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CLIN. CHEM.
32/10,
1797-1806
(1986)
The Chemistry of Aluminum as Related to Biology and Medicine
R. Bruce
MartIn
The increasing
number of roles discovered
for Al34 in
physiological
processes demands an understanding
of how
Al3’ interacts with compounds
in biological systems.
Al3” is
expected
to complex with oxygen donor ligands, especially
phosphates,
and it does so in soils, in the gastrointestinal
tract, and in cells. The stability of Al3’ complexes
has
generally been misjudged because of lack of recognition that
free, aqueous Al3” Is not the dominant
form in neutral
solutions 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 the
presence of inorganic phosphate,
the permitted free Al34 is
decreased further, through formation of insoluble aluminum
phosphate. This precipitate facilitates the elimination of Al3’
from the body. In contrast, citrate solubilizes Al34, and an
appreciable
fraction occurs as a neutral complex that may
pass through membranes
and provide a vehicle for Al34
absorption into the body. In the blood plasma the most likely
small-molecule
complex is that with citrate, while the only
competitive
protein complex is that with transferrin, a protein
built to transport Fe3’ but whose sites are only 30% occupied.
AdditIonal
Keyphrases:
complexes
kidney
disease
of Al with various anions,
Al-induced
.
compounds
disorders
Ionic spe-
cies of Al
metal-ion buffer systems
pAl
stability constants
ligand protonation
complex deprotonation
p/ciepliate
citrate
transferrin
ligand exchange rates
Intestinal absorption of dietary/medicinal
Al
fluoride
proteins
hemodialysis
effect on bone
.
.
.
.
.
Aluminum
has recently
been recognized
as a causative
in dialysis
encephalopathy,
osteodystrophy,
and rnicrocytic
anemia
occurring
in patients
with chronic
renal
failure
who undergo
long-term
hemodialysis
(1). Only a
small amount
of Al34 in dialysis
solutions
may give rise to
these disorders.
Encephalopathy
has also occurred
in children consuming
Al(OH)3
as a phosphate
binder
for renal
disorders
(2, 3). (Ca2”
would
have
been a more natural
choice.)
A13’ has also been
implicated
in neurotoxicity
associated
with
amyotrophic
lateral
sclerosis,
a form of
parkinsonism
with severe dementia,
in the indigenous
population
of Guam,
where the soils are high in Al34 and low in
Mg2’ and Ca24 (4), and in Alzheimer’s
disease
(5).
These developments
have led some investigators
to perform experiments
with Al34 without
appreciating
its solu-
agent
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 neutral
solution
and expect the aqueous
Al34 concentration
to be 1
mmol/L?
Is it appropriate
to use phosphate
buffers
with
Al34? In addition
to equilibrium
considerations,
are rate
effects
important?
Does the order
of mixing
of Al34 with
EDTA and F make a difference?
In the plasma
what is the
dominant
Al31 binder
among
small molecules
and among
proteins?
This
article
aims
to develop
the bioinorganic
chemistry
of Al34 so that answers
to such questions
become
evident
to readers.
Healthy
people
consume
Al34
in a variety
of ways.
Fortunately,
most foods do not dissolve
significant
amounts
of aluminum
from cookware.
However,
the attack
of food
juices on aluniinwn
vessels
varies
greatly,
depending
upon
pH, temperature,
and other substances
present.
Hot, acidic
fruit juices-in
the absence
of sugar-.corrode
aluminum
ware,
as do salt solutions
used in pickling
processes
(6).
These solutions
and conditions
should
probably
be avoided
by the prudent
homemaker.
Al34 is a component
of alumcontaining
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 related
A12(S04)3
have been added to drinking
water to improve
its
appearance.
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,
marashino
cherries,
and pickles
to improve
their appearance. Al34 salts are often added tops
cheeses
and to
beer.
Al2(S04)3
is the most
widely
used
coagulant
for
clarifring
turbid
drinking
water.
The success
of this treatment depends
mainly
upon precipitation
of Al(OH)3,
with
adsorption
of the turbidity.
Excess Al34 remains
in solution.
Thus a “dirty”
water
that contained
some sediment
(probably harmless)
and no or little
Al34 becomes
clear
and
contains
no or little sediment
and some Al34. It has been
suggested
that ingestion
of lead from plumbing
contributed
to the demise
of the Roman
empire
(7). Could it have been
due to A13 instead?
AI3
Complexes
The only
accessible
oxidation
state
for aluminum
in
biological
systems
is 3+.
Binding
of Al34 is primarily
electrostatic
(as opposed
to covalent)
and therefore,
in addition to charge,
ionic size is an important
parameter.
Effective ionic radii in angstrom
units (1 A = 10’
rim) in sixfold
coordination
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
CLINICAL CHEMISTRY, Vol. 32, No. 10, 1986
1797
fourfold
coordination
and Ca2
(1.12) larger
in its favored
eightfold
coordination
(8). Excluding
the unnatural
Ga34,
the radius
of Al3”
most
resembles
that
of Fe34.
Thus
appearance
of AL3” in Fe34 sites seems likely.
Except
with
F, Al3 forms weaker
complexes
than Fe3”’. The binding
of
Al3 and Fe3” to transferrin
is discussed
later.
Though
Mg
is somewhat
larger
than
A13”, displacement of the ubiquitous
Mg24 in biological
systems
by Al34
appears
likely.
Mg2
is often
associated
with phosphate
groups,
and comparison
suggests
that Al3” should also seek
those sites. In many physiological
systems
the Mg24 concentration
is about
1 to 2 mmoliL.
Al34 binds almost
i0 times
more strongly
to ATP
than does Mg24 (see below). Thus in
these
systems,
nanomolar
amounts
of Al34 can compete
with Mg2
for the phosphate
sites.
Ca2 not only is much larger than Al34, it also frequently
occupies
specific sites in proteins
(9). In its favored
eightfold
coordination
the volume
of Ca2
is nine times greater
than
Al3’ in sixfold coordination.
Since hole production
is energetically
costly, Al34 cannot
replace
Ca2’ in proteins
without substantial
readjustment
of liganding
groups.
Thus
competition
between
Al3” and Ca2” is less apt to be for
protein
binding
sites than
for small molecule
ligands
and
phosphate,
with which
both form insoluble
complexes.
In biological
systems
we anticipate
that Al3” will associate with oxygen
donor
ligands,
at least
some of which
should
be anionic
to help counter
the 3+ charge
on the
cation.
Carboxylate
and phosphate
groups,
inorganic
phosphate,
nucleotides,
and polynucleotides
meet this prescription.
Unless
carboxylate
groups
are arranged
to make
strong
chelation
possible,
as in citrate
(see below),
Al3”'
should prefer phosphate
binding.
Because
it carries
but one
carboxylate
group,
lactate
forms only weak complexes
with
Al3’. The stability
constants
remain
unknown,
witnessing
to the weakness
of the interaction.
Unless
other ligands
are
present,
dissolution
of Al3” as Al(lactate)3
should
yield
hydroxo
complexes
as the complex
dissociates
and weakly
bound
lactate
diffi.ises away.
A prominent
peak in an NMR
spectrum
of a pH 6 solution
containing
10 mmol of Al3”' and
30 mmol of lactate
per liter was assigned
to the 3:1 complex
(10). Under
these
experimental
conditions,
less than half
and possibly
much
less of the metal
ion is expected
to be
present
in a 3:1 complex.
Al3” forms only weak complexes
with amines
and sulfhydryl ligands.
There
is almost
no tendency
for Al34 to
complex
to sulfhydryl
groups.
Metal
ion hydrolysis
interferes
well before
A13” might
coordinate
to uni-dentate
amities.
With multi-dentate
amino-carboxylate
ligands
such
as EJYFA, Al34 forms strong
complexes
(11).
AI3
In Aqueous
Al(H20)63”
by common
=
consent
#
Al(H20)5(OH)24
we write
Al3’
K1
H20
=
it as the
H30”,
+
hydrolysis
+ H20
Al0H24
(ff’)[Al0H2”i/[Al341
+
reaction
H4
(1)
=
where the parentheses
signify
activity,
the brackets
concentration
and Kia is the associated
equilibrium
stant.
This reaction
is followed
by deprotonation
second Al3’-bound
water
molecule.
Al0H24
+
H20
Al(OH)2”'
H’
+
(H”)[Al(OH)2”j/[Al0H24]
=
signify
confrom a
(2)
10_56
=
Little
evidence
favors
significant
amounts
of soluble
Al(OH)3
in solution,
but deprotonation
from two more
bound waters
yield a soluble
tetrahydroxo
species.
Al(OH)24
+ 2H20
Al(OH)[
(H”)2[Al(OH)4]/[Al(OH)241
K
+
=
2ff”
10_121
(3)
The values
for the equilibrium
constant
at 25 #{176}C
assigned
to
each of the three
reactions
(and reaction
5 below)
are
converted
from a set of thermodynamic
values
(13) by a
developed
protocol
(14) to 0.16 ionic strength.
Figure
1 shows
the distribution
of soluble,
mononuclear
aluminum
ion species
in aqueous
solutions.
The octahedral
hexahydrate
Al(H20)634
dominates
at pH <5, and the
tetrahedral
Al(OH)4
at pH >6.2, while there is a mixture
of species
from 5 <pH
<6.2.
To find the relation
between
the free, non-hydrolyzed
aqueous
Al34 and the total soluble
aluminum
ion concentration, we define the mole fraction
of nonhydrolyzed
aqueous
ions as -y = [Al3” I/CM where
CM
=
[Al34]
Combination
+
with
lIy
=
[A1OH2”]
+
[Al(OH)2”]
1-3
equations
1 + [10551(H4)]
[102/(H”)4]
+
[Al(OH)4]
yields
+ [10”’/(H’)2]
+
(4)
The mole fraction
y is pH dependent
and must be calculated
for each pH. Substitution
into equation
4 yields,
at pH 6.5,
SolutIons
Whatever
other
ligands
may be present,
understanding
the state of Al34 in any aqueous
system
demands
awareness
of the species
that Al34 forms at different
pH values
with
the 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 yield
Al(OH)2”
and Al(OH)2”'.
Neutral
solutions
give an Al(OH)3
precipitate
that
redissolves
in basic
solutions,
owing
to
formation
of tetrahedral
Al(OH)[.
Polynuclear
species may
also form, their
compositions
being
time dependent
(12).
(Since this article
deals with relatively
low concentrations
of
total
Al34 in biological
systems
in the presence
of other
ligands,
polynuclear
species
are not considered.)
1198
Equilibria
among
mononuclear
Al3” species
in aqueous
solutions
may be described
by reactions
1-3 below.
For
convenience
we abbreviate
the hexahydrate
occurring
in
acid solutions
as Al3”'. This
abbreviation
disguises
the
actual
reaction,
deprotonation
of Al3 “-bound
water.
Although
the first deprotonation
reaction
is more accurately
written
as
CLINICAL CHEMISTRY, Vol. 32, No. 10, 1986
6
pH
Fig. 1. Distribution
aqueous solutions
of soluble,
mononuclear
na:
of aluminum
ion occumng as each designated
mole
fraction
any pH the indMdual mole fractions sum to unity
aluminum
ion species
in
species. At
1/y = 720 and, at pH 7.4, 1/y = 2.5 x 106. Because
at pH 7.4
virtually
all soluble
aluminum
ion occurs as Al(OH)4,
the
value 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 distribution
among
soluble,
mononuclear
aluminum
ion species
without
considering
the absolute
amounts
permitted
by the
limited
solubility
of Al(OH)3.
At reasonable
temperatures
the stable crystalline
phase of Al(OH)3
is the mineral
called
gibbsite.
The solubility
of Al(OH)3
from solid gibbsite
may
be described
as
Al(OH)3(gibbsite)
3H”'
[Al3”'J/(H”')3
+
K1’
=
Al3”’
=
1092
±
+
3H20
(5
The gibbsite
solution
reaction
could have been written
as
giving
Al34 + 30H”,
and the solubility
product
constant
expressed
correspondingly.
By using only H” and not 0H
in reactions
we avoid the need to calculate
the OH- concentration
or activity,
and for the H”' activity
we use
However,
equilibrium
is slowly achieved
with gibbsite,
and
the solubility
with respect
to amorphous
Al(OH)3 may be up
to 100-fold
greater.
The presence
of organic
ligands
such as
citrate
favors
formation
of non-crystalline
Al(OH)3
(15).
Because
there rarely
is equilibrium
with respect
to gibbsite
in biological
systems,
we employ
a more liberal
equilibrium
constant
for solubility
of amorphous
Al(OH)3:
K1
=
[Al34]/(H4)3
=
i0’#{176}’
(6)
Figure
2 shows the total molarity
of all soluble
aluminum
species
permitted
by both gibbsite
and a representative
amorphous
Al(OH)3
according
to equations
5 and 6. The
minimum
solubility
in both curves
occurs at pH 6.2.
From equation
6 we may estimate
the highest
permitted
hexahydrate
Al3” concentration
from amorphous
Al(OH)3
as [Al(H2O)634]
=
10107
x 103pH
At pH 7.4, the pH of
extracellular
fluids
such
as blood
plasma,
we obtain
[Al(H20)63’J
=
10_h15
molJL. This important
result
means
that, at the pH of blood plasma,
the highest
obtainable
free
Al3” concentration
allowed
by an amorphous
Al(OH)3
is 3 x
10’.’12 mol/L.
The negative
logarithm
of this value, designated as pAl (analogous
to pH), appears
in the first row of Table
1. The corresponding
total aluminum
concentrations
is
given by CM = [A13”']/y. At pH 7.4 we have already
found
that 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
.4
-J
0
z
D
-j
4
-J
.4
I0
I(0
0
pH
Fig. 2. Negative
logarithm
of total molar
concentration
of aluminum
allowed by Al(OH)3 solubility vs pH
Lower curve represents true equilibilum solubility from gibbets. Upper curve
depicts 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 the
minimum solubility occurs for both curves. From 5 <pH <6.2 there is a mixture of
soluble species, as shown in Fig. 1
Table
1. MaxImum Free Al3’ Molar ConcentratIons
Expressed
as pAl =
Iog[A134]
-
Complexon
Amorphous
Al(OH)3
Al(OH)2H2P04
Citrate, 0.1 mmol/Ld
Transferrin
d,
pH 4.0
pH 6.6
1.3
9.1
6.5a
8.3
pH 7.4
11.5
12.9c
14.0
14.6
125b
13.1
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.
This result deserves
emphasis.
For even though
the permitted total Al3
at pH 7.4 may reach 8 pinolJL,
most appears
as Al(OH)4”
and only 3 x 10’12
molJL as Al(H20)634.
If the
gibbsite
solubility
product
constant
were used, both allowed
concentrations
would be 1/30 as great.
Like any other ligand,
hydroxide
ion, by reactions
1-3 and
6, 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
containing Al34. Unless
a solution
is supersaturated
with respect
to
amorphous
Al(OH)3,
greater
than
nanomolar
concentrations of free Al3” in neutral
solutions
are unobtainable.
Upon
addition
of 1 mmol
of an Al3”’ salt per liter to a
solution
at pH 7.4 the free Al3”' concentration
is not 1
mmol/L
but only about
a miniscule
3 x 10
mol/L. The
predominant
water-derived
complex
is Al(OH)4
at 8
MmoIJL. Unless
the remainder
of the added Al34 has been
complexed
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” or
Al(H20)63”
concentration
rather
than
the much
greater
Al(OH)4”
concentration
that is the significant
quantity
in
neutral
solutions.
Because
they fail to incorporate
the basic ideas described
in this section,
many papers
in the literature
reach dubious
conclusions.
Dissociation
constants
for Al34 binding
that are
near to or greater
than the free Al3”' concentration
allowed
by the solubility
of Al(OH)3
are suspect.
In a study of Al3”
binding
to the important
calcium
regulatory
protein
calmodulin,
the authors
performed
equilibrium
dialysis
experiments
at pH 6.5 and calculated
the binding
constants
from
the presumed
total Al3”’ in solution
(16, 17). From equation
4 the free [Al3”'] is only 1/720 that of total [Al3”], and so their
binding
constants
must
be increased
by a factor
of 720.
However,
this investigation
was performed
near the minimum of Al(OH)3
solubility
(Figure
2), where
for the amorphous form equation
6 allows
1.6 nmol of free hexahydrate
Al3” 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. This
comparison
renders
suspect
the conclusion
that calmodulin
binds three
Al3”’ ions so strongly.
The result
ncieds to be
verified
at a pH removed
from the minimum
in Al(OH)3
solubility
and with an appropriate
metal-ion
buffer system
to control
the free Al3”’ concentration
(see below).
On a
structural
basis it seems most unlikely
that calcium-calniodulin
can bind
three
Al34
ions at anywhere
near
the
strengths
proposed
in these papers.
Suspect
also are inhibition
constants
in the miuimolar
range
for acetylcholinesterase
(EC 3.1.1.7)
activity
derived
from addition
of millimolar
amounts
of total Al3” at pH 7.5
CLINICAL
CHEMISTRY,
Vol. 32, No. 10,
1986
1799
(18),
total
where
Figure
Al3” per liter,
CondItIonal
Ligand
2 shows
a solubility
even from amorphous
StabIlity
of
Constants
Al3”'
L3
+
log K,
appears
=
=
K,
AlL#{176}
11.4 (19).
as a species
H”'
where PKa
For NTA
reaction
Al34
+
=
Ka
=
ligands
with amino
strengths
in neutral
in neutral
solutions,
to the deprotonated
taken
as a specific
to citrate.
constant,
K,, refers
[A1L#{176}].’[Al3”’][L3’1
In neutral
with a 2-
L3’
+
9.58.
in neutral
solutions
most of the
net charge.
We write
(H4)
solutions
we have
HL2
H4
+
[L31/[HL2’i
the
displacement
AlL#{176}
The concentration
of deprotonated
ligand
L3” available
to
the 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
pHdependent
stability
constant
is given by K’ = aIC,, or
log K’
=
log K,
-
P1a
-
log [(H4)+KaI
In the limit where
pH
piCa, unbound
ligand
is predominantly
in its basic form and log K’ = log K,. For pH 4 pK5
unbound
ligand
is predominantly
protonated
and log K’,, =
log K, -pK
+ pH. When
the pH is within
two log units of
pK5, the complete
equation
7 should be used. For NTA at pH
7.4 we have logK’
=
11.4
9.6 + 7.4 = 9.2. The conditional
stability
constant
provides
a measure
of complex
stability
under
specific
pH conditions
where
there
is a protonated
ligand.
By allowing
for withdrawal
of deprotonated
ligand
from solution
by protonation,
the value
of a conditional
stability
constant
becomes
less than that for the standard
stability
constant.
-
Complex
Deprotonation
The formulation
in terms
of the conditional
stability
constant
described
so far may be incomplete
if the complex
itself undergoes
one or more deprotonations
at a pH near to
or less than the pH of interest.
Al34 complexes
of both NTA
and citrate
deprotonate
in acidic solutions.
Continuing
with
the 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 from
an Al34-bound
water molecule
in the complex.
[More appropriately,
the left-hand
side of the reaction
could be written
with
the single
reactant
(H20)AlL#{176}.] The deprotonation
reaction
increases
the effective
stability
constant.
We enlarge
the scope of the conditional
stability
constant
described
above
to include
complex
deprotonation.
The
fraction
of non-deprotonated
complex
is given
by 13 =
(H”')/[(H4)
+ Kb], so that
0 <13 <1. The conditional,
pHdependent
stability
constant
is now given by K,, = aK,/(3 or
1800
logK,,
=
logK,
CLINICAL CHEMISTRY, Vol. 32, No. 10, 1986
+
log a
pH
+
+
log [(H4)
HL2
+
±
HOAIL”
+
Kb]
+
where
log a is evaluated
as in equation
7.
For NTA at pH 7.4 the conditional
stability
K,,
L”1/([Al3”i[HL”])
refers to the overall
reaction
Al34
Protonation
Stability
constants
for multi-dentate
groups overstate
their effective
binding
solutions.
Amino groups
are protonated
and tabulated
stability
constants
refer
ligand
(11). Nitrilotriacetate
(NTA)
is
example;
the model is easily transferable
For nitrilotriacetate,
L3, the stability
to the reaction
where
ligand
of only 10 pmol
Al(OH)3.
H”
=
(8)
[HOAl-
(9)
We now have logK,,
=
11.4
9.6 + 7.4 + 7.4
5.2 = 11.4.
By coincidence,
the conditional
and tabulated
stability
constants
are numerically
equal
at pH 7.4 only, because
the
stability-promoting
effect of a deprotonated
complex
exactly
offsets the destabilizing
effect of an appreciable
fraction
of
monoprotonated
unbound
ligand.
For citric acid the Al34 complex
also loses a proton,
but
from the citrate
ligand
according
to AlC#{176} AlCH.,,1
+ H4,
where
pKb = 3.4 (20). The three
successive
pK, values
for
citric 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’
the equilibrium
constant
log K,, = 8.0 + 7.4
3.4
citrate
complex
at pH 7.4
iO times
greater
than
owing
to deprotonation
solutions.
-
:±
-
AlCH1”
+
H4
K,,
=
[AlCH1]/([Al34][C3”])
and
12.0. In the case of the Al34the effective
stability
constant
is
the tabulated
stability
constant,
of the complex
in quite
acidic
=
It is occasionally
useful
to expand
further
the concept
of
conditional
stability
constant
to allow for metal-ion
hydrolysis, because
only the nonhydrolyzed,
aqueous
metal
ion is
considered
to be able to form complexes.
We have already
defined
the mole fraction
of nonhydrolyzed,
aqueous
ion as
y, and for Al34 its value at any pH may be calculated
from
equation
4. The conditional
stability
constant
that allows for
all 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
to
ATP at pH 6.95(21).
We find a from PKa =6.5,13=
1,7 from
equation
4, and calculate
the standard
stability
constant
logarithm
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
of
nonhydrolyzed
aqueous
Al34 to ATP.
There is a substantial 4.6 log unit allowance
for Al34 hydrolysis.
In comparison, for ATP4’
and Mg2”' we have log K, = 4.2 (22), 6.7 log
units weaker
than for Al3”'.
In terms of the standard
stability
constant,
the binding
of
Al3” to ATP4’
is significantly
stronger
(log K, = 10.9) than
to citrate
(log K, = 8.0). Mainly
because
the citrate
stability
constant
fails to allow for complex
deprotonation,
the conditional stability
constant
at pH 7.4 for citrate
becomes
log IC,,
=
12.0 as shown
above.
The corresponding
conditional
constant
at pH 7.4 for APP4’
is nearly
the same
as the
standard
constant,
log K,, = 10.8 (note that this is K,,, not
K’,,). This comparison
indicates
that citrate
complex
deprotonation
results
in a reversal
of the standard
order
for
stability
constants,
and we predict
that citrate
will extract
Al3”’ from ATP4’
in neutral
and weakly
acidic
solutions.
Experimentally,
citrate
has been used to remove
Al3”' from
ATP (23).
Whether
the standard
or conditional
stability
constants
are more useful depends
upon the situation.
In tabulations,
pH-independent
standard
stability
constants
are usually
used. On the other hand, the practicing
chemist
may argue
that, because
the pH region of overlap
of Al(H20)634
(Figure
1) and APP4’
with pK5 = 6.5 for ATPH3
is narrow
to nonexistent,
a conditional
constant
is more practical.
However,
-
conditional
constants
are pH-dependent,
and so must
be
calculated
for each pH. It is simpler
to calculate
a conditional stability
constant
from a standard
stability
constant
than
from another
conditional
constant
calculated
for another
pH. In addition,
only standard
stability
constants
give an
appropriate
comparison
between
the binding
affinities
of
two 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 APP
occurs
mainly
at the phosphate
chain,
with insignificant
interaction
at the nucleic
bases.
Transition
metal
ions also
interact
at the nucleic
bases
(24, 25). Nuclear
magnetic
resonance
spectra
indicate
that
aqueous
Al3”' ions form
several
complexes
with the phosphate
group
of APP and
undergo
exchange
on a millisecond
time scale (26). Direct
binding
of Al34 to N7 on the adenine
ring was proposed
for
one complex
on the basis of an upfield shift of H8. However,
metal-ion
binding
at N7 always
produces
downfield
shifts at
H8 (27). An upfield
H8 shift is attributed
to nucleic
base
stacking
or deprotonation
at Ni.
Consistent
with the localization
of Al3”' in the chromatin
in the cell nucleus
of neurofibrillary
tangles
(4, 28), the
slowly exchanging
Al3”' forms several
complexes
with DNA
(29). Some of the complexes
may be multinuclear
with two
or more
Al34 joined
by hydroxo
bridges.
The detailed
structures
for the Al3”'-DNA
interactions
remain
to be
definitively
specified.
In pH regions
similar
to those
in which
a metal
ion
deprotonates
a water
hydroxy
group,
the metal
ion may
interact
with and deprotonate
an alcoholic
hydroxy
group.
The ribose sugar contains
a pair of cis 2’,3’-hydroxy
groups
that may advantageously
form a chelate
with an appropriate metal
ion (24). In nucleoside
phosphates,
Al34 prefers
the basic phosphate
site. In nucleic
acid polymers,
however,
the negatively
charged
phosphate
on each residue
is not
basic.
In RNA the cis-hydroxy
groups
of ribose
provide
a
possible
binding
site at pH >5 for Al34, which
has little
tendency
to coordinate
to nitrogens
of the nucleic
bases.
For years it has been known
that certain
metabolites
such
as phosphate
and citrate
activate
yeastand brain-hexokinase (EC 2.7.1.1)
enzymes
at pH s7. The metabolites
thus
become
implicated
in regulating
hexokinase
function.
Recently,
however,
it has been shown
that the loss of hexokinase activity
at pH 7 is ascribable
to Al34 contamination
of commercial
preparations
of APP (23,30).
Phosphate
and
citrate
“activate”
by complexing
Al34 and freeing
the APP.
The variable
Al34 contamination
is low, of the order of only
mole per hundred
moles, and is the most common
metal-ion
contaminant
of APP preparations
(23,30,31).
Inhibition
by
Al34 shows
up because
the inactive
APP-Al3”’
complex
binds
to hexokinase
about
i0 times
more strongly
than
does APP-Mg2”'.
This enormous
hexokinase
binding
advantage
in favor
of APP-Al34
is the really
extraordinary
conclusion
of these investigations.
If other nucbeoside
phosphate-Al3
complexes
bind in a similarly
strong
fashion
to
even a fraction
of the wide range of enzymes
with nucleoside
phosphate
substrates
or cofactors,
then
Al34 (and other
metal ions?) becomes
potentially
toxic at low concentrations
by inhibiting
many key metabolic
processes.
Polymerization
of tubulin
to microtubules
benefits
from
the presence
of guanosine-5’-triphosphate
and Mg2”’. Recent
research
establishes
that
less-than-nanomolar
concentrations
of Al34 promote
polymerization,
making
Al3”' iO
times more effective
than Mg”’
(32).
Metal-Ion
Buffers
In conducting
experiments
with metal
ions it is often
necessary
to fix their
concentrations
reliably
at known
values.
Metal-ion
buffers
are analogous
to pH buffers
except
that it is the free metal-ion
concentration
that is controlled
in the presence
of excess ligand.
As with protonic
equilibria,
the buffering
is most effective
when the ratio of complexed
ligand
to total ligand
(molar
concentrations)
lies between
0.15 and 0.85. Because
ligand
concentration
exceeds
that of
tightly
bound
metal
ion, the total
metal
ion/total
ligand
molar
ratio,
R, should
also lie between
0.15 to 0.85. Not
included
in the treatment
are 1:2 Al34 to NTA or citrate
complexes,
because
analysis
indicates
that they do not form
to an appreciable
extent
with ligand
in concentrations
up to
10-2 molIL.
The development
of the previous
section
in terms
of the
conditional
stability
constant,
K,,, permits
us to formulate
directly
the equation
for metal
ion buffering.
For the NTAAl34 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 is
tightly
bound
so that
the concentration
of free Al3”’ is
negligible.
Substituting
in the expression
for the conditional
stability
constant,
rearranging,
and noting
that R = CM/CL
we obtain
[Al34]
This result
is a
tion is given by
constant
applies
ed. For systems
given by
K,,
=
R/(i
-
R)K,,
(10)
general
one. The free metal-ion
concentraequation
10, where the conditional
stability
to the specific pH for which it was calculatformally
analogous
to NTA and citrate,
K,, is
=
K,[1+Kb/(H”’)I/[i+(H”')/KaI
(11)
The (H”)/Ka
term
accounts
for protonated
ligand
and the
K,/(H4)
term for deprotonation
of Al3 “‘-bound
water
in the
complex.
The important
metal-ion
buffer
equations
10 and
ii
express
the free metal-ion
concentration
as a function
of
known
equilibrium
constants,
pH, and only the ratio (R) of
total metal ion to excess total ligand concentrations
and not
their
absolute
concentrations.
To vary the free metal-ion
concentration
at a fixed pH, it is necessary
to vary R. At a
fixed pH the right-hand
side of equation
10 describes
a
sigmoidal
curve
symmetrical
about
R = 0.5, similar
to a
titration
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-ion
concentrations
are very low, but controlled.
For the same R
value the pAl for the two buffer systems
differs by only 0.6
log units, or a factor of 4, in [Al34]. If one system
applies,
the
other
should
apply
also, and experiments
with
the two
different
kinds
of ligands
should
serve as a check on one
another.
Both buffer
systems
were used in this way in a
recent
evaluation
of the stability
constant
for Al3”’ binding
to a serum
protein,
transferrin
(33).
It is reasonable
to vary R from 0.15 to 0.85 to obtain
a
range of 1.5 log units in pAl. If a wider range is desired,
it is
desirable
to vary pH or find another
ligand
for a new metal
ion buffer system.
The condition
where
R = 1 or CM = CL
should
be avoided,
because
the solution
is unbuffered
with
wide swings
in the value
of pAl for even small
deviations
from exact equality
(analogous
to the endpoint
in a titration
curve).
CLINICAL
CHEMISTRY,
Vol. 32, No. 10, 1986
1801
With allowance
for an additional
ammomum
group deprotonation
in the free ligand,
the final equation
for EDTA
complexation
resembles
that for NTA. Substitution
of the
EDTA
stability
constants
(11) yields
at pH 7.4 and the
midpoint
of the metal-ion
buffer range,
pAl = 15.1. The free
Al34 concentration
in the presence
of EDTA
is 5000 times
less than that in the presence
of NTA. However,
equilibrium
in the ED’FA
complex
of Al34 is reached
slowly,
and
experiments
with it need careful
monitoring.
The [ethylenebis(oxyethylenemtrilo)]tetraacetic
acid (EGTA)
complex
of
Al34 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 between
0 and 1. If more than
one protonation
of the free ligand
occurs, or if more than one deprotonation
of a complex
takes
place, the expressions
for a and /3, respectively,
need to be
generalized
from the mono-proton
cases of NTA and citrate
considered
as prototypes
in this section.
Exchange
Stability
is not the only determining
ion reactions.
An important
but often
the rate of ligand
exchange
in and
coordination
sphere.
Ligand
exchange
importance
for Al3”’ because
they are
teristic
rate constant
for substitution
has been determined
for many
metal
water-exchange
rates follow the order
Al34
<Fe34
<Be2”'
<Ga3
4
parameter
in metaloverlooked
feature
is
out of the metal-ion
rates take on special
so slow. The characof inner sphere
water
ions (34). Increasing
Mg24
2+,
(]3+
<Ca24
Each
inequality
sign indicates
a 10-fold
increase
in rate
from about
1 s’
for Al34 and increasing
through
eight
powers
of ten to about
108 s”
for Ca2
at 25 #{176}C.
Although
these
specific
rate constants
represent
water
exchange
in
aquo metal
ions, they also reflect relative
rates of exchange
of other
uni-dentate
ligands.
Chelated
ligands
exchange
more 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
sufficient
reason
for Al3” inhibition
of enzymes
with Mg24 cofactors.
Any process
involving
rapid Ca24 exchange
obviously
would
be totally
thwarted
by Al34 substitution.
+
AI3
and
Phosphate
In the human
body, extracellular
fluids contain
about
2
mmol of total phosphate
per liter at pH 7.4 and intracellular
fluids about
10 mmol of total phosphate
per liter at pH 6.6.
Al34 forms an insoluble
salt with phosphate,
often designated as AIPO4, or sometimes
as A1PO4
2H20, corresponding
to the composition
of the mineral
variscite.
At 0.16 mol/L
ionic strength
the ptC values
for successive
deprotonations
from H3P04
-.
H2P04”
-.
HPO
P043”
are 2.0, 6.77,
and 11.6 (9). Thus P043” is the dominant
phosphate
species
atpH
>11.6 while equation
1 and Figure
1 shows that free
Al34 is the dominant
Al3”' species
only at pH <5.5. Thus
significant
amounts
of both Al3” and P04
are incompatible in solution
at any pH. If we seek the overall
neutral
complex
for which
there
is compatibility,
we note from the
PICa 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 to
6. For the purposes
of solution
chemistry
it is advantageous
to rewrite
AlP04
2H2O as Al(OH)2
H2P04.
The solubility
product
of variscite
has been reported
in
terms of the hypothetical
reaction
.
-*
.
1802
CLINICAL
CHEMISTRY,
Vol. 32, No. 10, 1986
Al(OH)2H2P04
for which
the
Al3”'
equilibrium
+
constant
20W
+
H2P04”
is
K9#{176}
=
(Al34)(OH’’)2(H2P04)
=
where the superscript0
designates
a thermodynamic
or zero
ionic strength
equilibrium
constant,
the parentheses
signify
activities,
and the value of plCg#{176}
=
30.5 refers to 25#{176}C
(35).
We immediately
recast
the solubility
equation
in terms
of
dissolution
by hydrogen
ion according
to
Al(OH)2H2P04
for which
the
+
2H4
equilibrium
±
Al3”
constant
K0#{176}
=
(Al34)(H2P04”)/(H”’)2
+
2H20
+
H2P04”
is
=
K90/K2
=
10-2.5
for the ion product
constant
of water pK = 14.0. We need to
find the concentration
equilibrium
constant
at 0.16 mol/L
ionic strength
and recast
the last equation
as
&
1o
0
-
r A 1341
-
L”
ru
flf
JYAlLL21P4
,,tx+2
-‘i
iY-”
-
/
-
ztloYAlY_
where
the brackets
signi
molar
concentrations.
The results now become
approximate
because
we can only estimate the activity
coefficients
at 0.16 mol/L ionic strength
as
=
0.14 and, for H2P04”,
y
=
0.73 according
to a
common
treatment
(36). Substitution
of these
values
into
equation
9 yields pK
=
-1.5.
From pH 3 to 11 there are
only two predominant
phosphate
species,
and the total
phosphate
concentration
is given
by T = [H2P04]
+
[IWO4].
The fraction
of the total
phosphate
that
is
H2P04”
is given by [H2P04iPF
=
(H4)/[(H4)
+ K,
where
plC2 = 6.77 at 0.16 ionic strength
(9). Substituting
these
considerations
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
where
(13)
the last equality
applies
to 0.16 molJL ionic strength.
13 furnishes
an estimate
of the free Al34
concentration
at 25#{176}C
allowed
by the solubility
of variscite
at 0.16
mol/L ionic strength
for a designated
3 <pH
<11
and
total phosphate
concentration,
T. (Supersaturation
and ion
pair formation
increase
the Al34 solubility.)
For intracellular fluids at pH 6.6 containing
10 mmol of total phosphate
per liter, equation
12 yields
for -1og1Al34]
=
pAl = 12.5,
while for extracellular
fluids at pH 7.4 containing
2 mmol of
total phosphate
per liter, pAl = 12.9. This pair of values
appears
in the second
row of Table
1 and represents
extremely
low maximum
free Al3”’ concentrations.
Under
the
same pair of conditions
the free Ca24 concentrations
permitted by insoluble
hydroxyapatite,
Ca5(P04)30H,
a principal
constituent
of bones and teeth,
are nearly
i0 times greater
(9). Thus, thermodynamically,
Al3”' easily
displaces
Ca24
from phosphate
binding.
Antacids
containing
Al(OH)3
deplete
phosphate
by precipitation,
tending
toward
a more negative
Ca24 balance
in
individuals
with low Ca2” intakes
(37). Dialysis
osteodystrophy
develops
with increased
Al34 concentrations
in the
presence
of normal
concentrations
of Ca24 and Mg24 in
plasma
(38). Even low accumulations
of Al3”' impair
bone
mineralization
(39). The presence
of an Al(OH)2H2P04
precipitate
may interfere
with the orderly,
kinetically
controlled deposition
of the bone mineral,
hydroxyapatite
(40).
Table
1 shows that in the presence
of typical
phosphate
concentrations
the limitation
on the free Al3”’ concentration
is not the solubiity
of amorphous
Al(OH)3;
rather,
it is the
more limited
solubility
of Al(OH)2H2P04.
Obviously,
experiEquation
ments
with Al34 in a phosphate
buffer should
be avoided.
When attempted
(41), such experiments
lead to unreliable
conclusions.
The difference
between
the free Al34 concentrations allowed
by Al(OH)3
and Al(OH)2H2P04
widens
as the
acidity
increases,
as illustrated
by the results
in Table
1 at
pH 4.0, where
the factor
becomes
iO.
Gastrointestinal
absorption
of Al34
occurs
from
Al(OH)3
but not from
Al(OH)2H2P04
antacids
(42).
There
seem to be two possible
explanations
for the extremely
low concentrations
of Al34 in living
organisms.
Either
the Al3”' locked in the earth’s
crust has been inaccessible to life, or biological
systems
have
evolved
to reject
Al34. If inaccessibility
is the answer,
release
of Al34 by acid
rain appears
more dangerous
than if some rejection
occurs.
Current
evidence
suggests
that resistance
to Al3”’ uptake
by
living
systems
results
accidentally
from the ubiquity
of
phosphate
and insolubility
of Al(OH)2H2P04,
and that when
this
system
becomes
ineffective,
little
defense
remains
against
Al34 uptake
in acidic solutions.
Plants
that accumulate Al34 do so in acidic soils (43), and evidently
detoxify
the
Al34 by possessing
basic
fluids
or organic
chelators.
We
examine
next a molecule
that solubiizes
Al34 so effectively
that it renders
phosphate
ineffective
in resisting
Al34 uptake.
AI3
and
Citrate
Because
about
0.1 mmol of citrate
is present
per liter of
blood plasma,
it becomes
the pre-eminent
small-molecule
plasma
binder
of a metal
ion such as Al34 that
prefers
oxygen
donor ligands.
In the following
discussion
we use a
typical
concentration
of citrate
in plasma,
0.1 mmol/L.
A
recent
assessment
of the citrate-Al3
system
recommends
the equilibrium
constants
that
are appropriate
for 25 to
37#{176}C
and 0.10 to 0.16 moIJL ionic strengths
with all plC5
values
in the commonly
used scale of activity
in hydrogen
ion (20). These
recommended
constants
are used in this
paper.
Slow polymerization
reactions
of Al3”’ complexes
requiring
hours to complete
(44) are excluded,
because
they
are unlikely
to occur
in a biological
milieu.
The three
successive
acidity
constants
(pICa values)
for citric
acid
under the above conditions
are 3.0, 4.4, and 5.8. Thus citric
acid occurs as the tricarboxylate
anion citrate
at pH 7.4 in
the plasma.
In acidic solutions
Al34 reacts
with citrate
di-anion,
M34
+ L112”
MLH”',
with log K0 = 4.7, and this complex
undergoes
deprotonation,
MLII”
±
ML + H4, with pKob =
2.5. Still, in acidic solutions
Al3”' reacts with the trianion
to
give an important
complex
of zero net charge,
M34 + L3
with K8 = [ML#{176}]/([M34][L3])
=
1080. This neutral
complex
undergoes
deprotonation,
ML#{176} MLH,,’
+ H4 in
quite acidic solutions,
with pKb = 3.4. The MLIL
complex
contains
three anionic
carboxylate
groups
and a deprotonated citi ate hydroxy
group.
Finally,
the neutral
complex
may
add a second citrate,
ML#{176}
+
L3”
ML’,
with log K2 = 5.0
(20).
Figure
3 shows
the species
distribution
of the several
complexes
in a solution
1 mol/L
in total
Al34 and 0.1
mmol/L
in citrate.
Only a little AlLH4
forms near pH 3
where the net neutral
A1L#{176}
rises to 0.4 mole fraction.
As the
pH increases,
AlL#{176}
is succeeded
by the strong
citrate
hydroxy deprothnated
complex
AlLH1’.
Even at this 100-fold
citrate
to Al34 mole ratio there
is little of the 2:1 complex
AlL23”. At pH values
only slightly
exceeding
the pH 7.4 of
plasma,
Figure
3 shows
that the water-derived
Al(OH)4
species
abruptly
intrudes
and dominates
at pH >8.
ML
C
0
H:,
0
1,
0
,I/\
MLH\LML’
I
2
3
4
5
6
7
8
9
pH
Fig. 3. Species disttibution for 1 tmol of AP” and 0.1 mmol of citrate
per liter, plotted as mole fraction (aluminum basis) vs pH
General features of the curve are largely independent of the metal-ion concentration, 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 pass
through
membranes.
Figure
4 shows how the mole fraction
of this single species depends
upon the citrate
concentration.
About half of all Al34 in solution
appears
as ML#{176}
near pH 3
in the presence
of 0.1 to 100 mmol
of citrate
per liter.
Significant
quantities
of net neutral
ML#{176}
occur even with
the lowest citrate
concentration
(10 mo1fL)
in Figure
4.
The significant
mole fraction
of net neutral
ML#{176}
from
pH
2 to 5 for a range
of citrate
concentrations
such as that
shown
in Figure
4 suggests
that
citrate
complexation
of
Al34 provides
an effective
means
for Al34 absorption
into
the body in the upper
region
of the gastrointestinal
tract.
This conclusion,
derived
from stability
constants,
is supported by experiments
with rats. Increased
Al34 concentrations
were found in both the brain
and bones of rats fed a diet
containing
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 humans
who are taking
an Al(OH)3-based
antacid
is enhanced
substantially
by concomitant
intake
of citrate
(47). Therefore, not only does citrate
solubilization
defeat both hydroxide and phosphate
precipitation
and elimination
of Al34, but
both equilibrium
arguments
and animal
experiments
reveal
that
absorption
occurs
as well. This
process
shows
that
people should
not take aluminum-containing
antacids
with
04
.0
0
0
LA,
a
02
2
4
3
5
pH
Fig. 4.
Neutral
citrate AIL#{176}
mole fraction vs pH
From left it, right, curves represent 100 to 0.01 mmol of cItrate per liter, M curves
are for I ,umol total AP per liter, but the general features of the curves depend
little upon the metai-ion
concentration.
The dashed
curie corresponds
to the
dashed curve in Fig. 3
CLINICAL
CHEMISTRY,
Vol. 32, No. 10,
1986
1803
fruit or juices.
Although
healthy
individuals
exclude
Al34 from their
systems,
solubiization
of Al34 by citrate
provides
a means
by which even the healthy
individual
may
absorb
Al34.
Table
1 summarizes
the permitted
free Al34 concentrations in the presence
of Al3-complexing
agents.
The results
are expressed
as pAl = -log
[Al34],
so that
the larger
numbers
represent
the lowest
free-Al34
concentrations.
At
all pH values
considered,
Al34 is removed
from solution
more 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 plasma, Al34 is most apt to be complexed
with citrate.
Further
analysis
shows that at lower pH values
such as might occur
in the stomach,
citrate
complexes
are again
favored,
with
some 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 no
stability
constants
have been determined.
In the contest
between
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
the
amount
of citrate
have
overlooked
a significant
variable
that may affect the conclusions
drawn.
Citrate
complexes
of Al34 inhibit
precipitation
of calcium
phosphate
at pH 7.4 (49). Under
the experimental
conditions the main complex
should
be AILW1
(Figure
3). On the
time scale of the precipitation
experiments
the complex
may
polymerize
to give Al3(OH)4L34
(44), a process that requires
base. Because
the total
amount
of Al34 present
is
small, this base-consuming
process should not interfere
with
the interpretation
of the base-consuming
precipitation
experiments.
citrus
Al3
and
Fluoride
Al34 forms relatively
strong
complexes
with fluoride
ion,
F-. Representative
successive
stability-constant
loprithms
for the addition
of 1 through
5 fluoride
ions to Al
at 2537 #{176}C
and 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
of
Al34 are stronger
than those of Fe34. Figure
5 shows the
distribution
curves
of aluminum-fluoride
complex
species
as a function
of pF = -log[F”’],
where
[F”’] is the molar
concentration
of free uncomplexed
fluoride
ion. From Figure
5 it can be seen the neutral
AIF3
complex
predominates
near
0.3 rnmol
of ambient
F”’ per liter,
but exists
in a
significant
mole fraction
from 0.02 to 5 mmol
of ambient
fluoride
per liter.
Ingesting
relatively
small amounts
of Al(OH)3
decreases
the absorption
ofF”
(51) from the intestine.
Along with the
concomitant
decrease
in phosphate
and Ca2 (52), Al34 thus
exerts
a threefold
adverse
effect on bone structures.
High
Al34 and low Ca24
and F
in home
water
supplies
of
dialysis
patients
leads to encephalopathy
and bone fractures
(53).
A surprising
linkage
ofF” with Al34 occurs in the story of
F’ activation
of the adenylate
cyclase
(EC 4.6.1.1)
enzyme
system.
Almoitt
since its discovery
it has been known
that
this enzyme
can be activated
by F as a nonphysiological
activator.
Later experiments
revealed
that the enzyme
has
two components,
and that
the F”’ acts on the guanine
nucleotide-binding
regulatory
component,
which
requires
both Me4
and a nucleotide
such as ATP for activation.
Most recently
a contaminant
in many
commercial
prepara+
1804
CLINICAL
CHEMISTRY,
Vol. 32, No. 10, 1986
tions of APP has been blamed
for the apparent
nucleotide
requirement.
As with hexokinase,
the contaminant
is Al3”
(54). Of several
metal
ions tested,
only
Be24 provided
similar
activation
of the regulatory
component.
Both Al34
and Be24 form a series
of F” complexes.
Although
the
anionic
complexes
AIF4
and BeF3
should
have been the
most prevalent
under
the experimental
conditions
(5 mmol
of added F per liter), there were also significant
amounts
of
the neutral
complexes
AIF3 and BeF2. Because
the regulatory component
is prepared
as a neutral
detergent
extract
from membranes,
I suggest
that it is passage
of the neutral
complexes
through
a hydrophobic
environment
that is responsible
for the activation.
In support
of this proposal,
investigators
in another
study find the most striking
effects
of A13 at the lowest F concentration,
0.2 mmol/L
in their
case (55), which
is at the optimum
for AIF3#{176}
formation
in
Figure
5.
However,
any interpretation
based
only on equilibrium
properties
presents
two difficulties.
First,
the experiments
(54) were performed
in the presence
of 1 mmol of EDTA per
liter, which
strongly
chelates
Al34, leaving
iO
times
as
much free Al34 as required
to combine
with F. Second,
of
the several
metal ions tested,
Sc”’ was ineffective
in activating the regulatory
component
(54), yet it forms a series of F
complexes
of the same strength
as Al3 and a strong
EDTA
complex (11). These two difficulties
may be explained
by the
introduction
of rate effects.
Al34 is known
to react slowly
with EIYFA.
In the rate
of water
substitution
from the
hydration
sphere,
Sc3”’ reacts
up to i0 times more rapidly
than Be24, which in turn reacts
102 times faster than Al34
(34). Rates of water
loss carry over to reactions
with other
ligands,
and so I account
for the experimental
results
by the
slow reaction
of Al34 with EDTA permitting
formation
of F”’
complexes,
of which the neutral
AIF3 is perhaps
the relevant
form.
Sc3”’ undergoes
relatively
rapid
sequestration
by
EDTA to form a strong
complex,
and F- is not thermodynamically
competitive.
Be24 reacts
at an intermediate
rate,
and in this case the 5 mmol of F” may compete
for the metal
ion with the 1 mmol of EDTA per liter. Compared
with Al34,
Be24 forms
a disproportionately
weaker
complex
with
EDTA
than with F.
This rate-dependent
hypothesis
predicts that prior incubation
of Al34 with EDTA at the pH 8.0
used in the experiment
will not yield activation.
Though
unpublished,
this experiment
had been performed,
and no
activation
occurred
(P. C. Sternweis,
personal
communication, August
1985), in agreement
with the prediction
of the
rate-dependent
hypothesis.
C
0
C’,
a
U-
0
pF
Fig. 5. Mole fraction of total aluminum
function of pF = -Iog[F],
where
found as fluoride complexes as a
[F-] Is the ambient F- molar
AI3
and
Proteins
Are any proteins
likely to be more effective
Al3”' binders
than is citrate?
The common
albumin
and globulin
proteins
of the plasma
bind metals
such as Al34 only weakly
and
nonspecifically.
Albumin
binds several
Ca2”', with a stability log K6
2, and Gd’
with log K8 = 3.9(56).
Competition
with
Chelex,
a cation-exchange
resin,
was
used
in an
attempt
to determine
Al34 binding
at pH 7.4 to two proteins
with reported
dissociation
constants
of 2.0 tmo1/L
for albumin
and a similar
0.52 moI/L
for transferrin
(57). Since
Al(OH)4”’
is the predominant
species,
these
conditional
dissociation
constants
require
a large y correction
according
to equation
4. However,
the similar
values
for the highly
dissimilar
proteins
in the Al34 binding
capacities
suggest
a
flaw in the method.
The 43 pg of total free Al34 per/liter
found
with
Chelex
at pH 7.4 corresponds
to a soluility
between
those for amorphous
Al(OH)3
and gibbsite
(Figure
2). Chelex
binds
Al34 much
too weakly
to provide
any
metal-ion
buffering
for competition
with transferri4.
The
suggestion
that Al34 binds strongly
at the amino
teiminal
site of human
serum
albumin
(58) is unconvincin.
Four
nitrogen
donors-consisting
of the human
albumin
amino
terminus,
two deprothnated
peptide
nitrogens,
and a histidyl ring nitrogen-tightly
chelate
Cu24 and also Ni24, with
spin pairing
of the latter
metal
ion (59). This
strongly
covalent,
quadri-dentate,
tetragonal
chelate
ring
system
cannot
be a strong
binding
site for Al34, which
depends
upon electrostatic
interactions
for binding.
Albumin
is much
too weak a metal
ion binder
to withdraw Al34 from any of the complexes
of Table 1, all of which
occur
in the plasma.
At the pH 7.4 of plasma,
albumin
cannot
compete
for Al34 with hydroxide
or phosphate
precipitation
and citrate
complexation.
If Al34 is to be protein
bound 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 a
normal
concentration
in plasma
of 3 g/L, with two metal-ion
binding
sites per 77 000-Da
protein
at only 30% site occupancy by Fe34 in the plasma,
transferrin
furnishes
unoccupied metal-ion
binding
sites at a concentration
of 50 mol/L.
Because
this
concentration
is half that
of citrate
in the
plasma,
to be competitive
transferrin
needs to bind Al34
twice as strongly
as does citrate
at pH 7.4.
Direct
competition
between
Al3
and Fe3” for binding
sites on transferrin
in combination
with an assumed
binding constant
for Fe3ttransferrin
were used in an attempt
to establish
the binding
constant
for Al34-transferrin
(60).
The key experiments
were conducted
with a large excess of
Al3”,
and solutions
were
examined
after
only 30 mm.
However,
it takes hours
for citrate-bound
Fe34 to displace
Al34 from
Al34-transferrin,
and therefore
the calculated
Al34 binding
constant
is much too high. The results
may be
fitted
by assuming
that
citrate-bound
Fe34 reacts
with
transferrin
100 times more rapidly
than does citrate-bound
Al34 (see Exchange
section
above) and insignificantly
in 30
mm with Al3”'-transferrin.
A recent
quantitative
spectroscopic
determination
of Al34
binding
to the two sites
of transferrin
yields
successive
stability
constants
of log K1 = 12.9 and log K2 = 12.3 under
blood-plasma
conditions
of pH 7.4 and 27 mmol of HCO3”’
per liter (33). For 1 mol
of total Al3 per liter of plasma,
the
free Al34’ concentration
permitted
by transferrin
is 10_146
mol/L. As indicated
in Table 1, this amount
is less than that
allowed
by insoluble
Al(OH)3,
Al(OH)2H2P04,
or by corn-“
plexation
citrate.
with
Thus
transferrin
is the
ultimate
car#{231}ierof Al’
in the plasma.
The same
study
found,
for
the successive
stability
conof Fe34 binding
to transferrin,
log K1 = 22.7 and log
K2 = 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 weaker
than
expected,
and it is suggested
that
the significantly
smaller
Al34 cannot
coordinate
to all the transferrin
donor
atom
available
to Fe3t
Qualitatively,
this
study
shows
that with a citrate/transferrin
molar
ratio
comparable
to
that found
in the plasma,
citrate
releases
both Al3”’ and
Fe34 to transferrin.
Because
of the nearly
10 log unit
difference
in their binding
strengths,
Al34 cannot
displace
Fe3”’ from transferrin.
However,
with most plasma
transferrin carrying
unoccupied
binding
sites, there are sufficient
resources
for binding
both of these metal
ions.
stants
I am
to Michael
grateful
for their
Bertholf
R. Wills, John Savory,
and Roger
encouragement
in my writing of this article.
L.
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