Ap Chapter 10 Bonding Ii
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Ap Chapter 10 Bonding Ii

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    Ap Chapter 10 Bonding Ii Ap Chapter 10 Bonding Ii Presentation Transcript

      • Bonding II
    • Bonding I – you learned..
      • Classifying Bonds
      • Calculate the Lattice Energy of Ionic Compounds
      • Write Lewis Structures
      • Formal Charges
      • Drawing Resonance Structures
      • Exceptions to the Octet Rule
        • The incomplete octet
        • Odd-electron molecules
        • The expanded octet
      • Using Bond enthalpies to Estimate the Enthalpy of a Reaction
    • Bonding II – you will learn…
      • Molecular Geometry – VSEPR models
        • Molecule in which the central atom has NO lone pairs
        • Molecule in which the central atom has lone pairs
        • Molecule in which there is more than one central atom
      • Predicting Dipole Moments
      • Hybridization of Atomic Orbitals
        • Hybridization of s and p orbitals
        • Hybridization of s, p and d orbitals
        • Hybridization in molecules containing double and triple bonds.
      • Molecular Orbital Diagrams
    • Molecular Geometry - VSEPR models
      • VSEPR models - Accounts for electron pairs around atoms.
        • Minimizes electron-pair repulsion.
      • Guidelines for Applying
        • Draw the Lewis Structures
        • Only consider electrons around the central atom
          • Account for both bonding and non-bonding (lone) pairs
          • Treat double and triple bonds as single bonds, Ex: CO 2 .. O=C=O.
        • Look at table 10.1 for overall arrangement of electrons.
        • In predicting bond angles
          • Lone pairs repel lone pairs and shared pairs more strongly than bond pairs of electrons.
          • There is no accurate way to predict exact bond angles when the central atom possesses one or more lone pairs.
    • Molecular Geometry - atom has NO lone pairs
      • General formula AB x where A is the central atom, B is/are the surrounding atoms and x is a number between 2 and 6…… most of the time.
      • Table 10.1 shows five possible arrangements of electron pairs around the central atom A .
        • Table 10.1 shows number of electron pairs, arrangement and molecular geometry.
        • Predict the geometry of CO 2 , SnCl 4 and NO 3 -1 , PF 5
      • Problems 10.8, 10.10, 10.12
    • Molecular Geometry – atom has lone pairs
      • General formula AB x E y where A is the central atom, B is/are the surrounding atoms, E is the number of lone pairs and x is a number between 2, 3.. and y 1,2,3, …...
      • Approach..
        • Count all electron pairs on the central atom
        • The number of electron pairs around the central atom determines the electron arrangement around the central atom
        • HOWEVER, the molecular geometry will NOT be the same as the electron arrangement. Geometry is based on atoms alignment, leaving out the lone pairs.
      • Table 10.2 – lone pair configuration
        • Class of molecule, i.e. AB x E y
        • # of electron pairs
        • Number of bonding pairs
        • Arrangement
        • Number of lone pairs
        • Molecular geometry.
      • Practice - Predict the geometry of O 3 , XeF 2 , IF 5
      • Problems 10.14
    • Molecular Geometry - more than one central atom
      • A Central Atom
        • Bonded to two or more atoms
      • Many molecules have more than one central atom.
      • Solve by making each of the central atoms the central atom.
        • Multi-step determination
        • C 2 H 8 and C 2 H 4 are examples
          • H 4 C-CH 4
          • H 2 C=CH 2
    • Predicting Dipole Moments
      • Two factors determine the if a molecule has a dipole moment.
        • Are the bonds in the molecule polar ?
          • Electron negativity determine if the bonds are polar.
          • Shift in electron density is symbolized by
        • Is the molecule polar?
          • Bond moment is a vector quantity and magnitude and direction.
          • Vector is the sum of the bond moment.
          • Check out the following; CO 2 CCl 4 , CCl 2 H 2
      • Problems 10.20, 10.22, 10.24
    • Valence Bond Theory (VB)
      • Introduced to explain chemical bond formation.
      • Describes covalent bonding as overlapping atomic orbitals
      • Orbitals share common regions of space.
      • VB uses the concept of hybridization
        • Blending/combining of two or more non-equal atomic orbits such as s and p to make a new hybrid sp orbit.
        • Hybrid orbitals overlap to create a covalent bond.
        • Hybrid orbitals allow paired electrons to become unpaired for bonding
          • Unpaired valence electrons do the bonding.
      • Steps for determining type of hybrid orbitals
        • Draw Lewis structures
        • Use VSPRE to determine electron pair arrangement (Table 10.1, 10.2)
        • Use Table 10.4 to determine hybrid state of central atom.
    • Valence Bond Theory – sp hybridization
      • sp hybrid orbit
        • Combines the s orbital and one p-orbital to form two equal obitals called sp-orbital, i.e. BeCl 2
        • Linear
      • sp 2 hybrid orbit
        • Combines the s orbital and two p-orbital to form two equal obitals called sp-orbital, i.e. BCl 3
        • Trigonal planar – 120 o angles
      • sp 3 hybrid orbit
        • Combines the s orbital and three p-orbital to form two equal obitals called sp-orbital, i.e. CH 4
        • Tetrahedron – 109.5 o
      • Problems 10.32, 10.34, 10.36
    • Valence Bond Theory – sp hybridization
      • sp 3 d hybrid orbit
        • 5 equivalent hybrid orbitals
        • Trigonal bipyramid, 120 o , 90 o
      • sp 3 d 2 hybrid orbit
        • 6 equivalent hybrid orbitals
        • octahedral, 90 o
    • Hybrid double and triple bonds
      • Determine the bonds that overlap with double and triple bonds
      • Two types of bonds
        • Sigma σ bond – end-to-end overlap
        • Pi bonds – side-to-side overlap