Chemical Properties of Minerals II

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. 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.

4. 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?

5. The Basics

6. Pauli: electron exists in two different states Insights into the Electronic Structure of Atoms Energy is being emitted discontinuously White lt Z Ca 20 Ba 56 26 Fe 3 Li Emission Spectra of Elements Intrinsic electron spin

7. 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.

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: 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

11. 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

12. 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.

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

15. 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

17. Subshell (s,p,d,f) Iron No. of occupying electrons s = 2 p = 6 d = 10 f = 14 At. No. = 26 At. Wt.= 55.85 1s22s22p63s23p64s23d6 Principal Quantum Number (n =1, 2, 3) Argon [Ar]4s23d6

18. Ground-state configuration Element (At. No.) 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

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. 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 K = 1s 2s2p3s3p3d 4s 2 2 6 2 6 1 Z = 19 Ca = 1s22s22p63s23p63d 4s2 Z = 20 The most stable form occurs when both metals lose their 4s electrons. Thus: K+ and Ca2+ [Ar]4s1 and [Ar]4s2

22. Macrominerals Microminerals First transition series 3d 4d 5d

23. Elements in the First Transition Series Sc Ti V Cr Mn Fe Co Ni Cu Zn 3d1 3d2 3d3 3d5 3d5 3d6 3d7 3d83d10 3d10 4s2 4s2 4s2 4s1 4s24s2 4s2 4s24s1 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

25. 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

26. Octahedral Square planar Tetrahedral

28. Fe element No. 26 1s2 2s2 2p6 3s2 3p6 4s2 3d6 Z Z Z Z Z Y Y X X X Y dxz dxy dyz X Y X d Y d Z2 X2-Y2

29. Metal Ion Antagonism Coord Ion Orb No. Cu+d10 4 sp3 tetrahedral 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+