Bonding

Bonding

1. Ionic

   • Structure
     • Ionic crystals are giant lattices of ions
       • "Giant" as made up of many repeating units
   • Properties
     • Conduct electricity when molten or dissolved
       • Ions are free to move and carry a charge
     • High melting points
       • Strong electrostatic forces need lots of energy to break bonds
     • Tend to dissolve in water
       • Polar water molecules pull ions aways from lattice causing it to dissolve
   • Bonding
     • Strong electrostatic forces of attraction between oppositley charged ions
       • Electrons transfered between one atom to another forming ions
         • Not just atons, compounds can form compound ions
           • Sulfate SO4 ^2-
           • Hydroxide OH-
           • Nitrate NO3 ^-
           • Carbonate CO3 ^2-
           • Ammonium NH4 ^+
           • Compounds have overall neutral charge

2. Covalent

   • Molecules held together by strong covalent bonds
     • Atoms share electrons to fully saturate both outer shells
       • Both nuclei attact to shared electrons electrostatically
         • Double or triple covalent bonds share multiple pairs or electrons
           • Covalent bonds shown by 'dash' e.g. C-H
             • Covalent bonds don't break during melting or boiling, other than giant covalent substances
               • To melt or boil, just overcome the intermolecular forces, dont need break covalent bonds
                 • Gives simple covalent compounds relatively low melting / boiling points
                   • So Cl2 is gas at room temp has low MP and BP
                     • Diamond does need the covalent bonds broken, sublimes at 3600 cel
                   • MP and BP determined by strength of attraction between particles
       • Dative covalent bond
         • Covalent bond in which both electrons come from the same atom
           • Shown by an arrow to replace the 'dash' N-->H
   • Giant Covalent structures (Macromolecular)
     • Many atoms covalently bonded
     • Diamond
       • Hardest known substance
         • Made of carbon, each bonded covalently with 4 other carbon atoms
           • Forms a tetrahedral shape with very strong covalent bonds
             • High melting point, extremely hard, good thermal conductor
               • Cannot conduct electricty (all e- held in bonds), cannot dissolve, refracts light
     • Graphite
       • Carbon arranged in sheets of flat hexagons with delocalised e-
         • Weak bonds between layers causes them to slide in pencils
           • Delocalised eletrons can move to carry a charge
             • Layers far apart compared to covalent bonds - gives a low density makes strong and lightweight
               • Strong covalent bonds in hexagon sheets causes high melting point
                 • Completely insoluble
       • one carbon bonded to 3 others

3. Shape of molecules

   • Shape of a molecule depends on the number or pairs of electrons
     • Electron pairs exist as charged clouds - there is a high probability of electron pair
       • As the clouds are all negative they repel each other
         • Shape affects how much it repels other clouds
           • LP:LP strongest repulsion, biggest angles
           • LP:BP Strong
           • BP:BP Weakest repulsion, smallest angles
           • Can work out the shape of molecule using VSEPR
             • Find number of electrons on the outer shell of central atom
               • Find how many electrons are used in bonding pairs
                 • Check for charge - positive = less electrons
                   • Sort any left over electrons as lone pairs
             • Shapes
               • 2 Pairs
                 • 2 BP Linear
               • 3 Pairs
                 • 3 BP Trigonal Planar
               • 4 Pairs
                 • 4 BP Tetrahedral
                 • 3BP 1LP Trigonal Pyramidal
                 • 2BP 2LP Bent
               • 5 Pairs
                 • 5 BP Trigonal Bipyramidal
                 • 4BP 1LP Seesaw
                 • 3BP 2LP T-shaped
               • 6 Pairs
                 • 6 BP Octahedral
                 • 4BP 2LP Square Planar

4. Polarisation and Forces

   • Intermolecular forces hold molecules together
     • Electrogenativity is the ablility of an atom to attact an electron pair
       • Fluorine, oxygen, chlorine all strongly electronegative
     • These forces are very weak, weaker than covalent, ionic and metallic bonds
       • vann der Waals
         • Perm. Dipole-Dipole
           • Hydrogen
             • Only happens between FON elements when covalently bonded to hydrogen
               • FON very electrogenative, draws BP electrons from H
                 • As polar bond and Hydrogen high charge density, hydrogen can form bonds with lone pairs on FON atoms
                   • Increases boiling, melting points. Water, HF higher boiling point in halides cause of h-bonds
           • Polar Bonds
             • Causing weak electrostatic fordes of attaction between partial charges
         • Cause all atoms and molecules to attract to each other
           • At any moment electrons likely to be on one side of the atom causing a temp. dipole
             • This can cause another temp. dipole opposite charged on the next atom, 2 dipoles attracted to each other
               • Temp. dipoles constantly created/destroyed overall attraction though
   • Covalent bonds polarised by differences in electronegativity
     • Covalent bonds between the same element are non-polar as same electronegativity
       • Water is polar and substances that are polar will dissolved well in it
   • Difference in electronegativity causes polar bond and a permanent dipole
     • Dipole is a difference in charge between 2 by shift in electron density

5. Metallic

   • Metals exist in giant metallic lattice structures
     • Outer most shell of electrons is delocalised
       • Leaving a positive metal ion that is attracted to d.electrons
         • Forms a lattice of closely packed positive ions
           • In a sea of delocalised electrons (metalic bonding)
             • Properties
               • High melting points as strong electrostatic attraction between electrons and ions
                 • More delocalised electrons per atom, stronger bonding, so higher melting point
                   • D.electrons can pass kinetic energy to each other making them good thermal conductors
                     • Good electrical conductors as d.electrons can move to carry a current / charge
                       • Insoluble (except liquid metals)
             • Materials (NOT RELATED TO METALLIC)
               • Covalent bonds don't break during melting or boiling
                 • To melt or boil, only break bond between molecules, inter force
                   • Simple covalent, easy to overcome these forces
                     • Giant covalent need to break covalent bonds
                       • thats why it never really melts, only sublimes at 3600 cel