VSEPR

• repulsion scale: lone pair - lone pair > lone pair - bond pair > bond pair - bond pair
• electron pair geometry = base geometry; can have variances; based on the number of electron pairs, both one and bond
• molecular geometry = specific variation of the electron pair geometry; based on the number of lone pairs

my scanner is a nub

oh man i hope i can stay awake in college.


the whole VSEPR lab. put it on your iPhones and you don't have to Biology (as much)





BONDS

Chemical bond: a net attractive force between atoms

• ionic = electrostatic forces that exist between ions of opposite charge
  
  • typically involes metal w/ nonmetal
  • Usually metallic element loses an electron(s) while the nonmetal element gains an electron(s)
  • arranged in a crystal lattice (a large network of the molecules) so a lot of attrative force between ions
  • compounds w/ these bonds are/have
    
    • conduct electricity in solution
    • high melting point
    • soluble in water
    • not so soluble in ethanol
    • insoluble in hexane
    • examples: NaCl, KI, Na2SO4
• covalent = results from the sharing of electrons between 2 atoms
  
  • typically involves nonmetal w/ nonmeta
  • coordinate covalent bonds are when both electrons come from the same atom
  • divided into polar covalent where electrons are unequally shared ...
    
    • compounds w/ these bonds are/have
      
      • nonconudctor
      • medium melting point
      • soluble in water
      • slightly soluble in ethanol
      • insoluble in hexane
      • examples: sucrose, water
    • ... and nonpolar covalent where electrons are equally shared
    • compounds w/ these bonds are/have
      
      • nonconudctor
      • low melting point
      • insoluble in water
      • not very soluble in ethanol
      • soluble in hexane
      • examples: iodine, vegetable oil
• metallic = electrostatic attraction between delocalized electrons and the metallic nuclei within metals
  
  • found in solid metals
  • each metal atom is bonded to several neighboring groups
  • bonding electors are free to move throughout the 3D structure
  • sea of electrons
  • examples: any metal (copper, iron, sodium)

Bond Order: Bond Order= # of shared pairs linking X and Y divided by # of X and Y links in molecule
Bond Length: distance between two bonded nuclei

• As you go down the periodic table, the bond length is bigger. As you go left to right across the table, the bond length gets smaller
• The higher the bond order, the bond length is smaller ( double bonds and triple bonds are smaller that single bonds)

Bond Dissociation Energy

• Enthalpy change for breaking a bond in a molecule with the reactants and products in the gas phase under standard condition
• Energy needed to break a bond/ energy received when made
• It's always positive-- always put in energy to break and release the same energy when the bond is broken

Ion Attraction and Lattice Energy

• The energy of attraction between ions of opposite charge depends on:
  • Magnitude of the ion charges (↑ion charge ⇒ ↑attraction)
  • Distance between ions (↑distance ⇒↓ attraction)

Isoelectronic Species

• Isoelectronic - having the same number of valence electrons and Lewis structures
  • Example: N2 and CO are isoelectronic; CO2, N20, N02, SCN- are isoelectronic (see book p. 389 for pics)
    

Octet Rule = chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons (bonding), has an octet of electrons in its highest occupied energy level

• exceptions //
  
  • hydrogen (2 electrons)
  • beryllium (4 electrons)
  • boron (6 electrons)

Free Radical

• Free radical = chemical species with an unpaired electron
• Examples: NO, NO2

Formal Charge

• Equation: # Valence - (# LPE + #bonds)
• LPE= lone pair electrons

Expanded Valence = can have more than 8 valence electrons

• who can do this?
  • nonmetal elements in period 3 or greater (have d-orbital)

Electroneutrality Principle

1. formal charges must add up to the overall molecule charge
2. formal charges should be as close to 0 as possible
3. negative charges should be placed on the most electonegative atom

POLARITY

• if a molecule is made up of entirely nonpolar bonds then... the molecule is nonpolar
• if a molecule is made up of entirely polar bonds then... the molecule can be either, depending of the shape & dipole moments
  • symmetric shape = nonpolar, nonsymmetric shape = polar
• if a molecule is made up of some polar and some nonpolar bonds... look at what the molecule is primarily composed of

Formal charge of an atom in a molecule or ion = group number - [LPE + (1/2)(BE)]

CHAPTER 10:

Key Terms:

• Hybrid Orbital: an orbital formed by mixing two or more atomic orbitals
• Orbital Hybridization: the combination of atomic orbitals to form a set of equivalent hybrid orbitals
• Sigma bond: A covalent bond that results from head-to-head overlap of atomic orbitals on two different atoms. A sigma bond concentrates electron density along the bond-axis
• Pi bond: A covalent bond that results from the sideways overlap of unhybridized p orbitals on two different atoms. A pi bond concentrates electron density above and below the bond axis.
• Valence Bond Theory: A theory of covalent bonding that pictures chemical bonds as being formed by the overlap of atomic orbitals in the valence shell of the combining atoms.

Valence Bond Theory -

1. qualitive
2. Overlap of atmoic orbitals in the valence shell
3. low potential energy