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Maximum number of bonds permitted using valence shell orbitals:

Maximum number of bonds permitted using valence shell orbitals:. N is positive to allow for overall positive charge. Could have made any atom positive. 14/2 = 7 bonds. Set-up single bond framework . 6 bonds used in single bond framework.

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Maximum number of bonds permitted using valence shell orbitals:

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  1. Maximum number of bonds permitted using valence shell orbitals: N is positive to allow for overall positive charge. Could have made any atom positive. 14/2 = 7 bonds

  2. Set-up single bond framework. 6 bonds used in single bond framework. 1 left to assign. It will be a pi bond, part of a double bond. The additional bond is used to provide maximum stabilization; completing the octet. Where does it go? Only two possibilities…. How are they interconverted?

  3. So what is the third structure? It won’t have the double bond. Will not obey the octet rule. Can collapse the pi bond into a lone pair. Could we have collapsed in other direction? No lone pair on this N. High energy structure; not very important.

  4. Our three structures….. Does not obey octet rule. Least stable, least important. Intermediate stable with positive charge on more electronegative oxygen. Most stable with positive charge on the electropositive atom, N.

  5. Examine one of the interconversions shown earlier. Looking at it in reverse. Higher energy structure Lower energy, one more bond. This is a common and significant resonance motion. A supply of electrons moving towards a place where needed and providing a new bond in the process.

  6. Exceptions to Octet Rule But the S violates the octet rule by having more than 8 electrons around the S. Valence shell is 3s and 3p. S now would have to use additional atomic orbitals This means S would have to use the 3d shell orbitals. Luckily, they are low in energy and can be utilized. Octet is expanded due to use of 3d orbitals.

  7. Donation from O to S. Observe that this is another example of donation of available electrons (lone pair) into a neighboring acceptor orbital. Empty 3d  Filled 2p

  8. Survey of Classes of Molecules • Hydrocarbons, C & H A Alkanes: only single bonds • Acyclic (no rings) Alkanes: CnH2n+2 suffix: -ane 2. Cyclic alkanes: example cyclohexane CnH2n+2-2(# rings) Presence of a ring eliminates two hydrogens

  9. Alkenes: double bonds, pi bonds • Suffix: -ene i acyclic ii cycloalkenes A pi bond eliminates two hydrogens C4H10 (butane)  C4H8 (butene) + 2 H • general formula: CnH2n+2-2(# rings) -2(#pi bonds)

  10. Alkynes: C to C triple bond = 2 pi bonds + sigma

  11. Oxygen Containing Molecules Incorporation of an oxygen into a molecule does not change the relationship between the number of carbons and number of hydrogens. CnH2n+2-2(#pi bonds) – 2(# rings)Oa Watch how this makes sense.

  12. Hydrogen Deficiency • - 2(#pi bonds) – 2(# rings) CnH2n+2 hydrogen deficiency. 2n + 2: Number of hydrogens for an acyclic compound with n carbons and no pi bonds. If a compound has a different number of hydrogens it must be due to the presence of pi bonds or rings. # pi bonds + # rings = 1/2 ( (2n + 2) - (actual number of hydrogens present) )

  13. Singly bonded Oxygen Alcohols, ROH, -ol Two OH groups: a diol

  14. Useful Classiification of Alcohols (and other things): Primary, Secondary, Tertiary Alcohols and many other groups may be classified as Primary: one carbon directly bonded to the C bearing the –OH. CH3CH2CH2OH Secondary: two carbons directly bonded to the C bearing the –OH. (CH3CH2)2CHOH Tertiary: three carbons directly bonded to the C bearing the –OH. (CH3CH2)3COH 1o, 2o, 3o

  15. Ethers: ROR, isomeric with alcohols Methoxy group (alkoxy)

  16. Doubly Bonded Oxygen Carbonyl Group This is a pi bond, resulting in the elimination of two hydrogens.

  17. Aldehydes, RCHO, -al 4 3 2 1 Still more complex, now we put a double bond in the parent chain. A more complex example Start numbering the chain at CHO Double Bonds part of Parent Chain 2-methylbutanal

  18. Ketones, R(CO)R, -one Both carbonyls must be included in parent chain –dione. Butyl Side chain. Simple example More complex example: Double bond included as part of parent chain 1

  19. Carboxylic Acids, RCO2H, -oic acid Another more complex problem Methyl group Both carboxylic groups must be in parent chain C O H C O H Ethyl Group 2 2 Vinyl side chain Start numbering here

  20. Attached to or part of the Parent Chain are: • Functional Groups: -OH, carbonyl (aldehydes, ketones), carboxylic acids • Substituents: alkyl groups, halogens, alkoxy groups

  21. Some Simple Substituents • Alkyl, -R: An alkane minus one hydrogen (providing the point of attachment to the parent chain). -CH2CH3: ethyl (-Et); -CH2CH2CH3: propyl; -CH(CH3)2: isopropyl • Alkoxy, -OR. –OCH2CH3: ethoxy (-OEt) • Halo, -X. –Cl: chloro

  22. More complex substituents Systematically named. Note that the atom directly attached to the parent chain is atom 1 in the substituent. Substituents on the butyl substituent!! 1 3 2 Substituent 4 6 Parent chain 6-(2-(chloromethyl)-4-methoxy-1-methylbutyl) Indicates position of attachment on parent chain

  23. Functional Groups Non-Functional Groups (alkyl, alkoxy, halide) will always appear as substituent and as a prefix. Functional Groups (pi bonds, alcohol, aldehyde, ketone, carboxylic acid) will determine the suffix of parent chain. If more than one functional group then we must prioritize. Highest priority determines suffix. The rest ( alcohol, aldehyde, ketone, carboxylic acid) appear a prefix. But note unsaturation (pi bonding) in the parent chain is always specified as a suffix.

  24. Functional Group Priorities Highest priority at bottom 1

  25. 2-(but-2-enyl)-3-(1-hydroxyethyl)-4-(2-oxoethyl)pent-2-enedioic acid One more example. 1 2 4 3 5 Highest Priority Functional Group: two carboxylic acids. Pentanedioic acid Double bond in parent chain: Pent-2-enedioic acid Ethyl side chain bearing a hydroxy at position 1: 3-(1-hydroxyethyl) Ethyl side chain bearing an oxo at position 2: 4-(2-oxoethyl) Butyl side chain with double bond at position 2: 2-(but-2-enyl)

  26. Polar Bonds and Electronegativity H Electronegativity increases towards the upper right. Electronegativity differences results in polar bonds. Example:

  27. Some Examples Polar Bonds can lead to Polar Molecules Combine the bond dipoles to get the overall dipole by means of vector addition. Bond dipoles are put tail to head to provide an overall result. Resultant Vector Combine these three vectors

  28. Predicting Molecular dipoles Clearly, molecular geometry is very important in determining molecular dipoles. Problem: Rank the following for size of molecular dipole. Bond angles are 120o. A > C > B(0)

  29. Molecular geometry (VSEPR) The number of groups of electrons around an atom determines the shape. Each lone pair, single, double, or triple bond counts as a group.

  30. Example 1 • CH4 • Four bonds = Four groups • Tetrahedral

  31. Two lone pairs Two bonds Example 2 Four groups of electrons. Tetrahedral geometry for groups. Bent Molecule

  32. Nature of Chemical Bonds So far, we have the octet rule which tells us how many bonds we can make. But how do we understand the nature of the bonds? Three models for bonding: ionic, valence bond, molecular orbitals.

  33. Ionic Bonding Requires very different electronegativities to make the complete transfer of electrons worthwhile. Not discussed further although it occurs in salts of organic acids. For example sodium acetate.

  34. Quantum or Wave Mechanics • Albert Einstein: E = hn (energy is quantized) • light has particle properties. • Erwin Schrödinger: wave equation • wave function, : A solution to a set of equations that depicts the energy of an electron in an atom. • each wave function is associated with a unique set of quantum numbers. • each wave function represents a region of three-dimensional space and is called an orbital. •  2 is the probability of finding an electron at a given point in space.

  35. Quantum or Wave Mechanics • Characteristics of a wave associated with a moving particle. Wavelength is designated by the symbol l .

  36. Quantum or Wave Mechanics • When we describe orbital interactions, we are referring to interactions of waves. Waves interact • constructively or • destructively. • When two waves overlap, if they are of the same sign then they combine constructively, build-up. Opposite sign overlap combines destructively, meaning they cancel.

  37. Shapes of Atomic s and p Orbitals • All s orbitals have the shape of a sphere with the center of the sphere at the nucleus. • Figure 1.8 (a) Calculated and (b) cartoon representations showing an arbitrary boundary surface containing about 95% of the electron density.

  38. Shapes of Atomic s and p Orbitals • Three-dimensional representations of the 2px, 2py, and 2pz atomic orbitals. Nodal planes are shaded.

  39. Shapes of Atomic s and p Orbitals 2px, 2py, and 2pz atomic orbitals.

  40. Molecular Orbital Theory • MO theory begins with the hypothesis that • electrons in atoms exist in atomic orbitals and • electrons in molecules exist in molecular orbitals.

  41. Molecular Orbital Theory • Rules: • Combination of n atomic orbitals (mathematically adding and subtracting wave functions) gives n MOs (new wave functions). • MOs are arranged in order of increasing energy. • MO filling is governed by the same rules as for atomic orbitals: • Aufbau principle: fill beginning with lowest energy orbital • Pauli exclusion principle: no more than 2e- in a MO • Hund’s rule: when two or more MOs of equivalent energy (degenerate) are available, add 1e- to each before filling any one of them with 2e-.

  42. Molecular Orbital Theory • MOs derived from combination by (a) addition and (b) subtraction of two 1s atomic orbitals.

  43. Covalent Bonding • Bonding molecular orbital: A MO in which electrons have a lower energy than they would have in isolated atomic orbitals. • Sigma (s) bonding molecular orbital: A MO in which electron density is concentrated between two nuclei along the axis joining them and is cylindrically symmetrical.

  44. Covalent Bonding • A MO energy diagram for H2. (a) Ground state and (b) lowest excited state.

  45. Covalent Bonding • Antibonding MO: A MO in which electrons have a higher energy than they would in isolated atomic orbitals.

  46. VB: sp3 Hybridization of Atomic Orbitals • The number of hybrid orbitals formed is equal to the number of atomic orbitals combined. • Elements of the 2nd period form three types of hybrid orbitals, designated sp3, sp2, and sp. • The mathematical combination of one 2s atomic orbital and three 2p atomic orbitals forms four equivalent sp3 hybrid orbitals.

  47. VB: sp3 Hybridization of Atomic Orbitals • sp3 Hybrid orbitals. (a) Computed and (b) cartoon three-dimensional representations. (c) Four balloons of similar size and shape tied together, will assume a tetrahedral geometry.

  48. VB: sp2 Hybridization of Atomic Orbitals • The mathematical combination of one 2s atomic orbital wave function and two 2p atomic orbital wave functions forms three equivalent sp2 hybrid orbitals.

  49. VB: sp2 Hybridization of Atomic Orbitals • Hybrid orbitals and a single 2p orbital on an sp2 hybridized atom.

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