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

2. = 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.

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

4. 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

5. Hybrid Orbitals - Carbon
Alternatively:
Atom:  |  |     Graphite:  |   | 
1s 2s p s 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)

6. Covalent Chemical Bonds
Graphite Structure

7. 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

8. 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.

9. 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.

10. 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-

11. 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

12. 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

13. 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)

14. 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

15. 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

16. Covalent Bonding of 2 H Atoms  H2 Molecule
Interaction Potential

17. 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

18. 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