Quantum theory gives us insight into
the electronic structure of atoms and allows us to rationalize the
biological behavior of minerals
Minerals are chemicals that function in a biological setting.
Why we need to know the principles of chemistry in a minerals course:
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?
Insights into the Electronic Structure of Atoms
Li
Fe
Ba
Ca
White lt
Emission Spectra of Elements
Z
20
56
26
3
Energy is being emitted discontinuously
Pauli: electron exists in two different states
Intrinsic electron spin
Electrons are arranged in a very specified manner around the nucleus of atoms
Conclusions:
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.
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 numberThe azimuthal quantum which determines the orbital shape and angular momentumThe magneto quantum number describes orientation of an orbital The spin quantum number describes electron spin
The following rules apply to orbitals
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: Orbitals are designated s, p, d, and f and adhere to the following:
s = spherical, 2 electronsp = sausage shape extending along x, y, and z axis, 6 electronsd = 5 degenerate orbitals along and between axes, 10 electronsf = (not a concern)
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: Atoms with a principle quantum number n = 1 have only a 1s orbital. Examples are hydrogen and helium.
Rule: Quantum states vary with atomic number, i.e., number of electrons
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
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.
Two Major Rules in Chemical Physics that impinge on the behavior of minerals
2s
X
Y
Z
1s
2px
2px
2py 2py
2pz
2pz
3s
1s
2s, 2p
n=1
n=2
n=3 3s, 3p, 3d
2s
(K shell)
(L shell)
(M shell)
Quant No. Configuration.
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
Shapes are the same, but differ in orientation
Principal Quantum Number(n =1, 2, 3)
Iron
1s22s22p63s23p64s23d6
No. of occupying electrons
Subshell(s,p,d,f)
s = 2p = 6d = 10f = 14
At. No. = 26At. Wt.= 55.85
Argon
[Ar]4s23d6
Element(At. No.)
Ground-stateconfiguration
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 are19 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
K = 1s 2s 2p 3s 3p 3d 4s
Ca = 1s2 2s2 2p6 3s2 3p63d 4s2
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
The most stable form occurs when both metals lose their 4s electrons. Thus:
K+ and Ca2+
2 2 6 2 6 1
[Ar]4s1 and [Ar]4s2
Z = 19
Z = 20
Elements in the First Transition Series
Sc Ti V Cr Mn Fe Co Ni Cu Zn
3d1 3d2 3d3 3d5 3d5 3d6 3d7 3d8 3d10 3d10
4s2 4s2 4s2 4s1 4s2 4s2 4s2 4s2 4s1 4s2
+3 +4 +3 +3 +2 +2 +1 +1 +1 +2 +4 +3 +3 +2 +2 +2 +5 +4 +4 +3 +3
Bio Essential Metals
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
Cu+ d10 4 sp3 tetrahedral
Coord Ion Orb No.
Zn2+ d10 4 sp3 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+
Metal Ion Antagonism