Chemical Properties of Minerals II
Basic Coordination Chemistry
Quantum theory
gives us insight into
the electronic
structure of atoms
and allows us to
rationalize the
biological behavior
of minerals
Why we need to know the principles of chemistry in a
minerals course:
Minerals are chemicals that function in a biological setting.
Minerals perform functions that are attuned to their chemical properties
Minerals are ions whose charge is determined by chemical principles
Recognizing that most minerals exist as complexes with proteins and
other organic molecules, their chemistry gives us insight into how
these interactions take place
In constructing life nature drew from a large pool of chemicals in an
attempt to find the ones that best fit the tasks that had to be performed.
Chemistry tells us how these decisions were made.
Electronic Structure is Behind Each of the Following Questions
1. Why do Ca2+, Mg2+, and Zn2+ exist only as +2 ions? Li+, Na+ and K+ as +1 ions?
2. Why are bio-complexes of iron red and potassium and sodium colorless?
3. Why are zinc complexes with proteins stable while sodium complexes fall apart?
4. Why was calcium chosen to become the crystalline component of bone?
5. Why is zinc able to block the absorption of copper in the intestine?
6. Why is arterial blood cherry red while venous blood is a darker red?
7. What makes carbon monoxide gas so deadly?
8. Why are plants green?
9. Why is a dangerous oxygen radical formed when iron reacts with hydrogen peroxide?
10. Will the same happen with zinc and hydrogen peroxide?
The Basics
Insights into the Electronic Structure of Atoms
Energy is being emitted discontinuously
White lt
Z
Ca
20
Ba
56
Fe
26
Li
3
Emission Spectra of Elements
Pauli: electron exists
in two different states
Intrinsic electron spin
Conclusions:
Electrons are arranged in a very specified
manner around the nucleus of atoms
Quantum theory: “the electronic energy of atoms are
quantizied….meaning they can only take on certain
discrete energy values”.
A direct indication of the arrangement of electrons
around a nucleus is ionization energies…the energy
required to remove an electron from a gaseous
neutral atom.
Some electrons are very labile
Some electronic states are
stable
1926 Erwin Schrödinger likened the motion of
electrons around a central nucleus of an atom
as having both a wave and particle character.
The energy associated with the electrons is
quantized or present in discrete energy
packets.
There are 4 quantum numbers that bear
directly on the position of electrons and their
energy:
The principle quantum number n, varies with atomic number
The azimuthal quantum which determines the orbital shape and angular momentum
The magneto quantum number describes orientation of an orbital
The spin quantum number describes electron spin
The following rules apply to orbitals
Rule: Orbitals are designated s, p, d, and f and adhere to the following:
s = spherical, 2 electrons
p = sausage shape extending along x, y, and z axis, 6 electrons
d = 5 degenerate orbitals along and between axes, 10 electrons
f = (not a concern)
Rule: At most, two electrons may occupy an orbital (or suborbital)
and they must be of opposite spin
Rule: s orbitals are spherical, with energy that varies only with distance from the
nucleus. At most 2 electrons may occupy an S orbital.
Rule: p orbitals extended along the major X, Y and Z axis designated px,
py and pz. Each holds 2 electrons, or 6 electrons to occupy the P orbital.
The energy varies with both distance and direction
Rule: d orbitals cover all space both along and between the axes.
Their configuration is that of 5 degenerate (equal energy) and
hold at most 10 electrons
The following rules apply to quantum states or atoms and orbitals
Rule: Quantum states vary with atomic number, i.e., number of electrons
Rule: Atoms with a principle quantum number n = 1 have only a 1s
orbital. Examples are hydrogen and helium.
Rule: Atoms with n = 2 have s and p orbitals
Rule: Atoms with n = 3 have s, p, and d orbitals
Rule: 4s orbitals are at a lower energy level than 3d and fill before 3d
Rule: Atoms with 4s and 3d orbitals when ionizing lose 4s first
Two Major Rules in Chemical Physics that impinge on the behavior of minerals
Hund’s rule: The lowest energy state of an atom is achieved when there is maximum
utilization of the surrounding space by the occupying electrons. Pairing of electrons in an
orbital is recognized as a higher energy state than single electrons of the same spin state
occupying the orbitals. This does not apply to s orbitals.
Pauli exclusion principle: No two electrons in an atomic orbital may share the same
set of quantum numbers. This rule led to the realization that electrons in the same
orbital must be of opposite spins.
Z
2pz
2px
2s
2py
2s
2py
Y
1s
3s
Quant No.
2px
X
2pz
Configuration.
n=1
1s
(K shell)
n=2
2s, 2p
(L shell)
n=3
3s, 3p, 3d (M shell)
Shapes are the same, but
differ in orientation
2p orbitals. At the second quantum level orientation also
becomes a factor in deciding orbital energy. Because there
are 3 orientations existing simultaneously, a p orbital can
hold a maximum of 6 electrons, 2 of opposite spin in each
Iron
At. No. = 26
At. Wt.= 55.85
No. of
occupying
electrons
s=2
p=6
d = 10
f = 14
1s22s22p63s23p64s23d6
Principal Quantum
Number
(n =1, 2, 3)
Subshell
(s,p,d,f)
Argon
[Ar]4s23d6
Element
(At. No.)
Ground-state
configuration
Abbreviated form
Sodium (11)
1s22s22p63s1
[Ne]3s1
Magnesium (12)
1s22s22p63s2
[Ne]3s2
Aluminum (13)
1s22s22p63s23p1
[Ne]3s23p1
Silicon(14)
1s22s22p63s23p2
[Ne]3s23p2
Phosphorus (15)
1s22s22p63s23p3
[Ne]3s23p3
Sulfur (16)
1s22s22p63s23p4
[Ne]3s23p4
Chlorine (17)
1s22s22p63s23p5
[Ne]3s23p5
Argon (18)
1s22s22p63s23p6
[Ne]3s23p6
Caution:
The 4s orbital is actually at a lower
energy level than the 3d. As a result 4s
orbitals will fill before 3d. But, when
ionized, electrons will be lost from the
4s before the 3d
Class Exercise
Atomic numbers of Potassium and Calcium are
19 and 20, respectively. Outer electrons are in the
M shell (n = 3).
Determine the electronic configurations of
potassium and calcium and determine their most
likely ionized form
Solution:
When n = 3, the atom must contain s, p, and d
subshells and 3 energy states. But, recall that
the 4s subshell with 2 electrons is of a lower
energy state than the 3d subshell and will fill
first
Z = 19
Z = 20
2
2
6
2
6
1
K = 1s 2s 2p 3s 3p 3d 4s
Ca = 1s2 2s2 2p6 3s2 3p63d 4s2
The most stable form occurs when both metals lose
their 4s electrons. Thus:
K+ and Ca2+
[Ar]4s1 and [Ar]4s2
Macrominerals
Microminerals
First transition series
3d
4d
5d
Elements in the First Transition Series
Sc
Cr
Mn
Fe
Co
Ni
Cu
3d1 3d2 3d3
3d5
3d5
3d6
3d7
3d8
3d10 3d10
4s2
4s1
4s2
4s2
4s2
4s2
4s1
4s2
+3
+2
+3
+4
+2
+3
+4
+1
+2
+3
+1
+2
+3
+1
+2
+2
+3
Ti
V
4s2 4s2
+4
+3
+4
+5
Bio Essential Metals
Zn
1. Why do Ca2+, Mg2+, and Zn2+ exist only as +2 ions? Li+, Na+ and K+ as +1 ions?
Ca = [Ar] 4s2
Mg = [Ne] 3s2
Zn = [Ar] 3d104s2
Li = [He] 2s1
Na = [Ne] 3s1
K = [Ar]4s1
Octahedral
Square
planar
Tetrahedral
Fe
element No. 26
1s2 2s2 2p6 3s2 3p6 4s2 3d6
Z
Z
Z
Z
X
Y
Y
X
dxy
Z
Y
X
dxz
dyz
X
dX2-Y2
Y
X
Y
dZ2
Ion
Cu+
Zn2+
Coord
Orb No.
d10
4
d10
4
Metal Ion Antagonism
sp3
sp3
tetrahedral
tetrahedral
Cd2+ d10
4
sp3
tetrahedral
Hg2+ d10
2
sp
linear
Cu2+ d9
4
dsp2
square planar
Ag2+
d9
4
dsp2
square planar
Fe2+
d6
6
d2sp3 octahedral
Prediction:
Zn2+ will interfere with Cu+
Cd2+ will interfere with Cu+ and Zn2+
Hg2+ interference will be minimal
Ag2+ will interfere with Cu2+ but not Zn2+