Bonding & forces Chem Unit 2 Pt 2

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Bonding & forces Chem Unit 2 Pt 2
2012-04-03 09:25:24
chemistry bonding forces unit

Starting with carbon structures!
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  1. What is an allotrope? How many allotropes does carbon have?
    • Are different forms of the same element in the same state.
    • Carbon has 3 allotropes - diamond, graphite and fullerenes.
    • Each allotrope has different giant molecular structure.
  2. Describe the structure of diamond.
    • Made up of carbon atoms - each carbon is covalently bonded with sigma bonds to 4 other carbon atoms.
    • Arranged in a tetrahedral shape - this is its crystal lattice structure.
  3. Describe the features of diamond because of its strong covalent bonds.
    • Very high melting point (actually sublimes)
    • Extremely hard - hardest known natural substance - used in diamond-tipped drills & saws
    • Poor conductor of heat and can't conduct electricity (all outer electrons held in localised bonds)
    • Doesn't dissolve in any solvent.
  4. Describe the structure of graphite.
    • Carbon atoms arranged in sheets/layers of flat hexagons
    • Carbon have 3 bonds each
    • The fourth outer electron of each carbon atom is delocalised. The sheets are therefore bonded together by weak intermolecular (London) forces.
  5. Describe features of graphite.
    • Weak bonds between layers in graphite - easily broken - so sheets can slide past each other (graphite feels slippery - used in pencils & dry lubricant)
    • Conducts electricity (delocalised electrons)
    • Layers are far apart compared to length of covalent bonds, so graphite is less dense than diamond (used to make strong, lightweight sports equipment)
    • Very high melting point - becuase of strong covalent bonds in hexagon sheets (sublimes)
    • Insoluble in any solvent (covalent bonds too difficult to break)
  6. What are fullerenes? Features and structure?
    • Shaped like hollow balls or tubes
    • Each carbon forms 3 covalent bonds (leaving free electrons that can conduct electricity)
    • Are nanoparticles
    • Soluble in organic solvents - form brightly coloured solutions
    • Because it's hollow, can be used to "cage" other molecules (structure forms around another molecule) - so can be used to deliver drug into specific cells of body.
  7. Describe a carbon nanotubes.
    • Like a single layer of graphite rolled up into a hollow cylinder.
    • Many covalent bonds make them very strong - used to make stronger lighter building materials
    • Conduct electricity - used as small wires in circuits for computer chips
    • End of tube can be "capped" or closed off to create large cage structure.
  8. Ionic and covalent bonding are the _______ of a _______ of bonding type.
    • extremes
    • continuum
  9. Define electronegativity.
    • The ability of an atom to attract the bonding electron pair in a covalent bond.
    • (Measured using Pauling scale)
  10. What is a dipole?
    • A difference in charge between two atoms caused by a shift in electron density in the bond.
    • (caused by polar bond)
  11. The greater the difference in _______, the more ______ the bond.
    • electronegativity
    • polar
  12. Describe how a polar bond CAN make polar molecules. In what situations doesn't it?
    • If the polar bonds create a net dipole on the molecule, the molecule itself has a permanent dipole and therefore is polar. (The polar bonds need to point roughly the same direction)
    • If the bonds forming the molecule are symmetrical, then these polar bonds will be "cancelled out" and therefore no polar molecule.
    • Lone pairs of electrons on central atom also has an effect (it may cancel out dipole created by other bonding pairs)
  13. The stronger the bond (the attraction between the atoms, higher enthalpy) what happens to the bond length?
    The stronger the bond, the shorter the bond length. More attraction will pull the repelling nuclei closer together.
  14. What are the only bonds that can be purely covalent?
    Only bonds between atoms of a single element (like diatomic gases) such as H2 or O2.
  15. What are the 3 intermolecular forces?
    • Instantaneous dipole-dipole/ London forces
    • Permanent dipole-dipole
    • Hydrogen bonding (strongest)
  16. Explain the trends in boiling temp of alkanes with increasing chain length.
    • Long-chain alkanes have higher boiling points than short-chain alkanes.
    • (More molecular surface area - more electrons to interact)
  17. Explain effect of branching in carbon chain to boiling temp of alkanes.
    • Straight-chain alkanes have higher boiling points than branched alkanes.
    • (smaller molecular surface area)
  18. Explain why instantaneous dipole-dipole forces occur between molecules.
    • Electrons and electron clouds are constantly fluctuating/moving back and forth within an atom and this causes an instantaneous (temporary) dipole (due to motion of electrons)
    • An instantaneous dipole in one molecule then induces a dipole in nearby atoms.
    • Dipoles are being created and destroyed all the time, but the overall effect is the atoms to be attracted to one another.
  19. Hydrogen bonding only happens when hydrogen is covanlently bonded what, what or what?
    • Fluorine, nitrogen or oxygen.
    • (HONF!)
  20. Stronger London forces require higher melting/boiling points. What a molecule have stronger London forces and thus higher melting temp?
    • Larger molecules have larger electron clouds - stronger London forces.
    • Molecules with greater surface area
    • [When you boil a liquid, you need to overcome the intermolecular forces, so greater London forces mean more energy is required to overcome stronger intermolecular forces. - same with melting solids]
  21. What experiment can you use to see if molecules are polar or not?
    • Put an electrostatically charged rod next to a jet of liquid.
    • If liquid is polar, the liquid will "bend" and move towards rod.
    • This is because the polar liquids contain molecules with permanent dipoles.
  22. For one substance to dissolve in another, what must happen?
    • Solute particles must be separated from each other.
    • Solute particles must become surrounded by solvent particles.
    • (Forces between solute-solvent must be strong enough to overcome solvent-solvent forces and solute-solute forces)
  23. What are the 2 main types of solvent?
    • Polar (eg. water)
    • Non-polar (eg. hexane)
  24. Even though Aluminium oxide (Al2O3) is an ionic compound, it is insoluble in water. Why?
    • Because their ionic bonding is too strong and therefore the solvent to solute forces would not overcome this. (Bonds between ions are stronger than those they would form with water molecules).
    • Al3+ has high charge density - very polarising
  25. In what solvents do ionic compounds dissolve in? How?
    • In polar solvents (eg. water)
    • Ions attracted to polar ends (dipoles) of water molecules
    • Water molecules separate the ionic lattice and surround ions (hydrated complexes) in a process called hydration.
    • Energy released called hydration energy.
  26. What solvent can alcohol dissolve in? How?
    • Polar solvents (eg. water)
    • Because the polar O-H bond in an alcohol is attracted to polar O-H bonds in water. Hydrogen bonds form between sigma- O and sigma+ H atoms.
  27. What effect does an increased carbon chain have on an alcohol's solubility in water?
    • Makes it less soluble
    • Because the carbon chain part of the alcohol isn't attracted to water.
  28. Can halogenoalkanes dissolve in water? Why?
    • They can't
    • They contain polar bonds but their dipoles aren't strong enough to form hydrogen bonds with water.
    • Hydrogen bonding between water is stronger than bonds that would be formed with halogenoalkanes, so they wont' dissolve.
  29. Substances usually dissolve best in solvents that have _______ _____/_____. _____ dissolves _____.
    Substances usually dissolve best in solvents that have similar bonds/forces. Like dissolves like.