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| Describe the trend in melting/boiling points of Period 2 and 3 elements |
The melting/boiling points of the Period 2 and 3 elements generally increase from the first to the fourth elements in the period, but then decrease from the fourth to the eighth elements
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Name the metals in Period 2 and 3
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PERIOD 2 = Li, Be PERIOD 3 = Na, Mg, Al |
| Describe and explain the melting and boiling points for the metals | For the metals (Li, Be, Na, Mg and Al) melting/boiling points increase across the period because the metal-metal bonds get stronger. The bonds get stronger because the metal ions have... 1. A greater charge 2. An increasing number of delocalised electrons 3. A decreasing ionic radius This leads to a higher charge density, which attracts the ions together more strongly |
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Name the elements with giant covalent structures in Period 2 and 3
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PERIOD 2 = B, C PERIOD 3 = Si |
| Describe and explain the melting and boiling points for the giant covalent structures | The elements with giant covalent structures (B, C and Si) have strong covalent bonds linking all their atoms together. A lot of energy is needed to break these bonds, therefore they have the highest melting/boiling points in their periods |
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Name the elements with simple covalent structures in Period 2 and 3
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PERIOD 2 = N, O, F (,Ne) PERIOD 3 = P, S, Cl (,Ar) |
| Describe and explain the melting and boiling points for the simple covalent structures | The covalent bonds between the atoms in a molecule are very strong but the melting/boiling points depend upon the strength of the London forces between their molecules. These forces are weak and easily overcome. More atoms in a molecule mean stronger London forces. E.g. in Period 3, sulfur (S₈) is the biggest molecule so it has a higher melting and boiling point than P₄ or Cl₂ |
| Describe and explain the melting and boiling points for the noble gases | The noble gases have the lowest melting and boiling points because they exist as individual atoms (they're monatomic), resulting in very weak London forces |
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