chapter 6 - shapes of molecules and intermolecular forces

• tetrahedral, 4 bonding. bond angle of 109.5

  examples include CH4 (methane)

  
• 2 bonding pairs, linear, bond angle = 180

  examples include CO2 and BeCl2

  
• trigonal planar, 3 bonding. bond angle = 120

  examples include BF3 (boron trifluoride)

  
• trigonal bipyramidal, 5 bonding bond angles = 90, 120

  examples include PCl5 (phosphorus pentachloride)

  A) 90

  B) 180

  
• 6 bonding, octahedron with bond angles of 90

  examples include, sulfur hexafluoride SF6

  
• 3 bonding, 1 lone pairs , bond angles=107

  examples include ammonia NH3

  
• 2 bonding, 2 lone pairs. non linear/bent with bond angles 104.5

  examples include water H2O

  
• 4 bonding, 2 lone pairs, square planar. bond angles=90

  examples include xenon tetrafluoride XeF4

  
• electron pair repulsion theory

  • molecules contain covalent bonds, as covalent bonds consist of electron piars they repel each other
  • bonds push each other as far apart as possible to reduce repulsive forces (if repulsion equal the bonds will be equally placed)
• electronegativity - attraction of a bonded atom for the pair of electrons in a covalent bond
• pure covalent bonds ( like H2 and Cl2) bonded electron pairs are shared evenly
• when molecules are of different elements the bonded electron pairs arent always shared equally

  • nuclear charges are different
  • atoms may be different sizes
  • shared pairs may be closer to one nucleus than the other
• electronegativity increases

  • nuclear charge increases
  • atomic radiation decreases

  large pauling value indicates very electronegative

  fluorine is the most electronegative (4.0)

  Noble gases aren't included because they don't form compounds

  non-metals (Cl, F, N, O) have most eletronegative

  group 1 metals are least electronegative
• electronegativity values can be used to estimate the type of bonding
• in non polar bonds, a bonded pair is shared equally between bonded atoms

  non-polar if

  bonded atoms are the same

  bonded atoms have the same/similar electronegativity
• polar covalent bonds

  when bonded atoms are different and have different electronegativity values

  such as HCl (hydrogen chloride)
• dipoles are formed when there is an unequal sharing of electrons within a bond
• the greater the difference in electronegativities, the more polar the bond will be
• polar bonds occur when there is an unequal sharing of electrons within a bond due to differences in electronegativities
• non polar molecules dont have dipoles because all bonds are symmetrical so no partial charges on either end
• Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond towards itself. The electron distribution in a covalent bond between elements with different electronegativities will be unsymmetrical.
• polar bonds form between atoms of elements with different electronegativities so that: The electrons are asymmetrically distributed between the atoms. A permanent dipole is present across the molecule. The higher the electronegativity difference, the more polar the bond.
• if polar bonds are arranged symmetrically so the diploes cancel out snd the molecule has no overall dipole and is therefore non polar
• if they are arranged so they dont cancel out then charge is distributed unevenly, so it will have an overall dipole. molecules with an overall dipole are polar
• types of intermolecular bonds

  london forces

  peremenant dipole dipole

  hydrogen bonding
• london forces are found between all atoms and molecules

  • due to random motion of electrons sometime syou have more in one space than another causing a difference in charge creating a temporary dipole.
  • this temporary dipole can induce opposite dipoles in other atoms/molecules which causes them to be attracted.
  • this causes a third dipole
  • because of random motion dipoles are created and destroyed constantly but overall atoms and molecules are attracted to eachother
• Hydrogen bonding occurs when hydrogen is bound to highly electronegative atoms such as F, O or N. It can form strong dipolar bonds with these atoms. Hydrogen bonding is stronger than London dispersion forces but weaker than ionic/covalent bonds.
• Permanent dipole-dipole interactions occur between polar molecules due to their partial charges. They attract each other because like poles repel and unlike poles attract. This force increases with increasing size of the molecule and decreasing distance between them.
• hydrogen bonding is the strongest intermolecular force

  • only tends to happen with very electronegative atoms such as F,O,N
  • this can have huge effects on the properties of compounds, they tend to be soluble in water and the boiling and freezing points are much higher than in molecules of the same size that don't have hydrogen bonds
• in ice molecules of h2o are held together in a lattice by hydrogen bonds.

  when ice melts hydrogen bonds are broken so that ice has more hydrogen bonds than liquid water. since hydrogen bonds are relatively long this makes ice less dense than water
• simple covalent bonds have low melting and boiling points

  intermolecular forces holding them together is weak and don't require much energy to break. they're often liquids or solids at room temperature. so as intermolecular forces get stronger, melting and boiling points also increase
• polar molecules arre soluble in water

  water is a polar molecule so it only tends to dissolve other polar substances. compounds with hydrogen bonds such as ammonia or ethanoic acid can form hydrogen bonds with water molecules so will be soluble. molecules that only have permanent dipole-dipole forces (like methane) will be insoluble
• simple covalent compounds don't conduct electricity

  even though some covalent molecules shave permanent dipoles, overall covalent molecules are uncharged. so they can't conduct electricity.
• e.g

  chlorine had stronger covalent bonds than bromine but under normal conditions, chlorine is a gas and bromine is a liquid. bomine has higher boiling points because its molecules are bigger giving stronger london forces.
• Electron pair repulsion theory


  Electron pairs repel each other so they are as far apart as possible
• Tetrahedral bond angle


  109.5

  
• Bonded pairs in tetrahedral shape


  4
• Lone pairs in tetrahedral shape


  0
• Pyramidal bond angle


  107
• Bonded pairs in pyramidal shape


  3