Chapter1 structure and bonding


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Chapter1 structure and bonding

  2. 2. COURSE SYNOPSIS• This course provides a chemicalbackground of sufficient depth to facilitatean understanding of the organic chemicalprocesses, which occur in industry. Topicscovered include organic nomenclature,reaction types and biomolecules.
  3. 3. COURSE OBJECTIVESUpon completion of this course, students shouldbe able to:1) Apply the concept of bonding between atomsin organic molecules2) Identify, name and describe the reactions oforganic compounds based upon theirfunctional activity3) Analyze chemical reactions and proposepossible chemical reaction mechanism
  4. 4. TEXTBOOK/REFERENCES1) Mc Murry, J.,2007. Organic Chemistry. 6thed. Belmont: Brooks Cole2) Solomons, T.W.G., 2008. OrganicChemistry. 9thed. Wiley3) Silberberg, M., 2008. Chemistry: TheMolecular Nature of Matter and Change.5thed. New York: McGraw-Hill.
  5. 5. CHAPTERSChapter 1 : Structure and BondingChapter 2 : Alkanes, Alkenes and AlkynesChapter 3 : Organic Reaction TypesChapter 4 : OrganohalidesChapter 5 : Benzene and AromaticityChapter 6 : Alcohols and CarbonylsChapter 7 : BiomoleculesChapter 8 : Polymers
  7. 7. Chapter 1:Strucutre & Bonding1.1 Atomic Structure: The Nucleus1.2 Atomic Structure: Orbitals1.3 Atomic Structure: Electron Configuration1.4 Development of chemical bonding theory1.5 The nature of chemical bonds: Valence bond theory1.6 sp3 Hybrid Orbitals1.7 sp2 Hybrid Orbitals1.8 sp Hybrid Orbitals1.9 Hybridization of Nitrogen, Oxygen, Phosphorus &Sulfur1.10 The nature of chemical bonds:Molecular orbital theorySummary
  8. 8. 1.1 Atomic Structure• Structure of an atom– Positively charged nucleus (very dense,protons and neutrons) and small (10-15m)– Negatively charged electrons are in a cloud(10-10m) around nucleus• Diameter is about 2 × 10-10m (200 picometers(pm)) [the unit angstrom (Å) is 10-10m = 100 pm]
  9. 9. Atomic Number and Atomic Mass• The atomic number (Z) is the number of protonsin the atoms nucleus• The mass number (A) is the number of protonsplus neutrons• All the atoms of a given element have the sameatomic number• Isotopes are atoms of the same element thathave different numbers of neutrons andtherefore different mass numbers• The atomic mass (atomic weight) of an elementis the weighted average mass in atomic massunits (amu) of an element’s naturally occurringisotopes
  10. 10. 1.2 Atomic Structure: Orbitals• Quantum mechanics: describes electronenergies and locations by a wave equation– Wave function solution of wave equation– Each wave function is an orbital,Ψ• A plot of Ψ 2describes where electron most likelyto be• Electron cloud has no specific boundary so weshow most probable area
  11. 11. Shapes of Atomic Orbitals forElectrons• Four different kinds of orbitals for electronsbased on those derived for a hydrogen atom• Denoted s, p, d, and f• s and p orbitals most important in organic andbiological chemistry• s orbitals: spherical, nucleus at center• p orbitals: dumbbell-shaped, nucleus at middle• d orbitals: elongated dumbbell-shaped, nucleusat center
  12. 12. p-Orbitals• In each shell there are threeperpendicular p orbitals, px, py, and pz,of equal energy• Lobes of a p orbital are separated byregion of zero electron density, anode
  13. 13. Orbitals and Shells• Orbitals are grouped in shells of increasing sizeand energy• Different shells contain different numbers andkinds of orbitals• Each orbital can be occupied by two electrons• First shell contains one s orbital, denoted 1s,holds only two electrons• Second shell contains one s orbital (2s) andthree p orbitals (2p), eight electrons• Third shell contains an s orbital (3s), three porbitals (3p), and five d orbitals (3d), 18electrons
  14. 14. 1.3 Atomic Structure: ElectronConfigurations• Ground-state electron configuration (lowestenergy arrangement) of an atom lists orbitalsoccupied by its electrons.• Rules:• 1. Lowest-energy orbitals fill first: 1s → 2s → 2p→ 3s → 3p → 4s → 3d (Aufbau (“build-up”)principle)
  15. 15. • 2. Electrons act as if they were spinning aroundan axis. Electron spin can have only twoorientations, up ↑ and down ↓. Only twoelectrons can occupy an orbital, and they mustbe of opposite spin (Pauli exclusion principle) tohave unique wave equations• 3. If two or more empty orbitals of equal energyare available, electrons occupy each with spinsparallel until all orbitals have one electron(Hunds rule).
  16. 16. HUND’S RULE• Oxygen =82ndShell 2p __ __ __2s __1stShell 1s __
  17. 17. • Atoms form bonds because the compound thatresults is more stable than the separate atoms• Ionic bonds in salts form as a result of electrontransfers (example: Na+Cl-)• Organic compounds have covalent bonds fromsharing electrons (G. N. Lewis, 1916)1.4 Development of ChemicalBonding Theory
  18. 18. • Kekulé and Couper independently observed thatcarbon always has four bonds• vant Hoff and Le Bel proposed that the fourbonds of carbon have specific spatial directions– Atoms surround carbon as corners of atetrahedron
  19. 19. 1) Lewis structures (electron dot) show valenceelectrons of an atom as dots– Hydrogen has one dot, representing its 1selectron– Carbon has four dots (2s22p2)Method of indicating covalent bonds
  20. 20. 2) Kekule structures (line-bond structures)have a line drawn between two atomsindicating a 2 electron covalent bond.
  21. 21. Note that:• Stable molecule results at completed shell, octet(eight dots) for main-group atoms (two forhydrogen)• Atoms with one, two, or three valence electronsform one, two, or three bonds.• Atoms with four or more valence electrons formas many bonds as they need electrons to fill thes and p levels of their valence shells to reach astable octet.• Carbon has four valence electrons (2s22p2),forming four bonds (CH4).
  22. 22. • Nitrogen has five valence electrons (2s22p3)but forms only three bonds (NH3).• Oxygen has six valence electrons (2s22p4) butforms two bonds (H2O)
  23. 23. Non-bonding electrons• Valence electrons not used in bonding are callednonbonding electrons, or lone-pair electrons– Nitrogen atom in ammonia (NH3)• Shares six valence electrons in three covalentbonds and remaining two valence electrons arenonbonding lone pair
  24. 24. Drawing Structures• Drawing every bond in organic moleculecan become tedious.• Several shorthand methods have beendeveloped to write structures.1)Condensed structures don’t have C-H orC-C single bonds shown. They areunderstood.2)Skeletal structures
  25. 25. 3 General Rules:1) Carbon atoms aren’t usually shown. Instead acarbon atom is assumed to be at eachintersection of two lines (bonds) and at the endof each line.2) Hydrogen atoms bonded to carbon aren’tshown.3) Atoms other than carbon and hydrogen areshown.Skeletal structure
  26. 26. 1.5 The Nature of ChemicalBonds: Valence Bond Theory• Covalent bond forms when two atomsapproach each other closely so that asingly occupied orbital on one atomoverlaps a singly occupied orbital on theother atom• Two models to describe covalentbonding.1)Valence bond theory2)Molecular orbital theory
  27. 27. Valence Bond TheoryElectrons are paired in the overlappingorbitals and are attracted to nuclei of bothatoms– H–H bond results from the overlap of twosingly occupied hydrogen 1s orbitals– H-H bond is cylindrically symmetrical,sigma (σ) bond
  28. 28. Bond Energy• Reaction 2 H· → H2 releases 436 kJ/mol• Product has 436 kJ/mol less energy than twoatoms: H–H has bond strength of 436 kJ/mol.(1 kJ = 0.2390 kcal; 1 kcal = 4.184 kJ)
  29. 29. Bond Length• Distance betweennuclei that leadsto maximumstability• If too close, theyrepel becauseboth arepositively charged• If too far apart,bonding is weak
  30. 30. 1.6 sp3Orbitals• Carbon has 4 valence electrons (2s22p2)• In CH4, all C–H bonds are identical (tetrahedral)• sp3hybrid orbitals: s orbital and three p orbitalscombine to form four equivalent, unsymmetrical,tetrahedral orbitals (sppp = sp3), Pauling (1931)
  31. 31. The Structure of Methane• sp3orbitals on C overlap with 1s orbitals on 4 Hatoms to form four identical C-H bonds• Each C–H bond has a strength of 436 (438)kJ/mol and length of 109 pm• Bond angle: each H–C–H is 109.5°, thetetrahedral angle.
  32. 32. The Structure of Ethane• Two C’s bond to each other by σ overlap of ansp3orbital from each• Three sp3orbitals on each C overlap with H 1sorbitals to form six C–H bonds• C–H bond strength in ethane 423 kJ/mol• C–C bond is 154 pm long and strength is 376kJ/mol• All bond angles of ethane are tetrahedral
  33. 33. 1.7 sp2Orbitals• sp2hybrid orbitals: 2s orbital combines withtwo 2p orbitals, giving 3 orbitals (spp = sp2). Thisresults in a double bond.• sp2orbitals are in a plane with120° angles• Remaining p orbital is perpendicular to the plane
  34. 34. Bonds From sp2Hybrid Orbitals• Two sp2-hybridized orbitals overlap to form a σbond• p orbitals overlap side-to-side to formation a pi(π) bond• sp2–sp2σ bond and 2p–2p π bond result insharing four electrons and formation of C-Cdouble bond• Electrons in the σ bond are centered betweennuclei• Electrons in the π bond occupy regions are oneither side of a line between nuclei
  35. 35. Structure of Ethylene• H atoms form σ bonds with four sp2orbitals• H–C–H and H–C–C bond angles of about 120°• C–C double bond in ethylene shorter andstronger than single bond in ethane• Ethylene C=C bond length 134 pm (C–C 154pm)
  36. 36. 1.8 sp Orbitals• C-C a triple bond sharing six electrons• Carbon 2s orbital hybridizes with a single porbital giving two sp hybrids– two p orbitals remain unchanged• sp orbitals are linear, 180° apart on x-axis• Two p orbitals are perpendicular on the y-axisand the z-axis
  37. 37. Orbitals of Acetylene• Two sp hybrid orbitals from each C form sp–sp σbond• pz orbitals from each C form a pz–pz π bond bysideways overlap and py orbitals overlap similarly
  38. 38. Bonding in Acetylene• Sharing of six electrons forms C ≡C• Two sp orbitals form σ bonds with hydrogens
  39. 39. Comparison of C-C and C-H bonds
  40. 40. 1.9 Hybridization of Nitrogen andOxygen• Elements other than C can have hybridizedorbitals• H–N–H bond angle in ammonia (NH3) 107.3°• C-N-H bond angle is 110.3 °• N’s orbitals (sppp) hybridize to form four sp3orbitals• One sp3orbital is occupied by two nonbondingelectrons, and three sp3orbitals have oneelectron each, forming bonds to H and CH3.
  41. 41. 1.10 The nature of chemicalbonds: Molecular Orbital Theory• A molecular orbital (MO): where electrons aremost likely to be found (specific energy and generalshape) in a molecule• Additive combination (bonding) MO is lower inenergy• Subtractive combination (antibonding) MO is higherenergy
  42. 42. Molecular Orbitals in Ethylene• The π bonding MO is from combining p orbitallobes with the same algebraic sign• The π antibonding MO is from combining lobeswith opposite signs• Only bonding MO is occupied
  43. 43. Summary• Organic chemistry – chemistry of carbon compounds• Atom: positively charged nucleus surrounded bynegatively charged electrons• Electronic structure of an atom described by waveequation– Electrons occupy orbitals around the nucleus.– Different orbitals have different energy levels and differentshapes• s orbitals are spherical, p orbitals are dumbbell-shaped• Covalent bonds - electron pair is shared betweenatoms
  44. 44. • Valence bond theory - electron sharing occurs byoverlap of two atomic orbitals• Molecular orbital (MO) theory, - bonds result fromcombination of atomic orbitals to give molecular orbitals,which belong to the entire molecule• Sigma (σ) bonds - Circular cross-section and areformed by head-on interaction• Pi (π) bonds – “dumbbell” shape from sidewaysinteraction of p orbitals
  45. 45. • Carbon uses hybrid orbitals to form bonds in organicmolecules.– In single bonds with tetrahedral geometry, carbon has foursp3hybrid orbitals– In double bonds with planar geometry, carbon uses threeequivalent sp2hybrid orbitals and one unhybridized porbital– Carbon uses two equivalent sp hybrid orbitals to form atriple bond with linear geometry, with two unhybridized porbitals• Atoms such as nitrogen and oxygen hybridize to formstrong, oriented bonds– The nitrogen atom in ammonia and the oxygen atom inwater are sp3-hybridized
  46. 46. TUTORIAL/ASSIGNMENTChapter 1 : Structure and BondingChapter 2 : Alkanes, Alkenes and AlkynesChapter 3 : Organic Reaction TypesChapter 4 : OrganohalidesChapter 5 : Benzene and AromaticityChapter 6 : Alcohols and CarbonylsChapter 7 : BiomoleculesChapter 8 : Polymers
  47. 47. Chapter 1: Tutorial/Assignment1) Give the ground state electron configuration foreach of the following elements:a) Potassium b) Arsenicc) Aluminium d) Germanium1) What is the hybridization of each carbon atomin acetonitrile, C2H3N?3) Fill in any nonbonding valence electrons thatare missing from the following structures?CH3SSCH3NH2CH3OO-CH3ODimethyldisulfide Acetamide Acetate ion
  48. 48. Chapter 1: Tutorial/Assignment4) Convert the following molecular formulas intoline-bond structures that are consistent withvalence rules:a) C2H6O b)C3H7Brc) C2H4O d) C3H9N5) What kind of hybridization do you expect foreach carbon atom in the following molecules?a) Propane b) 2-methylpropanec) 1- butane-3-yne c) Acetic acid
  49. 49. THAT’S ALL FOR TODAYThank you!!