CHY211- Speed of Chemical Process

Speed of Chemical Process Introduction and Background Chemical kinetic is a branch of chemistry that study how fast or slow the rate of reaction can occur. However, according to the collision theory, when two particles of a reactant (molecules or ions) are collides with each other with sufficient energy their old bond will breaks and new bonds will forms that result in increasing the concentration and temperature of the reactant particle and thus increasing the rate of reaction. Therefore, there are four factors that could affect the rate of reactions in a solution such as the natural of the reactants, the concentration of the reactants, temperature and catalysis of the reactants. Concentration has main effect on the rate of the reaction; therefore, increasing the concentration of a solution will increases the number of particles of collision per unit time and hence increasing the rate of reaction. Also, the temperature is one of the important factors that affect the rate of the reaction. As the temperature increases, the kinetic energy of the particles will increase too and so the number of collisions between the particles increases. As a result, the rate of reaction increases. However, the rate of reaction depends on the properties of the substances that are reacting such as state of matter, bond type and molecular size. For example, the rate of reaction of gases tends to react faster than liquid or solid. On the other hand, catalyst plays a big role in increasing the rate of reaction by brings particles into close connection for reaction to happen. The purpose of this lab is to determine experimentally how fast or slow the rate of reaction can take place with respect to time and to observe the effect of the concentration and temperature factors on the reaction by using constant temperature water bath and stop watch. This lab; however, divided into two parts. Firstly, observe the influence of the catalyst on chemical kinetics by adding thiosulphate to KI solution. Secondly, determine the effect of reactant concentration on the rate of reaction by preparing and measuring the volume of four fresh solutions (each solution contains (NH4)2S2O8, KI solution, Na2S2O3, KNO3 and one drop of EDTA solution) into 250 mL Erlenmeyer flasks once at a time. When ammonium peroxodisulphate was added to the mixture, the time was recorded when the blue-black color appeared. Also, the effect of the temperature on chemical kinetic was observed by putting solution number one of the 25 mL of (NH4)2S2O8 into a constant temperature bath of about 40oC for 10 minutes and then record the temperature values of each solution to calculate the rate of reaction. The following equation was used to compute the order of reaction and rate constant; Rate=k [A]x [B]y Where [A] and [B] are the molar concentration x and y are the reaction order K is the rate constant The slope of the rate of reaction was calculated by using this formula; Rate= [S2O82-]/t Where [S2O82-] is the concentration of the peroxodisulphate ions t is the changing in time of rate of reaction in seconds In conclusion to all the equations that are shown, the calculations, the observation and graphs will prove each one in their own way. Experiment Details Caution Potassium iodide 0.2 M (KI) is minor irritant. Avoid contact with eye and skin, it can cause eye and skin irritation. Eye and Skin contact: wash carefully for 15 minutes. Starch Solution 1% minor irritation; avoid contact with eyes, clothing and skin. Sodium thiosulphatecan be harmful. May cause irritation to skin and eye. Skin and Eye contact: wash thoroughly with water. Ammonium Peroxodisulphate 0.2 M (NH4)2S2O8cause eye and skin irritation, avoid contact with eye and skin. Skin contact: wash with water for 15 minutes. Eye contact: wash for 15 minutes. Potassium Nitrate 0.2M (oxide) may cause eye and skin irritation. Keep the container tightly sealed. Skin and Eye contact: wash immediately for about 15 minutes. EDTA 0.1 M might cause a little irritation avoids contact with eye, clothing and skin. Worker Personal Safety The following stuffs have to be followed in order to prevent any injures during the lab in the future; Dress constraint (lab coat): to avoid any direct body contact to chemical that may be irritation and harmful. Eye protection: eye goggles should be wearing during the lab to avoid any eye contact with chemicals. Gloves: prevent skin contact with corrosive and any chemical substance that might be dangerous. Materials Stop watch Burette dropper Constant Temperature Water Bath 250 mL Erlenmeyer Flasks Test tube 10mL graduated cylinder 100 mL Beaker Stopwatch 10mL Graduated Cylinder Test Tubes Constant Temperature Water Bath 100 mL Beaker Dropper 250mL Erlenmeyer flask Burette and clamps Solutions Four solutions had been prepared in order to perform this experiment. For each solution different amount of volume had been measured and used for each of Ammonium peroxodisulphate, potassium nitrate and potassium iodide of concentration 0.2 M. Solution 1 Flask A: Five different kind and volume of solutions had mixed together in flask A. 25.0 mL KI solution 1.0 mL starch solution of concentration (0.2%) 1.0mL Na2S2O3 48.0 mL KNO3 of concentration (0.2 M) 1 drop EDTA solution of concentration (0.1 M) Beaker B: 25.0 mL (NH4) S2O8 Solution 2 Flask A: 25.0 mL KI solution 1.0 mL starch solution 1.0 mL Na2S2O3 23.0 mL KNO3 1 drop EDTA solution Beaker B: 50.0mL (NH4) S2O8 Solution 3 Flask A: 50.0 mL KI solution 1.0 mL starch solution 1.0 mL Na2S2O3 23.0 mL KNO3 1 drop of EDTA solution Beaker B: 25.0mL (NH4) S2O8 Solution 4 Flask A: 12.5mL KI solution 1.0 mL starch solution 1.0 mL Na2S2O3 35.5 mL KNO3 1 drop of EDTA solution Beaker B: 50.0 mL (NH4) S2O8 Waste disposal Solutions Inorganic: The solutions written below were disposal in inorganic waste. Potassium iodide (KI), potassium Nitrate (KNO3), starch solution, ammonium peroxodisulphate (NH4)2S2O8, and EDTA. The solutions written below was disposal in inorganic waste Organic Waste: None Procedure There are two parts that was obtained in this experiment. In the first part, the effect of catalyst on the rate of chemical reaction had investigated and observed when the color had changed into blue-black color by adding thiosulphate to KI solution in a test tube. In the second part, the effect of reactant concentration on the rate of reaction determined by preparing four fresh solutions in 250 mL Erlenmeyer flasks. Also, the effect of temperature on the rate of the reaction observed by placing solution one into different temperature of 40oc (constant temperature bath) for about 10 minutes and record the temperature when it’s constant for 3 minutes. Part I: Preliminary Experiments In the first part, the effect of catalyst on the rate of reaction was observed when the color of the solution changed by adding thiosulphate to KI solution. 5mL of 0.2 M KI solution with 10 mL of water were diluted in a test tube and then 3 drops of starch solution were added to the mixture and mixed carefully. Then, 5mL of 0.2 M ammonium peroxodisulphate solution was added and mixed to observe any change in color of the solution. The above step was repeated when the solution color had changed. 4 drops of 0.4 M sodium thiosulphate solution were added and mixed to note the effect of adding thiosulphate has on the color of the solution. Part II: Kinetics Experiments Part A: Effect of Reactant Concentration on Rate In the second part, the effect of the reactant concentration on the rate of reaction was determined by preparing and measuring accurately the volume of four fresh solutions into 250 ml flasks. Two burettes were set up and held by a clamp on a retort stand to measure correctly the volume of the KI and KNO3 solutions. Then, two separate 1 mL pipettes were used for measuring the volume of the Na2S2O3 (sodium thiosulphate) and starch solution. Ammonium peroxodisulphate solution (NH4)2S2O8 was prepared in burettes on the front bench. Solution 1(Flask A) was prepared in a 250 mL dry, clean Erlenmeyer. 25.0mL of (NH4)2S2O8 solution was measured into a clean, dry 100mL beaker (Beaker B). The 25.0mL of (NH4)2S2O8 was quickly poured into solution 1 (Flask A) and swirled vigorously. The time was noted to the nearest second when the solutions are mixed. The contents of the flask were swirled throughout the experiment. The stopwatch kept running for (Aliquot 1). 2 x 10-4 mol of S2O82-was reacted when the blue-black color appears, at that instant. The time was recorded immediately and then 1mL aliquot of Na2S2O3 solution was quickly added and swirled till the color disappeared. The time was recorded for the reappearance of the blue-black color. Another 1mL aliquot of Na2S2O3 was added. The above producer was continued for five aliquots added to solution 1(Flask A). The above steps were repeated for each of solution 2, 3 and 4 except that 50.0mL of (NH4)2S2O8 solution (Beaker B) was added to Flask A of solutions 2 and 4 and 25.0mL of (NH4)2S2O8 solution was added to Flask A of solution 3. The temperature of each of the solutions was recorded. Part B: The Effect of Temperature on Rate In the second (B) part, the rate of reaction was conducted by determine the effect of the temperature on the rate using only solution one at a different temperature. Solution 1 and the 25.0mL of (NH4)2S2O8 solution were placed in a constant temperature bath of approximately 40oC for 10 minutes. Then, the temperature of each solution was recorded. (NH4)2S2O8 was added as quantitatively as possible to solution 1 when the temperatures were constant for 3 minutes. The time of addition of (NH4)2S2O8 was noted. Solution 1 was left in the water –bath for 10 minutes. The procedure used in part II (A) was repeated. Moles of S2O8 2- consumed Versus Time (s) Graph of solution 1 Moles of S2O8 2- Reacted Versus Time (s) Graph of Solution 2 Moles of S2O8 2- Reacted Versus Time (s) Graph of Solution 3 Moles of S2O8 2- Reacted Versus Time (s) Graph of Solution 2 Moles of S2O8 2- Reacted Versus Time (s) Graph of Solution 1 at (40oC) Discussion and Results The purpose of this lab was to determine the effect of each of the temperature, catalyst, natural of the reactant and concentration on the rate of reaction. Also, to investigate how fast the chemical kinetic can occur as each of the temperature and concentration increases. However, the collision theory states that, when the particles of the reactant in the reaction collide with each other will tend to increases the kinetic energy between the molecules, breaking the old bonds and forms new bonds, thus increasing the rate of the reaction. Though, during the lab, it was conduct that the speed of the rate of reaction depends on the number of particles that collides per unit volume. In part II (A) solution one, for example, the rate of reaction increases with respect to the time as the concentration of Na2S2O3 solution increases. In addition, when the solution in Beaker B was added to the solution mixture (KI, KNO3, Na2S2O3, starch and EDTA solution) in flask A, as the time was noted, it observed that the solution color changed from colorless to blue-black color when the solutions were swirled which indicated that the collision happened between particles thus increases the rate of reaction that evidence of the modification of the collision theory. Also, as evidence from the graph, it was investigated that the concentration of the rate is directly proportional as the slope of the best line fit was increasing linearly. In part I; however, the effect of catalyst on the speed of the reaction was investigated, as the color of the solution changed to blue-black color when sodium thisoulphate was added to the solution of KI, ammonium peroxodisulphate and water. Thus, catalyst is an important factor that increases the rate of reaction without causing any chemical changes. Catalyst lower the activation energy at a given temperature that because large fraction of molecules with enough energy had reacted and hence the catalyst increases the rate of forward and backward reaction. The activation energy can be computed by this equation; K=A e-EA/RT Where k is the rate constant A is constant E is the natural logarithms (e=2.718) R is the ideal gas constant (8.314J/mole K) T is the absolute temperature in K In the second part of the experiment, the effect of the temperature on chemical kinetic was observed as the temperature increase to 40oc using constant temperature water bath to solution one , the rate of reaction occur very fast compared with the temperature of solution one in part II (A) . However, from the data table, it was noticed that temperature has an influence on increasing the speed of the rate. The rate law equation shown below illustrations that the rate of the reaction depends on the concentration of the substance; Rate=k [A]x [B]y Where k is the magnitude of the rate constant that depend on the temperature. [A] And [B] are the concentration of the substances x and y are the order of the rate of the reaction The error that was observed in this experiment was that the rate of reaction not only depends on the concentration of the substance but also on the order of the reacted. During the calculation of the x and y value from the rate law formula, it was observed that as the value of y is equal to zero the concentration of the substance does not change. That means that the reactant with zero order, the concentration has no effect on the rate of the reaction. Also, it had been noticed that the value of the rate constant k varies with the temperature. In the solution one, the temperature at 40oC k is equal to 9.36 x 10 -3 is higher than at temperature 22.5 oC. That concludes, as the temperature of the rate of the reaction increase the rate constant increases too. By looking at the result, it had noticed and observed that the speed of reaction increases as the number of particles of the reactant that collide increases. Thus, this concludes that the rate of reaction is directly proportional to the rate constant, temperature, catalyst and concentration of the reaction. Conclusions: The rate of reaction is the result of the collision of the particles with each other with sufficient energy that occurred because of the breaking particles or molecules old bonds and forms new one. Otherwise, as the number of the particles that collides with enough energy increases, the temperature, kinetic energy between molecules and concentration increases thus, the speed of the reaction increases. However, due to the calculations, graphs and results that conduct in this lab, experiment shows that the rate of reaction on a solution depends on four factors such the natural of the reactant, the presence of the catalyst, temperature and concentration of the reactant. Determining the effect of the temperature on the rate of the reaction by using the same solution (solution one) using different temperature value had illustrated that the rate constant is directly proportional to the temperature. Thus, as the temperature of the substance increases, the rate of the reaction increases too. Overall, as the concentration of the substance with order of reaction one, two and three etc. increases, the speed of the reaction increases also. Catalyst is a substance and another factor that has effect on increasing the rate of the reaction by lowering the activation energy and speed the chemical kinetic processes without causing any chemical changes just causing physical changes such as the changing the color of the solution mixture. With the rate equation, Rate=k [X]m [Y]n , it was evidence that the relation between the rate of the reaction and concentration is linear and directly proportional . With the value of the rate constant k that derived and measured in this experiment in which k is increasing with increasing the temperature was as expected and this showed that the data that was computed in this lab was fairly accurate with small amount of uncertainly. References: General Chemistry (CHY211) Laboratory Lab Manual: Chemical Kinetic https://www.cdli.ca/sampleResources/chem3202/unit01_org01_ilo03/b_activity.html http://avogadro.chem.iastate.edu/MSDS/(NH4)2S2O8.html http://www.sasvrc.qld.gov.au/PersonalSafety.html http://sodiumthiosulfate.net/ http://en.wikipedia.org/wiki/Chemical_kinetics http://en.wikipedia.org/wiki/Collision_theory Calculations: The slope of rate of reaction has been computed from the graph as shown below: Solution 1 Convertion: 100mL to L 100 mL 1L =0.1 1000mL Slope= S2O82-/T = (10.0-2.0) x 10-4 mole (20.4-3.38) sec =4.68 x 10-5 mole .sec -1/0.1L = 4.68 x 10-4 mole /sec .L Solution 2 Slope= S2O82-/T = (10.0-2.0) x 10-4 mole (8.92-2.20) sec =1.19 x 10-4 mole .sec -1/0.1L = 1.19 x 10-3 mole /sec .L Solution 3 Slope= S2O82-/T = (10.0-2.0) x 10-4 mole (11.8-2.55) sec =8.48 x 10-5 mole .sec -1/0.1L = 8.48 x 10-4 mole /sec .L Solution 4 Slope= S2O82-/T = (10.0-2.0) x 10-4 mole (14.49-4.02) sec =7.68 x 10-5 mole .sec -1/0.1L = 7.68 x 10-4 mole /sec .L Solution 1 T2 Slope= S2O82-/T = (10.0-2.0) x 10-4 mole (6.65-1.56) sec =1.57 x 10-4 mole .sec -1/0.1L = 1.57 x 10-3 mole /sec .L