Dr. Gorring
GEOS 443 Mineralogy
Oct. 29, 1998
Atoms, Ions, and Bond Types in Minerals
1. Atoms- building blocks of minerals, smallest
subdivision of matter that still retain properties of an element.
-
Basic Structure
-
Nucleus containing neutrons (~1 amu, 0 charge) and protons (~1 amu, +)
surrounded by electrons (~0 amu, -); radius ranges from 0.5Å to 2.7Å.
-
Electrically neutral, # of protons = e-. Atomic # (Z) = # of
protons; Mass # (A) = protons (Z) + neutrons (N); isotopes of an element
have same Z, different N (i.e. 12C, 13C, 14C).
-
Atomic weight of an individual isotope = S of
atomic particles binding E; the atomic weight of an element (consisting
all isotopes) = S masses of all the isotopes
of that element weighted by abundance.
-
Bohr Model (1914)
-
Electrons are thought as "orbiting" around nucleus at certain distances,
energy levels, or "shells" designated by quantum numbers (n = 1, 2, 3,
.7) and by letters (K, L, M,
..Q), where n = 1 is the K shell, etc.
-
The energy levels are quantized (have discrete E associated with them)
given by the Einstein Equation. E = hc/l where
h is Planks constant (6.62517 x 1027 erg sec), c is the speed
of light, and l is the wavelength of emitted
light when an e- from a higher E level "drops down" to lower
E level.
-
Schrödinger Model (1926)
-
Bohr Model did not complete explain many observations of the behavior of
e-; (1) e- not only behaved as particles (Bohr),
but also as waves (de Broglie, 1923), (2) exact positions at a given time
were shown to be uncertain (Heisenberg Uncertainty); and (3) did not allow
for the possibility of different shaped e- orbits (i.e. elliptical).
-
Schrödinger Equation (wave eqn) relates the probability of finding
an e- at a given time and place to the mass and E of that e- at that time
and place.
-
Three Quantum Numbers (needed to fully describe atomic structure)
Principal (n): function of distance from nucleus; (K, L, M,
.);
E ~ -1/r2.
Azimuthal (l): determines shape of orbit; l = n-1 (0,1,2
);
s, p, d, f subshells.
Magnetic (ml): orientation of shape in x,y,z space;
determines # of orbitals in each subshells; m = (-l, 0, +l), for example,
when n =2, l = 1, and m = -1,0,1 such that the p subshell shell has 3 possible
orientations in x,y,z, space.
Spin (ms): (+½,-½
); e- can spin clockwise or counterclockwise; produces magnetic
field.
-
Pauli Exclusion Principle: no two e- have
the same quantum #s. Can only have 2 e- (with opposite spin)
in each orbital. Therefore, s subshell hold a max of 2e-; p
holds 6e-; d = 10e-; and f = 14e-
-
The Periodic Chart
-
Arranged by increasing atomic # (Z) and order in which e- fill outer shells.
-
Vertical columns are called Groups; the 8 columns correspond to # of e-
in outer shell (valence electrons).
-
Horizontal rows are called Periods; correspond to K, L, M, N,
. Shells
(principal quantum #s). Rows 4 and 5 have transition elements that fill
inner shells after outermost shells have been filled. For example, in row
4, the 4s shell is filled before the 3d shell. This is because 3d has higher
E than 4s). Lanthanide and Actinide elements are the same but fill inner
f shells, after outer s shells.
2. Ions
-
Elements may give up or borrow electrons, producing ions. Cations = loss
of electrons (X - ne- = Xn+); Anions = gain of e-
(i.e. X + ne- = Xn-). Ionic charge or valence
= # of protons - # of electrons.
-
Cations are on left side of Periodic Chart (metals); Group 1 = +1 valence;
Group 2 = +2. Anions on the right side of periodic chart (non-metals),
halogens, inert gases.
-
Ionization Potential (measured in electron volts, eV): measure
of energy needed to ionize. Metals are easily ionized and have low ionization
potential. Inert gases have the highest IP. Some elements can exist in
several oxidation states (i.e. Fe2+, Fe3+), lose
e- from the 4s and/or 3d orbitals.
-
Electronegativity (unitless #)- measure of ability of an atom to attract
e- to outer shell. Calculated from known bond strengths between
atoms. Elements with low EN are e- donors (Group 1 metals; Na,
K); high EN elements are e- acceptors (Group VII halogens; Cl,
F. Used to predict bond types between different atoms. Atoms with large
difference in EN form ionic bonds (Na,Cl); atoms with similar EN form covalent
bonds (Si,O).
3. Bond Types in Crystals
Bonds are electrical forces that bind atoms in crystalline solids.
Largely control physical properties (hardness, cleavage, melting T, and
conductivity). Four major types. Most minerals in fact are combinations
of the four types. Silicates are mostly ionic/covalent with some having
weak van der Waals/hydrogen bonds as well. For example, MICAS: within sheets
bonding is 50-50 ionic/covalent; bonds between sheets are mostly weak ionic
and/or van der Waals.
-
Ionic: One or more electrons are transferred
to the outer shell of another element so they both have filled outer shells
(inert configuration). Halite is Na+, Cl-; have electrical attraction that
produces IONIC BOND. Physical properties of minerals with ionic bonds:
soluble, moderate hardness and G, fairly high melting T, poor conductors
of heat and electricity. Strength of ionic bonding is proportional to 1/r
and q1q2, where r is interatomic distance between ions and q1, q2 is the
valence of the two ions.
-
Covalent (strongest chemical bond):
One or more electrons are shared in outer shells of two atoms to produce
the inert electronic configuration. Physical properties of minerals with
covalent bonds: insoluble, stable, moderate to very high hardness, very
high melting T, poor electrical conductors.
-
Metallic: Forces that arise from atoms that share
electrons that from a cloud of e- around the atoms. No e-
has affinity to any one atom and is free to move around in structure. Physical
properties of minerals with metallic bonds: soft, high plasticity, maleable,
very conductive.
-
van de Waals: Polarization of atoms causes a weak dipolar
attraction. Can bond together neutral molecules and essentially uncharged
structural units by developing small electrostatic forces on surfaces.
Defines zones of cleavage, and low hardness in some minerals (clays, micas).