Bonding in compounds

Bonding in compounds

1. Covalent bonding

   • Pure covalent
     • Non-polar
     • Equal sharing of electrons. (e.g both elements have a electronegativity of 3.0)
     • Present in all diatomic elements as they have the same electronegativity
   • Polar covalent
     • A molecule can have a difference in electronegatiity of greater than 0.3 and still not be polar covalent, if the electrons are evenly distributed around the molecule
     • The difference in electronegativtity must be greater than 0.3 for the molecule to be polar
     • A molecule can have a difference in electronegatiity of greater than 0.3 and still not be polar covalent if the electrons are evenly distributed around the molecule.
     • One end of an molecule which is covalently bonded will be δ− with more electons on that element and the other element will be δ+ with its electrons having been taken by the more electoregative element

2. Van der waals' forces

   • London dispersion forces
     • between all molecules and atoms
     • The only force present in non-polar molecules and monatomic elements
     • The weakest attractive force (lowest melting and boiling points as compounds)
     • between all halogens and between noble gas (monatomic) atoms
   • Hydrogen bonding
     • Between polar molecules (intermolecular) where an hydrogen atom is directly bonded to an nitrogen, oxygen or flourine atom
     • The strongest force of attraction( > both LDF and PD-PD interactions)
     • Water is an example of a compounds with hydrogen bonding between molecules
     • Highest melting and boinging points as compounds
   • Permanent dipole permanent dipole interactions
     • Stronger than LDF but weaker than hydrogen bonding
     • Between polar molecules

3. Ionic bonding

   • An electrostatic attraction between the positive ions of one element and the negative ion of another element
   • The greater the difference in electronegativity between two element, the less likely they are to share electrons
   • Element with greater electonegativity gains electrons to form negative ions and the element with the smaller electronegativity value loses electrons to form a positive ion
   • The thurther away two elements are in the periodic table the more likely they are to form ionic bonds
   • In practise, metals combining with non-metals form an ionic compound

4. Viscosity

   • a quantity expressing the magnitude of internal friction in a fluid (resistance to flow / thickness of a liquid)
   • The most viscous liquid in glycerol, then water , ethanol and diethyl ether
     • Liquids with hydrogen bonding between molecules have the highest Viscosity
       • The greater the number of hydroxyl groups the higher the viscosity

5. Solubility

   • Molecules with the same strength of intermolecular bonds are soluble
   • substances with different strengths of intermolecular bonds are insoluble
   • Polar substances are unable to dissolve in non-polar substances e.g. water (polar) is usble to dissolve wax (non-polar)

6. Miscibility

   • Miscible liquids mix thoroughly without any visable boundary between them
   • Oil and water are examples of two liquids which are immisable
     • If the type of bonding is similar they will be miscible e.g both polar molecules that can form hydrogen bonds

7. Thre bonding continuum

   • Places ionic bonding and pure covalent bonding at opposite ends, With polar covalent bonding in the middle
   • Implies that ionic compounds have stronger ionic properties than others. This is based on the electronegativity scale.