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

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?

The Basics

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.

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 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

First transition series

Macrominerals

Microminerals

3d

4d

5d

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

Tetrahedral

Square planar

Octahedral

Fe element No. 26

1s2 2s2 2p6 3s2 3p6 4s2 3d6

Z Z

Z Z

Z

X

X

X

X

XY

Y Y

Y

Y

dxy dyzdxz

d dX2-Y2 Z2

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