- CLIN. CHEM. 32/10, 1797-1806 (1986)

CLINICAL CHEMISTRY, Vol. 32, No. 10, 1986 1797

The Chemistry of Aluminum as Related to Biology and MedicineR. Bruce MartIn

The increasing number of roles discovered for Al34 inphysiological processes demands an understanding of howAl3’ interacts with compounds in biological systems. Al3” isexpected to complex with oxygen donor ligands, especiallyphosphates, and it does so in soils, in the gastrointestinaltract, and in cells. The stability of Al3’ complexes hasgenerally been misjudged because of lack of recognition thatfree, aqueous Al3” Is not the dominant form in neutralsolutions and that the solubility of Al(OH)3 limits the free Al3’at the plasma pH 7.4 to less than 1O� mol/L. In thepresence of inorganic phosphate, the permitted free Al34 isdecreased further, through formation of insoluble aluminumphosphate. This precipitate facilitates the elimination of Al3’from the body. In contrast, citrate solubilizes Al34, and anappreciable fraction occurs as a neutral complex that maypass through membranes and provide a vehicle for Al34absorption into the body. In the blood plasma the most likelysmall-molecule complex is that with citrate, while the onlycompetitive protein complex is that with transferrin, a proteinbuilt to transport Fe3’ but whose sites are only 30% occu-pied.

AdditIonal Keyphrases: kidney disease . Al-induced disorders

complexes of Al with various anions, compounds Ionic spe-

cies of Al . metal-ion buffer systems pAl . stability con-stants ligand protonation complex deprotonation p/�cie-pliate citrate . transferrin . ligand exchange rates In-testinal absorption of dietary/medicinal Al . fluoride proteins

hemodialysis effect on bone

Aluminum has recently been recognized as a causativeagent in dialysis encephalopathy, osteodystrophy, and rni-crocytic anemia occurring in patients with chronic renalfailure who undergo long-term hemodialysis (1). Only asmall amount of Al34 in dialysis solutions may give rise tothese disorders. Encephalopathy has also occurred in chil-dren consuming Al(OH)3 as a phosphate binder for renaldisorders (2, 3). (Ca2” would have been a more naturalchoice.) A13’ has also been implicated in neurotoxicityassociated with amyotrophic lateral sclerosis, a form ofparkinsonism with severe dementia, in the indigenous pop-ulation of Guam, where the soils are high in Al34 and low inMg2’ and Ca24 (4), and in Alzheimer’s disease (5).

These developments have led some investigators to per-form experiments with Al34 without appreciating its solu-

Chemistry Department, University of Virginia, Charlottesville,VA 22901.

Received April 11, 1986; accepted June 23, 1986.

tion chemistry. But certain questions require consideration.Can one add Al34 in a 1 mniol/L concentration to a neutralsolution and expect the aqueous Al34 concentration to be 1mmol/L? Is it appropriate to use phosphate buffers withAl34? In addition to equilibrium considerations, are rateeffects important? Does the order of mixing of Al34 withEDTA and F make a difference? In the plasma what is thedominant Al31 binder among small molecules and amongproteins? This article aims to develop the bioinorganicchemistry of Al34 so that answers to such questions becomeevident to readers.

Healthy people consume Al34 in a variety of ways.Fortunately, most foods do not dissolve significant amountsof aluminum from cookware. However, the attack of foodjuices on aluniinwn vessels varies greatly, depending uponpH, temperature, and other substances present. Hot, acidicfruit juices-in the absence of sugar-.corrode aluminumware, as do salt solutions used in pickling processes (6).These solutions and conditions should probably be avoidedby the prudent homemaker. Al34 is a component of alum-containing baking powders, more widely used in the U.S.A.than in Europe. Some antacid preparations contain Al3”.

Since Roman times or earlier, alums and the relatedA12(S04)3 have been added to drinking water to improve itsappearance. Alums are double sulfate salts of Al34 and Na4,K4, or NH�’ such as KAI(S04)2 12H20. Commercially,alums are added to foods such as frozen strawberries,maras�hino cherries, and pickles to improve their appear-ance. Al34 salts are often added top�s� cheeses and tobeer. Al2(S04)3 is the most widely used coagulant forclarif�ring turbid drinking water. The success of this treat-ment depends mainly upon precipitation of Al(OH)3, withadsorption of the turbidity. Excess Al34 remains in solution.Thus a “dirty” water that contained some sediment (proba-bly harmless) and no or little Al34 becomes clear andcontains no or little sediment and some Al34. It has beensuggested that ingestion of lead from plumbing contributedto the demise of the Roman empire (7). Could it have beendue to A13� instead?

AI3� Complexes

The only accessible oxidation state for aluminum inbiological systems is 3+. Binding of Al34 is primarilyelectrostatic (as opposed to covalent) and therefore, in addi-tion to charge, ionic size is an important parameter. Effec-tive ionic radii in angstrom units (1 A = 10’ rim) in sixfoldcoordination follow in parentheses after the ion: Be2” (0.45),A13� (0.54), Ga34 (0.62), Fe34 (0.65), Mg24 (0.72), Zn2”(0.74), and Ca24 (1.00). Be24 (0.27) is smaller in its favored

pH

6

1198 CLINICAL CHEMISTRY, Vol. 32, No. 10, 1986

fourfold coordination and Ca2� (1.12) larger in its favoredeightfold coordination (8). Excluding the unnatural Ga34,the radius of Al3” most resembles that of Fe34. Thusappearance of AL3” in Fe34 sites seems likely. Except withF, Al3� forms weaker complexes than Fe3”’. The binding ofAl3� and Fe3” to transferrin is discussed later.

Though Mg is somewhat larger than A13”, displace-ment of the ubiquitous Mg24 in biological systems by Al34appears likely. Mg2� is often associated with phosphategroups, and comparison suggests that Al3” should also seekthose sites. In many physiological systems the Mg24 concen-tration is about 1 to 2 mmoliL. Al34 binds almost i0� timesmore strongly to ATP� than does Mg24 (see below). Thus inthese systems, nanomolar amounts of Al34 can competewith Mg2� for the phosphate sites.

Ca2� not only is much larger than Al34, it also frequentlyoccupies specific sites in proteins (9). In its favored eightfoldcoordination the volume of Ca2� is nine times greater thanAl3’ in sixfold coordination. Since hole production is ener-getically costly, Al34 cannot replace Ca2’ in proteins with-out substantial readjustment of liganding groups. Thuscompetition between Al3” and Ca2” is less apt to be forprotein binding sites than for small molecule ligands andphosphate, with which both form insoluble complexes.

In biological systems we anticipate that Al3” will asso-ciate with oxygen donor ligands, at least some of whichshould be anionic to help counter the 3+ charge on thecation. Carboxylate and phosphate groups, inorganic phos-

phate, nucleotides, and polynucleotides meet this prescrip-tion. Unless carboxylate groups are arranged to makestrong chelation possible, as in citrate (see below), Al3”'should prefer phosphate binding. Because it carries but onecarboxylate group, lactate forms only weak complexes withAl3’. The stability constants remain unknown, witnessingto the weakness of the interaction. Unless other ligands arepresent, dissolution of Al3” as Al(lactate)3 should yieldhydroxo complexes as the complex dissociates and weaklybound lactate diffi.ises away. A prominent peak in an NMRspectrum of a pH 6 solution containing 10 mmol of Al3”' and30 mmol of lactate per liter was assigned to the 3:1 complex(10). Under these experimental conditions, less than halfand possibly much less of the metal ion is expected to bepresent in a 3:1 complex.

Al3” forms only weak complexes with amines and sulfhy-dryl ligands. There is almost no tendency for Al34 tocomplex to sulfhydryl groups. Metal ion hydrolysis inter-feres well before A13” might coordinate to uni-dentateamities. With multi-dentate amino-carboxylate ligands suchas EJYFA, Al34 forms strong complexes (11).

AI3� In Aqueous SolutIons

Whatever other ligands may be present, understandingthe state of Al34 in any aqueous system demands awarenessof the species that Al34 forms at different pH values withthe components of water. In solutions more acid than pH 5,Al34 exists as the octahedral hexahydrate, Al(H20)63”,often abbreviated as Al3t As a solution becomes less acid,Al(H20)634’ undergoes successive deprotonations to yieldAl(OH)2” and Al(OH)2”'. Neutral solutions give an Al(OH)3precipitate that redissolves in basic solutions, owing toformation of tetrahedral Al(OH)[. Polynuclear species mayalso form, their compositions being time dependent (12).(Since this article deals with relatively low concentrations oftotal Al34 in biological systems in the presence of otherligands, polynuclear species are not considered.)

Equilibria among mononuclear Al3” species in aqueoussolutions may be described by reactions 1-3 below. Forconvenience we abbreviate the hexahydrate occurring inacid solutions as Al3”'. This abbreviation disguises theactual reaction, deprotonation of Al3 “-bound water. Al-though the first deprotonation reaction is more accuratelywritten as

Al(H20)63” = H20 # Al(H20)5(OH)24 + H30”,

by common consent we write it as the hydrolysis reaction

Al3’ + H20 � Al0H24 + H4K1� = (ff’)[Al0H2”i/[Al341 = (1)

where the parentheses signify activity, the brackets signifyconcentration and Kia is the associated equilibrium con-stant. This reaction is followed by deprotonation from asecond Al3’-bound water molecule.

Al0H24 + H20 � Al(OH)2”' + H’= (H”)[Al(OH)2”j/[Al0H24] = 10_56 (2)

Little evidence favors significant amounts of solubleAl(OH)3 in solution, but deprotonation from two morebound waters yield a soluble tetrahydroxo species.

Al(OH)24 + 2H20 � Al(OH)[ + 2ff”K� (H”)2[Al(OH)4]/[Al(OH)241 = 10_121 (3)

The values for the equilibrium constant at 25 #{176}Cassigned to

each of the three reactions (and reaction 5 below) areconverted from a set of thermodynamic values (13) by adeveloped protocol (14) to 0.16 ionic strength.

Figure 1 shows the distribution of soluble, mononuclearaluminum ion species in aqueous solutions. The octahedralhexahydrate Al(H20)634 dominates at pH <5, and the

tetrahedral Al(OH)4 at pH >6.2, while there is a mixtureof species from 5 <pH <6.2.

To find the relation between the free, non-hydrolyzedaqueous Al34 and the total soluble aluminum ion concentra-tion, we define the mole fraction of nonhydrolyzed aqueousions as -y = [Al3” I/CM where

CM = [Al34] + [A1OH2”] + [Al(OH)2”] + [Al(OH)4]

Combination with equations 1-3 yields

lIy = 1 + [10551(H4)] + [10”’/(H’)2] +

[10�2/(H”)4] (4)

The mole fraction y is pH dependent and must be calculatedfor each pH. Substitution into equation 4 yields, at pH 6.5,

Fig. 1. Distribution of soluble, mononuclear aluminum ion species inaqueous solutions�na�: mole fraction of aluminum ion occumng as each designated species. Atany pH the indMdual mole fractions sum to unity

.4-J0

z

D-j

4-J.4I-0I-

(00

pH

Fig. 2. Negative logarithm of total molar concentration of aluminumallowed by Al(OH)3 solubility vs pHLower curve represents true equilibilum solubility from gibbets. Upper curvedepicts representative solubility from amorphous AI(OH)3. AP� is the predomI�nant soluble aluminum species at pH <5 and Al(OH)4 at pH >62, where theminimum solubility occurs for both curves. From 5 <pH <6.2 there is a mixture ofsoluble species, as shown in Fig. 1

CLINICAL CHEMISTRY, Vol. 32, No. 10, 1986 1799

1/y = 720 and, at pH 7.4, 1/y = 2.5 x 106. Because at pH 7.4virtually all soluble aluminum ion occurs as Al(OH)4, thevalue of 1/y also gives the molar ratio of[Al(OH)4]I[Al34] =

2.5 x 106.To this point we have described the equilibria and distri-

bution among soluble, mononuclear aluminum ion specieswithout considering the absolute amounts permitted by thelimited solubility of Al(OH)3. At reasonable temperaturesthe stable crystalline phase of Al(OH)3 is the mineral calledgibbsite. The solubility of Al(OH)3 from solid gibbsite maybe described as

Al(OH)3(gibbsite) + 3H”' �± Al3”’ + 3H20K�1’ = [Al3”'J/(H”')3 = 1092

The gibbsite solution reaction could have been written asgiving Al34 + 30H”, and the solubility product constantexpressed correspondingly. By using only H” and not 0Hin reactions we avoid the need to calculate the OH- concen-tration or activity, and for the H”' activity we useHowever, equilibrium is slowly achieved with gibbsite, and

the solubility with respect to amorphous Al(OH)3 may be upto 100-fold greater. The presence of organic ligands such ascitrate favors formation of non-crystalline Al(OH)3 (15).Because there rarely is equilibrium with respect to gibbsitein biological systems, we employ a more liberal equilibriumconstant for solubility of amorphous Al(OH)3:

K�1 = [Al34]/(H4)3 = i0’#{176}’� (6)

Figure 2 shows the total molarity of all soluble aluminumspecies permitted by both gibbsite and a representativeamorphous Al(OH)3 according to equations 5 and 6. Theminimum solubility in both curves occurs at pH 6.2.

From equation 6 we may estimate the highest permittedhexahydrate Al3” concentration from amorphous Al(OH)3as [Al(H2O)634] = 10107 x 103pH At pH 7.4, the pH ofextracellular fluids such as blood plasma, we obtain[Al(H20)63’J = 10_h15 molJL. This important result meansthat, at the pH of blood plasma, the highest obtainable freeAl3” concentration allowed by an amorphous Al(OH)3 is 3 x10’.’12 mol/L. The negative logarithm of this value, designat-ed as pAl (analogous to pH), appears in the first row of Table1. The corresponding total aluminum concentrations isgiven by CM = [A13”']/y. At pH 7.4 we have already foundthat 1/y = 2.5 x 106, so that CM = (3 x 10�) x (2.5 x 106)

= 8 x 106 mol/L = 8 �mol/L as the permitted total

Table 1. MaxImum Free Al3’ Molar ConcentratIonsExpressed as pAl = - Iog[A134]

Complexon

Amorphous Al(OH)3Al(OH)2H2P04Citrate, 0.1 mmol/LdTransferrin d,

pH 4.0 pH 6.6 pH 7.4

1.3 9.1 11.56.5a 125b 12.9c8.3 13.1 14.0

14.6

1 b 10, and C2 mmol of total phosphate per liter, d �mol of total At3* per

liter. ‘Under plasma conditions with 50 �mol of unoccupied sites per liter.

aluminum ion concentration from an amorphous Al(OH)3.(5�� This result deserves emphasis. For even though the permit-

ted total Al3� at pH 7.4 may reach 8 pinolJL, most appearsas Al(OH)4” and only 3 x 10’12 molJL as Al(H20)634. If thegibbsite solubility product constant were used, both allowedconcentrations would be 1/30 as great.

Like any other ligand, hydroxide ion, by reactions 1-3 and6, withdraws Al34 from solution. In aqueous solutions,regardless of the other ligands present, reactions 1-3 occur

and the species distribution shown in Figure 1 prevails.These equilibria must be considered in all solutions contain-ing Al34. Unless a solution is supersaturated with respect toamorphous Al(OH)3, greater than nanomolar concentra-tions of free Al3” in neutral solutions are unobtainable.Upon addition of 1 mmol of an Al3”’ salt per liter to asolution at pH 7.4 the free Al3”' concentration is not 1mmol/L but only about a miniscule 3 x 10� mol/L. Thepredominant water-derived complex is Al(OH)4 at 8

MmoIJL. Unless the remainder of the added Al34 has beencomplexed by other ligands, it will form insoluble Al(OH)3(Figure 2). When Al3” binds to other ligands or proteins,Al34, not Al(OH)[, is bound, and it is the free Al3” orAl(H20)63” concentration rather than the much greaterAl(OH)4” concentration that is the significant quantity inneutral solutions.

Because they fail to incorporate the basic ideas describedin this section, many papers in the literature reach dubiousconclusions. Dissociation constants for Al34 binding that arenear to or greater than the free Al3”' concentration allowedby the solubility of Al(OH)3 are suspect. In a study of Al3”binding to the important calcium regulatory protein calmo-dulin, the authors performed equilibrium dialysis experi-ments at pH 6.5 and calculated the binding constants fromthe presumed total Al3”’ in solution (16, 17). From equation4 the free [Al3”'] is only 1/720 that of total [Al3”], and so theirbinding constants must be increased by a factor of 720.However, this investigation was performed near the mini-mum of Al(OH)3 solubility (Figure 2), where for the amor-phous form equation 6 allows 1.6 nmol of free hexahydrateAl3” per liter, corresponding to 1.1 �Lmol of total Al34(mainly Al(OH)4, Figure 1) per liter. Based on total Al34,their dissociation constants span a range from 0.1 to 1.2.tmolJL (16, 17), corresponding to solubilities between those

for amorphous Al(OH)3 and gibbsite in Figure 2. Thiscomparison renders suspect the conclusion that calmodulinbinds three Al3”’ ions so strongly. The result ncieds to beverified at a pH removed from the minimum in Al(OH)3solubility and with an appropriate metal-ion buffer systemto control the free Al3”’ concentration (see below). On astructural basis it seems most unlikely that calcium-calnio-dulin can bind three Al34 ions at anywhere near thestrengths proposed in these papers.

Suspect also are inhibition constants in the miuimolarrange for acetylcholinesterase (EC 3.1.1.7) activity derivedfrom addition of millimolar amounts of total Al3” at pH 7.5

1800 CLINICAL CHEMISTRY, Vol. 32, No. 10, 1986

(18), where Figure 2 shows a solubility of only 10 pmol oftotal Al3” per liter, even from amorphous Al(OH)3.

CondItIonal StabIlity Constants

Ligand Protonation

Stability constants for multi-dentate ligands with aminogroups overstate their effective binding strengths in neutralsolutions. Amino groups are protonated in neutral solutions,and tabulated stability constants refer to the deprotonatedligand (11). Nitrilotriacetate (NTA) is taken as a specificexample; the model is easily transferable to citrate.

For nitrilotriacetate, L3, the stability constant, K,, refersto the reaction

Al3”' + L3� AlL#{176} K, = [A1L#{176}].’[Al3”’][L3’1

where log K, = 11.4 (19). In neutral solutions most of theligand appears as a species with a 2- net charge. We write

H”' + L3’ Ka = (H4) [L31/[HL2’i

where PKa = 9.58.For NTA in neutral solutions we have the displacement

reaction

Al34 + HL2 � H4 + AlL#{176}

The concentration of deprotonated ligand L3” available tothe metal ion is reduced by occurrence of protonated species.The fraction of unbound, deprotonated ligand is given by a= K,/[(H4)+Kj so that 0 <a <1. The conditional pH-dependent stability constant is given by K’� = aIC,, or

log K’� = log K, - P1�a - log [(H4)+KaI

In the limit where pH � piCa, unbound ligand is predomi-nantly in its basic form and log K’� = log K,. For pH 4 pK5unbound ligand is predominantly protonated and log K’,, =

log K, -pK� + pH. When the pH is within two log units ofpK5, the complete equation 7 should be used. For NTA at pH7.4 we have logK’� = 11.4 - 9.6 + 7.4 = 9.2. The conditionalstability constant provides a measure of complex stabilityunder specific pH conditions where there is a protonatedligand. By allowing for withdrawal of deprotonated ligandfrom solution by protonation, the value of a conditionalstability constant becomes less than that for the standardstability constant.

Complex Deprotonation

The formulation in terms of the conditional stabilityconstant described so far may be incomplete if the complexitself undergoes one or more deprotonations at a pH near toor less than the pH of interest. Al34 complexes of both NTAand citrate deprotonate in acidic solutions. Continuing withthe NTA description, we write

AIL#{176}+ H20 � HOAIL + H”Kb (H”)[HOAILJ/[AlL#{176}]

where pKb = 5.2(19). The reaction refers to proton loss froman Al34-bound water molecule in the complex. [More appro-priately, the left-hand side of the reaction could be writtenwith the single reactant (H20)AlL#{176}.] The deprotonationreaction increases the effective stability constant.

We enlarge the scope of the conditional stability constantdescribed above to include complex deprotonation. Thefraction of non-deprotonated complex is given by 13 =

(H”')/[(H4) + Kb], so that 0 <13 <1. The conditional, pH-dependent stability constant is now given by K,, = aK,/(3 or

logK,, = logK, + log a + pH + log [(H4) + Kb] (8)

where log a is evaluated as in equation 7.For NTA at pH 7.4 the conditional stability K,, = [HOAl-

L”1/([Al3”i[HL�”]) refers to the overall reaction

Al34 + HL2� �± HOAIL” + H” (9)

We now have logK,, = 11.4 - 9.6 + 7.4 + 7.4 - 5.2 = 11.4.By coincidence, the conditional and tabulated stability con-stants are numerically equal at pH 7.4 only, because thestability-promoting effect of a deprotonated complex exactlyoffsets the destabilizing effect of an appreciable fraction ofmonoprotonated unbound ligand.

For citric acid the Al34 complex also loses a proton, butfrom the citrate ligand according to AlC#{176}� AlCH.,,1 + H4,where pKb = 3.4 (20). The three successive pK, values forcitric acid relevant to biological systems are 3.0,4.4, and 5.8,so that, at pH 7.4, a = 0.975 and log K’,, = log K, = 8.0(20).For the overall reaction occurring at pH 7.4

Al34 + C3’ �:± AlCH1” + H4

the equilibrium constant K,, = [AlCH1]/([Al34][C3”]) andlog K,, = 8.0 + 7.4 - 3.4 = 12.0. In the case of the Al34-citrate complex at pH 7.4 the effective stability constant isiO� times greater than the tabulated stability constant,owing to deprotonation of the complex in quite acidicsolutions.

It is occasionally useful to expand further the concept ofconditional stability constant to allow for metal-ion hydroly-sis, because only the nonhydrolyzed, aqueous metal ion isconsidered to be able to form complexes. We have alreadydefined the mole fraction of nonhydrolyzed, aqueous ion asy, and for Al34 its value at any pH may be calculated fromequation 4. The conditional stability constant that allows forall possibilities is K’,, = ayK,//3.

An apparent or conditional stability constant of log K’,, =

6.2 was determined by a kinetic method for Al34 binding toATP at pH 6.95(21). We find a from PKa =6.5,13= 1,7 fromequation 4, and calculate the standard stability constantlogarithm from log K, = log K’,, - log a + log 13- log ‘y = 6.2+ 0.1 + 0.0 + 4.6 = 10.9. This last value refers to binding ofnonhydrolyzed aqueous Al34 to ATP�. There is a substan-tial 4.6 log unit allowance for Al34 hydrolysis. In compari-son, for ATP4’ and Mg2”' we have log K, = 4.2 (22), 6.7 logunits weaker than for Al3”'.

In terms of the standard stability constant, the binding ofAl3” to ATP4’ is significantly stronger (log K, = 10.9) thanto citrate (log K, = 8.0). Mainly because the citrate stabilityconstant fails to allow for complex deprotonation, the condi-tional stability constant at pH 7.4 for citrate becomes log IC,,= 12.0 as shown above. The corresponding conditionalconstant at pH 7.4 for APP4’ is nearly the same as thestandard constant, log K,, = 10.8 (note that this is K,,, notK’,,). This comparison indicates that citrate complex depro-tonation results in a reversal of the standard order forstability constants, and we predict that citrate will extractAl3”’ from ATP4’ in neutral and weakly acidic solutions.Experimentally, citrate has been used to remove Al3”' fromATP (23).

Whether the standard or conditional stability constantsare more useful depends upon the situation. In tabulations,pH-independent standard stability constants are usuallyused. On the other hand, the practicing chemist may arguethat, because the pH region of overlap of Al(H20)634 (Figure1) and APP4’ with pK5 = 6.5 for ATPH3 is narrow to non-existent, a conditional constant is more practical. However,

CLINICAL CHEMISTRY, Vol. 32, No. 10, 1986 1801

conditional constants are pH-dependent, and so must becalculated for each pH. It is simpler to calculate a condition-al stability constant from a standard stability constant thanfrom another conditional constant calculated for anotherpH. In addition, only standard stability constants give anappropriate comparison between the binding affinities oftwo different metal ions, as the comparison of Al34 and

Mg2”’ demonstrates.Similar to the alkali and alkaline-earth metal ions,

binding of Al34 to nucleoside triphosphates such as APPoccurs mainly at the phosphate chain, with insignificantinteraction at the nucleic bases. Transition metal ions alsointeract at the nucleic bases (24, 25). Nuclear magneticresonance spectra indicate that aqueous Al3”' ions formseveral complexes with the phosphate group of APP andundergo exchange on a millisecond time scale (26). Directbinding of Al34 to N7 on the adenine ring was proposed forone complex on the basis of an upfield shift of H8. However,metal-ion binding at N7 always produces downfield shifts atH8 (27). An upfield H8 shift is attributed to nucleic basestacking or deprotonation at Ni.

Consistent with the localization of Al3”' in the chromatinin the cell nucleus of neurofibrillary tangles (4, 28), theslowly exchanging Al3”' forms several complexes with DNA(29). Some of the complexes may be multinuclear with twoor more Al34 joined by hydroxo bridges. The detailedstructures for the Al3”'-DNA interactions remain to bedefinitively specified.

In pH regions similar to those in which a metal iondeprotonates a water hydroxy group, the metal ion mayinteract with and deprotonate an alcoholic hydroxy group.The ribose sugar contains a pair of cis 2’,3’-hydroxy groupsthat may advantageously form a chelate with an appropri-ate metal ion (24). In nucleoside phosphates, Al34 prefersthe basic phosphate site. In nucleic acid polymers, however,the negatively charged phosphate on each residue is notbasic. In RNA the cis-hydroxy groups of ribose provide apossible binding site at pH >5 for Al34, which has littletendency to coordinate to nitrogens of the nucleic bases.

For years it has been known that certain metabolites suchas phosphate and citrate activate yeast- and brain-hexoki-nase (EC 2.7.1.1) enzymes at pH s7. The metabolites thus

become implicated in regulating hexokinase function. Re-cently, however, it has been shown that the loss of hexoki-nase activity at pH �7 is ascribable to Al34 contaminationof commercial preparations of APP (23,30). Phosphate andcitrate “activate” by complexing Al34 and freeing the APP.The variable Al34 contamination is low, of the order of onlymole per hundred moles, and is the most common metal-ioncontaminant of APP preparations (23,30,31). Inhibition byAl34 shows up because the inactive APP-Al3”’ complexbinds to hexokinase about i0� times more strongly thandoes APP-Mg2”'. This enormous hexokinase binding advan-tage in favor of APP-Al34 is the really extraordinaryconclusion of these investigations. If other nucbeoside phos-phate-Al3� complexes bind in a similarly strong fashion toeven a fraction of the wide range of enzymes with nucleosidephosphate substrates or cofactors, then Al34 (and othermetal ions?) becomes potentially toxic at low concentrationsby inhibiting many key metabolic processes.

Polymerization of tubulin to microtubules benefits fromthe presence of guanosine-5’-triphosphate and Mg2”’. Recentresearch establishes that less-than-nanomolar concentra-tions of Al34 promote polymerization, making Al3”' iO�times more effective than Mg�”’ (32).

Metal-Ion Buffers

In conducting experiments with metal ions it is oftennecessary to fix their concentrations reliably at knownvalues. Metal-ion buffers are analogous to pH buffers exceptthat it is the free metal-ion concentration that is controlledin the presence of excess ligand. As with protonic equilibria,the buffering is most effective when the ratio of complexedligand to total ligand (molar concentrations) lies between0.15 and 0.85. Because ligand concentration exceeds that oftightly bound metal ion, the total metal ion/total ligandmolar ratio, R, should also lie between 0.15 to 0.85. Notincluded in the treatment are 1:2 Al34 to NTA or citratecomplexes, because analysis indicates that they do not formto an appreciable extent with ligand in concentrations up to10-2 molIL.

The development of the previous section in terms of theconditional stability constant, K,,, permits us to formulatedirectly the equation for metal ion buffering. For the NTA-Al34 system of equation 9 the total ligand concentration CL= [ffl)] + [HOA1L”] and the total Al34 concentration CM= [HOA1L”] as ligand occurs in excess and the metal ion istightly bound so that the concentration of free Al3”’ isnegligible. Substituting in the expression for the conditionalstability constant, rearranging, and noting that R = CM/CLwe obtain

[Al34] = R/(i - R)K,, (10)

This result is a general one. The free metal-ion concentra-tion is given by equation 10, where the conditional stabilityconstant applies to the specific pH for which it was calculat-ed. For systems formally analogous to NTA and citrate, K,, is

given by

K,, = K,[1+Kb/(H”’)I/[i+(H”')/KaI (11)

The (H”)/Ka term accounts for protonated ligand and theK�,/(H4) term for deprotonation of Al3 “‘-bound water in thecomplex.

The important metal-ion buffer equations 10 and iiexpress the free metal-ion concentration as a function ofknown equilibrium constants, pH, and only the ratio (R) oftotal metal ion to excess total ligand concentrations and not

their absolute concentrations. To vary the free metal-ionconcentration at a fixed pH, it is necessary to vary R. At afixed pH the right-hand side of equation 10 describes asigmoidal curve symmetrical about R = 0.5, similar to atitration curve. For the midpoint at R = 0.5, [Al34] = i/K,,.At pH 7.4 and R = 0.5 we obtain, for NTA, pAl = -log[Al34] = 11.4 and, for citrate, pAl = 12.0. Both metal-ionconcentrations are very low, but controlled. For the same Rvalue the pAl for the two buffer systems differs by only 0.6log units, or a factor of 4, in [Al34]. If one system applies, theother should apply also, and experiments with the twodifferent kinds of ligands should serve as a check on oneanother. Both buffer systems were used in this way in arecent evaluation of the stability constant for Al3”’ bindingto a serum protein, transferrin (33).

It is reasonable to vary R from 0.15 to 0.85 to obtain arange of 1.5 log units in pAl. If a wider range is desired, it isdesirable to vary pH or find another ligand for a new metalion buffer system. The condition where R = 1 or CM = CLshould be avoided, because the solution is unbuffered withwide swings in the value of pAl for even small deviationsfrom exact equality (analogous to the endpoint in a titrationcurve).

Al(OH)2H2P04 � Al3”' + 20W + H2P04”

1802 CLINICAL CHEMISTRY, Vol. 32, No. 10, 1986

With allowance for an additional ammomum group depro-tonation in the free ligand, the final equation for EDTAcomplexation resembles that for NTA. Substitution of theEDTA stability constants (11) yields at pH 7.4 and themidpoint of the metal-ion buffer range, pAl = 15.1. The freeAl34 concentration in the presence of EDTA is 5000 timesless than that in the presence of NTA. However, equilibriumin the ED’FA complex of Al34 is reached slowly, andexperiments with it need careful monitoring. The [ethylene-bis(oxyethylenemtrilo)]tetraacetic acid (EGTA) complex ofAl34 is poorly defined, and EGTA should not be used as an

Al34 buffer.The ideas developed here apply to other systems as well.

The basic equation is K,, = aK�//3, with both a and f3 between0 and 1. If more than one protonation of the free ligandoccurs, or if more than one deprotonation of a complex takesplace, the expressions for a and /3, respectively, need to begeneralized from the mono-proton cases of NTA and citrateconsidered as prototypes in this section.

Exchange

Stability is not the only determining parameter in metal-ion reactions. An important but often overlooked feature isthe rate of ligand exchange in and out of the metal-ioncoordination sphere. Ligand exchange rates take on specialimportance for Al3”’ because they are so slow. The charac-teristic rate constant for substitution of inner sphere waterhas been determined for many metal ions (34). Increasingwater-exchange rates follow the order

Al34 <Fe34 <Be2”' <Ga3� 4 Mg24 � �2+, (]�3+ <Ca24

Each inequality sign indicates a 10-fold increase in ratefrom about 1 s’ for Al34 and increasing through eightpowers of ten to about 108 s” for Ca2� at 25 #{176}C.Althoughthese specific rate constants represent water exchange inaquo metal ions, they also reflect relative rates of exchangeof other uni-dentate ligands. Chelated ligands exchangemore slowly. The slow ligand exchange rate renders Al +

useless as a metal ion at the active sits of enzymes. The io�times faster exchange rate of Mg� provides sufficientreason for Al3” inhibition of enzymes with Mg24 cofactors.Any process involving rapid Ca24 exchange obviously wouldbe totally thwarted by Al34 substitution.

AI3� and Phosphate

In the human body, extracellular fluids contain about 2mmol of total phosphate per liter at pH 7.4 and intracellularfluids about 10 mmol of total phosphate per liter at pH 6.6.Al34 forms an insoluble salt with phosphate, often designat-ed as AIPO4, or sometimes as A1PO4 . 2H20, correspondingto the composition of the mineral variscite. At 0.16 mol/Lionic strength the ptC� values for successive deprotonationsfrom H3P04 -. H2P04” -. HPO -* P043” are 2.0, 6.77,and 11.6 (9). Thus P043” is the dominant phosphate speciesatpH >11.6 while equation 1 and Figure 1 shows that freeAl34 is the dominant Al3”' species only at pH <5.5. Thussignificant amounts of both Al3” and P04 are incompati-ble in solution at any pH. If we seek the overall neutralcomplex for which there is compatibility, we note from thePICa values that H2P04 dominates from pH 2 to 6.8 and,from Figure 1, Al(OH)24 is a principal species from pH 5.5 to6. For the purposes of solution chemistry it is advantageousto rewrite AlP04� 2H2O as Al(OH)2 . H2P04.

The solubility product of variscite has been reported in

terms of the hypothetical reaction

for which the equilibrium constant is

K9#{176}= (Al34)(OH’�’)2(H2P04) =

where the superscript0 designates a thermodynamic or zeroionic strength equilibrium constant, the parentheses signifyactivities, and the value of plCg#{176}= 30.5 refers to 25#{176}C(35).

We immediately recast the solubility equation in terms ofdissolution by hydrogen ion according to

Al(OH)2H2P04 + 2H4 �± Al3” + 2H20 + H2P04”

for which the equilibrium constant is

K�0#{176}= (Al34)(H2P04”)/(H”’)2 = K90/K�2 = 10-2.5

for the ion product constant of water pK� = 14.0. We need tofind the concentration equilibrium constant at 0.16 mol/Lionic strength and recast the last equation as

& 0 - r A 1341 ru flf� -‘i ,,tx+�2 -

�1o - L�” JYAl�LL21���P4 iY-” / - ztloYAlY_where the brackets signi� molar concentrations. The re-sults now become approximate because we can only esti-mate the activity coefficients at 0.16 mol/L ionic strength as

= 0.14 and, for H2P04”, y = 0.73 according to acommon treatment (36). Substitution of these values intoequation 9 yields pK� = -1.5. From pH 3 to 11 there areonly two predominant phosphate species, and the totalphosphate concentration is given by T� = [H2P04] +

[IWO4]. The fraction of the total phosphate that isH2P04” is given by [H2P04iPF� = (H4)/[(H4) + K�, whereplC2 = 6.77 at 0.16 ionic strength (9). Substituting theseconsiderations into equation 12 and solving for the free Al5”’molar concentration, we obtain

[Al3”’] = K10(H4)[(H”’) + K2]fr�= (H”')[(H”) + 106’77j/30 T� (13)

where the last equality applies to 0.16 molJL ionic strength.Equation 13 furnishes an estimate of the free Al34

concentration at 25#{176}Callowed by the solubility of varisciteat 0.16 mol/L ionic strength for a designated 3 <pH <11 andtotal phosphate concentration, T�. (Supersaturation and ionpair formation increase the Al34 solubility.) For intracellu-lar fluids at pH 6.6 containing 10 mmol of total phosphateper liter, equation 12 yields for -1og1Al34] = pAl = 12.5,while for extracellular fluids at pH 7.4 containing 2 mmol oftotal phosphate per liter, pAl = 12.9. This pair of valuesappears in the second row of Table 1 and represents ex-tremely low maximum free Al3”’ concentrations. Under thesame pair of conditions the free Ca24 concentrations permit-ted by insoluble hydroxyapatite, Ca5(P04)30H, a principalconstituent of bones and teeth, are nearly i0� times greater(9). Thus, thermodynamically, Al3”' easily displaces Ca24from phosphate binding.

Antacids containing Al(OH)3 deplete phosphate by precip-itation, tending toward a more negative Ca24 balance inindividuals with low Ca2” intakes (37). Dialysis osteodys-trophy develops with increased Al34 concentrations in thepresence of normal concentrations of Ca24 and Mg24 inplasma (38). Even low accumulations of Al3”' impair bonemineralization (39). The presence of an Al(OH)2H2P04precipitate may interfere with the orderly, kinetically con-trolled deposition of the bone mineral, hydroxyapatite (40).

Table 1 shows that in the presence of typical phosphateconcentrations the limitation on the free Al3”’ concentrationis not the solubiity of amorphous Al(OH)3; rather, it is themore limited solubility of Al(OH)2H2P04. Obviously, experi-

C0

0

1�,

0

ML H:,

,I/\

MLH�\LML�’

I 2 3 4 5 6 7 8 9

0�4

.0

0

0

LA,

a

� 02

2 3 4 5

pH

Fig. 4. Neutral citrate AIL#{176}mole fraction vs pHFrom left it, right, curves represent 100 to 0.01 mmol of cItrate per liter, M curvesare for I ,umol total AP� per liter, but the general features of the curves dependlittle upon the metai-ion concentration. The dashed curie corresponds to thedashed curve in Fig. 3

CLINICAL CHEMISTRY, Vol. 32, No. 10, 1986 1803

ments with Al34 in a phosphate buffer should be avoided.When attempted (41), such experiments lead to unreliableconclusions. The difference between the free Al34 concentra-tions allowed by Al(OH)3 and Al(OH)2H2P04 widens as theacidity increases, as illustrated by the results in Table 1 atpH 4.0, where the factor becomes iO�. Gastrointestinalabsorption of Al34 occurs from Al(OH)3 but not fromAl(OH)2H2P04 antacids (42).

There seem to be two possible explanations for the ex-tremely low concentrations of Al34 in living organisms.Either the Al3”' locked in the earth’s crust has been inacces-sible to life, or biological systems have evolved to rejectAl34. If inaccessibility is the answer, release of Al34 by acidrain appears more dangerous than if some rejection occurs.Current evidence suggests that resistance to Al3”’ uptake byliving systems results accidentally from the ubiquity ofphosphate and insolubility of Al(OH)2H2P04, and that whenthis system becomes ineffective, little defense remainsagainst Al34 uptake in acidic solutions. Plants that accumu-late Al34 do so in acidic soils (43), and evidently detoxify theAl34 by possessing basic fluids or organic chelators. Weexamine next a molecule that solubiizes Al34 so effectivelythat it renders phosphate ineffective in resisting Al34 up-take.

AI3� and Citrate

Because about 0.1 mmol of citrate is present per liter ofblood plasma, it becomes the pre-eminent small-moleculeplasma binder of a metal ion such as Al34 that prefersoxygen donor ligands. In the following discussion we use atypical concentration of citrate in plasma, 0.1 mmol/L. Arecent assessment of the citrate-Al3� system recommendsthe equilibrium constants that are appropriate for 25 to37#{176}Cand 0.10 to 0.16 moIJL ionic strengths with all plC5values in the commonly used scale of activity in hydrogenion (20). These recommended constants are used in thispaper. Slow polymerization reactions of Al3”’ complexesrequiring hours to complete (44) are excluded, because theyare unlikely to occur in a biological milieu. The threesuccessive acidity constants (pICa values) for citric acidunder the above conditions are 3.0, 4.4, and 5.8. Thus citricacid occurs as the tricarboxylate anion citrate at pH 7.4 inthe plasma.

In acidic solutions Al34 reacts with citrate di-anion, M34+ L112” � MLH”', with log K0 = 4.7, and this complexundergoes deprotonation, MLII” �± ML + H4, with pKob =

2.5. Still, in acidic solutions Al3”' reacts with the trianion togive an important complex of zero net charge, M34 + L3 �� with K8 = [ML#{176}]/([M34][L3]) = 1080. This neutralcomplex undergoes deprotonation, ML#{176}� MLH,,�’ + H4 inquite acidic solutions, with pKb = 3.4. The MLIL� complexcontains three anionic carboxylate groups and a deprotonat-ed citi ate hydroxy group. Finally, the neutral complex mayadd a second citrate, ML#{176}+ L3” � ML�’, with log K2 = 5.0(20).

Figure 3 shows the species distribution of the severalcomplexes in a solution 1 �mol/L in total Al34 and 0.1mmol/L in citrate. Only a little AlLH4 forms near pH 3where the net neutral A1L#{176}rises to 0.4 mole fraction. As thepH increases, AlL#{176}is succeeded by the strong citrate hy-droxy deprothnated complex AlLH1’. Even at this 100-foldcitrate to Al34 mole ratio there is little of the 2:1 complexAlL23”. At pH values only slightly exceeding the pH 7.4 ofplasma, Figure 3 shows that the water-derived Al(OH)4species abruptly intrudes and dominates at pH >8.

pH

Fig. 3. Species disttibution for 1 �tmol of AP” and 0.1 mmol of citrateper liter, plotted as mole fraction (aluminum basis) vs pHGeneral features of the curve are largely independent of the metal-ion concentra-

tion, The dashed curve corresponds to the dashed curve in Fig. 4

The net neutral ML#{176}citrate complex is of special interest,because it provides a means by which Al34 may passthrough membranes. Figure 4 shows how the mole fractionof this single species depends upon the citrate concentration.About half of all Al34 in solution appears as ML#{176}near pH 3in the presence of 0.1 to 100 mmol of citrate per liter.Significant quantities of net neutral ML#{176}occur even withthe lowest citrate concentration (10 �mo1fL) in Figure 4.

The significant mole fraction of net neutral ML#{176}from pH2 to 5 for a range of citrate concentrations such as thatshown in Figure 4 suggests that citrate complexation ofAl34 provides an effective means for Al34 absorption intothe body in the upper region of the gastrointestinal tract.This conclusion, derived from stability constants, is support-ed by experiments with rats. Increased Al34 concentrationswere found in both the brain and bones of rats fed a dietcontaining aluminum citrate, or even just citrate (45, 46).The citrate alone evidently chelates trace Al34 in the diet.Moreover, the Al34 concentration in the blood of humanswho are taking an Al(OH)3-based antacid is enhancedsubstantially by concomitant intake of citrate (47). There-fore, not only does citrate solubilization defeat both hydrox-ide and phosphate precipitation and elimination of Al34, butboth equilibrium arguments and animal experiments revealthat absorption occurs as well. This process shows thatpeople should not take aluminum-containing antacids with

pF

1804 CLINICAL CHEMISTRY, Vol. 32, No. 10, 1986

citrus fruit or juices. Although healthy individuals excludeAl34 from their systems, solubiization of Al34 by citrateprovides a means by which even the healthy individual may

absorb Al34.Table 1 summarizes the permitted free Al34 concentra-

tions in the presence of Al3�-complexing agents. The resultsare expressed as pAl = -log [Al34], so that the largernumbers represent the lowest free-Al34 concentrations. Atall pH values considered, Al34 is removed from solutionmore effectively by phosphate than as amorphous Al(OH)3(or even gibbsite), and most effectively of all by citrate.Therefore, of all the non-protein components of blood plas-ma, Al34 is most apt to be complexed with citrate. Furtheranalysis shows that at lower pH values such as might occurin the stomach, citrate complexes are again favored, withsome formation of Al3�-oxalate complexes near pH 2-3.Salicylate complexes (48) are insignificant. The neuraminic(sialic) acid (9) in gastric juice might bind some Al3”', but nostability constants have been determined. In the contestbetween Al34 precipitation by phosphate and elimination,and Al34 solubiization by citrate and possible absorption,equilibrium arguments indicate that solubilization wins.Studies of Al34 ingestion that do not measure or control theamount of citrate have overlooked a significant variablethat may affect the conclusions drawn.

Citrate complexes of Al34 inhibit precipitation of calciumphosphate at pH 7.4 (49). Under the experimental condi-tions the main complex should be AILW1 (Figure 3). On thetime scale of the precipitation experiments the complex may

polymerize to give Al3(OH)4L34 (44), a process that re-quires base. Because the total amount of Al34 present issmall, this base-consuming process should not interfere withthe interpretation of the base-consuming precipitation ex-periments.

Al3� and Fluoride

Al34 forms relatively strong complexes with fluoride ion,F-. Representative successive stability-constant loprithmsfor the addition of 1 through 5 fluoride ions to Al + at 25-37 #{176}Cand 0.16 ionic strength are 6.4, 5.2, 3.8, 3.3, and 1.3(50). Unlike almost all other ligands, the F complexes ofAl34 are stronger than those of Fe34. Figure 5 shows thedistribution curves of aluminum-fluoride complex speciesas a function of pF = -log[F”’], where [F”’] is the molarconcentration of free uncomplexed fluoride ion. From Figure5 it can be seen the neutral AIF3 complex predominatesnear 0.3 rnmol of ambient F”’ per liter, but exists in asignificant mole fraction from 0.02 to 5 mmol of ambientfluoride per liter.

Ingesting relatively small amounts of Al(OH)3 decreasesthe absorption ofF” (51) from the intestine. Along with theconcomitant decrease in phosphate and Ca2� (52), Al34 thusexerts a threefold adverse effect on bone structures. HighAl34 and low Ca24 and F in home water supplies ofdialysis patients leads to encephalopathy and bone fractures(53).

A surprising linkage ofF” with Al34 occurs in the story ofF’ activation of the adenylate cyclase (EC 4.6.1.1) enzymesystem. Almoitt since its discovery it has been known thatthis enzyme can be activated by F as a nonphysiologicalactivator. Later experiments revealed that the enzyme hastwo components, and that the F”’ acts on the guaninenucleotide-binding regulatory component, which requiresboth Me4 and a nucleotide such as ATP for activation.Most recently a contaminant in many commercial prepara-

tions of APP has been blamed for the apparent nucleotiderequirement. As with hexokinase, the contaminant is Al3”(54). Of several metal ions tested, only Be24 providedsimilar activation of the regulatory component. Both Al34and Be24 form a series of F” complexes. Although theanionic complexes AIF4 and BeF3 should have been themost prevalent under the experimental conditions (5 mmolof added F per liter), there were also significant amounts ofthe neutral complexes AIF3 and BeF2. Because the regula-tory component is prepared as a neutral detergent extractfrom membranes, I suggest that it is passage of the neutralcomplexes through a hydrophobic environment that is re-sponsible for the activation. In support of this proposal,investigators in another study find the most striking effectsof A13� at the lowest F concentration, 0.2 mmol/L in theircase (55), which is at the optimum for AIF3#{176}formation inFigure 5.

However, any interpretation based only on equilibriumproperties presents two difficulties. First, the experiments

(54) were performed in the presence of 1 mmol of EDTA perliter, which strongly chelates Al34, leaving iO� times asmuch free Al34 as required to combine with F. Second, ofthe several metal ions tested, Sc�”’ was ineffective in activat-ing the regulatory component (54), yet it forms a series of Fcomplexes of the same strength as Al3� and a strong EDTAcomplex (11). These two difficulties may be explained by theintroduction of rate effects. Al34 is known to react slowlywith EIYFA. In the rate of water substitution from thehydration sphere, Sc3”’ reacts up to i0� times more rapidlythan Be24, which in turn reacts 102 times faster than Al34(34). Rates of water loss carry over to reactions with otherligands, and so I account for the experimental results by theslow reaction of Al34 with EDTA permitting formation of F”’complexes, of which the neutral AIF3 is perhaps the relevantform. Sc3”’ undergoes relatively rapid sequestration byEDTA to form a strong complex, and F- is not thermody-namically competitive. Be24 reacts at an intermediate rate,and in this case the 5 mmol of F” may compete for the metalion with the 1 mmol of EDTA per liter. Compared with Al34,Be24 forms a disproportionately weaker complex withEDTA than with F. This rate-dependent hypothesis pre-dicts that prior incubation of Al34 with EDTA at the pH 8.0used in the experiment will not yield activation. Thoughunpublished, this experiment had been performed, and noactivation occurred (P. C. Sternweis, personal communica-tion, August 1985), in agreement with the prediction of therate-dependent hypothesis.

C0

C’,a

U-

0

Fig. 5. Mole fraction of total aluminum found as fluoride complexes as afunction of pF = -Iog[F], where [F-] Is the ambient F- molar

AI3� and Proteins

CLINICAL CHEMISTRY, Vol. 32, No. 10, 1986 1805

Are any proteins likely to be more effective Al3”' bindersthan is citrate? The common albumin and globulin proteins

of the plasma bind metals such as Al34 only weakly andnonspecifically. Albumin binds several Ca2”', with a stabil-ity log K6 -“ 2, and Gd�’ with log K8 = 3.9(56). Competitionwith Chelex, a cation-exchange resin, was used in an

attempt to determine Al34 binding at pH 7.4 to two proteinswith reported dissociation constants of 2.0 �tmo1/L for albu-min and a similar 0.52 �moI/L for transferrin (57). SinceAl(OH)4”’ is the predominant species, these conditionaldissociation constants require a large y correction accordingto equation 4. However, the similar values for the highlydissimilar proteins in the Al34 binding capacities suggest aflaw in the method. The 43 pg of total free Al34 per/literfound with Chelex at pH 7.4 corresponds to a solu�ilitybetween those for amorphous Al(OH)3 and gibbsite (F�igure

2). Chelex binds Al34 much too weakly to provide anymetal-ion buffering for competition with transferri4. Thesuggestion that Al34 binds strongly at the amino tei�minal

site of human serum albumin (58) is unconvincin�. Fournitrogen donors-consisting of the human albumin amino

terminus, two deprothnated peptide nitrogens, and a histi-dyl ring nitrogen-tightly chelate Cu24 and also Ni24, with

spin pairing of the latter metal ion (59). This stronglycovalent, quadri-dentate, tetragonal chelate ring systemcannot be a strong binding site for Al34, which dependsupon electrostatic interactions for binding.

Albumin is much too weak a metal ion binder to with-draw Al34 from any of the complexes of Table 1, all of whichoccur in the plasma. At the pH 7.4 of plasma, albumincannot compete for Al34 with hydroxide or phosphate pre-cipitation and citrate complexation. If Al34 is to be proteinbound in plasma, it must be linked to a much stronger Al3�binder than albumin.

With a pair of sites that avidly bind Fe34’, transferrin

stands as the leading plasma protein for Al3� binding. At anormal concentration in plasma of 3 g/L, with two metal-ionbinding sites per 77 000-Da protein at only 30% site occu-pancy by Fe34 in the plasma, transferrin furnishes unoccu-pied metal-ion binding sites at a concentration of 50 �mol/L.Because this concentration is half that of citrate in theplasma, to be competitive transferrin needs to bind Al34twice as strongly as does citrate at pH 7.4.

Direct competition between Al3� and Fe3” for bindingsites on transferrin in combination with an assumed bind-ing constant for Fe3t�transferrin were used in an attemptto establish the binding constant for Al34-transferrin (60).The key experiments were conducted with a large excess ofAl3”, and solutions were examined after only 30 mm.However, it takes hours for citrate-bound Fe34 to displaceAl34 from Al34-transferrin, and therefore the calculated

Al34 binding constant is much too high. The results may befitted by assuming that citrate-bound Fe34 reacts withtransferrin 100 times more rapidly than does citrate-boundAl34 (see Exchange section above) and insignificantly in 30mm with Al3”'-transferrin.

A recent quantitative spectroscopic determination of Al34binding to the two sites of transferrin yields successivestability constants of log K1 = 12.9 and log K2 = 12.3 underblood-plasma conditions of pH 7.4 and 27 mmol of HCO3”’per liter (33). For 1 �mol of total Al3� per liter of plasma, thefree Al34’ concentration permitted by transferrin is 10_146

mol/L. As indicated in Table 1, this amount is less than thatallowed by insoluble Al(OH)3, Al(OH)2H2P04, or by corn-

plexation with citrate. Thus transferrin is the ultimatecar#{231}ierof Al�’ in the plasma.

The same study found, for the successive stability con-stants of Fe34 binding to transferrin, log K1 = 22.7 and logK2 = 22.1 (33). These values agree closely with a revision,obtained by equilibrium dialysis, of those in the literature.By comparison, the Al34 stability constants are weakerthan expected, and it is suggested that the significantlysmaller Al34 cannot coordinate to all the transferrin donoratom� available to Fe3t Qualitatively, this study showsthat with a citrate/transferrin molar ratio comparable tothat found in the plasma, citrate releases both Al3”’ andFe34 to transferrin. Because of the nearly 10 log unit

difference in their binding strengths, Al34 cannot displaceFe3”’ from transferrin. However, with most plasma trans-ferrin carrying unoccupied binding sites, there are sufficientresources for binding both of these metal ions.

I am grateful to Michael R. Wills, John Savory, and Roger L.Bertholf for their encouragement in my writing of this article.

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