Crystal Chemistry – chemical composition, internal structure and physical properties
The Atom – protons (+), neutrons (0) and electrons (-)
-
atomic radii are expressed in angstroms (Å)
1Å=10-8cm
smallest atom is H = 0.46 Å; largest atom is Cs = 2.72 Å
atom is electrically neutral (#P = #e’s)
mass of atom is in nucleus (#P + #N), because weight of electron is 1/1837 of a
Proton
e’s move rapidly around the nucleus to give rise to large effective diameters (up to
100,000 the nucleus diameter)
Neils Bohr model
Atomic Number of element (Z) = # P in the nucleus
Atomic Mass of element = #Ps + #Ns in the nucleus
Isotopes – of an element have different masses = same #Ps + different #Ns
e.g. Oxygen
- 18O 8P+
- 17O 8P+
- 16O 8P+
10N
9N
8N
8e8e8e-
A.M. = 18
A.M. = 17
A.M. = 16
0N
1N
2N
1e1e1e-
A.M. = 1 (Hydrogen)
A.M. = 2 (Deuterium)
A.M. = 3 (Tritium)
e.g. Hydrogen
- 1H
- 2H
- 3H
1P+
1P+
1P+
Periodic Table – elements are arranged in order of increasing atomic number (Z) to
show their period repetition of chemical and physical properties.
- the number of Protons in the nucleus (Z) and the number of Electrons, determine
the order
- Vertical Columns = Groups – I, II, III, IV, V, VI, VII, and VIII
- Horizontal Rows = Periods – 1, 2, 3, 4, 5, 6, 7
- Table is subdivided into metals, non-metals and gases.
- Metals – most conductive, metallic luster, ductile
- Nonmetals – brittle, insulators
- Gases – inert, noble
- 2 rows below the rest of the table are the Lathanide Series (Rare Earth Elements –
REE – Z of 58 to 71) and the Actinide Series (Z of 90 to 103)
Atomic Structure – Atoms can be considered as shells (orbitals) in which electrons orbit
the nucleus.
- Shell name
Max # eQuantum #
K
2
1
amount of energy necessary
L
8
2
to bring 1 electron together
M
18
3
with the nucleus
N
32
4
↓
O
50
5
P
72
6
Q
98
7
Each Shell has subshells of increasing energy s, p, d, f, g
max.
s – 2ep – 6ed – 10ef – 14eg – does not get filled
(from Klein, 2002, p. 49)
-
for the most part, each subshell will be filled in order
at higher atomic numbers, orbitals with higher quantum numbers start to fill
before orbitals with lower quantum numbers are completely filled e.g. 4s fills
before 3d
(from Klein, 2002, p. 50)
- f shells are not important for bonding of atoms to form minerals
Example: 1s2 2s2 2p6 3s2 3p1
shorthand – KL3s23p1
- has 2 electrons in each s subshell and 6 electrons in the 2p shell.
In the Periodic Table
- Roman numerals above each column (group) indicate the # of e- in an outer shell
- each column contains elements that have similar chemical behavior
- Horizontal tiers are periods, each having the same outer shell filled, such that the s
and p orbitals are completely filled on the right hand side of the table.
Nobel gases (in group VIII) have p and s orbitals filled and therefore do not gain
or loose electrons
- O and F, and the elements immediately below them, like to gain electrons to fill
orbitals when bonding and become anions: O2-, F- Elements below N will give off or add 3 electrons (+5, -3)
- Elements below H, Be will give off electrons to become cations
- Electrons that are given off or taken are called Valence Electrons
Ionization Potential
- is the amount of energy required to pull off the first, second,…. electron from an
atom
- Ionization potentials increase left-to-right in the periodic table because with
increasing charge in the nucleus, the electrons are more tightly packed around the
nucleus
(from Klein, 2002; Fig. 3.15a)
Electronegativity
-
-
The ability of an atom in a crystal structure or molecule to attract electrons to its
outer shell.
Low = electron donors (<1.9) (Noble gases = 0)
High = electron acceptors (>2.1)
in a specific period, electronegativity values rise as a function of increasing
atomic #
electonegativity values of elements in columns (e.g. Column I: H, Li, Na, K, Rb)
decrease with increasing Z. (Same is true for ionization potential)
(from Klein, 2002; Fig. 3.15b)
Therefore:
-
Bond Strength (binding energy) between the nucleus and the first valence electron
of an element (in a specific group) decreases as the volume of the atom increases.
Large atoms hold their outer valence electrons more loosely than do smaller
atoms
Electonegativity (EN) is useful in assessing bond type
Atoms with similar EN values will form covalent bonds
Chemical Bonding
Ionic Bonding – occurs when the electronegativity of one atom exceeds that of the other
- the more electronegative element will attract an electron into the outer shell
-
the attraction is non-directed (produces higher symmetry in crystals) Na+Cl-
(from Klein, 2002)
Characteristics of Minerals with Ionic Bonds
- e.g Halite
- most common type of bond in minerals
- moderate hardness and specific gravity
- fairly high melting points
- poor conductors of electricity and heat
- non-directional bonding = high symmetry
- cleavage
Covalent Bonding – occurs between atoms with high EN
- atoms share electrons such that orbitals overlap
- produces the strongest chemical bonds but lower symmetry crystals
- e.g. F2, O2 and N3 (A.N. 9, 8, 7)
Characteristics of Minerals with Covalent Bonds
- minerals with covalent bonds include Diamond
- strongest chemical bond
-
generally insoluble
very high melting temperatures
poor to non-conductors
lower symmetry than minerals with ionic bonds
Hybrid Bonding
Covalent bond where electrons are redistributed among subshells
-
e.g. Diamond (ground state of carbon versus diamond)
-
No bonds are totally ionic (e.g. in NaCl the transferred electron will tend to spend
more time between the atoms)
Amount of Ionic character = 1- e-1/4 (Xa-Xb)2 Xa-Xb=difference in electronegativity
-
Metallic Bonding
-
-
bonding may be attributed to attractive forces between nuclei with filled electron
orbitals plus a cloud of loosely bound electrons (these are free to move from atom
to atom, or even out of structure – photoelectric effect).
The weak bonds explain plastic, tenacious, ductile + conductive characteristic of
metal
Natural examples are the native elements – all have isometric symmetry
Van der Waals (or Residual) Bonding
-
very weak residual type bond
ties uncharged or neutral molecules together into a cohesive unit using small
residual charges on their surface
atoms behave as weak dipoles (electrons concentrated on one side of an atom in
order to avoid each other as much as possible)
common in organic compounds and solidified gases, not often seen in minerals
when seen in minerals is produces good cleavage and low hardness
e.g. Graphite – covalent sheets joined by v.d.W. bonds
e.g Sulfur – covalently bonded rings of sulfur that are held together by v.d.W.
bonds , explains why Sulfur melts at 112.8°C and has hardness of 1.5 to 2.5.
Hydrogen Bonding
- polar molecules can form crystalline structures by the attraction between the
oppositely charged ends of molecules.
- Occurs when H is ionically bonded (donates its electron) to a more
electronegative ion such as O2. This causes the proton to become unshielded and
the molecule to develop a dipolar charge
- Hydrogen bonds are weak but stronger than van der Waals bonds
- Hydrogen bonding common in hydroxides, micas and clays
Atomic and Ionic Radii
-
distance between centers of two bonded atoms = bond length
Atomic radii – neutral atoms (Table 3.10, pg. 65)
-
Ionic radii – charged atoms (Table 3.11 pg. 67)
(1)
(2)
(3)
Bond lengths of an ion can vary depending on:
co-ordination number
% of covalent, ionic or metallic character of bond
charge (higher charge = smaller radius)