Analysis Of Heats Of Reaction And Calorimetry Biology Essay

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The objective of the experiments provided is to exploit investigations that incorporate analysis of heats of solution and dilution and these studies incorporate oxidoreductase reactions, organic reactions, reactions that have precipitation of a solid or release gas and to finish, one reaction in which a complex ion is formed. Calorimetric experimentations described are deliberate for more concrete perception on thermodynamic changes including free energy, entropy and entropy. The calorimeter constant to account for both the physical and chemical heat changes was determined to be 3.48J/K, CuSO4 with NH3 had ∆H of -13.15KJ/mol with 34.25% error , FeSO4(aq) with NaOH(aq) had a ∆H of -51.09 KJ/mol with 88.88% error. Whereas, MgSO4(aq) with NaOH(aq) had a ∆H of -5.07 KJ/mol with 99.37% error. ∆H for anhydrous sodium acetate was 20.30KJ/mol with 76.36% error and enthalpy for trihydrated sodium acetate was -70.77KJ/mol with 95.15% error. For the Heats of Dilution enthalpies varied for the different volumes of ethanol: 5mLs (∆H of -9.20 KJ/mol), 10mLs (∆H of -6.82KJ/mol), 25mLs (∆H of -8.56KJ/mol), 50mLs (∆H of -6.05KJ/mol), 100mLs (∆H of -2.20KJ/mol) and finally for 150mLs (∆H of -0.77KJ/mol). Reaction of acetic anhydride with NaOH(aq) had ∆H of -41.14 KJ/mol with percent error of 186.7%. The results obtained from the described experiments had large errors; nonetheless, the experiments served the purpose of successfully displaying exothermic and endothermic reactions, relationship between internal energy, equivalence of heat and enthalpy under constant units of pressure and also provided relations of the first law of thermodynamics.

Introduction:

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Calorimetry has an incredibly extensive amount of applicability, varying from drug design in therapeutics, to controlling chemical reactions inclusive of waste and disposal by the constructive and nuclear industries, and also the study of metabolic rates in biological systems which mainly include metabolic analyses.[1] When a reaction takes place, the associated energy change involves a transformation of heat; the study of heat change in chemical reactions is formally known as Thermochemistry. A calorimeter (a simple illustration is shown in Image 2) can be used to determine the heat discharged or captivated by a reaction. The chemical reaction itself is referred to be the 'system' where as the 'surroundings' include the thermistor, the air around the calorimeter and etc.[5] When two compounds of different temperatures are put together, heat will be relocated from high temperature to the mixture with lower temperature till both are at the same temperature. The compound at higher temperature liberates heat whereas compound at lower temperature captivates it, so that: Equation 1. [12] The negative heat change signifies that the heat is flowing out of the reaction which in turn is equivalent to the heat streaming into the system. Although most of the heat is absorbed by the contents in the surroundings during an exothermic reaction but some of the heat is absorbed by the calorimeter itself and therefore, calorimeter constant generally accounts for both the physical and chemical components of heat deviations.[4]

The notion of calorimetry embodies the law of conservation of heat energy. For the sake of simplicity and not considering any supplementary methods of energy change, the first law of thermodynamics upholds that the adjustment in internal energy (∆U) of a system is equivalent to the heat (q) supplied by the system minus the work(p∆V) performed on the system Equation 2.[16]

The heat capacity is dependent on how much substance is available, considered to be an extensive property and it is the total amount of heat required for increasing the temperature of the whole calorimeter by 1K.[2] The specific heat(Csp) of a compound is an important variable when considering temperature change because the amount of heat applied will be dependent on the specific heat of that particular compound.[10] To calculate the heat evolved from a reaction, consider equation 3. Equation 3 [7]

Measurements to obtain heat change can be taken at constant volume or at constant pressure. At constant volume the p∆V=0, indicates that no work is performed by the system and thus the specific heat capacity is equivalent to the change in internal energy. On the other hand, when considering constant pressure the specific heat capacity is generally greater than its component at constant volume and this is due to the supplementary energy required to carry out work. At constant pressure, the heat changes are correspondent to the constant enthalpy (∆H).[9]

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Moreover, enthalpies of measured reactions and Hess's Law can be used to assess enthalpy difference for a reaction that may not be explicitly detected. For instance, it may not be possible to determine enthalpy of Equation 4 but if enthalpies of intermediates are known then Hess's Law provides a way for the enthalpy of equation 4 to be determined in the following way:[3]

Enthalpy, much like internal energy, is a state function.[15] As shown in Image 1, the difference in internal energy in route 1 going from A to B is the same as the internal energy in route 2 going from B to A. Hence, ∆U1 and ∆U2 are equivalent to each other regardless of the route used. For the purposes of enthalpy, it is notable that when heat is transferred to the surroundings the reaction is exothermic, ∆Hrxn<0. For the purposes of endothermic reactions, the heat is absorbed by the system from the surroundings, ∆Hrxn>0.

Image 1: Serves to display that ∆U (internal energy) is a state function.

Image 2: Displays a simple example of a calorimeter.

Procedure: from:-

Pattison, Dexter B.; Miller, John G.; Lucasse, Walter W.Simplified colorimetric studies of various types. Journal of Chemical Education, vol. 20, issue 7, p. 319-325

Note that the measurements of heat change are initially taken in mAmps using a thermistor and consequently it is necessary to convert current into temperature (K). Before starting each experiment allow the calorimeter to equilibrate to room temperature and make sure that the apparatus is dry before each use. For maximum accuracy, calibrate the calorimeter at temperature spans ranging from ice cold water to boiling point. For the purposes of obtaining an adequate calibration constant, first add cold water (0 °C) and measure the temperature. After the calorimeter has had enough time to equilibrate at room temperature, measure temperature of water that is in equilibrium with the surroundings (room-temperature); similarly, measure the temperature of the water at boiling point.

The phase transition illustrates transformations of the solid, liquid and gaseous states of a material. For the phase transition of water, add 100mLs of ice cold (0 °C) to the calorimeter and measure the current for approximately a minute. Add 100mLs of boiling water into the calibrator and start measuring the temperature.

Although each calorimetric study has its own standardized procedure but generally all of the studies presented in this experiment can follow a conventional pathway. Take in to consideration that both solutions for a reaction should be in equilibrium with the surroundings before the reaction takes place and thus some solutions will have to be prepared 24hours in advance to allow enough time for the equilibration process. The subsequent steps for reaction of CuSO4 with NH3 can also be carried out for reactions of: MgSO4(aq) with NaOH(aq), FeSO4(aq) with NaOH(aq), and acetic anhydride with NaOH(aq). For the Reaction of Copper Sulfate with Excess Ammonia, add 100mLs of Cu(SO4)aq to the calorimeter, measure the temperature for 5 minutes. At the 5th minute add100mLs of the ammonia to the calorimeter Cu(SO4) solution(CuSO4 should already be in the calorimeter), close the top with a cork and thoroughly mix by carefully swirling the calorimeter while allowing the thermistor apparatus to take measurements. After the mixing of two compounds, the temperature is read every twenty seconds till the 10th minute or for a prolonged period until the heat exchange has leveled off.

The procedures for reactions in Heats of Dilution and Heats of Solution should be slightly adjusted. The chemical reaction for Heats of Dilution will only have one liquid. Measure 100mLs of diluted water at room temperature; weigh out 6.8 grams of anhydrous sodium acetate (NaC2H3O2(s)) (for dissolution of sodium acetate trihydrate, weigh out 8.2 grams of (NaC 2H3O2x3H2O(s)). Be sure to close the lid of the container to avoid having the anhydrous sodium acetate absorb moisture from the atmosphere. This process is immensely time sensitive so quickly place the solid in the calorimeter and add 100mLs of diluted water. Stir gently and take the temperature. As for Heats of Dilution varying volumes of water and alcohol must be used (volumes are listed in Table 1 below).

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Table 1: The Heat of Dilution of Ethyl Alcohol by Water

5mLs of 95.6% alcohol, 200mLs of water

10mLs " " 200mLs " "

25mLs " " 175mLs " "

50mLs " " 150mLs " "

100mLs " " 100mLs " "

150mLs " " 50mLs " "

Discussion/Results:

As mentioned before, some of the heat is absorbed and in some cases released by the calorimeter itself; heat capacity of the calorimeter was determined to be 3.48J/mol. For reaction of CuSO4 with NH3 ∆H was -13.15KJ/mol, which indicates that the reaction was exothermic and thus the system released heat into the surroundings. The percent error for the reaction of CuSO4 with NH3 was 34.25%.

For the reactions of MgSO4(aq) with NaOH(aq)and FeSO4(aq) with NaOH(aq) a precipitate was produced. Although these two reactions stated above are quite analogous but they showed discrepancy for the heats of reaction. For instance, FeSO4(aq) with NaOH(aq) had a ∆H of -51.09 KJ/mol with 88.88% error; whereas, MgSO4(aq) with NaOH(aq) had a ∆H of -5.07 KJ/mol with 99.37%. The percent error for the preceding reactions was quite large. ∆H for heats of solution of water with anhydrous sodium acetate is 20.30KJ/mol with 76.36% and enthalpy for trihydrated sodium acetate is -70.77KJ/mol with 95.15%. As seen in Image 3 and Image 4 the enthalpy signs for heats of anhydrous and trihydrated sodium acetate are of opposite signs. For the Heats of Dilution enthalpies varied for the different volumes of ethanol: 5mLs (∆H of -9.20 KJ/mol), 10mLs (∆H of -6.82KJ/mol), 25mLs(∆H of -8.56KJ/mol), 50mLs(∆H of -6.05KJ/mol), 100mLs (∆H of -2.20KJ/mol) and finally for 150mLs (∆H of -0.77KJ/mol). The results shown in Image 5 show a large deviation from literature values; the image shows -∆H of ethanol and water mixture against # of moles of H2O in 95.9% ethanol. The end results of the alcohol and water mixture could have been better if pure alcohol was used and also poor results are due to using damp beakers for holding ethanol. Since water was already present in the alcohol the results were inconsistent.

The organic reaction of acetic anhydride with NaOH(aq) had ∆H of -41.14 KJ/mol; yet again with enormous percent error of 186.7%. The energy for the preceding reactions was shifted from the system to the surroundings and thus they were exothermic. However, as seen for the reaction of acetic acid with NaOH(aq)percent error is larger than 100% which indicates that the results were inaccurate for most of the experiments because the law of conservation of energy specifies that heat of the system and the surrounding must remain constant.[14] According to equation 1 the heat lost by the system must be equivalent to the heat gained by the surroundings but the data indicates otherwise.

Overall, even with large errors in calculations this experiment in fact served to successfully correlate heat changes and flow of heat to chemical reactions. The energy transformation that was detected is considered both chemical and physical and can be seen in the graphs included in Image 3 and Image 4. Image 3 serves to show that when enthalpy of the reactants is larger than that of the products and thus Image 3 is a representation of exothermic reactions because heat is flowing from the system to the surroundings. On the other hand, Image 4 functions to represent endothermic reactions where the heat flows from the surroundings to the system and enthalpy of products is larger than that of the reactants.

Experimental error could have occurred due to factors such as negligent handling of equipment and measuring volumes imprecisely. The main source of error was due to not closing the mouth of the calorimeter with the cork quickly enough to avoid losing heat and additionally, heat might have escaped because possibly the calorimeter was not well insulated which would have played a role in the determining the differences of the temperatures. The ∆T would have reformed due to the feeble insulation of the calorimeter and this would consequently lead to bigger calculation errors.[11] It is noteworthy that even in the instructions provided in Simplified Calorimetric Studies of Various Types: section with the reaction of acetic anhydride with sodium hydroxide listed that the literature values provided are rather long dated and even errors small as one percent "… in the measurement of heat of combustion," would explain the large difference obtained at the end because as mentioned before a small error could eventually lead to a large difference in calculation.[17]

Image 3 Image 4

Image 5

Conclusions:

The calorimeter was 3.48J/K, CuSO4 with NH3 had ∆H of -13.15KJ/mol with 34.25% error , FeSO4(aq) with NaOH(aq) had a ∆H of -51.09 KJ/mol with 88.88% error. Whereas, MgSO4(aq) with NaOH(aq) had a ∆H of -5.07 KJ/mol with 99.37% error. ∆H for anhydrous sodium acetate was 20.30KJ/mol with 76.36% error and enthalpy for trihydrated sodium acetate was -70.77KJ/mol with 95.15% error. For the Heats of Dilution enthalpies varied for the different volumes of ethanol: 5mLs (∆H of -9.20 KJ/mol), 10mLs (∆H of -6.82KJ/mol), 25mLs (∆H of -8.56KJ/mol), 50mLs (∆H of -6.05KJ/mol), 100mLs (∆H of -2.20KJ/mol) and finally for 150mLs (∆H of -0.77KJ/mol). Reaction of acetic anhydride with NaOH(aq) had ∆H of -41.14 KJ/mol with percent error of 186.7%. As mentioned before the results obtained from the described experiments were particularly erroneous; nonetheless, the experiments were consistent and served the purpose of effectively demonstrating exothermic and endothermic reactions, relationship between internal energy, equivalence of heat and enthalpy under constant units of pressure and also provided relations of the first law of thermodynamics.

These laboratory methods can be redirected to reflect upon and study systems that can provide a higher yield of work. It would be convenient to determine how to improve heat transformation that yields a small loss of energy and provides more energy to do work. To improve results obtained from this particular experiment, in the future it would be wise to consider the change in enthalpy caused to molecular interactions hence using heats to dilution to find the actual change in enthalpies for each reaction.[6]