1. Chemical Properties of Minerals II Basic Coordination Chemistry

2. Quantum theory gives us insight into the electronic structure of atoms and allows us to rationalize the biological behavior of minerals

3. 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.

4. Electronic Structure is Behind Each of the Following Questions 1. Why do Ca 2+, Mg 2+, and Zn 2+ 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?

5. The Basics

6. 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

7. 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.

8. Some electrons are very labile Some electronic states are stable

9. 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

10. 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 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: 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

11. 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

12. 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

13. 2s X Y Z 1s 2px 2py 2pz 3s 1s 2s, 2p n=1 n=2 n=3 3s, 3p, 3d 2s (K shell) (L shell) (M shell) Quant No. Configuration.

15. 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

17. Principal Quantum Number (n =1, 2, 3) Iron 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 No. of occupying electrons Subshell (s,p,d,f) s = 2 p = 6 d = 10 f = 14 At. No. = 26 At. Wt.= 55.85 Argon [Ar]4s 2 3d 6

18. Element (At. No.) Ground-state configuration Abbreviated form Sodium (11)1s 2 2s 2 2p 6 3s 1 [Ne]3s 1 Magnesium (12)1s 2 2s 2 2p 6 3s 2 [Ne]3s 2 Aluminum (13)1s 2 2s 2 2p 6 3s 2 3p 1 [Ne]3s 2 3p 1 Silicon(14)1s 2 2s 2 2p 6 3s 2 3p 2 [Ne]3s 2 3p 2 Phosphorus(15)1s 2 2s 2 2p 6 3s 2 3p 3 [Ne]3s 2 3p 3 Sulfur (16)1s 2 2s 2 2p 6 3s 2 3p 4 [Ne]3s 2 3p 4 Chlorine (17)1s 2 2s 2 2p 6 3s 2 3p 5 [Ne]3s 2 3p 5 Argon(18)1s 2 2s 2 2p 6 3s 2 3p 6 [Ne]3s 2 3p 6

19. 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

20. 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

21. K = 1s 2s 2p 3s 3p 3d 4s Ca = 1s 2 2s 2 2p 6 3s 2 3p 6 3d 4s 2 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 Ca 2+ 226261 [Ar]4s 1 and [Ar]4s 2 Z = 19 Z = 20

22. First transition series Macrominerals Microminerals 3d3d 4d4d 5d5d

23. Elements in the First Transition Series Sc Ti V Cr Mn Fe Co Ni Cu Zn 3d 1 3d 2 3d 3 3d 5 3d 5 3d 6 3d 7 3d 8 3d 10 3d 10 4s 2 4s 2 4s 2 4s 1 4s 2 +3 +4 +3 +3 +2 +2 +1 +1 +1 +2 +4 +3 +3 +2 +2 +2 +5 +4 +4 +3 +3 Bio Essential Metals

25. 1. Why do Ca 2+, Mg 2+, and Zn 2+ exist only as +2 ions? Li +, Na + and K + as +1 ions? Ca = [Ar] 4s 2 Mg = [Ne] 3s 2 Zn = [Ar] 3d 10 4s 2 Li = [He] 2s 1 Na = [Ne] 3s 1 K = [Ar]4s 1

26. Tetrahedral Square planar Octahedral

28. Fe element No. 26 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 Z Z ZZ Z X X X X X Y YY Y Y d xy d yz d xz d d X 2 -Y 2 Z2Z2

29. Cu + d 10 4sp 3 tetrahedral Coord Ion Orb No. Zn 2+ d 10 4sp 3 tetrahedral Cd 2+ d 10 4sp 3 tetrahedral Hg 2+ d 10 2splinear Cu 2+ d 9 4dsp 2 square planar Ag 2+ d 9 4dsp 2 square planar Fe 2+ d 6 6d 2 sp 3 octahedral Prediction:Zn 2+ will interfere with Cu + Cd 2+ will interfere with Cu + and Zn 2+ Hg 2+ interference will be minimal Ag 2+ will interfere with Cu 2+ but not Zn 2+ Metal Ion Antagonism