Chemical Kinetics of bromate-bromide reaction Introduction Chemical reactions occur all the time and interfere in our everyday life in countless aspects. In order to understand how reactions (or in general, processes) take place, it is essential to study chemistry. For any process there are two main fields of study: Thermodynamics which is concerned with the direction a process will take and all the energy changes between the initial and final states and Kinetics which deals with how fast a certain process will be. In general, chemical reactions occur through the breaking and formation of chemical bonds between atoms in the compounds involved in the reactions. Experimental investigation reveals that different reactions take place at different speeds, or rates. In this context the term “rate” is taken to mean the quantity of a specific reactant which disappears or product which appears in unit time at a given temperature. Reaction rates thus expressed have units (mole-1 s-1). For example, the following reaction + →+ (1) will have the following rates’ relationships −1 [] = −1 [] = 1 [] = 1 [] (2) And any expression of these is related to the rate of a reaction as a whole. If a chemical reaction is to take place between two reactive species, these species must collide with each other in order that bond breaking-forming can take place and lead to the formation of products. There are different factors affecting number and effectiveness of such collisions such as number of reactant molecules in the reaction vessel, the temperature at which the reaction is taking place and the various orientations of the molecules during the collisions. These factors, as a result, will have their effect on the rate of the reaction. Effect of Concentration on Reaction Rates Rates of reactions are affected by changing the concentrations of reacting species since increasing the concentration will lead to higher number of collisions. The quantitative influence of concentration on the rate is best determined experimentally and it can be shown that for kinetically simple cases, the rate is proportional to some power of the concentration of a given reactant. For example, for our earlier reaction between A and B it might be shown in separate experiments that, ∝ [] ∝ [] where x and y are the order of reaction rate corresponding to reactant A and order of reaction rate corresponding to reactant B. the overall order of the reaction is x+y. the order of the reaction – especially for simple kinetics – is related to the number of species in the rate determining step (RDS) which is the slowest step in the reaction mechanism. We can therefore conclude that for such a reaction, = [] [] (3) Where k is the constant of proportionality and is called the rate constant for the reaction. Effect of Temperature on Reaction Rates The rates of most reactions are known to increase significantly with rising temperature. A good rule of thumb is that the rate roughly doubles for every 10 oC rise. There are exceptions to this but on the whole it is a fair guide. In 1889 Arrhenius proposed that the temperature dependence of the rate of reaction is governed by the equation which now bears his name, = − ⁄ (4) Where k is the rate constant, Ea is the activation energy, R is the gas constant and T the temperature in degrees Kelvin. The pre-exponential term A is the property of the particular reaction related to the collision frequency of the reactive species. We might thus expect A to be itself temperature dependent and this is indeed the case. However, in equation (4), the dependence of k on temperature is dominated by the strong exponential term, so, in the analysis of experimental data, the dependence of A on temperature is usually ignored as a first approximation. Equation (4) can be rewritten in logarithmic form, = − (5) A graph of lnk against 1/T is therefore expected to be linear with negative slope given by –Ea/R. Measurement of a reaction rate at various temperatures therefore provides us with a means of determining the activation energy for the reaction. The reaction to be studied in this experiment is between bromate and bromide ions in acidic medium and occurs according to the equation. In an acidic medium the bromated ion oxidizes the bromide ion to give bromine and water as the following reaction 3− + 5 − + 6 + → 3 2 + 3 2 This reaction is relatively slow and as a result the kinetics of this reaction can be studied using simple methods. The only difficulty in studying this reaction is how to determine the concentration changes for this reaction. One method overcome this difficulty is introducing other reaction by which exact amounts can be determined. If a small known quantity of phenol and an indicator are added to the reaction mixture, the liberated bromine will react very quickly with the phenol to produce tribromophenol: 6 5 + 2 → 6 2 + + + − Once all phenol has been brominated a slight excess of Br2 is sufficient to immediately decolorize the methyl orange indicator. Thus the quantity of phenol added to the reaction represents the amount of 3− consumed during the time elapsed between the mixing of the reactants until the disappearance of the indicator. The overall reaction becomes: The rate equation of the bromate-bromide reaction takes the form: −[3− ] = [3− ] [ − ] [ + ] (6) In this experiment the order with respects to the bromate, bromide and H+ will be determined. In addition the rate constant at two different temperatures will be evaluated to calculate the activation energy of this reaction, Experimental: Determination of the reaction orders The method of initial rates is used to determine the order for each reactant. This method involves measuring the rate of reaction at very short times before any significant changes in concentrations of reactants occur. Two or more kinetic experiments are examined in which only one reactant concentration is change while the other remains constant and then investigate the change in the rate of the reaction. Experimentally, prepare the following mixtures in two test tubes using accurate volume measuring glassware: Test tube 1 3− Test tube 2 0.33 M − 0.62 M 3 Water 1 10.0 mL 10.0 mL 20.0 mL 2 10.0 mL 20.0 mL 3 20.0 mL 4 10.0 mL # + Phenol Methyl 0.50 M 0.06 M orange 25 mL 20.0 mL 10.0 mL 5.0 mL 7.0 mL 28 mL 20.0 mL 10.0 mL 5.0 mL 10.0 mL 14.0 mL 21 mL 20.0 mL 10.0 mL 5.0 mL 10.0 mL 0.0 mL 25 mL 40.0 mL 10.0 mL 5.0 mL 0.50 M As an example, for the first run, mix the amounts of reagents indicated in the two test tubes, inside the fume hood mix the content of the two tubes and start the stopwatch. Once the pink color of the methyl orange disappears stop the stop watch and record the time. Dispose of all solutions in the waste bottle inside the fume hood. In the same manner perform the other trials. Examining the table, we can see that trials 1+2 used to determine the exponent y for − , trials 1+3 used to determine the exponent x for 3− and trials 1+4 used to determine the exponent z for + . Determination of the activation energy of the reaction Repeat the first trial in the previous part a two other temperature, such as in ice bath (T= 0oC) and in thermostatic water bath set at T= 35 oC. Doing so, you will end up with trial one performed at T= room temp, 0 oC and 35 oC. Determination of the Specific rate constant In order to determine the specific rate constant, prepare the following mixtures in two test tubes using accurate volume measuring glassware. These mixtures are identical except for the amount of the phenol added. Mix the contents of the two test tubes and measure the time required to decolorization to occur. Test tube 1 Run 3− Test tube 2 0.33 M − 0.62 M 3 0.50 M 1 10.0 mL 10.0 mL 2 10.0 mL 3 + Phenol Methyl 0.50 M 0.06 M orange 20.0 mL 20.0 mL 4.0 mL 5.0 mL 31.0 mL 10.0 mL 20.0 mL 20.0 mL 6.0 mL 5.0 mL 29.0 mL 10.0 mL 10.0 mL 20.0 mL 20.0 mL 8.0 mL 5.0 mL 27.0 mL 4 10.0 mL 10.0 mL 20.0 mL 20.0 mL 10.0 mL 5.0 mL 25.0 mL 5 10.0 mL 10.0 mL 20.0 mL 20.0 mL 12.0 mL 5.0 mL 23.0 mL # Water Calculations: Exponents of the rate equation Calculate the initial concentrations of the reactants. Each mole of 3− reacts with bromide − and acid to give 3 moles of bromine (Br2). Now each mole of bromine Br2 reacts with one mole of phenol C6H5OH producing bromophenol. Assume the concentration of phenol to be 1.0 M, then the methyl red will be decolorized when a/3 M of 3− are consumed. Assuming that in the initial stages of the reaction the decrease in 3− concentration is linear, we can write: = −[3− ] = (/3) = [3− ] [ − ] [ + ] (7) where t is the time (in seconds) needed for the decolonization of the indicator, and a is the initial concentration of 3− . Assuming that the exponents are small whole numbers, use the results of your experiments to find x, y and z and write the whole rate equation for this reaction. Comment on the rate equation and suggest a suitable mechanism for the reaction depending on this rate equation. Specific rate constant In order to obtain the specific rate constant (k), we need to represent amounts 3− consumed with time. A plot of 0.033-(a/3) vs time helps to obtain an initial value for −[3− ] . The initial rate can be determined as the slop of the straight line tangent to the curve at the beginning of the curve (see the Figure 1). Knowing the initial concentrations of the reactants and the rate equation (order of each reactant) the value of k can be calculated. Finally, the sodium nitrate is added to the mixtures to have a constant ionic strength for all experiments, you should calculate the ionic strength of your reaction mixtures and state it with your answer. Question: do you think that k could vary by varying the ionic strength? Figure 1: schematic graph representing the relation between the amount of bromate consumed with time. Activation energy According to Arhenius equation (2), there is an exponential relationship between the rate constant and temperature. Taking the logarithm form (equation (3)) plot lnk vs 1/T where T in Kelvin, we should obtain a linear line with slope equal –Ea/R and the intercept is lnA. Obtain both Ea and A for such a graph. Show that under the conditions used in this experiment, the value of t is inversely proportional to the specific rate constant k. Chemical Kinetics of bromate-bromide reaction Data sheet Name : Partner’s name: Lab section: Part A: Estimating the reaction order Volume of Volume of Volume of − − ( )M ( )M ( )M Ambient temperature= Volume of Water Volume of Volume of Volume of + Phenol Methyl ( )M ( )M ºC Time orange 1 2 3 4 Part B: Estimating activation energy Volume of Volume of Volume of − − ( )M ( )M ( )M temperature 2= Volume of Water Volume of Volume of Volume of + Phenol Methyl ( )M ( )M ºC Time orange 1 temperature 3= Volume of Volume of Volume of − − ( )M ( )M ( )M Volume of Water Volume of Volume of Volume of + Phenol Methyl ( )M ( )M ºC Time orange 1 Part C: Estimating the rate constant Volume of # ( − )M Volume of ( − )M Volume of ( )M Ambient temperature= Volume of Water Volume of ( + )M Volume of Volume of Phenol Methyl ( )M orange 1 2 3 4 Instructor’s signature: Date: ºC Time

© Copyright 2020