experiment 11 chem113 PDF

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experiment 11 chem113...

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Finding Enthalpy of Reactions Using Calorimetry Intro: The purpose of this experiment is to measure the enthalpy for three acid-base reactions using calorimetry and Hess’s law after measuring temperature change. The change in enthalpy will give the heat of reaction for three acid-base reactions. To find the change in enthalpy using Hess’s law, the total enthalpy is calculated from rearranging multiple reactions in order to sum all of the enthalpies reach one total enthalpy. To find change in enthalpy, temperature change must be measured using a calorimeter. The calorimeter will measure the change in heat after “two or more substances are combined” (French et al. 60). To minimize heat lost to the surroundings, it is best to use a Styrofoam cup that is placed inside of a beaker to add insulation. Before finding the enthalpy of the acid-base reactions, hot and cold water need to be used in the calorimeter to find the heat capacity and to “determine the amount of heat lost to calorimeter”; the specific heat of water will be the same for the other acid-base reactions because they are “aqueous solutions” and because the same calorimeter is used; this is called the calibration step (French et al. 61). The heat equation that can be used if it is assumed that “no heat is lost to the surroundings” is shown below: q cal=−q hw −q cw (French et al. 61). The heat (q) found from the equation above can used to plug into either of the equations below, depending on if mass is known or not, to find the specific heat (Cs) of the aqueous solutions. The heat (q) found is also equal to

∆ H, which is the enthalpy of the system.

q=C ∆ T ∆ H=q=m C s ∆ T

(French et al. 61). When two substances are reacting, heat is lost or gained from the reaction and “will be absorbed by either the solution or the calorimeter” (French et al. 62). If temperature change is measured, then the amount of heat absorbed by both the solution and the calorimeter can be evaluated by the equation below: q rxn=−q soln−qcal (French et al. 62). If the equation above is expanded with the change in temperature of the solution, the change in temperature of the calorimeter, the specific heat of water, and the mass of the solution it will resemble the equation below: q rxn=−m soln s ∆ T −C cal ∆T (French et al. 62). After the heat is calculated using data from the experiment, the enthalpy can be found by dividing the change in heat of the reaction by the “number of moles in the limiting reagent” (French et al. 62). The number of moles can be found from the balanced net ionic equations of the limiting reagent. If

∆H

is a negative value then the reaction is exothermic, which means

that heat is released to the surroundings. If

∆H

is a positive value then the reaction is

endothermic, which means that the system absorbs heat from surroundings. This calculation can be done using the equation below: ∆ H=

q n (French et al. 62).

The first chemical reaction will be made from sodium hydroxide and hydrochloric acid, which yields “water and aqueous sodium chloride” (French et al. 63). The balanced chemical equation is shown below: 1 HCl ( aq ) +1 NaOH ( aq ) → 1 NaCl ( aq ) +1 H 2 O (l) . The second chemical reaction is produced from sodium hydroxide and ammonium chloride, “yielding aqueous ammonia (NH3 (aq)), water, and aqueous sodium chloride” (French et al. 63). The balanced chemical equation is shown below: 1 NaOH ( aq ) +1 N H 4 Cl ( aq) →1 H 2 O ( l ) + 1 NaCl ( aq )+1 N H 3 (aq) . The final chemical reaction is made from mixing hydrochloric acid and aqueous ammonia, “yielding aqueous ammonium chloride” (French et al. 63). Aqueous ammonia (NH3 (aq)) is used because it is easier to use in a lab setting since it is a liquid and because it changes temperature faster than most solids. The balanced chemical equation is shown below: 1 HCl ( aq ) +1 N H 3 ( aq) → N H 4 Cl (aq) .

Materials: 

Styrofoam cup



2.0 M hydrochloric acid



2.0 M sodium hydroxide solution



2.0 M ammonium hydroxide solution



100 ml beaker



150 ml beaker



250 ml beaker



50 mL graduated cylinder



Hot plate



Thermometer



Ring Stand



Iron clamp



Cardboard square



MeasureNet system



Temperature probe

Methods: 1. Gather Styrofoam cup and 100mL beaker and record the mass of both. 2. Set up ring stand with an iron clamp for the temperature probe. Connect temperature probe to MeasureNet and set up the “Temp vs. Time” graph with the limits x=0-60 and y=0-100, then press “Enter”. 3. Place Styrofoam cup inside 250mL beaker with cardboard square over the top of the cup. 4. Place 25mL cold water in Styrofoam cup and record mass. Place 25mL of hot water in 100mL beaker and record mass. 5. Record initial temperature of the hot and cold water. 6. Place Styrofoam cup of cold water under the ring stand and set up temperature probe in the water. 7. Start the time vs. temperature graph of the cold water for 20 seconds then gently but quickly add the hot water to the Styrofoam cup. Place cardboard square back over the top. 8. Stop the graph after it reaches a plateau, around 60seconds. Save the file as “010”. 9. Record final temperature of the solution.

10. Repeat steps 1-8 for 25mL 2.0 M hydrochloric acid and 2.0 M sodium hydroxide. Place the hydrochloric acid in the Styrofoam cup and then add sodium hydroxide during the graphing process. Save the file as “001” when graphing is complete. 11. Repeat steps 1-8 for 25mL 2.0 M ammonium chloride and 2.0 M sodium hydroxide. Place the ammonium chloride in the Styrofoam cup and then add sodium hydroxide during the graphing process. Save the file as “002” when graphing is complete. 12. Repeat steps 1-8 for 25mL 2.0 M hydrochloric acid and 2.0 M ammonium hydroxide. Place the hydrochloric acid in the Styrofoam cup and then add ammonium hydroxide during the graphing process. Save the file as “003” when graphing is complete.

Discussion: The purpose of this experiment is to find the enthalpy of a reaction by using calorimetry and Hess’s law after measuring the temperature change for the reaction. The change in enthalpy will give the heat of reaction for three acid-base reactions. The enthalpy for the reaction between sodium hydroxide and hydrochloric acid is calculated to be -107 kJ/mole. The enthalpy for the reaction between sodium hydroxide and ammonium chloride is calculated to be -10.132 kJ/mole. The enthalpy for the reaction between hydrochloric acid and ammonia if calculated using mass, specific heat, and change in temperature is -94.122 kJ/mole. The percent error for this calculation is 77.55%. The enthalpy for the same reaction if calculated using Hess’s law is -96.868 kJ/mole. The percent error for this calculation is 82.735%.

The results support the purpose of this experiment because the enthalpy was found for each reaction using both calorimetry and Hess’s law. Although the purpose was supported, the percent error was extremely high which would suggest that this experiment was not accurate. The first potential source of error is that the calorimeter was not fully insulated. This would mean more heat would be lost to the surroundings and the enthalpy change calculated wouldn’t be accurate. This could be prevented by using a more insulated vessel for the solutions to perform the calorimetry with. Another source of error could have been not measuring as close to 25mL as possible for the reactants, which would mean that the amount of reactants used were not balanced. This would affect the results because if there was significantly more of one reactant then the data would not be as accurate as what it could be if the two reactants were balanced. This could be prevented by taking the time to accurately measure 25mL of each reactant. A third source of error could have been reading the volume wrong on the graduated cylinder. This would affect the enthalpy calculations because the volume recorded would not be accurate. This could be prevented by not rushing when reading the volume and making sure the reading is as accurate as possible.

Conclusion: During this experiment, I have learned to measure temperature change using a calorimeter and use that data to calculate the enthalpy of a reaction. I also learned how to find the heat of reaction using Hess’s law in a real-life experiment setting; before this lab I only knew how to calculate enthalpy with Hess’s law from given reactions. Finding enthalpy is a very common real-life situation for people or companies that want to efficiently use energy, an example of this is car manufacturers finding the energy that the fuel releases so they can make

sure the car burns energy and fuel proficiently. This experiment could have been improved with having more chemicals in several different fume hoods so that there was no wasted time standing in line waiting to measure chemicals. This was not only an issue because of wasted time but there was also added pressure to pour the chemicals quickly which may have led to measurements not close to 25mL.

Work Cited: French, April N., Allison Soult, Stephen Testa, Meral Savas, Francois Botha, Carolyn Brock, Charles Griffith, Darla Hood, Robert Kiser, Penny O’Conner, William Plucknett, Donald Sands, Diane Vance, William Wagner. “Experiment Eleven: Enthalpy of Reaction.” Chem 113 General Chemistry II Lab Manual. Plymouth, M: Hayden-Mcneil, 2018. 59-64. Web. 7 February 2020. https://www3.chem21labs.com/labfiles/36194_43_Exp%2010_FrenchA %202187-1%20W20..pdf?rf=6771...