Exp 4 Lab Report Joe CHM 113 (1)

pdf

School

University of Miami *

*We aren’t endorsed by this school

Course

113

Subject

Chemistry

Date

Apr 3, 2024

Type

pdf

Pages

Uploaded by PresidentElementStingray8

Aaron Joe CHM 113 Lab Report 2/27/24 Chemical Thermodynamics: Heat of Formation of MgO(s) Part A: Introduction: This experiment aims to determine the heat, or enthalpy, involved in the production of MgO(s). In this experiment, the enthalpy values were determined using a basic calorimeter. To calculate the amount of heat released, utilize calorimetry. Hess's law was utilized to determine the enthalpy of formation using three distinct chemical reactions in order to achieve the enthalpy values. According to Hess's law, the enthalpy of a reaction that is carried out in steps will equal the total of the enthalpy changes for each step. The three distinct equations must be applied since the air contains N2. Mg and N2 would react to generate Mg3N2. This measure could only be implemented in a pure oxygen environment, which is not possible in a laboratory setting. Procedure: 1. Set up the calorimeter. 2. Prepare 250mL of 1M HCl solution. 3. Using a graduated cylinder, transfer exactly 100mL of the dilute HCl solution into the calorimeter. 4. Clean a piece of Mg ribbon by rubbing steel wool to remove the layer of oxide on the surface. Weigh the Mg ribbon to approximately 0.25 grams and record the measurement. 5. Wind the magnesium ribbon around the loop of the stirrer. And record the initial temperature of the HCl solution.

6. Immerse the stirrer in the HCl solution and stir gently to ensure the Mg never comes above the surface of the solution. 7. Monitor the temperature rise and record the highest value reached. 8. Repeat steps 1 to 3. In step 4, use MgO instead and measure the weight to around 0.4 grams. 9. Record the initial temperature of HCl solution. Add the MgO and stir until completely dissolved. 10. Monitor and record the highest temperature reached. Equations: 2 reactions being measured: Mg(s) + 2HCL (aq) → MgCl2(aq) + H2(g) MgO(s) + 2HCL (aq) → MgCl2(aq) + H2)(g) Heat of formation of MgO: Mg(s) + 1/2O2(g) → MgO(s) q=mc(delta T) qsol=-qrxn Observations: Table of Physical Data: Mg Ribbon MgO Ribbon Weight (g) 0.334 0.403 Atomic Weight (g/mol) 24.035g/mol 41.035g/mol Number of Moles (mol) 0.013896 0.010048 Table of Temperature Changes and Heat Released:

For the MgO Ribbon reaction, we may have not given it enough time to heat up, and therefore obtained a very small temperature change. Mg Ribbon MgO Ribbon Initial Temperature (C) 23 C 23 C Final Temperature (C) 34 C 26 C Temperature Change (C) 11 C 3 C Heat given off (kcal) 1.1 kcal .3 kcal Enthalpy of Reaction (kcal/mol) -79.1595 kcal/mol -29.857 kcal/mol Since both ribbons display a rise in temperature, we can conclude that both reactions are exothermic. Calculations: The amount of heat released, q(heat)=mc(delta T), divided by the number of moles, yields the enthalpy of reaction. For Mg Ribbon: q(H2O)=(1.0cal/gC) x (100g) x 11.0 C= 1100 cal qsol=-qrxn 1100cal=-1100cal -1.1kcal Enthalpy= -1.1kcal/0.013896 moles = -79.1595 For MgO, we must utilize Hess’ Law in order to calculate the heat of formation : -79.1595kcal/mol - ( -29.857kcal/mol) - (-68.3kcal/mol) = -117.6025 kcal/mole MgO Enthalpy of Formation in (kJ/mol): -119.6025 kcal/mole x 4.184 kJ= -492.047 kJ/mol

Your preview ends here

Eager to read complete document? Join bartleby learn and gain access to the full version

Access to all documents
Unlimited textbook solutions
24/7 expert homework help

Part A Discussion: The Hf of MgO was successfully calculated using the data from this experiment. -492.047 kJ/mol was the experiment's outcome. Since MgO is created by burning Mg (s), which is an exothermic reaction, the figure appears to be accurate. Since it is losing heat into its surroundings, this suggests that the enthalpy value is negative. The calorimeter is the primary cause of error in this experiment. It is a basic calorimeter that isn't entirely sealed off. This is a serious problem since air contains mostly nitrogen, which reacts with magnesium to make Mg3N2, which creates new compounds. Furthermore, the system might have lost some heat in the thermometer hole, which would explain the lower enthalpy value. By adding more substance, such as PlayDoh, to adequately cover the hole or precisely measuring the hole to fit the thermometer, this inaccuracy might be avoided. Moreover, it's possible that the solid was partially dissolved because you couldn't see it in the Styrofoam, which could have resulted in a guessed amount of time. This would have caused inaccuracies in our final temperature, and as a result inaccuracies in the calculation of the heat given off. (Most likely our value of q would be higher). This error could be prevented by waiting longer after the reaction is completed, using a different cup or lid to visually see inside, or using mixing tools to ensure it is fully stirred and dissolved. A human mistake that occurred was that we did not stir for the first few minutes of the experiment. This may have caused the magnesium to not dissolve as well, and perhaps create inaccuracies in the temperature changes. The MgO heat of production was measured with a basic Styrofoam calorimeter. Mg(s) and MgO(s) were dissolved, and the data collected were used to calculate each of the enthalpies of formation. Next, the three chemical equations were subjected to Hess's law in order to

calculate the overall enthalpy of production of MgO, which came out to be -492.047kJ/mol. The lab's purpose was effectively accomplished by figuring out the enthalpy value. Part B: Introduction: The purpose of this experiment is to find the specific heat capacity of the unknown metal. We will do this by utilizing calorimetry as well as the ability to calculate both q values. We will then utilize these in the q=mc(delta T) to isolate for the specific heat capacity of the unknown. Metals are distinguished by their intense qualities, such as their specific heat, which can be used to identify a metal by its distinct specific heat capacity. Therefore by calculating the specific heat, we are able to determine the identity of the unknown metal. Equations: qlost= -qgained qlost= sp.h. X no. grams x (delta T) for the metal Heat gained= no. gra,s x 1cal/gC x (delta T) for calorimeter and contents Specific heat can be calculated with above equations Procedure: 1. Determine the weight of the unknown rod and record it in the data table. 2. Using a copper wire, suspend the unknown metal into a beaker of boiling water. 3. Measure exactly 50.0mL of water with a graduated cylinder and transfer it into the dry, double stacked Styrofoam cup. Record the initial temperature of the water. 4. When the metal rod, has reached the temperature of boiling for approximately 10 minutes, transfer it into the Styrofoam cup quickly. 5. Swirl the cup gently and record the highest temperature reached.

6. Repeat the experiment and determine the average specific heat Observations: Water Unknown Metal Weight (g) 50g 12.0570g Initial Temperature (C) 20 C 100 C Equilibrium Temperature (C) 25 C 25 Change in Temperature (Delta T) (C) 2 C -75 C Specific Heat Capacity (cal/gC) 1.0 cal/gC 0.1106 cal/gC Calculations: To calculate heat of water, we use this equation since it gained heat: Heat gained= no. gra,s x 1cal/gC x (delta T) for calorimeter and contents: (50cal/gC)(2 C)= 100 cal The heat lost by the metal is calculated by this equation: Heat lost= sp.h. X no. grams x (delta T) for the metal (Specific Heat of Unknown) x (12.0570g) x (-75 C) To find specific heat of unknown metal, we can use the following: qlost= -qgained, plug in knowns and then solve for the specific heat (12.0570g) x (-75 C) x (Unknown Specific Heat) = -100 cal Specific Heat of Unknown = 0.1106 Conclusion: The specific heat capacity of the unknown metal was successfully determined by the experiment, and the calculations were accurate because the specific heat capacity result returned

Your preview ends here

Eager to read complete document? Join bartleby learn and gain access to the full version

Access to all documents
Unlimited textbook solutions
24/7 expert homework help

as a positive number. Furthermore, the metal's specific heat capacity was lower than that of water, which makes sense given that the metal's heat fluctuated more dramatically than water's. The experiment yielded a result of 0.1106. Heat loss during the metal's transit from the boiling water to the Styrofoam cup was identified as the experiment's source of mistake. If the amount of external water is decreased, the inaccuracy is disclosed with the metal and the specific heat capacity would be lower than anticipated. Furthermore, the Styrofoam cup may have lost some heat to its surroundings as it is not a completely closed system with a hole on the top for the thermometer. This can be fixed by using PlayDoh or another substance to better cover the hole or creating a smaller hole for the thermometer to be inserted in. The error that this possesses on the experiment is a lower calculated specific heat capacity of the actual metal. Unfortunately, due to our errors, the specific heat capacity calculated for the unknown metal was incorrect. The specific heat options given were Sn = .056 cal/gC and Al = .215 cal/gC. Since our specific heat capacity does not match with either, it is probable that a substantial error was made. In order to fix this error, we must address the errors mentioned previously, as well as address some human errors. One significant one was that the unknown metal may have not been hot enough. Because we took it and transferred it early, this caused a less significant change in temperature. To fix this, we simply have to wait longer before transferring to the calorimeter. Another error was that the calorimeter cap had slipped off when inserting the thermometer. This would cause significant loss of heat transfer, and interfere with the data as well. Using a Styrofoam calorimeter, the experiment ascertained the specific heat capacity of an unidentified metal. The unknown metal was heated before being added to the water at room

temperature. The quantity of heat absorbed by the water and the amount of heat lost by the metal were calculated using the data gathered. The metal's specific heat capacity (0.1106) was then calculated using these values.

Related Documents

Lab Report 1.pdf

Lab Report 1.docx

Lab Report 5.pdf

Experiment 2 Lab Report.pdf

Experiment 3 Lab Report.pdf

Expt 3ab Report Joe.pdf

Spinach Lab.pdf

Experiment 6 Lab Report.pdf

Practice Problems_Chapter4_Alkenes.pdf

Experiment_4_Carbohydrates_Group_8_Membe.pdf

Heat of RXN - In class assignment WEEK 2.pdf

IMG_5055.jpeg

Recommended textbooks for you

Chemistry & Chemical Reactivity

Chemistry

ISBN:9781133949640

Author:John C. Kotz, Paul M. Treichel, John Townsend, David Treichel

Publisher:Cengage Learning

Chemistry & Chemical Reactivity

Chemistry

ISBN:9781337399074

Author:John C. Kotz, Paul M. Treichel, John Townsend, David Treichel

Publisher:Cengage Learning

Chemistry

ISBN:9781305957404

Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Publisher:Cengage Learning

Chemistry: An Atoms First Approach

Chemistry

ISBN:9781305079243

Author:Steven S. Zumdahl, Susan A. Zumdahl

Publisher:Cengage Learning

Chemistry

ISBN:9781133611097

Author:Steven S. Zumdahl

Publisher:Cengage Learning

Introductory Chemistry: A Foundation

Chemistry

ISBN:9781337399425

Author:Steven S. Zumdahl, Donald J. DeCoste

Publisher:Cengage Learning

SEE MORE TEXTBOOKS

Recommended textbooks for you

Chemistry & Chemical Reactivity
Chemistry
ISBN:9781133949640
Author:John C. Kotz, Paul M. Treichel, John Townsend, David Treichel
Publisher:Cengage Learning
Chemistry & Chemical Reactivity
Chemistry
ISBN:9781337399074
Author:John C. Kotz, Paul M. Treichel, John Townsend, David Treichel
Publisher:Cengage Learning
Chemistry
Chemistry
ISBN:9781305957404
Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste
Publisher:Cengage Learning
Chemistry: An Atoms First Approach
Chemistry
ISBN:9781305079243
Author:Steven S. Zumdahl, Susan A. Zumdahl
Publisher:Cengage Learning
Chemistry
Chemistry
ISBN:9781133611097
Author:Steven S. Zumdahl
Publisher:Cengage Learning
Introductory Chemistry: A Foundation
Chemistry
ISBN:9781337399425
Author:Steven S. Zumdahl, Donald J. DeCoste
Publisher:Cengage Learning