CH2 - Bonding (1)

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  • CH2 - Bonding (1)
    • Chemical Bonding
      • Ionic Bonding
        • Ions form when atoms lose or gain electrons.
          • Positive ions form through electron loss.
          • Negative ions form through electron gain.
        • The Ionic Bond is the attraction between the positive and negative ions.
          • Generally ionic compounds contain a metal ion.
          • Certain amount of repulsion between ions of the same charge.
            • This is minimised by the ions arranging themselves so that they are close only to those of an opposing charge.
              • This attraction outweighs any repulsion.
        • Ionic bonds are electrostatic attraction between oppositely charged ions.
          • E.G. Na2+ O-
          • E.G. Li+ F-
        • Ionic compounds exist as Giant Ionic Lattices.
          • Properties
            • Crystalline solids.
            • High melting point
              • High boiling point.
            • Soluble in water.
            • Electrical conductors when molten or in aqueous solution.
        • When drawing diagrams to represent ionic bonding only show the outer shells.
          • Show the atoms before and after electron transfer.
          • Dot & Cross Diagrams.
      • Covalent Bonding
        • Occurs when 2 atoms share a pair of electrons between them.
          • Electrons involved occupy the same orbital but have different spins.
            • Some repulsion will occur.
          • The attraction of each nucleus for the shared electron pair is the covalent bond.
            • This attraction is greater than the repulsion between the bonding pair of electrons.
          • 2 atom can be the same or different.
            • This is why some covalent substances are elements while others are compounds.
        • Covalent bonds are a shared pair of electrons between two atoms.
          • E.G. C02
          • E.G. H2O
        • Results in the formation of simple molecules.
          • Also results in the formation of macromolecules.
      • Coordinate Bonding
        • A type of covalent bonding as the bond is a shared pair of electrons between 2 atoms.
        • A coordinate bond is a covalent bond in which both shared electrons originate from the sae atom - the donor atom.
        • E.G. Coordinate bonding of NH3 molecule.
          • On the nitrogen atom there is a pair of electrons that have not been used in bonding.
            • This is a lone pair of electrons.
            • These electrons can be donated to an electron deficient atom to form a coordinate bond.
              • An example of an electron deficient atom is Boron Trifluoride.
                • There are only 6 electrons in the outer shell of the boron atom so a donor atom is required to donate a pair of electrons and form a coordinate bond.
    • Electro - negativity and Polar Bonds
      • Electro - negativity is the ability of an atom to attract the shared pair of electrons in a covalent bond.
        • Electro - negativity decreases as you travel down a group.
          • This is because the increasing nuclear charge is greatly outweighed by the distance of the outer shell from the nucleus.
        • Electro - negativity increases as you move across periods.
          • Higher value of electro -negativity on the right side of the periodic table.
        • Group 0 is ignored because those elements do not form bonds.
        • Fluorine is the most electro - negative element.
          • The larger the numerical value the more electronegative that element is.
      • Polarisation of Covalent Bonds
        • e.g. Cl-Cl
          • Identical values = same electro - negative value.
          • Both atoms have an equal attraction for the electron pair that makes the covalent bond.
          • Therefore the bonding electron pair is exactly halfway between both atoms when drawn.
          • This is a non-polar bond.
        • e.g. HCl
          • Cl has a higher electro - negativity than H.
          • This makes the Cl more able to attract the shared pair of electrons than H.
          • Shared electron pair is closer to the Cl atom than the H when drawn.
          • This is a polar bond.
          • It is the difference in the electro - negativities of the two bonded atoms that will influence the amount of ionic or covalent character that the resulting bond posseses.
            • A small difference between the values would suggest that the electron pair is more equally shared.
            • A large difference in the two values would suggest a more polar bond - one atom has a greater ability to attract the electron pair.
            • Not all covalent bonds between two different atoms will be considered as polar.
            • The more polar the bond the greater the partial negative and positive charges of each atom.
    • Forces Between Molecules
      • Intermolecular Forces
        • Forces between molecules
          • Hugely influence the physical properties of substances including: melting point, boiling point, solubility and volatility.
        • 3 Types
          • Permanent Dipole
            • Molecules that contain a permanent polar bond attract each other this way.
            • Exist permanently because of the differences in electro-negativity of the atoms involved in the bond.
            • Stronger than Induced Dipole to Induced Dipole attractions.
              • This is because these bonds are permanent.
          • Induced Dipole
            • Induced = made to happen
            • A very weak force that exists between all single atoms and molecules.
              • Occur because of electron movements within the atom / molecule.
              • E.G. He (Helium)
                • Helium has an extremely low melting point meaning that the forces between the atoms are very weak.
                  • If solid He at 273 degrees (0K / absolute zero) is warmed by one degree it melts.
                    • This means enough energy has been supplied to overcome the forces between the atoms.
            • What is the reason for the need of energy to break the forces?
              • Each He atom contains 2 electrons in constant motion - at any instant they could both be on one 'side' of the atom.
                • This makes one 'side' of the atom slightly positive and the other slightly negative.
                  • This is an INDUCED DIPOLE
                  • Another He atom will have an induced dipole 'set up' causing the two to attract one another.
              • E.G. He (Helium)
                • Helium has an extremely low melting point meaning that the forces between the atoms are very weak.
                  • If solid He at 273 degrees (0K / absolute zero) is warmed by one degree it melts.
                    • This means enough energy has been supplied to overcome the forces between the atoms.
            • A dipole is a separation of charge so that one end of the particle is positive and the other is negative.
            • The induced dipole - induced dipole attractions are stronger between chlorine atoms than He atoms because there are more electrons in the chlorine atom therefore the dipoles are stronger.
              • Therefore chlorine has the higher melting / boiling points because more energy is needed to overcome the ID - ID.
              • On  descending group 7 the melting / boiling points increase because of the increasing number of electrons in each element so there are a greater number of ID - ID attractions.
          • Hydrogen Bonding
            • The strongest type of permanent dipole - permanent dipole attraction.
              • Strongest type of intermolecular force but is only 10% the strength of a covalent bond.
            • Occurs between molecules that contain H bonded to one of the three most electro-negative elements - N, O & F
              • If a molecule contains one of these bonds:
                • H - O
                • H - F
                • H - N
                • It will be able to 'hydrogen bond' to a similar molecule.
            • When  bonded to N, O or F the  two bonding electrons are held so closely  to the more electro-negative atom H exists almost as a bare proton because it has only one electron..
            • The Influence of Hydrogen  Bonding  on Boiling Point
              • The strength of the the Van Der Waals forces increases with increasing Mr.
              • As molecules increase in Mr so does the boiling point.
              • This is because there are a greater number of electrons present in larger molecules.
              • HF, H2O and NH3 all have much higher boiling points.
                • This is because of these molecules having a far stronger attraction  between them.
                  • This increased attraction is caused by hydrogen bonding.
      • IntramolecularForces
        • Forces within molecules
          • E.G Covalent bonds
  • When drawing diagrams to represent ionic bonding only show the outer shells.
    • Dot & Cross Diagrams.
    • Covalent Bonding
      • Occurs when 2 atoms share a pair of electrons between them.
        • Electrons involved occupy the same orbital but have different spins.
          • Some repulsion will occur.
        • The attraction of each nucleus for the shared electron pair is the covalent bond.
          • This attraction is greater than the repulsion between the bonding pair of electrons.
        • 2 atom can be the same or different.
          • This is why some covalent substances are elements while others are compounds.
      • Covalent bonds are a shared pair of electrons between two atoms.
        • E.G. C02
        • E.G. H2O
      • Results in the formation of simple molecules.
        • Also results in the formation of macromolecules.
    • Show the atoms before and after electron transfer.

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