1. Covalent Bonds
Chapter 8

2. Review
Define: Valence electrons Cation Anion Octet rule Ionic bond Ionic compound Draw Lewis structures for Carbon, Hydrogen, and Chlorine Write compound formed when boron and sulfur combine. Name it. Write compound formed when calcium and iodine combine. Name it. Write compound formed when barium and a phosphate ion combine. Name it.

3. Covalent bonds Sharing electrons to achieve octet (noble gases)
Molecule: atoms joined together with covalent bond
Diatomic molecule: molecule of 2 atoms
H2, N2, O2, F2, Cl2, Br2, I2
Usually the same atom, but CO is also considered diatomic molecule
Compound: two or more DIFFERENT atoms held together by covalent OR ionic bond

4. Molecular compounds
Lower melting and boiling points than ionic compounds
Many are gases or liquids at room temperature
2 or more nonmetals/metalloids (Groups 4A-7A)
Molecular formula: shows how many atoms of each element a molecule contains
H2O is 2 hydrogen atoms and 1 oxygen atom
Whereas, NaCl is a 1:1 ratio of sodium atoms to chloride atoms

5. Types of covalent bonds
Single covalent bond: sharing a single pair of electrons (H2)
Double covalent bond: sharing 2 pairs of electrons (O2)
Triple covalent bond: sharing 3 pairs of electrons (N2)

6. Coordinate covalent bonds
One atom contributes both bonding electrons
The shared electron pair comes from one of the bonding atoms
Carbon monoxide and polyatomic ions

7. Naming covalent bonds Name elements in order listed in formula
Use prefix to indicate number of each kind of atom
Omit mono- when formula contains one atom in the first position
Suffix of second element is –ide
H2O: Dihydrogen monoxide
CO: Carbon monoxide
Table 9.4 on page 269

8. Bond dissociation energies
Energy required to break covalent bond
Large bond dissociation energy = strong covalent bond
C-C 347KJ/mol
C=C 657KJ/mol
C≡C 908KJ/mol
Strong carbon-carbon bonds help explain stability of carbon compounds
Compounds with C-C or C-H single covalent bonds (methane) are unreactive because dissociation energy is so high

9. Resonance
Resonance structure: occurs when it is possible to draw 2 or more valid electron dot structures that have same # of electron pairs for a molecule or ion
Electrons do not change back and forth
Ozone (O3)

10. Exceptions Total # valence electrons is an ODD number.
There are also molecules in which atoms have fewer, or more, than 8
NO2, ClO2, and NO
Several with EVEN # of valence electrons
Atom acquires less than octet
Some atoms, S and P, sometimes expand octet to include 10 to 12 electrons
PCl5 and SF6

11. Molecular orbitals
When 2 atoms combine, atomic orbitals overlap to produce molecular orbitals
Atomic orbital belongs to an atom and molecular orbital belongs to the molecule
2 electrons in each atomic orbital and 2 electrons therefore in molecular orbital

12. Valence-shell electron-pair repulsion theory (VSEPR)
The repulsion between electron pairs causes molecular shapes to adjust so that electrons stay as far apart as possible
Unshared pairs are important in predicting shapes of molecules
Tetrahedral: Methane (CH4)
Pyramidal: Ammonia (NH3)
Bent: Water (H2O)
Linear: Carbon dioxide (CO2)

13. Review Define electronegativity and describe trend
Table 6.2 p177
Electronegativity difference tells you what kind of bond is likely to form

14. Polar bonds and molecules
Nonpolar covalent bond: atoms pull equally and electrons are equally shared
H2, O2, Cl2
Polar covalent bond: electrons are shared unequally
More electronegative atom attracts electrons more strongly and gains a slight negative charge. The less electronegative atom has a slightly positive charge.
Increase electronegativity, increase polarity of bond
Electronegativity greater than 2.0=ionic bond

15. Polar molecules
One end of the molecule is slightly - and other end is slightly +
Molecule with 2 poles is called a dipole
Effect of polar bonds on the polarity of entire molecule depends on shape and orientation of polar bonds
Bent=polar
Linear=nonpolar
Tetrahedral=nonpolar
Pyramidal=polar

16. Attractions between molecules
Intermolecular attractions are weaker than either ionic or covalent bonds
Van der Waals forces (weakest)
Dipole interactions and dispersion forces
Hydrogen bonds

17. Dipole interactions
Occur when polar molecules are attracted to one another
Attractions occurs between oppositely charged regions of polar molecules

18. Dispersion forces Occurs between polar or non-polar molecules
Caused by movement of electrons
Weakest of all molecular interactions

19. Hydrogen bonds
Attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom
Must involve H because it is only chemically reactive element with valence electrons that are not shielded from nucleus by other electrons
Bonded with oxygen, nitrogen, or fluorine
Strongest intermolecular force
Surface tension
Cohesion/adhesion

20. Hydrogen bonding

21. Surface tension “skin” on surface of water
Water molecules at surface cannot form H-bonds with air molecules so they are drawing into body of the liquid producing surface tension
Water is pulled together creating the smallest surface area possible

22. Cohesion and adhesion
Cohesion: water attracted to other water molecules Adhesion: water attracted to other materials *Capillary action*