Transcript Mineralogy Lecture 07

Introduction to Mineralogy
Dr. Tark Hamilton
Chapter 3: Lecture 7
The Chemical Basis of Minerals
(sizes, shapes & directions)
Camosun College GEOS 250
Lectures: 9:30-10:20 M T Th F300
Lab: 9:30-12:20 W F300
Controls of Size in Atoms & Ions
• In lone atoms (few and far between):
– Electrons are attracted by Z = Atomic #
•
92U
1s
< 1H1s
– Electrons in same shell repel each other
•
9F
> 3Li
– Heavier atoms/ions are bigger down groups
•
56Ba
> 12Mg
• +/- charges of nearest neighbours count
– Al+3 is smaller in (AlO4)-5 than with (AlO6)-9
Ice - Covalent Bonded
Mineral & Rock: H2O
H-O-H bond angle
104.5°
(Tetrahedral = 109.5°
1.8 Angstrom
Weak Hydrogen bonds
between waters
1 Angstrom
Strong molecular bonds
within waters
Open Channels
(low density)
Low O
High O
There are many Higher Pressure Polymorphs!
Each Oxygen is
Tetrahedrally coordinated
by 4 Hydrogens
Sub-Atomic Physics
for
Dummies & Mineralogists
The Bohr Hydrogen Atom
Better than
Rutherford’s
Plum Pudding
Model
Explains H emission lines
Classical orbits decay
Elements & Isotopes
Carbon
Atomic # 6
At.wt. 12.011
Nitrogen
Atomic # 7
At.wt. 14.0067
Oxygen
Atomic # 8
At. wt. 15.9994
Stable Isotopes
Geologically Common & Important
Elements in Minerals
Major Elements: Al, Ar, Ca, C, H, K, Fe,
Mg, Mn, N, Na, O, P, S, Si, Ti
(16)
Minor Elements: As, B, Ba, Be, Bi, Br, Ce,
Cd, Cs, Cl, Cr, Co, Cu, F, He, La, Pb, Hg, Mo,
Nb, Ni, Ra, Rn, Sc, Se, Sr, Sn, W, V, Y, Zn, Zr
(32)
Trace Elements: Ag, Au, Dy, Eu, Ga, Gd,
Hf, Ho, I, Ir, Li, Lu, Nd, Os, Pd, Pt, Rb, Rh,
Sm, Ta, Tb, Te, Tl, Th, Tm
(25)
Electrons Locations as
Schroedinger’s Wave Functions
S Orbital
S Orbital Radius
Subshell Orbital Geometries
P2 x y z Orbitals
S1,2,3 Orbitals
D3 Orbitals
Orbital Quantum Numbers
& Order of Filling Sub-shells
Spectral Line Names:
s – Sharp
p – Principal
d – Diffuse
f - Fundamental
Quantum Numbers for Dummies &
Mineralogists
n Principal Quantum # = size of atom
& row in periodic table
K, L, M: important shells for
X-ray generation & identification
s, p, d, f – Electron energy levels
orbital shapes & directions
Every electron in every element
has a unique set of quantum numbers
& an energy level all its own.
Electron jumps = Δ energy, photons
Quantum Numbers = Electron E’s
Magnetic Spin Quantum #
Stern-Gerlach Experiment
Spin +/- ½
Deflected Ag beam
by strong
Magnetic field
Gave 2 spots
Half filled 5s1
Magnetic Field & Induction
Pauli Exclusion Principle:
No 2 electrons may have identical values
for all quantum numbers
Schroedinger’s Quantum Numbers
• n – principal shell (distance from nucleus)
• l – azimuthal, shape of orbital
• ml – magnetic, direction/orientation of orbital
Spin Quantum Number (empirical)
• ms – spin magnetic
• Pairs of electrons in subshells: s, p, d, f differ in
the sign of their spin magnetic #
Elements and Ions
Low Ionization potential forms Ions easily
Electronegativity attracts electrons
to atom in covalent bonds
Lone electrons ionize easily
Noble gases need X-rays to ionize
Geeze-Louise! This looks like the
list of major & minor elements!