Chemistry aqa Unit 2

Physical properties at room temperature: (all have 'swimming pool' smell)

Fluorine- pale yellow gas

Chlorine- greenish gas

Bromine- red-brown liquid

Iodine- black solid

Melting and boilling points increases, more electrons making van der Waals forces between molecules stronger <---Size of atoms increases---> Electronegativity decreases down group due to less atttaction between nucleus and bonding electrons in the outer shell caused by shielding.

F-F bond is unexpectedly weak as small size of atom leads to repulsion.

Oxidising ability increases as we go up the group... Fluorine is a powerful oxidising agent (electron acceptor).

Displacement reactions: Metal halides will be displaced by a more reactive halogen i.e. chlorine will displace bromide ions, iodine will not-

Cl2 (aq) + 2Na (aq) + 2Br- (aq) --> Br2 (aq) +2Na+ (aq) + 2cl- (aq)

Extraction of bromine from sea water Cl2 (aq) + 2Br- (aq) --> Br2 (aq) + 2cl- (aq)  

Extraction of Iodine from kelp 2I- + MnO2 + 4H+ --> Mn2+ 2H2O + I2

Fluorine cannot be investigated in (aq) solution because it reacts with water.

Reducing ability increases as we go down the group... Iodine is a strong reducing agent, the larger the ion, the more easily it loses an electron.

Reactions of solid NaX with conc. H2SO4:

General reaction for all solid halides: NaX (s) + H2SO4 (l) --> HX(g) + NaHSO4(s) i.e.

Sodium Chloride NaCl (s) + H2SO4 (l) --> HCL(g) + NaHSO4 (s) steamy fumes

Sodium Bromide (strong enough to reduce further) steamy fumes & brown fumes of Br                    2H+ +2Br- + H2SO4 (l)  --> SO2 (g) + 2H2O (l) + Br2 (l)

Sodium Iodide (stronger than Br- so reduce even further) Bad egg smell, black solid, steamy fumes, yellow solid sulfur 

8H+ +2I- + H2SO4 (l)  --> H2S (g) + 4H2O (l) + I2 (s)

Metal halides with silver ion:

Dilute HNO3 added to rid soluble carbonate or OH- impurities. Few drops of AgNO3 added to form precipitates. Silver Fluoride, no ppt (soluble in water), Silver chloride, white ppt. Dissolves in dilute NH3., Silver bromide, cream ppt. Dissolves in conc. NH3., Silver Iodide, pale yellow ppt. Insoluble in conc. NH3

Chlorine in industry:        +1                -1

Cl2 (g) + H20 (l) <--> HCLO (aq) + HCL (disproportionation reaction)

2CL2 + 2H20 --> 4HCL + 02 Pale green to colourless in sunlight, so chlorine rapidly lost from swimming pools so, alternative is used: (kept slightly acidic) NaClO + H20 <--> NaOH + HCLO

Active ingredient in household bleach: Cl2 + 2NaOH --> NaCLO + NaCl
+ H20

Physical properties:

High melting points with typical giant metallic structure.  

Magnesium has lowest melting point as the lattice arrangement of atoms are different making them easier to separate.

Atoms get bigger as we go down group so atomic (metallic) radius increases --> electrons in sea of delocalised electrons further from positive nuclei so strenghth of metallic bonds decreases down group--> melting points decreases slightly down group.

Metals get more reactive as we go down group as ionisation energies decrease as less energy required to remove electrons as they are shielded from positive nucleus.

Metals go from oxidation state 0 to +2 by losing 2e-.

Hydroxides:

Reaction with water: Mg(s) + 2H20(l) --> Mg(OH)2 (aq) + H2 (g)

Reaction with steam: Mg(s) + H20(l) --> MgO (s) + H2 (g)

Mg(OH)2 used to neutralise stomach acid (almost insoluble). Ca(OH)2 used in agriculture to treat acidic soil (sparingly soluble). Strontium (more soluble) and Barium (dissolves to produce strongly alkaline solution- Ba(OH)2 + aq --> Ba2+ + 2(OH)- )react more vigorously.

Sulfates:

Less soluble as we go down the group. Barium sulfate is virtually insoluble, used in barium meals and absorbs xrays. Also used to test for sulfates. Solution acidified with HCL to remove carbonate ions, barium chloride added and while ppt. forms if barium sulfate is formed.

Sulfide ores to oxides: by roasting (heating in air). ZnS + 1 1/2 O2 --> ZnO + SO2 

SO2 by-product, sold as sulfuric acid... SO2 + H2O + 1/2 O2 --> H2SO4

Iron: reduction by carbon. 2000K in blast furnace. C + O2 --> CO2 then CO2 + C --> 2CO      Fe2O3 + 3 CO --> 2Fe + 3CO2

Reduction of Mn and Cu: CuCO3 --> CuO + CO2  then  2CuO + C --> 2Cu + CO2

Aluminium by electrolysis: oxide dissolved in molten cryolite, Na3AlF6.

Negative electrode (steal casing): 2Al3+ + 6e- --> 2Al  

Positive electrode (made of carbon): 3O 2-  --> 11/2 O2 + 6e-  carbon electrodes regularly replaced as O2 burns C to CO2

Combination of the two half equations: Al2O3--> 11/2 O2 

Titanium: if reduced with carbon, formation of titanium carbide makes metal brittle.

TIO2 + 2C + 2Cl2 --> TiCl4 (l) + 2CO   TiCl4 is purified by distillation.

TiCl4 (l) + 4Na (l) --> Ti (l) + 4NaCl (l) Reduced with molten sodium under argon atmosphere to prevent metals reacting with nitrogen and oxygen in the air. Mg is an alternative to Na.

Tungsten: WO3 + 3H2 --> W + 3H20  H2 is a flammable gas.

Recycling Iron reduces scrap iron in landfills. Easily separated because it is magnetic. Melting scrap iron does not produce CO2, however energy required to melt scrap iron does produce CO2.

Reduction with scap iron, more economic: Cu2+ + Fe --> Cu + Fe2