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+ 10N 8e- A.M. = 18- 17O 8P+ 9N 8e- A.M. = 17- 16O 8P+ 8N 8e- A.M. = 16

e.g. Hydrogen

- 1H 1P+ 0N 1e- A.M. = 1 (Hydrogen)- 2H 1P+ 1N 1e- A.M. = 2 (Deuterium)- 3H 1P+ 2N 1e- A.M. = 3 (Tritium)

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 # e- Quantum #K 2 1 amount of energy necessaryL 8 2 to bring 1 electron togetherM 18 3 with the nucleusN 32 4 ↓O 50 5P 72 6Q 98 7

Each Shell has subshells of increasing energy s, p, d, f, g max.s – 2e-

p – 6e-

d – 10e-

f – 14e-

g – 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)

- Bond lengths of an ion can vary depending on:(1) co-ordination number(2) % of covalent, ionic or metallic character of bond(3) charge (higher charge = smaller radius)