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# How does the Activation Enthalpy and the Rate of Reaction vary with the Concentration of Reactants, Catalysts, and the Presence of Different Catalysts?

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Introduction

How does the Activation Enthalpy and the Rate of Reaction vary with the Concentration of Reactants, Catalysts, and the Presence of Different Catalysts? Aim To use the iodine clock experiment to test the benefit of some of the transitional metal catalysts and to explain how they work. The iodine-clock reaction. In this the oxidation of iodide ions to iodine molecules occurs, which are soluble in water and will show up as a pale brown colour. However, if starch is added to the reaction mixture the colour change is to a dark-blue colour, this is because the starch molecules form a complex with iodine. The reaction can be represented by the following half-equation: Equation 1 2 I-(aq) I2 (aq) + 2 e- E /V = +0.54 There is a range of oxidation agents available to carry out this reaction. In fact, almost all transition metal ions may possibly be used to oxidise iodide ions. A look through the table of electrode potentials will give a list of possible oxidation agents. However, in order to study the catalytic properties of the transition metal ions, another type of oxidation agent had to be chosen. The standard iodine-clock experiment uses the reduction of the peroxy-disulphate ions. The half-equation of this reaction is: Equation 2 2 e- + S2O82-(aq) 2 SO42-(aq) E /V = +2.01 Therefore, the overall equation of the reaction is: Equation 3 2 I-(aq) + S2O82-(aq) I2 (aq) + 2 SO42-(aq) As we are using Potassium Salts of both anions, the equation with the spectator ions (K+) would be: Equation 4 2 KI(aq) + K2S2O8(aq) I2 (aq) + 2 K2SO4(aq) This reaction involves two anions having adequate energy to overcome the repulsive forces between them and colliding in the right fashion. This occurs without a catalyst, but the speed at which this happens can be significantly increased with the addition of a transitional metal catalyst. ...read more.

Middle

These are: - * The solution should not be ingested or injected. * Spillages should be cleaned up with an adequate amount of water. * Take extra precautions with salts such as Cupric Sulphate as contact with the skin could be harmful, however at 0.1M the risk associated with the use of such solutions is negligible, unless the is a prolonged contact with the skin. * Mercury salts have a tendency to be poisonous, so a solution more concentrated than 0.002M should be labelled so. The long-term effects of Mercury exposure: - * Metal retardation * However the procedure I intend to carry out does not involve prolonged usage, and the small amount that I will be using can be washed a way down the sink with an ample amount of water when I am finished with it, or if a leakage or spillage should occur. K2S2O8 is an irritant in its solid state so this represents a health hazard. This indicates than when making up my standard solution gloves and eye protection should be worn at all times as this should help to prevent any likely body contact. If a spillage should occur, like all of the above plenty of water should be used to clean it up. In the cases of where water cannot be used due to a hazard (i.e. the electronic balance) use a vacuum cleaner and then wipe thoroughly with a damp clothe to clean. At 0.1M Silver Nitrate is an irritant and will stain clothes. Protective clothes should therefore be worn along with a lab coat. NB. None of the above chemicals should be ingested. If gloves are not used for the dilute solution, hands should be washed thoroughly with soap and warm water at the end of the practical session in order to minimise the risk of ingestion after the lab session. Apparatus * Beakers (50ml & 100ml) * Conical flasks (125ml) * Stop clock * 3-figure balance * Standard bottles (250ml) * Water bath (Room temp-70oC) ...read more.

Conclusion

The catalyst takes part in the reaction but is always returned back to its original state and oxidation number. Orders of reactions KI was found to be of zero order, but I believe that this is due to an excess of iodide ions, meaning that there is always more than enough iodide ions to use up all of the Thiosulphate. But I do believe that my experiments on the order of KI did not go down to a small enough value. Peroxydisulphate was found to be of order 1.118. This unusual value is most likely due to inaccuracies in my method or in the equipment. Further research needs to be done as to what the value should be. If something is of order zero then no mater what concentration is used the reaction rate will not change. Order one shows that if the concentration is doubled the reaction rate is doubled, if the concentration is halved the rate is halved. Another order is of order two, this means that if the concentration is doubled the rate is quadrupled, concentration halved the rate is quartered. Temperature Changes It is very obvious, and unsurprising, that the graph and table results show that an increase in temperature lowers the length of the reaction. As a general rule a rise of 10*C will double the reaction rate. This can be explained with the collision theory; the higher the temperature the faster the particles are moving, this means that they collide more often. Also each particle has more energy and so is more likely to over-come the activation enthalpy when it does collide. These two factors put together mean that a rise in temperature has significant effects on the rate. Resources Ted Lister and Janet Renshaw, Chemistry for A-level Third Edition. Pg369-378 G. F. Liptrot, J. J. Thompson and G. R. Walker, Modern Physical Chemistry. Pg443-452 Ann Lainchbury, John Stephens and Alec Thompson, ILPAC 9 second edition. Pg61-99 (chapters 9-11) Nuffield Advanced Science, Chemistry, Students' book II, Topics 13 to 19. Pg121-153 (chapter 16) Salters Advanced Chemistry 1994 data book. Table 19. Hazcards. ...read more.

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