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)

VII. Other VSEPR ShapesSquare Pyramidal

See-Saw

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