Chapter 2. Chemical Properties II. Determinants of
Biological Properties
K
ey :
To fully comprehend the biological properties of minerals will require knowledge of
their chemistry. Valence states, solubility, oxidation-reduction capacity, proteins and nucleic
acid interactions, discussed in chapter 2 are a start. In this chapter we continue the importance
of chemical properties by keying on the arrangement of electrons around the nucleus. As we
will see, this will determine how a mineral interacts with a protein, nucleic, etc. and can often
determine color, electron exchange, and other facets of mineral function. This is a first step
towards understanding the basis for a mineral’s adeptness to perform a specific biological
function.
O
bjectives:
1. To learn the basis of quantum theory for determining electronic configurations
2. To associate electron configurations of minerals with stereochemical properties
3. To be able to predict a mineral’s ionic state from its electron distribution
4. To learn the role ligands play in deciding a mineral’s function and properties
I. Quantum Theory in Chemistry
Quantum theory is the basis of modern chemical thought. According to the theory
electrons are distributed around a central nucleus of an atom in orbitals that have discrete
energy states or quanta. The German physicist Shrödinger described orbitals as wave functions
and solved a very complicated equation to arrive at wave numbers and electron density. What
prompted Shrödinger to postulate electron motion as waves was a need to solve the mystery
of ionization energies (the energy required to remove an electron). As one moves horizontally
through the rows of the periodic table of elements, ionization energy increases in a cyclic
fashion culminating at its highest level at the noble gases in each row. Why some elements
were quick to lose their outer shell electrons whereas others retained them tenaciously
required an explanation. A second pressing and equally profound observation in need of an
explanation was the unique banding pattern of an electronic spectrum that showed patterns
typically seen when emitting light is manifested as waves. Shrödinger’s resolved the dilemma
by postulating electrons as particles with a wave-like motion. These concepts were the first to
apply quantum theory to atomic structure and the two have since been forever linked.
2. Quantum Theory Explains Electron Distribution
As one moves through the periodic table starting with hydrogen, each increment in
electrons is tabulated by an increase in atomic number (Z). As Z grows, so do the value of
quantum numbers that basically describe the energy status of atomic orbitals, their size, shape
and orientation, whether or not they are empty or occupied with electrons. Older ideas
classified orbital energy states as shells, K, L, M, etc. that electrons occupied. All were basically
spherical in shape which meant energy was basically determined only by the distance from the
nucleus; the greater the distance, the higher the energy of the electron. Geometrical shape
and other factors were ignored in these early theories. Quantum and wave theory, however,
considers not only distance but orbital shape, occupying electrons and spin states, i.e., four
quantum numbers must be taken into account. The principal quantum number n specifies the
major energy level of the electron in the atom, which is an index of the energy of the orbitals
within that level. For the hydrogen atom, for example, n = 1 which basically says only one
energy level is permissible. Second, third and fourth energy levels bring additional quantum
numbers: the azimuthal that specifies angular momentum or shape, the magnetic quantum
number that specifies orbital orientation away from the central nucleus and the spin quantum
number that prevents duplicating the energy of the two electrons within the same orbital.
Thus, the lightest element, hydrogen, has the lowest energy orbital, 1s, within which there is
one electron. Filling the 1s orbital with its full complement of 2 electrons gives rise to the noble
gas, helium. The hydrogen atom, therefore, is 1s1 and helium is denoted as 1s2. The
nomenclature for tabulating electrons in orbitals is shown below.
principal
quantum
number
No. of
electrons
1 s2
orbital
(subshell)
A quantum theory’s view of an atomic structure, therefore, is to consider orbitals as
regions of highest probability for finding electrons as they traverse their way around the
central nucleus of atoms. Individual atomic orbitals such as 1s in hydrogen have at most two
electrons of opposite spin. Elements higher than hydrogen and helium move into the second
energy level where an additional orbital, the p orbital is present. P orbitals feature at most 6
electrons distributed in 3 equal energy atomic orbitals of two electrons each, viz., px, py, and pz
(Fig. 2.1). A d orbital comes into play at the third energy level. As discussed below, d orbitals
accommodate a maximum of 10 electron distributed spatially in 5 atomic orbitals. Biologically
important metals extend into the 2nd, 3rd, and 4th quantum energy levels, displaying s, p, d, and f
subshells.
3. Quantum Theory Predicts Orbital Shape
All s orbitals regardless of quantum energy level are spherically symmetric whereas
those called p are cylindrically symmetric along the long each axis (Fig. 2.1). Note that all p
orbitals have the same basic shape, but differ in their spatial orientation. A p designation for
any value for n will always refer to three orbitals of equal energy status oriented along the
three long axes and accommodating at most 6 electrons. A d designation will relate to 5
orbitals of equal energy state accommodating at most 10 electrons. As will be discussed
below, the d orbital arises by taking into account all of the possible ways to redistribute 10
electrons in a spatial arrangement around the nucleus in such a way that electron-electron
interaction are minimized and no single electron energy is unique.
Px
Figure 2.1. Shape of P Orbitals. Typical of a p orbitals is a sausage-shaped extension along each
long axis out from the nucleus and a node near the origin. P orbitals can accommodate as many as
six electrons which distribute as px2 py2 pz2 as shown in the figure
4. The Fourth Quantum Number …Spin State
Finally, according to Pauli’s exclusion principle, no two electrons in an atomic orbital of
an atom can share the same set of quantum numbers. Consequently, two electrons that share
3 of the quantum numbers must still be different. This gives rise to the notion of spin state,
which basically says an electron spins around its axis like a top. Thus, those electron’s with 3
identical quantum numbers must be opposite in their spin states. Spin states may be either +1/2
or -1/2, the choice is arbitrary and simply denotes that spin is in opposite directions. These
states are generally designated spin up and spin down as represented by the notations
respectively. To avoid countering Pauli’s principle, Hund’s rule recognized that lower
energy states of orbits are reached when there is the greatest interaction between electrons in
neighboring orbitals. Taken another way, Hund essentially recognized the importance of
geometric shape of orbitals with emphasis on using as much space outward as possible. This
idea gave substance to notion that the spin states for two electrons must not only be different,
but single electrons occupying a multi-electron orbital favors one in each orbital. To help you
understand Hund’s rule, think of three electrons occupying the three p orbitals. Hund’s rule
recognizes that the lowest energy, most favorable state occurs when there is no pairing, i.e.,
one electron in each of the px, py, and pz. Moreover, the electrons must be of the same spin,
thereby assuring maximum clearance.
5. 3d ORBITALS
The five orbitals that comprise the 3d subshell are shown in Figure 2.2. When combined
they give rise to 5 atomic orbitals as shown in Fig. 2.4. Special recognition in mineral nutrition
must be given to elements whose electrons extend into the third and fourth quantum levels
and occupy 3d and 4s orbitals. Elements that fit this category are very familiar, such as iron,
copper, zinc, manganese, all known to be essential in biological systems. 3d elements make
their appearance in the fourth row of the periodic table beginning with potassium (Z = 19) and
calcium (Z = 20), both beyond argon [Ar]. Potassium and calcium however, fill the 4s orbital
before filling the 3d. Hence ionization occurs with the loss of electrons in the 4s subshell, which
reverts to the closed shell and highly stable argon configuration. Losing one 4s electron gives
rise to K+ and two lost gives Ca2+ respectively, which are giving the designation “closed shell”
ions. This means there are no occupied 3d or 4s orbitals in K+ or Ca2+. The highest filled orbital
for the two is 3p6. To understand why 4s is loaded first, consider that a 4s orbital is actually at a
lower energy state than a 3d. Consequently, as atomic number increases and more electrons
are added, the 4s is filled before the 3d. This important property gives rise to a series of
elements called the transition elements that are characterized as having partially filled 3d
orbitals and filled 4s. Chromium (Cr) and copper (Cu) are exceptions to this rule in that both
have one, not two, 4s electron in place before filling their 3d orbitals. Upon ionization,
however, 4s electrons are lost before 3d. Thus, adding electrons puts 4s orbitals at a lower
energy state than 3d, but discharging them makes 4s the higher energy state.
Z
Z
Y
Y
X
X
dxz
Z
Z
Y
X
Y
d
2 2
X -Y
X
d Z2
Figure 2.2. Shape and Orientation of 5, 3d orbitals. See Fig. 2.4 for a composite picture.
6. The First Transition Series Elements
The 10 elements scandium (Sc) through zinc (Zn) that represents the first transition
series are perhaps the most important for biological purposes. Within this series are 8
nutritionally essential metals (Table 2.1). Transition elements are characterized by highly
colorful complexes that are formed in the presence of neutral or charged ligands. We will see
shortly that the electronic basis for the red color of iron, the blue color of copper, the green
color of nickel, etc. is associated with interaction of light with the 3d orbitals of these atoms.
One element in particular stands out. That element is Zn2+, which in the ionized state has lost
its two 4s electrons but retained the full complement of 10 electrons in its 3d subshell. Because
the shell is filled, electronic transitions between orbitals in response to external light energy is
prohibited. Consequently, Zn2+ complexes are colorless and proteins containing bound zinc
have no special spectra properties. As noted above, the 4s energy state is filled before the 3d.
This will be clear to you by examining Table 2.1 below. In the table, note that the lower energy
orbitals, i.e., lower quantum states, are filled first (represented by the electronic configuration
for Argon) before electrons start to occupy the 3d. The 4s is filled first for all but Cr and Cu.
Scandium
Titanium
Vanadium*
Chromium*
Manganese*
Iron*
Cobalt*
Nickel*
Copper*
Zinc*
[Ar]4s23d1
[Ar]4s23d2
[Ar]4s23d3
[Ar]4s13d5
[Ar]4s23d5
[Ar]4s23d6
[Ar]4s23d7
[Ar]4s23d8
[Ar]4s13d10
[Ar]4s23d10
Table 2.1. Elements in the first transition series. An (*) is placed after those known to be
nutritionally essential. [Ar] is the abbreviation for the electronic configuration of argon.
7. 3d Orbital Splitting
The 5, 3d orbitals as a composite picture are shown in Fig. 2.4. On careful inspection,
one can see that three of the orbitals, viz., dxy, dyz, and dxz extend outward with major lobes
falling between the axes. Only the dx2-y2 and dz2 orbitals extend along the X, Y and Z axes. This
feature signifies that a ligand approaching the metal center will cause greater perturbation to
orbitals oriented along the axes than those between axes. As a consequence of the geometry
the orbitals are split with the dx2-y2 and the dz2 bearing the brunt of the change. The 5 equalenergy orbitals are now replaced by two of higher energy and 3 or lower energy (Fig. 2.3). The
overall energy of the d orbitals does not change, i.e., as much as is elevated is compensated for
the same amount that is decreased when the orbitals are split. For octahedral complexes, the
high energy orbital is designated eg, and the low t2g. These are two new orbitals for
accommodating electrons. What soon will become clear is that the energy differential
between the two split states (o) will depend on the type of ligand. Basically, all ligands are not
the same when orbitals are subject to splitting. Some cause a very large energy differentials
others subtle changes. This is explained below.
eg
x2-y2
xy
xz
yz
x2-y2
z2
o
z2
t2g
xz
xy
yz
Figure 2.3 Octahedral complexes are created by six ligands positioning around a central metal
creating
It can be said that a metal ion with no ligand present has all 5, 3d orbitals of equal energy
status. Such a situation is not possible in the biological world. Even water itself can be a ligand
to a metal ion and cause orbital splitting. Depending on number of ligands bound, different
geometrical complexes can form. Fig. 2.4 show that 6 ligands give rise to an octahedral
Octahedral
Square
planar
Tetrahedral
Figure 2.4 Complexes that form from interaction with 3d orbitals. Upper left shows the 5
degenerate orbitals before being approached by a ligand. All are of the same energy state.
complex (eight-sided figure) typical of heme and vitamin b12. If four ligands are bound, the
complex can be either in a tetrahedral or a square planar arrangement. The electron splitting
noted in Fig. 2.3 applies only to octahedral complexes, however. Table 2.3 shows some of the
more common low molecular weight ligands that bind to 3d transition metal ions in biological
systems.
Table 2.3. Common Ligands. Underline shows atom that contacts the metal ion.
Ligand
Name
FClBrICNNCSSCNOHO2ONOCO
H2O
NH3
Fluoride
Chloride
Bromide
Iodide
Cyanide
Isothiocyanate
Thiocyanate
Hydroxide
Oxide
Nitrite
Carbon monoxide
Water
Ammonia
Name as ligand
Fluoro
Chloro
Bromo
Iodo
Cyano
Isothiocyanato
Thiocyanato
Hydroxo
Oxo
Nitro
Carbonyl
Aqua
Amine
8. High Spin-Low Spin States
As noted in Fig. 2.3, o signifies the energy differential between the two split d-orbitals
in an octahedral complex. The magnitude of this differential depends on the strength of
attacking ligand to forcefully displace the electrons to other orbitals. Weak ligands such as Cland H2O cause only slight perturbations; stronger ones like CN- , NO2- and CO are major
displacers. Judgments made experimentally have resulted in the following series:
Cl- < F- < H2O < NH3 < NO2- < CN- < CO
Moving from left to right in the above series gives rise to ligands with increasingly greater dorbital splitting energy (larger o). Splitting energy is not wasted, however, but instead is fed
back to the electrons occupying the orbitals. Basing orbital splitting on ligand strength helps
explain why a metal ion may exhibit more than one color or even no color because of the ligand
to which it is attached. Strong splitting ligands have sufficient energy to force a pairing or
electrons in orbitals in violation of Hund’s rule. Weak splitting allows the electrons to adhere to
Hunds rule, i.e., exist in separate orbitals with the same spin state but no pairing. The latter
condition is referred to as a “high spin” state; the former, a low spin state. This will become
clearer to you by examining the consequence of orbital splitting, which you now know is
strongly dependent on the type of ligand.
The term “high spin” congers up the image of outer electrons basically existing unpaired, whereas
“low spin” brings to mind paired electrons closer to the center nucleus”
In Figure 2.5 below, note the distribution of the 3d electrons in the split orbitals. A 3d4 ion split
by a weak ligand gives rise to a high spin state of electron distribution. That same ion split by a
high energy ligand is low spin because energy is sufficient to force all four 3d electrons to
occupy the low energy t2g split orbital and induce pairing. Only weak ligands, therefore, can
give rise to a Hund’s rule of electron distribution and only d4, d5, d6, and d7 ions can have a low
spin state.
Note that the limitations do not designate the names of metals with that property. This is because a
3d metals can exist in multiple valence states, such as iron existing both as Fe(II) d6 and Fe(III) d5.
This means the primary valence of the metal ion is a factor in whether it will have high or low spin
potential.
d4
High Spin
Low Spin
Caution: Remember, the high spin state is the LOW energy state and the low spin is the HIGH
energy state. With this in mind its easy to see that a ligand the causes a strong energy differential
between split orbitals will be the one that gives a low spin state, i.e., electron pairing is the outcome
of a high energy endeavor.
d5
High Spin
Low Spin
d6
High Spin
Low Spin
Figure 2.5. d-Electron Configurations for high and low spin octahedral Complexes.
It should also be apparent that a greater preponderance of singular electrons in the high spin state
makes an ion paramagnetic. This contrasts with diamagnetic ions which tend to be void of singular
electrons because of pairing.
9. Ionization
Atoms gain and lose electrons to reach a stable electronic configuration. Metals have a
propensity to lose the outer, higher energy electrons giving rise to ions deficient in one of more
electrons. In viewing the periodic table it is clear that those elements in the first and second
column are prone to lose 2s and 3s electrons, respectively, giving rise to monovalent and
divalent cations. The electrons lost result in a stable electron environment and hence cations
that are formed this way include Na+, K+, Mg2+, Ca2+ are highly stable and have the surrounding
electrons basically in a spherical configuration. Spherical implies no directionality or equal in all
direction. This closed shell configuration for the ions does not permit coordinate-covalent
sharing of electrons with ligands. Hence, engaging a ligand is primarily through electrostatic
interactions and not as geometrically ordered complexes seen with the 3d transition metals.
The closed shell is also prohibitory to electron transitions between orbitals and hence the ions,
regardless of the ligand that binds to them, tend to be colorless.
10. Summary
Quantum theory of atoms helps us correlate electronic states with mineral properties.
For our purposes it gives us insight into the limitations of these properties. Monovalency and
divalency of important minerals can be rationalized by knowing why one or two electrons,
respectively, are lost from the atom when it ionizes. The importance of valence is made even
more relevant when considering metal ions that comprise the list of micronutrients. Metals
such as iron, copper, zinc, manganese, etc. form well-defined geometrical complexes with
proteins and other ligands. The ability of a protein to accommodate a metal ion will depend on
whether that protein can be manipulated to form a binding site that can accommodate the
metal ion. Gaining and losing electrons is a major property of redox minerals. Insights into
electronic structure show us why macro-minerals cannot be used as redox metals and resigned
more to be stable ions. In essence the electronic state of a mineral has a rationale and from
this we can gain some understanding of what a mineral can and cannot do in a living system.
PROBLEMS:
1. Draw the electron configuration of sodium and explain why Na+ is the only ionized
form.
2. What electrons are lost when calcium ionizes?
3. Name all atoms who principle quantum number n = 2
4. Iron exists as 2+ and 3+. Explain how this happens.
5. What is the minimum principle quantum number for an atom with p orbitals? d
orbitals?
6. Why is the color of a metal ion complex dependent on the ligand that engages the
metal?
7. Can an octahedral complex be formed with a metal that has no d orbitals?
8. Show a high spin configuration for the electrons in manganese ion (Mn2+).
9. Which ion has paramagnetic properties, Cu+ or Cu2+?
10. Explain why Zn2+ has no paramagnetic properties.
11. Which of the following would be considered “closed shell” ions: Ca2+, Mg2+ K+, Zn2+
12. Identify the elements that correspond to the following electron configurations:
a. 1s22s22p2
__________________________
b. 1s22s22p6
__________________________
c. 1s22s22p64s23d5
__________________________
d. [Ar]4s23d10
__________________________
e. [Ne]2p63s2
__________________________
13. Predict the valence and identify the element with the following configuration:
a. 1s22s22p63s2
__________________________
b. 1s22s22p64s23d6
__________________________
14. What is the electronic configuration of the following ions: (Hint: ion’s and atoms will
not have the same configuration. Your want the ion’s configuration)
a. H+
__________________________
b. Mg2+
__________________________
c. Cu2+
__________________________
d. Cu+
__________________________
e. Fe3+
__________________________
f. Co3+
__________________________
15. To illustrate the importance of “orbital splitting” in biology, can you explain why
arterial blood is redder in color than venous blood? You may also want to explain why a sign of
a person poisoned by carbon monoxide is a very rich red color to their blood.
16. Look at Fig. 2.5. and explain why in aqueous solution [Fe(H2O)6]2+ is paramagnetic
and [Fe(CN)6]-4 is not. Defend your answer. (Hint: what are the electronic requirements to be
paramagnetic, i.e., influenced by a magnetic field).
17. Explain why Ni2+, Cu2+ and Zn2+ can never form a high spin ion. Based on your answer
what is the maximum number of 3d electrons that will still allow a ligand to create a high spin
ion.
18. For each octahedral configuration below, complete the electron distribution in
accordance with the column heading.
High Spin
Low Spin