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Lab 7 Enthalpy and Hess’s Law Sydney Boucher CHEM 113L Professor Dunn March 28th, 2021

Sydney Boucher

2 Abstract

In this lab, part A experimentally determined the enthalpy of formation for MgO. To do this, Hess’ law is used for the following formation reaction: Mg (s) + ½ O2 (g) → MgO (s). The results of part a show that the enthalpy of formation for MgO is -317.6 kJ/mol. Part C found the heat of two different solutions containing salt, as the salt dissolves into the solution, one increases the temperature, while the other decreases the temperature. The two salts were CaCl2 and KNO3. Calcium chloride yielded a positive change in temperature and potassium nitrate yielded a negative change in temperature. With many of the experiments yielding high percentages of error, all results are valid when compared to the theoretical values. Introduction Lab 7, Enthalpy and Hess’s Law focuses on determining the enthalpy of a reaction by using Hess’ Law. This law was used to determine the enthalpy of formation of MgO, the enthalpy of a neutralization reaction, the heat of a solution and reaction between a salt and a base. While each value was found experimentally and calculated correctly, high percentage errors were yielded. This lab utilizes Hess’ Law to find the correct values for the change in T, q, m, and n to use in the equation, q = m* Cs * change in T, to find the enthalpy of each reaction. Data and Calculations Each part of this lab focused on finding the enthalpy of three different types of reactions. The enthalpy of formation, neutralization reaction, and the reaction between a salt and a base. The results of Part A, the enthalpy of formation of MgO, showed a positive enthalpy and the salt-base reaction of part C yielded both positive and negative enthalpies.

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Part A: Hess’s Law and Determination of ΔHf。MgO (kJ/mol) **only one trial was taken for each solute per instructors request. Parameters for the ΔHfo of MgO Mg

MgO

A

ΔT (。C)

7.1

3.5

B

Q soln (J)

246.6

121.5

C

n (mol)

.033

.005

D

ΔHrxn (kJ/mol)

-7.47

-24.3

E

Ave Δ Hrxn (kJ/mol) --------------------------

F

ΔHf。MgO (kJ/mol)

-317.6

G

Error %

47.2%

--------------------------

The equation used to find q is q=m·Cs·∆T where q is heat, m is mass of substance, Cs is the specific heat of the substance, and ∆T is the change in temperature. Part B: Enthalpy Change of Neutralization; omitted this semester

Part C: Enthalpy of Solution Parameters for the ΔHsol of CaCl2 and KNO3, salts CaCl2

KNO3

A

ΔT (。C)

4.5

-5.1

B

Q soln (J)

216.5

-245.4

C

n (mol)

.0095

.103

D

ΔHrxn (kJ/mol)

-22.7

-23.8

E

Error %

72.5%

166%

The equation used to find q is q=m·Cs·∆T where q is heat, m is mass of substance, Cs is the specific heat of the substance, and ∆T is the change in temperature.

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Results and Discussion The results, for the most part, were not valid or reliable. This is due to the large percent error that almost every experiment yielded. Part A obtained a 32% error and part C was found to have 72.5% and a 166% error for the two salts. The lowest of all of these is with the CaCl2 in part C with 72.5% error, which is still considered relatively high. I was surprised to see the trend of such high percentages of error, especially one as high as part C, at 166%. Conclusion The steps I would take to improve the data would first be to redo this lab. I believe that it would be of great importance to complete the lab with more intention and care, in an attempt to pinpoint the exact cause of such high error. Also, it would be beneficial to receive instructor or peer feedback in future labs to yield better, more accurate results. Other experiments that should be included in this class are labs that have bigger reactions. Something that could be done is a reaction that would cause a lot of foam, or light, or fire to be produced. Similar to the lab that was done in class with the salt sprays and bunsen burner, specifically the color of the flame they produced was the most enjoyable to conduct. The evidence showing the high percent error could derive from countless reasons. It leads one to think that there were consistent errors made in the experimenting process, due to the lack of errors in the calculations process. I believe that a main source of error came from the calorimeter and electronic instruments. Classified as a type of systematic error, this comes from misuse of the instruments, which in this case is the calorimeter. This means that I would have done or not done something that is key to collecting accurate data. The main way that I would go about improving this lab would be to give the theoretical values to the experimenters. This would be so that they can adjust or redo their experiment or

Sydney Boucher

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calculations throughout the lab to obtain more accurate results. It would also show the students where their errors may be coming from throughout the experiments, and teach them to have more caution when approaching specific areas of the experiments. Overall, the remainder of the lab is straight forward, easy to follow, and relatively simple....