The purpose of this experiment was to determine the rate law graphically from the rate of disappearance and the x y values also the specific rate constant (k). Activation energy was also determined, and the effect of catalyst was evaluated in the reaction between peroxodisulphate ion S2O82-, and iodide ion, I-. S2O82-(aq) + 3 I-(aq) --> 2 SO42-(aq) + I3(aq)
The general expression for the rate law, given this overall reaction, is: rate of disappearance of S2O82- = k[S2O82-]m[I-]n
Chemical Kinetics is the chemistry of how fast or slow something is. These rates of reaction or the speed depends on many factors. Experiments show that the rate of homogenous reactions depend upon: The nature of the reactants, the concentration of the reactants, the temperature and the catalysis (Department of Chemistry and Biology, 2012). In this experiment the affect of temperature and concentration on the rate of reaction was determined.
Temperature is the measure of average kinetic energy. An increase in this energy would in turn increase the collisions of the particles. So higher temperature implies higher average kinetic energy of molecules and more collisions per unit time. A general rule of thumb for most chemical reactions is that the rate at which the reaction continues will approximately double for each 10°C increase in temperature.
Changing the concentration of a solution alters the number of particles per unit volume. A higher concentration of reactants leads to more effective collisions per unit time, which leads to an increasing reaction rate. In this experiment as the concentration of S2O82- , the rate of the reaction also increases.
A catalyst lowers the activation energy of a chemical reaction and increase the rate of a chemical reaction without being consumed in the process (McMurray J et al, 2012). A lower activation energy for the catalyzed reaction means that, at a given temperature, a larger fraction of the molecules will possess enough energy to reach, and thus the catalyst increases the rate of both the forward and backward reactions (Department of Chemistry and Biology, 2012).
In this experiment EDTA was added to uncatalyze the reaction, and to remove any catalytic effects of copper. This also increases the activation energy.
The activation energy is the energetic barrier that has to be overcome between the reactants and the products. At the top of this energetic barrier, there is a transition state complex that forms instantaneously before the products are formed. The difference in energy between the initial reactants and the activated complex C is the activation energy, Ez of the reaction (Department of Chemistry and Biology, 2012). Arhenius’s equation can be used to determine the activation energy.
The rate of a reaction is determined by observing the rate of disappearance of the reactants. Rate of disappearance of A = change in concentration of A/time reacquired for change = -∆A/∆T Rate of appearance of C = change in concentration of C/time reacquired for change = ∆C/∆T Procedure
Effect of Concentration
4 reaction solutions were prepared in 250 mL flask. In flask A 25.0 mL KI solution was added in a flask using a graduated cylinder. 2 separate 1 mL pipettes were used for measuring 1.0 mL Na2S2O3 and starch solutions (Department of Chemistry and Biology, 2012). These 2 solution were also added to flask A. 48.0 mL KNO3 solution was added into Flask A. 25.0 mL of (NH4)S2O8 (ammonium peroxodisulphate) was measured in a 100 mL dry beaker, named Beaker B respectively.
25.0 mL of (NH4)S2O8 was poured into flask A containing 25.0 mL of KI solution, and the time was noted to the nearest second at each aliquot. When the solution turned blue-black colour, time was recorded indicating that 2.0 x 10^-4 moles had reacted. Each time the color appeared the time was recorded as the next aliquot. Up to 5 aliquots were recorded to 10.0 x 10^-4 moles.
Solution 2, 3 and 4 were prepared in the same...
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