• Join over 1.2 million students every month
  • Accelerate your learning by 29%
  • Unlimited access from just £6.99 per month
Page
  1. 1
    1
  2. 2
    2
  3. 3
    3
  4. 4
    4
  5. 5
    5
  6. 6
    6
  7. 7
    7
  8. 8
    8
  9. 9
    9
  10. 10
    10
  11. 11
    11
  12. 12
    12
  13. 13
    13
  14. 14
    14
  15. 15
    15
  16. 16
    16
  17. 17
    17

A-Level Investigation - Rates of Reaction – The Iodine Clock

Extracts from this document...

Introduction

MCB A-Level Investigation Rates of Reaction - The Iodine Clock Aim: carry out investigations to determine the rate equation for the oxidation of iodide ions by hydrogen peroxide in acid solution. Investigate the effect of various transition metal catalysts on the reaction rate. Introduction to research The research of this project involved investigation of a range of scientific literature available from Oxford University's online science library, various chemistry texts and other Internet sources. The Iodine clock reaction This reaction has been used repeatedly down the years in the study of reaction rates. Most famously by Harcourt and Esson in the late 18th - early 19th century. The reaction between hydrogen peroxide and iodide ions is still known as the Harcourt-Esson reaction. The Iodine Clock reaction can be represented by the half equations: H2O2 (aq) + I- (aq) --> IO- (aq) + H2O (l) Slow [1] IO- (aq) + H3O+ (aq) - HIO (aq) + H2O (l) Fast [2] HIO (aq) + H3O+ (aq) + I- (aq) --> I2 (aq) + 2H2O (l) Fast [3] Overall: 2H+ + H2O2 + 2I- --> 2H2O + I2 Iodide ions are oxidised to diatomic iodine molecules. These are soluble in water and will turn it a pale brown color. In this experiment the presence of starch will emphasize the color change at the end-point of the reaction by forming a deep blue complex with the iodine. Use of Sodium Thiosulphate [ Na2S2O3 ] Sodium Thiosulphate(VI) is a reducing agent. In the iodine clock reaction, it plays the following part: Half equations; 2S2O32- (aq) --> S4O62- (aq) + 2e- I2 (aq) + 2e- --> 2I- (aq) Overall: 2S2O32- (aq) + I2 (aq) --> S4O62- (aq) + 2I- (aq) The Thiosulphate has the effect of reducing iodine to iodide ions. Since the Iodine Clock Reaction is basically the conversion of iodide ions to iodine, the Thiosulphate actually does the reverse. ...read more.

Middle

This is signified by a change in the solution color from clear to deep blue. (8). Immediately as the reaction reaches this point, stop the stopwatch. Record the time to the nearest second. (9). Repeat this reaction using the same volumes of reactants to obtain a second result. If there were massive variations (>20seconds) it would be advisable to repeat again to eliminate a possible rogue result. (10). Repeat steps 1 to 9 for the remaining volumes of independent variable (in this case 0.1 Molar KI - 4ml, 6ml, 8ml, 10ml and 12ml) adjusting the distilled water content of the solution to keep the overall solution volume at a constant 200ml (respectively 175ml, 173ml, 171ml, 169ml and 167ml). (11). In the same fashion, perform sets of experiments to investigate the effects of varying the concentration of the other reactants: Potassium Iodide and Sulphuric Acid. The various concentrations and volumes for these investigations are listed in Tables 1.2 and 1.3 above. (12). Once all values have been obtained it is necessary to plot graphs of rate against concentration (or volume) limiting reagent (independent variable). To calculate reaction rate: Reaction Rate = 1/REACTION TIME To calculate reactant concentration in the solution: using n = cv Concentration reagent (in solution) = Moles reagent Total solution volume (constant 0.2dm) Where: Moles reagent = Concentration x Volume Units: Concentration: moldm-3 Volume: dm Rate: seconds-1 (13). The graph Rate against Concentration (or volume) should be plotted with ?-axis representing Concentration and y-axis representing Reaction Rate x103 (by raising the reaction rate it will make the numbers easier to work with). (14). To determine the order of the reaction with respect to the concentration of a particular the shape of the graph it produces should be compared with the example below: Figure 2.1 Green line represents 2nd order Red line represents 1st order Blue line represents 0th order The overall order can then be deduced by addition of the singular orders w.r.t. ...read more.

Conclusion

+ e- - Fe2+ (s) Eo /V = +0.77 The following then occurs: Fe3+ (aq) + e- - Fe2+ (s) Eo /V = +0.77 H2O2(aq) + 2H+(aq) + 2e- - 2 H2O (l) Eo /V = +1.77 Hence the catalyst, Fe3+, is regenerated. According to this rule Zn2+, Cr3+, Fe2+ and Cu2+ would not catalyse the oxidation of Iodide ions since they are not strong enough oxidising agents. However it is possible that Fe2+ worked using this reaction path: Fe3+ (aq) + e- - Fe2+ (s) Eo /V = +0.77 H2O2(aq) + 2H+(aq) + 2e- - 2 H2O (l) Eo /V = +1.77 Iron (II) ions are initially oxidised by the Hydrogen Peroxide, which creates water and Iron (III) ions: I2 (aq) + 2e- - 2I- (aq) Eo /V = +0.54 Fe3+ (aq) + e- - Fe2+ (s) Eo /V = +0.77 The Iron (III) ions then oxidise Iodide ions back to iodine whilst regenerating the catalyst. Despite the Eo value suggesting that Fe2+ will not function as a catalyst, the reaction steps simply occur initially in a different order. To explain the fact that Cu2+ ions worked as a catalyst: When considering tabulated electrode potential values one must keep in mind that for the published values it is assumed that the ions are in aqueous solution in their aqua-complex, and that the values are performed at exactly STP. If these ions are bonded differently or have a different ligand in place of the aqua group then the Eo value will be altered dramatically. It is therefore possible that the Eo value of Cu2+ was changed so much by the temperature in the laboratory and the presence of non-aqua ligands that it enabled these ions to catalyse the reaction by increasing their oxidising power. It is more likely that the Cu2+ ions formed a complex with iodide ions which accounts for this change in Eo value - perhaps the complex [CuI4]2-. Mg2+ ions, Zn2+, and Cr3+ ions did not function as catalysts since their Eo values lie outside the range defined by the Iodine clock reaction, 0.54V --> 1.77V. ...read more.

The above preview is unformatted text

This student written piece of work is one of many that can be found in our GCSE Patterns of Behaviour section.

Found what you're looking for?

  • Start learning 29% faster today
  • 150,000+ documents available
  • Just £6.99 a month

Not the one? Search for your essay title...
  • Join over 1.2 million students every month
  • Accelerate your learning by 29%
  • Unlimited access from just £6.99 per month

See related essaysSee related essays

Related GCSE Patterns of Behaviour essays

  1. Marked by a teacher

    Reaction Rates Investigation

    3 star(s)

    This helped us decide how long to test the reaction for. Data recorded for a 3cm piece of Magnesium ribbon at 90 Seconds Time (Seconds) Molar (Hydrochloric Acid) 10 20 30 40 50 60 70 80 90 1 3 11 23 26 34 37 41 45 49 1.2 9 22

  2. Rates of Reaction - The Iodine Clock

    the reaction rate - it is said to be first order with respect to the concentration of that chemical. If a doubling of concentration causes the rate to quadruple - then it is referred to as second order. Similarly, if the increase in concentration has no effect, then it is known as zeroth order with respect to that reactant.

  1. The aim of the investigation is to examine the kinetics involved in the reactions ...

    A graph to show the distribution of energy: Ref: Nuffield Advanced Chemistry Students Book, Longman, p257. The graph shows the number of particles in each range of kinetic energy. The area beneath the curve is directly proportional to the total number of particles.

  2. The Iodine Clock Investigation

    The rate of change of concentration of substance B, rB may also have been measured. The rate of change of the concentration follows a general mathematical expression in the form: Rate = k [A]a[B]b[C]c This is known as the concentration rate equation where: * The rate equation has units of mol dm-3s-1 (other units may be used).

  1. How does the activation enthalpy and the rate of the iodine-clock reaction vary with ...

    Stannous ions (a group 4 'metal' with multiple oxidation states) Mercuric ions (transition metal in +2 oxidation state) Silver ions (transition metal in +1 oxidation state) (3) Analyse the results with the different catalysts - and produce alternative hypothesis. As we are expecting the Activation enthalpy to be similar, an

  2. The Iodine Clock

    When carrying out the preliminary investigation we found that the only cross that would disappear in the allowed amount of time was a florescent yellow cross; out of the possible blue, pink & yellow crosses. During the preliminary investigation, we tested out which was the best way to collect the reactants with was.

  1. An Investigation into the effect of concetration on the Rate of Reaction Between Potassium ...

    A stop clock was taken and as the two solutions were poured in to the beaker the stop clock was started. A person observed above the beaker until the two solutions had reacted and the cross had disappeared and then the stop clock was stopped.

  2. To determine the rate law for a chemical reaction among hydrogen peroxide, iodide and ...

    It was pour into the conical flask. 8. 1 g of solid KI (record the exact mass) and 3 drops of ammonium molybdate catalyst were added into the conical flask. 9. The solution mixture was stirred until the KI dissolves.

  • Over 160,000 pieces
    of student written work
  • Annotated by
    experienced teachers
  • Ideas and feedback to
    improve your own work