Orbital Bonding Theory Review Orbital 3 dimensional space

Orbital Bonding Theory

�Review: �Orbital: 3 dimensional space 1 or 2 electrons occupy around the nucleus of an atom Basic Shapes:

Properties of Waves Frequency - the number of waves per second Wavelength - distance between each wave Amplitude - size of each wave above (+) or below (-) the axis Node – point where the amplitude = zero

Wave Interactions Reinforcement and Interference of Waves Reinforcement: overlapping waves that are in phase (both + or both -)that increase the amplitude Interference: out of phase waves (one + and one -) that reduce the amplitude of both waves

Relationship to Bonding �For stable bonding to occur overlapping orbitals must be IN PHASE giving rise to reinforcement and a boding molecular orbital. �Out of phase orbitals result in interference and an antibonding molecular orbital.

�Example: H + H = H 2 �Each atom contains one electron in the 1 s orbital. �As the atoms approach each other the orbitals overlap in a reinforcing pattern and form a bonding molecular orbital that encompasses both nuclei and contains the two electrons (a shared pair)

�The result of the reinforcing orbitals is an area of high electron density between the nuclei. This prevents the positively charged nuclei from repelling each other and the atoms stay bonded together. �This molecular orbital is cylindrically symmetrical around the axis (it can rotate without breaking) between the two hydrogen atoms and is called a sigma (σ) molecular orbital or a sigma (σ) bond.

�Bonding of Out of Phase Orbitals �Out of phase orbitals, when they interact, produce a node of low electron density between the nuclei. �Since there is little negative charge due to the lack of electron density, the positive nuclei will repel each other resulting in a higher energy, less stable bond. �This is called an antibonding orbital or a sigma star (σ*) bond

Sigma bond formation in p sublevels �p sublevels can also form sigma bonds (and antisigma bonds) when their orbitals align and overlap in a horizontal axis

Carbon Orbital Hybridization �Carbon: 1 s 2 2 p 2 �The 1 st level electrons are not used for bonding because the energy level is full and stable. �The second energy level has four electrons, and thus, it seems, would be able to bond four times to fulfill the octet rule. �However, the two electrons in the 2 s sublevel are fairly stable and would not normally bond. �Thus the orbitals can hybridize or blend the 2 s orbital and the three 2 p orbitals into sp orbitals

Carbon Orbital Hybrids �sp 3: when carbon forms 4 single bonds �sp 2: when carbon forms 1 double bond called a pi bond �sp: when a carbon forms 1 triple bond or 2 double bonds on the same carbon

sp 3 bonding Hybridization of the 2 s orbital and the three 2 p orbitals result in four equal orbitals called sp 3 orbitals. This spreads the four valence electrons around the nucleus equally and allows for the carbon to bond four times forming the tetrahedron with the equivalent bond angles of 109. 5 o.

Area used for bonding

sp 3 bonding �As the sp 3 orbitals approach other orbitals (s, sp 3), they will form sigma bonds when they overlap. �If with an s = sp 3 -s sigma bond �If with another sp 3 = sp 3 -sp 3 sigma bond �These are all single bonds and can rotate freely around the axis of the bond.

sp 3 bonding in methane �The 1 s orbitals of hydrogen overlap with the sp 3 orbitals of the carbon to form 4 sp 3 -s sigma bonds.

sp 3 bonding in ethane C 2 H 6 �Each carbon forms 3 sp 3 -s sigma bonds with hydrogen. The last bond of each carbon forms between the two carbons making one sp 3 -sp 3 sigma bond. sp 3 -sp 3

sp 2 bonding �In this state carbon hybridizes two of its p orbitals with the 2 s orbital to create three sp 2 hybrids. �These three orbitals spread out evenly and form a triangular plane with the bond angles = 120 o. �This is known as a trigonal carbon or trigonal planar.

�The remaining p orbital does not hybridize but exists perpendicular to the trigonal planar carbon sp 2 orbitals.

�If two sp 2 carbons or another atom with a p orbital line up so the p orbitals form a double bond by overlapping side to side rather than end to end like in a sigma bond. �This bond is called a pi bond or a p bond.

Pi bond formation

�Bonding in Ethene (C 2 H 4) �Each of the two carbons form three sp 2 orbitals and have 1 unhybridized p orbital. �Each carbon binds with two hydrogens forming a total of 4 sp 2 -s bonds. �The carbons bond to each other two ways: � 1) the first bond is the sp 2 -sp 2 sigma bond between the remaining sp 2 orbitals � 2) the second is the pi bond between the unhybridized p orbitals.

Bonding in Ethene sp 2 -sp 2 σ bond

Characteristics of pi bonds 1) Are not cylindrically symmetrical like sigma bonds. -implication: these bonds cannot rotate or the bond will break. - locks side groups into a particular geometry (cis and trans)

2) pi bonds have more energy than sigma bonds and are less stable - implication: when a double bond breaks to become a single bond, it is the pi bond that breaks, not the sigma 3) pi bonds exist farther from the nucleus and are more exposed - implication: pi bonds are more “vulnerable” to attack by another atom or molecule making it the site of chemical reactivity = location of chemical reactions

Nucleophilic Addition Reaction

sp orbitals �Carbon hybridizes the 2 s and one 2 p orbital to form two sp orbitals. �The other 2 p orbitals remain unhybridized. �The sp orbitals will form sigma bonds and the 2 p orbitals will each form a pi bond - Result: triple bond formed from one sigma and 2 pi bonds.

Bonding in Ethyne �Ethyne = C 2 H 2 �Each carbon forms one bond with a hydrogen �Hydrogen s orbital and Carbon sp orbital = sp-s sigma bond. �Carbons bond to each other three times in two ways � 1) sp-sp sigma bond � 2) 2 pi bonds that are perpendicular to each other

Propadiene