Group 7, Equilibria and Reaction Rates

• Chlorine- pale yellow gas; pale yellow-green solution; pale yellow-green hydrocarbon solution.
• Bromine- red-brown liquid (very volatile, gives red-brown gas); red-brown solution; red-brown hydrocarbon solution.
• Iodine: grey-black solid (sublimes on heating to give purple gas); pale yellow solution; red-pink hydrocarbon solution; yellow-brown potassium iodide solution. 

• The halogens are oxidising agents, with their strength decreasing down the group.
• A halogen can oxidise other halide ions if the halide is below it in the group. 
• Chlorine disporportionates in cold dilute alkali, forming a mixture of chlorate (ClO-) and chloride (Cl-) ions (and water).

• Different halide solutions form different coloured precipitates with silver nitrate solution.
• Adding ammonia solution confirms the test, because halides have different solubilities in ammonia solution.
• Chloride ions give a white precipitate, which dissolves in dilute ammonia solution
• Bromide ions give a cream precipitate, which dissolves in concentrated ammonia solution.
• Iodide ions give a yellow precipitate, which does not dissolve in any ammonia solution. 

• All halogens (potassium halides) react with concentrated sulphuric acid to produce a hydrogen halide. 
• Down the group, the strength of the halide ion as a reducing agent (itself losing electrons) increases.
• This means their reactions with concenetrated sulphuric acid will give slightly different products.
• Concentrated acid means it is liquid, not aqueous. 

• KCl reacts to give HCl (steamy fumes) and potassium hydrogensulphate, KHSO4, which is aqueous.
• KBr reacts to give HBr initially, however the HBr is oxidised by the sulphuric acid to give bromine (liquid), while the sulphuric acid is reduced to give sulphur dioxide (gas). This turns acidified potassium dichromate paper from yellow to green.
• KI forms HI, which is immidiately oxidised to form iodine (solid), and sulphuric acid is further reduced to hydrogen sulphide (rotten egg smell)

• To make pure hydrogen bromide or hydrogen iodide from their potassium salts, we have to use a weaker oxidising agent than conc sulphuric acid.
• We use phosphoric (V) acid.
• All 3 hydrogen halides fume in moist air.
• All 3 are extremely soluble in water, forming acidic solutions.
• All 3 react with ammonia gas (NH3) to form the corresponding ammonium halide (eg NH4Cl).
• They all appear as white smoke (solid). 

• Given our knowledge of chlorine, bromine and iodine, we can work out the chemical/physical properties of flourine and astatine.
• Flourine is a gas and astatine a solid, because number of electrons (and therefore Van der Waals forces) increases down the group.
• For the same reasons, Flourine has the lowest boiling point and Astatine the highest.

• Flourine is the most oxiding and the most electronegative.
• Both HF and HAt are soluble in water to form acidic solutions.
• Flourocarbons will be the most stable, astacarbons the least stable, because bond enthalpy decreases down the group. 

• The concentration of an iodine solution can be found by titrating it with sodium thiosulphate of known concentration.
• The thiosulphate ion is S2O3 2- (thio means replace O with S in sulphate ion).
• The reaction between iodine and thiosulphate is a redox reaction.
• Iodine is reduced to iodide ions and thiosulphate is oxidised. 

The factors which influence rate of reaction are:

• Concentration: increasing the concentration of reactants increases the frequency of collisions between particles, leading to more successful collisions, which increases the rate of reaction
• Pressure: increasing the pressure of a gas reaction has the same effect as increasing the concentration, for the same reasons. 

• Temperature- Increasing the temperature increases the collision frequency, but there is also a more important effect that increases the rate of reaction.
• Surface area- Increasing the surface area of a solid increases the rate of reaction.
• Catalysts- increase rate of reaction. 

For a reaction to occur 3 conditions are needed:

• A collision between the reacting particles.
• The reacting particles must have sufficient energy to break their existing bonds so that new bonds can form
• They must collide in the correct orientation.

The energy needed to break the existing bonds is called the activation energy for a reaction. 

• At a given temperature, the particles in a reaction do not have 1 particular energy, but a range of energies.
• When temperature increases, the distribution of molecular energies changes.
• The distribution of molecular energies is known as the Maxwell-Boltznann distribution, which is shown graphically. 

• x axis measures energy, y axis measures fraction of particles with a particular energy.
• No particles have 0 energy.
• At higher temps the peak of the graph is lower and moves to the right, and the area of graph beyond the activation energy is greater.
• This shows there are more particles with energy greater than the activation energy. This is why rate of reaction increases at higher temperatures. 

• A catalyst increases rate of reaction by providing an alternate route for a reaction, with a lower activation energy.
• The effect of a catalyst can be shown on a reaction profile diagram. 
• These are enthalpy level diagrams that show activation energy. 

Chemical equilibria can only be established in closed systems, which have 3 features. They are:

• reactions that do not go to completion- there are always products and reactants.
• reversable reactions- reactants form products, which form reactants, ect.
• dynamic- when a reaction appears to have finished, the chemical continue to react, but the rate of the forward reaction = rate of the reverse reaction. This is called dynamic equilibrium.

• The effect of changes in conditions on chemical equilibria are determined by applying Le-Chatelier's principle.
• This states that: when a change is imposed on a chemical equilibrium, the reaction responds in such a way to oppose the change. As a result, the position of equilibrium changes. 
• The conditions that change are temperature, pressure and concentration. 

• If the temperature of a reaction is increased when the forward reaction is exothermic, the reaction will respond in the direction that lowers the temperature, so the reverse (endothermic) reaction is favoured.
• The forward reaction is favoured if temperature is decreased.
• This is the opposite for an endothermic forward reaction. 

• The pressure of a gas is its number of particles per unit volume.
• If the pressure of a gaseous equilibrium reaction is increased, the reaction will will favour the direction that reduces the pressure- the direction that reduces the number of gaseous molecules.
• Pressure only affects equilibrium reactions where there is a change in the number of gaseous molecules. 

• Lowering the concentration of a reactant will make the reaction favour the reverse reaction, so that more reactant will be formed. 
• Increasing the concentration of a reactant will make the reaction favour the forward reaction, so that less reactant will be formed.