1. Chemistry Review Atom Nucleus = protons + neutrons
(~ all mass of the atoms)
p + n (variable)  atomic weight (isotopes)
Atomic weight number is the average atomic weight and includes the number of isotopes of an element.
An isotope of an element has the same number of protons (positive charged) but a different number of neutrons (neutral).
Number of protons and electrons (negative charged) define the atomic number of an element Z.

2. Composition of the Earth’s Crust
Most common silicates are from these
O alone = 94 vol. % of crust
Perhaps good to think of crust as a packed O array with interspersed metal cations in the interstices!
Analogy works for minerals too (they make up the crust)

3. Chemistry Review Electron cloud around the nucleus
give its statistical size
Atomic radii are in the range Å
1 Å = m
The SI unit is the nanometre = 10-9 m K L M
Bohr atomic model: electrons travel along specific orbits
with fixed energy levels: K,L,M,N shells

4. Chemistry Review Quantum mechanics, Erwin Schrödinger
Occurrence of electrons are expressed in wave equations ().
The position of electron in space can be specified by four quantum numbers:
Principal quantum number: n
= reflects effective volumes of electron orbits and their energy levels (in Bohr model = shells)
2. Azimuthal quantum number l (orbital shapes); Bohr model = subshells; for each shell there are l = n - 1 subshells
3. Magnetic quantum number, m
4. Electron spin s, spin of electrons only in two directions,
+1/2, -1/2

5. Chemistry Review Quantum numbers / shells
Principal quantum No. n Azimuthal quantum No. l
shells subshells or orbitals
innermost K (n = 1) 0 = s
lower energy L (n = 2) 0,1 = s, p
M (n = 3) 0,1,2 = s, p, d
outermost N (n = 4) 0,1,2,3 = s, p, d, f
higher energy

6. s-orbitals, max. 2 electrons
p orbitals, each dumbbell
= 2 electrons, max. 6 electrons

7. d orbitals, each dumbbell 1 electron = max. 10 electronsz x y d xz z y x d xy x y z d yz z y x d z2 y x z d
x2-y2

8. Quantized energy levels
Chemistry Review
f orbitals =
max. 14 electrons
Quantized energy levels s p d f
Relative Energy
Note that the energy does not necessarily increase K  L  M  N etc.
4s < 3d
n = 1 K L 3 M N O P Q

9. Progressive filling of orbitals as energy increases
Again, energy does not increase regularly K  L  M  N etc.
Some complications with spin V  Cr, etc.

10. Pauli exclusion principle:
There are no two electrons in any one atom,
which have all four quantum numbers the same.
Hund’s rule:
The number of unpaired electrons
in a given energy level is a maximum
Electron configuration of N     
1s2, 2s2, 2p3

11. Al = 13 electrons: 1s2 2s2 2p6 3s2 3p1
Fe = 26 electrons: 1s2 2s2 2p6 3s2 3p6 4s2 3d6

12. It is the outermost shell or valence electrons that are fundamental
Elements are part of Groups in the Periodic Table.
Groups are arranged in vertical directions
Elements of the same group = same number of valence electrons.
In a chemical bond, an element tries to achieve a stable electron configuration, which can be (but must not be) the inert noble gas configuration.

13. Chemical properties of elements

14. Most common cations are Na+, K+, Ca2+, Mg2+, Fe2+, Fe3+, Al3+, Si4+
Compatibility with these elements requires:
Valence state difference of maximum 1 (only when charge balance can be provided)
Similar ionic size
Electron configuration which allows similar
coordination environment
Highly compatible are the cations
V3+, Cr3+, Mn2+, Co2+,Co3+, Ni2+, Ni3+, Zn2+, Sr2+
Less compatible: Ti4+,Li+, B3+, Cu2+, Ba2+, Rb+, Cd2+

15. Electronegativity is the ability of an atom in a crystal structure to attract electrons into its outer shell
In general, electronegativity increases in a period from left to right and in a group from bottom to top.
(except for inert gases which are very low)

16. Elements are classified as:
Metals w/ e-neg < 1.9 thus lose e- and  cations
Nonmetals > 2.1 thus gain e- and  anions
Metalloids intermediate (B, Si, Ge, As, Sb, Te, Po..)

17. Occurrence of Minerals
Rock-forming minerals:
Minerals that form major components of common rocks; e.g. quartz, feldspar, calcite.
Accessory minerals:
Minerals that usually occur in minor amounts in rocks; e.g. zircon, sphene
Ore minerals:
Minerals that ore sources of metals, even hypothetically; e.g. sphalerite→ zinc, galena→ lead, pyrite→ iron, sulfur.
Industrial minerals:
Minerals used more or less for their own properties; e.g barite, graphite, gypsum
Gemstones:
Minerals with sufficient beauty to be used for ornamental purposes; e.g. diamond, ruby (corundum), amethyst (quartz)

18. How the elements occur in minerals
Major elements – form essential portions of abundant minerals; e.g. silicon, oxygen, potassium.
Mineral elements – occur usually in a few weight percent in abundant minerals; e.g. titanium, manganese, nickel.
Trace elements – occur at the ppm level in abundant minerals; e.g. strontium, rubidium, gallium

19. silicate
iron
sulfide

20. Goldschmidt’s Classification of the elements
V.M. Goldschmidt, noting that the elements were concentrated in different portions of the mineral components of iron meteorites, devised a classification based on these components where the elements were concentrated. The idea is based on the idea that iron meteorites may represent planetary cores and are relatively primitive representatives of the early formation of the Earth.

21. The classification is four-fold:
Lithophile elements that concentrate in silicates;
e.g. lithium, beryllium, aluminum
2. Siderophile elements that concentrate in iron;
e.g. platinum, iridium, nickel
3. Chalcophile elements that concentrate in sulfides;
e.g. copper, zinc, arsenic
4. Atmophile elements that concentrate in gas; e.g.hydrogen, nitrogen, helium