10.5-10.7 Molecular Shape, Polarity and Valence Bond Theory

Molecular Shape and Polarity A Lewis structure can help us determine whether a molecule is polar or non-polar. If dipole moments of individual polar bonds sum up to a net dipole moment, the molecule is polar. Ex. HCl If the dipole moments of individual polar bonds cancel each other (sum is zero), the molecule is non-polar. Ex. CO2 Vectors (arrows) can be drawn to show the direction of electron pull and magnitude of partial charge (length of vector). Electron potential maps can also be used to show electron rich (red), electron moderate (yellow), and electron poor (blue) areas.

Let’s Try a Practice Problem! Determine if CCl4 is polar. Draw the Lewis structure, determine if there are polar bonds and if so, superimpose a vector on each bond, and then determine if the polar bonds add together to form a net dipole moment. No, carbon tetrachloride is a nonpolar molecule since the dipole moments of the individual polar bonds cancel each other out.

Valence Bond Theory The valence electrons of the atoms in a molecule reside in an atomic orbital (s, p, d, and f) or they may be in hybrid combinations of these. The overlap of two half-filled orbitals results in a covalent bond. However, the overlap of a filled orbital with an empty orbital is also possible. This is called a coordinate covalent bond. The geometry of overlapping orbitals determines the shape of the molecule. I will draw atomic orbital diagrams of both hydrogens that make up a molecule of hydrogen to show this overlap.

Let’s Try a Practice Problem! The answer to the question “what is a chemical bond?” depends on the bonding model. Answer these questions: a.) What is a covalent bond according to the Lewis model? b.) What is a covalent bond according to valence bond theory? a.) According to the Lewis model, a covalent bond exists when two atoms of a molecule share electrons; this is represented by dots. b.) According to valence bond theory, a covalent bond exists when half-filled atomic orbitals overlap

Hybridization A mathematical procedure in which the standard atomic orbitals are combined to form new atomic orbitals called hybrid orbitals that correspond more closely to the actual distribution of electrons in chemically bonded atoms. (These orbitals have different shapes and energies). Molecules with 4 electron pairs (remember this means shared or unshared electrons) around the central atom are said to be sp3 hybridized and have bond angles of 109.5o. Molecules with three electron pairs around the central atom are sp2 hybridized with 120o bond angles. Molecules with two electron pairs around the central atom are sp hybridized with 180o bond angles. There are additional hybridized orbitals, but these are controversial, and are not included in the AP curriculum.

An Example of sp3 Hybridization If you think about carbon’s electron configuration (not its Lewis structure), how many bonds would you expect carbon to form? 2, because carbon’s electron configuration is 1s22s22p2, so it has two half filled p orbitals. But we know that carbon forms four bonds, as in methane. The reason that carbon can form four bonds is because it’s 2s and 2p sublevels hybridize (or combine), to form a new type of or orbital with a new shape. This allows for a more stable molecule. Carbon now has 4 sp3 hybridized orbitals.

SP2 Hybridization and Double Bonds The hybridization of one s and two p orbitals is sp2 hybridization. ..\Pictures, Videos, and Links for PowerPoints\ethene Formation of a double bond Clip.avi Molecules with three electron pairs around the central atom are sp2 hybridized with bond angles of 120o. Let’s consider formaldehyde (H2CO) This structure has trigonal planar geometry and the C-H bonds are single bonds with end-to-end orbital overlap. This type of overlap is called a sigma (σ)bond. The C=O double bond is formed between a p overlap σ bond, and a parallel p overlap pi (π) bond.

A single bond, a bond order of 1, is always made up of a σ-bond and a double bond, a bond order of 2, is always made up of one σ-bond and one π-bond. π-bonds are weaker than σ-bonds leading to σ-bonds having higher bond energies. When a double bond is formed, rotation around the bond is restricted, this leads to structural isomers existing. For example, 1,2-dichloroethene can exist as either, cis-1,2-dichloroethene, or trans-1,2-dichloroethene. Cis- means same side and trans-means opposite sides.

sp Hybridization and Triple Bonds Hybridization of one s and one p orbital results in sp hybridization. Molecules with two pairs of electrons around the central atom are said to be sp hybridized (linear molecules) with bond angles of 180o. A triple bond, bond order of 3, contains one σ-bond and two π-bonds and also prevents rotation. An example of a molecule that has sp hybridized orbitals is acetylene (aka ethyne). HC≡CH.

Let’s Try a Practice Problem! Write the correct hybridization and indicate how many sigma and pi-bonds exist in a molecule of HCN (cyanic acid). First draw the Lewis structure Now determine the electron geometry: Linear Now select the correct hybridization of the central atom: Carbon is sp hybridized Now determine how many sigma and pi bonds exist. 2 sigma bonds and 2 pi bonds exist in the molecule above. One single bond exists between the H-C and C≡N triple bond consists of 1 sigma and two pi bonds.

10.5-10.7 pg. 476 #’s 62 (Write hybridization and number of sigma and pi bonds), 66 (Write hybridization of each interior atom and number of sigma and pi bonds) and 68. Read 10.8 pgs. 459-471