Chapter 9 Molecular Shape

VSEPR Theory Valence Shell Electron Pair Repulsion Pairs of valence electrons are arranged as far apart from each other as possible. Determines the shape of a molecule.

Hybrid Orbitals Orbitals of electrons in a bond Combination of shapes and properties of the original atomic orbitals sp, sp2, & sp3 hybridized

Bonding Orbital To determine the number of bonding orbitals in a molecule, count the number of bonds around the CENTRAL ATOM of the molecule (things attached to the central atom) Each bond will count for one bonding orbital In this case multiple bonds will only count for one bonding orbital

Nonbonding Orbital To determine the number of nonbonding orbitals in a molecule, count the number of lone electrons pairs around the CENTRAL ATOM of the molecule Each PAIR of electrons will count for one nonbonding orbital

Steps for determining shape: Draw the Lewis Dot Structure Count the number of bonding orbitals Count the number of nonbonding orbitals Determine the shape Redraw the Lewis Dot Structure to represent the shape

VSEPR Shapes We will learn about five shapes: Linear Trigonal Planar Tetrahedral Pyramidal Bent

1. Linear CO2 O = C = O 2 bonding orbitals 0 nonbonding orbitals Bond angle = 180o sp hybridized

1. Linear

2. Trigonal Planar BCl3 Cl B Cl Cl 3 bonding orbitals 0 nonbonding orbitals Bond angle = 120o sp2 hybridized

2. Trigonal Planar

3. Tetrahedral F CF4 F C F 4 bonding orbitals 0 nonbonding orbitals Bond angle = 109.5o sp3 hybridized

3. Tetrahedral

4. Pyramidal NH3 H N H H 3 bonding orbitals 1 nonbonding orbitals Bond angle = 107o sp3 hybridized

4. Pyramidal

5. Bent H2O H O H 2 bonding orbitals 2 nonbonding orbitals Bond angle = 105o sp3 hybridized

5. Bent

Shape # orbitals b.o. nb. o. Bond < linear 2 180o trigonal planar 3 120o bent 4 105o pyramidal 1 107o tetrahedral 109.5o

Lewis Structures in the Ghetto Carbon makes 4 bonds, no unshared pairs Nitrogen: 3 bonds, 1 unshared pair Oxygen: 2 bonds, 2 unshared pairs Hydrogen: 1 bond, no unshared Halogens: 1 bond, 3 unshared pairs

Bond Length Different pairs of atoms form different bond lengths As you go down a group, bond length increases b/c atoms are getting larger Multiple bonds are shorter than single bonds Single > double > triple

Polar Covalent Bonds Polar Bond - atoms don’t share electrons equally. Example: When O bonds with H, there is a difference in electronegativity between O & H which results in more of the electrons surrounding the O more frequently. This causes the O end of the bond to be more negative and the H end of the bond to be more positive.

δ- O – H δ+ Polar Covalent Bonds lower case Greek letter delta indicates charge This bond contains a dipole, which merely means a positive and a negative end.

Electronegativity difference In order to determine the polarity of a bond, calculate the difference in electronegativity Electronegativity difference Type of Bond >2.0 Ionic 2.0 – 0.4 Polar Covalent <0.4 Nonpolar

Polar molecules Molecules can be polar just like bonds The polarity of a molecule is determined by: The polarity of the individual bonds The shape of the molecule

Polar molecules If a molecule is symmetrical (in all 3 planes) it is nonpolar If a molecule is asymmetrical (not symmetrical) it will be polar

Polar Mcl Exception Tetrahedral Mcls are nonpolar when all four atoms attached to the central atom are the same. CH4 – nonpolar CH3Cl - polar

Polar Mcl Exception In the Ghetto Tetrahedral, linear, and trigonal planar are usually nonpolar if the same thing is attached all the way around (like CH4 ) Pyramidal and bent are inherently asymmetrical and are usually polar

Shape Exceptions B & Al will form Trigonal Planar molecules BF3 & AlH3 Be will form linear molecules BeCl2