Presentation on theme: "Molecular Shape and Theory of Chemical Bonding"— Presentation transcript:
1Molecular Shape and Theory of Chemical Bonding Shapes of Molecules and Polyatomic IonsPolar and Nonpolar MoleculesBonding TheoryMolecular Orbital MethodDelocalized ElectronsBand Theory of Bonding in Solids
2Shapes of molecules and polyatomic ions Molecules and polyatomic ions are not all ‘flat’ structures.Many have a three dimensional arrangement that helps account for their various chemical and physical properties.Several models are used to help predict and describe the geometries for these species.One model is called the Valence Shell Electron Pair Repulsion model (VSEPR)
3VSEPR modelAccording to this model, for main group elements, electron pairs will be as far apart from each other as possible.This occurs in three dimensional space.Both bonded and unshared pairs will occupy space with unshared pairs taking up more space.The geometry is based on the total number of electron pairs - total coordination number.
4VSEPR shapes 2 2 0 AB2 Linear 3 3 0 AB3 Trigonal planar 2 1 AB2 Bent Coordination Electron pairs GeneralNumber Bonding Unshared Formula ShapeAB LinearAB Trigonal planarAB BentAB TetrahedralAB Trigonal pyramidalAB BentAB Linear
5Molecular geometry Molecules have specific shapes. Determined by the number of electron pairs around the central speciesBonded and unshared pairs count.Multiple bonds are treated as a single bond for geometry.Geometry affects factors like polarity and solubility.
6Some common geometries e- pairs aroundShape central atom ExampleLinear BeH2Trigonal planar BF3Tetrahedral CH4Trigonal pyramidal NH3Bent H2O
20Geometry and polar molecules For a molecule to be polar- must have polar bonds- must have the proper geometryCH4 non-polarCH3Cl polarCH2Cl2 polarCHCl3 polarCCl4 non-polarWHY?
21Polar and nonpolar molecules Polarity is an important property of molecules.It affects physical properties such as melting point, boiling point and solubility.Chemical properties also depend on polarity.Dipole moment, m, is a quantitative measure of the polarity of a molecule.
22Dipole moment This property can be measured by placing molecules in an electrical field. Polar molecules will align whenThe field is on. Nonpolar molecules will not.+-+-
23Polar and nonpolar molecules Most bonds between atoms of dissimilar elements in a molecule are polar. That does not mean that the molecule will be polar.Electronegativities:Oxygen = 3.5Carbon = 2.5Difference 1.0(polar bond)O = C = OThe electronegativity valuesShow that the C-O bond would be polar with electronsBeing pulled towards the oxygens. However, due toThe geometry, the pull happens in equal and oppositedirections.
24Polar and nonpolar molecules For a molecule to be polar, the effects of bond polaritymust not cancel out.One way is to have a geometry that is not symmetrical like in water.H HO..Electronegativitydifference = 1.3Here, the effects of the polar bonds do notcanceled so the molecule is polar.
25Polar and nonpolar molecules A molecule is nonpolar if the central atom is symmetrically substituted by identical atoms.CO2, CH4 , CCl4A molecule will be polar if the geometry is not symmetrical.H2O, NH3, CH2Cl2The degree of polarity is a function of the number and type of polar bonds as well as the geometry.
26Bonding theoryTwo methods of approximation are used to describe bonding between atoms.Valence bond methodBonds are assumed to be formed by overlap of atomic orbitalsMolecular orbital methodWhen atoms form compounds, their orbitals combine to form new orbitals - molecular orbitals.
27Valence bond methodAccording to this model, the H-H bond forms as a result of the overlap of the 1s orbitals from each atom.74 pm
28Valence bond methodHybrid orbitals are need to account for the geometry that we observe for many molecules.Example - CarbonOuter electron configuration of 2s2 2px1 2py1We know that carbon will form2 four equivalent bonds - CH4, CH2Cl2 , CCl4.The electron configuration appears to indicate that only two bonds would form and they would be at right angles -- not tetrahedral angles.
29HybridizationTo explain why carbon forms four identical single bonds, we assume the the original orbitals will blend together.Unhybridized Hybridizedenergy2s2p2sp3
30HybridizationIn the case of a carbon that has 4 single bonds, all of the orbitals are hybrids.sp325% s and 75% p character1+ 34s p sp3
31Ethane, CH3CH3 bond sp3 hybrids s bond - formed by an endwise 1s orbitalof H bondsp3hybridss bond - formedby an endwise(head-on) overlap.Molecules areable to rotatearound singlebonds.
35sp2 hybrid orbitalsTo account for double bonds, a second type of hybrid orbital must be pictured. An sp2 hybrid is produced by combining one s and 2 p orbitals. One p orbital remains.2p2penergy2sp22sUnhybridized Hybridized
36sp2 hybrid orbitalsThe unhybridized p orbitals are able to overlap, resulting in the formation of a second bond - p bond.A p bond is asideways overlapthat occurs bothabove and below theplane of the moleculeParts of the moleculeare no longer able torotate about the bond.C
45Other hybrid orbitalsd orbitals can also be involved in the formation of hybrid orbitals.Hybrid Shapesp Linearsp2 Trigonal planarsp3 Tetrahedralsp3d Trigonal bipyramidalsp3d2 Octahedral
46Molecular Orbital Method When atomic orbitals combine to form molecular orbitals, the number of molecular orbitals formed must equal the number of atomic orbitals mathematically combined.Example - H2Two 1s orbitals will combine forming two molecular orbitals. The overall energy of the new orbitals is the same as the original two 1s. However, they will be at different energies.
47H2 molecular orbital diagram 1ss1sH2s*1senergyOrbital shapes
48Molecular orbitalsWhen two atomic orbitals combine, three types of molecular orbitals are produced.Bonding orbital - s or pThe energy is lower than the atomic orbitals and the electron density overlaps.Antibonding orbital - s* or p*The energy is higher than the atomic orbitals and the electron density does not overlap.Nonbonding - nElectron pairs not involved in bonding.
49Homonuclear diatomic molecules These molecules are simple diatomics where both atoms are of the same element.Energy diagrams for these types of molecules are similar to the one for H2.We can develop energy diagrams for a range of molecules or possible molecules to see if they bond and how.
50MO diagram of helium s*1s s1s If we develop a diagram for helium we see that botha bonding andantibondingorbital will befilled.The result is thatit is no more stablethan the unbondedform -- it will notbondHe1ss1sHe2s*1senergy
51Molecular orbital bonding For a molecule to be stable, you must have more electrons in bonding orbitals than in antibonding orbitals.The bonded form will be at a lower energy so will be more stable.Bonding and antibonding orbitals for both s and p bonds must be considered.Lets look at the MO diagram for O2.
53MO diagram for O2 Each oxygen atom has 8 electrons for a total of 16. We can now place 16 electrons into the MO diagram and see what happens.Remember, don’t pair electrons unless you need to and fill a lower energy orbital before proceeding to the next higher one.O2 will form if we have more bonding than antibonding electrons.
55Heteronuclear diatomic molecules Molecular orbital diagrams become more complex when bonding between two nonidentical atoms is considered.The atomic energy levels are not the same and there are differing numbers of electrons.A simple example is NO where the orbitals are similar but not identical.
56MO diagram for NO 1s 2s 2px 2py 2pz s*2px s2px s2s s*2s s*1s s1s p2py p2pzp*2pyp*2pzNNOO
57Delocalized electrons MO diagrams for polyatomic species are often simplified by assuming that all s and some p orbitals are localized -- shared between two specific atoms.Resonance structures require that electrons in some p orbitals be pictured as delocalized.Delocalized - free to move around three or more atoms.
58Delocalized electrons Benzene is a good example of delocalized electrons.We know that the bonding between carbons has an order of 1.5 and that all of the bonds are equal.=
59Aromatic hydrocarbons p orbitals overlapsidewise all aroundthe ring. No localizeddouble bonds.
60Band theory of bonding in solids This is an extension of delocalized orbitals.Each atom interacts with all of the others in the crystal, resulting in an enormous number of ‘molecular orbitals.’3s9 Na3s9 Na
61Band theory of bonding in solids A group of very closely spaced energy levels.Energy gapThe difference in energy between the bonding and antibonding orbitals.Forbidden bandsA ‘space’ that separates bands.
62Band theory of bonding in solids The s and p bands ofGroup II (2)metals overlap.pEnergysInternuclear distance
63Band theory of bonding in solids ConductorA material with a partially filled energy band.InsulatorThe highest occupied band is filled or almost completely filled. The forbidden band just above the highest filled is wide.SemiconductorThe gap between the highest filled band and the next higher permitted band is relatively narrow.
64Band theory of bonding in solids EmptyEnergyForbidden, wideFilledinsulatorEmptyEnergyForbidden, narrowFilledsemiconductorEnergyNoForbiddenconductor