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Experiment No. 1 CALORIMETRY When a chemical reaction occurs, heat is exchanged between the system and the surroundings. This heat exchange is called the heat of reaction, qrxn. In an exothermic reaction, heat is released from the system to the surroundings (qrxn < 0). On the other hand, heat is absorbed by the system from the surroundings (qrxn > 0) in an endothermic reaction. At constant pressure, the heat of reaction of is equal to the enthalpy change, ? H, which is measured through calorimetry.

This process includes the use of a calorimeter, a device that traps the heat absorbed or released. In this experiment, a calorimeter was made by inserting a test tube with a cork into a styroball. The temperature was measured by allowing NaOH to react with HCl. After calibrating the calorimeter, the heats of reaction for given reaction systems were then measured. This experiment aims to understand how calorimetry is done by creating a self-made calorimeter. Answers to Questions: Calibration of the calorimeter: ) Give the net ionic thermochemical equation of the reaction used to calibrate the calorimeter NaOH(aq) + HCl(aq) > NaCl(aq) + H2O(l) OH-(aq) + H+(aq) > H2O(l) a. Is the reaction endothermic or exothermic? The reaction is endothermic. b. Which is the limiting reactant? The limiting reactant is HCl. c. How much (in moles) limiting reactant was used? 5. 00 mL x 1L1000mLx 1M=0. 005 mol d. How much heat was generated (or absorbed) by the reaction? -157 J0. 005 mol=-31. 4 KJ 2) Relate the sign of the ? T to the ? H of the reaction used for the calibration.

As the temperature increases, the enthalpy also increases. This is due to the fact that enthalpy is proportional to temperature. If ? T is positive, then ? H is also positive and vice versa. 3) What is the heat capacity of the calorimeter? Relate its sign to the sign of ? T. qrxn = -qcalCcal= qcalm? T= 157 J(15 g)(2. 50°C) =4. 19 J/g°C -157 J = 157 JA positive ? T gives a positive heat capacity. ?T = 2. 50°C 4) In the appendix, show the derivation to obtain the equation used to calculate the heat capacity of the calorimeter. Determination of Heats of Reaction: ) Give the net ionic reaction for each reaction. | Reaction System| Net Ionic Equation| 1| 10 mL of 1M NH3 + 5 mL of 1M HCl| NH3(aq) + H+(aq) > NH4+(aq)| 2| 10 mL of 1M NaOH + 5 mL of 1M CH3COOH| OH-(aq) + H+(aq) > H2O(l)| 3| 10 mL of 1M NH3 + 5 mL of 1M CH3COOH| NH3(aq) + H+(aq) > NH4+(aq)| 4| 10 mL of 1M NaoOH + 5 mL of 1M HNO3| OH-(aq) + H+(aq) > H2O(l)| 5| 15 mL of 1M HCl + 0. 05 g Mg| Cl-(aq) + Mg(s) > MgCl(aq)| 6| 15 mL of 1M CH3COOH + 0. 05 g Mg| CH3COO-(aq) + Mg(s) > MgCH3COO(aq)| 7| 15 mL of 1M CuSO4 + 0. 5 g Zn| SO42-(aq) + Zn(s) > ZnSO4(aq)| 8| 10 mL of 0. M Na2CO3 + 5 mL of 1M CaCl2| Na2CO3(aq) + Cl2(aq) > 2NaCl(aq) + CO32-(aq)| 6) Determine the limiting reactant and the amount of the limiting reactant in each of the reaction performed. Net Ionic Equation| Limiting Reactant| NLR (mol)| NH3(aq) + H+(aq) > NH4+(aq)| HCl| 5×10-3| OH-(aq) + H+(aq) > H2O(l)| CH3COOH| 5×10-3| NH3(aq) + H+(aq) > NH4+(aq)| CH3COOH| 5×10-3| OH-(aq) + H+(aq) > H2O(l)| HNO3| 5×10-3| Cl-(aq) + Mg(s) > MgCl(aq)| Mg| 2×10-3| CH3COO-(aq) + Mg(s) > MgCH3COO(aq)| Mg| 2×10-3| SO42-(aq) + Zn(s) > ZnSO4(aq)| Zn| 8×10-3|

Na2CO3(aq) + Cl2(aq) > 2NaCl(aq) + CO32-(aq)| CaCl2| 5×10-3| 7) Calculate for the theoretical and experimental enthalpy of each reaction. a. Determine whether the reaction is endothermic or exothermic. b. Give the % error of your experimental values. | Experimental ? H (kJ/mol)| Theoretical ? H (kJ/mol)| % Error| Endothermic or Exothermic| 1| 35. 2| | | Endothermic| 2| 28. 0| | | Endothermic| 3| 38. 4| | | Endothermic| 4| 50. 2| | | Endothermic| 5| 276| | | Endothermic| 6| 281| | | Endothermic| 7| 8. 13| | | Endothermic| 8| -2. 6| | | exothermic| 8) Relate the sign of ? T to the sign of the experimental ?

H. As seen in the data collected, of ? T is positive then ? H is also positive. The energy in the system depends on the temperature, so as the temperature rises then energy in the system also rises. This is also true for a decrease in temperature which leads to a decrease in the energy thus resulting in a negative ? H. 9) For reactions 1-4, which pair gave the most and the least exothermic (or endothermic) reaction? Explain the observation. (Use both theoretical and experimental values to answer this. ) 10) For reactions 5 and 6, which gave the most exothermic reaction? What is the theoretical yield of each? 1) For reactions 7 and 8, what are the solid products of the reaction? What is the theoretical yield of each? 12) Reaction 8 is a synthesis (combination) reaction. Using the theoretical ? H of the reaction, deduce the relative magnitude of the energy of bond breaking and bond formation during the reaction. 13) In the appendix, show the equation used to determine the heats of reaction for reactions 1-6 and reactions 7 and 8. 14) Tabulate the possible sources of errors and their effect to the following parameters: ? T, Ccal, and ? H. accompany each with a reason. Source of error| Effect on ?

T| Effect on Ccal| Effect on ? H| Wrong concentration of solution| Can increase/decrease| Can increase/decrease| Can increase/decrease| Failing to immediately cover the test tube| Decrease| Decrease| Decrease| Wrong calibration| Wrong measurement| Wrong measurement| Wrong measurement| Appendix: References: Petrucci et al. (2011). General Chemistry: Principles and Modern Applications. 10th Edition. Pearson Education, Canada, 2011. Academic Group – Institute of Chemistry (2011). General Chemistry II– Laboratory Manual. Institute of Chemistry, University of the Philippines—Diliman. Name: Marion Clarisse L.

GelidoDate Performed: April 19, 2013 Co-workers: ClassDate Submitted: April 22, 2013 Experiment No. 2 COUPLED REACTIONS A non-spontaneous reaction (? °Grxn > 0) needs a spontaneous reaction (? °Grxn < 0) to achieve an overall spontaneous reaction. These reactions are known as coupled reactions. It is suitable to consider them as a single system, where the sum of each of the ? °Grxn of the pair is equal to the negative ? °Grxn of the system. In this experiment, two slabs of dry ice were used wherein a hole was gauged in the middle with four canals running to the edges of the ice.

The cavity was filled with Mg ribbon which was then ignited, after which it was immediately covered with the other slab. The set-up glowed brightly for a few seconds, and it was observed after that C(s) and MgO(s) were formed from the reaction. This experiment aims to study how coupled reactions work. Answers to Questions: 1) Give the balanced equation for the observed reaction in the experiment. 2Mg(s) + CO2(g) > 2MgO(s) + C(s) 2) Describe the physical properties of the products produced and relate it to the balanced equation of the reaction.

Black and white substances were formed from the experiment. Based on the equation of the reaction, the white substances were the MgO(s) while the black substances were carbon. There were more MgO(s) produced, which agrees with the balanced equation obtained. 3) What is the (approximate) Gibb’s Free Energy of the reaction observed? Assume that the reaction temperature is about 1000K. Use the Ellingham diagram provided in the manual to calculate for this. The approximate Gibb’s Free Energy is -520 kJ at 1000K. Since ? °G is less than zero, this means that the reaction was spontaneous. ) Why should the Mg ribbon be filed prior to use? It is important for the Mg ribbon to be filed to ensure that no MgO is included in the experiment which could lead to errors. 5) What factors contributed for the slow/delayed ignition of Mg ribbons? Explain how each retarded the ignition. The ice chamber provided a cool temperature which contributed to the slow ignition of the Mg ribbons. Some of the heat that was used to light up the Mg was absorbed by the dry ice. For the Mg ribbons to light up, much heat was required to get the reaction started.

The presence of CO2 also contributed to the delayed ignition the lessened amount of oxygen present in the environment. 6) What is the possible side reaction in the experiment? Give the balanced chemical equation for this side reaction. 2Mg(s) + O2(g) > 2MgO(s) This is because of the presence of oxygen in the environment which could react with Mg. 7) What is the theoretical combined mass of the products in the experiment? Compare this to the one obtained in the experiment. Account for any difference. The mass of the products was not obtained. 8) What are the possible sources of errors in the experiment?

Explain their effect to the yield. Improper filing/not filing of the Mg ribbons could affect the mass of the products. Another source of error is if the other slab of ice was not immediately used to cover the reaction which could lead to Mg reacting with oxygen, creating more MgO(s) products. References: Petrucci et al. (2011). General Chemistry: Principles and Modern Applications. 10th Edition. Pearson Education, Canada, 2011. Academic Group – Institute of Chemistry (2011). General Chemistry II– Laboratory Manual. Institute of Chemistry, University of the Philippines—Diliman.

Name: Marion Clarisse L. GelidoDate Performed: April 19, 2013 Co-workers: Anna Ebuen, Nikki Macasaet, Bianca BautistaDate Submitted: April 22, 2013 Experiment No. 3 CHEMICAL KINETICS Chemical kinetics is involved with how chemical reactions are measured, how they can be predicted, and how to deduce possible reaction mechanisms from the reaction-rate data. There are three factors that affect the rate of reaction: concentration, temperature and the presence of a catalyst. The experiment was divided into three parts, each designed to investigate the different factors that affect the rate of reaction.

Answers to Questions: Determination of Rate Law of the Reaction Between Thiosulfate and Hydronium Ion 1) Give the balanced chemical equation for the reaction under study. Na2S2O3(aq) + 2HCl(aq) > 2NaCl(aq) + S(s) + SO2(g) + H2O(l) 2) Describe the physical properties of the products produced and relate it to the balanced equation of the reaction. After a certain amount of time, it was observed that the X mark was no longer visible due to the cloudiness produced by the solution. This cloudiness was caused by the sulfur that precipitated which agrees with the balances equation of the reaction. ) Of the species involved in the reaction, which chemical species served to determine the “end” of the reaction? What is the physical manifestation of the “end of the reaction”? The “end” of the reaction was determined by the formation of the sulfur which caused the cloudiness in the solution. 4) Explain why the rate of reaction was approximated to be equal to 1/t. The equation for the rate of reaction is rate of reaction=? [product] x 1? T. Since the concentration in each run is kept constant, it is safe to approximate the rate of reaction to 1/t. ) Discuss the significance of using the same timer when measuring the time it takes for the reaction to be completed. Determine the effect of not using the same timer to the rate of reaction. It is important to use the same timer when measuring the time it takes for the reaction to be completed so that the data collected will more or less be constant. Not using the same timer could give different times affecting the overall data. 6) Discuss the significance of having only one person to judge the completion of the reaction. Determine the effect of having multiple individuals judging the “end” of the reaction.

Having multiple people judge the “end” of the reaction would be very confusing since different people might have different interpretations of the “end” of the reaction. Having only one person judge would therefore be more efficient. 7) Discuss the significance of adding the acid last in runs 1-3 and adding thiosulfate last in runs 4-6. It is important to add acid and thiosulfate last since a large amount of heat is released when these are mixed with water. If water is added after acid and thiosulfate, it would only result to the solution having a higher concentration of acid or thiosulfate. ) Determine the rate order with respect to the reactants in the reaction. Show in simple calculations the complete solution. Thiosulfate: Rate 2=k[S2O32-]2m[H+]2n Rate 3=k[S2O32-]3m[H+]3n 0. 035=k[0. 05]m[0. 4]n 0. 099=k[0. 025]m[0. 4]n -1. 50=[2]m m=log2(3. 353)=-1. 50 H+: Rate 4=k[S2O32-]4m[H+]4n Rate 5=k[S2O32-]5m[H+]5n 0. 053=k[0. 1]m[0. 6]n 0. 042=k[0. 1]m[0. 4]n 1. 26=[1. 5]n n=log1. 5(1. 26)=0. 57 m+n=-1. 5+0. 5=1 Therefore, the rate of reaction is rate= k[S2O32-] 9) What does the order calculated imply about the molecularity of the reaction? This implies that the molecularity of the reaction depends on the thiosulfate. 0) Set up the rate law. rate= k[S2O32-] 11) Based on the rate law, suggest a mechanism for the reaction between thiosulfate and hydronium ion. S2O32-(aq) + 2H+(aq) > H2S2O3(aq) H2S2O3(aq) > S(s) + SO2(aq) + H2(g) S2O32-(aq) + 2H+(aq) > S(s) + SO2(aq) + H2O Temperature-dependence of the Reaction Rate: The Arrhenius Equation 12) Based on the data obtained, how is the temperature related to the rate of reaction? Explain the observed trend in the experiment using theories in kinetics. It was observed that as the temperature increases, the faster the rate of reaction.

This is because k=Ae-Ea/RT, which shows that temperature is directly proportional to temperature. 13) Explain why 1/t is taken to be the rate constant. Since the initial concentrations are taken to be constant, 1/t can be taken to be the rate constant. 14) Construct a –ln t vs 1/t graph. Report the equation of the best fit line and the linearity coefficient. What does the linearity coefficient suggest about the relationship of –ln t with 1/t? Figure 1. Graph of ln t vs 1/t Theoretically, the rate of reaction should be directly proportional to the concentration of the reactants. 5) Calculate for the Ea of the system. Discuss the significance of the sign of the Ea and explain why is it so. 16) Determine the ? H of the reaction and construct an energy profile for the reaction. Catalysis: 17) Give the balanced equation for the reaction of permanganate with oxalate. S2O82-(aq) + 2I-(aq) > I2(aq) + 2SO4 18) What is the role of Cu2+? Cu2+ was used as a catalyst and as an indicator of the “end” of the reaction. 19) What is the role of starch in the reaction? Starch was used as a determinant for the amount of iodine in the solution. 20) Draw an energy diagram for this exothermic reaction.

Overlay the plot for the uncatalyzed reaction with the catalyzed one (use dashed lines). 21) Give the balanced equation for the reaction of permanganate with oxalate. 5C2O42-(aq) + 2MnO4-(aq) + H+ > 10CO2(g) + 8H2O(l) + 2Mn+(aq) 22) Why was the H2SO4 used to acidify the solution? What other acids can be used in replacement of H2SO4? H2SO4 was used since a strong acid was needed to convert MnO4 into Mn. If anything other than a strong acid is used, it is possible that another form of Mn would be produced. Other acids that can be used are HCl, HNO3, HClO4, H2SO4, HBr, and HI. 3) What determined the “end” of the reaction? Aside from the one used in the experiment, what phenomenon would you use to signal the end of the reaction? The decolorization of the solution signaled the “end” of the reaction. Another phenomenon that can determine the end of a reaction is through the formation of a precipitate. 24) Among the species involved in the reaction, what is the catalyst? Mn is the catalyst. 25) Explain autocatalysis. Autocatalysis occurs when the reaction produces a compound that can be used as catalysis for a second reaction.

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