Covalent Bonds

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  • Covalent Bonding
    • 'A covalent bond is the attraction of a shared pair of electrons for the nuclei of two atoms'
    • A covalent bond is often shown by  a single line between two symbols e.g. H-H
      • Some atoms can also share two pairs of electrons shown by two lines e.g. O=C
        • Triple bonds can also occur e.g between nitrogen atoms shown by  a triple bond
          • Double bonds are stronger than single bonds, and triple bonds stronger still
    • Lone Pairs
      • Outer shell electrons not involved in bonding also form pairs called lone pairs
    • Exceptions to the '8 electrons' rule
      • Group 3 elements have only three outer shell electrons, so can only form 3 normal covalent bonds. They therefore have to have less than 8 electrons in their outer shell
      • There are many compounds where the dot and cross structures show more than eight electrons in the outer shell of an atom but only of the element is in perid three or below
    • Bonding in Compound ions
      • Hydroxide ions are compound ions: the oxygen and hydrogen are covalently bonded together and an electron is gained form elsewhere resulting in a charged ion.
    • 'A dative covalent bond is one in which both of the electrons in the covalent bond are provided by the same atom'
    • Polar and Non-Polar Covalent Bonds
      • In a molecule e.g. H2, the pair of electrons is equally shared between the two identical atoms
        • However, when the bonding atoms are different, the pair of electrons is not equally shared
      • 'Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond'
        • Electronegativity increases from left to right across the periodic table
        • Electronegativity decreases on going down a group in the periodic table
        • The three most electronegative elements are Fluorine, Nitrogen and Oxygen
    • Shapes of Molecules
      • 'Electron pair repulsion theory states that the electron pairs in the outer shell of the central atom in a molecule repel each other and therefore arrange themsleves to be as far away form each other as possible'
      • Linear e.g. Carbon Dioxide
        • Two regions of electron density
        • Two double bonds and no lone pairs
        • Bond angle is 180
      • Trigonal Planar e.g. Boron Triflouride
        • Three regions of electron density
        • Three bonded paris and no lone pairs
        • Bond angle is 120
      • Tetrahedral e.g. Methane
        • Four regions of electron density
        • Four bonded pairs and no lone pairs
        • Must be drawn 4 dimensionally
        • Bond angle is 109.5
      • Octahedral e.g. Sulphur Hexafluoride
        • Six regions of electron density
        • Six bonded paris amd nop lone pairs
        • Bond angle is 90
      • Pyramidal e.g. Ammonia
        • Four regions of electron density
        • Three bonded pairs and one lone pair
        • Bond angle is 107
      • Non-Linear e.g. Water
        • Four regions of electron density
        • Two bonded pairs and two lone pairs
        • Bond angle is 104.5


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