1. Ch. 8 Covalent Bonding

2. 8.1 Molecular Compounds

3. Sharing Electrons
Atoms held together by covalent bonds (sharing electrons)
Molecules are neutral groups of atoms
Diatomic molecules contain only two atoms

4. Representing Molecules
Molecular Formula: chemical formula of a molecule, shows the number of each type of element in the molecule
Structural Formula: shows the bonds between atoms
Space-filling Model: shows the arrangement between atoms
Perspective Drawing: shows the arrangement between atoms
Ball and Stick Molecular Model: shows the arrangement between atoms

5. Representing Molecules
What is molecular formula? Structural formula?

6. Comparing Molecular and Ionic Compounds
Representative units
Molecules (I.e. C6H12O6 or C3H6O)
Number of atoms of each element
Composed of nonmetals
Formula units (I.e. NaCl or MgCl2)
Ratio of ions of each element
Composed of a metal cation and a nonmetal anion

7. 8.2 The Nature of Covalent Bonding

8. Octet Rule
Electrons are shared (not lost or gained) to achieve the electron configuration of a noble gas (full octet unless H or He)
Single bond: one pair of electrons shared
Double bond: two pairs of electrons shared
Triple bond: three pairs of electrons shared
In a structural formula, shared pairs of electrons are signified by dashes

9. Octet Rule Unshared pairs of electrons are left as dots
Also called “lone pairs”

10. Drawing Dot Structures for Molecules
H2O H2
HCl F2 O2

11. Coordinate Covalent Bonds
Coordinate covalent bond: one atom contributes both bonding electrons
NH3  NH4+
Often form polyatomic ions
Polyatomic ions: group of atoms with a net positive or negative charge
SO SO42-

12. Exceptions to the Octet Rule
Octet rule cannot be satisfied if a molecule has an odd number of valence electrons
NO2, NO, ClO2
Phosphorus and sulfur expand the octet to ten or twelve electrons

13. Bond Dissociation Energies
Energy required to break a bond
Large bond dissociation energies denote strong covalent bonds

15. Resonance
Resonance structures are multiple possible dot structures for the same molecule or polyatomic ion
Actual bonding is a hybrid of both, used when neither is a sufficient description

16. 8.3 Bonding Theories

17. Molecular Orbitals Formed when atoms bond
Belong to the molecule as a whole (not to one atom)
If occupied by electrons of a covalent bond it is called a bonding orbital

18. Bonding Orbitals Sigma bond: combine end-to-end
Pi bond: combine side-to-side

19. VSEPR Theory Valence-Shell Electron-Pair Repulsion Theory
Explains three-dimensional shapes of molecules
Lone pairs of electrons strongly repel bonding pairs of electrons
3D shapes of molecules

21. Hybrid Orbitals
Hybridization provides information about molecular bonding and shape
Describes double and triple covalent bonds

22. 8.4 Polar Bonds and Molecules

23. Bond Polarity
When atoms share electrons equally, a nonpolar covalent bond is formed
When atoms do not share electrons equally, a polar covalent bond is formed
The more electronegative atom is partially negative
The less electronegative atom is partially positive
Polar molecules are formed when one end of the molecule is partially positive and the other end is partially negative
A molecule with two poles is called a dipole

24. Attractions Between Molecules
Intermolecular attractions are weaker than ionic or covalent bonds
Hydrogen bonds- hydrogen bonded to very electronegative atom is also weakly bonded to another highly electronegative atom (F, O, N)
Van der Waals forces
Dipole interactions- based on partial charges
Similar to, but weaker than ionic bonds
Dispersion forces- based on movement of electrons
Increase with number of electrons in a molecule

25. Intermolecular Attractions and Molecular Properties
Varying intermolecular attractions explain the variety of physical properties among covalent compounds
Network solids- all atoms are covalently bonded to each other