Chapter 10: Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory Flashcards Preview

Gen Chem 01 > Chapter 10: Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory > Flashcards

Flashcards in Chapter 10: Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory Deck (26):

VSEPR theory

Valence Shell Electron Pair Repulsion theory

Based on electron groups repeling one another through coulombic forces


Electron groups


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)
Trigonal planar

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
sp2 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
sp3 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
sp3d 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
sp3d2 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

B2, C2, N2 = pi 2p then sigma 2p

O2, F2, Ne2 = 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