Chemistry

  • Created by: jordx24
  • Created on: 08-04-18 22:37
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  • Topic 2 - (Bonding, Structure and Properties of matter)
    • Ionic Bonding
      • Ions are electrically charged particles formed when atoms lose or gain electrons.
        • The loss or gain leaves a complete highest energy level, so the electronic structure of an ion is the same as that of a noble gas.
        • Metal atoms lose the electron, or electrons, in their highest energy level and become positively charged ions.
        • Non-metal atoms gain an electron, or electrons, from another atom to become negatively charged ions.
        • The number of charges on an ion formed by a metal is equal to the group number of the metal.
        • The number of charges on an ion formed by a non-metal is equal to the group number minus 8.
        • Hydrogen forms H+ Ions.
    • Covalent Bonding
      • Sharing Electrons
        • When non-metal atoms bond together, they share pairs of electrons to make covalent bonds.
        • Atoms only share electrons in their outer shells (highest energy levels).
        • Each atom involved generally makes enough to covalent bonds to fill up its outer shell.
        • Covalent bonding happens in compounds of non-metals and in non-metal elements.
        • You can use dot and cross diagrams to draw this.
        • Hydrogen atoms have just one electron. They only need one more to complete the first shell.
        • Chlorine needs one more electron, Oxygen needs two more electrons and Nitrogen needs three more electrons.
        • Simple molecular stuctures - the forces of attraction is very weak, to melt or boil a simple molecular compound, you only need to break these feeble intermolecular forces and not the covalent bonds. So the melting and boiling points are very low, easily parted.
    • Polymers
      • Long chains of repeating units
      • In a polymer, lots of small units are linked together to form a long molecule that has repeating sections. All the atoms in a polymer are joined by strong covalent bonds.
      • To find the molecular formula of a polymer, write down the molecular formula of the repeating unit in brackets. and put a 'n' outside. So for the poly(ethene), the molecular formula of the polymer is (C2H4)n*
      • The intermolecular forces between polymer molecules are larger than between simple covalent molecules, so more energy is needed to break them. This means most polymers are solid at room temperature.
      • The intermolecular forces are still weaker than ionic or covalent bonds, so they generally have lower boiling points than ionic or giant molecular compounds.
    • Giant Covalent Bonding
      • Macromolecules
      • All the atoms are bonded together by strong covalent bonds.
      • They have very high melting point and boiling points as a lot of energy is needed to break the covalent bonds between the atoms.
      • Don't contain charged particles, so they don't conduct electricity.
      • The main examples are diamond and graphite, which are both made from carbon atoms only, and silicon dioxide (silica).
    • Allotropes of Carbon
      • Allotropes are just different structural forms of the same element in the same physical state, e.g. they're all solids. Carbon has quite a few allotropes with lots of different properties.
      • Diamond is very hard, it has a giant covalent structure, made up of carbon atoms that each form four covalent bonds. This makes diamond really hard.
      • Graphite contains sheets of Hexagons, each carbon atom only forms three covalent bonds, creating sheets of carbon atoms arranged in hexagons. There aren't any covalent bonds between the layers - they're only held together weakly, so they're free to move over each other. This makes graphite soft and slippery, so its good for lubricating. It also has a high melting point - the covalent bonds in the layers need loads of energy to break.
      • Graphene is one layer of Graphite, Graphene is a sheet of carbon atoms joined together in hexagons. The sheet is just one atom thick, making it a two-dimensional compound. The network of covalent bonds make it very strong. Its also incredibly light, so can be added to composite materials to improve their strength without adding much weight. Like graphite, it contains delocalised electrons so can conduct electricity through the whole structure. This means it has the potential to be used in electronics.
      • Fullerenes form spheres and tubes, fullerenes are molecules of carbon, shaped like closed tubes or hollow balls. They're mainly made up of carbon atoms arranged in hexagons, but can also contain pentagons (rings of five carbon atoms) or heptagons (rings of seven carbon atoms). Buckminster fullerene was the first fullerene to be discovered. It's got the molecular formula C'60 and forms a hollow sphere containing 20 hexagons and 12 pentagons.
        • Fullerenes can be used to 'cage' other molecules. The fullerene structure forms around another atom or molecule, which is then trapped inside. This could be used to deliver a drug in the body. Fullerenes have a huge surface area, so they could help make great industrial catalysts - individual catalyst molecules could be attached to the fullerenes (the bigger the surface area the better). Fullerenes also make great lubricants.
          • Fullerenes can form nanotubes - tiny carbon cylinders. The ratio between the length and the diameter of nanotubes is very high. Nanotubes can conduct both electricity and thermal energy (heat). They also have high tensile strength (they don't break when they are stretched). Technology that uses very small particles such as nanotubes is called nanotechnology, Nanotubes can be used in electronics or to strengthen materials without adding much weight, such as in tennis racket frames.
    • Metallic Bonding
      • Metallic bonding involves Delocalised Electrons, metals also consist of a giant structure, The electrons in their outer shell of the metal atoms are delocalised (free to move around). There are strong forces of electrostatic attraction between the positive metal ions and the shared negative electrons.  These forces of attraction hold the atoms together in a regular structure and are known as metallic bonding. Metallic Bonding is very strong.
      • Substances that are held together by metallic bonding include metallic elements and alloys. Its the delocalised electrons in the metallic bonds which produce all the properties of metals.
      • Most metals are solid at room temperature. The electrostatic forces between the metal atoms and the delocalised sea of electrons are very strong, so needs lots of energy to be broken. This means that most compounds with metallic bonds have very high melting and boiling points so they are generally solid at room temp.
      • Metals are good conductors of Electricity and Heat. The delocalised electrons carry electrical current and thermal energy through the whole structure, so metals are good conductors of electricity and heat.
      • Most metals are Malleable, the layers of atoms in a metal  can slide over each othe, making metals malleable - this means that they can be bent or hammered or rolled into flat sheets.
      • Alloys are harder than pure metals, Pure metals aren't quite often right for certain jobs - they're often too soft when they're pure so are mixed with other elements to make them harder. Most of the metals we use everyday are alloys - a mixture of two or more metals or a metal and another element. Alloys are harder and so more useful than pure metals.
      • Different elements have different sized atoms. So when another element is mixed with a pure metal, the new atoms will distort the layers of metal atoms, making it more difficult for them to slide over each other. This makes alloys harder than pure metals.

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