Crystal Binding (Bonding) Continued More on Covalent Bonding Part V

= 1s2 2s2 2 p6 3s2 3p5 This is the covalent or shared electron bond. Consider 2 close Cl atoms. Each has electronic shells = 1s2 2s2 2 p6 3s2 3p5 If they move close until their outer orbitals overlap, the atoms can share 2 e- & "fill" the remaining 3p shell of each Cl. The electronic energy there is lowered, which causes the orbitals to stay overlapped; resulting in a strong bond in Cl2 . This is the covalent or shared electron bond.

Covalent Chemical Bonds Hybrid Orbitals - Carbon Atom:  |  |   Diamond: |  |    1s 2s 2p 1s 2(sp3) Diamond C-C-C angle = 109o 28’

Larger Overlap  Stronger Bond Covalent Bonds are Directional The 2(sp3) orbital is tetrahedrally shaped. Larger Overlap  Stronger Bond Covalent Bonds are Directional Each C is tetrahedrally coordinated with 4 others (& each of them with 4 others...) The C-C-C bond angle is fixed at 109o 28' (max. overlap) Note the Face-Centered Cubic lattice The directional character of the bonds,  lower coordination & symmetry, density

Hybrid Orbitals - Carbon Alternatively: Atom:  |  |     Graphite:  |   |  1s 2s 2p 1s 2(sp2) 2p As many know, C is “flexible” in the sense that it can participate in many different kinds of bonding. In fact, many atoms in the center of the Periodic Table with partially filled valence shells are variable in how they bond (this includes Si)

Covalent Chemical Bonds Graphite Structure

Belongs to the Hexagonal Crystal Class The 3 2(sp2) orbitals are coplanar & 120o apart The orbital overlap is similar to that in diamond within the planes (strong too!). Belongs to the Hexagonal Crystal Class Note the p-bonding between the remaining 2p's This results in delocalized e- 's in 2p orbitals which results in electrical conductivity only within sheets. There are other hybrids as well (dsp2 in CuO- planar X) e- may resonate in bonds of non-identical atoms & give a partial ionic character if one much more e-neg than other

Covalent-Network and Molecular Solids Diamonds are an example of a covalent-network solid in which atoms are covalently bonded to each other. They tend to be hard and have high melting points.

Covalent-Network and Molecular Solids Graphite is an example of a molecular solid in which atoms are held together with van der Waals forces. They tend to be softer and have lower melting points.

Covalent Bonds occur between atoms that are “sharing” electrons: Form covalent compounds. There is a “tug of war” for the electrons. There can be single, double & triple covalent bonds: Single bond – a bond in which 2 atoms share a pair of electrons. Double bond – bond that involves 2 shared pairs of e-. Triple bond – bond that involves 3 shared pairs of e-

Combinations of atoms of non-metallic atoms are likely to form covalent bonds Groups 4A, 5A, 6A, and 7A Summarized by G. Lewis in the octet rule sharing of e- occurs if atoms achieve noble gas configuration, H2 is an exception to this rule

Column (Group) Trends Halogens form single covalent bonds in their diatomic molecules (ex: F – F) Chalcogens form double covalent bonds in their diatomic molecules (ex: O = O) Phicogens form triple covalent bonds in their diatomic molecules ( N = N )\ The Carbon group tends to form 4 bonds with other atoms

Double and triple covalent bonds As we just briefly saw for C, covalent bonding can be explained using electron configurations and orbital boxes Double and triple covalent bonds Oxygen forms a double bond in a diatomic molecule It is an exception to the octet rule, 2 unpaired e-. Nitrogen forms a triple bond in a diatomic molecule Satisfies the octet rule, all e- are paired Multiple covalent bonds can form between unlike atoms (ex: CO2, CH3OH)

Molecular Orbitals As we briefly showed, when 2 atoms covalently bond, their atomic orbitals overlap to produce molecular orbitals (orbitals that apply to the entire molecule) The molecular orbital model of bonding requires that the number of molecular orbitals equal the number of overlapping atomic orbitals When 2 atomic orbitals overlap, 2 molecular orbitals are created One is called a bonding orbital, the other is called an anti-bonding orbital The anti-bonding orbital has a higher energy that the atomic orbitals from which it formed

Molecular Orbitals When H2 forms, the 1s atomic orbitals overlap 2 electrons are available for bonding (see next slide) The energy of the e- in the bonding molecular orbital is lower than the e- in the atomic orbitals of the separate H atoms Electrons seek the lowest energy level, so they fill the bonding molecular orbital This makes a stable covalent bond between the H atoms The anti-bonding orbital is empty Sigma and pi bonds are caused by the overlapping of “s” and “p” orbitals

Covalent Bonding of 2 H Atoms  H2 Molecule Interaction Potential

Hybrid Orbitals In orbital hybridization, several atomic orbitals mix to form the same total number of equivalent hybrid orbitals One 2s and three 2p orbitals of a carbon atom overlap to form an sp3 hybrid orbital These are at the tetrahedral angle of 109.5o Four sp3 orbitals of carbon overlap with the 1s orbitals of the four hydrogen atoms This allows for a great deal of overlap, which results in the formation of 4 C-H sigma bonds These are unusually strong covalent bonds

Bond Polarity Covalent bonds involve the sharing of electrons However, they can differ in how the bonds are shared Depends on the kind and number of atoms joined together When electrons are shared equally, a nonpolar covalent bond is formed When the atoms share the electron unequally, a polar covalent bond is formed The more electronegative element will have the stronger electron attraction and will acquire a slightly negative charge The less electronegative element will acquire a slightly positive charge