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Chapter Ten

Chapter Ten. Chemical Bonding II : Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory. Contents. Artificial Sweeteners: Fooled by Molecular Shape VSEPR Theory: The Five Basic Shapes VSEPR Theory: The Effect of Lone Pairs VSEPR Theory: Predicting Molecular Geometries

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Chapter Ten

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  1. Chapter Ten Chemical Bonding II : Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory

  2. Contents • Artificial Sweeteners: Fooled by Molecular Shape • VSEPR Theory: The Five Basic Shapes • VSEPR Theory: The Effect of Lone Pairs • VSEPR Theory: Predicting Molecular Geometries • Molecular Shape and Polarity • Valence Bond Theory: Orbital Overlap as a Chemical Bond • Valence Bond Theory: Hybridization of Atomic Orbitals • Molecular Orbital Theory: Electron Delocalization

  3. Artificial Sweeteners: Fooled by Molecular Shape • The sugar molecule (the key) enters the active site (the lock─ sugar receptor protein) of taste cell, resulting in ion channels opened, nerve signal transmission, and reaching the brain a sweet taste. • Artificial sweeteners such as aspartame and saccharin bind to the active more strongly than sugar. • Similarities in the shape of sucrose and artificial sweeteners give those sweeteners the ability to stimulate a sweet taste sensation.

  4. ***** • Valence Shell Electron Pair Repulsion (VSEPR) theory: A theory that allows prediction of the shapes of moleculesor polyatomic ion based on the idea that electrons˗ either as lone pairs or as bonding pairs ˗ repel one another. • Electron geometry: The geometrical arrangement of electron groups in a molecule. • Molecular geometry: The geometrical arrangement of atoms in a molecule.

  5. ***** • VSEPR theory proceeding • Write a best Lewis structure • Determine VSEPR notation: AXmEn: A: Central atoms X: Terminal atoms E: Lone pairs electrons H2O for example: AX2E2

  6. ***** • Determine the electron geometry • An electron group can be: - either single bond or a multiple bond - a (resonance) hybrid bond - a lone pairs of electron - a unpairedsingle-electron • Repulsion force in general: LP vs. LP > LP vs. BP > BP vs. BP * Lone Pairs (LP), Bonding Pairs (BP) Angle for repulsion forces: 90° > 120° > 180° • For central (interior) atom belong to third-period or higher element with VSEPR notation such as AX5, AX4E, AX3E2, AX6, AX5E, AX4E2 require an expanded octet such as 3d orbital. • Multiple bond occupy more space than single bond

  7. ***** • Determine the molecular geometry • Structures for the central atom without lone-pair electrons (AXn type), electron geometry and molecular geometry are identical. • Structures for the central atom with lone-pair electrons (AXnEm type) type), electron geometry and molecular geometry are different.

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  10. ***** * Count only electron groups around the central atom. Each of the following is considered one electron group: a lone pair, a single bond, a double bond, a triple bond, or a single electron.

  11. VSEPR Theory: The Five Basic Shapes (All electrons around the central atom are bonding group) • Two Electron Groups (AX2): Linear

  12. Three Electron Groups (AX3): Trigonal Planar * double bond contains more electron density than the single bond

  13. Four Electron Groups (AX4): Tetrahedral • Five Electron Groups (AX5): Trigonal Bipyramidal

  14. Six Electron Groups (AX6): Octahedral

  15. Example 10.1VSEPR Theory and the Basic Shapes Determine the molecular geometry of NO3−. Solution NO3− has 5 + 3(6) + 1 = 24 valence electrons. The Lewis structure has three resonance structures: Use any one of the resonance structures to determine the number of electron groups around the central atom. The nitrogen atom has three electron groups. The electron geometry is trigonal planar: The molecular geometry is also trigonal planar.

  16. VSEPR Theory: The Effect of Lone Pairs (Some electrons around the central atom are lone pairs) • Four Electron Groups with Lone Pairs AX3E AX2E2

  17. * Effect of Lone Pairs on Molecular Geometry

  18. Five Electron Groups with Lone Pairs AX4E AX3E2 AX2E3

  19. Six Electron Groups with Lone Pairs AX5E AX4E2

  20. 4. VSEPR Theory: Predicting Molecular Geometries Example 10.2Predicting Molecular Geometries Predict the geometry and bond angles of PCl3. Solution Step 1PCl3 has 26 valence electrons. Lewis structure for the molecule: Step 2The central atom (P) has four electron groups. Step 3Three of the four electron groups around P are bonding groups and one is a lone pair. Step 4The electron geometry is tetrahedral The molecular geometry is trigonal pyramidal

  21. Example 10.3Predicting Molecular Geometries Predict the geometry and bond angles of ICl4−. Solution Step 1ICl4− has 36 valence electrons. Lewis structure for the molecule: Step 2The central atom (I) has six electron groups. Step 3Four of the six electron groups around I are bonding groups and two are lone pairs. Step 4The electron geometry is octahedral The molecular geometry is square planar.

  22. Representing Molecular Geometries on Paper Examples:

  23. Predicting the Shapes of Larger Molecules Example:

  24. Example 10.4Predicting the Shape of Larger Molecules Predict the geometry about each interior atom in methanol (CH3OH) and make a sketch of the molecule. Solution The Lewis structure of CH3OH. Three-dimensional sketch of the molecule:

  25. ***** • Memo for VSEPR • Without lone-pair electrons

  26. ***** • With lone-pair electrons

  27. 5. Molecular Shape and Polarity • Bond dipole versus Molecular dipole • Bond dipole: A separation of positive and negative charge in an individual bond. • Molecular dipole: • For diatomic molecule: molecular dipole is identical to bond dipole. • For a molecule consisted by three or more atoms, molecular dipole is estimated by the vector sumofindividual bond dipole moment (net dipole moment).

  28. Polar molecule versus Nonpolar molecule • Polar molecule: A molecule in which the molecular dipole is nonzero. • Nonpolar molecule: A molecule in which the molecular dipole is zero. • Molecular polarity prediction • Draw the Lewis structure for the molecule and determine its molecular geometry. • Determine if the molecule contains polar bonds by electronegativity values. • Determine if the polar bonds add together to form a net dipole moment.

  29. ***** • Examples CO2 Molecular geometry: linear (net) dipole moment: m = 0 D Nonpolar molecule H2O Molecular geometry: bent (net) dipole moment: m = 1.84 D Polar molecule

  30. Example 10.5Determining if a Molecule Is Polar Determine if NH3 is polar. Solution Lewis structure: Determine if the molecule contains polar bonds. The electronegativities of nitrogen and hydrogen are 3.0 and 2.1, respectively. Determine if the polar bonds add together to form a net dipole moment. The three dipole moments sum to a net dipole moment. Ans: The molecule is polar.

  31. Polarity effects of the intermolecular forces • Example 1: For H2O • Example 2: Like dissolves like Polar molecules interact strongly with other polar molecules excluding the nonpolar molecules and separating into distinct regions.

  32. Quantum-Mechanical Approximation Technique • Perturbation theory (used in valence bond theory): A complex system (such as a molecule) is viewed as a simpler system (such as two atoms) that is slightly altered or perturbed by some additional force or interaction (such as the interaction between the two atoms). • Variational method (used in molecular orbital theory): The energy of a trial function (educated function) within the Schrodinger equation is minimized.

  33. Schrodinger equation revisited H = E • H (Hamiltonian operator), a set of mathematical operations that represent the total energy (kinetic and potential) of the electron within the atom. • E is the actual energy of the electron. • is the wave function , a mathematical function that describes the wavelike nature of the electron. • Perturbation theory: Approach by small changes to a known system in which Hamiltonian operator is modified. • Variational method: Approach by combining systems of comparable weighting in which wave function is modified.

  34. Valence bond theory versus molecular orbital theory • Valence bond theory (VB): An advanced model of chemical bonding in which electrons reside in quantum-mechanical orbitals localized on individual atoms that are a hybridized blend of standard atomic orbitals; chemical bonds result from an overlap of these orbitals. • Molecular orbital theory (MO): An advanced model of chemical bonding in which electrons reside in molecular orbitals delocalized over the entire molecule. In the simplest version, the molecular orbitals are simply linear combinations of atomic orbitals.

  35. 6. Valence Bond Theory: Orbital Overlap as a Chemical Bond • Valence bond theory describes that covalent bonds are formed when atomic orbitals on different atoms overlap. • Simple Atomic Orbitals (AO’s) Overlap Bonding in H2 for example • A covalent bond is formed by the pairing of two electrons with opposing spins in the region of overlap of atomic orbitals between two atoms. • This overlap region has a highelectron charge density. • The overall energy of the system is lowered.

  36. Acceptable simple Atomic Orbitals (AO’s) Overlap Bonding in H2S for example • Predicted H˗S˗H angle is 90o, actual H˗S˗H angle is 92o, therefore, the simple AO overlap is acceptable for H2S molecule.

  37. Unacceptable simple Atomic Orbitals (AO’s) Overlap Example 1: CH4 * Actually, the central atom of H2S, H2O, NH3, and CH4, are sp3 hybridization Example 2: NH3 and H2O

  38. ***** 7. Valence Bond Theory: Hybridization of Atomic Orbitals • Hybridization: A mathematical procedure in which standard atomic orbitals are combined to form new, hybrid orbitals. • Hybridizing is mixing different types of orbitals in the valence shell to make a new set of degenerate orbitals such as sp, sp2, sp3, sp3d, sp3d2. • Hybrid orbitals minimize the energy of the molecule by maximizing the orbital overlap in a bond. • Those central atoms are available hybridized, however, those terminal atoms are supposed to be unhybridized.

  39. ***** • General statements regarding hybridization • Hybridization is employed for central atom only, thus, the hybrid orbital describes the electron geometry for central atom. • Number of hybrid orbitals = Number of standard atomic orbitals combined = Number of σ bond+ Number of lone pairs. • Number of hybridization obitals of a central atom = 2 → sp; = 3 → sp2; = 4 → sp3; = 5 → sp3d; = 6 → sp3d2. • Hybrid orbitals may overlap with standard atomic orbitals or with other hybrid orbitals to form σ bond. • Molecular geometry is described by the relative atomic position around central atom.

  40. sp3 hybridization (C for example) one s orbital with three p orbitals combine to form four sp3 hybrid orbitals (degenerate).

  41. ***** Examples of sp3 hybridization (for central atom) Standard orbitals Central atom Mole- cule Hybrid Orbital Geometry σ σ σ σ C 2s 2p sp3 σ σ σ lone N 2s 2p sp3 σ σ lone lone O 2s sp3 2p

  42. sp2 hybridization (B for example) one s orbital with two p orbitals combine to formthree sp2 hybrid orbitals

  43. ***** Examples of sp2 hybridization (for central atom) Standard orbitals Unhybridized Orbital Central atom Mole- cule Hybrid Orbital σ σ σ B 2s 2p 2p sp2 σ σ σ π C 2p 2s 2p sp2 σ σ π lone 2p 2s N 2p sp2

  44. sp hybridization (Be for example) one s orbital with one p orbitals combine to formtwo sp hybrid orbitals

  45. ***** Examples of sp hybridization (for central atom) Standard orbitals Unhybridized Orbital Central atom Mole- cule Hybrid Orbital σ σ Be 2s 2p 2p sp π π σ σ C 2s 2p 2p sp

  46. ***** • About Multiple Covalent Bond • σ(sigma) bond: The first covalent bond formed by end-to-end overlap of standard or hybridized orbitals between the bonded atoms: s + s, s + p, p + p (end-to-end), s + hybrid orbital p + hybrid orbital, hybrid orbital + hybrid orbital • π(Pi) bond: The second (and third, if present) bond in a multiple bond, results from side-by-side overlap of unhybridized p orbitals: p + p (side-by-side) • Summary: • Single bonds: oneσ bond • Double bond: one σ bond and one π bond • Triple bond: one σ bond and two π bonds

  47. Sigma Bonding and Pi Bonding

  48. VB theory of bonding in ethylene (H2C=CH2) example of sp2 hybridization and a double bond • Lewis structure • A π-bond has two lobes (above and below plane), but is one bond, side-by-side overlap of 2p–2p

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