Liquids and Solids

Unit 04: BONDINGIB Topics 4 & 14Text: Ch 8 (all except sections 4,5 & 8)Ch 9.1 & 9.5Ch 10.1-10.7

My Name is Bond. Chemical Bond

1PART 3: Hybridization & Delocalization of Electrons

Hybridization Hybridization: a modification of the localized electron model to account for the observation that atoms often seem to use special atomic orbitals in forming molecules. This is part of both IB and AP curricula.

BeF2The VSEPR model predicts that this molecule is linear --- which of course it is.In fact, it has two identical Be-F bonds.

F Be - FBeF21s2s2pENERGYF Be - FBe1s22s2

OK, so where do the fluorine atoms bond?BeF21s2s2pENERGYexcitation1s2s2pF Be - FBe1s22s2BeF21s2s2pENERGYexcitation1s2s2phybridizationtwo sp hybrid orbitalsF Be - FBe1s22s22pBeF2 sp hybridization

sp hybrid orbitals

BF31s2s2pENERGYexcitation1s2s2phybridizationthreesp2 hybrid orbitalsB1s22s22p12p

BF3 sp2 hybridization

sp2 hybrid orbitals

CH41s2s2pENERGYexcitation1s2s2phybridizationfoursp3 hybrid orbitalsC1s22s22p2

CH4 sp3 hybridization

CH4 sp3 hybridizationsp3 hybrid orbitals

sp3 hybrid orbitals

H2O1s2s2pENERGYhybridizationfoursp3 hybrid orbitalsO1s22s22p4lonepairsavailable for bonding

H2O sp3 hybridization

What about hybridization involving d orbitals?PF53s3pENERGYexcitationhybridizationfivesp3d hybrid orbitalsP1s22s22p63s23p3To simplify things, only draw valence electrons3d3s3p3d

PF5 sp3d hybridization

3sp3d hybrid orbitalsNH31s2s2pENERGYhybridizationfoursp3 hybrid orbitalsN1s22s22p3lonepairavailable for bonding

NH3 sp3 hybridization

Something to think about: is hybridization a real process or simply a mathematical device (a human construction) weve concocted to explain how electrons interact when new chemical substances are formed?Valence electron pair geometry# of orbitalsHybrid orbitalsElectron density diagramExamples

Linear

2

Trigonal planar

3

Tetrahedral

4

Trigonal bipyramidal

5

Octahedral

6

spsp2sp3sp3dsp3d2BF2HgCl2CO2BF3SO3CH4H2O NH4+PF5SF4BrF3SF6XeF4 PF6- and bondsIn Hybridization Theory there are two names for bonds, sigma () and pi ().

Sigma bonds are the primary bonds used to covalently attach atoms to each other.

Pi bonds are used to provide the extra electrons needed to fulfill octet requirements.

and bondsEvery pair of bonded atoms shares one or more pairs of electrons. In every bond at least one pair of electrons is localized in the space between the atoms, in a sigma () bond.The electrons in a sigma bond are localized in the region between two bonded atoms and do not make a significant contribution to the bonding between any other atoms.

and bondsIn almost all cases, single bonds are sigma () bonds. A double bond consists of one sigma and one pi () bond, and a triple bond consists of one sigma and two pi bonds.Examples:

H HC CH HH H:N N:One bondOne bond and one bond.One bond and two bonds. bondsA Sigma bond is a bond formed by the overlap of two hybrid orbitals through areas of maximum electron density. This corresponds to the orbitals combining at the tips of the lobes in the orbitals.

bondsA Pi bond is a bond formed by the overlap of two unhybridized, parallel p orbitals through areas of low electron density. This corresponds to the orbitals combining at the sides of the lobes and places stringent geometric requirements on the arrangement of the atoms in space in order to establish the parallel qualities that are essential for bonding.

Remember bonds are unhybridized

strawberry pierhubarb piestrawberry-rhubarb pieXBond StrengthSigma bonds are stronger than pi bonds.

A sigma plus a pi bond is stronger than a sigma bond. Thus, a double bond is stronger than a single bond, but not twice as strong.

and bondsWhen atoms share more than one pair of electrons, the additional pairs are in pi () bonds. The centers of charge density in a () is above and below (parallel to) the bond axis.

Ethene: C2H4

35Ethyne: C2H2H C C - H

36Delocalized ElectronsMolecules with two or more resonance structures can have bonds that extend over more than two bonded atoms. Electrons in pi () bonds that extend over more than two atoms are said to be delocalized. Example: Benzene (C6H6)

Example: Benzene bonds (12) electrons in sp2 hybridized orbitals bonds (3) electrons in unhybridized p-orbitals

Close enough to overlapDelocalization of ElectronsDelocalization is a characteristic of electrons in pi bonds when theres more than one possible position for a double bond within the molecule.

Example: ozone (O3)These two drawn structures are known as resonance structures.

Example: ozone (O3)They are extreme forms of the true structure, which lies somewhere between the two.

Evidence that this is true comes from bond lengths, as the bond lengths for oxygen atoms in ozone are both the same and are an intermediates between an O=O double bond and an O-O single bond.

Example: ozone (O3)Resonance structures are usually drawn with a double headed arrow between them.

Note that benzene (C6H6) has six delocalized electrons. Since the p-orbitals overlap (forming three pi bonds, every-other-bond around the ring) all six electrons involved in pi bonding are free to move about the entire carbon ring.

sigma bonding in benzene(sp2 hybrid orbitals)

p orbitals6 delocalized electronspi bonding in benzene(unhybridized p orbitals)Formal ChargeA concept know as formal charge can help us choose the most plausible Lewis structure where there are a number of possible structures. This is not part of the IB curriculum, but it is part of the AP curriculum. This theory certainly has its critics; however, it has been included in this section of the course as it may help you in determining the most likely structure.

Definition of formal charge:

Formal Charge# valence es on the free atom# valence es assigned to the atom in the structureRules Governing Formal ChargeTo calculate the formal charge on an atom:Take the sum of the lone pair electrons and one-half the shared electrons. This is the number of valence electrons assigned to the atom in the molecule.Subtract the number of assigned electrons from the number of valence electrons on the free, neutral atom to obtain formal charge.The sum of the formal charges of all atoms in a given molecule or ion must equal the overall charge on that species.If nonequivalent Lewis structures exist for a species, those with formal charges closest to zero and with any negative formal charges on the most electronegative atoms are considered to best describe the bonding in the molecule or ion.

Example: CO2Possible Lewis structures of carbon dioxide:

O = C = O :O C O:.. .. .. .. .. .. Valence e- 6 4 6 6 4 6(e- assigned to atom)6 4 6 7 4 5Formal Charge0 0 0 -1 0 +1Example: NCO-For example if we look at the cyanate ion, NCO-, we see that it is possible to write for the skeletal structure, NOC-, CNO-, or CON-. Using formal charge we can choose the most plausible of these three Lewis structures.

Example: NCO-Find formal charge

Valance Electrons546# electrons assigned to atom646-10 0

Example: NCO-Find formal charge

Valance Electrons456# electrons assigned to atom646-2+1 0

Example: NCO-Find formal charge

Valance Electrons465# electrons assigned to atom666-20 -1

Example: NCO-Thus, the first structure is the most likely

-1 0 0-2 +2 -1-2 +1 0