Chapter 10: Chemical Bonding II: Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory Flashcards
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)
- # of orbitals combined = # of hybrid orbitals formed
- Bonds = two half-filled orbitals with spin-pairing electrons overlapped (or a filled orbital over an empty)
- Interaction = either (a) alignment along axis between atoms or (b) parallel to each other & perpendicular to the interatomic axis
- Maximizes bonding and stability, minimizes energy
- 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)
- Covalent bond that results when the interacting atomic orbitals point along the axis connecting the two bonding nuclei
- End-to-end overlap
- Single bond = sigma bond
Multiple bond = 1 sigma bond + x pi bonds - Stronger than pi bonds
- Free bond rotation is allowed (single bond)
Pi bond (5)
- Covalent bond that results when the bonding atomic orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nuclei
- Side-by-side orbital overlap
- pi bond = multiple bonds
- Weaker than sigma bonds
- 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)
- # of MOs formed = # of atomic orbitals combined
- The more stable the bonding MO, the less stable the antibonding MO
- The filling of MOs proceeds from low to high energies (aufbau principle)
- Each MO has max of two electrons (Pauli exclusion principle)
- Use Hund’s rule when adding electrons to MOs of the same energy
- # 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
