Group 2 and Group 7 Chemistry

Group 2
1. Electron Configuration:
  1. Always end in the s2 orbital
    1. This is because the highest energy subshell that its most outer electron is in is the s orbital
2. Atom size
  1. Increases as you go down the group because of more shielding, more distance.
    1. This is because there are more electrons which require more shells
3. Bonding
  1. Usually have metallic bonding between each other
    1. Electrostatic forces of attraction between cations and a sea of delocalised electrons
4. Melting Points:
  1. Decrease as you go down the group
    1. Cations have the same charge, but shielding and distance increase
      1. This means there is a weaker nuclear attraction
        This means, as you go down the group, it's easier to overcome the electrostatic forces of attraction
5. Ionisation energies
  1. Decrease as you go down the group
    1. This is because shielding increases, and so does distance
      1. This leads to a weaker nuclear attraction and makes it easier to get rid of an electron
6. Reaction with water
  1. As you go down the group, reactivity increases:
    1. This is because there's more shielding, more distance, weaker nuclear attraction and so it's easier to lose electrons
  2. Group 2 elements will always react with water to give a metal hydroxide and hydrogen gas
  3. Mg + H2O(l) -----> Mg(OH)2 + H2(g)
    1. Magnesium with steam will give you a different product
      1. Mg(s) + H2O(g) -------> MgO + H2(g)
    2. This reaction is not so vigorous
  4. Ca + H2O(l)----> Ca(OH)2 + H2(g)
    1. And so on.. The reactions get more and more vigorous as you go on
  5. Solubility of the group 2 hydroxides
    1. As you go down the group, the group 2 hydroxides get more and more soluble
      1. Mg(OH)2 is the most insoluble, it forms a white precipitate, and is known as a suspnsion in water
      2. Ba(OH)2 is the most soluble and forms a clear solution
      3. The reason for this is because as you go down the group, there is weaker nuclear attraction between the hydroxide ion and the group 2 ion and so they seperate easily in water
        As you go down the group, the solution formed gets more alkaline because there are more OH ions free in the solution
  6. Uses of the hydroxides:
    1. The hydroxides can be used to neutralise acids to give a salt and water
      1. Magnesium hydroxide is used to neutralise stomach acid.
        A weak hydroxide is used because we don't want to neutralise all the stomach acid
        Magnesium carbonate does the same job but leads to the build up of gas in the stomach
      2. Calcium hydroxide can be used to neutralise acidity in soil by farmers
        It's a stronger alkaline than Mg(OH) so does a faster job
        AKA Slaked lime
7. Solubility of the group 2 sulphates
  1. They get more and more insoluble as you go doen the group
    1. i.e. Magnesium sulphate is the most soluble and forms a clear solution
      1. Barium sulphate is the most insoluble and form a white precipitate
        We can use BaSO4 as a Barium meal
        It's swallowed and because it's dense, it is able to absorb x-ray.
        This helps to show soft tissue such as your gut
        Although barium ions are toxic, it is still used because its so hard to seperate it from the sulphate ion
  2. Test for Sulphate ions:
    1. You can add something that will give you barium ions such as BaCl2
      1. It will form a white precipitate of BaSO4 if the sulphate ions are present
        Similarly, this can be done the opposite way to test for Ba2+ ions as it will give the same result
        Problem: BaCO3 also gives a white precipitate
        This can be solved by adding dilute nitric acid first to get rid of the carbonate and hydroxide ions
        It will give carbon dioxide and water as a product
Group 7
1. Electron Configuration
  1. Always end in p5
2. F2-- Yellow vapour
3. Cl2- Green gas
4. Br2- Orange liquid and Brown gas
5. I2- Black solid, Purple gas
6. Size of atom: Increases as you go down the group
  1. More electrons, more shells, bigger atom
7. Electronegativity:
  1. Decreases as you go down the group
    1. This is because there's more shielding and more distance, weaker nuclear attraction, so weaker ability to pull electron towards it
8. Melting points:Increase as you go down the group
  1. More electrons, bigger atoms, more van der Waals formed, more energy required to overcome these van der Waal's
9. Bond Strength: Decreases as you go down a group
  1. Because of more shielding and weaker nuclear attraction which means it's easier to break the bond
  2. Exception: Fluorine
    1. It's a very small atom and so you have more electrons within a smaller space
      1. There is repulsion and so the bond breaks easily
10. Reactions:
  1. Displacement/ Oxidising ability
    1. If something is an oxidising agent, then the element itself is going to be reduced.
      1. In other words, it's going to gain electrons
        The ability of a group 7 element to gain an electron decreases as you go down the group
        This is because, there is more shielding and weaker nuclear attraction and you go down the group
        Displacement reactions are redox reactions where the halogen acts as the oxidisng agent
        A more reactive halogen will displace a less reactive halogen
        e.g. 2NaBr + Cl2-------> 2NaCl + Br2
        e.g. NaBr + I2-------> NaBr + I2
        I2 is less reactive so there will be no change
        Cl2 is more reactive than Br so it will be displaced
  2. Reducing ability
    1. Here, the halide ions act as reducing agents, meaning they make something else gain electrons so they themselves must have lost electrons
      1. As you go down the group, the ability to lose electrons gets stronger as nuclear attraction gets weaker
        In other words, as you go down the group, the reducing ability of the halides gets stronger
    2. We can test the reducing ability of a group 7 element by reacting it with concentrated sulphuric acid
      1. When we react H2SO4 with NaCl, we get a normal acid/base reaction:
        H2SO4 + NaCl-----> HCl + NaHSO4
        The oxidation state of the sulphur hasn't changed: It starts off as 6 and ends as 6. All the oxidation states are the same
        This proves that the chloride ion is not a good reducing agent
      2. When we react with a bromide ion we get:
        H2SO4 + NaBr ------> HBr + NaHSO4
        HBr gives off misty fumes
        But then the HBr reacts further with the H2SO4
        HBr + H2SO4 -------> H2O + Br2 + SO2
        The oxidation state of the Br has gone from -1 to 0. It has lost electrons i.e. been oxidised
        The oxidation state of the sulphur has gone from +6 to +4 meaning it has gained 3 electrons i.e. been reduced
        Half Equations:
        2Br- ------> Br2 + 2e-
        H2SO4 + 2e- + 2H+ --------> SO2 + 2H2O
        Overall Equation
        H2SO4 + 2H+ +2Br- ------> SO2 + 2H2O + Br2
        Observations:
        HBr- Mistry fumes
        Br2- Brown gas, orange liquid
        SO2- Misty fumes
      3. When we react with the Iodide ion: We get the same thing as the bromide ion (Further reaction)
        However, for the products we get H2S
        Half Equations:
        2I- -----> I2 + 2E-
        H2SO4 + 8H+ + 8e- -----> H2S + 4H2O
        Overall equation:
        8I- + H2SO4 + 8H+-----> 4I2 + H2S + 4H2O
        Observation:
        HI - Misty fumes
        H2S - Rotten egg smell
        I2- Black solid, purple vapour
        You may occasionally get S (yellow solid)
        Iodide is a better reducing agent than Br because it made S gain more electrons
11. Testing for Halide ions:
  1. 1. Add dilute nitric acid
    1. 2. Add Silver nitrate (AgNO3)
      1. AgCl- A white precipitate
      2. AgBr- A cream precipitate
      3. AgI- A pale yellow precipitate
      4. Further Test:
        Add dilute ammonia
        AgCl will dissolve, the others will not
        Add conc ammonia:
        Both AgBr and AgCl will dissolve, AgI will not
        AgI is insoluble in ammonia
    2. To get rid of any hydroxide or carbonate ions
12. Uses of Chlorine:
  1. Reaction with water:
    1. H2O + Cl2-----> HClO + HCl
      1. HClO is a disinfectant
        An oxidisng agent
      2. This is a disproportionation reaction where the Cl has been both oxidised and reduced
    2. In the presence of UV light:
      1. 2H2O + 2Cl2-----> 4HCl + O2
        Isn't safe in swimming pools due to acidity
  2. Reaction with NaOH:
    1. NaOH + Cl2-----> NaClO + NaCl + H2O
      1. Also a disproportionation reaction
      2. Sodium chlorate: active ingredient in bleach
        The oxidising agent
      3. Safer to use in pools because:
        NaClO + H2O-----> NaOH + HClO

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Group 2 and Group 7 Chemistry

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	Created by bilkis-21 over 10 years ago