Atomic Bonding in Solids

• Lecture 5 discusses interatomic and atomic bonding in solids

Recap - Atomic Bonding in Solids

• Interatomic forces (between atoms) binds atoms together
• Negligible interaction occurs if the distance between atoms is large
• When atoms come closer, each atom exerts forces on the other atoms
• These forces are attractive and repulsive and developed between them
• Magnitude of each force is a function of interatomic distance
• FA depends on the bonding between the two atoms
• FR comes into play causing outer shells to overlap as r becomes smaller
• State of equilibrium is achieved at r = ro and FN = 0

Covalent Bonds

• Similar electronegativities share electrons in covalent bonds
• Bonds involve valence electrons – normally s and p orbitals
• H2 is an example of a covalent bond
• Each Hydrogen atom has 1 valence electron and needs 1 more
• Electronegativities are the same in Hydrogen
• Covalent bonding is directionally sensitive
• Atoms must come at a specific angle

Bond Hybridization

• Carbon can form sp³ hybrid orbitals from overlapping s and p orbitals

Covalent Bonding: Carbon sp³

• CH4 is an example of a covalent bond with carbon
• Carbon has 4 valence electrons and needs 4 more
• Hydrogen has 1 valence electrons and needs 1 more
• Electronegativities are comparable for Carbon and Hydrogen
• Electrons are shared in covalent bonds for Caron-Hydrogen

Examples of Covalently Bonded Molecules

• CH4, H2O, HNO3 and HF are covalently bonded
• Covalent bonding is found in elemental solids
• Diamond (carbon), silicon, and germanium
• Solid compounds composed of elements on the right-hand side of the periodic table are covalently bonded
• Gallium arsenide (GaAs), indium antimonide (InSb), and silicon carbide (SiC) are covalently bonded

Covalent Bonding

• The number of covalent bonds possible for a particular atom is determined by the number of valence electrons
• For N' valence electrons, an atom can covalently bond with at most 8 - N' other atoms
• Chlorine: N' = 7 for chlorine and 8 - N' = 1, so one Chlorine atom can bond to only one other atom, as in Cl2
• Carbon: N'=4, so each carbon atom has 8 - 4 = 4 electrons to share
• Diamond is a three-dimensional interconnecting structure where each carbon atom covalently bonds with four other carbon atoms

Covalent Bonding Strength

• Covalent bonds can be very strong (ex. diamond), which are very hard and have melting temps greater than > 3500 degrees Celsius
• Covalent bonds can also be weak, such as bismuth that melts at ~270 degrees Celsius
• Polymeric materials typify this bond, the basic molecular structure being a long chain of carbon atoms that are covalently bonded together with two of their available electrons

Ionic-Covalent Mixed Bonding

• It is possible to have interatomic bonds that are partially ionic and partially covalent
• Very few compounds exhibit pure ionic or covalent bonding
• The degree of either bond type depends on the relative positions in the periodic table and the difference in their electronegativities
• The wider the separation from the lower left to the upper-right-hand corner, the more ionic the bond
• Conversely, the closer the atoms are together, the greater the degree of covalency
• % ionic character = (1-e^(-(XA-XB)^2)/4 ) x(100%)
• where XA & XB are Pauling electronegativities
• Example: MgO has a ionic character of 70.2%
• XMg = 1.3
• XO = 3.5

Metallic Bonding

• Metallic bonds are found in metals and their alloys
• Valence electrons are not bound to any 1 atom in the solid and are free to drift throughout the entire metal to form a "sea of electrons" or an "electron cloud"
• Metallic bond is electrostatic attraction between +ve ion cores and delocalized electrons
• Remaining non-valence electrons and atomic nuclei form ion cores
• Ion cores possesse a net positive charge equal in magnitude to the total valence electron charge per atom
• Free electrons act as a glue to hold the ion cores together
• Bonding strength varies
• Energies range from 68 kJ/mol (0.7 eV/atom) for mercury to 850 kJ/mol (8.8 eV/atom) for tungsten
• Respective melting temperatures vary
• Some general behaviors of the various material types (i.e., metals, ceramics, polymers) may be explained by bonding type

Secondary Bonding (weak bonds)

• Due to interaction between dipoles due to asymmetry in electron cloud distribution
• Fluctuating dipoles contribute to secondary bonding forces
• Permanent dipoles-molecule contribute to secondary bonds
• Van der Waals bonding occurs
• Physical bonds are weak in comparison to the primary or chemical bonds
• Bonding energies typically on the order of only 10 kJ/mol (0.1 eV/atom)
• Occurs between virtually all atoms or molecules
• The presence of van der waals forces may be obscured if any of the three primary bonding types is present
• Forces arise from atomic or molecular dipoles
• Bonding results from the coulombic attraction between the positive end of one dipole and the negative region of an adjacent one
• Hydrogen bonding, a special type of secondary bonding, is found to exist between some molecules that have hydrogen as one of the constituents

Properties From Bonding

• Bond length, r, defines distance between atoms
• Bond energy, Eo defines size of the "bond energy"
• Melting temperature, Tm, increases if Eo is larger
• Coefficient of thermal expansion, α and α is larger if E is smaller

Summary - Bonding Types

• Ionic - Large bond energy, Nondirectional (ceramics)
• Covalent - Variable large-Diamond small-Bismuth, Directional (semiconductors, ceramics polymer chains)
• Metallic - Variable large-Tungsten small-Mercury, Nondirectional (metals)
• Secondary - smallest bond energy, Directional inter-chain (polymer) inter-molecular

Summary - Primary Bonds

• Ceramics: Large bond energy (Ionic & covalent bonding), large Tm, large E, small α
• Metals: Variable bond energy (Metallic bonding), moderate Tm, moderate E, moderate α
• Polymers: Directional Properties, Secondary bonding dominates (Covalent & Secondary), small Tm, small E, large α