Chemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals Chapter 10

Valence Shell Electrons Valence Shell Electron Pair Repulsion (VSEPR) Theory Valence Shell Electrons the outer shell electrons of an atom, the ones involved in bonding for a given elements, # of valence electrons = Group Number C: Group 4A, 4 valence electrons O: Group 6A, 6 valence electrons Draw Lewis structure NH3

– Key: Electrons are all negatively charged (e-) – Action: electron pairs (bonding pairs & lone pairs ) around the central atom repel each other to keep themselves as far away as possible – Outcome: Maximum separation Minimum repulsion Symmetry (equal repulsion) – Applications: predict the electron pair geometry predict the bond angle predict the molecular geometry predict the hybridization of the central atom predict the polarity of the molecule – Electron pair geometry : the arrangement of all electron pairs (bonding pairs & lone pairs ) surrounding the central atom of the molecule. – Molecular geometry: the arrangement of only bonding pairs surrounding the central atom of the molecule

VSEPR CO2 Valence Shell Electron Pair Repulsion Example? O=C=O Class AB2 2 Class # of bonding pairs # of lone pairs Electron pair geometry Molecular Geometry linear Total # of electron pairs Example? CO2 O=C=O

VSEPR # of bonding pairs # of lone pairs Total # of electron pairs geometry Molecular Geometry Class trigonal planar trigonal planar AB3 3 3 trigonal planar bent AB2E 2 1 3

VSEPR # of bonding pairs # of lone pairs Total # of electron pairs geometry Molecular Geometry Class tetrahedral tetrahedral AB4 4 4 tetrahedral trigonal pyramidal AB3E 3 1 4 tetrahedral bent AB2E2 2 2 4

# of bonding pairs # of lone pairs Total # of electron pairs geometry Molecular Geometry Class trigonal bipyramidal trigonal bipyramidal AB5 5 5 trigonal bipyramidal seesaw AB4E 4 1 5 trigonal bipyramidal T-shaped AB3E2 3 2 5 trigonal bipyramidal linear AB2E3 2 3 5

VSEPR # of bonding pairs # of lone pairs Total # of electron pairs geometry Molecular Geometry Class octahedral octahedral AB6 6 6 octahedral square pyramidal AB5E 5 1 6 octahedral square planar AB4E2 4 2 6

Bond angle Periodic table # of valance shell electrons Electron pair geometry Molecular geometry Lewis structure Total # of electron pairs around central atom Hybridization of the central atom What is the electron pair geometry, bond angle, molecular geometry and hybridization of the central atom of NH3? electron pair geometry: tetrahedral bond angle: 109.5o molecular geometry: trigonal pyramidal hybridization: sp3

bonding-pair vs. bonding Repulsive force bonding-pair vs. bonding pair repulsion lone-pair vs. lone pair repulsion lone-pair vs. bonding >

Dipole Moments and Polar Molecules Predicting Polarity of Molecules: Dipole Moments and Polar Molecules m = Q x r Q is the charge r is the distance between charges 1 D = 3.36 x 10-30 C m H F electron rich region electron poor d+ d- • A measure of the polarity of a molecule • The dipole moment can be determined experimentally • Measure in Debye units • Predict polarity by taking the vector sum of the bond dipoles

Molecules containing net dipole moments are called polar molecules Molecules containing net dipole moments are called polar molecules. Otherwise they are called nonpolar molecules because they do not have net dipole moments. Diatomic molecules: Determined by the polarity of bond Polar molecules: HCl, CO,NO Nonpolar molecules: H2,F2,O2 Molecules with three or more atoms: determined by the polarity of the bond and the molecular geometry Dipole moment is a vector quantity, which has both magnitude and direction.

• Symmetric molecules such as these are nonpolar because the bond dipoles cancel • All of the basic shapes are symmetric, or balanced, if all the domains and groups attached to them are identical O C O C O O Linear molecule Nodipolar moment bent molecule Net dipolar moment

Which of the following molecules have a dipole moment? H2O, CO2, SO2, and CH4 O H S O dipole moment polar molecule dipole moment polar molecule C H C O no dipole moment nonpolar molecule no dipole moment nonpolar molecule

Molecular Geometry & Hybridization • Lewis structures and VSEPR do not tell us why electrons group into domains as they do • How atoms form covalent bonds in molecules requires an understanding of how orbitals interact Molecular Geometry & Hybridization Covalent Bonding Theories • Valence bond (VB) theory – Bonding is an overlap of atomic orbitals • Includes overlap of hybrid atomic orbitals • Molecular orbital (MO) theory – Bonding happens when molecular orbitals are formed

Hybridization – mixing of two or more atomic orbitals to form a new set of hybrid orbitals. Mix at least 2 nonequivalent atomic orbitals (e.g. s and p). Hybrid orbitals have very different shape from original atomic orbitals. Number of hybrid orbitals is equal to number of pure atomic orbitals used in the hybridization process. Covalent bonds are formed by: Overlap of hybrid orbitals with atomic orbitals Overlap of hybrid orbitals with other hybrid orbitals

Predict the Hybridization of the Central Atom Total # of electron pairs Electron pair geometry Hybridization Examples 2 linear sp BeCl2 3 trigonal planar sp2 BF3 4 tetrahedral sp3 CH4, NH3, H2O 5 trigonal bipyramidal sp3d PCl5 6 octahedral sp3d2 SF6 Total # of electron pairs = number of hybrid orbital (sum of the superscripts!)

Formation of sp3 Hybrid Orbitals

Formation of sp Hybrid Orbitals

Hybridization in molecules containing double and triple bonds Two Kinds of covalent Bonds • Sigma bond(σ): bonding density is along the internuclear axis – head to head overlap of two hybrid orbitals – head to head overlap of p + hybrid • any s character to the bond = sigma bond • Pi bond(π): bonding density is above and below the internuclear axis – Sideway overlap p + p

(σ): Sigma bond

(π): Pi bond

Multiple bonds • Double bond: One sigma and one pi bond between the same atoms –Example: ethylene • Triple bond: One sigma and two pi bonds between the same atoms –Example: acetylene

Three Types of bonds: 1. Sigma bond(σ): e-density (overlap region) is along the internuclear axis 2. Pi bond(π): e-density (overlap region) is above and below the plane of the internuclear axis. Relative reactivity of bonds: • Pi bonds are more reactive than sigma bonds • Less energy is needed to break a pi bond than a sigma bond 3. Delocalized Bonds • are not confined between two adjacent bonding atoms, but actually extend over three or more atoms. • less reactive than normal pi bonds • Example: benzene

Sigma (s) and Pi Bonds (p) Framework of the molecule is determined by the arrangement of the sigma-bonds; Hybrid orbitals are used to form the sigma bonds and the lone pairs of electrons. Sigma (s) and Pi Bonds (p) Single bond 1 sigma bond Double bond 1 sigma bond and 1 pi bond Triple bond 1 sigma bond and 2 pi bonds How many s and p bonds are in the acetic acid (vinegar) molecule CH3COOH? C H O s bonds = 6 + 1 = 7 p bonds = 1

Valence bond theory O No unpaired e- Should be diamagnetic But experiment show that there are two unpaired electrons—paramagnetic. Molecular orbital theory – bonds are formed from interaction of atomic orbitals to form molecular orbitals.

Molecular Orbital (MO) Theory • Bonds are formed from interaction of atomic orbitals to form molecular orbitals. MO theory takes the view that a molecule is similar to an atom • The molecule has molecular orbitals that can be populated by electrons just like the atomic orbitals in atoms • Number of MOs formed = number of atomic orbitals combined

Energy levels of bonding and antibonding molecular orbitals in hydrogen (H2). A bonding molecular orbital has lower energy and greater stability than the atomic orbitals from which it was formed. An antibonding molecular orbital has higher energy and lower stability than the atomic orbitals from which it was formed.

Molecular Orbital (MO) Configurations -MOs follow the same filling rules as atomic orbitals The number of molecular orbitals (MOs) formed is always equal to the number of atomic orbitals combined. The more stable the bonding MO, the less stable the corresponding antibonding MO. The filling of MOs proceeds from low to high energies. Each MO can accommodate up to two electrons. Use Hund’s rule when adding electrons to MOs of the same energy.(Electrons spread out as much as possible, with spins unpaired, over orbitals that have the same energy) The number of electrons in the MOs is equal to the sum of all the electrons on the bonding atoms.

( ) - bond order = 1 2 Number of electrons in bonding MOs Number of electrons in antibonding MOs ( - ) bond order ½ 1 ½