₂) × [H⁺]^-1, in

agreement with the empirical rate law suggested by the

bilogarithmic representation of k_fvs [H⁺]. The deviations

from the straight line for the data points collected at p[H]

values close to 7 are equally well explained by considering

the reaction of either Pb²⁺with the deprotonated ligand

species L (k₁) or the monohydrolyzed Pb(OH)⁺cation with

the monoprotonated macrocycle LH⁺(k₄) as the additional

pathway. Unfortunately, the appearance of precipitates above

p[H] 7 precluded a more detailed investigation of the

complexation mechanism in alkaline media.

concentration is expressed by eq 7, where â₀₁₁, â₀₁₂, and â_10-1

stand for the first and second cumulative protonation constant

of the ligand and for the first hydrolysis constant of Pb²⁺

(10^-7.32for I ) 0.1 NaClO₄, T ) 298 K),⁵³respectively.

[L]_tot) [L] + [LH⁺] + [LH₂²⁺] + [Pb(L)²⁺] ≈

[L](1 + â₀₁₁[H⁺] + â₀₁₂[H⁺]²)

[L]_tot

[L] )

(6)

(1 + â₀₁₁[H⁺] + â₀₁₂[H⁺]²)

[Pb²⁺]_tot) [Pb²⁺] + [Pb(OH)⁺] + [Pb(L)²⁺] )

[H⁺] + â_10-1

[Pb²⁺]

+ [Pb(L)²⁺]

[H⁺]

(

)

[H⁺]

[Pb²⁺] ) ([Pb²⁺]_tot- [Pb(L)²⁺])

(7)

+

(

)

[H ] + â_10-1

As far as the reverse reactions are concerned, significant

k_dvalues increasing linearly with the proton concentration

were only obtained for L¹below p[H] 5 (Figure 6b).

According to eq 10, the slope of the best-fit line passing

through the origin corresponds to the sum k_-3+ k_-6) 8.9(4)

× 10⁴M^-1s^-1. Since the forward rate constant k₆was

inaccessible, it is reasonable to assume also that k_-6≈ 0.

The absence of any significant ordinate at the origin further

suggests that k_-1+ k_-4is negligible as well as k_-2and k_-5,

which would lead to a nonlinear dependence of k_dvs [H⁺].

Conversely, none of the six reverse rate constants k_-icould

be determined directly for L²and L³under the experimental

conditions used herein. Taking into account the equilibrium

constants reported in Table 2, the computed estimates of k_-1

) k₁/â₁₁₀(3.0 × 10^-2, 1.3 × 10^-3, and 3.4 × 10^-4M^-1s^-1

for L¹, L², and L³, respectively) and k_-3) k₃× K₀₁₁/â₁₁₀

(1.1 × 10⁵, 2.9 × 10⁴, and 8.2 × 10³M^-1s^-1for L¹, L²,

and L³, respectively) are in good accordance with the above-

mentioned results.

Substituting the concentrations given by eqs 6 and 7 into eq

5, the rate law rewrites as eq 8, k and kh being defined by

eqs 9 and 10, respectively.

d[Pb(L)²⁺]

V )

) k[L]_tot[Pb]_tot- (k[L]_tot+ kh)[Pb(L)²⁺] (8)

dt

k ) {k₂â_10-1+ (k₁+ k₄â₀₁₁

(k₃â₀₁₁+ k₆â₀₁₂

_10-1)[H⁺]²+ k₅â₀₁₂[H⁺]³}/

{(1 + â₀₁₁[H⁺] + â₀₁₂[H⁺]²)([H⁺] + â_10-1)} (9)

k_-2

â

_10-1)[H⁺] +

â

K

[H⁺]

kh ) k_-1+ k_-4

+

^w+ (k_-3+ k_-6)[H⁺] + k_-5[H⁺]²(10)

Integration of the corresponding differential, eq 11, yields

eq 12, which is mathematically equivalent to the empirical

monoexponential fitting function used to model the experi-

(53) Cuenot, F.; Meyer, M.; Guilard, R. Unpublished results.

7904 Inorganic Chemistry, Vol. 44, No. 22, 2005

Cyclam DeriWatiWes

Table 4. Values of Second-Order Formation Rate Constants k₁, k₃, and k₄Fitted by Nonlinear Least Squares for Various Models^a

model

k₁(M^-1s^-1

)

k₃(M^-1s^-1

)

k₄(M^-1s^-1

)

ø²

Pb²⁺+ L¹and Pb²⁺+ L¹H⁺

10(2) × 10⁶

25(3) × 10⁷

23(2) × 10⁷

13.6(5) × 10⁵

98(1) × 10⁵

78(1) × 10⁵

0.0083

Pb²⁺+ L¹H⁺and Pb(OH)⁺+ L¹H⁺

Pb²⁺+ L²and Pb²⁺+ L²H⁺

7(1) × 10⁶

90(9) × 10⁵

67(6) × 10⁵

0.00091

0.00076

Pb²⁺+ L²H⁺and Pb(OH)⁺+ L²H⁺

Pb²⁺+ L³and Pb²⁺+ L³H⁺

Pb²⁺+ L³H⁺and Pb(OH)⁺+ L³H⁺

^a[MES] ) 0.025 M; I ) 0.1 M NaClO₄; T ) 298.2(2) K.

Table 5. Nonlinear Least-Squares Adjusted Values of k_d, k^H₁, k^H₂, and

K for the H⁺-Assisted Demetalation Kinetics^a

Pb(L¹)²⁺

Pb(L²)²⁺

Pb(L³)²⁺

k_d(s^-1

K (M^-1

k^H₁(s^-1

k^H₂(M^-1s^-1

)

9(4)

1.6(9)

0

)

2.2(3) × 10²

2.2(1) × 10²

5.2(2) × 10³

1.6(1) × 10²

1.18(5) × 10²

3.56(6) × 10³

0.44(9) × 10²

1.9(3) × 10²

3.4(2) × 10³

)

^aI ) 0.1 M (H,Na)ClO₄; T ) 298.2(2) K.

adjusted parameters (K, k_d, k^H, and k^H) are presented in

Table 5.

1

2

Figure 7. Variation of the pseudo-first-order rate constant k_obsfor the

dissociation of the lead(II) complexes as a function of proton concentration.

Solid lines were drawn according to eq 15 and the refined parameters

2

k_d+ k^HK[H⁺]_tot+ k^HK[H⁺]_tot

collected in Table 5. (0) [Pb(L¹)²⁺

]

_tot) 1.84 × 10^-5M; (O) [Pb(L²)²⁺

]

tot

1

2

) 1.80 × 10^-5M; (4) [Pb(L³)²⁺

]

tot

) 1.13 × 10^-5M; I ) 0.1 M

k_obs

)

(15)

1 + K[H⁺]_tot

(H,Na)ClO₄; T ) 298.2(2) K.

Discussion

Dissociation Kinetics. The proton-assisted dissociation

processes of the three lead(II) complexes have also been

studied by stopped-flow absorption spectrophotometry. The

overall reaction writes as eq 13.

Solution Thermodynamics. In the investigated p[H] range

(1.8-11.5) the three tetraamines behave as a moderately

strong diprotic base. As a common feature among cyclam

derivatives,⁵⁴the tri- and tetraprotonated species do not form

in any appreciable amount, even in strongly acidic aqueous

solutions, as a consequence of positive charge accumulation.

The difference ∆log K ) log K₀₁₁- log K₀₁₂directly reflects

the electrostatic repulsion energy between the ammonium

centers once the statistical factor has been factored out

(∆log K ) log 8/3 ≈ 0.43 for a receptor possessing four

identical and independent binding sites).⁵⁵For each macro-

cycle, the much larger than expected ∆log K value experi-

mentally observed (Table 2) reveals strong Coulombic

interactions in the diprotonated form, preventing further

proton addition.

Data in Table 2 also indicate an increase of the overall

basicity when the primary acetamide groups (L¹) are replaced

by either N,N-dimethyl- (L²) or N,N-diethyl- (L³)

carbamoylmethyl arms. The 3 orders of magnitude difference

between the â₀₁₂values recorded for primary and tertiary

amides can be understood in terms of intramolecular

hydrogen-bond formation between the macrocyclic nitrogen

and amidic hydrogen atoms of L¹. Involvement of each

nitrogen free-electron doublet in a five-membered ring, as

observed in the crystal structure, hinders the protonation

process and lowers by several orders of magnitude the proton

affinity of a tertiary amine in water, as demonstrated by

several authors.^56-61Obviously, such N-H‚‚‚N interactions

Pb(L)²⁺+ 2H⁺f Pb²⁺+ LH₂

(13)

2+

Kinetic runs were performed under pseudo-first-order condi-

tions by mixing the lead(II) complex (1-2 × 10^-5M) with

perchloric acid in large excess (2 × 10^-4M < [H⁺]_tot< 5

× 10^-2M). At any acid concentration the reaction was

quantitative, as evidenced by the total disappearance of the

UV absorption band. Moreover, experimental absorbance

data, recorded as a function of time at 274 nm, evidenced

no significant amplitude loss during the mixing time and were

fitted as a single-exponential decay with excellent confidence.

Thus, the proton-assisted demetalation reaction follows first-

order kinetics with respect to the complex. The apparent

dissociation rate law is expressed by the differential eq 14,

where k_obsrepresents the pseudo-first-order dissociation rate

constant.

d[Pb(L)²⁺]

V ) -

) k_obs[Pb(L)²⁺]

(14)

dt

Plots of k_obsversus [H⁺]_totfor the three complexes (Figure

7; the raw data may be found in the Supporting Information)

show a saturation behavior at low acid concentrations ([H⁺]_tot

< 0.02 M) followed by a quasi-linear increase at higher

concentrations. Nonlinear least-squares analysis provided an

excellent agreement between the data plotted in Figure 7 and

the rate expression given by eq 15. Best estimates of the

(54) Bencini, A.; Bianchi, A.; Garcia-Espana, E.; Micheloni, M.; Ramirez,

J. A. Coord. Chem. ReV. 1999, 188, 97-156.

(55) Perlmutter-Hayman, B. Acc. Chem. Res. 1986, 19, 90-96.

Inorganic Chemistry, Vol. 44, No. 22, 2005 7905

Cuenot et al.

cannot occur in L²or L³, so that protonation constants remain

typical for tertiary amines bearing electron-withdrawing

groups.⁶²In comparison, K₀₁₁and K₀₁₂values reported for

1,4,8,11-tetramethyl-1,4,8,11-tetraazacyclotetradecane (TMC)

are approximately 10 times higher (log K₀₁₁) 9.39 and

log K₀₁₂) 9.05; I ) 0.1 NaNO₃; T ) 298.2 K).⁶³

Since the three macrocycles exhibit different acid-base

properties, speciation calculations provide an accurate means

to discuss their lead binding affinity.²¹As can be seen in

Figure 3d, the efficiencies of the primary and tertiary amides

are essentially identical, although the latter species form

complexes that are 3 orders of magnitude more stable than

the former. Comparison of the â₁₁₀and â₀₁₂values found in

Table 2 indicates that the higher stability of Pb(L²)²⁺and

Pb(L³)²⁺over Pb(L¹)²⁺parallels the higher basicity exhibited

by the free bases L²or L³. Hence, the lower affinity of L¹

for lead(II) is almost integrally compensated by its more

acidic character, leading to a similar lead affinity at any given

p[H] value. Besides carboxylic, phosphonic, and phosphinic

tetraamino derivatives, our series provides an additional

example where the stability of the metal complexes is

essentially governed by the ligand’s overall basicity.⁶⁴

Figure 8. Comparison of the apparent pseudo-first-order formation rate

constant k_obscalculated for [L]_tot) 10^-3M as a function of p[H]. I ) 0.1

M NaClO₄; T ) 298.2(2) K. The portion of the dotted line drawn for L¹

shows that the corresponding rate constants become too fast to be measured

by a stopped-flow mixing device.

valuable clues as the lead-uptake rate increases in the order

L³≈ L², L¹. The superior kinetic performances evidenced

by L¹over L²and L³in weakly acidic media are essentially

related to its less basic character as shown in Figure 3c,

whereas the tertiary amides are essentially present below p[H]

7 in their almost unreactive diprotonated form.

From a mechanistic point of view, the p[H] dependency

of the second-order formation rate constants k_fcould be

satisfactorily rationalized in terms of two parallel paths,

although proton ambiguity hampers the definitive assignment.

Nevertheless, data in Table 4 show that Pb²⁺and the

monoprotonated macrocycle are the main reacting species

Formation Kinetics. Our kinetic investigations have

evidenced a single rate-limiting step process leading to the

fast formation of the lead(II) complexes, although the

medium’s acidity level exerts a marked influence on the

overall reaction rate. This has some important practical

consequences if one seeks to implement related amide-

substituted cyclam sequestrants in solid/liquid extraction

processes.²¹Keeping in mind that the ultimate goal of this

study was to provide an appraisal of the kinetic lead-uptake

performances of ligands L¹-L³, only the apparent complex-

ation rate constants (k_obs) under precisely defined p[H]

conditions and concentrations enable a direct comparison.

Hence, each p[H]-dependent curve plotted in Figure 8 has

2+

(k₃) below p[H] ≈ 5.5, although the LH₂form predomi-

nates in solution. Close to neutrality an alternative pathway

becomes operative but the reaction of Pb²⁺with L (k₁) is

kinetically indistinguishable from that of Pb(OH)⁺with LH⁺

(k₄). As found for several main-group and transition-metal

cations,⁶⁷it might be argued that the monohydrolyzed lead

species Pb(OH)⁺is more reactive than the aquated Pb²⁺

cation due to labilization of the water molecules in the first

solvation shell by the hydroxo ligand, leading to a partial

charge transfer to the metal. The anticipated rate-enhance-

ment factor might be on the order of 10-10⁴, whereas the

k₄/k₃ratio does not exceed 5 in the case of L¹and is even

less than unity for the two other ligands (Table 4). Con-

versely, the refined k₁values corresponding to the reaction

of Pb²⁺with the free-base ligands are systematically higher

by a factor of 7-30 than those found for k₃. This result is

fully consistent with the more favorable electrostatic inter-

action in the early stage of the complexation process when

the divalent lead cation reacts with an uncharged rather than

a monopositively charged macrocycle (vide infra). Relying

on these arguments, participation of Pb(OH)⁺in the mech-

anism depicted in Scheme 1 can be ruled out, thus solving

the proton ambiguity.

been calculated for a total ligand concentration of 10^-3

M

by combining eqs 9 and 10 with the rate law expression, eq

3. Although it should be stressed that lead speciation in tap

water at pH values between 6 and 9 differs noticeably from

the actual solution compositions used in this study due to

the predominance of soluble Pb(CO₃),^65,66Figure 8 provides

(56) Serratrice, G.; Mourral, C.; Zeghli, A.; Be´guin, C. G.; Baret, P.; Pierre,

J. L. New J. Chem. 1994, 18, 749-758.

(57) Caris, C.; Baret, P.; Pierre, J. L.; Serratrice, G. Tetrahedron 1996, 52,

4659-4672.

(58) Cohen, S. M.; Meyer, M.; Raymond, K. N. J. Am. Chem. Soc. 1998,

120, 6277-6286.

(59) Xu, J.; O’Sullivan, B.; Raymond, K. N. Inorg. Chem. 2002, 41, 6731-

6742.

(60) Kotek, J.; Lubal, P.; Hermann, P.; Cisarova, I.; Lukes, I.; Godula, T.;

Svobodova, I.; Taborsky, P.; Havel, J. Chem.sEur. J. 2003, 9, 233-

248.

(61) Wang, Y. M.; Li, C. R.; Huang, Y. C.; Ou, M. H.; Liu, G. C. Inorg.

Chem. 2005, 44, 382-392.

(62) Smith, R. M.; Martell, A. E.; Motekaitis, R. J. NIST Critically Selected

Stability Constants of Metal Complexes Database, release 6.0; NIST

Standard Reference Data No. 46: NIST: Gaithersburg, MD, 2001.

(63) Grzejdziak, A. Monatsh. Chem. 1994, 125, 107-117.

(64) Lukes, I.; Kotek, J.; Vojtisek, P.; Hermann, P. Coord. Chem. ReV.

2001, 216-217, 287-312.

(65) Hem, J. D.; Durum, W. H. J.sAm. Water Works Assoc. 1973, 65,

562-568.

(66) Hem, J. D. Geochim. Cosmochim. Acta 1976, 40, 599-609.

The kinetic parameters derived in the present work

evidence a noticeable influence of rather minor structural

changes in the ligands L¹-L³on the reaction rate. Consider-

ing the second-order rate constants k₃that mirror the intrinsic

reactivity of the monoprotonated macrocycles, the methylated

L²H⁺and ethylated L³H⁺tetraamides react about 7 times

more rapidly with Pb²⁺than the less crowded acetamide

(67) Lincoln, S. F.; Merbach, A. E. AdV. Inorg. Chem. 1995, 42, 1-88.

7906 Inorganic Chemistry, Vol. 44, No. 22, 2005

Cyclam DeriWatiWes

Table 6. Intrinsic Formation Constants (k′) of Lead(II) Tetraaza-

f

macrocyclic Complexes in Water at 298.2 K^a

ligand

EDTA

HEDTA

EDDDA

EGTA

1,3-PDTA

HEIDA

NTA

k′_f(M^-1s^-1

)

k′_f/k_ex(M^-1

)

I (M)

0.3

ref

LH^3-6.1 ×10¹¹

81

32

96

LH^2-2.4 ×10¹¹

LH^3-2.5(2) ×10¹⁰3.3

0.4-0.8 80

0.3 96

0.4-0.8 80

LH^3-7.5 ×10¹⁰

10

LH^3-6.8(3) × 10¹⁰9.1

L^2-

L^3-

2.0 × 10¹⁰

2.7

0.3

96

79

97

91

85

86

24

<1.5 × 10¹¹<20

LH^2-4.4(9) × 10⁶5.9 × 10^-4

[18]aneO₆

L

5.92 × 10⁶

7.9 × 10^-40.1

2.3 × 10^-70.2

3.1 × 10^-20.2

1.9 × 10^-100.2

Figure 9. Linear free energy relationship between the second-order rate

constants k₃and the first protonation constants â₀₁₁. I ) 0.1 M NaClO₄; T

) 298.2(2) K.

Ac₂[18]aneN₂O₄LH₂1.7 × 10³

LH^-2.30 × 10⁸

cyclen

DOTA

LH⁺8.3(1) × 10⁵1.1 × 10^-40.2

2+

LH₂1.4(2)

derivative L¹H⁺. The empirical linear free-energy relationship

between log k₃and log â₀₁₁values displayed in Figure 9

provides a reasonable explanation for this significant effect.

The weaker basicity of L¹compared to its analogues L²and

L³has been ascribed to the formation of intramolecular

hydrogen bonds engaging the free electron pair of each

tertiary amine and its appended acetamide group in a five-

membered ring. Provided such interactions might also occur

in the monoprotonated state, the decreased reactivity of L¹H⁺

can directly be related to the necessity to break these

hydrogen bonds prior to Pb-N bond formation and the

concomitant chelate ring closure.^60,68Finally, increased steric

hindrance provided by the N-terminal ethyl groups can be

put forward to explain the slightly slower formation of

Pb(L³)²⁺(k₃) 7.8(1) × 10⁶M^-1s^-1) with respect to

Pb(L²)²⁺(k₃) 9.8(1) × 10⁶M^-1s^-1).

LH^3-1.8(4) × 10⁹2.4 × 10^-10.1

LH₂7(3) × 10³

9 × 10^-7

0.1

2-

DOTAM (L⁴)

LH⁺1.19(4) × 10⁵1.6 × 10^-50.1

LH⁺1.36(5) × 10⁶1.8 × 10^-40.1

LH⁺9.8(1) × 10⁶1.3 × 10^-30.1

LH⁺7.8(1) × 10⁶1.0 × 10^-30.1

L¹

L²

L³

this work

^aEDTA, ethylenediamine-N,N,N′,N′-tetraacetate; HEDTA, N-(2-hy-

droxyethyl)ethylenediamine-N,N′,N′-triacetate; EDDDA, ethylenediamine-

N,N′-diacetate-N,N′-dipropionate; EGTA, 1,2-bis(2-aminoethoxyethane)-

N,N,N′,N′-tetraacetate; 1,3-PDTA, 1,3-propylenediaminetetraacetate; HEIDA,

N-(2-hydroxyethyl)iminodiacetate; NTA, nitrilotriacetate; [18]aneO₆,

1,4,7,10,13,16-hexaoxacyclooctadecane; Ac₂[18]aneN₂O₄, 1,4,10,13-

tetraoxa-7,16-diazacyclooctadecane-N,N′-diacetate; cyclen, 1,4,7,10-

tetraazacyclododecane; DOTA, 1,4,7,10-tetraazacyclododecane-N,N′,N′′,N′′′-

tetraacetate.

Eigen-Wilkins theory for a dissociative D or I_dmechanism

involving unidentate ligands.^73,74For the reaction scheme

outlined in eq 16, the rate constant of the second-order

process is controlled by the rate of solvent exchange in the

first solvation sphere of the metal cation (k_ex) and is expressed

by the product k_EW) K_OS× k_ex, assuming a steady-state

regime.

Moreover, the great sensitivity of the formation rate

constant k₃to structural factors such as the macrocyclic ring

size is also evidenced by comparing the data measured for

the 14-membered cyclam and 12-membered cyclen tetra-

acetamide derivatives L¹and L⁴. While the overall reaction

mechanism remains the same, enlarging the cyclen frame-

work of L⁴H⁺(k₃) 1.19(4) × 10⁵M^-1s^-1)²⁴by two carbon

atoms, as found in L¹H⁺(k₃) 1.36(5) × 10⁶M^-1s^-1),

induces a 10-fold increase of the complexation rate. Since

both ligands bear the same chelating side chains, this

difference reflects the enhanced rigidity and the higher steric

strain exhibited by the smaller macrocycle L⁴as it wraps

around the Pb²⁺cation. Nevertheless, this trend of higher

rates when the ring size changes from 12 to 14 atoms suffers

from numerous exceptions.^69-71For instance, no pronounced

differences in the kinetic parameters were noticed for the

insertion of Ca²⁺, Zn²⁺, or Ni²⁺in the tetraacetate analogues

of L¹and L⁴, namely TETAH^3-and DOTAH^3-.⁶⁸

K_OS

k_ex

MS_x+ L y z {MS_x‚‚‚L}

9

8 MS_x-1L + S ^+L8

MS_x-2L₂+ S ^fast8 ML_x(16)

steps

K_OScorresponds to the equilibrium constant for the fast

formation of an outer-sphere complex. In practice, estimates

of K_OSare computed with the Fuoss equation.⁷⁵The water-

exchange rate constant for Pb²⁺at 298 K has been measured

by Yasunaga and Harada.⁷⁶These authors proposed a value

of 7.5 × 10⁹s^-1based on ultrasonic relaxation experiments.

Accordingly, the kinetic parameters are expected to be

insensitive to the chemical nature of the entering ligand since

the key step is the displacement of a water molecule in the

first solvation shell. Unfortunately, a detailed discussion of

the lead-association mechanism and rates is hampered by

the scarcity of kinetic data available in the literature.

Formation rates that are pertinent to the following discussion

are reviewed in Table 6.

Identification of the actual position of the rate-determining

step is one point of interest in discussing the complex

formation mechanism with multidentate cyclic chelators.⁷²

To that end, the experimental second-order rate constants

can be compared with the limiting values predicted by the

The largest body of information has been collected for

polyaminocarboxylic open-chain chelators. Overall, these

compounds display the highest reaction rates with intrinsic

(68) Kasprzyk, S. P.; Wilkins, R. G. Inorg. Chem. 1982, 21, 3349-3352.

(69) Izatt, R. M.; Bradshaw, J. S.; Nielsen, S. A.; Lamb, J. D.; Christensen,

J. J.; Sen, D. Chem. ReV. 1985, 85, 271-339.

(70) Izatt, R. M.; Pawlak, K.; Bradshaw, J. S.; Bruening, R. L. Chem. ReV.

1991, 91, 1721-2085.

(73) Eigen, M.; Wilkins, R. G. AdV. Chem. Ser. 1965, 49, 55-67.

(74) Wilkins, R. G. Acc. Chem. Res. 1970, 3, 408-416.

(75) Fuoss, R. M. J. Am. Chem. Soc. 1958, 80, 5059-5061.

(76) Yasunaga, T.; Harada, S. Bull. Chem. Soc. Jpn. 1971, 44, 848-850.

(71) Izatt, R. M.; Pawlak, K.; Bradshaw, J. S.; Bruening, R. L. Chem. ReV.

1995, 95, 2529-2586.

(72) Elias, H. Coord. Chem. ReV. 1999, 187, 37-73.

Inorganic Chemistry, Vol. 44, No. 22, 2005 7907

Cuenot et al.

rate constants (k′_f) that vary between 0.2 and ∼6 × 10¹¹M^-1

s^-1and thus fall within the expected range for diffusion-

limited rate constants (10¹⁰-10¹¹M^-1s^-1).⁷⁷The favorable

electrostatic outer-sphere interactions experienced by the

3-fold negatively charged species tend to hasten the overall

formation rates by a factor of 10-100 with respect to water

exchange, depending upon the distance of closest approach

(a) of both partners forming the ion pair. A typical example

is provided by EDTAH^3-for which K_OS) k′_f/k_ex≈ 81.

Solving the Fuoss equation yields an approximate value of

5.1 Å for the parameter a. In Table 6 most of the data

pertaining to EDTA-type ligands correspond to reaction of

aquated lead(II) with the monoprotonated species. Rabenstein

and co-workers proposed a stepwise formation of the lead

complex that proceeds by dissociation of water molecules

followed by bond formation to yield a triscoordinated

intermediate in which one end of the partially attached ligand

is blocked by the proton residing on the unbound nitrogen

atom. Proton transfer to a carboxylate oxygen atom or to a

solvent molecule is required before the water substitution

proceeds to completion. While proton migration was initially

sought to be the rate-limiting step in order to explain the

much lower than predicted second-order rate constant found

for NTAH^2-,^78,79these authors later assumed that it should

correspond more likely to formation of one of the first three

bonds.⁸⁰Correlating the kinetic data collected for Ni²⁺, Cu²⁺,

Zn²⁺, Cd²⁺, and Pb²⁺, Margerum suggested an alternate but

more probable path in which the proton is shifted from the

ammonium to a carboxylate oxygen atom before reaction

with the aquated cation.⁷⁷Hence, proton transfer is no longer

rate determining, while anchoring to the metal occurs through

the protonated carboxylate group. Fuhr and Rabenstein also

pointed out that six-membered chelate ring formation was

significantly slower compared to five-membered ring clo-

sure.⁸⁰

the ion pair at 5 Å, K_OS) 0.315 and 0.048 M^-1for

{Pb²⁺‚‚‚L} and {Pb²⁺‚‚‚LH⁺}, respectively. The latter figure

is in excellent agreement with the experimental value

+

measured by Palanche´ et al. for the association of Fe(OH)₂

+

with the LH₆form of azotobactin δ.⁸⁴Under these condi-

tions, the product K_OS× k_exis on the order of 2.4 × 10⁹(L)

or 3.6 × 10⁸M^-1s^-1(LH⁺). In comparison, the experimental

bimolecular rate constants k₁and k₃are about 10 (L²and

L³), 250 (L¹), and 3000 times (L⁴) lower.

It is well known that attachment of ligating side chains to

a cyclic polyamine can dramatically accelerate the

complexation reaction.⁷²For lead(II) this behavior is best

exemplified by comparing the rates, despite the charge

difference, reported for cyclenH⁺(k′_f) 8.3(1) × 10⁵M^-1

s^-1)⁸⁵and DOTAH^3-(k′_f) 1.8(4) × 10⁹M^-1s^-1).⁸⁶Of some

relevance to the present study, Moreau et al. recently

summarized the prolific literature related to the metalation

kinetics of DOTAH₂.

^{2- 87}Although some mechanistic details

still remain a matter of debate, there is general agreement

on the rapid formation of an intermediate protonated species

in which a lanthanide cation is only bound by the four

carboxylate oxygen atoms located above the ring, before slow

reorganization leads to inclusion of the metal ion into the

macrocyclic cavity. Concerted proton migration from both

ammonium sites to bound carboxylate oxygen atoms and

concomitant formation of two nitrogen-metal bonds has

been envisioned as potential rate-limiting steps affording a

second intermediate.⁸⁷In the final step both carboxylic

protons are lost, while creation of the last pair of lanthanide-

nitrogen bonds requires several weeks.

Accordingly, assistance of the appended amide side chains

in the early stage of the Pb²⁺binding process is also expected

for L¹-L³, although no kinetic data are available to the best

of our knowledge for the incorporation of Pb²⁺into cyclam.

This assumption is also supported by previous observations

on neutral cyclam derivatives substituted by 2-hydroxyalkyl

groups.⁸⁸However, due to the much lower metal and proton

binding affinities of carbamoyl versus carboxylate oxygen

atoms, a mechanism similar to that mentioned above for

DOTA seems unlikely. While the first two water molecules

are certainly displaced by the entering carbonyl oxygen

atoms, subsequent interactions with the adjacent tertiary

amines bearing both involved arms enable the rapid closing

of one and then a second five-membered chelate ring.

Rahardjo and Wainwright demonstrated that at least two

dangling hydroxyethyl functions are required for accelerating

metal insertion into a cyclam framework.⁸⁹Such a sequential

arm-by-arm chelation process is also supported by the general

In contrast to open-chain chelates, formation of macro-

cyclic lead complexes is several orders of magnitude slower.

As expected, cryptands with their more rigid polycyclic

framework are less reactive.^81,82All the available data suggest

that ion pairing followed by desolvation of the Pb²⁺cation

is not rate limiting.^72,77,83Indeed, the unfavorable electrostatic

interactions expected for L¹-L⁴cannot solely account for

the second-order rate constants appearing in Table 6 that are

lower than those predicted by the Eigen-Wilkins theory.

Hence, the rate-determining step is taking place after

formation of the first metal ion-donor atom bond. Setting

the distance of closest approach of both partners forming

(77) Margerum, D. W.; Cayley, G. R.; Weatherburn, D. C.; Pagenkopf, G.

K. Kinetics and Mechanisms of Complex Formation and Ligand

Exchange; ACS Monograph 174; American Chemical Society: Wash-

ington, DC, 1978.

(84) Palanche´, T.; Blanc, S.; Hennard, C.; Abdallah, M. A.; Albrecht-Gary,

A. M. Inorg. Chem. 2004, 43, 1137-1152.

(85) Kodama, M.; Kimura, E. J. Chem. Soc., Dalton Trans. 1977, 2269-

2276.

(78) Rabenstein, D. L.; Kula, R. J. J. Am. Chem. Soc. 1969, 91, 2492-

2503.

(86) Pippin, C. G.; McMurry, T. J.; Brechbiel, M. W.; McDonald, M.;

Lambrecht, R.; Milenic, D.; Roselli, M.; Colcher, D.; Gansow, O. A.

Inorg. Chim. Acta 1995, 239, 43-51.

(79) Rabenstein, D. L. J. Am. Chem. Soc. 1971, 93, 2869-2874.

(80) Fuhr, B. J.; Rabenstein, D. L. Inorg. Chem. 1973, 12, 1868-1874.

(81) Cox, B. G.; Garcia-Rosas, J.; Schneider, H. NouV. J. Chim. 1982, 6,

397-399.

(82) Cox, B. G.; Garcia-Rosas, J.; Schneider, H.; van Troung, N. Inorg.

Chem. 1986, 25, 1165-1168.

(83) Wilkins, R. G. Kinetics and Mechanism of Reactions of Transition

Metal Complexes, 2nd ed.; VCH: Weinheim, Germany, 1991.

(87) Moreau, J.; Guillon, E.; Pierrard, J. C.; Rimbault, J.; Port, M.;

Aplincourt, M. Chem.sEur. J. 2004, 10, 5218-5232.

(88) Dey, B.; Coates, J. H.; Duckworth, P. A.; Lincoln, S. F.; Wainwright,

K. P. Inorg. Chim. Acta 1993, 214, 77-84.

(89) Rahardjo, S. B.; Wainwright, K. P. Inorg. Chim. Acta 1997, 255, 29-

34.

7908 Inorganic Chemistry, Vol. 44, No. 22, 2005

Cyclam DeriWatiWes

Scheme 2

formation mechanism of metal-peptide complexes.⁹⁰Once

the cation has been brought in the vicinity of the cyclam

ring, the fifth bond formation could involve a nitrogen atom

located on the opposite side of the already bound amide arms.

Deprotonation of the ammonium group and binding of the

appended amide substituent is supposed to occur in the last

stages, leading to the final octacoordinated complex. In all

likelihood, one of those late steps should then be rate

determining.

to a hexacoordinated Pb²⁺species. The subsequent

dissociation of the remaining donor atoms is controlled by

the first-order rate constant k^H. The value of the preequi-

1

librium constant K is sensitive to the electronic features (σ-

donation) of the macrocycle, although there is no direct

proportionality between the complex and the free ligand

protonation constants. In contrast, the first-order rate con-

stants k^H₁associated with dissociation of intermediate B does

not depend on the nature of the ligand.

Dissociation Kinetics. The proton-concentration depen-

dence exhibited by the pseudo-first-order rate constants k_obs

displayed in Figure 7 clearly illustrates the increasing lability

Alternatively, complex B can also dissociate according to

a bimolecular reaction with a second proton (k^H), directly

2

releasing a Pb²⁺ion and one molecule of diprotonated ligand

LH₂²⁺. This latter pathway becomes operative for total acid

concentrations above ca. 0.01 M and predominates at

concentrations higher than 0.02 M. The kinetic data in Table

5 show a regular decrease of the second-order rate constant

of the complexes in the order Pb(L³)²⁺< Pb(L²)²⁺

,

Pb(L¹)²⁺, which corresponds to the reverse sequence of the

thermodynamic â₁₁₀stability constants (Table 2). Moreover,

the observed rate law (eq 15) suggests a complex reaction

scheme with multiple dissociation paths involving rapid

formation of a protonated intermediate. The experimental

data support the mechanism depicted in Scheme 2.

k^Has the steric bulk of the N-terminal amide substituents

2

increases. Accordingly, Pb(L¹H)³⁺dissociates about 1.5 times

faster than Pb(L²H)³⁺or Pb(L³H)³⁺, confirming the protec-

tive role exerted by the alkyl groups against the proton attack.

According to the kinetic parameters collected in Table 5,

the less efficient pathway corresponds to the spontaneous

self-dissociation (k_d) of the octacoordinated Pb(L)²⁺com-

plexes labeled A in Scheme 2, followed by the fast

protonation of the liberated ligand. The associated first-order

rate constants k_dprimarily reflect the inertia of the complexes

in the absence of protons (solvent-assisted dissociation). As

expected, the k_dvalue was the highest for the less stable

Pb(L¹)²⁺species (k_d) 9(4) s^-1), while it became statistically

insignificant in the case of Pb(L³)²⁺.

The second demetalation route assumes the fast protona-

tion of complex A which is in equilibrium (K) with a

monoprotonated intermediate B. Since amines are much

stronger bases than amidic oxygen atoms, it is assumed that

the entering proton first breaks one metal-nitrogen bond,

giving rise to an ammonium site.⁶⁰It is anticipated that the

protonated nitrogen atom will move away from the coordina-

tion center as much as possible due to charge repulsion,

inducing the complete dechelation of the aminocarbamoyl

arm for steric reasons. Intermediate B is therefore assimilated

In related work the same rate law and mechanism were

deduced for the acid-promoted dissociation of a pentacoor-

dinated 1,4,8,11-tetraazacyclotetradecane-1,8-bis(methylphos-

phonato)copper(II) complex by Kotek et al.⁶⁰and of the lead

complex formed with 1,4,7,10-tetraoxa-7,16-diazacyclo-

octadecane-N,N′-diacetic acid by Laing et al.⁹¹The latter

authors reported similar kinetic parameters to those found

herein (K ) 100(40) M^-1, k^H₁) 18(11) s^-1, and k^H₂) 9(4)

M^-1s^-1; T ) 298 K; I ) 0.2 (Li,H)ClO₄). In contrast, the

lifetime of Pb(DOTA)^2-(k_obs) 6.15 × 10^-5s^-1)⁸⁶is much

longer in acidic conditions ([H⁺] ) 0.02 M; I ) 0.1 M; T )

298 K) than those measured for the cyclam-based tetraamides

(k_obs) 277 s^-1for Pb(L¹)²⁺). This difference highlights the

extraordinary inertness generally observed for all metal

complexes incorporating a cyclen unit appended with acetate

groups.^92-94Moreover, the presence of additional protonation

(91) Laing, J. L.; Taylor, R. W.; Chang, C. A. J. Chem. Soc., Dalton Trans.

1997, 1195-1200.

(92) Toth, E.; Bru¨cher, E.; Lazar, I.; Toth, I. Inorg. Chem. 1994, 33, 4070-

4076.

(90) Metal Ions in Biological Systems. Concepts on Metal Ion Toxicity;

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J. F. Inorg. Chem. 1992, 31, 1095-1099.

Sigel, H., Ed.; Marcel Dekker: New York, 1986; Vol. 20.

Inorganic Chemistry, Vol. 44, No. 22, 2005 7909

Cuenot et al.

sites also has some important consequences from a mecha-

nistic point of view. As shown by Pippin et al.,⁸⁶formation

of the Pb(DOTAH) intermediate (K ) 2.0(5) × 10³M^-1)

most probably involves the initial protonation of a carboxyl-

ate oxygen atom. The proton then transfers to an amine (K

) 5.1(5) M^-1) before a second proton is rapidly taken up.

The rate-limiting step has been assigned to dissociation of

the neutral Pb(DOTAH₂) complex, which is characterized

by a first-order rate constant of 6.7(3) × 10^-4s^-1. Since

amide groups are much more acidic than carboxylates, it is

anticipated that Pb(L⁴)²⁺would be even more inert,²⁴this

assumption being supported by the fact that this complex

keeps its integrity in 0.5 M HCl as stressed by Maumela et

al.,⁹⁵whereas Pb(DOTA)^2-is fully decomposed at pH values

below 2.

determining step. The two first bond formations likely

involve an oxygen atom from two amide groups, which are

immediately followed by chelate ring closure through

nitrogen-lead interactions. The lower intrinsic reactivity of

the acetamide ligand L¹H⁺compared to the tertiary amides

L²H⁺and L³H⁺is easily understood by considering that

intramolecular hydrogen bonds are preserved in the mono-

protonated state. However, due to the lower basicity of L¹,

the apparent pseudo-first-order complexation rate constants

are about 10 times higher than those measured for L²and

L³in the investigated p[H] range 4-7. Finally, the proton-

mediated dissociation studies of the complexes evidenced a

fast release of the lead cation in a single rate-limiting step.

The observed rate law supports a demetalation mechanism

with three parallel pathways. Besides spontaneous self-

decomplexation of the reagent, the kinetics is controlled in

weakly acidic conditions by the first-order dissociation of a

rapidly formed monoprotonated intermediate and at higher

acid concentrations by the bimolecular attack of a second

scavenging proton.

Conclusion

A combination of potentiometric and absorption spectro-

scopic data allowed us to characterize the formation and the

dissociation mechanism of octacoordinated lead complexes

of primary and tertiary tetrakiscarbamoylmethyl cyclam

derivatives. In solution, the Pb²⁺cation is entrapped in a

molecular cage constituted by the four macrocyclic nitrogen

and the four amidic oxygen atoms. The weaker proton and

lead binding constant found for L¹in comparison to L²and

L³can be explained by the existence in solution of hydrogen

bonds between each lone pair of electrons belonging to an

amine and one terminal NH₂hydrogen atom of the appended

substituent, as found in the crystal structure for the free-

base macrocycle L¹. However, changes in the lead complex

stabilities and ligand basicities are integrally compensated,

so that all three macrocycles exhibit a similar uptake capacity

for identical proton, metal, and ligand concentrations. The

complexation reaction proceeds according to an Eigen-

Winkler mechanism with a single, albeit late, rate-

Acknowledgment. This paper is dedicated to Dr. Anne-

Marie Albrecht-Gary on the occasion of her birthday. The

work was carried out within the framework of the French

governmental project “Re´duire la teneur du plomb dans

l’eau” (contract no. 99V0367-Ministe`re De´le´gue´ a` la

Recherche et aux Nouvelles Technologies and Ministe`re de

l’Equipement, des Transports et du Logement) and also

supported by the Centre National de la Recherche Scienti-

fique (CNRS) and the Conseil Re´gional de Bourgogne.

F.C. thanks the Ministe`re De´le´gue´ a` la Recherche et aux

Nouvelles Technologies for a Ph.D. scholarship. We also

acknowledge A. Bucaille for providing some compounds.

Supporting Information Available: ¹H-¹H NMR COSY

spectrum of [Pb(L²)](ClO₄)₂in DMSO-d₆, tables of the pseudo-

first-order rate constants (k_obs) for the formation and proton-assisted

dissociation of Pb(L¹)²⁺, Pb(L²)²⁺, and Pb(L³)²⁺, and X-ray

crystallographic files in CIF format for the structure determinations.

This material is available free of charge via the Internet at

http://pubs.acs.org.

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IC0508019

7910 Inorganic Chemistry, Vol. 44, No. 22, 2005

Article Doi

Synthesis, characterization, and x-ray crystal structures of cyclam derivatives. 8. Thermodynamic and kinetic appraisal of lead(II) chelation by octadentate carbamoyl-armed macrocycles

DOI: 10.1021/ic0508019

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Article abstract of DOI:10.1021/ic0508019

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