• Mineral º …defined, but generally not fixed, composition…

- a review of basic chemistry
 

• Chemical elements:

- protons, neutrons (nucleus), electrons

Nuclear chemistry:

 - atomic number (Z) = number of protons
- specific for particularly elements (periodic table)
- neutrons @ protons weight, differences in numbers create isotopes
- Atomic weight = sum of protons and neutrons
- written as superscript in front of element symbol

Example: Potassium (Z = 19)
- 40K has 21 neutrons
- 39K has 20 neutrons

Electrons

• uncharged atoms, number of electrons = number of protons
• orbit nucleus in systematic way
 - organized according to energy levels
- energy depends on quantum number (n, l, ml, ms)
- quantum number unique for each element
 

• n = higher energy (similar to shells K, L, M…)

• l refers to shape of region where electron likely to be found (similar to subshell s, p, d, f)

• ml and ms have to do with orbitals within subshells and spin of electrons

-important of magnetic properties
 

• Specific energy corresponds to electron quantum number

• Energy of different subshells overlap

• Electrons fill subshells systematically in order of energy level

• Configuration of electrons:

 - core: all orbital position of any individual shell filled with electrons
 - valence: electrons in shells that are not completely filled
 

Formation of Ions

Ions º excess of deficiency of electrons relative to protons

Anions – net negative charge
Cations – net positive charge

Valence or Oxidation state is the value of the charge on an ion.
 

· Configuration of valence electron controls  whether gain or lose electron

- metals – typically have one or two valence electrons, form cations
- non-metals – typically require few electrons to fill valence shells, form anion

· Systematic pattern for filling valence shells

- 1 – 20 and 31 – 38
- between 20 and 31 the shells fill from internal shells
- elements may have differing numbers of shells filled
 - e.g. Ferrous and Ferric iron
 

· Electronegativity º Propensity of element to gain or lose electron

- Based on arbitrary scale: Li = 1, C = 2.5, F = 4
- Low number lose electrons, high numbers gain electrons
 

Earth Abundances of Elements

· Determine abundance of elements in earth on basis of measurements of large number of rocks

· 8 common elements: O, Si, Al, Fe, Ca, Na, K, and Mg
 - make up most of earth
 - make up most minerals

· These elements are mainly crustal

· Determination of Earth abundances difficult, can’t sample core or mantle

· Estimates made by:
 - mass and density based on geophysics
- compositions of magmas from mantle and mantle xenoliths
- compositions of meteorites
 - models
 
 

Chemical Bonding

Two categories of bonds:
(1) involve valence electrons: ionic, covalent, and metallic
(2) don’t involve valence electons: van der Waals and hydrogen
 

Valence-related bonds:

• common characteristic – bonding fills valence shells
 
 

Ionic bonds
• transfer of electron from one element to another element
• results in filled valence shells
• distance between ions depends on attractive forces (Coulomb law) and repulsive forces (Born repulsion)
• Ions bond in ratio so that positive = negative charges
 - NaCl (Halite), CaF₂ (fluorite)
• Ions act like spheres
 - alternating cations and anions
 - strong bonds
 - brittle because like ions repel, cleavage common

Covalent Bonds

• Form with overlap of orbitals as mechanism for sharing electrons
• Examples: Diamond, Graphite (within sheets)
 
 

Metallic Bonds

• a type of covalent bond
• electrons shared without systematic change in orbitals
 - free to move throughout crystal structure
•Formation:
- low electronegativity (weakly held valence electrons)

Relation between valence dependent bonds

· Most bonds not purely ionic, covalent or metallic
· Amount of bond type depends on electronegativity
- greater difference in Electronegativity = ionic
· Of 8 common elements, only one is anion – O
- Electronegativity of O = 3.5
- Electronegativity of others range 0.8 (K) to 1.8 (Si)
- Ionic bonding range 50% (Si-O) to 80% (K-O)

· Native elements – (e.g. S, Fe, Au, etc.)
- Electronegativity differences are zero
- Bonding intermediate between covalent and metallic
- Low electronegativity values favor metallic
- High electronegativity values favor covalent

· Covalent – ionic
- Statistical: where are the ions located?
 - covalent = equal probability of ion on either ion
 - ionic = infinitely high probability of finding electron on one ion – not really possible

· Covalent – Metallic
- Covalent, ions shared over only 2 atoms
- Metallic, ions shared over all atoms

· Metallic – ionic
 - few shared characteristics
 - only in metal alloys (2 or more metals)

Valence bonds and physical properties

· Ionic and Covalent – little electrical conductance
· Metallic – high conductance

· Solubility – ionic highly soluble

Non-valence bonds

· Occur because of asymmetric charge distributions
 - create electrostatic forces
 - Two types: van der Waals and Hydrogen
 

Hydrogen

Ice example:
· H2O is polar molecule
- O is more electronegative than H
- “claims” more of electron
- net negative charge on O side of molecule

· The asymmetric charges allow solidifying liquid when T < 0
 

van der Waals

Carbon example

· Graphite – carbon bonded in sheets with covalent bonds
- over time electrons evenly distributed
- at a given time, excess electrons on one side of sheet
- creates weak electrostatic attraction
· Physical properties
 - typically soft
 - graphite good lubricant

· Note: solids typically have many different bond types

Atom and Ion Size

· Assume that atoms are spheres
 - clearly simplification, electron distributions not spherical
 - assumption works well for arrangement in solids

· Atoms pack together in regular arrangement

· For solids assume spheres
- effective radius – based on distance between adjacent atoms in solid
- essentially the radius if the atom is sphere

· Bond length – sum of effective radius of two adjacent atoms
- metallic bonds: all same effective radius
- ionic bonds: effective radius different between two atoms
 

X-Ray diffraction
· technique to determine bond lengths
 - metallic bonds: identified effective length
 - ionic bonds: can’t separate effective length for each ion
 - determine lengths on basis of comparison between many bonds
 

· Primary variables for effective ionic radius:

 - oxidation state, i.e. charge on ion
 - coordination number, i.e. number of ions surrounding central ions
 
 

· Oxidation state
 - cations smaller than anions
- positive charge holds electrons closer
 

· Coordination
- Think of solids as large anions surrounding small spaces filled by cations
- Size of space determined by effective radius of anions
- Cation effective radius changes to fill space