Structure and Bonding

• ionic bond - a strong electrostatic attraction between oppositely charged ions in a compound.
• covalent bond - a strong electrostatic attraction between positively charged nuclei and a shared pair of electrons.
• metallic bond - a strong electrostatic attraction between positive metal ions and a sea of delocalised electrons
• dative (coordinate) covalent bonding - a shared pair of electrons, where both electrons are supplied by 1 atom. These bonds are shown by an arrow.

  
• period 3 elements exclusively can exceed their 8 electron occupancy.
• ionic structures properties

  • high mp - lots of energy is required to break strong ionic bonds.

  
  

  • don't conduct in solid but do conduct in molten - ions become free to move.

  
  

  • brittle - electrostatic repulsion of liked charged ions when layers slide cause them to break.

  
  

  • most can dissolve in water, as water molecules are polar, and can attract the positive and negative ions to break up the structure.
• in ionic structures, ionic bond strength:

  • increases with charge density of both the cation and the anion (high charge and low ionic radius)

  
  

  YOU HAVE TO MENTION BOTH!
• bond strength is inversely proportional to bond length.
• a polar liquid can only dissolve a polar solid (i.e. water dissolves NaCl)

  
• metallic structures properties:

  • high mp - lots of energy is required to overcome strong metallic bonds.
  • malleable - layers of metal ions can slide over each other. Alloys aren't malleable as different sized particles disrupt the layer structure, meaning layers can no longer slide readily.
• More metallic structures:

  • as the charge on a metal ion increases, number of delocalised electrons also increases, as more electrons are being lost to form the ion. This thus makes bonds stronger and increases its mp.

  
  

  • down a group, mp decreases since the atomic radius is larger, thus the electrostatic attraction of delocalised electrons to the now less charge dense cation decreases.
• giant covalent structure properties (diamond or silicon dioxide or regular silicon):

  • tetrahedral shape

  
  

  • very high mp and sublimation point - lots of energy is required to break many strong covalent bonds.

  
  

  • Insulates (electricity) - all of its valence electrons are localised in covalent bonds.

  
  

  • conducts heat - due to its tightly packed, rigid arrangement.
• giant covalent structure properties (graphite):

  • conductor - each carbon leaves 1 delocalised electron which is free to move.

  
  

  • high mp - lots of energy is required to overcome many strong covalent bonds.

  
  

  • soft - weak dispersion force between layers, allowing layers to slide.

  
  

  • low density - layers are far apart in comparison to covalent bond length
• bond length - the distance between the centres of two covalently bonded atoms.
• Valence shell electron pair repulsion theory (VSEPR) - electrons will arrange themselves so they are as far apart as possible.
• Repulsion

  bond pair:bond pair<bond pair: lone pair< lone pair: lone pair
• each lone pair reduces the bond angle by 2.5 degrees because of their higher repulsion.
• Explaining molecular shape steps:

  1. state the number of bond and lone pairs.
  2. state the valence shell electron pair repulsion theory.
  3. If there are no lone pairs, state that electron pairs repel equally.
  4. If there are lone pairs, state that lone pairs repel more.
  5. state the actual shape and bond angles.
• Molecular Shapes (2 total e- pairs)

  • 2 bond pairs and 0 lone pairs = linear.
• Molecular Shapes (3 total e- pairs)

  • 2 bond pairs and 1 lone pair = bent

  
  

  • 3 bond pairs and 0 lone pairs = trigonal planar
• Molecular Shapes (4 total e- pairs)

  • 4 bond pairs and 0 lone pairs - tetrahedral

  
  

  • 3 bond pairs and 1 lone pair - trigonal pyramidal.

  
  

  • 2 bond pairs and 2 loan pairs - bent
• Molecular Shapes (5 total e- pairs)

  • 2 bond pairs and 3 lone pairs - linear

  
  

  • 3 bond pairs and 2 lone pairs - t-shaped (90 degrees bond angle goes to 87.5)

  
  

  • 4 bond pairs and 1 lone pair - see saw

  
  

  • 5 bond pairs and 0 lone pairs - trigonal bipyramidal
• Molecular Shapes (6 total e- pairs)

  • 4 bond pairs 2 lone pairs - square planar (bond angle remains unchanged as the two lone pairs repel each other equally. This is because they are placed opposite to each-other to minimise repulsion)

  
  

  • 5 bond pairs 1 lone pair - square pyramid

  
  

  • 6 bond pairs 0 lone pairs - octahedral
• Bond angles

  linear - 180 degrees

  trigonal planar - 120 degrees

  tetrahedral - 109.5 degrees

  trigonal bipyramidal - 90 degrees (between axial and equitorial) and 120 degrees (between 2 equitorials)

  octrahedral - 90 degrees.
• lone pairs occupy equitorial positions to minimise repulsion.
• Images for Molecular Shapes 2

  
• Images for Molecular Shapes 1

  
• Intermolecular forces only apply to simple molecular structures.
• electronegativity - ability of an atom in a covalent bond to attract the bonding pair of electrons.

  These atoms are at the top right of the periodic table.
• ionic bond - large electronegativity dif

  covalent bond - small electronegativity dif
• polar molecule - unequal distribution of electrons (dipole) due to atoms having different electronegativities. The more electronegative atom pulls electrons towards itself.

  
  

  atoms with the same electronegativity value are not polar, and the shared pair of e- sit in the middle (hydrocarbons)
• polar molecules must be asymmetrical.
• symmetrical molecules have no lone pairs.
• a symmetrical molecule will not be polar, even if individual bonds within the molecule are polar, as dipoles cancel out.

  
• only asymmetrical molecules with polar bonds will be polar:

  
• 3 types of intermolecular force:

  • London dispersion forces
  • permanent dipole:permanent dipole forces
  • hydrogen bonds.
• London dispersion forces (all molecules)

  the random movement of electrons (instantaneous and induced dipoles) makes a slight electron density on one side of the molecule.

  
  

  the bigger the molecules/atom, the more electrons it has and hence the stronger the London forces it has since it has a bigger electron cloud.(e.g. why HI has a higher bp than HCl)

  
• Factors affecting dispersion forces

  • number of electrons - more electrons = stronger forces.
  • shape of molecules (in isomers when number of electrons is the same). The branched isomer has a lower bp and thus weaker dispersion forces, because the SA of contact between molecules is lower in the branched isomer.

  
• Permanent dipole:permanent dipole forces (polar molecules only). These molecules also have dispersion forces.

  
• Hydrogen bonding - polar molecules containing hydrogen atoms covalently bound to F, O or N (the most electronegative and thus form the biggest dipole).