Chapters 4 and 5

Ionic Compounds

• ionic bonding
  • forms ionic compounds
  • network lattice
  • electrons are transferred between atoms of different elements to form ions
  • NaCl
• covalent bonding
  • forms molecules
  • atoms make chemical bonds by sharing electrons between each other
  • HCl
  • a region of e- density between the nuclei arises from this sharing
  • bond formation results in a release of energy
  • Types of covalent bonds
    • Pure/Non-polar covalent (equal sharing)
      • EN 0-0.4
    • Polar covalent (unequal sharing)
      • EN 0.4-1.8
    • ionic 
      • EN >1.8
• Bond Dipole Moments
  • un-even sharing of electrons in a bond leads to polar bond and a dipole moment 

Naming Compounds

• Ionic bonds
  • cation + anion
  • numbers of each ion implied by typical ion charges
• Multiple charged ionic (transition metals)
  • cation(charge) + anion
• Molecular
  • requires a different system due to enormous variation in the ratios in which atoms are combined
  • more metallic + less metallic
  • use prefixes to specify the numbers of atoms of each element
• Binary Acids
  • H + nonmetal
  • hydro + nonmetal + ic acid
• Oxyacids
  • H + polyatomic anion
  • anion + ic acid

Drawing Lewis Structures

• Lewis structures are used to indicate the connectivity of atoms in a molecular compound
  • shows the valence electrons
• Octet rule
  • atoms like to form bonds to achieve 8 valence e-
  • the number of bonds for main group atoms can be predicted from the number of valence e- required to achieve an octet
• Steps
  • Determine the total number of valence electrons
    • for cations, subtract one electron for each positive charge
    • for anions, add one electron for each negative charge
  • Draw a skeleton structure of the molecule/ion, with the least electronegative element in the center. Connect each atom to the central atom with a single bond (2 electrons)
  • Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen), completing an octet around each atom
  • place all remaining electrons on the central atom
  • Minimize Formal Charge by rearranging electrons of the outer atoms to make multiple bonds with the central atom
    • formal charge is the hypothetical charge the atom would have if we could redistribute the electrons in the bonds evenly between the atoms
      • number of valence electrons - number of non-bonding electrons - number of bonding electrons/2
      • have to satisfy the octet rule first
• Exceptions
  • Odd e- molecules 
    • free radicals
    • usually groups of 3 or 5
  • Electron deficient molecules
    • Groups 2 and 12
  • Hypervalent molecules
    • n > 3
    • exceeds normal valence
• Resonance
  • sometimes a single Lewis structure is not adequate to describe a molecule, with the other possibilities only differing by the location of e- (atoms connected the same)
  • resonance structures
    • when a bond can be represented multiple ways
  • the delocalization of e- in Lewis structures
  • each discrete Lewis structure for the same molecule that only differs by the locations of e- is called a resonance form

VSEPR Theory and Molecular Shape

• Assumes that electron pairs in the valence shell of a central atom repel one another, and can be used to predict molecular shape
• e- pairs include bonding e- and lone pair e-
• https://www.google.com/url?sa=i&rct=j&q=&esrc=s&source=images&cd=&cad=rja&uact=8&ved=2ahUKEwi2tZW4zJjfAhXxm-AKHYQ4CeYQjRx6BAgBEAU&url=https%3A%2F%2Fsites.google.com%2Fa%2Fcoe.edu%2Fcourtney-s-chemistry%2Fhome%2Fvsepr-theory&psig=AOvVaw2B50NRkJPUgLrwNTJtRnts&ust=1544645871316887
• electron pairs repel one another
• Bond polarity and dipole moment
  • polar covalent bonds have unequal distribution of e- in a bond, giving a bond dipole moment, represented by mu
  • mu = Qr
    • Q = magnitude of partial charge
    • r = distance between charges
    • mu units = D = Debye
  • 3D shape must be considered to predict the molecular dipole moment
    • sum up the vectors (direction and magnitude) for each bond dipole

Sigma and Pi Bonds

• Valence Bond Theory (VBT)
  • describes a covalent bond as the overlap of half-filled valence atomic orbitals that yield a pair of electrons shared between the two bonded atoms
  • Bond Types
    • sigma bonds
      • single bond
      • head on orbital overlap
      • e- density concentrated along internuclear axis
    • pi bonds
      • double bonds
      • side on orbital overlap
      • e- density concentrated orthogonal to internuclear axis with p orbitals
        • internuclear axis is a node
  • double bonds
    • composed of one sigma bond and one pi bond
    • results from one p orbital and one hybridized orbital
  • triple bonds
    • composed of one sigma bond and two pi bonds
    • results from two p orbitals and one hybridized orbital

Hybridization with VSEPR Theory

• Hybridization of standard AOs creates a new hybrid orbitals that achieve a better picture of the experimental bonding
• Used in VBT to predict accurate bond angles and geometry
• sp3 hybridization
  • results from average of one s and three p orbitals
  • forms 109.5 degree bond angles (tetrahedral)
  • identified by 4 things around central atom
• sp2 hybridization
  • results from average of one s and two p orbitals
  • forms 120 degree bond angles (trigonal planar)
  • identified by 3 things around central atom
• sp hybridization
  • results from average of one s and one p orbital
  • forms 180 degree bond angles (linear)
  • identified by 2 things around central atom

Hybridization with SPE