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Chapter 6.2 : Covalent Bonding and Molecular Compounds
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Chapter 6.2 : Covalent Bonding and Molecular Compounds

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    Chapter 6.2 : Covalent Bonding and Molecular Compounds Chapter 6.2 : Covalent Bonding and Molecular Compounds Presentation Transcript

    • Chapter 6.2
      Covalent bonding and molecular compounds
    • Objectives:
      Define molecule and molecular formula.
      Explain the relationships between potential energy, distance between approaching atoms, bond length, and bond energy.
      State the octet rule
      List the six basic steps used in writing Lewis structures.
      Explain how to determine Lewis structures for molecules containing single bonds, multiple bonds, or both.
      Explain why scientists use resonance structures to represent some molecules
    • Molecular compounds
      Molecule –
      Neutral group of atoms that are held together by covalent bonds
      Compose most living things
      • Molecular compound–
      • Chemical compound whose simplest units are molecules
    • Molecular compounds
      Chemical formula –
      Indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts
      H2O CO2 C12H22O11
      • Molecular formula –
      • Shows the types and numbers of atoms combined in a single molecule of a molecular compound
      • H2O CO2 C12H22O11
      • Diatomic molecule –
      • Molecule containing only two atoms
      • H2 O2 N2
    • Formation of a Covalent Bond
      Nature favors bonding because
      Puts atoms at lower potential energy
      • Approaching nuclei and electrons
      • Attractedto each
      • Decrease in potential energy
      • At the same time, both nuclei and two electrons
      • repel each other
      • Increase in potential energy
      Electron Negative
      Positive Nucleus
      • Potential energy is minimized when attractive forces are equal to the repulsive forces
    • Characteristics of Covalent Bond
      Bond length –
      The distance between two bonded atoms at their minimum potential energy, that is, the average distance between two bonded atoms
      • Bond Energy –
      • The energy required to break a chemical bond and form neutral isolated atoms
      • As the bond energy increases-
      • The bond length decreases
    • Bond length
    • Octet Rule
      Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level.
      • Hydrogen is an exception because it is stable with 2 electron in outer shell.
      • The eight electrons come from the main-group energy levels being filled.
      • s2p6 totals 8 electrons
    • Bonding and octet rule
      F __ __ __ __ __
      1s 2s 2p
      F __ __ __ __ __
      1s 2s 2p
      Bonding electron pair in overlapping orbitals
      Li __ __
      1s 2s
      F __ __ __ __ __
      1s 2s 2p
    • Exceptions to the Octet Rule
      Most main-group elements
      Tend to form covalent bonds according to octet rule
      • Exceptions
      • Hydrogen – forms bonds where it is surrounded by only two electrons
      • Boron - has just 3 valence electrons, so it tends to form bonds in which it is surrounded by six electrons
      • BH3
      H
      B
      H H
    • Electron-Dot Notation
      Electron-configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the elements symbol
      Dots – valence electrons
      Symbol – nucleus and inner-shell electrons
      I
      7 Valence electrons
      53 protons
      73 neutrons
      36 inner shell electrons
    • Electron-Dot Notation
      First three rows of periodic table
      Group
      1 2 13 14 15 16 17 18
    • Lewis Structures
      Formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pair or dashes and dots between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons.
      • Unshared pair ( lone pair ) –
      • Pair of electrons that is not involved in bonding and that belongs exclusively to one atom
      • Structural formula –
      • Indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule
      • Single Bond –
      • Covalent bond produced by the sharing of one pair of electrons between two atoms.
      B F
      B – F
    • Draw the Lewis Structure of iodomethane, CH3I
      1. Determine the type and number of atoms in the molecule
      1 C 3 H 1 I
      2. Write the electron-dot notation for each type of atom in the molecule
      C H I
      3. Determine the total number of valence electrons in the atoms to be combined
      C 1 x 4e-= 4e-
      H 3 x 1e- = 3e-
      I 1 x 7e- = 7e-
      14e-
      Total valence electrons
    • 4. Arrange the atoms to form a skeleton structure for the molecule. IfCarbonis present, it is the central atom. Otherwise, the least-electronegative atom is central (except for hydrogen, which is never central). Then connect the atoms by electron-pair bonds.
      H
      H C I
      H
      5. Addunshared pairs of electrons so that each hydrogen atom shares a pair of electrons and each other nonmetal is surrounded by eight electrons
      H
      H C I
      H
      6. Count the electrons in the structure to be sure that the number of valence electrons used equals the number available. Be sure the central atom and other atoms besides hydrogen have an octet.
      14 e- so this is correct!!!
    • Multiple Covalent Bonds
      Double bond –
      Covalent bond produced by the sharing of two pairs of electrons between two atoms
      H H
      C C
      H H
      H H
      C C
      H H
      OR
      • Triple bond –
      • Covalent bond produced by the sharing of three pairs of electrons between two atoms
      N N
      N N
      OR
    • Draw the Lewis Structure for methanal, CH2O, which is also known as formaldehyde.
      1. Determine the type and number of atoms in the molecule
      1 C 2 H 1 O
      2. Write the electron-dot notation for each type of atom in the molecule
      C H O
      3. Determine the total number of valence electrons in the atoms to be combined
      C 1 x 4e-= 4e-
      H 2 x 1e- = 2e-
      O 1 x 6e- = 6e-
      12e-
    • 4. Arrange the atoms to form a skeleton structure for the molecule. IfCarbonis present, it is the central atom. Otherwise, the least-electronegative atom is central (except for hydrogen, which is never central). Then connect the atoms by electron-pair bonds.
      H
      H C O
      5. Addunshared pairs of electrons so that each hydrogen atom shares a pair of electrons and each other nonmetal is surrounded by eight electrons
      H
      H C O
      6. Count the electrons in the structure to be sure that the number of valence electrons used equals the number available. Be sure the central atom and other atoms besides hydrogen have an octet.
      H
      H C O
      H
      H C O
      OR
      14 e- = 12 e-
      SO
    • Resonance Structures
      Bonding in molecules or ion that cannot be correctly represented by a single Lewis structure.
      O O O
      O O O
      Or