Chapter 14: Covalent bonds

1. Single covalent bonds
2. Double and triple covalent bonds
3. Covalent compounds
4. Coordinate covalent bonds
5. Resonances (different dot structures for same compound)
6. Exceptions to the octet rule
7. Molecular orbitals
8. VSEPR Theory (unbonded electron pairs influence molecule shape)
9. Hybrid orbitals (sp¹, sp², etc.)
10. Polar bonds
11. Polar molecules
12. Bond disassociation energy (energy needed to break one single bond) [Reference]
13. Intermolecular attractions
14. Properties of molecular substances

Chapter 14

14-1 Single Covalent Bonds

- When two atoms trying to gain electrons share them, a covalent bond is made. If two people have 4 puzzle pieces each, but 8 are needed to complete the puzzle, it only makes sense that they would share.

- Lewis dot structures can be used to show this: X:X (the two electrons in the middle are being shared by the two atoms X and X).

14-2 Double and Triple Covalent bonds

- Two atoms can share more than one pair of electrons. For example, here is a complete Lewis Structure: :N:::N: (the N's are sharing 6 electrons).

14-3 Covalent Compounds

- Understanding covalent compounds is easy using Lewis dot structures.

14-4 Coordinate Covalent Bonds

- When one atom supplies both electrons in a covalent bond, it a coordinate covalent bond.

14-5 Resonance

- Resonance occurs when more than 1 dot structure is valid for a molecule.

14-6 Exceptions to the Octet Rule

- Hydrogen only needs two electrons to complete it's outer energy level.

- Sometimes it is impossible to make a complete Lewis structure.

14-7 Molecular Orbitals

- Quantam mechanics provides a different model for molecular compounds. It says that the two orbitals being shared overlap to create the bonding molecular orbital and the antibonding molecular orbital.

14-8 VSEPR Theory

- Molecules are shaped so that valence electron pairs are as far away from each other as possible. Unshared electron pairs are een more repulsive, and the molecule will bend itself away from them.

- See Chempire's Online Molecule Shapes Guide.

14-9 Hybrid Orbitals

- Two orbitals mix and form a hybrid where the molecule is sharing electrons.

14-10 Polar Bonds

- If the two atoms sharing electrons are the same, the sharing is equal. But if the two atoms are not the same, one (the one with the higher electronegativity) is going to pull harder than the other.

- See Chempire's Online Electronegativity Chart.

14-11 Polar Molecules

- A molecule is polar when one end is more negative than the other.

- A dipole is a molecule with two distinct polar ends (one more positive, one more negative).

14-12 Bond Dissociation Energies

- A bond dissociation energy is how much energy is needed to break a bond.

- See Chempire's Online Bond Dissociation Energies Chart.

14-13 Intermolecular Attractions

- Intermolecular attractions are forces that act between molecules.

- The two weak intermolecular forces are called the van der Waals forces. 1) Dispersion forces: At a point in time, most of an atoms could by chance be gathered all at one end of a molecule. For that split second, that end will be more negative. This generates a dispersion force. 2) Dipole interactions: The negative end of one polar molecule is attracted to the positive end of another polar molecule. This is a dipole interaction.

- Hydrogen bonds are the strongest intermolecular forces. If a hydrogen atom is bonded to an extremely electronegative atom, the latter will have the shared electrons most of the time. When it does, the hydrogen atom is, for all practical purposes, a naked proton. As you can guess, this unshielded proton will make a very polar molecule.

14-14 Properties of Molecular Substances

- Most stable molecular substances exist as network solids, where every atom is bonded together. Diamond is a network of carbon atoms.

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