Chapter 2. Chemical Properties I. General
K
ey :
In this chapter we will explore in a summary way the chemical properties of minerals.
Our goal is to understand how minerals fit into the framework of living systems keying on the
properties that underlie their functions.
O
bjectives:
1. To define 5 fundament chemical properties of minerals related to function,
2. To comprehend how minerals behave in biological settings.
I. CHEMICAL PROPERTIES
The chemical properties of minerals give us insight into the form of a mineral in its
natural environment as well as a rationale for selectivity, bioavailability, and antagonism. Many
minerals, particularly metal ions, exist as complexes, a large number of which are with
enzymes. Are these complexes constructed by accident or do they adhere to a fixed chemical
design? To answer the question, consider that complexes of trace metals show uniqueness as
to numbers of attaching molecules (ligands) and spatial geometry. Clearly, bond angles and
distances must be considered when dealing with such complexes. Indeed, chemical properties
of minerals cannot be separated from biological functions. Five major properties that
contribute to biological selection of minerals are:
1) charge or valence state of the ion
2) solubility
3) redox property
4) coordination geometry
5) preference towards specific ligands
We will explore each of these to as they related to the macro- and microminerals.
1. Valence
Minerals in living systems exist as positive or negative ions. The ions of sodium,
potassium, calcium and magnesium have alkaline properties and bear positive charges. In
contrast non-metals such as chloride, sulfates and phosphates have negative charges and acidic
properties. For the most part macrominerals have only one stable valence state, whereas
microminerals exist in multiple valence states, zinc (Zn2+) being an exception to this rule. As
shown in Table 2.1, macrominerals have only one valence state whereas microminerals such as
copper and iron have two or more stable cations, Cu+ or Cu2+ and Fe2+ or Fe3+, respectively.
Table 2.1. Valence state of macro- and microminerals.
Macrominerals
Valence
Sodium
Potassium
Magnesium
Calcium
Chlorine
Na+
K+
Mg2+
Ca2+
Cl-
Microminerals
Iron
Zinc
Copper
Manganese
Cobalt
Nickel
Molybdenum
Iodine
Valence
Fe2+, Fe3+
Zn2+
Cu+, Cu2+
Mn2+, Mn3+, Mn4+
Co+, Co2+, Co3+
Ni+, Ni2+
Mo4+, Mo5+, Mo6+
I-
Having multiple stable valence states imparts a freedom to move between states which allow
the mineral to exchange electrons during reaction or assume unique multiple spatial
geometries. Losing or gaining electrons is a very important chemical property of metals ions
designated as “redox” metals as discussed below.
2. Solubility
Water is the solvent of life. This is tantamount to saying most reactions of life take
place in aqueous media. For minerals, solubility is manifested as an equilibrium between the
constituent ions and solid form, a reaction that strongly depends on the pH of the medium.
Hydrated ions or polar complexes are capable of free movement. Sparingly soluble ones
require water-soluble proteins to transport them. As shown below, the ionization constant
which measures the ratio of free ion to solid becomes smaller as atomic number increases.
Monovalent macrominerals have the greatest propensity to exist as free ions whereas divalent
macrominerals are two orders of magnitude or more less in free ion quantity.
Na+, K+
10-1M
Mg2+, Ca2+
10-3M
Zn2+
10-9M
Cu2+
10-12M
Fe3+
10-17M
Thus charge is not the only determinant of water solubility. Zinc, copper and iron salts are
sparingly soluble in water at neutral pH despite a strongly positive charge on their ions.
Insolubility in this instance is due to a hydrolysis reaction in which the metal ion forms a bond
with oxygen of H2O, discharging a proton and forming an insoluble hydroxide such as Fe(OH)3
or Cu(OH)2 that separates from solution. Increasing the acidity by lowering the pH shifts the
equilibrium away from the hydroxide and towards the free ion and water as illustrated in the
figure below.
H+
Fe(OH)3 (solid)
Fe3+ (aqueous) + 3H2O
Zn(OH)2 (solid)
Zn2+(aqueous) + 2H2O
Ferrous ion (Fe2+) is more soluble in water and is favored over ferric (Fe3+) for absorption.
Adding vitamin C or a comparable reducing agent to the diet will therefore enhance the
solubility of iron and promote its uptake by intestinal cells. Not to be overlooked is the acidic
environment of stomach acid, which also promotes solubility and absorption. Ca3(PO4)2 is
moderately ionized in aqueous solution and only a small fraction is absorbed as the free ion
from the intestinal track. Insolubility in this instance is caused by a strong association between
the phosphate and the calcium, an interaction that is only partially be overcome by water
molecules. Insolubility impedes a mineral’s ability to exist in the aqueous phase and hence can
have an impact on its bioavailability.
3. Oxidation-Reduction
Metal ions with multiple valences have the additional property of donating or accepting
electrons during reactions. Reactions in which electrons are transferred are called oxidationreduction reactions and metals capable of these are referred to as redox metals. Redox metals
exist in a minimum of two stable valence states and movement between these states presents
an energy differential that is referred to as a redox potential. The value obtained is a measure
of the affinity of the metal ion to retain its valence electrons and thus allows one to predict
electron flow from donor to receiver when two ions at equal concentration are present in
water. This is illustrated in the figure below. Oxidation or loss of electrons occurs with a gain in
the positive charge; reduction is a loss of positive charge as illustrated by ferrous iron being
oxidized by Cu2+. Reading from left to right, Fe2+ (ferrous form) is oxidized
Cu2+ + Fe2+
Cu+ + Fe3+
to Fe3+ concomitant with Cu2+ (cupric form) reduced to Cu+ (cuprous form). A change in
valence signifies a single electron has been transferred between the two ions. The reduced
copper is less positive and the oxidized iron more positive. Reading the equation from right to
left, Fe3+ would be considered the electron acceptor or oxidant and Cu+ the electron donor or
reductant. Thus redox minerals in biological systems have the potential to serve as both
oxidants and reductants, not only among themselves, but against organic molecules as well. In
so doing the redox metals behave as both oxidants and antioxidants in a biological system.
4. Coordination Geometry
Charge is not the only factor that determines the behavior of a metal ion in a complex.
Important to consider are specific geometric configurations with precisely defined spatial and
stereo geometries. As illustrated by iron and copper in Fig. 2.1, ionic radii vary with the
oxidation state of the metal ion, which is particularly important for selectively. Seven
biologically essential elements are part of the first transition series of elements. With the
exception of zinc, each element in the series is characterized by partially filled 3d orbitals that
give the metal ion complex a unique structure. The coordination number or “secondary
valence” gives the total number of atoms (donor atoms or ligands) that can bind to the metal.
Binding of two or more must satisfy the spatial geometry of the complex, a situation very
similar to hybrid orbital states of carbon, sp, sp2, sp3, for linear, square planar, and tetrahedral,
respectively. Some complexes are shown in Fig 2.2.
3d10
3d9
Figure 2.1. Ionic radii as a function of charge on the ion. Numbers indicate radius length in
Angstroms.
Coordination Complexes
Table 2.1 summarizes chemical parameters for a number of metal ions, drawing
attention to similarities in factors that relate to complex formation. Based on three
parameters, it is possible to discern and predict metal ion antagonism and to show why only
one metal ion will be suited for a particular functions. For example, Zn2+ and Cu+ prefer
tetrahedral complexes that engage four ligands. Assuming their ionic radii are comparable,
these two metals have the potential to compete for similar binding sites on proteins. It is
known that a high level of dietary zinc will suppress copper uptake. Similarly, Cd2+ competes
with Zn2+ for binding sites on the metal storage protein metallothionein. Hg2+ is unique in
showing no such interaction with zinc, copper or cadmium. Instead, mercury prefers a linear
complex and hence will show minimal interference with any of the metals because it cannot
accommodate the same stereochemical site as these other metals. Ag2+ ions are strong
antagonists of copper. Both ions can engage in a square planar arrangement. Iron, however, is
at its lowest energy state when it forms an octahedral complex, thus making it less likes to
antagonize zinc or copper. In its ionized state, Fe2+ exists in an octahedral arrangement with 6
binding ligands (Fig. 2.3). This chemical signature is not duplicated by copper or zinc, which
have more of a tendency to form square planar or tetrahedral complexes that bind at most 4
ligands. Failure to form octahedral complexes lessens the likelihood that copper and zinc will
appear in hemoglobin, of if so to duplicate the functions of iron in that protein. Likewise, iron
is unable to bind to sites designated for copper or zinc. Specificity in electronic configuration is
one of
Table 2.2. Coordination Complexes of Transition Metal Ions
Ion
Orbital
Cu+
d10
(tetrahedral)
Zn2+
d10
(tetrahedral)
Cd2+
d10
(tetrahedral)
Hg2+
d10
(linear)
Cu2+
d9
(square pl.)
Two
Ag2+coordination
d9 complexes
(square pl.)
Fe2+
d5
(octahedral)
(CH3)3Si
Configuration
Coordination No.
sp3
4
sp3
4
sp3
4
sp
2
dsp2
4
dsp2
4
d2sp3
6
NCH3
Si(CH3)3
N-Co-N
CH3N
NCH3
Cu
Si(CH3)3
(CH3)3Si
(a)
CH3N
(b)
Co
NH2
H2N
Zn
Fe
NH2
H2N
NH2
H2N
CH3
H3C
Se
NH2
H2N
CH3
(c)
Co
(d)
Figure 2.2. Stereo complexes of metals with different ligands.
(e)
strongest determinants for selection, which ordains that only the metal ion in question is able
to partake in a biologically active form and perform the particular function. It also forms the
basis for metal ion interactions which can be both antagonistic and synergistic.
5. Preference for Ligands
As their name implies, ligands are basically groups of atoms that form chemical bonds
with metal ions. They literally tie up the metal ion, capturing it to render it either biologically
active, more water soluble, or part of a complex. Ligands that form biological complexes with
metals tend to favor N, O, and S atoms in the group as shown in Fig. 2.2. These atoms are
chosen because they can provide electron pairs that form coordinate-covalent bonds with the
metal ion. Only rarely does a metal-carbon bond occur in biology, the most familiar example
being the bond between cobalt and the corrin ring of vitamin B12. Ligands-binding groups on
proteins tend mostly to be arise from the carboxyl groups of aspartate and glutamate, the
imidazole group of histidine, and the sulfur groups of cysteine and methionine. These amino
acids must be in the proper order to form a pocket in the protein for the metal to fit. Metals
too large or too small, or incapable of assuming the proper three-dimensionality or the pocket
are excluded. Only those that meet the strict structural criteria will engage the protein or
complex.
D. SUMMARY
The minerals in living systems are not alien to the laws of chemistry. The ease with
which particular minerals moves though a system, cross over membranes, engage other
components, perform chemical reactions, all depend on chemical properties of the mineral.
The environment of the internal milieu is largely aqueous, which makes solubility in water at
neutral pH an important determinant of ease of movement and absorption by cells. Minerals
that have high water-soluble properties tend to be more easily absorbed in the intestine and at
the cell membrane. Those insoluble at neutral pH tend to require ligands such as protein to
obtain solubility properties. Minerals that have ease of donating or receiving electrons have
the potential to participate in oxidation-reduction reactions associated with antioxidant activity
and enzyme catalysis. In their interaction with organic components minerals form specific
complexes with defined geometries. This property underscores mineral selectivity and assures
that only the mineral that allows a proper fit to be selected. The chemical properties of
minerals extend into other phenomena that are in the realm of the transition metals, of which
biological microminerals make up a large fraction. Insight into how chemical properties
translate into these important functions will be discussed in the following chapter.
PROBLEMS
1. List 3 instances where the solubility of a mineral is important to its behavior in a biological
system.
2. Based on the ionization constants shown in page 2, determine the ratio of free ion to
complex (insoluble) ion for each of the metals. Based on your results what would be the
concentration of free calcium in plasma? What would be the consequences if calcium in the
plasma was to increase ten times?
3. Suppose you mixed copper ions with zinc ions under conditions normally found in plasma
(37oC, pH 7.4). Which ion pair would donate and which would receive electrons. (HINT: look at
Table 2.1 before answering).
4. As shown in Fig. 2.1, which ion of copper is larger? Offer an explanation based on chemistry
why this is the case?
5. In Table 2.2, In the treatment of Wilson’s disease, a condition characterized by an excess
accumulation of copper in the liver and brain, which metal salt would you added to the diet to
prevent a Wilson patient from absorbing too much copper? It is known that the cuprous form
of copper is favored in absorption across the intestine.
6. Would mercury ions be expected to interfere with the absorption of iron? Explain your
answer.
7. Explain how a genetic mutation in which an alanine substitutes for a histidine could affect
mineral metabolism.