VSEPR theory

Valence Shell Electron Pair Repulsion theory

Based on electron groups repeling one another through coulombic forces

Electron groups

Includes:

Lone pairs

Single bonds

Multiple bonds

Single electrons

Effect of lone pairs

Occupy more space on the central atom because electron density is exclusively on the central atom

Make bonding pair angles smaller

Bonding v. lone pair repulsion

Bonding-bonding < Lone-bonding < Lone-lone

Hatched wedge

Bond "goes into" page

Solid wedge

Bond is "coming out of" the page

Polar molecule (4)

Must have:

Polar bonds (bond dipole moments)

Unsymmetrical shape (vector addition)

*Lone pairs affect polarity

*Polarity affects intermolecular forces of attraction ("like dissolves like," boiling points)

Dipole moment

µ = qr = partial charges * distance btwn charges

Represented with a vector arrow (pointing toward negative charge)

Electron geometry and polarity

__ Nonpolar__ (when identical terminals)

Linear

Trigonal planar

Tetrahedral

**Polar**

Bent

Trigonal pyramidal

Valence bond theory (5)

1. # of orbitals combined = # of hybrid orbitals formed

2. Bonds = two half-filled orbitals with spin-pairing electrons overlapped (or a filled orbital over an empty)

3. Interaction = either (a) alignment along axis between atoms or (b) parallel to each other & perpendicular to the interatomic axis

4. Maximizes bonding and stability, minimizes energy

5. Makes new set of degenerate orbitals

Hybrid orbitals

Mixture of multiple standard atomic orbitals that correspond more closely to the actual distribution of electrons in chemically bonded atoms

Sigma bond (5)

1. Covalent bond that results when the interacting atomic orbitals point along the axis connecting the two bonding nuclei

2. End-to-end overlap

3. Single bond = sigma bond

Multiple bond = 1 sigma bond + x pi bonds

4. Stronger than pi bonds

5. Free bond rotation is allowed (single bond)

Pi bond (5)

1. Covalent bond that results when the bonding atomic orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nuclei

2. Side-by-side orbital overlap

3. pi bond = multiple bonds

4. Weaker than sigma bonds

5. Bond rotation is restricted (double bond)

Geometry/angles/hybridization: 2 electron groups (6)

2 bonding groups

0 lone pairs

linear **electron** geometry

linear **molecular** geometry

180° bond angles

sp hybridization

Geometry/angles/hybridization: 3 electron groups (8)

trigonal planar **electron** geometry

sp^{2} hybridization

3 bonding groups/0 lone pairs:

trigonal planar **molecular** geometry

120° bond angles

2 bonding groups/1 lone pair:

bent **molecular** geometry

<120° bond angles

Geometry/angles/hybridization: 4 electron groups (11)

tetrahedral **electron** geometry

sp^{3} hybridization

4 bonding groups/0 lone pairs:

tetrahedral **molecular** geometry

109.5° bond angles

3 bonding groups/1 lone pair:

trigonal pyramidal

<109.5° bond angles

2 bonding groups/2 lone pairs:

bent **molecular** geometry

<109.5° bond angles

Geometry/angles/hybridization: 5 electron groups (14)

trigonal bipyramidal **electron** geometry

sp^{3}d hybridization

5 bonding groups/0 lone pairs:

trigonal bipyramidal **molecular** geometry

120° (equatorial) & 90° (axial) bond angles

4 bonding groups/1 lone pair:

seesaw **molecular** geometry

<120° (equatorial) & <90° (axial) bond angles

3 bonding pairs/2 lone pairs:

t-shaped **molecular** geometry

<90° bond angles

2 bonding pairs/3 lone pairs:

linear **molecular** geometry

180° bond angles

Geometry/angles/hybridization: 6 electron groups (11)

octahedral **electron** geometry

sp^{3}d^{2} hybridization

6 bonding groups/0 lone pairs:

octahedral **molecular** geometry

90° bond angles

5 bonding groups/1 lone pair:

square pyramidal **molecular** geometry

<90° bond angles

4 bonding pairs/2 lone pairs:

square planar **molecular** geometry

90° bond angles

Molecular orbital theory

Applies Schrödinger's wave equation to a molecule to calculate a set of molecular orbitals

Electrons belong to whole molecule, as do orbitals (delocalization of electrons)

Linear combination of atomic orbitals (LCAO)

Method of adding together atomic orbitals to make molecular obitals using a weighted average

Since orbitals are wave functions --> constructive & destructive interference

Bonding molecular orbital

Occurs when wave functions combine constructively

Has lower energy, greater stability than atomic orbitals from formation

Most electron density is between the nuclei

Antibonding molecular orbital

Occurs when wave functions combine destructively

Has higher energy and lower stability than atomic orbitals from formation

Most of electron density is outside the nuclei (node between nuclei)

Writing MO diagrams (6)

1. # of MOs formed = # of atomic orbitals combined

2. The more stable the bonding MO, the less stable the antibonding MO

3. The filling of MOs proceeds from low to high energies (aufbau principle)

4. Each MO has max of two electrons (Pauli exclusion principle)

5. Use Hund's rule when adding electrons to MOs of the same energy

6. # electrons in MOs = # electrons on bonding atoms

Bond order

Difference between # valence electrons in bonding and antibonding orbitals (divided by two)

Can = 1/2

Higher bond order = stronger (stability & bond energy) and shorter bonds (lenght)

Bond order = 0 = unstable (no bond will form)

MO diagrams for second-period homonuclear diatomic molecules

B_{2}, C_{2}, N_{2} = **pi** 2p **then** **sigma** 2p

O_{2}, F_{2}, Ne_{2} = **sigma** 2p **then pi** 2p

MO diagrams for second-period heteronuclear diatomic molecules

More electronegative atom:

lower atomic orbitals

closer to molecular orbitals

nonbonding orbitals remain localized here