Lecture 3
Crystal Chemistry Part 2:
Bonding and Ionic Radii
Salt, Calcite and Graphite models
Chemical Bonding in Minerals
• Bonding forces related to electrically
charged particles – negative attracts positive
• Bond strength controls most physical and
chemical properties of minerals
•In general, the stronger the bond, the harder
the crystal, higher the melting point, and
the lower the coefficient of thermal expansion
Chemical Bonding in Minerals
•Five general types bonding types:
Ionic
Covalent
Metallic
van der Waals
Hydrogen
Commonly different bond types occur in
the same mineral
Chemical Bonds
Electrical in nature- responsible for most mineral properties
1) Ionic
Na: low 1st Ionization Potential 1s2 2s2 2p6 3s1
Na  e- + Na+ (Sodium ion has a Neon configuration) 1s2 2s2 2p6
Cl: high e-neg takes the e-  Cl- (Cl- ion has Argon configuration)
Now they have opposite charges & attract = bond
Bonding is strong (e.g. Salt has high melting point)
But easily disrupted by polarized solvents (e.g. water)
Poor electrical conductors; electron strongly held by anion
Strength  (1/bond length) & valence
Also non-directional so symmetrical packing is possible
(Isometric crystal system is common in Alkali Metal – Halogen Salts).
If electronegativity of anion and cation differs by 2.0 or more
will be mostly ionic , say about 70%.
Halite (NaCl)- An Example of Ionic
Bonding
Na fits into interstices
+
Na+ lost an electron shell, smaller; Cl- gained an electron, repels nucleus, larger
Ionic
Bonding
Example: NaCl
Na (1s22s22p63s1) –> Na+(1s22s22p6) + eCl (1s22s22p63s23p5) + e- –> Cl- (1s22s22p63s23p6)
Problem 1

Write down the electron configuration for
neutral Chlorine Cl and for Chloride Ion Clusing the info from lecture 2.
Chemical Bonds
2) Covalent
Consider 2 close Cl atoms, each = 1s2 2s2 2 p6 3s2 3p5
If draw closer until overlap an outer orbital, can
share whereby 2 e- "fill" the remaining 3p shell of
each Cl
Low energy condition causes electrons to stay
overlapped; results in a strong bond  Cl2
This is the covalent or shared electron bond
Usually stronger than Ionic bond
Covalent bonding – sharing of
valence electrons
Cl:1s2 2s2 2p6 3s2 3p5 so 7 electrons in outer shell
“The sharing of an electron pair … constitutes a single bond” S&P p54.
Chemical Bonds
3) Metallic Bonding
Metals have few, loosely held valence electrons
If closely pack them can get up to 12 nearest neighbors
This causes a high density of valence e- around any given
atom & also a high density of neighbor atoms around the
loose valence eThese become a sea of mobile electrons
Metals are excellent conductors
Chemical Bonds
4) Van der Waals Bonds
Weakest bond – due localized excess charge
Usually between neutral molecules (even large
ones like graphite sheets)
Weakness of the bond is
apparent in graphite cleavage
Caused by momentary correlations in the
charge polarity of adjacent atoms
More Detail

Now let’s look at the bond types in more
detail
Ionic Bonds Dominate Most Mineral Geometry
Most minerals
have a strong
ionic
component.
 Mostly covalent
Ion complexes
SiO4 -4, CO3 --,
etc. are ionically
bonded to metal
ions to achieve
neutrality.

Calcite CaCO3
Ionic Bond Properties



Results in minerals displaying moderate degrees
of hardness and specific gravity, moderately
high melting points, high degrees of symmetry
Poor conductors
Strength of ionic bonds are related to:
1) the spacing between ions
2) the charge of the ions
Stronger bond has a higher melting point
Compound Bond Strength = Melting Point
vs. interionic distance, ionic charge
9
17
35
53
Sodium Na+ with various anions
Small inter-ionic distance = higher melting point
12
20
38
56
+2 cations
A (ångström) = 10 -10m
Small inter-ionic distance = higher melting point
3
11
19
37
+1 cations
Li F is an exception
Interionic Distance vs. Hardness
4
12
20
38 56
22
21
12
11
Closer Interionic Distance =
Increased Bond Strength
(Hardness)
Covalent Bonding







formed by sharing of outer shell
electrons
strongest of all chemical bonds
most covalent minerals are
insoluble in acids
high melting points,
hard, nonconductive
have low symmetry due to
multi-directional bonding.
common among elements with
high numbers of vacancies in
the outer shell (e.g. C, Si, Al, S)
Diamond
Tendencies for Ionic vs. Covalent Pairing
Ionic Pairs
Covalent
Pairs
Si-O, C-O, S-O, N-O, P-O
Covalent-Ionic continuum
Difference in electronegativity of the
elements involved tells us if one member
is more attractive to electrons i.e. forms
ionic bonds. F to Na 4.1 – 1 = 3.1, very
different, so Na-F bond very ionic in
character.
Si-O difference 3.5-1.8 = 1.7 ~ 50% covalent
Covalent
Ionic
Metallic Bonding
Atomic nuclei and inner filled electron
shells in a “sea” of electrons made up of
unbound valence electrons.
 Typical of elements with low ionization
potential. Valence electrons easily
stripped.
 Yields minerals with minerals that are soft,
ductile/malleable, highly conductive (due
to easily mobile electrons).
 Non-directional bonding produces high
symmetry

Van der Waals (Residual) Bonding



created by weak bonding of oppositely
depolarized electron clouds
commonly occurs around covalently bonded
elements
produces solids that are soft, very poor
conductors, have low melting points, with low
symmetry crystals and strong cleavage.
Hydrogen Bonding example ICE
Electrostatic
bonding between an
H+ ion with an anion
or anionic complex
or with a polarized
molecules
Weaker than ionic
or covalent;
stronger than
Van der Waals
H+
Close packing of
polarized molecules
polarized H2O
molecule
Anions
Ice
One Hydrogen bond shown as red line above
Summary of Bonding Characteristics
Crystal Chemistry
Crystals can be classified into 4 types:
1. Molecular Crystals
Neutral molecules held together by weak van der Waals
bonds
Rare as minerals
Mostly organic
Weak and readily
decompose, melt,
cleave, etc.
Example: graphite
Crystal Chemistry
2. Covalent Crystals
Atoms of similar high e-neg and toward right side of
Periodic Table
Also uncommon as minerals (but less so than molecular)
Network of strong covalent
bonds with no weak links
Directional bonds  low
symmetry and density
Example: diamond
Crystal Chemistryhard-sphere model
The diamond structure
All carbon atoms in IV coordination
FCC unit cell
ball-and-stick model
polyhedral model
blue C only
Crystal Chemistry
3. Metallic Crystals
Atoms of similar e-neg and toward left side of Periodic Table
Metallic bonds are directionless bonds  high
symmetry and density
Pure metals have same sized atoms
Closest packing  12 nearest mutually-touching neighbors
Cubic Closest Packing (CCP) abcabcabc stacking = FCC
cell (face-centered cubic AKA cubic close packed)
Hexagonal Closest Packing (HCP) ababab = hexagonal
cell
Also BCC in metals, but this is not Closest Packing
More on coordination and closest packing next time
Crystal Chemistry
4. Ionic Crystals
Most minerals
First approximation:
 Closest-packed array of oxygen atoms
 Cations fit into interstices between oxygens,
balance the negative charges. Negative charges
mostly due to oxide ions O Different types of interstitial sites available
 Cations occupy only certain sites where can fit
 Only enough cations to attain electrical
neutrality
Multiple Bonding in Minerals

Graphite – covalently bonded
sheets of C loosely bound by
Van der Waals bonds.

Mica – strongly bonded silica
tetrahedra sheets (mixed
covalent and ionic) bound by
weak ionic and hydrogen
bonds

Calcite: Cleavage planes
commonly correlate to planes
of weak ionic bonding versus
strong covalent bonds
in CO3--
