To determine the rate law for a chemical reaction among hydrogen peroxide, iodide and acid, specifically by observing how changing each of the concentrations of H2O2, I- and H+ affects the rate of reaction,and examine the effects of temperature and a catalyst on the rate of reaction.
A description of the sequence of steps by which a chemical reaction takes place is called a reaction mechanism. Many studies go into trying to determine possible mechanisms. The rates at which reactions take place provide important insights.
The rate law for a chemical reaction is a quantitative expression involving constants related to the nature of the chemical reaction and the concentrations of reactants. In order for reactants to react, they must come in contact with one another (or at least come nearby). The probability of a collision is related to a function of the concentration. So, for the general rate law: Rate=k [A]a[B]b[C]c
[A], [B] and [C] represent concentrations of reactants and catalysts, and a, b, and c represent exponents that may or may not be related to the coefficients of the corresponding balanced chemical equation.
The quantities in brackets are read as concentration and are raised to an appropriate power. Multiplied together with the constant (k), they give the rate of the reaction. The numerical values of a, b, and c must be determined by experimentation. These numbers determine the order of the reaction. Added together they give the over-all order of the reaction. It was the purpose of this experiment to determine the order of H2O2, a reactant in the iodine clock reaction.
The reaction to be studied in this experiment was the acid buffered oxidation of iodide to triiodide by hydrogen peroxide H2O2 (aq) + 3 I- (aq) + 2 H+ (aq) → I3- (aq)+ 2 H2O (aq) --- (I)
I3- (aq) + 2 S2O32- (aq) → 3 I- (aq) + S4O62- (aq) --- (II)
2 I3- (aq) + starch → starch-I5-complex + I- (aq) --- (III)
The first equation indicates that, in an acidic solution, iodide ions were oxidized by hydrogen peroxide to triiodide ions. These triiodide ions were reduced back to iodide ions by thiosulfate ions, as indicated in equation (II). This reaction was much faster than the reaction of equation (I); it consumed triiodide ions as fast as they were formed. This prevented any readily apparent reaction of equation (III). However, after all the thiosulfate ions had been consumed by the reaction of equation (II), triiodide ions reacted with starch to form the blue starch-pentaiodide complex.
By varying the concentration of each of the three reactants (H2O2, I- and H+), we will be able to determine the order of the reaction with respect to each reactant and the rate law of the reaction, which of the form: Rate = k [H2O2]x[I-]y[H+]z
By knowing the reaction times (Δt) and the concentrations of H2O2 of two sepatate reaction mixtures (mixtures A & B), the reaction order of H2O2, x, can be calculated. x = log(Δt2/Δt1) / log ([H2O2]1/[H2O2]2)
The same method is used to obtain the reaction order with respect to I- (mixture A & C), and H+ (mixture of A & D).
Part I – Standardization of H2O2 Solution
1. A burette was rinsed with deionized water and 0.05M Na2S2O3 solution. 2. The stopcock of the burette was closed and the sodium thiosulphate solution was poured into it until the liquid level was near to the zero mark. The stopcock of the burette was opened to allow the titrant to fill up the tip and then adjust the liquid level near zero. 3. The initial burette reading was recorded in Table 1. The reading should be accurate to 2 decimal places. 4. 1.00cm3 of the ~0.8M H2O2 solution was pipetted into a clean 125 cm3 conical flask. 5. 25cm3 of the deionized water was measured with a 50cm3 measuring cylinder. It was poured into the conical flask. 6. 10cm3 of 2.0M sulphuric acid was measured with a 10cm3 clean measuring cylinder. It was poured into the conical...