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Ch. 8 Covalent Bonding8.1 Molecular Compounds
I. Molecules• A. Neutral groups of atoms
joined by covalent bonds
• B. Covalent bonds: atoms share electrons instead of giving/taking when electronegativity similar
Ex. CH4, CO2, H2O, NH3
• C. Diatomic molecule: pairs of same atoms (O2)
II. Covalent Properties• A. Weaker than ionic bonds
• B. Lower melting/boiling pts.
• C. Between non-metals
• D. Usually gas or liquid at room temp.
8.2 The Nature of Covalent Bonding
I. Single Covalent Bond• A. Sharing of a single electron pair between
atoms
• B. Can represent e- pair as line
• C. Unshared e-s called “lone pair”, e-s between atoms represent the bond
II. Octet Rule• A. Still follow rule that atoms are most stable
with 8 valence electrons
• B. Octet comes from atoms own e- and shared ones
III. Multiple Bonds• A. Double covalent bond: sharing two PAIRS
of electrons
• B. Triple covalent bond: sharing three PAIRS of electrons
Ex. N2
IV. Lewis Dot Models
• A. Count up valence electrons of all atoms
• B. Usually first atom in formula is in center
• C. Start with single bonds (two e-), see if remaining e- can be used to make all stable
• D. Try double or triple bonds if not
• E. Molecule charges mean missing or extra collective electrons (use brackets around molecule)
V. Coordinate Covalent Bonds• A. When one atom contributes both electrons
for a bond
• B. Ex.: Carbon Monoxide (CO)If electrons shared equally:
Since not stable, you get:
VI. Resonance Structures• A. Have two or more equally stable
electron dot structures
• B. Rotate between structures by e- moving
Example:
Ozone
VII. Octet Exceptions• A. The octet rule cannot be met in molecules with
odd numbers of e-
• B. Some atoms can have more or less than 8 valence e- and be stable
Boron trifluoride
VIII. Formal Charge• A. If two different e- dot structures can both be
stable, formal charge can determine which is more stable
• B. Formal charge = valence e- - e- in molecule for atom
• C. E- for atom = lone pairs + ½ of bonding e-
• D. More stable = less charges and any negatives on most electronegative atoms
IX. Bond Dissociation Energies• A. Energy needed to break apart two covalently
bonded atoms• B. Stronger for multiple bonds
Ex. C – C = 347 kiloJoule (kJ)
C = C = 614 kJ
C C = 839 kJ• C. Reaction energy = sum of energy to break bonds
(reactants) – sum of energy to make bonds (products)
8.3 Bonding Theories
I. Hybrid Orbitals• A. Combination of orbitals of an atom • B. Names based on orbitals and # of e- involved
• C. When 4 atoms/lone pairs attached to central atom there is one s and three p orbitals used (sp3)
II. Double/Triple Bond Hybrids• A. When 3 atoms/lone pairs
attached to central atom use one s and two p (sp2)
• B. Left-over P orbital perpendicular to rest
• C. 2 atoms/lone pairs use 1 s and 1 p (sp), two p perpendicular
III. More Than an Octet
• A. To exceed 8 valence electrons, atoms use d orbitals
• B. 10 e- (sp3d)
• C. 12 e- (sp3d2)
• D. Double and triple bonds need the same hybrid orbitals as single bonds and lone pairs
IV. Molecular Orbitals• A. E- fit in orbitals in atoms, when molecules
combine molecules have own orbitals
• B. Two S orbitals overlap to form sigma (σ) orbitals with σ bonds
• C. P-orbitals can align or be parallel to each other
• D. Aligning form σ orbitals, parallel form Pi () orbitals
• E. σ bonds stronger than bonds
• F. Single covalent bonds are made of σ bonds
• G. Double bonds = 1 σ and 1 , triple = 1 σ and 2
Hybrid orbitals attach to form molecular orbitals
Ex. C2H4
Ex. C2H2
V. VSEPR Theory• A. Shows 3-D structure of molecules
• B. Valence-Shell Electron-Pair Repulsion: because e- repel each other, 3-D shape puts pairs farthest apart
• C. Use shape that puts E- PAIRS and BONDED ATOMS farthest apart
VI. Examples• A. 4 bonds = tetrahedral
• B. 3 bonds, 1 lone pair = pyramidal
• C. 2 bonds, 2 lone pairs = bent
• D. 2 bonds = linear (CO2)
8.4 Polar Bonds and Molecules
I. Bond Determination
• B. Difference 2 or more = ionic bond
• C. 0.4 or less difference = covalent bond
• D. 0.4 - 2 difference = polar covalent bond
• A. Difference between ionic and covalent bond is electronegativity of atoms
II. Polar Bond• A. Unequal sharing of electrons
• B. Cl (3) pulls e- closer to it than H (2.1), Cl gets slight negative charge (-), H slight positive (+)
III. Polar Molecule• A. When polar bond
makes entire molecule polar (“dipole”)
• B. Sometimes polar bonds don’t make molecule polar
IV. Intermolecular Attractions• A. Molecules attracting each other
• B. Dipole attractions: charge of one molecule attracted to opposite charge on anotherEx. Dipole-Dipole
V. Induced dipoles• When electrons temporarily shift positions to form artificial dipoles (“Van der waals Forces”/ “London Dispersion Forces”)
VI. Hydrogen Bonds• Hydrogen in a polar molecule attracted to
electronegative atom on neighboring molecule
• Example: H2O
Hydrogen Bond
VII. Network Solids• A. Most covalent molecules melt at low temp.
• B. Network solids are very stable covalently bonded molecules with high melting pts.
Vaporizes at 3500ºC
Silicon Carbide melts at 2700ºC