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10 - 1 Molecular Shape and Theory of Chemical Bonding Shapes of Molecules and Polyatomic Ions Polar and Nonpolar Molecules Bonding Theory Molecular Orbital.

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Presentation on theme: "10 - 1 Molecular Shape and Theory of Chemical Bonding Shapes of Molecules and Polyatomic Ions Polar and Nonpolar Molecules Bonding Theory Molecular Orbital."— Presentation transcript:

1 Molecular Shape and Theory of Chemical Bonding Shapes of Molecules and Polyatomic Ions Polar and Nonpolar Molecules Bonding Theory Molecular Orbital Method Delocalized Electrons Band Theory of Bonding in Solids

2 Shapes 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. Valence Shell Electron Pair Repulsion model (VSEPR) One model is called the Valence Shell Electron Pair Repulsion model (VSEPR)

3 VSEPR model According 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.

4 VSEPR shapes Coordination Electron pairs General NumberBonding Unshared Formula Shape AB 2 Linear AB 3 Trigonal planar 2 1 AB 2 Bent AB 4 Tetrahedral 3 1 AB 3 Trigonal pyramidal 2 2 AB 2 Bent 1 3 AB Linear

5 Molecular geometry Molecules have specific shapes. Determined by the number of electron pairs around the central species Bonded and unshared pairs count. Multiple bonds are treated as a single bond for geometry. Geometry affects factors like polarity and solubility.

6 Some common geometries e - pairs around e - pairs around Shape central atom Example Linear 2 BeH 2 Trigonal planar 3 BF 3 Tetrahedral 4 CH 4 Trigonal pyramidal4 NH 3 Bent 4 H 2 O

7 Linear - CO 2

8 Trigonal planar, BCl 3

9 Bent, H 2 O

10 Pyramidal, NH 3

11 Tetrahedral, CH 4

12 Molecular geometries based on tetrahedral Bent and pyramidal are actually tetrahedral but some of the electron pairs are not bonded. H N HH Pyramidal H C HH H Tetrahedral Bent H O H

13 Other geometries. Other shapes are also observed. Five bonds or lone electron pairs Trigonal bipyramidal Seesaw T-shaped Linear Six bonds or lone electron pairs Octahedral Square pyramidal Square planar

14 VSEPR shapes Coordination Electron pairs General NumberBonding Unshared Formula Shape AB 5 Trigonal bipyramidal 4 1 AB 4 Seesaw 3 2 AB 3 T-shaped 2 3 AB 2 Linear AB 6 Octahedral 5 1 AB 5 Square pyramidal 4 2 AB 4 Square Planar

15 Trigonal bipyramidal

16 Square planar

17 Octahedral

18 Molecular geometry As molecules get larger, the rules regarding molecular geometry still hold. H C H H H C H H

19 Ethane

20 Geometry and polar molecules For a molecule to be polar - must have polar bonds - must have the proper geometry CH 4 non-polar CH 3 Clpolar CH 2 Cl 2 polar CHCl 3 polar CCl 4 non-polarWHY?

21 Polar 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 Dipole moment, , is a quantitative measure of the polarity of a molecule.

22 Dipole moment This property can be measured by placing molecules in an electrical field. Polar molecules will align when The field is on. Nonpolar molecules will not. +-+-

23 Polar 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. O = C = O Electronegativities: Oxygen = 3.5 Carbon = 2.5 Difference 1.0 (polar bond) Electronegativities: Oxygen = 3.5 Carbon = 2.5 Difference 1.0 (polar bond) The electronegativity values Show that the C-O bond would be polar with electrons Being pulled towards the oxygens. However, due to The geometry, the pull happens in equal and opposite directions.

24 Polar 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. Electronegativity difference = 1.3 Here, the effects of the polar bonds do not canceled so the molecule is polar. H O..

25 Polar and nonpolar molecules A molecule is nonpolar if the central atom is symmetrically substituted by identical atoms. CO 2, CH 4, CCl 4 A molecule will be polar if the geometry is not symmetrical. H 2 O, NH 3, CH 2 Cl 2 The degree of polarity is a function of the number and type of polar bonds as well as the geometry.

26 Bonding theory Two methods of approximation are used to describe bonding between atoms. Valence bond method Bonds are assumed to be formed by overlap of atomic orbitals Molecular orbital method molecular orbitals When atoms form compounds, their orbitals combine to form new orbitals - molecular orbitals.

27 Valence bond method According to this model, the H-H bond forms as a result of the overlap of the 1s orbitals from each atom. 74 pm

28 Valence bond method Hybrid orbitals are need to account for the geometry that we observe for many molecules. Example - Carbon Outer electron configuration of 2s 2 2p x 1 2p y 1 We know that carbon will form2 four equivalent bonds - CH 4, CH 2 Cl 2, CCl 4. The electron configuration appears to indicate that only two bonds would form and they would be at right angles -- not tetrahedral angles.

29 Hybridization To explain why carbon forms four identical single bonds, we assume the the original orbitals will blend together. Unhybridized Hybridized energy 2s2s 2p2p 2sp 3

30 Hybridization In the case of a carbon that has 4 single bonds, all of the orbitals are hybrids. sp 3 25% s and 75% p character + 3 s p sp 3 1 4

31 Ethane, CH 3 CH 3 1s orbital of H  bond sp 3 hybrids  bond  bond - formed by an endwise (head-on) overlap. Molecules are able to rotate around single bonds.  bond  bond - formed by an endwise (head-on) overlap. Molecules are able to rotate around single bonds.

32 Ethane, CH 3 CH 3

33 Rotation of single bond

34 Rotation of single bond

35 sp 2 hybrid orbitals To account for double bonds, a second type of hybrid orbital must be pictured. An sp 2 hybrid is produced by combining one s and 2 p orbitals. One p orbital remains. Unhybridized Hybridized energy 2s2s 2p2p 2sp 2 2p2p

36 sp 2 hybrid orbitals The unhybridized p orbitals are able to overlap, resulting in the formation of a second bond -  bond. A  bond is a sideways overlap that occurs both above and below the plane of the molecule Parts of the molecule are no longer able to rotate about the bond. A  bond is a sideways overlap that occurs both above and below the plane of the molecule Parts of the molecule are no longer able to rotate about the bond. C C

37 Ethene

38 Bonding in ethene 1s orbital  overlap  bond  bond  overlap sp 2 hybrids sp 2 hybrids

39 Bonding in ethene

40 sp hybrid orbital Forming a triple bond is also possible. This requires that two p orbitals remain unhybridized. Unhybridized Hybridized energy 2s2s 2p2p 2sp 2p2p

41 sp hybrid orbital Now two p orbitals are available to form  bonds.

42 Ethyne

43 Bonding in ethyne sp hybrid p overlaps

44 Bonding in ethyne

45 Other hybrid orbitals d orbitals can also be involved in the formation of hybrid orbitals. HybridShape spLinear sp 2 Trigonal planar sp 3 Tetrahedral sp 3 dTrigonal bipyramidal sp 3 d 2 Octahedral

46 Molecular 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 - H 2 Two 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.

47 H 2 molecular orbital diagram HH 1s1s 1s1s 1sH21sH2 1s1s Orbital shapes energy

48 Molecular orbitals When two atomic orbitals combine, three types of molecular orbitals are produced. Bonding orbital -  or  The energy is lower than the atomic orbitals and the electron density overlaps. Antibonding orbital -  * or   The energy is higher than the atomic orbitals and the electron density does not overlap. Nonbonding - n Electron pairs not involved in bonding.

49 Homonuclear 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 H 2. We can develop energy diagrams for a range of molecules or possible molecules to see if they bond and how.

50 MO diagram of helium He 1s1s 1s1s  1s He 2 1s1s energy If we develop a diagram for helium we see that both a bonding and antibonding orbital will be filled. The result is that it is no more stable than the unbonded form -- it will not bond

51 Molecular 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  and  bonds must be considered. Lets look at the MO diagram for O 2.

52 MO diagram for O 2 1s1s 2s2s 2p x 2p y 2p z 2s2s 1s1s  * 2pz  2pz 2s2s  2s  1s 1s1s  2px  2py  2px  2py

53 MO diagram for O 2 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. O 2 will form if we have more bonding than antibonding electrons.

54 MO diagram for O 2 1s1s 2s2s 2p x 2p y 2p z 2s2s 1s1s  * 2pz  2pz 2s2s  2s  1s 1s1s  2px  2py  2px  2py

55 Heteronuclear 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.

56 MO diagram for NO NO 1s1s 2s2s 2p x 2p y 2p z 2s2s 1s1s  * 2px  2px 2s2s  2s  1s 1s1s  2py  2pz  2py  2pz NO

57 Delocalized electrons MO diagrams for polyatomic species are often simplified by assuming that all  and some  orbitals are localized -- shared between two specific atoms. Resonance structures require that electrons in some  orbitals be pictured as delocalized. Delocalized Delocalized - free to move around three or more atoms.

58 Delocalized 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. =

59 Aromatic hydrocarbons p orbitals overlap sidewise all around the ring. No localized double bonds. H H HH HH

60 Band 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.’ 3s 9 Na 3s 9 Na

61 Band theory of bonding in solids Band A group of very closely spaced energy levels. Energy gap The difference in energy between the bonding and antibonding orbitals. Forbidden bands A ‘space’ that separates bands.

62 Band theory of bonding in solids Energy Internuclear distance s p The s and p bands of Group II (2) metals overlap. The s and p bands of Group II (2) metals overlap.

63 Band theory of bonding in solids Conductor A material with a partially filled energy band.Insulator The highest occupied band is filled or almost completely filled. The forbidden band just above the highest filled is wide.Semiconductor The gap between the highest filled band and the next higher permitted band is relatively narrow.

64 Band theory of bonding in solids insulator semiconductor conductor Energy Empty Forbidden, wide Filled Empty Forbidden, narrow Filled No Forbidden


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