Group 2 and Group 7 Chemistry

Group 2 and Group 7 Chemistry

1. Group 2

   • Electron Configuration:
     • Always end in the s2 orbital
       • This is because the highest energy subshell that its most outer electron is in is the s orbital
   • Atom size
     • Increases as you go down the group because of more shielding, more distance.
       • This is because there are more electrons which require more shells
   • Bonding
     • Usually have metallic bonding between each other
       • Electrostatic forces of attraction between cations and a sea of delocalised electrons
   • Melting Points:
     • Decrease as you go down the group
       • Cations have the same charge, but shielding and distance increase
         • 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
   • Ionisation energies
     • Decrease as you go down the group
       • This is because shielding increases, and so does distance
         • This leads to a weaker nuclear attraction and makes it easier to get rid of an electron
   • Reaction with water
     • As you go down the group, reactivity increases:
       • This is because there's more shielding, more distance, weaker nuclear attraction and so it's easier to lose electrons
     • Group 2 elements will always react with water to give a metal hydroxide and hydrogen gas
     • Mg + H2O(l) -----> Mg(OH)2 + H2(g)
       • Magnesium with steam will give you a different product
         • Mg(s) + H2O(g) -------> MgO + H2(g)
       • This reaction is not so vigorous
     • Ca + H2O(l)----> Ca(OH)2 + H2(g)
       • And so on.. The reactions get more and more vigorous as you go on
     • Solubility of the group 2 hydroxides
       • As you go down the group, the group 2 hydroxides get more and more soluble
         • Mg(OH)2 is the most insoluble, it forms a white precipitate, and is known as a suspnsion in water
         • Ba(OH)2 is the most soluble and forms a clear solution
         • 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
     • Uses of the hydroxides:
       • The hydroxides can be used to neutralise acids to give a salt and water
         • 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
         • 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
   • Solubility of the group 2 sulphates
     • They get more and more insoluble as you go doen the group
       • i.e. Magnesium sulphate is the most soluble and forms a clear solution
         • 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
     • Test for Sulphate ions:
       • You can add something that will give you barium ions such as BaCl2
         • 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

2. Group 7

   • Electron Configuration
     • Always end in p5
   • F2-- Yellow vapour
   • Cl2- Green gas
   • Br2- Orange liquid and Brown gas
   • I2- Black solid, Purple gas
   • Size of atom: Increases as you go down the group
     • More electrons, more shells, bigger atom
   • Electronegativity:
     • Decreases as you go down the group
       • This is because there's more shielding and more distance, weaker nuclear attraction, so weaker ability to pull electron towards it
   • Melting points:Increase as you go down the group
     • More electrons, bigger atoms, more van der Waals formed, more energy required to overcome these van der Waal's
   • Bond Strength: Decreases as you go down a group
     • Because of more shielding and weaker nuclear attraction which means it's easier to break the bond
     • Exception: Fluorine
       • It's a very small atom and so you have more electrons within a smaller space
         • There is repulsion and so the bond breaks easily
   • Reactions:
     • Displacement/ Oxidising ability
       • If something is an oxidising agent, then the element itself is going to be reduced.
         • 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
     • Reducing ability
       • Here, the halide ions act as reducing agents, meaning they make something else gain electrons so they themselves must have lost electrons
         • 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
       • We can test the reducing ability of a group 7 element by reacting it with concentrated sulphuric acid
         • 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
         • 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
         • 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
   • Testing for Halide ions:
     • 1. Add dilute nitric acid
       • 2. Add Silver nitrate (AgNO3)
         • AgCl- A white precipitate
         • AgBr- A cream precipitate
         • AgI- A pale yellow precipitate
         • 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
       • To get rid of any hydroxide or carbonate ions
   • Uses of Chlorine:
     • Reaction with water:
       • H2O + Cl2-----> HClO + HCl
         • HClO is a disinfectant
           • An oxidisng agent
         • This is a disproportionation reaction where the Cl has been both oxidised and reduced
       • In the presence of UV light:
         • 2H2O + 2Cl2-----> 4HCl + O2
           • Isn't safe in swimming pools due to acidity
     • Reaction with NaOH:
       • NaOH + Cl2-----> NaClO + NaCl + H2O
         • Also a disproportionation reaction
         • Sodium chlorate: active ingredient in bleach
           • The oxidising agent
         • Safer to use in pools because:
           • NaClO + H2O-----> NaOH + HClO