Reaction Rates and Equilibrium OBJECTIVE: To acquire an understanding of the factors that influence reaction rates and equilibrium conditions. INTRODUCTION: It is often useful to manipulate chemical systems in order to achieve optimal results for a given process. An understanding of the factors that influence the rate of a chemical reaction as well as those influencing the equilibrium position can make it possible to do reactions that are useful but were impractical. An example is the development of the Haber process for fixing nitrogen. Nitrogen gas is abundant but quite inert, making it difficult to convert the abundant atmospheric nitrogen into nitrogen compounds for biological organisms and industrial processes. Prior to 1905, the nitrogen compounds used to manufacture fertilizer and explosives came primarily from natural nitrate-containing deposits. In 1905 ammonia was first synthesized for atmospheric nitrogen and hydrogen, using the Haber process. The reaction itself appears simple:

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N2(g) + 3H2(g)⇔ 2NH3(g) However, the process was made practical only after an understanding was gained of the factors that influence the reaction rate and equilibrium conditions. Kinetics In order for a chemical reaction to occur several conditions must be met: the reactants must collide with each other, they must do so in the proper orientation and the collision must occur with enough energy to break and reform chemical bonds. Any change to the reaction system that increases the likelihood of these conditions being met will increase the rate of the reaction. Higher concentrations of reactants increases the probability of molecular collisions, higher temperatures both increase the probability of collision and the energy of the collision. Changing reaction rates and temperatures therefore has an impact on the reaction rate. The addition of substances called catalysts is another way to increase the rate of the reaction. Catalysts provide an alternate, faster pathway for a reaction by forming intermediate compounds with one or more of the reactants (homogeneous catalysts) or providing surfaces on which the reactants adsorb and react more easily (heterogeneous catalysts). The catalyst significantly increases the rate of the chemical reaction by decreasing the activation energy but is not itself used up in the reaction. The iodine clock reaction is a convenient reaction for observing kinetics. The reaction is between potassium iodate, KIO3, and sodium bisulfite, NaHSO3; the net ionic reaction is given by the following equation.

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IO3−(aq) + 3HSO3

−(aq)⇔ I−(aq) + 3SO42−(aq) + 3H +(aq)

Equilibrium In many chemical reactions the reactants are not totally converted to the products because of a reverse reaction - the products react to form the original reactants. Such reactions are said to be reversible. The reaction proceeding the right is called the forward reaction; that to the left, the reverse reaction. Both reactions occur simultaneously. When the rate of the forward reaction is equal to the rate of the reverse reaction, a condition of chemical equilibrium exists. At equilibrium the products react at the same rate as they are produced. Thus the concentrations of the substances in equilibrium do not change, but both reactions, forward and reverse, are still occurring. Le Chatelier’s Principle relates to systems in equilibrium and states that when the conditions of a system in equilibrium are changed the system reacts to counteract the change and reestablish equilibrium. In this experiment we will observe the effect of changing the concentration of one or more substances in a chemical equilibrium. Consider the hypothetical system:

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A + B⇔C + D When the concentration of any one of the species in this equilibrium is changed, the equilibrium is disturbed. Changes in the concentration of all the other substances will occur to establish a new position of equilibrium. For example, when the concentration of B is increased, the rate of the forward reaction increases, the concentration of A decreases, and the concentration of C and D increase. After a period of time the two rates will become equal and the system will again be in equilibrium, even though the concentrations of each individual component have been altered from their initial values. Evidence of a shift in equilibrium by a change in concentration can easily be observed if one of the substances involved in the equilibrium is colored. The appearance of a precipitate or the evolution of a gas can sometimes be used to detect a shift in equilibrium. The net ionic equation for the equilibrium system to be studied is given below. Iron (III) chloride with potassium thiocyanate:

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Fe3+(aq) + SCN−(aq)⇔ Fe(SCN)2+(aq)(yellow) (clear) (red)

PROCEDURE: Influences on Reaction Rates We can monitor the rate of reaction by the disappearance of the bisulfite in the iodine clock reaction to determine the rate of the reaction. We do so by adding more IO3

- than HSO3

- at the start of the reaction. When we have used up all the bisulfite, there is still some iodate left. This will then react with the product iodide, I-, and results in the formation of I2.

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IO3(aq) + 5I−(aq) + 6H +(aq)⇔ 3I2(aq) + 3H2O(l) We can detect the appearance of iodine with the aid of starch indicator; this reagent forms a blue complex with iodine. The time it takes for the blue color to suddenly appear indicates when all the bisulfite was used up in the first reaction. That’s why the name: iodine clock. Thus you should measure the time (with a stopwatch) elapsed between mixing the two solutions and the appearance of the blue color. Make up two solutions, A and B. Solution A should be made up according to the table below. For Solution B, make up nine identical solutions of 5 mL 0.025 M NaHSO3, 2 mL 0.5% (m/v) starch, and 18 mL water. Place the reactants in two separate 250-mL beakers according to the outline below (volumes given are in mL). Simultaneously pour the two reactants into a third beaker and time the appearance of the blue color. Repeat the experiment with the other trial conditions. In order to fully analyze the influence of temperature on reaction rate, room temperature should be recorded. It is assumed that the solutions are initially in thermal equilibrium with the room air, determine room temperature in order to record the temperature for those reactions. Place solutions in appropriate waste container once the reaction is complete. Trial 0.1 M KIO3 Water 0.1 M H2SO4 Temperature 1 10 15 0 Room 2 8 17 0 Room 3 6 19 0 Room 4 4 21 0 Room 5 2 23 0 Room 6 6 19 0 80°C 7 6 19 0 40°C 8 6 19 0 10°C 9 6 17 2 Room Table 1. Quantities of reactants needed in solution A for each trial. Influences on equilibrium positions

Set up 8 clean, small test tubes in a test tube rack. To each of the test tubes add 20 drops of 0.01 M Fe(NO3)3, 20 drops of 0.01 M KSCN, and 20 drops of water. Note the color. Do not add any additional chemicals to the first test tube since it will be used as a control in order to accurately determine color change upon addition of a stressor. After addition of the following, note any change in appearance of the solution relative to the blank. Place solutions in appropriately labeled waste container. Tube Solution added 2 10 drops 0.1 M Fe(NO3)3 3 10 drops 0.1 M KSCN 4 5 drops 3 M NaOH 5 5 drops 3 M HCl 6 5 drops 0.1 M Na2HPO4 7 Heated in hot water bath for 10 minutes 8 Cooled in ice bath for 10 minutes