Bonding and structure

Bonding and structure

1. Ionic bonding

   • Electrostatic attraction between oppositely charged ions.
   • Metal + non-metal.
   • Electrons transferred to the non-metal, forming oppositely charged ions that attract.

2. Covalent bonding

   • A bond formed by a shared pair of electrons.
   • Lone pairs aren't used in the covalent bond.
   • Dative covalent bonds is a shared pair of electrons which has been provided by one atom.

3. Hydrogen bonding

   • Molecules containing O-H and N-H bonds are polar with permanent dipoles. Dipoles are strong.
   • A strong dipole-dipole attraction between an electron deficient hydrogen atom and a lone pair of electrons on a highly electronegative atom.
   • Properties of water
     • Ice is less dense than water - Ice has an open lattice with hydrogen bonds holding the water molecules apart, when ice melts the hydrogen bonds collapse allowing H20 molecules to move closer.
     • High melting and boiling points - strong bonds, extra bonds over and above van der Waal's forces, a lot of energy needed to overcome strong bonds.
     • Extra intermolecular bonding from hydrogen bonds explains high surface tensions and viscosity of water.

4. Metallic bonding

   • Electrostatic attraction between positive metal ions and delocalised electrons.
   • Positive ions are fixed in a lattice. Outer shell electrons are delocalised, shared between all atoms in the structure.
   • Properties of giant metallic lattices
     • High melting and boiling points - attraction between positive ion and negative delocalised electrons is strong, high temps needed to overcome them.
     • Electrical conductivity - delocalised electrons can move and carry charge.
     • Malleability and ductility - delocalised electrons can move, allowing layers to slide.
   • Alloys are mixtures of metals, in different proportions, allows modification of a metals properties.

5. Ionic charges

   • Ionic charges can be predicted by how many electrons are needed to reach noble gas electron configuration.
   • Be, B, C, Si don't form ions, too much energy is needed to transfer the outer-shell electrons.
   • Molecular ions - groups of covalently bonded atoms lose or gain electrons to form ions.
   • Predicting ionic formulae
     • total + = total -
     • Overall charge is 0. How many + are needed to balance the - charges.

6. Structure of ionic compounds

   • Each ion is surrounded by oppositely charged ions, attracting each other forming a giant ionic lattice.
   • Properties
     • High melting and boiling points -- strong electrostatic forces, greater charge, stronger forces.
     • Electrical conductivity - when solid, ions are in a fixed position so electrical charge cannot be carried. When melted/dissolved the lattice breaks down and ions can move and carry electrical charge.
     • Solubility - dissolves in polar solvents, breaks down, water molecules surround ions.

7. Structures of covalent compounds

   • Simple molecular structures
     • 3D structure of molecules, bonded by weak intermolecular forces.
     • Atoms of each molecule bonded by strong covalent bonds.
     • Properties
       • Low melting and boiling points - weak, van der Waals' intermolecular forces, little energy needed to overcome them.
       • Electrical conductivity - no free charged particles to carry charge.
       • Solubility - soluble in non-polar solvents because of van der Waals' forces between simple molecular structure and the non-polar solvent, weakening the lattice.
   • Giant covalent structures
     • Diamond - strong covalent, no delocalised electrons so not conductive, hard due to tetrahedral structure as external forces are spread. Graphite - delocalised electrons between layers so conductive, soft due to weak forces between layers allowing sliding.
     • Properties
       • High melting and boiling points - high temps needed to break strong covalent bonds.
       • Electrical conductivity - no free charged particles (except graphite) so not conductive.
       • Solubility - insoluble in polar and non-polar solvents, covalent bonds are to strong to be broken by solvents.

8. Electronegativity and polarity

   • Polar and non-polar bonds
     • Non-polar is when the electrons in the bond are evenly distributed. A polar bond has an overall dipole.
     • Dipoles cancel out in symmetrical molecules.
   • Electronegativity is a measure of the attraction of a bonded atom for the pair of electrons in a covalent bond.
   • The bonding atom with a greater attraction for the electron pair is more electronegative.
   • Small difference in electronegativity, fair share of electrons, polar covalent. Large difference in electronegativity, unfair share, ionic bond.
   • Electronegativity increases to the top right of the periodic table.

9. Intermolecular forces

   • Intermolecular forces are weak because they are cause by weak attractive forces between small dipoles in different molecules.
   • Hydrogen bonds, permanent dipole-dipole forces, van der Waals' forces.
   • Permanent dipole-dipole forces
     • Weak attractive force between permanent dipoles in neighbouring polar molecules.
     • Polar molecules have permanent dipoles. The permanent dipole of one molecule attracts one in another polar molecule, forming weak permanent dipole-dipole forces.
   • Van der Waals' forces
     • Weak, small, temporary attractive forces between induced dipoles in neighbouring molecules between polar and non-polar molecules.
     • Boiling points - electrons increase, van der Waals' forces increase, boiling points increase.
     • Caused by movement of electrons in shells, unbalancing charge distribution in shells, causes instantaneous dipoles, induce a dipole in the neighbouring molecule.
     • Small induced dipoles attract each other causing weak intermolecular forces.