Calorimetry and Hess’s Law Finding for the ΔHrxn Combustion of Magnesium PDF

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Chemistry Lab 1105

April 23, 2021 Spring 2021

Calorimetry and Hess’s Law Finding for the ΔHrxn Combustion of Magnesium I.

Introduction In this experiment, calorimetry, as well as Hess’s law, were studied. Therefore, the heat

formation and the transfer are studied in chemical processes by the determination of all the surroundings that is where all the heat changes happened, as well as, the changes of the temperature. Heat can be either absorbed or released. When heat is released, the process is known as exothermic, now if the heat is observed, the process is known as endothermic. These processes happened at the same one for the system and the other for the surroundings. For instance, when the system releases heat, the surroundings gain heat or vice-versa. There are also other terms to describe heat and heat transfer such as enthalpy. Enthalpy is the heat that is being observed or leased by the system. Enthalpy can be calculated, however, in order to calculate the heat that is generated, the mass, the change in temperature as well as the specific heat are needed. Once the amount of heat that is generated is collected, the enthalpy can then be calculated by dividing qp which is the heat that is being absorbed or released at a constant pressure by the number of moles of the material. Now, the specific heat is the heat that is required to raise the temperature of one gram of the material by one degree Celsius or kelvin. The way that is used is to calculate how much heat must be added to a substance in order to raise the temperate. Now some of the main objectives of this experiment were to be able to determine the heat capacity of a coffee-cup calorimeter, to be able to use the techniques of the calorimeter to be able to measure the reaction enthalpies, to be able to apply Hess’s law in the calculation to be able to determine the enthalpy of the reaction of Magnesium burning in Oxygen, and finally, to be able to combine all the concepts of calorimeter and Hess’s law to be able to measure ΔHrxn of the exothermic reaction. The calorimetry and Hesse’s law were used to determine the surrounding conditions and by using that information determine the heat that was released. These individual calculations 1

and heat released were convinced with Hess’s law and then manipulated to determine ΔHrxn for the Magnesium burning in Oxygen. The enthalpy of the reaction was determined by using a calorimeter by exchanging the heat with a calibrated calorimeter. The change in temperature of the measuring part of the calorimeter was converted into the amount of heat. Now the calorimeter constant was the heat capacity of the calorimeter in J/k. Hess’s law is defined as no matter of the multiple stages or steps of a reaction, the total enthalpy change for the reaction is the sum of all changes. It was useful in this experiment because two reactions were conducted with the ΔHrxn being calculated for both of them. Using the chemical equations, these were then manipulated to add and cancel out to get the desired results. Now, in this experiment, there were some chemicals and procedural hazards that were dealt with. HCl was carefully transferred from one place to another hence is a very coercive substance, also, the hot substances, and the hot plate. In order to avoid these hazards, proper PPE was wearing the whole time. II. Results AnalysisA. Data and Calculations Part I Heat Capacity of the Cup Calorimeter Exact volume of cold water (to the nearest 0.1mL) (D.1)

50.0 mL

Temperature of cold water (in cup) (D.2)

21.18℃

Exact volume of hot water (to the nearest 0.1mL) (D.3)

50.1mL

Temperature of hot water (in-cylinder) (D.4)

50.0℃ 35.97℃

Tf from the graph by extrapolation (D.5) Table I Calculations I 50.0 mL =

1𝑔 = 50.0g 1𝑚𝐿

50.1 mL =

1𝑔 = 50.1g 1𝑚𝐿

ΔTHW = 35.97℃-50.0℃ = -14.03℃ 2

ΔTHW = 35.97℃-21.18℃ = 14.79℃ qHW = (50.1gHW)(4.184

𝐽 )(-14.03℃ ) = -2940.95 J 𝑔•℃

qcw = (50.0gHW)(4.184

𝐽 )(14.79℃ ) = 3094.07 J 𝑔•℃

qCal = |-2940.946J| - |3094.068 J| = -153.12 J CCAL =

−153.12𝐽 𝐽 = 10.35 14.79℃ ℃ Graph I

Part II Determination of the Heat of Reaction of Mg in HCl Description of Sample

Pieces of metal

Exact volume of HCl (to the nearest 0.1mL) (D.6)

50.1mL

Initial temperature of HCl (in cup) (D.7)

20.90℃

Exact mass of Mg (to the nearest 0.001g) (D.8)

0.180g

Tf from graph by extrapolation (D.9)

36.83℃ Table II

Write the thermochemical equation which took place in the cup: Mg (s) + 2HCl (aq) → MgCl2 (aq) + H2 (g)

ΔHrxn1 = -472.89

𝐾𝐽 𝑚𝑜𝑙

Calculations II: 3

50.1 mL =

1𝑔 = 50.1g 1𝑚𝐿

ΔTsoln = 36.83℃ - 20.90℃ = 15.93℃ qsoln = (50.1g)(4.184

qCAL =( 10.35

𝐽 )( 15.93℃) = 3339.22 J 𝑔•℃

𝐽 )(15.93℃) = 164.88 J ℃

qrxn1 = -(3339.22 J+ 164.876 J)) = -3504.10J n = 0.180g

ΔHrxn1 =

1𝑚𝑜𝑙𝑒 = 7.41 x 10-3Mg moles 24.304𝑔 −3504.096 𝐽 −3

7.41 𝑥 10 𝑚𝑜𝑙

ΔHrxn1 =-472887.99

= -472887.99

𝐽 𝑚𝑜𝑙

𝐽 1𝐾 𝐽 𝐾𝐽 = = -472.89 𝑚𝑜𝑙 1000𝐽 𝑚𝑜𝑙 Graph II

Part III Determination of Heat of Reaction of MgO in HCl Description of Sample

Powder

Exact volume of HCl (to the nearest 0.1mL) (D.10)

50.0mL

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Initial temperature of HCl (in cup) (D.11)

21.35℃

Exact mass of MgO (to the nearest 0.001g) (D.12)

0.250g

Tf from graph by extrapolation (D.9)

24.31℃

Table III Write the thermochemical equation which took place in the cup: MgO (s) + 2HCl (aq) → MgCl2 (aq) + H2O (l)

ΔHrxn2 =-104.82

𝐾𝐽 𝑚𝑜𝑙

Calculations III ΔTsoln = 24.31℃ - 21.35℃ = 2.96℃ qsoln = (50.0g)(4.184

qCAL = (10.35

𝐽 𝑔•℃

) (2.96℃) = 619.23 J

𝐽 ) ( 2.96℃) = 30.64J ℃

qrxn2 = -(619.23J + 30.64J) = -649.87 J n = 0.250g

ΔHrxn2 =

1𝑚𝑜𝑙𝑒 = 6.20 x 10-3Mg moles 40.304𝑔 −649.87 𝐽 −3

6.20 𝑥 10 𝑚𝑜𝑙

ΔHrxn2 = -104817.74

= -104817.74

𝐽 𝑚𝑜𝑙

𝐽 1𝐾 𝐽 𝐾𝐽 = = -104.82 𝑚𝑜𝑙 1000𝐽 𝑚𝑜𝑙 Graph III

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Part IV Determination of the Enthalpy of Combustion of Mg Using Hess’s Law Calculations IV Rxn 1) Mg (s) + 2 HCl(aq) →MgCl2 (aq) + H2 (g) Rxn 2) MgO (s) + 2 HCl(aq) →MgCl2 (aq) + H2O (l) Rxn 3) H2 (g) + ½ O2(g) → H2O (l)

𝐾𝐽 𝑚𝑜𝑙 𝐾𝐽 ΔHrxn2 = -104.82 𝑚𝑜𝑙 𝐾𝐽 ΔHF = -285.84 𝑚𝑜𝑙 ΔHrxn1 = -472.89

Manipulate Rxns 1, 2, and 3, and combine them to achieve the desired equation below. Mg (s) + 2 HCl(aq) →MgCl2 (aq) + H2 (g)

ΔHrxn1 = -472.89

MgCl2 (aq) + H2O (l) →MgO (s) + 2 HCl(aq)

ΔHrxn2 = 104.82

H2 (g) + ½ O2(g) → H2O (l) Mg (s) + ½ O2 (g) → MgO (s)

𝐾𝐽 𝑚𝑜𝑙

𝐾𝐽 𝑚𝑜𝑙 𝐾𝐽 ΔHF = -285.84 𝑚𝑜𝑙

ΔHrxn= -653.91

𝐾𝐽 𝑚𝑜𝑙

Theoretical ΔHrxn [ See Appendix II, Page 8 of Tro (2017 or latest edition)] = -601.6

% error

𝐾𝐽 𝑚𝑜𝑙

|601.6 − 653.9 | ×100 = 8.69% 601.6 6

B. Discussion Part IHeat Capacity of the Cup Calorimeter In this part of the experiment, the heat capacity of the calorimeter was calculated. A calorimeter is an object used to measure the heat of a chemical reaction as well as the heat capacity. In this experiment, the heat capacity of the calorimeter was necessary to be determined because the heat capacity was later used to determine different enthalpies in the reactions. It is important in calculating the qcal as well as the qrxn1 for the cup, this was leather used to find the ΔHrxn for each of the reactions. The heat capacity of the calorimeter was determined from the amount of water and the temperature change that underwent by water. By simply using the specific heat of water, the mass of the water, which was obtained by multiplying the volume of water by the density, and finally the change in temperature to obtain the specific heat of the cup. By simply analyzing the graph that was obtained using Vernier Software, when hot water was added to the cold water, the temperature of the cold water started to increase as it can be observed in “Graph I” until it got to a point where it stayed stable and after a few minutes started to decrease. The cold water absorbed 14.03-degree celsius of heat because it went from 21.18 celsius to 35.97 celsius. The heat capacity for this part of the experiment was CCAL =

−153.12𝐽 𝐽 = 10.35 , It was calculated by using the mass of the hot water, the mass of the 14.79℃ ℃ cold, the temperature of the water, the change in temperature for water, the density, as well as the specific heat of the water. The rest of the calculations can be found in “Calculations I”. Part II Determination of the Heat of Reaction of Mg in HCl In this part of the experiment, the heat reaction of Mg in HCl was tested. For this part of the experiment, 50 mL of 1M HCl was measured and it was poured into a dry calorimeter cup, the

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temperature of the liquid was measured using the Vernier probe. After that, 4 pieces of Mg ribbon were then weighted, and before adding the Mg to the HCl, the start button was clicked to get an accurate reading for the temperature of the HCl. After a minute the weighted pieces of Mg were put inside the calorimeter cup containing the 50 mL of acid. Then the solution was mixed continuously by swirling the beaker. The temperature then rose to a point where it was stable which took about 6 to 8 mins. The graph increased from a temperature of 20.90-degrees Celsius to 36.83 degrees after adding the Mg ribbon indicating that it was an exothermic reaction. The enthalpy of the reaction was ΔHrxn1 =

−3504.096 𝐽 −3

7.41 𝑥 10 𝑚𝑜𝑙

x

1𝐾𝐽 𝐾𝐽 = -472.89 , again, it is 1000𝐽 𝑚𝑜𝑙

negative because the reaction is an exothermic reaction. The enthalpy was calculated by dividing qrxn1 which was -3504.10J by the number of moles. However, the joules were later converted to KJ to have the units as KJ/mol for more detailed calculations on how that value was achieved can be found in “Calculations II”. The balanced thermochemical equation was, Mg (s) + 2HCl (aq) → MgCl2 (aq) + H2 (g)

ΔHrxn1 = -472.89

𝐾𝐽 𝑚𝑜𝑙

Part III Determination of the Heat of Reaction of MgO in HCl In this part of the experiment, the heat reaction of MgO in HCl was tested. The procedure was similar to the one in the previous part. 50 mL of 1M HCl was first measured and then it was poured in a dry calorimeter cup, the temperature of the acid was then measured by using the Vernier probe and the software. After that, 0.250 grams of pure MgO was weighed using the valance. The MgO was then added to the added acid, however, before adding the MgO to the acid, the start button was first clicked to get an initial reading for the temperature of the acid. After the Vernier probe was reading the temperature of the acid for about a minute or two the mass of MgO was then added to the calorimeter containing the acid. Then the mixture was then 8

mixed continuously by swirling the beaker. Between the tests, the calorimeter was cleaned thoroughly so that way the residues from the previous test would not affect the reading from the present results, that way the results were more accurate. It was also important to continuously stir the solution that way it would occur faster and the reading for the temperature was more accurate, and also to make sure that the solution was completely dissolved. By looking at graph III, it can be noticed that the temperature of the water was slowly increased when the mass of MgO was added to the acid. That behavior indicated that the reaction was an exothermic reaction same as the previous one. The enthalpy of the reaction was ΔHrxn1 =

−649.87 𝐽 −3

x

6.20 𝑥 10 𝑚𝑜𝑙𝑒𝑠

1𝐾𝐽 𝐾𝐽 = -104.82 , again, the reason why it is a negative value is that the reaction is an 1000𝐽 𝑚𝑜𝑙 exothermic reaction. The enthalpy was calculated by dividing qrxn2 which was -649.87 J by the number of moles or MgO. However, the joules were later converted to KJ to have the units as KJ/mol for more detailed calculations on how that value was achieved can be found in “Calculations III”. The balanced thermochemical equation was, MgO (s) + 2HCl (aq) → MgCl2 (aq) + H2O (l)

ΔHrxn2 =-104.82

𝐾𝐽 𝑚𝑜𝑙

PART IV Determination of the Enthalpy of Combustion of Mg Using Hess’s Law Using part II and part III of the experiment, the enthalpy for the combustion of one mole of magnesium in oxygen using Hess’s Law was ΔHrxn= -653.91

𝐾𝐽 , this was accomplished by 𝑚𝑜𝑙

rearranging some values in order to add or cancel out. As shown, Mg (s) + 2 HCl(aq) →MgCl2 (aq) + H2 (g) MgCl2 (aq) + H2O (l) →MgO (s) + 2 HCl(aq)

𝐾𝐽 𝑚𝑜𝑙 𝐾𝐽 ΔHrxn2 = 104.82 𝑚𝑜𝑙

ΔHrxn1 = -472.89

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H2 (g) + ½ O2(g) → H2O (l)

ΔHF = -285.84

𝐾𝐽 𝑚𝑜𝑙

ΔHrxn= -653.91

𝐾𝐽 𝑚𝑜𝑙

The Balance chemical equation is shown below, Mg (s) + ½ O2 (g) → MgO (s)

A better understanding of these calculations can be found in “calculations IV” The theoretical enthalpy for the reaction was -601.6

𝐾𝐽 The percent error that was calculated 𝑚𝑜𝑙

in this experiment was 8.69%. The accuracy of the results was not exactly as it was expected because an 8.69% error was still very high. However, the possible high percent error can be due to the measurements of the acids by using a graduated cylinder, it could also be due to systematic error when measuring the temperature of the water in the first part of the experiment. III. Conclusions The lab techniques that were used in collecting the experimental data were to properly use the calorimeter as well as how to properly operate the Vernier software and to be able to read the exact temperatures in the graph. The Hess’s Law was used to manipulate the data that was collected for the reactions by letting combine these chemical equations that way achieve the experimental goal. The reactions that were studied in this experiment were, Rxn 1) Mg (s) + 2 HCl(aq) →MgCl2 (aq) + H2 (g) Rxn 2) MgO (s) + 2 HCl(aq) →MgCl2 (aq) + H2O (l) Rxn 3) H2 (g) + ½ O2(g) → H2O (l)

𝐾𝐽 𝑚𝑜𝑙 𝐾𝐽 ΔHrxn2 = -104.82 𝑚𝑜𝑙 𝐾𝐽 ΔHF = -285.84 𝑚𝑜𝑙

ΔHrxn1 = -472.89

The first two were collected from the experimental, and graph II and graph III, it can be noticed that the temperature increased for both of them meaning that the reaction was an exothermic reaction, also, since the reactions were exothermic, the values that were calculated were also negative as it can be noticed in the values above. The enthalpies for the reactions indicated that

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all the reactions were exothermic reactions as well as the final. Based on my calculated percent error, it can be inferred that by using calorimetry and Hess’s Law carefully the enthalpy of a highly exothermic reaction can be determined, however, the experiment, must be carefully made due to the fact that small sources of error can cause a large percent error at the end of the experiment. The accuracy and precision were achieved by measuring almost exact amounts of acid and water as well as writing the exact values of the mass of solutions that were required. The way that it can be improved is by having a better thermometer that measures values more accurately. The application of calorimeter and Hess’s Law can be applied in many practical applications, for instance, some industries that research different methods in some production use this law to determine if their method is the most effective one to produce the products on the energy that is needed or the energy that is released. They measured how much energy is released or needed so that they can make the choice.

POST LAB QUESTIONS

1. What is calorimetry and Hess’s Law? A calorimeter is an object used to measure the heat of a chemical reaction as well as the heat capacity. Now, Hess’s Law is defined as no matter of the multiple stages or steps of a reaction, the total enthalpy change for the reaction is the sum of all changes. 2. What makes it difficult to measure the enthalpy of reaction of Mg burning in oxygen, and what reactions will you perform to determine it? Getting the exact values when measuring made it a bit challenging because, in order to

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get the desired results, the measurements had to be almost perfect. The reactions that were performed were Mg reacting in HCl and MgO reacting in HCl. 3. What is the heat capacity of your coffee-cup calorimeter, and why do you need to determine it? The heat capacity of the coffee cup calorimeter was 10.35J/℃ this was later used to find the qcal and for the rest of the calculation in part II and part III. 4.

What is the enthalpy of reaction of Mg reacting in HCl? The enthalpy for the reaction of Mg reacting in HCl was ΔHrxn1 = -472.89

𝐾𝐽 𝑚𝑜𝑙

5. What is the enthalpy of reaction of MgO reacting in HCl? The enthalpy for the reaction of MgO reacting in HCl was ΔHrxn2 = -104.82

𝐾𝐽 𝑚𝑜𝑙

6. How will you manipulate the three equations and their enthalpies to achieve the desired overall equation and calculate its ΔHrxn? The way that these three equations were manipulated was by rearranging the values to later be able to cancel out similar values and then, add the negative values and positive values to get the desired value of ΔHrxn which was ΔHrxn= -653.91

𝐾𝐽 𝑚𝑜𝑙

7. What is your overall enthalpy of reaction for burning one mole of Mg in O2? The overall enthalpy of the reaction for burning one mole of Mg in O2 was ΔHrxn= -653.91

𝐾𝐽 𝑚𝑜𝑙

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References New Jersey City University. (2013). CHEM 1105: General Chemistry I laboratory and recitation: miniscale experiments. Boston, MA: Pearson Learning Solutions. Tro, N. (2017). Chemistry: A molecular approach (4th ed.). Boston, MA: Pearson. Tro, N. (2017). Chapter 4, Sections 6.4, 6.6, 6.7, 6.8, and Appendix II. Chemistry: A molecular approach (4th ed.). Boston, MA: Pearson.

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